A. L. DENTAND R. J. KOKES
3772 R-au ligands and the highly conjugated phthalocyanine and tetraphenylporphine ligands is most probably due to the deprotonated nitrogen sites in the equivalent CuN2- chromophore. The near equivalence of the four nitrogen atoms in the C U N ~ ~chromophore suggests that a delocalized model employing both inplane and out-of-plane n-bonding may be appropriate. Indeed, the agreement of the qualitative HMO
calculations with the covalency changes upon protonation, the infrared spectra, and structures suggested by analogy with those known for substituted biguanide complexes support this conclusion. The low gll and large -4.1,values found for the R-au complexes appear to be characteristic of complexes containing the equivalent CuN2- chromophore and are a result of the strong covalency in the complexes.
Hydrogenation of Ethylene by Zinc Oxide. I.
Role
of Slow Hydrogen Chemisorption by A. L. Dent and R. J. Kokes Department of Chemistry, The Johns Hopkins Unhersity, Baltimare, Maryland
61618 (Received March 64, 1969)
An experimentalbasis is provided for dividing hydrogen adsorption on zinc oxide into two types. Type I is rapid and reversible; type I1 is irreversible and occurs rapidly initially but slowly in the latter stages. Type I hydrogen is responsible for the observed OH and ZnH bands in the infrared and the principal source of hydrogen for the hydrogenation of ethylene. Type I1 does not participate in the hydrogenation of ethylene at room temperature but modifies the catalyst and enhances its activity. Neither the hydrogen adsorption nor ethylene adsorption is markedly affected by pretreating the catalyst with oxygen. The catalyst is poisoned for ethylene hydrogenation by exposure of the catalyst to oxygen provided chemisorbed hydrogen is present before the pretreatment. If care is taken to exclude chemisorbed hydrogen, the effect of oxygen pretreatment is relatively slight; hence, water (or its precursor) appears t o be the poison. I t is suggested that nonstoichiometry is not required for the development of active sites for hydrogenation. Instead, it is proposed that the strained sites, perhaps formed by dehydration, as suggested by Weller and Volta, are the active sites.
Introduction Ethylene hydrogenation over metals is a classical example of heterogeneous catalysis. Studies of this reaction and related phenomena strongly suggest that the reaction scheme includes the following surface processes
C2H5
+ H -+
CzH6(gas) (slow)
(2)
Such a scheme is consistent with a wide variety of data including kinetic,’ exchange,z chemisorption and infrared studies,3t4but for our purposes, we wish to focus on only two of these aspects. First, in accord with this scheme, when light ethylene reacts with deuterium over metals, a distribution of ethanes with the formula C2Ha-zDz (0 5 z 5 6) is obtained. Second, the fact that active hydrogenation catalysts also are effective hydrogen-deuterium equilibration catalysts is consistent with (but not required by) the implied dissociative adsorption of hydrogen. It should be noted, The Journal of Physical Chemistry
however, that while equilibration and hydrogenation are related reactions, the former is often more rapid than the latter. Metal oxides are also catalysts for both hydrogenation of ethylene5 and hydrogen-deuterium exchange.6 Studies of hydrogenation over oxides are scarcer than such studies over metals, but the definitive work of Burwell and coworkers’ shows that hydrogenation over one oxide, chromia, is simpler than that over metals: for example, addition of deuterium to light ethylene (1) For recent reviews, see: (a) G. C. Bond, “Catalysis by Metals,” Academic Press, London, 1962; (b) G. C. Bond and P. B. Wells,
Aduan. Catal., 5 , 2 5 7 (1967). (2) T. I. Taylor, “Catalysis,” Vol. V, P. H. Emmett, Ed., Reinhold Publishing Corp., New York N. Y , 1957 p. 257. (3) R. P. Eischens and W. A. Pliskin, Aduan. Catal., 10,l (1958). (4) L. H. Little, “Infrared Spectra of Adsorbed Species,” Academic Press, New York, N. Y., 1966, p 100-137. (5) D. L. Harrison, D. Nicholls, and H. Steiner, J . Catal., 7, 359 (1967). (6) D. A. Dowden, N. MacKenzie, and B. M. W. Trapnell, Proc. Roy. Soc., A237,245 (1956).
