Pheloux, A., Compt. Rend. 224,2238 (1957). Pings, C. J., Nebeker, E. B., IND.ENG.CHEM.FUNDAMENTALS 4, 376 (1965). Prigogine, I., "Introduction to the Thermodynamics of Irreversible Processes," 2nd ed., Interscience, New York, 1961. Rastogi, R. P., et al., Physi'ca 27, 265 (1961). Rastogi, R. P., et al., Tram. Faraday SOL.61, 854 (1964). Van Rysselberghe, P., IND.ENG.CHEM.FUNDAMENTALS 4, 142 ( 1965a). Van Rysselberghe, P., IND. ENG.CHEM.FUNDAMENTALS 5 , 152 (1966). Van Rysselberghe, P., J . (Them. Phys. 36,1329 (1962).
Van Rysselberghe, P., J . Chem. Phys. 43, 3422 (1965b). Wei, J., Zahner, C., IND.ENG.CHEM.FUNDAMENTALS 5 , 151 (1966). Wei, J.,Zahner, C., J . Chem. Phys. 43, 3421 (1965). Wendt, R. P., J . Phys. Chem. 66, 1219 (1962). Wilks. S. S.. "Mathematical Statistics." Princeton University Press, Princeton, N. J., 1944. RECEIVED for review September 12, 1966 ACCEPTEDMay 5, 1967 Contribution No. 1479 from Indiana University Chemistry Department.
C A T A L Y S I S OF A L C O H O L A N D E T H E R D E H Y D R A T I O N ON G A M M A - A L U M I N A HERMAN J. SOLOMON,1 HARDING BLISS, AND JOHN B. BUTT Department of Engineering and Applied Science, Yale University,
New
Haven, Conn.
A recycling reactor was used for the experimental measurement of partial pressures of all products vs. time in the dehydration of alcohol and ether over y-alumina catalysts. Data were observed at 220' to 374" C. with alcohol and ether feeds under such conditions that diffusional effects were not important. A kinetic model is proposed and is shown to be effective in explaining the experimental observations. This model, the rate constants and activation energies of which are reported, is consistent with a single site interaction on the surface. The relationship of the model to recent mechanistic proposals is discussed.
THE dehydration of ethanol with alumina catalysts, despite the large amount of literature on the subject, remains a most interesting problem. One challenging aspect of this is the fact that two principal over-all reactions occur simultaneously (ether and ethylene formation) ; the first predominates at low temperatures, the second a t higher ones. As a result the over-all system involves a combination of parallel and simultaneous rea.ctions. Partial explanations of this behavior have been given by Butt, Bliss, and Walker (Z),and more recently by Knozinger and Kohne (77), but it is clear that further investigation could be expected to be fruitful only if the range of experimental variables studied, particularly temperature, were greatly expanded. Such a detailed and extensive study of a single alcohol dehydration system would provide data pertinent to recent mechanistic studies which have employed a very wide range of reactant alcohols but only small ranges of experimental conditions for the individual investigations. Accordingly, the first purpose in the present research was to measure partial pressure-time data for the reaction over a wide temperature range, employing a recycling reactor system. Since the behavior of ether, one of the products of alcohol dehydration, is coupled to that of alcohol in the over-all scheme, ether was also iemployed as reactant in many runs. A second purpose was to establish a kinetic model capable of correlating the experimental data with particular regard to the effects of temperature and feed composition and to relate such a kinetically based model interpretation to various mechanistic proposals for the dehydration catalysis. The range of experimental variables involved in this study is given in Table I. Present address, Esso Research and Engineering Co., Linden, N. J.
Experimental
(a,
Apparatus. T h e recycling reactor (see Fi ure 1) used was a modification of that described previously consisting of a loop of 35-mm. borosilicate glass tubing with associated sample port, pressure tap, and vacuum connections. The pump used, contained within the tubing loop, was a Model 19A-1352 Vanaxial blower made by the Globe Equipment Co., Cleveland, Ohio, and gold-plated to prevent catalytic effects. This pump could not be operated above 150' C . ; hence it was located in that part of the glass system above the salt bath level where such temperature control was possible. Since small radial temperature gradients occurred with a catalyst bed of 35-mm. diameter, the catalyst was located
Table I. Experimental Conditions
Catalyst
5.002 grams of 7-alumina, Type 1404, Harshaw Chemical Co., supplied as I / B X '/B inch cylinders, cut in half axially and radially so that their maximum dimensions were l / 1 6 inch. This catalvst was shown to exhibit tvDical x-rav diffraction pattern for ?-alumina. NazO = 0.047c, CaO = 0.97, by weight 190 s q . meters/g. 220' C., 290 mm. Hg initial pressure 254' C., 155, 230, 310, 316, 319,s and 321 mm. HF 291 O-C., 308 mm. Hg 306' C., 326 mm. Hg 329 ' C., 304 mm. Hg 374" C., 182 mm. Hg 241 ' C., 261 mm. Hg 291 ' C., 183 mm. Hg 306" C., 207 mm. Hg 329' C., 191 mm. Hg I
BET area Alcohol feed
Ether feed
a
,
Catalyst uncut in this run.
