3781
HYDROGENATION OF ETHYLENE BY ZINC OXIDE
Hydrogenation of Ethylene by Zinc Oxide.
11.
Mechanism and Active Sites by A. L. Dent and R. J. Kokes Department of Chemistry, The Johns Hopkins University,Baltimore, Maryland
81818 (Received March 84, 1060)
The rate of hydrogenation of ethylene over zinc oxide at room temperature has been determined for hydrogen pressures ranging from 23 to 583 mm and ethylenepressures ranging from 5 to 330 mm with the ratio of hydrogen to ethylene ranging from 1to 50. Initial rates were proportional to the hydrogen pressure to the one-half power and showed a slight dependence on ethylene pressure ascribable to the extent of the surface covered with ethylene. The kinetic isotope effect was also determined for a fixed set of conditions. These measurements,supplemented by infrared studies of the sample either with chemisorbed gases (CZH~, 02, CO) or under reaction conditions and an examination of poisoning effectsof water vapor provide the basis for a model of the active site consisting of zinc ions imbedded in a close-packed oxide layer. This tentative model leads to a mechanism consistent with the kinetics and provides a plausible explanation for most of the experimental observations. It further suggests that there should be striking similarities between hydrogenation over oxides and hydrogenation over homogeneous catalysts. Introduction Nearly 30 years ago, Woodman and Taylor' discovered that zinc oxide is an effective catalyst for the hydrogenation of ethylene. Since this discovery, a number of reports'-' have appeared dealing with this aspect of the chemistry of zinc oxide, but the only thorough studies of kinetics have been carried out by Teichner's g r ~ u p ~ -at ~ Jelevated temperatures, Le., roughly 100 to 400". Under these conditions, the reaction is quite complex: orders and activation energies change from one temperature range to a n ~ t h e r and ,~ formation of ethylene residuess with consequent poisoning' is evident. Since these studies span a range of temperatures in which the adsorption of hydrogen is a complex phenomenon,*-1° complexities of the related hydrogenation kinetics are to be expected. Nevertheless, the results are such that they frustrate attempts to put forth a coherent picture for the mechanism of ethylene hydrogenation. Recently, it has been shown in this laboratory1l!l2 that room temperature catalysis of hydrogenation by ~~ zinc oxide is similar to that of ~ h r o m i a . ' ~In~ particular, although the mechanism of hydrogenation over these oxides may involve the same surface processes believed to occur with metal catalysts,15i.e. C2H4
+ €3
CzHs 4- H
-
C2H.5
C2He,
these oxides lead to CzH4D2 on deuterogenation rather found with metals. Presumably, than the CZHO-~D, this means the first step is irreversible over oxides, i.e., "alkyl reversal" does not occur, whereas on metals the first step is quite reversible. It has been suggested that this behavior of oxides is a consequence of the noninteraction of widely separated active sites. l 2 t l 3
The clean, 1,Baddition of deuterium to olefins over oxides is not found at elevated temperatures.14 Perhaps the specificity observed stems from the factI6 that at room temperature hydrogenation involves only the rapid reversible hydrogen chemisorption (type I). The irreversible chemisorption (type 11))also present at room temperature but not effective in hydrogenation, may play a more important role at elevated temperatures and lead to a process involving two types of reactants, type I and type I1 chemisorption. I n this event, one would expect the observed complex kinetics. It is the purpose of this paper to examine in detail the kinetics and mechanism of hydrogenation of ethylene over zinc oxide. The major emphasis is on the reaction (1) J. F.Woodman and H. S. Taylor, J. Amel,. Chem. SOC.,62, 1393 (1940). (2) E.H. Taylor and J. A. Wethington, ibid., 76,971 (1954). (3) J. Aigueperse and 8. J. Teichner, Ann. Chim. (Paris), 7, 13 (1962). (4) J. Aigueperse and S. J. Teichner, J . Catal., 2, 359 (1963). (5) F. Boson-Verduraz, B. Arghiropoulos, and 8. J. Teichner, Bull. SOC.Chim. Fr., 2854 (1967). (6) D. L. Harrison, D. Nicholls, and H. Steiner, J. Catal., 7, 359 (1967). (7) F.Bozon-Verduraz and 9. J. Teichner, ibid., 11,7 (1968). (8) H. 8. Taylor and C. 0. Strother, J . Amer. Chem. SOC.,56, 586 (1934). (9) V. Kesavulu and H. A. Taylor, J. Phys. Chem., 64, 1124 (1960). (10) M.J. D. Low, J . Amer. Chem. SOC., 87,7 (1965). (11) W. C. Conner, R. A. Innes, and R. J. Kokes, ibid., 90, 6858 (1968). (12) W. C. Conner and R. J. Kokes, J . Phys. Chem., 73, 2436 (1969). (13) R. L. Burwell, Jr., A. B. Littlewood, M. Cardew, G. Pass, and C. T. H. Stoddard, J . Amer. Chern. SOC.,82,6272(1960). (14) A. B. Littlewood and R. L. Burwell, Jr., ibid., 82, 6287 (1960). (15) G. C. Bond, "Catalysis by Metals," Academic Press, London, 1962,pp 258-270. (16) A. L. Dent and R. J. Kokes, J . Phys. Chem., 73,3772 (1969). Volume 7.9,Number 11 November 1060
A. L. DENTAND R. J. KOKES
3782 at room temperature since this is the temperature where the aforementioned simplicity is obtained. The classical techniques for kinetic studies are supplemented with infrared techniques in an attempt to gain some information on the nature of the active sites. Thus, this paper represents an extension of the results reported in part I of this series.