HYDROGENATION OF ETHYLENE BY ZINC OXIDE yields only CzH4Dz;hence, there is no smear of the deuterium distribution in the product ethane. This result implies that the first step in the above reaction scheme is irreversible; the relatively low rate observed for isomerization of higher olefins during hydrogenation supports this conclusion. I n part, this difference from metals can arise from the wide separation and lack of interaction of active sites on oxides, a suggestion first put forth by Burwell which has received some support from results recently reported from this laboratory.8 Attempts have been made to correlate the catalytic activity of metal oxides with the configuration of d electrons in the metal ion.5,6 Such correlations cannot be as general as for metals,' because the oxides active for ethylene hydrogenation include aluminumg~10 and zinc oxidesl1-l6for which the metal ions have empty and filled d shells, respectively. Correlation to d shell structure alone ia rendered still more suspect by recent observations that the addition of either deuterium*' or deuterium-hydrogen mixtures8 to ethylene over zinc oxide yields products quite similar to those obtained over chromia. On the basis of these results, together with quoted17 unpublished results of Ozaki on cobalt oxides, it has been suggested17that the relative simplicity of the deuterogenation of ethylene over chromia may be a general characteristic of oxide hydrogenation catalysts that is independent of the d shell structure of the metal ion. It is the purpose of this research to shed some light on the mechanism of hydrogenation over one of these oxides in the hope that it will have general validity for all of these oxides. Our candidate for this prototype is zinc oxide for a variety of reasons. First of all, it is relatively transparent in the infrared so that the kinetic runs can be supplemented by infrared studies. This aspect is particularly attractive insofar as Eischens, Pliskin and Low1* have shown that hydrogen chernisorbed on zinc oxide shows strong bands in the infrared and Bozon-Verduraz, Arghiropoulos, and Teichner19 have studied the infrared of ethylene chemisorbed on zinc oxide. Secondly, whereas chromium oxide can show a variety of oxidation states depending on pretreatment,20 zinc oxide is limited to one principal oxidation state. Mechanistic studies have been carried out on zinc oxide by Teichner and coworker^.'^-^^,^^ For the most part, however, these studies have been carried out at elevated temperatures, 100 to 400°, and results with chromia' suggest that the simplicity characteristic of oxides is largely lost at elevated temperatures. The first part of this work deals with the role of fast and slow hydrogen chemisorptionz1 in hydrogenation reactions and the effect of oxygen pretreatment. These are questions that have some bearing on the behavior of chromia. Parravano and Boudart2' state: "It is important to note that the principal features of
3773 hydrogen chemisorption, which are summarized above, apply equally well to other adsorbents than zinc oxide, for instance, to chromium oxide. A satisfactory explanation, therefore, must not depend on the specific properties of zinc oxide." In like vein, Weller and VoltzZ2have noted that exposure to oxygen poisons chromia catalysts, a result also reported for zinc o ~ i d e . ' ~ . - ' With ~ zinc oxide it has been suggested that such oxygen poisoning occurs because it renders the catalyst more nearly stoichiometric and nonstoichiometry is required for activity.12-16 Such an explanation cannot be invoked for chromia since exposure to oxygen promotes nonstoichiometry.zz The second part of this work deals with the mechanism of the hydrogenation reaction and the possible nature of the active sites.
Experimental Section Materials. Tank hydrogen (or deuterium) was purified by slow passage through degassed charcoal a t - 195'. Tank ethylene and oxygen were purified by alternate condensation, vaporization and pumping followed by distillation. Samples of hydrogen purified in this manner were shown by mass spectrographic analysis to contain less than 0.01% oxygen; gas chromatographic analysis demonstrated that the purity of the ethylene was comparable to that of the hydrogen. The oxygen purity was not checked, but prior to use it was freed from water by circulation through a liquid nitrogen trap. The zinc oxide used for these studies was Kadox 25 supplied by the New Jersey Zinc Co. Infrared Studies. Samples of zinc oxide weighing about 0.5 g were pressed into circular disks 20 mm in (7) (a) R. L. Burwell, Jr., A. B. Littlewood, M. Cardew, G. Pass, and C. T. H. Stoddart, J . Amer. Chem. Soc., 82, 6272 (960); (b) G. Pass, A. B. Littlewood, and R. L. Burwell, Jr., ibid.,82, 6281 (1960); (c) C. T. €3. Stoddart, G. Pass, and R . L. Burwell, Jr., ibid., 82, 6284 (1960); (d) A. B. Littlewood and R. L. Burwell, Jr., ibid., 52, 6287 (1960); (e) M.Cardew and R. L. Burwell, Jr., ibid.,82, 6289 (1960). (8) W. C. Conner and R. J. Kokes, J . Phys. Chem., 73, 2436 (1969). (9) J. H. Sinfelt, ibid.,65,232 (1964). (10) Y .Amenomiya, J. H. B. Chenier, and R. J. Cvetanovic, J . Catal., 9 , 2 8 (1967). (11) J. F. Woodman and H. S.Taylor, J . Amer. Chem. Soc., 62, 1393 (1940). (12) E. Molinari and G. Parravano, ibid.,75,5233 (1953). (13) E. H. Taylor and J. A. Wethington, Jr., ibid.,76, 971 (1954). (14) J. Aigueperse and S. J. Teichner, Ann. Chim. (Paris), 7 , 13 (1962). (15) J. Aigueperse and S. J. Teichner, J. Catal., 2 , 3 5 9 (1963). (16) F. Bofion-Verduraz and 8. J. Teichner, a'bid., 11, 7 (1968). (17) W. C. Conner, R. A. Innes, and R. J. Kokes, J . Amer. Chem. SOC., 90,6858 (1968). (18) R. P. Eischens, W. A . Pliskin, and M. J. D. Low, J . Catal., 1, 180 (1962). (19) F. Boeon-Verdurafi, B. M.Arghiropoulos, and S. J. Teichner, Bull. SOC.Chim. Fr., 2864 (1967). (20) Charles P. Poole, Jr., and D. S. MacIver, Adoan. Catal., 17, 223 (1967). (21) G. Parravano and M.Boudart, ibid., 7, 47 (1955). (22) S. W. Weller and S. E. Voltz, J . Amer. Chem. Soc., 76, 4695 (1954).