VOL. 6
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AUGUST 1 9 6 7
325
DIRECTION OF GAS FLOW C
HEATING
STAINLESS STEEL PIPE
GLASS TUBING
SAMPLE PORT,n.
Y
B
OPCOCK
-1 CATALYST
B
I
I
L-v
----e-
TO PRESSURE GAUGE
II
V4"STAINLESS ST.EEL TUBING
Figure 1 .
Recycling reactor system
inside a cylindrical section of 18-mm. glass tubing, 4 inches high. A few 2-mm. glass beads were placed in the annulus between the 18- and 35-mm. tubing in the reactor section; there was no gas flow i n this annulus. A thermocouple was placed in the center of the bed. The preheat-recycle section of the reactor system was immersed in a molten salt bath in which the temperature was controlled to 3~0.1' C. with a bimetallic regulator. Several sections of '/r-inch 0.d. stainless steel tubing were placed in the horizontal run of the immersed section to promote heat transfer to the recycling gases. The glass tubing above the salt bath was heated with heating tapes to prevent condensation. System pressure was measured with a Heise pressure gage, Model 8127, with a range of 0 to 1500 mm. of Hg, divisions of 1 mm., and a precision of 0.25 mm. of Hg. Analysis. Gas samples were taken from the reactor in a heated syringe and transferred to a Beckman Model GC-2 gas chromatograph. The column (Perkin-Elmer Type W) was operated at 20 p.s.i.g. and 130' C., and separated air (leakage) and ethylene (combined in one peak), ether, ethanol, and water. The effluent from this column was passed through a cold trap to remove ether, ethanol, and water and then fed to a Perkin-Elmer Model 154-C Fractometer (Type J column) operated at 20 p.s.i.g. and 80' C. This column separated air and ethylene. Reagents employed in the experimentation and for chromatograph calibration were 200-proof reagent grade ethanol, reagent grade diethyl ether, laboratory-distilled water, and Mathieson C.P.grade ethylene. Procedure. The salt bath was brought to the desired temperature and held overnight. For the first run of the day a charge of reactant was added, the blower turned on, and the system operated for 5 minutes; for subsequent runs this was not done. The reactor was then evacuated to below 2 mm. of Hg absolute and held for 15 minutes. The feed was injected through a sample port after the reactor had been isolated from the vacuum system and the blower started. About 5 seconds were generally required for this, although the time of starting the blower was set as zero time. Pressure and catalyst temperature were read a t frequent intervals during the course of each run. Samples of the reaction mixture were taken after 3 or 6 minutes, and every 6 minutes thereafter. Run lengths varied from 20 minutes to 5 hours, depending on the magnitude of the reaction rates under the conditions of the particular run. Catalyst Temperatures. Catalyst temperature variation with axial or radial position never exceeded 0.2' C.; thus reported temperatures were measured with a thermocouple located half-way along the axis. Since the pump was located in a relatively cool part of the system, a small drop in bed 326
I&EC FUNDAMENTALS
temperature (nominally 1' C.) was ordinarily noted at the start of a run. The duration of this deviation was not sufficient to affect results. Catalyst Stability a n d Reproducibility. Fresh catalyst demonstrated some dehydrogenation activity; however, after an hour of operation noticeable aldehyde formation ceased. Accordingly, the catalyst was stabilized by 5 hours' reaction before being used. Reproducibility of the stabilized catalyst was checked by making three runs at 254' C. with ethanol feed, each separated from the others by several runs a t other conditions. The reproducibility observed is shown in Figure 2 ; these results agree with previous reports (2, 8, 9) concerning the stability of y-alumina for dehydration reactions. Similarly, the rate of adjustment of the catalyst surface to the transient conditions within the reactor system was verified to be rapid in comparison with measured conversion rates, as discussed previously (2). I t was shown by charging the empty system with ethanol at 375 mm. of Hg and 310' C. and holding for 25 hours that the reactor itself was not catalytic. Diffusional Effects. Since the superficial velocity through the catalyst bed could not be varied easily, the effects of boundary layer diffusion were determined by calculating the rate of diffusion of alcohol to the catalyst surface under the most severe condition (initial rates at 374' C.). The j D correlation of Hougen and Wilke (6) was used, employing a value of 0.39 sq. foot per hour (0' C. and 1 atm.) corrected for temperature and pressure as the diffusivity of alcohol in air. For typical velocities (0.7 s.c.f.m.) the rate of diffusion at the highest temperature was calculated to be five times the initial rate of reaction, and 15 times this value at the lowest temperature, so the effect is very small. Pore diffusion effects were studied experimentally by slicing catalyst particles in half axially and radially and comparing the results of runs made with cut and uncut catalyst, as shown in Figure 3. The difference between the two sets of results was small enough to justify the assumption that pore diffusion was not significant. All data reported are with the smaller catalyst particle unless noted otherwise, and on the basis of the observations above it is reasonable to believe that these data are representative of the surface reactions alone. Interpretation of Data and Results
Treatment of Pressure. Since there was a small amount of air leakage into the system in some runs, the measured total pressure could be in error; therefore total pressures were
Figure 2.