1.00
1 o
!
.-c
E
Experimental Section Experimental details regarding materials, rate measurements, and infrared studies were described in a previous paper. Measurement of the effect of adsorbed water on the rate and adsorption were carried out on sample Zn0-10 in the circulating system. This sample weighed 9.5 g; it was slightly sintered in experiments preceding this so that the area was about 5 m2/g compared to 10 mz/g for a fresh catalyst. Preliminary experiments showed that when the activated catalyst was exposed to about 2 cc of water vapor at room temperature, the pressure rapidly dropped to zero. Helium was then admitted to the circulating system and circulated through a liquid nitrogen trap in series with the sample. Contents of the trap were checked by evacuation and warmup every 30 min to see if water evolved. When this was done as a function of catalyst temperature, it was found that there was little water evolution below 300". By heating between 300 and 450" all of the water could be recovered. For runs reported herein, about 3 cc of water vapor was admitted to the activated catalyst. The catalyst was then conditioned by circulating helium over the catalyst of 300" for 1 hr. Since previous experiments have shown that water desorbed at 300", this treatment should yield a uniform distribution of water throughout the catalyst bed. The catalyst was cooled to room temperature, exposed to hydrogen at about 200 mm, and the rate of adsorption was followed for 5 min. After this the catalyst was degassed for 2 hr and the rate of ethylene hydrogenation was measured for CzH4: Hz = 110:310 mm. At this point, the catalyst was degassed for 2 hr at room temperature and the temperature was raised and the water evolution measured. This sequence was repeated until all of the water was removed. I n another sequence, ethylene chemisorption and rate were measured as a function of adsorbed water by the same procedures. Results Kinetics. Rates were determined from the initial slopes of plots of conversion us. time and correlated to initial pressures of reactants. Tests of order based on the changes in the reactant pressure during the course of the reaction were not attempted because the continuous promotion by type I1 chemisorption would be expected to frustrate such attempts. l6 Typical plots of conversion us. time are shown in Figure 4 (to be disThe Journal of Physical Chf?mi&Tg~
\
0.50 (u
+ 0 [L
0.0 I
'
0.0
I
I
10
I
I
I
20
Figure 1. Order with respect to hydrogen: 0, sequence SA, P Q H ~= 12.7 mm, scale factor X 1; 0, sequence 9B, P c ~ H =~ 27.4 mm, scale factor X 1.12; A, sequence 9C, P c ~ = H 71 ~ mm, scale factor X 0.90; A, sequence QD, PC~H( = 129 mm, =~ 24 mm, scale scale factor X 0.87; 0 sequence 6A, P c ~ H factor X 0.85.
cussed later). It can be seen that such initial rates are well defined. During the course of these runs the catalyst was accidentally poisoned several times and reactivation was required. Such reactivation caused a nonsystematic shift in the standard activity by as much as 20%, but the mean of standard activities after all reactivations showed a mean deviation of less than 10%. The order with respect to hydrogen, determined for several different activation sequences at fixed ethylene pressures is illustrated by Figure 1. Since there is some dependence on ethylene pressure and the standard activity varies somewhat from one sequence to another, we have multiplied the rates for each sequence by a scale factor. These sequences, which represent a 10fold range of ethylene pressures and a 25-fold range of hydrogen pressures, combine to give a well-defined straight line passing (nearly) through the origin when the rates are plotted us. the square root of the hydrogen pressure. It is noteworthy that this plot includes data for two different samples of catalyst, Zn0-6 and Zn0-9. The dependence of rate on ethylene pressure is shown in Figure 2. In sequence 9 (circles) hydrogen pressure was not kept constant but was varied from 70 to 314 mm. Dependence on hydrogen was factored out by dividing the rate by the square root of the hydrogen pressure. We have also included in this plot (triangles) average values of rate/(PHZ)"' for other sequences (here no scale factor is used to bring the rates into agreement). The first four triangular points represent the sequences 9A, 9B, 9C, 9D that were used in Figure 1 to establish the order in hydrogen. As the error bars show, the average of deviations in the average values of rate/(PHJ1" were of the order of 6%. I n
HYDROGENATION OF ETHYLENE BY ZINC OXIDE
3783
7,
6.0.
c .-
E
. -. \ N
-c E E
4.0 .
c
N
I
N
a”
‘
W
I
c
a,
I I :
X
2.0 -
c
0
10
LL
0
0
IO0
200
300
X 0 0
0
Pc2H4mm
Figure 2. Order with respect to ethylene: 0 , sequence 9; A, average values for different sequences shown in Figure 1.
toto, then, this plot not only establishes the order with respect to ethylene, but furnishes a test of the order with respect to hydrogen. These studies do not include reactant pressures with ethylene in excess, but the ratio of hydrogen to ethylene varied from 1: 1 to 50: 1. The pressure ranges studied vary from 5 to 330 mm for ethylene and from 23 to 583 mm for hydrogen. The curve in Figure 2 is the same shape as an adsorption isotherm. Measurements of the adsorption isotherm for ethylene were reported in the first paper of this series (Figure 4 of ref 16). From these data and Figure 2 we can replot rate/(PHJ1” us. amount of adsorbed ethylene. Such a plot (Figure 3), although limited in range, is well represented by a straight line through the origin. Thus the empirical kinetic expression for hydrogenation of ethylene is
I
0.40
v
I 0 60
cc/g
Figure 3. Order with respect to adsorbed ethylene: 0, sequence 9; A, average values for different sequences shown in Figure 1.