Volume 73, Number 11 November 1960
A. L. DENTAND R. J. KOKES
3774 diameter under a pressure of 25,000 psi in a stainless steel ring. This disk was centered in a modified cylindrical gas cell with NaCl windows. The central part of the cell could be heated by nichrome windings while the windows were kept at room temperature by water-cooled copper coils wrapped around the ends of the cell. A measure of the sample temperature was provided by a thermocouple introduced into the cell via a glass-to-metal seal; the configuration was such that the junction formed a pressure contact with the zinc oxide disk. The cell was connected via tubulations on either side of the disk to a circulation loop with a magnetically driven pump so that reacting gases could be circulated over the catalyst disk. The rest of the vacuum system was a conventional B E T apparatus. After assembly of the cell, the sample was degassed first at about loo", then for 3 hr at 400" (the standard degassing procedure). If any bands due to C-H species were observed, oxygen was circulated over the catalyst for 3 hr at 400" with a liquid nitrogen trap in the line. Then, after brief degassing, the temperature was lowered to 300" and dried hydrogen was circulated over the catalyst for 3 hr. The above oxidation-reduction cycle followed by standard degassing always removed species giving rise to CH bands. Spectra of the disk were determined with a model 112C Perkin-Elmer infrared spectrometer (modified for double pass) with NaCl optics. Normally, spectra were recorded with a spectral slit width varying from the order of 15 cm-' at 3500 cm-l to 3.5 cm-' at 1200 em-'. Kinetic Runs. Ethylene hydrogenation was carried out in a closed system with forced circulation of gases by a magnetically driven pump. Hydrogen-ethylene or deuterium-ethylene mixtures were made up as desired in the BET system and toeplered into the circulation loop. Circulation was started with the catalyst bypassed to mix the gases before the catalyst was cut into the reaction loop. The volume of the reaction system was 109 cc and the circulation rate (judged by the reaction rate with very fast catalysts) was about 300 cc/min. The fastest rate measured corresponded to about 3% per pass. The reaction was followed by periodic chromatographic analysis of small samples (-1%) of the gas in the circulation loop. All rates were computed from the initial slopes of plots of ethane formed os. time. Usually, the first few points were indistinguishable from a straight line through the origin (cf. Figure 5). Samples used for kinetic runs were first pressed into disks, then broken up t o improve the flow characteristics. Samples 6 and 9 weighed about 1 g; sample 10 weighed about 10 g. The initial activation consisted of (a) degassing while raising the temperature to 300", (b) circulation of hydrogen (-100 mm) over the catalyst at 300" and through a liquid nitrogen trap for 30 min, and (c) degassing for 3 hr at 450". This treatment yielded a The Journal of Physical Chemistry
sample with an area of 9.6 m2/g. In reactivation procedure (a) and (b) were repeated and the final degassing (c) was carried out at 400". (In a few runs the above activation procedure was carried out with oxygen rather than hydrogen; such a pretreatment is referred to as 0 2 activation.) Normally, the catalyst was degassed 1.5 hr at room temperature between runs and reactivated when periodic standard activity runs showed a decline in activity over 10%. During the course of this work catalysts 9 and 10 were subjected to prolonged treatment with oxygen at 500". After this treatment, the areas of these catalysts were 5.0 to 6.0 m2/g. Catalysts that were thus sintered bear the added label S, and in their reactivation steps a-b and c were carried out at 400 and 450", respectively. A comparison of the rates for different samples pretreated as described is given in Table I. Each of the ~~~~
Table I : Catalytic Activity of ZnO
Sample
PH?,mm
Zn0-6 Zn0-9 ZnO-9-S Zn0-10-S
244 130 130 300
Ratea/g Ratea/cm2 of cat., of cat., molecules/sec molecules/seo P c ~ Hmm ~, X 1018 X 10-11
24 13 13 110
3 04 2 97 1 82 1 96
3.2b 3.1 3.0 3 9"
a The experimentally determined orders (see later) were used to correct these rates to PH, = 130 mm and P c ~ H = ~13 mm. The area was not determined. Since the pretreatment was comparable to that of Zn0-9, the area was assumed to be the same. This catalyst, sintered by previous experiments to an area of 5 m2/g and was hydrogen activated and degassed 16.5 hr a t 450" prior to these runs.
listed numbers is the average value obtained for all runs on the freshly activated samples. There were nonsystematic variations in activity from one activation to another by as much as 20%, but the mean deviation from the average for each sample was less than 10%. The agreement of rates/cm2 for different samples (after corrections for differences in area, sample size and reactant pressures) demonstrates the reproducibility of these rate measurements. Adsorption measurements were made on a 10-g sample (Zn0-10) with a standard BET system. Adsorbed volumes were measured to within +0.002 cc/g unless otherwise specified.
Results Chemisorption of Hydrogen. Adsorption of hydrogen at room temperature and 128 mm on freshly activated zinc oxide is represented by the uppermost plot in Figure 1. Rapid initial adsorption is followed by a slow process which is detectable for several days. (After 24 hr the amount of hydrogen adsorbed is 0.30 to 0.34
HYDROGENATION O F
3775
ETHYLENE BY Z I N C OXIDE
I
0
006 0
1
I
I
20
40
60
1 80
T I M E , rnin. Figure 0, run A, run D, run
1. Adsorption of hydrogen vs. time ( P = 128 mm): 1, virgin cat,alyst; 0, run 2, after run 1 2 min degas; 3, after run 2 15 rnin degas; 5, after run 4 (16 hr adsorption) 30 rnin degas.