Reproducibility of alcohol conversion at 254' C
Figure 3.
TIME (MINUTES1 Conversion for alcohol feed
Figure 4.
Effect of particle size
219.8' C., 290 mm. Hg
determined from stoichiometry and composition measurement. The stoichiometry of these reactions with alcohol feed may be described by two independent reactions:
+ HzO = CzH50CzHj + H z 0
CzHE0H = CzH4 2CzH5OH
The relation between initial pressure, P I , total pressure, PT, and mole fraction of olefin in the reactor (air-free basis), xo, is:
Pp = P1
+ xo P*
(1)
The mole fraction of oslefin is determined from the chromatographic analysis. Since
Pi
=
X(
PT
(2)
the partial pressure of a.ny component, i, may be computed as Figure 5.
TIME (MINUTES) Conversion for ether feed 241
(3)
C., 260.5 mm. Hg
The calculated total pressure from Equation 1 was used instead of measured values in computing each P,, but calculated and measured total pressures agreed within 0.5% in runs where air leakage was small. Equation 1 may be written in the same way for ether feed, where the independent .reactions are now:
+ CzH4 2CzH50H = CzHsOCzHj + HzO CzH50Cz:Hs = CzHsOH
Presentation of Reciults. I t is convenient to divide the partial pressures of each component by the initial pressure, PI, and the resulting normalized pressures, Pf/P1, are shown as points in Figures 4 to 13. The lines on these figures were computed with the kinetic model discussed below. These normalized partial pressures vary from 0 to 1 except for ethylene with ether feed, which is 0 to 2 . Effect of Pressure. The plot of normalized partial pressures us. time for runs made with alcohol feed a t 254' C. and 155, 230, and 316 mm. of Hg pressure shows negligible effect of pressure; so in the work to follow the absolute level of pressure does not appear, and the interpretation is directed toward relating the behavior of the normalized partial pressure us. time data to temperature and feed composition.
Figure 6.
TIME (MINUTES) Conversion for alcohol feed 254' C., 31 6 mm. Hg
Kinetic Model
The stoichiometry of the reactions by which ethanol or ether may be converted to ethylene, water, and ether or ethanol may be described in terms of two reactions as noted VOL. 6
NO. 3 A U G U S T 1 9 6 7
327
TIME (MINUTES) Conversion for alcohol feed
Figure 7.
Figure 9.
TIME (MINUTES) Conversion for alcohol feed 306' C., 325.5 mm. Hg
291 C., 307.5 mm. Hg
I
Figure 8.
TIME (MINUTES) Conversion for ether feed
Figure 10.