80
60
(D
I
0”
40
8?
20
0
wherein 0 represents the equilibrium ethylene chemisorption, i.e., in the absence of hydrogen, corresponding to the ethylene pressure. We recognize, of course, that ethylene adsorption under reaction conditions need not be the same as the defined 0. Isotope Effects. Figure 4 shows a sequence of standard runs made to determine the kinetic isotope effect. Runs 57 and 59 were carried out at slightly different pressures; after correction for these differences, the average of these rates if 0.665 =k 0.017 mm/min. The corresponding rate for deuterium is 0.290 mm/min. Thus, the rate with deuterium is roughly half of the corresponding rate with hydrogen. Parravano, Friedrik and Boudart” have measured the rates of adsorption of deuterium and hydrogen on zinc oxide. These measurements pertain to the L‘slow” adsorption and show that there is no kinetic isotope effect. Thus, our observations of a normal isotope effect for the hydrogenation reaction provides additional evidencela that the rate of “slow” adsorption
I
0.20
5
10
15
M i NUTES
Figure 4. Hydrogen isotope effects: 0, run 57, hydrogen; 0 run 58, deuterium; 0, run 59, hydrogen; A, run 60, deuterium.
cannot be the rate controlling step in the hydrogenation. Infrared Xtudies. As shown by Eischens, Pliskin, and Low1*the reversible chemisorption of hydrogen on degassed zinc oxide gives rise to two infrared bands of comparable intensity at 3500 and 1710 cm-l; these bands are ascribed to OH and ZnH stretching bands formed by the dissociative adsorption of hydrogen. Degassed zinc oxide shows at least three bands in the 3500 cm-1 region, presumably due to structural OH groups in the solid, which bracket the chemisorbed OH (17) G.Parravano, H.G. Friedrik, and M. Boudart, J . Phys. Chem., 63, 1144 (1959).
(18) R.P.Eischens, W. A. Pliskin, and M. J. D. Low, J. Cutal., 1, 180 (1962). Volume 78,Number 11
November 1969
A. L. DENTAND R. J. KOKES
3784 H
2
2 cn I
H
c p
c
m
v -
l
a [1I
+
I
A
Figure 5, Effect of carbon monoxide on deuterium bonds: A, OD band, 84 mm Dt; B, ZnD band, 84 mm Dz; C,OD band, 84 mm Dt after adding 35 mm CO; D, ZnD band, 84 mm DZ after adding 35 mm CO. The bars indicate a Av of 10 cm-l. Transmission scales are the same for OD band but for drawings B and D, they are shifted.
band. Studies of chemisorption of hydrogen are not greatly disturbed by this overlap; nevertheless, for this reason, it is more convenient to make chemisorption studies with deuterium. Chemisorption of deuterium gives rise to bands at 2610 and 1227 cm-l corresponding to OD and ZnD (Figure 5A and B). In this region, the background due to degassed zinc oxide is smooth. At room temperature, adsorption of ethylene gave rise to no bands that were not masked by or ascribable t o bands of the unsaturated hydrocarbon. When the gas phase was removed by brief evacuation, no bands due t o chemisorbed hydrocarbon were observed. In particular, there was no evidence for the bands of chemisorbed ethylene observed by Bozon-Verduraz, Arghiropoulos, and Teichner6 at slightly higher temperatures. Our result is consistent with the adsorption experiments, which show rapid, reversible chemisorption of ethylene at room temperature. Exposure of the catalyst to carbon monoxide at room temperature for about 1 hr followed by brief evacuation of the gas phase gives rise to bands due to chemisorbed carbon monoxide similar to those observed by Amberg.1Q,20Preadsorption of such carbon monoxide followed by adsorption of deuterium yields ZnD and OD bands in the usual positions. Quantitative intensity comparisons were not made, but auxiliary experiments suggest that there was little effect on band intensities provided the catalyst was not exposed to carbon monoxide beyond 1 hr. Thus, strongly chemisorbed carbon monoxide, like chemisorbed oxygen, has little or no effect on the chemisorption of hydrogen. If, however, gaseous carbon monoxide and gaseous deuterium are both present, an effect is observed (Figure 5C and D). The band due to ZnD broadens and perhaps splits into several bands; because of the extreme broadening, it is difficult to say if the intensity deThe Journal of Physical Chemistry
Figure 6. Deuterium bands under reaction conditions: A, OD band before and after reaction; B, ZnD band before and after reaction; C, OD band during reaction; D, ZnD band during reaction. The bars indicate a Av of 10 ern-'. Transmission scales differ for OD and ZnD bands.