+
+-
+
cc/g.) Similar behavior has been noted by others, both for zinc oxide28--25 and chromia.26 Since the fast and slow processes appear to involve different modes of c h e m i ~ o r p t i o n , ~ we ~ - ~ ~looked for an experimental method of separating the fast and slow chemisorption. After the last point for run 1 was obtained, the sample was evacuated for 2 min, hydrogen was readmitted, and the amount of readsorption was measured as a function of time (curve 2). After the rapid readsorption of about 0.075 cc/g, the slow process resumes a t a rate equal to that when run 1was terminated. This process was followed for about 1 hr; a t this point the catalyst was evacuated for 15 rnin and the readsorption was measured (curve 3). Once again, after a rapid readsorption, the slow process resumes at a rate comparable to that a t the end of run 2. The catalyst was then exposed to hydrogen and the adsorption was followed for 20 hr; at this point, the slow process was undetectable over a period of 1 hr. Then, the catalyst was evacuated for 0.5 hr and the readsorption was followed (curve 5). I n this run slow adsorption is not observed; hence, the amount of rapid adsorption is the total adsorption. The amount of rapid readsorption estimated from run 5 (0.112 cc/g) is quite close to the zero-time intercept for run 3 (0.113 cc/g) but substantially above the zero-time intercept for run 2 (0.074 cc/g). The above experiments suggest: (a) fast chemisorption is removed by evacuation for more than 15 min; (b) the fast chemisorption is independent (compare runs 3 and 5) of the amount of slow chemisorption. Further support for conclusion (a) is supplied by the data in Figure 2 obtained as follows. First, the catalyst was “saturated” with slow hydrogen chemisorption by adsorption for 24 hr; total adsorption at this point was about 0.35 cc/g. Then, the catalyst was evacuated for the periods indicated and the readsorption determined
I
I
20
I
I
40
I 60
T I M E min.
Figure 2. “Fast” hydrogen adsorption vs. degassing time ( P = 135 mm): 0, after 16-36 hr adsorption; A, intercepts of Figure 1, runs 2 and 3.
from data similar to that shown for run 5 (Figure 1). These data clearly support the conclusion that the “fast” hydrogen is removed by evacuation for 15 to 60 min without removal of the “slow” hydrogen. With the above procedures we can experimentally define the amount of “fast” and “slow” chemisorption. It seems clear, however, that some of the “slow” hydrogen occurs rapidly. For example, the amount of adsorption on a freshly activated catalyst after 2 min (Figure 1) is 0.154 cc/g. Of this adsorption (in terms of our operational definitions) at most 0.112 cc/g is classified as “fast” chemisorption; hence, at least 0.042 cc/g of “slow” chemisorption occurs in 2 min. Accordingly, to avoid a misnomer, we shall designate the “fast” chemisorption as type I and the “slow” chemisorption as type 11. It should be emphasized that whereas all slow chemisorption is of type 11, not all type I1 chemisorption occurs slowly; in contrast to this, all type I chemisorption is rapid. We can utilize the above procedure to obtain isotherms for the type I hydrogen. The catalyst is first saturated with type I1 hydrogen by overnight adsorption. Then, after evacuation for 0.5 hr, an isotherm is measured. This measurement is completed within 1 hr; hence, in principle, the slow process should not contribute. Figure 3 shows such a plot. Open circles are adsorption points obtained by increasing the pressure from 35 to 370 mm; solid points are obtained by decreasing the pressure from 275 to 7 mm. The agreement of adsorption and desorption points suggests that type I adsorption, as defined, can be treated as a weak reversible chemisorption with a saturation coverage (reached above 40 mm) corresponding to about 5% of the V , value. The above experiments provide an operational basis (23) H. S. Taylor and C. 0. Strother, J . Amer. Chem. Soc., 5 6 , 686 (1934). (24) V. Kesavulu and H. S. Taylor, J . Phys. Chem., 64, 1124 (1960). (25) M. J. D. Low, J . Amer. Chem. Soc., 87,7 (1965). (26) R. L. Burwell, Jr., and H. 8. Taylor, ibid., 58, 697 (1936). Volume 78, Number 11 November 1960
A. L. DENTAND R. J. KOKES
3776
0
IO0
200
300 0
Prnrn H,
Figure 3. "Fast" hydrogen isotherms: 0, hydrogen activation; A, oxygen activation (solid symbols denote desorption points).