0
0 It
II
A1
- OH + C ~ H S O H= A1 - OCzHs + HzO I1
The complex may then decompose and regenerate the site
0
II
A1
0
- OCzHs
= CzH4
11
+ A1 - OH
or it may react with more alcohol to form ether and regenerate the site 328
l&EC FUNDAMENTALS
I
I
t
I
TIME (MINUTES) Conversion for ether feed 306' C., 207 mm. Hg
291 C., 182.5 mm. Hg
above. Stoichiometry, however, has little do do with rates of reaction or with what actually takes place on the catalyst, and more than two reactions may, in fact, occur. There are conflicting views as to what happens on the alumina surface in catalysis of dehydration reactions, but some of these are helpful in formulating a kinetic model. Mechanistic Studies. Ipatieff in 1936 (7) proposed that alcohol dehydration involved the formation with A1-OH sites of an ethoxy complex on the surface by the reaction
A ALCOHOL
0
(I
A1
0
- OCzHS + CsHsOH
ii
=
C2HbOCzHs
+ A1 - OH
Topchieva, Yun-Pin, and Smirnova (78) wrote the first two steps in essentially the same way, but wrote the third as I
1
'/z
HzO
I
1
+ A1 - OCzHs = '/*C2HjOCzHs + A1 - OH I
I
The site, written slightly differently by Topchieva et a / . , is again regenerated. There is ample evidence from infrared studies for the existence of the hydroxyl groups on y-alumina essential to this mechanism and Peri's (4, 12) surface structure model demonstrates several methods of which the necessary sites are generated on dehydration of the alumina. Heiba and Landis (5) have shown that various aluminum alkoxides decompose to yield products similar to those which result from decomposition of ethers over alumina and suggest that alcohols acd ethers react with alumina to form surface alkoxide compounds,
A ALCOHOL 0 ETHER
+
Figure 1 1.
OLEFlN
TIME (MINUTES) Conversion for alcohol feed
Figure 12.
TIME (MINUTES) Conversion for ether feed 329" C., 191 mm. Hg
329' C., 304 mm. Hg
which decompose to yield the observed products. Some conflicting evidence, ho\vever, has been reported by Jain and Pillai ( 8 ) . The twci ethoxy mechanisms may be concisely written as: Ipatieff.
A A
+ S eC +W + C eE + S
+S
C S 0 Topchieva.
+S C 2C + W A
(Reaction 1) (Reaction 2 ) (Reaction 3)
+W +S S 2E + S
S C 20
(Reaction 1) (Reactions 3) (Reaction 4)
in which A, E, W, and 0 denote alcohol, ether, water, and olefin.
I
AI-0-C
I
S is used to denote A1-OH
and C to designate
I
&I 5 .
TIME (MINUTES)
I
An alternative mechmism postulates carbonium ion formation and has been considered by several workers. Brey and Krieger ( 7 , 79)have written for alcohol dehydration -0H-
CzH60H
11
C2Hs' I
-H'
C2H4
C
-CZHSO-
C zHsOCZHS and Winfield (20) has presented a very detailed consideration of this mechanism in terms of ions and vacancies on the alumina surface. Although conversion data such as those reported herein cannot provide a complete basis for evaluation of such detail, if one writes the scheme in terms of a generalized acidic site, the Brey and Krieger mechanism may be written:
H+S
+
+
e
+ +
H'S CzHs +S HOH CZIls +S S C Z H ~ H + S CZHSOCZHS 2 CzHs +S CzHsOH
CzHsOH
+
These equations are essentially identical to those of the Ipatieff mechanism, the only difference being in the way in which site and complex are characterized. Obviously, other carbonium ion schemes could be written to correspond to the T'opchieva mechanism. Thus, both of these specific mechanistic postulates are described by equations of the same form which lead to similar
Figure 13.
Conversion for alcohol feed 374' C., 182 mm. Hg
rate equations, and it will be difficult to discriminate between them on the basis of rate or conversion measurements alone. O n the other hand, difference between the two ethoxy mechanisms can be distinguished, since different rate equations are involved. Other Mechanisms. T h e ethoxy and carbonium ion mechanisms are similar in that only one type of site is involved in each. A complex is formed on that site which may react in a variety of ways, as indicated in Reactions 1 to 4. Some recent work tends toward a mechanistic interpretation of dehydration catalysis involving two sites. T h e internal elimination mechanism of Pillai and Pines (73), for example, requires one part of the alcohol molecule to be held at a basic site and another at an acid site. It is thus a concerted mechanism involving a pair of sites. The probable interaction of ether with the two sites would also yield olefin via this elimination reaction. A different type of two-site mechanism is proposed by Shchekochikhin and Makarov (75), in which certain sites (claimed to be confirmed by infrared studies) experience ethoxy formation a t lower temperatures yielding ethers. Other sites are described as VOL. 6
NO. 3
A U G U S T 1967
329
CHI
o=c
/
I
A1
+0