creases. By way of contrast, the OD band is virtually unaffected. Exposure to water vapor followed by evacuation, results in an increase in the intensity of the structural OH bands. Bands due to chemisorption of hydrogen (or deuterium) are not observed following this treatment. Moreover, the activity for ethylene hydrogenation on this sample is reduced by about one order of magnitude. Degassing at about 400°, however, brings about a decrease in the intensity of the structural OH bands, restores the activity of the catalyst, and the sample again shows OH and ZnH bands on chemisorption of hydrogen. As noted earlier, the cell was attached to a circulating system so that infrared observations could be made during the course of catalytic hydrogenation. In a typical kinetic run, a 10:1 mixture of deuterium and ethylene (total pressure, 220 mm) was introduced into the circulation loop with the catalyst bypassed. The ethylene was frozen out by circulation through a liquid nitrogen trap. Then the catalyst was exposed to deuterium and the spectra shown in Figure 6A and B were recorded. At this point, with circulation through the cell, the ethylene was vaporized and hydrogenation of ethylene began, The course of the reaction was monitored by observations of bands due to gaseous ethane and periodic chromatographic analysis of gas samples. Figure 6C and D shows spectra obtained during the hydrogenation. The band corresponding to ZnD broadens and shifts from 1232 to 1195 cm-' but the integrated intensity stays the same; the band cor(19) J. H. Taylor and C. H. Amberg, Can. J . Chem., 39, 535 (1961). (20) C. H. Amberg and D. A, Seanor, Proceedings of the 3rd International Congress of Catalysis, Amsterdam, 1964.
3785
HYDROGENATION OF ETHYLENE BY ZINC OXIDE responding to OD shifts from 2610 to 2624 cm-l and the integrated intensity drops to 60% of the initial value. The spectra remain the same during most of the reaction. At the last stages of the reaction, however, the spectra start to change back to their initial appearance. When the reaction is complete, the spectra are identical with those shown in Figure 6A and B. It is also found found that if at any point during the reaction the hydrocarbons are frozen out, the original spectra are obtained; when the hydrocarbons are revaporized, the spectra characteristic of reaction (Figure 6C and D) are obtained. It should be recalled that the intensity of the infrared bands is relatively independent of pressure in this The fact that intensities and appearances of the spectra are the same in 200 mm of deuterium (before reaction) as in 180 mm of deuterium and 20 mm of ethane (after reaction) simply means ethane has no effect on the deuterium bands. Poisoning Experiments. In these experiments, after an initial study of the chemisorption and activity, about 10 g of catalyst was exposed t o 2.3 cc of water, which was circulated over the catalyst at 300” for 1 hr. Adsorption and activity were then measured as a function of the amount of water removed by circulation of dry helium through a liquid nitrogen trap in series with the heated catalyst. Experiments with dried catalyst showed that the water collection procedure following the sequence of activity and adsorption experiments always yielded about 0.2 cc of water. This water is presumed to come from reaction of type I1 hydrogen to form water in the collection procedure. (Preliminary experiments show that either the slow adsorption during adsorption measurements or the hydrogen uptake during reaction would be expected to contribute 0.2 to 0.3 cc.16 The combined effect of these experiments in sequence has not been experimentally examined.) With the assumption that all type I1 hydrogen was converted to water, the amount of water added agreed with that removed within 1%. The effect of water adsorption on rate for two sequences of water adsorption and removal is shown in Figure 7. The area of this catalyst, which was sintered by previous experiments, was ab2ut 5 m2/g. If one assigns an effective area of 10.6 A2 to a water molecule,21monolayer coverage corresponds to 14 cc. It is clear from Figure 7 that coverage of about 5% reduces the rate by about 50%. Figure 8 is a plot of rates us. adsorption for a catalyst poisoned with water to various degrees. Amounts of type I adsorption could not be estimated as in the previous runs for this would add of the order of 1 cc of type I1 hydrogen, convertible to water, in each run. Instead, the hydrogen adsorption after 2 min was measured. Preliminary experiments showed this included about 0.3 cc of type I1 adsorption. Accordingly, the total amount of hydrogen adsorbed in 2 min minus
0
0
I .o
2 .o
3.0
cc H,O Figure 7. Activity us. water adsorption: 0, hydrogen adsorption sequence; 0, ethylene adsorption sequence.
‘O’OI i
c c H, c c C,H,
Ads fast Ads X IO-’
Figure 8. Correlation of rate with adsorption: 0, hydrogen adsorption sequence; 0, ethylene adsorption sequence.
0.3 cc was taken as a rough measure of the amount of type I chemisorption. It appears that the rate reduction brought about by water adsorption arises primarily from the associated reduction of type I chemisorption. Figure 8 is consistent with this picture insofar as it implies that the rate reduction due to water adsorption parallels a concomitant reduction of type I chemisorption, whereas ethylene chemisorption is only little affected by water chemisorption. This conclusion does not depend on the procedure used to assess the amount of type I chemisorption; at most, a more precise assessment would bring only a scale shift in the abscissa of Figure 8.
Discussion Infrared and adsorption studies impose boundary conditions on the nature of the active sites for the (21) 9. J. Gregg and K. El. W. Sing, “Adsorption, Surface Area and Porosity,” Academio Press, New York, N.Y., 1967, p 82.