for the separation of the two types of chemisorption but, for information on structure, we must turn to infrared studies. Eischens, Pliskin, and I,ow18 have shown that chemisorption of hydrogen on zinc oxide at room temperature produces two bands, one at 3510 cm-' and the other at 1710 cm-l. The former can be assigned to OH and the latter to ZnH. We have also observed these bandsz7and, as Eischens, et al.,'8 have reported, we find: (a) the ratio of OH to ZnH band intensity is a constant except at the highest coverage where the OH band becomes relatively more intense; (b) the bands disappear on brief evacuation; and (c) the bands increase in intensity with increasing pressure up to 40 mm; at higher pressures, they do not show a pronounced increase in intensity. With procedures established by the chemisorption experiments, it is possible to examine the infrared spectra due to both types of chemisorption. In the absence of type I chemisorption but when type I1 chemisorption is about 0.24 cc/g ; i.e., after adsorption for 24 hr followed by a 30 min evacuation, no bands are observed, but when about 0.10 cc/g of type I hydrogen is present, the bands exhibit their maximum intensity. This observation together with the pressure dependence of the band intensity compared to that for type I chemisorption (Figure 3) provides compelling evidence that the bands arise solely from type I chemisorption as suggested by Eischens, et al.Is When a degassed sample of zinc oxide is exposed to oxygen at room temperature, the background transmission increases by a factor of 2 to 4 in the wavelength region of interest. This background transmission remains unchanged after evacuation for 1 hr at room temperature ; hence, this change is due to chemisorption of small amounts of oxygen (-0.005 cc/g). Thomasz8 has shown that for single crystals, the absorption is directly proportional to the number of carriers. Accordingly, this change in transmission is consistent with the observations that chemisorption of small amounts of oxygen brings about a dramatic reduction in the conductivity.29 Exposure to hydrogen at about The Journal of Physical Chemistry
250" makes the sample opaque even after it is cooled to room temperature. This high-temperature chemisorption is that responsible for the high-temperature maximum in the isobarz3and has been designated as type B by Kesavulu and TayloraZ4In accord with infrared results, type B chemisorption increases the ~ o n d u c t i v i t y . ~By ~ way of contrast, neither type I nor type I1 chemisorption produces changes in the background transmission. Thus, it seems clear that the slow process at room temperature is different from the high temperature type B chemisorption. Some support of this view is afforded by the reportz4that desorption of all the hydrogen chemisorbed at room temperature can be effected by evacuation at 150" whereas desorption of type B chemisorption requires temperatures in excess of 250". I n conclusion, then, we disagree with the conthat "The protonic adsorpclusions of Eischens, et tion which is responsible for the high-temperature maximum of the isobar occurs to a limited extent at room temperature as a slow chemisorption." It should be noted, however, that this disagreement is based on further data and does not result from conflicting data. The effect of oxygen on hydrogen chemisorption is of interest because of the reported poisoning effect of oxygen on the catalytic activity of zinc oxide for hydrogenation13-ls and hydrogen-deuterium exchange.12 It was, therefore, somewhat surprising to find that although chemisorption of oxygen at room temperature had a pronounced effect on the conductivity, it had little effect on the integrated intensity of the infrared bands. This result is confirmed by the chemisorption studies. Figure 3 includes points for an isotherm obtained after the catalyst was subjected to dry oxygen treatment for 2 hr at 400" followed by 0.5
Table 11: Effect of Pretreatment on Hydrogen Adsorption Pretreatment
VI, c c / g a
VI[! cc/gb
Activation Degas 1hr 400°,2 hr O24OOo, 0 . 6 hr degas 300' Degas 1 hr 400°, 2 hr 0 2 400°, 0.5 hr degas 150' Degas 1 hr 475O, activation, 02 25O, 0.5 hr degas 25"
0.114
0.072
0.110
0.082
0.110
0.054
0.095
O.OG2
Volume adsorbed a t 100 mm after saturating with type 11Hz. adsorbed 15 min after pretreatment minus the amount in column 2. This figure is thus a rough gauge of the amount of rapidly occurring type I1 chemisorption. a
' Total volume
(27) W e have also observed the ZnH band on the SP 500 zinc oxide from New Jersey Zinc Co. (28) D.G.Thomas, J . Phgs. Chern. SoZids, 10,47(1959). (29) See, for example, R. Glemza and R. J. Kokes, J. Phys. Chern., 69, 3254 (1965). (30) Y.Kubokawa and 0. Toyama, (bid.,60,833 (1956).
3777
HYDROGENATION OF ETHYLENE BY ZINCOXIDE Table 111: Participation of Type I1 Hydrogen in Hydrogenation 87
64
Pretreatment hr. Initial GH,, cc Initial HB,cc Ethane formed, cc Adsorbed Dz, ccC Dz in ethane DBin hydrogen Dz on surface
16, 0 . 5 1.86 3.65 0.73 0.20 0.014 0.043 0 . 132d
16, 1.5 1.50 16.69 0.87 0.20 0,005 0.027 0.178
16, 0 . 5 1.51 16.43 0.89 0.20 0.017 0.036 0 . 15e
3 . 5 , 0.25 1.87 4.10 0.77 0.14 0.012
... ...
0.5b 1.71 19.25b 0.61 N O . 08 0.61
...
0.028d
a The first figure pertains to the exposure time to deuterium prior to the run and the second to the evacuation time prior to the run. In this run the catalyst was activated and a 0.5-hr hydrogenation run was carried out. After evacuation for 1.5 hr, the amount of deuterium on the surface was determined by exchange. c This figure is an estimate of the amount of type-I1 adsorption based on data for a These figures were dedifferent larger sample. Since the catalyst in runs 87-89 was sintered the type I1 adsorption is probably less. termined by the measured exchange after 24-hr contact with 200-500 mm of Hz. e These figures were determined by mass balance.