These sites yield olefins a t higher temperatures. Ether chemisorption is not discussed by these workers. Knozinger (70)has also postulated two types of active sites, oxygen ions and hydroxyl groups. Ethoxy groups and hydrogen-bonded molecular alcohol are postulated as surface intermediates. Knozinger has presented a reaction scheme also postulated by us ( Z ) , describing gas-phase reactants and products as 2A A E
eE +W
0 fW +2 0 W E + A f0
+
(Reaction (Reaction (Reaction (Reaction
la) 2a) 3a) 4a)
These reactions bear some similarity to the one-site mechanisms-for example, Reaction l a is the sum of Ipatieff's 1 and 2-and have indeed been analyzed in such terms previously (2). There is, however, no doubt that any of the twosite mechanisms would yield an array of rate equations differing considerably from the one site cases. Proposed Kinetic Model. The summary above is representative of mechanistic proposals for dehydration catalysis by alumina. Our principal efforts in the present study are directed toward the evaluation of kinetic models based on onesite mechanisms, since these will be involved even in the more complex two-site proposals. The model which we have employed here is a combination of Reactions 1 to 4,and is given by
+ e + + + +
A S C W E+S e C +A 0 S C E 25 e 2C W
(Reaction (Reaction (Reaction (Reaction
1) 2) 3) 4)
(Since alcohol synthesis from ethylene and water does occur a t high pressure, Reaction 3 must be reversible under such conditions.) This model has been thoroughly tested with the methods described below. Variants on the above explored included the omission of Reaction 4, substituting Reaction 4 for 2, making 2 irreversible to the left, etc. Eight changes of this sort were tried, as well as three other variations involving a change in the number of sites per unit weight of catalyst due to hydration by product water. T h e rates of these reactions are described by the following equations: (4)
(5)
in which the partial pressure of each component is designated by PA, etc., Cis the concentration of the complex on the surface, and S is the concentration of available sites. We have little information concerning the magnitude of C and S other than the relation
so-c+s 330
I&EC FUNDAMENTALS
(8)
which requires the open sites plus the complex-occupied sites to be constant. The best data on this are given by Pines and Haag ( 1 4 , who measured the adsorption of trimethylamine on various aluminas at 150 mm. of Hg, and 25' to 300' C. The amount adsorbed decreased over this temperature range, and the absolute amount adsorbed a t 300' C. was of the order of 1 X 1013 molecules (acid sites) per sq. cm. I t is advantageous to normalize these equations by dividing the rates and the partial pressures by P1 and the values of C and S by SO. Hence
IFSO SO) 0s A'
-
RI
=
Rq
= (k4FSo2)Os2 E'
(kiRSo) Oc W'
- ( k q B So2) Ocz W'
(9)
(12)
in which A , etc., designate normalized partial pressures of the constituents, Oc is the fraction of the surface occupied by the complex, and Os is the fraction of the surface which is unoccupied. The parenthetic terms are the lumped constants which must be determined. As noted, So is unknown, but it is now associated with a rate constant. The values of Os and Oc are also unknown, but they must vary between 0 and 1. Their values are computed by the integration method to be described. Initial values are Os = 1, Oc = 0, A' = 1 (alcohol feed), E' = 1 (ether feed), and W' = 0. An Equilibrium Restriction. Cope (3) has shown that the reaction 2C2H60H = C Z H S O C Z H ~ HzO is reversible and has reported values of the equilibrium constant, K,. Typical values are 6 at 240' C. and 4 at 300' C. Since this reaction is the sum of Reactants 1 and 2 in our scheme, they must also be considered reversible and
+
This restriction must be accounted for in the analysis of the conversion data. Integration of Rate Equations. The model rate equations were integrated to allow comparison with partial pressuretime data. Equation 9 to 12 were integrated numerically, using a Runge-Kutta technique for simultaneous equations, with trial values of the parameters. The measured partial pressure us. time curves were then compared with those calculated. An iterative procedure was employed until a reasonably good fit was obtained. The process was repeated for all runs with both ether and alcohol feeds and a t all temperatures. Assistance in the trial and error procedure was provided by making Arrhenius plots for each parameter, so that trial values for other temperatures were thus somewhat restricted. The equilibrium relation of Equation 13 was also helpful in the trials. Full details of the procedure are described by Solomon (76),but a few points should be noted : 1. Certain constants are most important in fitting certain regions of the conversion-time curves. For example, kiF& is very important in determining initial slopes of alcohol, ether, and water curves with alcohol feed, kzFSo is very important in determining initial slopes of ether and alcohol curves with ether feed, and klRSo is most important in fitting final slopes of the alcohol curve with alcohol feed. 2. The search for these constants was carried out with a number of objectives in mind:
