Oxygen adsorption on silver catalysts during the course of partial oxidation of ethylene

Applied Catalysis A: General 304 (2006) 93–108 www.elsevier.com/locate/apcata

Oxygen adsorption on silver catalysts during the course of partial oxidation of ethylene Akimi Ayame *, Shigeru Eimaeda 1, Lin Feng 2, Hirofumi Hayasaka 3 Department of Applied Chemistry, Muroran Institute of Technology, 27-1 Mizumoto-cho, Muroran, Hokkaido 050-8585, Japan Received 8 September 2005; received in revised form 4 February 2006; accepted 11 February 2006 Available online 29 March 2006

Abstract To explore oxygen adsorption kinetics on silver catalyst surfaces during the course of ethylene oxidation, relations between the amount of oxygen consumed by reaction with ethylene (q) and oxygen partial pressure ( pO2 ) at a given ethylene pressure ( pC2 H4 ), i.e., kinetic data, were measured using a pulse reaction technique on K2SO4-promoted Ag/a-Al2O3 catalyst and a flow reaction method on Cs and Re-copromoted Ag/(aAl2O3-crystal) catalyst at 453–529 K. The kinetic data were analyzed by the use of three kinetic models derived on the basis of a redox model assuming that, under steady states, no oxygen desorption occurred and the rate of oxygen adsorption on silver surface was equal to the rate of surface reactions consisting of adsorbed oxygens and gaseous ethylene. One was the model that two-step consecutively dissociative oxygen adsorptions forming oxygen adatom and then admolecule were rate-determining (Model 1), and others were molecular oxygen adsorption ratedetermining model (Model 2) and surface reaction rate-determining model when one-step dissociative oxygen adsorption forming oxygen adatom only lay in a complete equilibrium state (Model 3). The Model 1 showed the most excellent suitability to the kinetic data and reproduced accurately the distinctive features for experimental q
pO2 curves. The coverage of oxygen admolecule estimated from the Model 1 was 0.48–0.96 and 0.12– 0.34 on the K2SO4-promoted Ag/a-Al2O3 catalyst at pC2 H4 2 0:03 atom and on the Cs and Re-copromoted Ag/(a-Al2O3-crystal) catalyst at pC2 H4 3 0:30 atom, respectively. They were about thousand times bigger than those for oxygen adatom. These results indicated that the first dissociative adsorption on vacant active center was much slower than the second dissociative oxygen adsorption on adatom; the adatom seemed likely to act as active center for the second oxygen adsorption. Selectivity to ethylene oxide elevated almost linearly with increase in the coverage of oxygen admolecule. Activation energies and heats of adsorption obtained for dissociative oxygen adsorption were quite reasonable values compared to literature values and remarkedly reduced with increasing pC2 H4 . # 2006 Elsevier B.V. All rights reserved. Keywords: Silver catalyst; Ethylene; Partial oxidation; Dissociative oxygen adsorption; Superoxide; Redox model; Adsorption kinetics

1. Introduction Many investigations concerning mechanism and kinetics of ethylene partial oxidation on silver catalyst and of oxygen adsorption on silver have been carried out [1–5]. Kinetic studies can be fundamentally classified into Langmuir–Hinshelwood mechanism [6–12] and Rideal–Eley mechanism [13–17]. It is,

* Corresponding author. Tel.: +81 143 86 9060; fax: +81 143 86 9060. E-mail address: [email protected] (A. Ayame). 1 Present address: Oji Paper Co. Ltd., 5-12-8 Ginza, Chuo-ku, Tokyo 1040061, Japan. 2 Present address: Kyatarah Industrial Co. Ltd., 7800 Chihama, Daito-Cho, Ogasa-Gun, Shizuoka 437-1412, Japan. 3 Present address: Kyowa Chem. Eng., Ltd., 6-1-5-1 Kitano, Kiyota-ku, Sapporo, Hokkaido 004-0866, Japan. 0926-860X/$ – see front matter # 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.apcata.2006.02.031

however, very difficult to find the distinct and consistent reaction orders for partial pressures of ethylene and oxygen, because those have been measured employing a large variety of silver catalysts and reaction conditions. As origins of such uncertainty, the followings are considered: (i) the active state of silver catalyst surface is strongly susceptible to reaction conditions, especially feed gas composition, (ii) on any oxygen coverage, the epoxidation and the complete oxidation of ethylene take place competitively and/or simultaneously, and also the consecutive oxidation of ethylene oxide proceeds at the same time, and (iii) it is difficult to represent the epoxidation rate and the combustion rate as simple function of the partial pressure (or coverage) of oxygen, because the adsorption behaviors of oxygen on the silver surface, especially just during the course of the partial oxidation, are unclear yet. Ayame et al. [16] have previously derived a set of available empirical rate