b
hr degassing at 400" ; within experimental error the points fall on those for the hydrogen activated sample. In an effort to see some effect, the sample was subjected to more and more intensive oxygen treatment. These results are summlarized in Table 11. There is a gradual decline in adsorption (ascribable to sintering) but the effect of oxygen pretreatment on hydrogen chemisorption is negligible for type I adsorption and not very large for type I1 adsorption; hence, the reported poisoning effect of oxygen cannot be ascribed t o inhibition of hydrogen chemisorption. (It is worth noting that hydrogen activation followed by exposure to oxygen at 25" produced the biggest effect on type I adsorption. This decrease (last entry, Table 11) may be due to gradual sintering or it could be evidence of the interaction of residual adsorbed hydrogen with oxygen to block out type I sites.) Ethylene Chemisorption. Adsorption of ethylene on zinc oxide at room temperature is rapid and reversible; an isotherm is shown in Figure 4. On the time scale of these experiments (ca. 1 hr), we found no evidence for the dimerization reported by Ozaki and coworkers. From the amounts of adsorption a t higher pressures, i t appears that the number of sites for ethylene chemisorption is about five times the number available for type I hydrogen chemisorption, (The solid line indicates the extent of type I chemisorption.) The effect of oxygen pretreatment on ethylene chemisorption was found to be small (Figure 4); hence, the reported poisoning effects of oxygen cannot be ascribed to inhibition of ethylene chemisorption. The Role of Hydrogen in Hydrogenation. Hydrogenation of ethylene over zinc oxide at room temperature occurs at rates of the order of 0.05 cc/min. This is slower than the rate of type I adsorption and comparable t o the initial rate of type I1 adsorption; hence, on this basis alone, either or both types of hydrogen chemisorption could be intermediates in the hydrogenation of ethylene.
r
0.40
m \ u
V
0.20
Y
-
0
100
200
300
400
Prnm C,H4
Figure 4. Ethylene Isotherms: 0, hydrogen activation: (The horizontal line indicates the amount of type I hydrogen adsorption.)
0, oxygen activation.
The relative participation of type I and type I1 chemisorption can be assessed by tracer techniques. Consider the data for run 64 in Table 111. The catalyst was first exposed to deuterium for 16 hr then evacuated for 0.5 hr. This leaves about 0.20 cc of type I1 Dz on the surface. Then a regular hydrogenation run is made with Hz:CzH4 = 27 mm: 14 mm. At the end of this run 0.73 cc of hydrogen had reacted, but even though this represents a turnover of reacting hydrogen on the surface nearly fourfold greater than the type I1 deuterium adsorption, analyses of the ethane shows only 1.9% of the product (0.014 cc) forms by reaction with Dz. Had all of the adsorbed deuterium reacted, 27.4% of the product (as CzH4Dz) would have formed via reaction with Dz. Analyses of the unreacted hydrogen showed the deuterium content was about, 1.5%; hence, most of the ethane produced could have (31) A. Ozaki, H. Ai, K. Kimura, presented at the Fourth International Congress on Catalysis, Moscow, 1968.
Volume 78, Number 11 November 1060
A. L. DENTAND R. J. KOKES
3778 resulted from exchange of deuterium with hydrogen and its subsequent reaction. Exhaustive exchange of the catalyst with hydrogen after the reaction was complete accounted for 0.132 cc/g of deuterium, presumably still present as type I1 after reaction is complete. (Note that the sum of the last three entries in column 2 is 0.191 cc/g, in good agreement with the amount of type I1 chemisorption, 0.20 cc/g, estimated from adsorption experiments.) Run 88 (Table 111)was similar to run 64 but here the hydrogen pressure was increased to see if the type I1 chemisorption was labilized by the higher hydrogen pressure (120 mm). Once again, about 1.9% of the ethane formed from Dz even though there was a turnover of reacting hydrogen on the surface more than four times the amount of type I1 chemisorption. Run 89 shows that if the catalyst is evacuated 1.5 hr after saturating with deuterium, the amount of incorporation of type I1 deuterium in the product is reduced; run 61 shows that if the amount of t,ype I1 adsorption and the degassing time is decreased the incorporation of type-I1 deuterium in the products is still comparable to that for run 64. Run 87b represents a basically different kind of isotopic experiment. Here a run with pure deuterium was carried out, and after evacuation the amount of type I1 adsorption was measured by exhaustive exchange. I n the absence of ethylene, the amount of type I1 adsorption in 0.5 hr required for a run would be about 0.08 cc/g. I n the presence of ethylene, the amount measured by exchange is 0.028 cc/g. Thus, although type I1 adsorption occurs in the presence of ethylene the amount is smaller. This reduction in type I1 adsorption may be assigned (in part) to the incorporation to form ethane. In sum, the above data show that the predominant pathway to the hydrogenation of ethylene involves type I chemisorbed hydrogen and that the rate of type I1 chemisorption is reduced by a factor of 3 by the presence of ethylene. Hydrogen Promotion. Although type I1 hydrogen
Table IV : Hydrogen Promotion Run no.
Rateb mm/min
Pretreatmenta
9-62
Hp activation
9-63 9-S-75
Hz promotion €12 activation
+
0 2
R.T.
0.673 0.233
94-76 9-5-80
1% promotion 0%activation
+
0 2
R.T.
0.314 0.220
9-5-81
132 promotion
a
See text.
0.533
0.270
Rate under standard conditions;
= 130:13mm.
The Journal of Physical Chemistry
60
0
0 - 1
.-4
0-3
A-6
4 0
8
I2
4
8
T I M E , min.