Fitting simultaneously all four components with either feed. Approximately constant deviations for each component a t all times. Fitting the measured maxima. Fitting initial rates. Approximately the same deviations for all components. The values so determined will differ in some cases from those evaluated by means of nonlinear least squares. The Langmuir-Hinshelwood equations for surface reaction employed for correlation in prior work (2) were also tested in early trials on an analog computer. These forms were not able to fit the observed selectivity of olefin formation with temperature. Values of the rate 'constants for the model are shown in Table 11. It may be confirmed that Equation 13 holds for these values of the parameters and it should be noted especially that these final resulh show Reaction 4 to be irreversible. The agreement between data and calculations is illustrated in Figures 4 to 13, in which the points represent experimental data and the curves ,are calculated results. A few sample calculations are given here to demonstrate the agreement between rates calculated from the lines on Figures 4 to 13 and those calculated from the various constants. For example, from Equations 9 and 10
dA' = R I + R2 = dt
(kipS0)
OsA'
+
(ki~S0) Oc
( k 2 ~ S o6sE' ) dA ' - (initial) = dt
W'
-
Though exact agreement cannot be expected, these figures show that the numerical values of the rate constants as reported are consistent with the curves plotted. Behavior of ec a n d es. The fraction of the catalyst surface covered and that unoccupied were calculated in the machine computation. Since there are no data to compare with these calculated quantities, they are not reported in detail. However, OC goes through a maximum if the product ether (alcohol feed) o r alcohol (ether feed) goes through a maximum; otherwise it does not. Values of & were computed as high as 0.18; in general, Oc is higher with ether feed than with alcohol. Temperature Dependence of Rate Constants. Arrhenius plots for the rate constants are given in Figures 14 to 16 and corresponding activation energies are reported in Table 111.
Table 111.
Activation Energies CaI./G. Mole 22,400 7,300 26,000 6,600 46,400 15,200
+
( k z ~ s oOCA' )
- (klF,YO)
since Os = A' = 1 and E' = Bc = W' = 0 At 254' C., Table 11, = -0.0017 X 5 grams = -0.0085 At 254' C., Figure 6.. dA'/dt (initial) = -0.014 Similarly, from Equations 10 and 12
= - ( k ~ p S o ) OsE' f ( k 2 ~ S oOCA' )
-
(k4pS;)
+
@s2E'
(k4RS02) ec2
dE - (initial) = dt
W'
- (k~p!?O) - (kdFSO2)
since Os = E' = 1 and Oc = A' = W' = 0 1-5
-
At 241' C., Table 11, = (-0.00065 0.00027) x 5 grams = -0.0046 At 241 ' C., Figure 5, dE'/dt (initial) = -0.0044 from graph
Table It. Temp., C . 219.8 241 254 291 306 329 3 74 Value for k&
klFS0 0.00026 0.00076 0.00170 0.00580 0.00840 0.0196 0.0620
and not k&/P1.
kid0
0.0208 0.0296 0.0384 0.0560 0.0660 0.0803 0.0930
Figure 14.
and
1.6
2.0
2.i
(%-1,
Arrhenius plot for K1rSo, KZFSO,
K4FSO'
Values of Constants kZFS0 k2RSO 0.00019 0.00065 0.001 30 0.0176 0.0110 0.0262 0.1040
1.9
1.7
i/T x lo3
0.110 0.148 0.176 0.280 0.310 0.380 0.450
k3SOa
0,00116 0.0100 0.0310 0.552 2.00 8.78 74.2
k4~So' 0.00013 0.00027 0.00040 0.00180 0.00144 0.00220 0.00380
Units of parameters are: kl&, (g.cat.)-' (min.)-I kl&, (g.cat.)-' (min.)-l kzFSO, (g.cat.)-' (mzn.)-l kzRSO, (g. cat.)-' (min.)-l k S o , (g. cat.)-' (min.)-' (mm. Hg) k a ~ s o ' ,(g. cat.)-' (min.)-l
VOL. 6
NO. 3
AUGUST 1967
331
”*
.io0
3 I
Figure 16. Figure 15.
While the temperature dependence of So has not been reported, Pines and Haag (74) have shown that the number of sites on alumina a t 300’ C. is about 0.6 that at 200’ C. Correlating this ratio with an exponential inverse temperature dependence leads to an “activation energy” for SO of about -3000 cal. Thus, the “true” activation energies for the reactions involved should be about 3000 cal. greater than shown for all reactions except that involving Soz, in which case the “true” figure should be about 6000 cal. greater than shown. The activation energies of Table I11 are children of the kinetic model and, as such, cannot be compared directly with those obtained from less detailed analysis. However, the activation energy of the main alcohol reaction, 22,400 3000, is in qualitative agreement with that reported by Brey and Krieger (I), 38,000, by Kabel and Johanson ( g ) , 30,500, and by Stauffer and Kranich (77), 30,800.