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equations, based on Rideal–Eley mechanism, for an integral fixed bed reactor packed with the K2SO4-promoted Ag catalyst supported on a-Al2O3, which was well-stabilized using the reactant gas mixture of C2H4/O2/N2 = 2.5/20/77.5 by volume [15,16]. The rate equations reproduced the kinetic data with good precision in the range of 1–10 atm under the fixed feed composition. The formation rate of ethylene oxide depends on the 1/2 power of the oxygen partial pressure, which was also reported by Todes and Andrianova [14] and Yasumori et al. [18]. The reaction order with respect to oxygen suggests that ethylene oxide is produced by the reaction of ethylene with an oxygen atom but does not indicate whether it is an oxygen adatom or a terminal atom of superoxide ion. On silver catalyst surfaces, it is well known that there are at least three kinds of adsorbed oxygen species; molecular, atomic, and subsurface oxygens [1,3,4]. Smeltzer et al. found first the presence of the three kinds of adsorbed species on silver catalysts [19], which were later ascertained by Czanderna [20], Kilty et al. [17], and Kondarides et al. [21]; dissociatively adsorbed oxygen with very fast adsorption rate (non-active adatom), molecularly adsorbed oxygen (active admolecule), and dissociatively adsorbed oxygen (active adatom) including surface diffusion and migration into the bulk phase, the adsorption processes of which had the adsorption activation energies of 4–12, 33–42, and 58–100 kJ mol
1, respectively [17,19–21]. Backx et al. first observed the presence of adsorbed oxygen molecule on Ag(1 1 0) plane by HREELS [22]. The presence of molecular oxygen and/or superoxide ion (O2
) has been demonstrated by the direct observations using ED [23,24], ESR [25–27], UPS [28,29,32,33], XPS [28,30–32], HREELS [33,34], and SERS [35]. Since these experiments were carried out under UHV, the oxygen admolecule was completely desorbed above 170 K. Under experimental conditions near to the practical epoxidation, Gerei et al. [36] and Kilty et al. [17,37] observed the IR band 870 cm
1 assigned to n(O–O) of  CH2–CH2–O–O–Ag. Grant and Lambert assigned the TPD peak at 380 K to oxygen admolecule perpendicular to silver surface in the experiments using AES, XPS, and oxygen isotope exchange reactions [38,39]. The presence of the intermediate (CH2–CH2–O–O–Ag) in ethylene oxide production process has been confirmed through many valuable investigations with ESR [40,41], SERS [42], and high-sensitive SERS [43,44]. Nakatsuji et al. studied the mechanism of the silver-catalyzed partial oxidation by theoretical calculations using the ab initio second-order MFller-Plesset (MP2) method combined with the dipped adcluster model (DAM) [45–47]. They concluded that the active species was superoxide ion, which was molecularly adsorbed in a bent end-on geometry on the silver surface, ethylene reacted with one terminal atom of the superoxide, and the subsequent reaction proceeded smoothly without a large energy barrier to yield ethylene oxide, while the atomically adsorbed oxygen was not necessarily selective for the epoxidation of ethylene. On the other hand, many researches for oxygen adatom using HREELS [22,48], LEED [48–50], AES [49], XPS [51– 53], and STM [54] were also reported. But, directly