Figure 5 . Ethane formation vs. time ( P H ~ : P c ~=H300: ~ 115 mm); 0, run 1, Hp activation; A, run 2, repeat; 0 , run 3, after hydrogen promotion; 0, run 4, repeat; . , run 5, after hydrogen promotion; A, run 6, repeat. (The lower abscissa is for runs 4, 5, and 6.)
chemisorption does not participate in the hydrogenation of ethylene it does promote the rate of reaction. This effect is documented in Table IV, which lists rates as a function of pretreatment. In column 2 “activationJ’ denotes a high temperature pretreatment (see Experimental Section) ; the term “Hz promotion” denotes saturating the surface with type-I1 chemisorption; the term “ 0 2 R.T.” denotes exposing the catalyst to dry oxygen at room temperature followed by 0.5 hr degassing at room temperature. Data in Table IV show that hydrogen promotion increases the rate by about 30%. This promotion is evident for a catalyst subjected to standard pretreatment (9-62, 9-63), for a catalyst with chemisorbed oxygen (9-S-75, 9-5-76), and for a catalyst activated in oxygen (9-S-80, 9-S-81). Moreover, the effect seems to be about the same if the promotion is carried out with deuterium rather than hydrogen. Although the effect of hydrogen promotion is reproducible, it only lasts for one run. Figure 5 shows a sequence of six runs on sample Zn0-10 carried out at effectively the same pressures of hydrogen and ethylene. The numbers denote the order in which the experiments were performed. Run 1 was immediately after activation and run 2 mas a repeat run made after the standard run with room temperature evacuation between runs. Promotion Prior to run 3, the catalyst was saturated with type 11 hydrogen, i e . , promoted. The rate in run 3 is about 1.33 25y0 greater than the average for unpromoted runs. Then, after standard evacuation, run 4 was performed and the rate, within experimental error, returned to 1.26 that characteristic of unpromoted runs. Promotion with hydrogen prior to run 5 again yielded the promoted 1.23 rate, but the repeat run 6 yielded the unpromoted rate. Thus, although saturation of the surface with type 11 i.e., Hz:CZH~ hydrogen promotes the rate for the following run, the effect does not persist thereafter even though only
3779
HYDROGENATION OF ETHYLENE BY ZINC OXIDE 40
.-cE
30
\
E E W
ta U
20
X
0
-
0
IO
-LLL R U N NO.
68697(
727374
757677
78
84
Figure 6. Effect of oxygen pretreatment on rate. Each arrow represents hydrogen activation. Cross-hatched columns indicate rates for a freshly activated catalyst or one subjected to standard degassing. Open columns that are labeled represent rates after treatment with dry oxygen for 1 hr a t the indicated temperature, cooling to room temperature, and degassing 30 min. The unlabeled open column represents the rate after hydrogen promotion.
about 10% of the type I1 hydrogen is removed by reaction with ethylene. Oxygen Poisoning. Oxygen poisoning effects are complex and depend on the details of pretreatment. This is illustrated by the bar graph in Figure 6 which summarizes rates for a sequence of experiments defined by run numbers. After the initial activation, two runs on a partly sintered catalyst under standard conditions (68, 69) yielded an average rate of 0.39 i 0.02 mm/min. This is a factor of 3 greater than that for a catalyst subsequently treated with oxygen at room temperature (70). Oxygen treatment a t successively higher temperatures for 1 hr by circulation over the catalyst and through a liquid nitrogen trap revealed that the rate decreased further after 250" treatment (to 0.075 mm/min) but increased a t still higher temperatures. This suggests that the poisoning may arise in part from the reaction of adsorbed hydrogen with oxygen to form water. In the next sequence, the freshly activated catalyst, presumably sintered by the 500" oxygen treatment, was exposed to oxygen at room temperature. This catalyst may have some chemisorbed hydrogen due to the pretreatment but should have less than that used in run 70 on which two hydrogenation runs had been made. Despite the presumed sintering, the reaction rate in this run (75) was nearly twice that in run 70 with similar oxygen pretreatments. To estimate the effect of interaction of oxygen with chemisorbed hydrogen, the catalyst was exposed to hydrogen overnight and run 76 was made. This yielded the already noted promotion by hydrogen insofar as the rate was increased by about 30%. This rate is a factor of more
than 3 greater than the rate after subsequent exposure to oxygen (77). Thus, the poisoning effect of oxygen is much greater in the presence of chemisorbed hydrogen than in its absence. Following this sequence, two more runs were made on a freshly activated catalyst (78,84). These measured rates are consistent with the supposition that treatment of the catalyst with oxygen at 500" in run 74 resulted in some sintering. The above experiments strongly suggest that oxygen poisoning stems mainly from interaction with chemisorbed hydrogen. Possibly this interaction involves the formation of water, known to be a strong poison." Based on this sequence alone, the effect of oxygen alone is somewhat uncertain because of the possibility that some hydrogen is retained on the catalyst by the activation procedure. I n one set of experiments prior to the sequence in Figure 6, however, the catalyst was first activated by circulation of dried oxygen over the catalyst at 505" for 2 hr and cooling to room temperature in dried oxygen. With this pretreatment there should be no chemisorbed hydrogen, although sintering would be expected to occur. After degassing a t room temperature for 30 min the rate on this catalyst was found to be 0.378 mm/min. The catalyst was then subjected to the standard hydrogen activation. Since this activation was at lower temperatures than the oxygen activation, no sintering was expected. The rate on this catalyst was 0.370 mm/min. Thus, it appears that the poisoning due to dry oxygen is almost entirely due to interaction with chemisorbed hydrogen; if chemisorbed hydrogen is not present, poisoning by exposure to oxygen a t elevated temperatures or room temperature is not observed.