Volume 78, Number 11
November 1060
3786 hydrogenation of ethylene. The simultaneous development of ZnH and OH bands suggests the sites consist of a surface Zn-0 pair. Adsorption studiesl6 show type I chemisorption occurs only on 5% of the total surface; hence, the number of these pair sites is limited. Aloreover, since oxygen chemisorption has no effect on the type I adsorption or the associated infrared bands, the active sites are not associated with lack of stoichiometry. Accordingly, it is of interest to consider the surface structures expected for stoichiometric zinc oxide. In zinc oxide the oxide ions are arranged in positions corresponding to hexagonal closest packing with zinc ions filling one-half the tetrahedral holes (see Figure 9a).22 The idealized structure can be viewed in terms of spheres of zinc ions and oxide ions, both i c fourfold coordination, with a radius of 0.70 and 1.33 A, respectively. (These radii are 95% of the common octahedral radii.) If the spheres corresponding to oxide ions were strictly close packed, Le., if the oxide-oxide distance were fixed at 2.66 8,the zinc ions would have to squeeze into tetrahedral holes with a radius of 0.30 A. The computed molar volume for such a structure is 8.0 cc compared to the experimental molar volume 14.9 cc. If, instead, we assume the c1ose;packed layers of hard oxide spheres of radius 1.33 A expand isotropically until the tetrahedral holes c y accommodate hard zinc spheres with a radiusoof 0.70 A, the oxide-oxide distance would now be 3.30 A rather than 2.66 8. Then, the computed molar volume would be 15.2 cc in good agreement with the experimental value. Such a structure would be very open; only 45% of the space would be filled by spheres compared to 75% for the strictly closepacked structure. In addition to the tetrahedral holes in the oxide layers in zinc oxide, there are also octahedral and trigonal holes. The latter constitute a passageway from tetrahedral to octahedral sites. The assumed isotropic expansion increases the radius of the octahedral hole from 0.55 t o 1.0 8 and incyases the radius of the trigonal hole from 0.30 to 0.58 A. Thus, in the expanded structure which approximates the real structure, the trigonal hole can almost accommodate the zinc ion. It is of some interest to note that because of the hexagonal packing, the octahedral sites lie in a straight line perpendicular to the close-packed layer. Thus, we have a straight vhannel of such sites, each 2.0 8 in diameter, separated by a trigonal “squeeze point” 1.2 A in diameter. Figure 9ri provides a diagrammatic view of the layer structure in zinc oxide. The heavy solid lines represent the clcse-packed oxide layer; circles represent the interlayer zinc ions; the lighter lines represent the four “bonds” zinc ions make to oxide ions. Interchange of all of the zinc with all of the oxide ions yields the same structure; hence, positions of zinc ions like those of oxide ions, corresponded to an expanded, hexagonal The Journal of Phy8ical Chemistry
A. L. DENTAND R. J. KOKES
Figure 9. Structure of ZnO: A, schematio view perpendicular to the c axis. Heavy lines represent a nearly close-packed layer of oxide ions; circles represent zinc ions in tetrahedral holes; light lines represent tetrahedral “bonds;” dotted lines represent surfaces. (B)2lose-packed layers in 0001 surface (zinc uppermost) or 0001 surface (oxide uppermost). Larger spheres represent oxide ions. C, Layer structure for 10x0 surface. Larger spheres represent oxide ions.
close packing. The same drawing would also apply for the closely related zinc blende structure which differs from zinc oxide insofar as the anions show cubic rather than hexagonal close packing. The latter structure is commonly found for III-V compounds such as GaAs. For these compounds, development of close-packed planes occurs so that atoms on the surface have one rather than three dangling bond^.^^,^* Cuts yielding such surfaces are indicated by the dashed lines in Figure 9a. Studies with etched, polished single crystals reveal that on GaAs (or GaSb) one close-packed layer (111) contains only gallium atoms; the opposite one (111)contains only arsenic atoms. Because of the strong similarity in structure a similar situation holds for zinc there will be one close-packed layer containing only zinc ions (0001) and one containing only oxide ions (OOOT). The close-packed layers for either structure are shown in Figure 9b. (22) A. F. Wells, “Structural Inorganic Chemistry,” Oxford University Press, London, 1962, p 63. (23) H.C. Gatos, Science, 137,311 (1962).
(24) D. D. Pretzer and H. D. Hagstrum, Surfuce Sci., 4, 266 (1966).