+
Conclusions
Partial pressure us. time data, observed through a wide range of temperature with both alcohol and ether feed, exhibited a considerable variation in behavior, including the appearance of maxima in certain products and temperaturedependent selectivity for olefin formation. Practically complete conversion was obtained in some cases. No effect of pressure on these results was found, but there were substantial effects of temperature and initial reactant. This wide range of behavior for all products can be interpreted in terms of a relatively simple kinetic model based on a series of four surface reactions:
+ S e C + H10 + C e S + CzH50CzH5 C C2H4 + S CzHsOCzH5 + 2 s 2C + H20 C2H5OH CzHsOH
-+.
-+.
This scheme indicates only one type of active center and it is inviting to identify the site and complex with an ethoxy mech-
I
I and C2H50-Al,
I 332
I
I
\A
\
I
I
-
Arrhenius plot for K3So
Arrhenius plot for KlRSo and
~ R S O
anism (HO-A1
I
l&EC FUNDAMENTALS
respectively), but other
I
mechanistic postulates involving a simple site will result in similar rate equations, and it is also possible to write (under certain assumptions) the concerted mechanisms in an analogous array (though not as shown above), in which S would refer to the acid-base combination on the surface. The model is shown to provide an excellent fit for the results of experimentation-the distribution of four different products among six different reactions for two possible feed materials over a temperature range exceeding 150’ C. Acknowledgment
The financial assistance of E. I. du Pont de Nemours & Co. and of the Monsanto Co. is gratefully acknowledged. We also thank the Harshaw Chemical Co. for supplying the catalyst. Nomenclature
A’ C
= normalized partial pressure alcohol, PA/P1 = complex concentration, (9. moles complex) (g.
catalyst) -l normalized partial pressure of ether, PE/P1 klF, klR, k2F, k2R = forward and reverse reaction velocity constants of Reactions 1 and 2, (min.)-l (g. moles sites) -1 ka = reaction velocity constant of Reaction 3, (mm. Hg) (min.) -1 (g. moles sites) -1 kdF, k b R = forward and reverse reaction velocity constants for Reaction 4, (g. catalyst) (min.)-l (g. moles sites) -z = equilibrium constant in terms of partial pressures KP = partial pressure, mm. H g P = initial pressure, mm. Hg Pl = total pressure, mm. Hg PT rl, rz, ra, 14 = rates of Reactions 1, 2, 3, and 4, (mm. Hg) (min.) L1 (g. catalyst) R1, Rz,R3, R4 = rates of Reactions 1, 2, 3, and 4 riormalizedLe., divided by initial pressure, (min.)-l (g. catalyst) -1 S = concentration of unoccupied sites, (g. mole sites) (g. catalyst) -l so = original concentration of unoccupied sites, (g. moles sites) (g. catalyst)-l x = mole fraction = normalized partial pressure of water, P,/Pi W’
E‘
=
GREEK 8c 8s
=
fractional coverage of complex, CIS0
= fraction of unoccupied sites, SI&
SUBSCRIPTS A = alcohol C = complex E = ether 2 = general 0 = olefin S = unoccupied sites TY = water literature Cited (1) Brey, \Y. S.,Krieger, K. A., J . A m . Chem. SOC. 71, 3637 (1949). (2) Butt, J. B., Bliss, H., Walker, C . A., A.Z.Ch.E. J . 8, 42 (1962). f 3 1 Coue. C. S.. A.Z.Ch.E. J . 10. 277 11964). ( 4 j Grknler, K. G., J . Chem. Piys. 37, 2094 (1962). ( 5 ) Heiba, E. I., Landis, P. S., J . Catalysis3, 471 (1964). ( 6 ) Hougen, 0. A., IVilke, C. R., Trans. A.Z.Ch.E. 45, 445 (1945).