experimental evidences, such as IR and SERS, for the interactions between ethylene and oxygen adatom seem likely to be poor. Backx et al. [22] and Campbell [31,52] demonstrated the existence of subsurface oxygen using HREELS and XPS, respectively. Ertl et al. reconfirmed the presence of subsurface oxygen just below the first silver surface layer and oxygen species migrated into bulk silver using UPS, XPS, ISS and Raman spectroscopy [55,56]. The subsurface oxygen promotes the formation of ethylene oxide [38,57,58]. Furthermore, when the alkali metals are added into silver catalysts as promoter, the presence of other kinds of surface oxygen species are easily stipulated. One of the authors observed in the dynamic measurements by in situ XPS method that the five kinds of surface oxygen species (oxygen admolecule, superoxide ion, subsurface oxygen, oxygen adatom, and oxydic oxygen atom) existed simultaneously on the self-supporting disks of pure silver powder and of the Cs and Re-copromoted silver powder during the exposure to O2-gas jet of 5  10
6 Pa in UHV at 483 K [59–61]. Consequently, the oxygen species and oxygen adsorption behaviors on silver catalysts are very complex and a number of unclear problems for their details are left. The subjects that one desires earnestly to know are related to the oxygen adsorption behaviors and the relative amounts of oxygen species on the silver surfaces during the course of the partial oxidation of ethylene. However, the direct measurement of oxygen adsorption behaviors during the reaction is very, very difficult. Yokoyama and Miyahara reported that no isotope exchange between N216O and 18O2 occurred during the ethylene oxidation [62]. Yong and Cant also reported that: (i) the reaction of ethylene with 18O2 was at least five times faster than with N216O, (ii) when ethylene was absent the oxygen isotope exchange between 18O2 and N216O was detectable but the rate was much slower than that of ethylene oxidation under similar conditions, and (iii) no oxygen isotope exchange occurred during ethylene oxidation with the mixture of 18O2 and N216O [63]. These results are anticipated to suggest that, during the oxidation, the recombination and desorption of adsorbed oxygens do not take place and the rate of dissociative oxygen adsorption is equal to or slower than the overall rate of surface reactions between adsorbed oxygens and ethylene. In the present work, based on the results reported by Yokoyama and Miyahara [62] and Yong and Cant [63], an attempt was challengingly, though it was classically, made to explore oxygen adsorption behaviors and coverages of adsorbed oxygen species on K2SO4-promoted Ag/a-Al2O3 catalyst and highly selective Cs and Re-copromoted silver/(aalumina crystal) catalyst just during the ethylene oxidation took place. The K2SO4-promoted Ag/a-Al2O3 catalyst resulted in good catalytic performance under conditions of the partial pressure ratios of ethylene to oxygen of 3–5/20 at 453–523 K [15,16]. However, even the steady activity state obtained after the reaction for 30–50 h sometimes varied for a long time when the partial pressure ratio was altered. Consequently, kinetic data for the K2SO4-promoted Ag/a-Al2O3 catalyst were measured using a conventional pulse reaction technique, since the contact

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time between reaction gas mixture and catalyst surface was very short and then the occurrence of catalytic activity changes was very hard. On the other hand, the Cs and Re-copromoted silver/(a-alumina crystal) catalyst showed very stable activity and selectivity under conditions of the partial pressure ratio of ethylene to oxygen higher than 30/8 at 453–493 K [67,68] and the process time required to reach steady activity states under different conditions was only 1 or 2 h. So, measurement of kinetic data for this catalyst was carried out using a fixed bed flow reaction. If the results obtained on the two types of silver catalysts, which indicated good catalytic performances under reaction conditions of largely different partial pressure ratios, agree with each other, the conclusions derived from the present work are expected to reflect more precisely the feature of truth. 2. Rate equations and oxygen adsorption models 2.1. Oxygen adsorption rate Catalytic partial oxidation of ethylene on silver catalysts is one of the simplest complex reaction, which comprises of epoxidation of ethylene, consecutive oxidation of ethylene oxide, and direct combustion of ethylene. The net total reaction can be expressed by Eq. (1). XC2 H4 þ qO2 ! ! ! xC2 H4 O þ 2yðCO2 þ H2 OÞ

(1)

The q is given by Eq. (2) using total conversion of ethylene (X = x + y) and selectivity to ethylene oxide (S = x/X). q ¼ 12ð6
5SÞX

(2)

As the oxidation proceeds through reactions between gaseous ethylene and oxygen adsorbed on silver surface (Rideal–Eley mechanism), the amount of ethylene converted ( pC2 H4 X) and that of oxygen consumed ( pC2 H4 q) can be correlated to the total rate of surface reactions consisting of adsorbed oxygen and ethylene in gas-phase (i.e., reduction rate of the silver surface; rred) and the rate of oxygen adsorption on the surface (rads) under steady state, respectively, as follows: rred ¼ c1 pC2 H4 X

(3)

rads ¼ c2 pC2 H4 q þ rdes

(4)

where c1 and c2 are rate constants and rdes is the rate of oxygen desorption. 2.2. Redox model As no oxygen desorption of Eq. (4) occurs during ethylene oxidation according to the experimental results by Yokoyama and Miyahara [62] and Yong and Cant [63], the rdes is able to be treated as zero. Consequently, under steady states, such a redox model as shown in Fig. 1 and Eq. (5) is established. rred ¼ rads (5)

Fig. 1. Relationships among the oxidation rate by oxygen adsorption (rads), the reduction rate by ethylene (rred), and the oxygen surface coverage (uS) on silver catalyst surface under a steady state (concept on the redox model).