HYDROGENATION OF ETHYLENE BY ZINCOXIDE Planes other than close-packed can be formed with only one (or less) dangling bond per surface atom. For zinc oxide, these planes are parallel to the c axis and can be designated as 1070 or the equivalent. Such a plane is pictured in Figure 9c. This surface consists of lines of raised zinc oxide pairs in which both the zinc and oxygen have a dangling bond. Between each line of pairs there is a groove defined by a subsurface line of pairs containing no dangling bonds. The analog of this surface in zinc blende is formed by a cut along the 110 planes. Clearly, these faces are not polar; they contain equal numbers of cations and anions. on single crystals of GaAs show that LEED the 110 plane at the surface has the same structure as in the bulk, but the 111 and 111planes show rearrangement in the surface planes. The nature of these surface reconstructions are indicated by guidelines in Figure 9b. The upper right hand corner shows the ideal two-dimensional unit cell, a 1 X 1 net. Reconstruction on the gallium surface expands this repeat unit to a 2 X 2 net (lower right) whereas reconstruction on the arsenic surface expands the repeat unit to a 3 X 3 net (left side).25 [Consider the gallium (111) surface to be represented by the drawing in Figure 9b with small spheres uppermost whereas the arsenic (111) surface has large spheres uppermost.] The driving force for the rearrangement of the 111 surface appears to be the tendency of the more electropositive component, which presumably has an unfilled dangling bond,23 to form an sp2 rather than sp3 bond. This can be approached if the gallium atoms in the surface move toward trigonal holes in the underlying arsenic layers. In GaAs the size of the trigonal holes is small compared to the size of the atoms; hence, such motion is accompanied by considerable strain. MacraeZ5proposes this strain results in loss of a superficial gallium atom or its replacement by an arsenic atom. Figure 9b shows how the systematic loss or replacement of the cross-hatched atom gives rise to a 2 X 2 net on the surface. Macrae25 has suggested that the formation of a 3 X 3 net on the 111 surface stems from the systematic introduction of strains at the points marked by x’s. It is believed23 that the more electronegative atoms have a filled dangling bond and therefore maintain sp3 bonding at the surface; hence, rearrangement is not the result of a driving force for rehybridization but stems from more subtle factors. The results described above for GaAs are also obtained for the isostructural GaSb; hence, it does not seem unreasonable to extrapolate these results to the nearly isostiuctuiral zinc oxide. On this basis we would expect no surface reconstruction on the 1070 planes, which are similar to the 110 surface in zinc blende, but might expect surface reconstruction on the close-packed planes. The 0001 planes, consisting of zinc ions with unfilled dangling bonds, will again tend to settle into
3787 the underlying trigonal holes and thereby form sp2 rather than sp3 bonds. Because the zinc ions can almost fit into the trigonal holes, the strain will be much less than that for GaAs; hence, the surface reconstruction found for GaAs may not be needed. Accordingly, we picture the 0001 planes as a layer of zinc ions, still uppermost, but closer to the underlying oxide layer than they are in the ideal structure. The OOOT planes, consisting of oxide ions with filled dangling bonds, will be subject to the same kinds of forces present in GaAs. R!tacrae2bhas suggested that the 3 X 3 can be the result of inclusions at the x’s in Figure 9b. We suggest that it is zinc ions that actually occupy these trigonal sites on the O O O i face of zinc oxide. These ions should not be viewed as deviat,ions from stoichiometry; they result from reconstruction of the surface to stabilize the 0007 face. Accordingly, zinc ions located at x’s in Figure 9b would not be removed by high-temperature pretreatment with oxygen if the OOO’T faces are to persist as a reconstructed surface phase. The picture of the active site as a zinc ion imbedded in a close-packed oxide layer has much to recommend it. The 3 X 3 repeat sequence would lead to sites that are not adjacent, and hence, are essentially noninteracting;12J3 moreover, it would assure that only a fraction of the surface contains active sites. Hydrogen adsorption on these sites occurs as I 1
I
-Zn-0-
+ H2(g)
H H I I 1 -Zn-0-
(In our model based on spheres, the outermost plane tangent to oxide spheres would be 0.53 A above the surface of the zinc sphere. To emphasize this point, we have shown the dangling bond of the zinc shorter than that of the oxide ion.) If, however, water is adsorbed on the zinc oxide first, we expect the following reaction to occur
H I
I _
I
-Zn-0-
1
H
? I + H20 +-Zn--O-
Adsorption of water vapor was shown to be strong; hence, preadsorption of water would be expected to prevent the occurrence of the first reaction, as observed. (We shall term the OH formed as a pair from water adsorption as hydroxyl in order to distinguish it from the OH formed by hydrogen chemisorption.) Hydroxylation would be expected to occur not only at active sites but also on the lines of ZnO pairs on 1OTO type faces (9c). The latter pairs, because of their protuberance above the surface, would be expected to retain hydroxyls more tenaciously than the active site wherein (25) A . U.MacRae, Surface Sei., 4,247 (1966).
Volume 73, N U T ~ ~ 11E TNovember 1060
A. L. DENTAND R. J. KOKES
3788 the zinc ion is less accessible because it is shielded by the larger oxygen anions. I n line with this three hydroxyl bands are seen even for a well degassed catalyst that shows an infrared spectra due to chemisorbed hydrogen. We would expect three hydroxyl bands from the line of hydroxylated ZnO pairs on the lOi0 faces corresponding to 0, 1, and 2 hydroxylated neighbors.26 Since these hydroxyls occur on different planes from the OD groups formed by deuterium chemisorption, exchange is very slow. I n addition, all hydroxyl bands including those associated with the active site occur at different frequencies from the OH bands simply because the environment is different ; e.g., hydroxyl groups occiir in pairs. The effect of chemisorbed carbon monoxide on the hydrogen chemisorption is also consistent with this picture. Irreversibly, slowly chemisorbed, carbon monoxide shows bands at 1575 and 1340 cm-l identified as carboxylate groups which are a precursor to reduction. l o r20 Under conditions of these experiments, strongly bound carbon monoxide covers about 5% of the oxide surface.20 If this chemisorption occurs randomly on the oxide sites of all exposed planes, it would be expected to have little effect on that 5% of the surface on which type I hydrogen chemisorption occurs. This agrees with our infrared data. There is also a reversible form of chemisorbed carbon monoxide, which also covers about 5% of the surface, and shows a band at about 2174 cm-l.zo This species is presumed to be found by a polar interaction. As such, it might be expected to concentrate a t the zinc ion of the active site. I n line with this, the ZnH band for chemisorbed hydrogen is dramatically influenced by this type of chemisorption with little effect on the associated OH band. Chemisorption of hydrogen can be represented by the set of equations