Discussion From these experiments a rather detailed picture of the hydrogenation process emerges. Adsorption of type I hydrogen, rapid and reversible, occurs on the special sites constituting about 5% of the available surface. (This assay is based on the assumption that chemisorbed hydrogen and physisorbed nitrogen have the equal effective cross sections.) The nature of these sites is open to question but adsorption on these sites gives rise to Zn-H and OH bands in the infrared. Moreover, it is this hydrogen species that is primarily responsible for the hydrogenation of ethylene. Type I1 hydrogen adsorption exceeds the amount of type I adsorption by a factor of a t least 2; e.g., after overnight exposure to hydrogen, type I adsorption is 0.11 cc/g and type 11is 0.23 cc/g. Precise assessment of the capacity of the catalyst for type I1 hydrogen is not possible because adsorption continued for days and saturation was never achieved. Type I1 hydrogen contributes little to the hydrogenation of ethylene and does not give rise to Zn-H and OH bands. Moreover, although it may be an intermediate in the formation of the species responsible Volume 75,Number 11 November lQ69
A. L. DENTAND R.J. KOKES
3780 for the high-temperature maximum in the isobar, it is experimentally distinguishable from this species. Even though type I1 chemisorption is unreactive, it does modify the catalyst and promotes the rate of reaction. This promotional effect does not persist after exposure to ethylene in a hydrogenation run even though on!y about one-fourth of the type I1 chemisorption is removed by such treatment. There is no experimental basis for speculation on the mechanism of this promotion but similar promotional on poisoning effects have been noted for metallic catalysts.a2 During the course of a hydrogenation run, type-I1 adsorption occurs concurrently. This may result in a continual promotional effect which extends the initial linear portion of plots of yield vs. time. Such effects will result in the apparent order of the reaction changing during the course of the reaction, a result noted by Aigueperse and Teichner.16 It seems clear that if the order for the unpromoted catalyst is desired, the order must be based on initial rates and initial pressures. It should be emphasized that the above conclusions are restricted to room temperature. Participation of type I1 chemisorption may be expected to be more important a t higher temperatures. Perhaps such participation, the extent varying with temperature, is the source of the complex kinetics observed by BozonVerduraa and Teichner.Ip Oxygen chemisorption at room temperature or elevated temperatures clearly decreases the conductivity (as reflected by the background transmission). Previous workers have reported that oxygen poisons hydrogenation reactions;12-14 for example, Taylor and Wethingtonla note that exposure to oxygen at 360’ decreased the rate to less than 1%of the value before such treatment. Understandably, this has led these authors to the c o n c l u ~ i o n “It , ~ ~ is clear that catalytic activity, like electrical conductivity is closely related to oxygen deficiency in the oxide.” Later experiments by Aigueperse and Teichner14 showed that doping with lithia or gallia had little effect on the activity, a result that led to the conclusions, “The electronic structure of the catalysts is therefore without influence on the catalytic activity.]’ However, on the basis of poisoning experiments with oxygen, they also noted that “The formation of a nonstoichiometric oxide is a necessary condition for the catalytic activity in the hydrogenation of ethylene.” Our results, however, show that exposure of the catalyst to dry oxygen at high temperatures or room temperature has a very small effect on the cata-
The Journal of Phyadcal Chemistrzl
lytic activity at room temperature provided the amount of chemisorbed hydrogen on the catalyst is small during oxygen pretreatment. The fact that oxygen pretreatment of the catalyst has little effect on either type I hydrogen adsorption or ethylene adsorption is in line with these activity effects. Since oxygen pretreatment is supposed to reduce the nonstoichiometry, this clearly implies that nonstoichiometry is not a requirement for catalytic activity. It seems more likely to us that the reported poisoning with oxygen is due to interaction with adsorbed hydrogen to form water at the active sites. Thus water (or its precursor) is the poison rather than the oxygen. Similar observations have been made by Hindin and Weller on aluminaaa insofar as they find water poisons the hydrogenation activity whereas dry oxygen has little effect. The active sites for hydrogenation on chromia and alumina have been associated with “strained sites” created by dehydration of a hydroxylated oxide surface.22*saCreation of these sites for a surface terminating in close-packed hydroxyl groups can be represented as H
H
H
H
I I I I 0 0 0 0 I I I I
-M-M-M-M-
--j
H
H
I 0 0 I I
0
--M-M--M--R/I--
I I
+ H2O
Thus, these strained sites can be viewed as surface anion vacancies. There is reason to believe the active site on zinc oxide is similar to the one pictured above. Such a site would be isolated and noninteracting and would be poisoned by exposure to water by reversal of the foregoing equation. Some aspects of the chemistry discussed in the next paper, however, suggest the surface between active sites consists of oxide ions rather than hydroxyl ions. Such a picture has the same qualitative consequences insofar as rehydroxylation by water would still destroy the active site, i.e., poison the catalyst, and the sites would be isolated. Perhaps this modification applies to chromia and alumina as well.
Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for support of this research. (32) W. K. Hall and P. E.Emmett, J.Phw. Chem., 63, 1102 (1959). (33) 8.G. Hindin and S. W. Weller, Advan. Catal., 9,70 (1957).