(7) Ipatieff, V., “Catalytic Reactions at High Temperatures and Pressures,” p. 552, Macmillan, New York, 1936. (8) Jain, J. R., Pillai, C . N., Tetrahedron Letters 11, 675 (1965). (9) Kabel, R. L., Johanson, L.N., A.Z.Ch.E. J . 8,621 (1962). (10) Knozinger, H., Z. Physik. Chem. ( N e w Folge) 48, 151 (1966). (11) Knozinger, H., Kohne, R., J . Catalysis 5 , 264 (1966). (12) Peri, J. B., J . Phys. Chem. 69,211 (1965). (13) Pillai, C. N., Pines, H., J . Am. Chem. Soc. 83, 3274 (1961). (14) Pines, € I , , Haag, W. O., Zbid., 82, 2471 (1960). (15) Shchekachikhin, Y . M., Makarov, A. D., Kinetika i Katalir 5 , 568 (19154). (16) Solomon, H. J., D. Eng. dissertation, Yale University, New Haven, Conn., 1966. (17) Stauffer, J. E., Kranich, W. L., IND. ENG. CHEM.FUNDAMENTALS 1. 107 (1962). (18) Topchikva, K.V.,’ Yun-Pin, K., Smirnova, I. V., Adoan. Catalysis 9, 799 (1957). (19) Whitmore, F. C., J . A m . Chem. SOC. 54, 3274 (1932). (20) Winfield, M. E., “Catalysis,” P. H. Emmett,. ed... Vol. VII. Chap. 2, Reinhold,’New York,.l960.
RECEIVED for review September 8, 1966 ACCEPTEDMarch 28, 1967
CATALYTIC ADDITION OF HYDROGEN CHLORIDE
T O VINYIL CHLORIDE Studies in a Stirred Reactor R O B E R T G. R I N K E R A N D W I L L I A M H. CORCORAN Chemical Engineering Laboratory, California Institute of Technology, Pasadena, Calif.
The catalyzed reaction between hydrogen chloride and vinyl chloride in a stirred reactor was studied at temperatures in the range of 1 6 4 ” to 299” F., mass-flow ratios of vinyl chloride to hydrogen chloride from 2.1 6 to 2.46, and catalyst exposure times up to 1 8 0 0 hours. Zinc chloride carried on Celite was the catalyst, and regimes of fouling and nonfouling were observed. For the nonfouling case the rate of formation, rd, of the 1 ,I -dichlciroethane was found to follow the equation: rd = C(ph p o - pd/KP), where Cis an empirically determined rate constant, K p is the equilibrium constant referred to partial pressures, and ph, po, and pd are the partial piressures, respectively, of hydrogen chloride, vinyl chloride, and 1,l -dichloroethane. A fourcenter reaction appeared to be the most probable mechanism.
et al. (1928, 1931, 1932) made a systematic study of the addition of hydrogen halides to vinyl halides. They used a tubular reactor packed with a solid catalyst and obtained conversion data in the temperature range of 25’ to 200’ C. The work showed that ferric chloride catalyzed the reaction of hydrogen chloride with vinyl chloride to give 1 , l dichloroethane but had a much lower activity than either HgC12 or ZnC12. There was no correction, however, for the loss of activity due to fouling. Independent of the temperature and type of catalyst, the only product of the reaction between hydrogen chloride and vinyl chloride was 1,l-dichloroethane. The result suggested that the course of the addition reaction was governed by directed adsorption on the catalyst surface and was not a chain reaction. Between 1928 and 1040 Kharasch e t al. (1928, 1931, 1933, 1934, 1937, 1939, 1940) conducted a long series of comparative experiments on the mechanisms of the addition reactions of hydrogen chloride, hydrogen bromide, and hydrogen iodide with vinyl chloride. Kharasch arrived a t several generalizations from the experiments. First, he concluded that the additions of hydrogen chloride and hydrogen iodide IBAUT
to double bonds, although different in rate, occur by similar mechanisms which do not involve chain reactions. O n the other hand, the addition of hydrogen bromide involves chains in which Br is the chain carrier. Secondly, the effect of peroxides and metallic halides on the mechanisms and rates of addition was explained. Particularly the addition of HC1, with or without added ZnC12 or FeC18, and the addition of hydrogen bromide, with added FeC13, to vinyl chloride were mechanistically equivalent to the normal, uncatalyzed addition of hydrogen bromide to vinyl chloride to give l-bromo-lchloroethane. Therefore, the addition of hydrogen chloride to any ethylene derivative with or without catalysis would give only the normal addition product. Extensive experiments on the rate of thermal decomposition of 1,l- and 1,2-dichloroethane in a nonpacked tubular flow reactor were conducted by Barton (1949). I n a series of reactions involving the addition of small amounts of oxygen or chlorine to the feed, he observed that the rate of pyrolysis of the 1,2-dichloroethane was greatly increased, whereas that of the 1,l-dichloroethane was unaffected. H e concluded that in the case of the former compound, the mechanism of VOL. 6
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