H H I -Zn-0-
H
I + -0-
H
I
1_ -Zn-0-
+ -0-
H H I --Zn-0-
H I
1
-Zn-0-
I
Jr-Zn-0-
H
I
I
+ -0I
(3)
1
-Zn-0-
1
I
+ -0-
l
at moderate concentrations, the OH band intensity would be proportional to the ZnH band intensity, but as saturation is approached, i.e., as ZnH approaches Cot the OH band intensity would increase dramatically relative to the ZnH band intensity. This is consistent with the infrared observations. Qualitative aspects of this result still follow even if reactions 3 and/or 4 take place only to a limited extent. This follows because as saturation is approached, the concentration of ZnO goes to zero and drives reaction 4 to the right. Infrared studies of the adsorption of ethylene at room temperature on zinc oxide revealed no bands that could be ascribed to opening of the double bond with formation of a saturated species. Accordingly, it appears that chemisorption occurs by interaction of a n bond with the surface. For steric reasons we would not expect the ethylene to approach the imbedded zinc ion of the active site close enough for such interaction; hence, we believe ethylene adsorption is confined primarily to oxide sites. Since, however, ethylene covers about 40% of :he surface (assuming an effective cross section of 23 A2)27 we can expect some chemisorption on oxide sites adjacent t o the imbedded zinc ion. These expected modes of adsorption are
HzC=CH2
I
+ -0-
H&=CH2
I
_r -0-
(54
represents the active site
the oxide ions which form a wall separating and -0active sites. (This is an oversimplification in the The Journal of Physical Chemistry
If Kz is sufficiently small, this expression predicts that
(4)
/
I n the above, -Zn-0-
sense that the zinc of the active site is likely to have several equivalent oxide nearest neighbors; nevertheless, rewriting the sequence in terms of an active site such as ZnOa would only greatly expand the number of equations without introducing significantly new features. For the sake of simplicity, therefore, we shall couch our discussion in terms of the simpler picture.) The first equation represents the initial adsorption act; the next equation provides for the site-to-site migration required for hydrogen-deuterium exchange which occurs readily on this catalyst.12 Steps 3 and 4 provide a mechanism whereby hydrogen attached t o zinc can exchange with hydrogens attached to oxide ions; studies by Eischens, et. uZ.,‘~ of the exchange process seem to require such a pathway. I n terms of these equilibria, we can compute the dependence of the total OH concentration (OH) on the concentration of ZnH (ZnH). As might be expected, this depends on the number of active sites (Co) and the number of oxide sites (02-). If the latter are only sparsely covered and all equilibria lie to the right except for (2), this leads to the relation
(26) For an analogous situation on aluminas, see J. B. Peri, J. Phys. Chem., 69,220 (1965). (27) 8. J. Gregg and K. S. W. Sing, “Adsorption, Surface Area and Porosity,” Academic Press, New York, N. Y., 1967, p 80.
3789
HYDROGENATION OF ETHYLENE BY ZINC OXIDE H&=CH2 I
H2C:=CH2
+ -Zn-0-
1
I
_C -Zn-0-
1
(5b)
H2C=CH2 H I -Zn-0-
I
+ H2C=CH2
/
(54
-Zn-0-
Since adsorption of ethylene occurs on oxide sites, ethylene adsorption should inhibit adsorption of hydrogen on these sites. Moreover, since the coverage with ethylene is relatively high, at least one of the oxide sites next to a ZnH should be occupied with a resulting perturbation of the ZnH band. Thus, these equations are consistent with the observation that admission of ethylene to a sample with chemisorbed hydrogen decreases the intensity of the OH band and perturbs the ZnH band without much effect on its intensity. Rates of adsorption of ethylene and hydrogen and the establishment of stable infrared spectra are rapid compared with the hydrogenation reaction;16 hence, it seems reasonable to assume that equilibrium is established in reactions 1 through 5. We shall further assume as a first approximation that adsorption of ethylene on oxide sites is indiscriminate and that all reactions listed as eq 5 have the same equilibrium constant. Reaction t o form ethane can occur stepwise as
+-Zn-0-
-zn-OI
CHz I -Zn-O--
R
= ICs
d F aK4
or R
(ZnO-CzH4)(H2(g))'/* (11)
= Ic'0'(H2(g))'/a
where 0' represents the adsorption of ethylene on an oxide site adjacent to the exposed zinc. (Even if the three reactions listed for ethylene chemisorption have different equilibrium constants, the form of eq 11 persists with a different combination of constants for k'.) Although 0' should depend on the pressure of hydrogen as well as ethylene, the evidence from the infrared that ethylene displaces hydrogen from oxide sites suggests this dependence may be slight. Accordingly, it is reasonable as a rough approximation to assume that the adsorption of ethylene on all sites in the presence of hydrogen shows the same pressure dependence as is found in the absence. On such a basis, eq 11 is in complete accord with the experimentally determined kinetics which yield a rate proportional to the square root of the hydrogen pressure and a dependence on ethylene pressure similar to that found for ethylene adsorption. (It should be noted that eq 11 is based solely on the law of mass action and does not require assumptions regarding the form of the isotherm. Strictly speaking, k6 should be written as k ~ / *, r where y is the activity coefficient of the activated complex, and 0' and (Hz(g)) should be regarded as activities; the latter will be well approximated by the pressure.) The assumption of equilibrium prior to the ratecontrolling step means that the kinetic isotope effect depends only on the partition functions of the activated complex and that for gaseous reactants. The activated complex is essentially an adsorbed alkyl radical. If we assume the reaction coordinate is the stretching of the newly formed C-H bond, the bending frequencies will be most important in determining the isotope effect. I n CHCla and CDC& the degenerate bending frequencies are 1210 and 908 cm-', respectively. From these and the zero point energies for hydrogen and deuterium, we can estimate the isotope effect by procedures outlined by Bigeleisen and Wolfsberg.2s Such a computation yields a rate with deuterium that is 69% that with hydrogen. The assumption of a slight loosening of the bending frequencies will decrease this value to that observed experimentally ( 4 0 % ) . Thus, the isotope effect is roughly consistent with the proposed mechanism. By way of contrast, one would expect significantly larger isotope effects or none were reaction 3 or sa, respectively, the rate controlling step.
*
I
CH,
Substitution of eq 10into 9 yields
H
I
+ --O-+ I
CzH6
+ --Zn--O-
1
+ -0-I
(7)
Since addition o§ deuterium to ethylene results solely in CzH4Dz,"the first step is irreversible. Bonding of an alkyl radical to a metal atom seems more likely than to an oxide ion; hence, the reactive sites have the form shown. Presumably, the steric factors that prevent 7r bonding with the zinc would be inoperative in the more directed u bond of the alkyl radical. I n terms of these equations, steady-state production of ethane, R, is given by
R = kG(H-znO-Ci") Use of eq 5 with the equilibrium assumptions yields
R = k6K6(H ZnO) (C2Hqg)) =
Equations 1through 4 can be combined to yield
(8)
(28) J. W. Bigeleisen and M. Wolfsberg, Advan. Chem. Phys., I, 15 (1958). Volume 78, Number 11 November 1969
E. JAMES NOWAK
3790 We feel the type I1 hydrogen adsorption may involve partial penetration into the bulk along the octahedral channels referred to earlier. The trigonal squeeze point with a diameter of 1.20 A does not seem prohibitively small compared to the size of a hydrogen molecule. Thus, we tentatively suggest that type I1 hydrogen may occuey octahedral holes, with a nominal diameter of 2.0 A, near the surface. Since the slow step for such sorption is likely to be expansion of the trigonal sites due to lattice vibrations, such a mechanism is consistent with the observation that the rate of type I1 sorption is the same with hydrogen and deuterium.'' Although the above picture is plausible, it does not provide an explanation for the promotional effects of hydrogen. In conclusion, let us note that we have tried to fit the data to other models. None we have considered, however, yield such a clean-cut, half-prder dependence on hydrogen and are also consistent with the auxiliary observations. It should be emphasized, however, that the essentials of the model we propose are that the active sites are isolated zinc oxide pairs with a pathway for limited hydrogen migration between sites. The
proposed picture of the surface structure that can give rise to these sites must be regarded as speculative even though many of the consequences seem to fit auxiliary observations. The essential features, however, bear many similarities to homogeneous catalysts which, because they are effective in dilute solution, can also be regarded as isolated, noninteracting sites. We can hope to find many analogies, therefore, in the behavior of oxides and homogeneous hydrogenation catalysts. It must, however, be emphasized that the model we propose and the relative simplicity of the kinetics would not be expected to obtain at higher temperatures where type I1 hydrogen plays a role. Thus, if analogies with homogeneous catalysts are to be sought, care should be taken that the temperature of the study is in a region such that the above model is applicable.
Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for support of this research. The writers also acknowledge the benefit of many helpful discussions of this work with Professor F. S. Stone.
Catalysis of the Reduction of Supported Nickel Oxide by E. James Nowakl Department o j Chemical Engineering, The University of New Mexico,Albuquerque, New Mexico (Received March 34, 1969)
87108
Small amounts of metallic platinum or palladium, incorporated in alumina-supportednickel oxide catalysts, have been observed to enhance activation of the catalysts by hydrogen. By contrast, small amounts of gold and silver, which do not readily adsorb hydrogen as atoms, also do not enhance activation of the catalysts. Hydrogenolysis, isomerization,and aromatization reactions of n-heptane with hydrogen, used for characterizing catalytic activity, were carried out in a pulse-flow reactor coupled with a gas chromatograph. Platinum or pallawere sufficient to cause large increases in conversion dium, in atom ratios to nickel (as oxide) as low as 5 X over the nickel component, while gold and silver, in atom ratios to nickel as high as 1 x 10-lproduce no such effect. There is now a preponderanceof evidence favoringreduction by surface-diffusing hydrogen atoms as the mechanism for activation of supported nickel oxide by platinum and palladium.
Introduction Commercial nickel catalysts are usually supported by refractory metal oxides such as alumina or silicaalumina, and they are activated by heating in a hydrogen-rich atmosphere which presumably reduces some or all of the nickel to its metallic form. Activation requires temperatures considerably higher than the temperature for reduction of pure bulk nickel oxide, apparently because the reducibility of nickel oxide is decreased by its strong interaction with the support The Journal of Physical Chemistry
materiaL2 Such high activation temperatures are undesirable, because they may cause sintering of the metallic nickel. It is apparent, then, that catalyzed reduction of supported nickel oxide could result in increased activity for the final catalyst by maximizing the extent of reduction a t a temperature low enough to avoid appreciable loss of area due to sintering. Ac(1) Sandia Corporation, P. 0. Box 969, Livermore, Calif. 94560. (2) F. N. Hill and P. W. Selwood, J . Amer. Chem. Soc., 71, 2522 (1949).