The surface oxygen coverage, us, corresponding to a pair of pO2 and pC2 H4 and reaction temperature is simultaneously fixed. Here, the rads, which is defined by Eq. (4) when rdes = 0, is also represented by a function of pO2 and us. Although the rred can be similarly given by terms of pC2 H4 and us, the rred is treated using Eq. (3), according to the aim of the present work. Hereafter, kinetic data measured are analyzed on the basis of the assumption that the oxygen adsorption step is a ratedetermining step and a pseudo(partial)-equilibrium state is kept between the oxygen partial pressure in gas-phase and the amount of oxygen adsorbed on the silver surface [64]. 2.3. Two-step dissociative oxygen adsorptions: Model 1 It has been well known that oxygen admolecule and oxygen adatom are simultaneously present on silver surface together with subsurface oxygen atom. If the surface reactions are very fast and the dissociative oxygen adsorption proceeds ratedeterminingly, the following two-step dissociative adsorption model of oxygen molecule can be considered. K1 and K2 are oxygen adsorption equilibrium constants at the pseudo(partial)adsorption equilibrium state based on the steady state hypothesis [64]. K1

O2 þ 2u
!2uO K2

(6)

O2 þ 2uO
!2uOO

(7)

O2 þ u þ uO ! uO þ uOO

(8)

The u, uO, and uOO are defined as the coverage of vacant active site, oxygen adatom, and oxygen admolecule, respectively.

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Since Eq. (8) can be also derived from Eqs. (6) and (7), Eqs. (6) and (7) are main adsorption processes of oxygen. Since ethylene oxide is synthesized from ethylene in gas-phase and oxygen admolecule; C2 H4 þ uOO ! C2 H4 O þ uO and the oxidation proceeds with keeping the pseudo-equilibrium states of oxygen adsorptions, Eqs. (9)–(11) are accessible from the mass balance of uO, i.e., uO ¼ ðK1 pO2 Þ1=2 u
ðK2 pO2 Þ1=2 uO uO ¼

ðK1 pO2 Þ1=2 1 þ ðK2 pO2 Þ1=2

uOO ¼ ðK2 pO2 Þ b¼

1=2

(9)

u ¼ bu

uO ¼ ðK2 pO2 Þ

(10) 1=2

bu

ðK1 pO2 Þ1=2 1=2

1 þ ðK2 pO2 Þ

(11) (12)

where the parameters with upper circle mark are their values corresponding to pO O2 . Consequently, the rrel value can be determined by measuring q and qO. The rads can be also given as the sum of the forward reaction rates of Eqs. (6) and (7): rads ¼ a11 pO2 u2 þ a12 pO2 ðuO Þ2 ¼

a11 pO2 ð1 þ a12 b2 =a11 Þs 2 ð1 þ ðK1 pO2 Þ1=2 Þ2

;

(17) where a11 and a12 are adsorption rate constants and the term of a11(1 + a12b2/a11)s2 is almost constant. So, Eq. (17) is converted to Eqs. (18) and (19). rrel ¼

k1 ¼

k1 pO2

(18)

ð1 þ ðK1 pO2 Þ1=2 Þ2

1=2 2 ð1 þ ðK1 pO Þ O2 Þ O pO2

ðconstantÞ

(19)

The establishment of Eq. (12) in the linear form implies that b is nearly constant, as
1=2 K2 gradient ¼ 1 ¼ b (13) K1

The relative rate constant, k1, and the equilibrium constant, K1, are determined by analysis of the linear form of Eq. (18) using the least square method. The estimated value of k1 is also confirmed by Eq. (19).

Furthermore, if the coverages of CO2, H2O, C2H4O, and irreversibly adsorbed organic residues [13,65] are expressed by uC, uW, uEO, and uresi, respectively, the mass balance of u is expressed by

2.4. Molecular oxygen adsorption: Model 2

u ¼ 1
uO
uOO
uC
uW
uEO
uresi ¼ ð1
uC
uW
uEO
uresi Þ
uO
uOO

(14)

The term of (1
uC
uW
uEO
uresi) is here postulated to be almost constant, because the catalyst surface was stabilized for a long time using the flow reaction and also the steady activity under a standard feed gas (described later) did not changed even after the measurements of kinetic data using feeds of different gas compositions. Consequently, Eq. (14) is reduced to Eq. (15) if the total number of vacant active sites is given by s (constant). s u ¼ s
uO
uOO ; u ¼ (15) 1 þ ðK1 pO2 Þ1=2 This indicates that u is formally fixed by the first dissociative oxygen adsorption step. If K1 is determined by an adequate method, the u, b and K2, uO, and uOO are able to be estimated using Eqs. (15), (12), (10) and (11), respectively. In this work, when a maximum oxygen partial pressure ( pO O2 ) is adopted as a reference pressure in each experimental series at a pair of constant pC2 H4 and reaction temperature, a relative reaction rate (rrel) is defined using Eqs. (3)–(5) at rdes = 0, as follows: rrel ¼

rred rads c2 pC2 H4 q q ¼ O ¼ ¼ O O O q rred rads c2 pC2 H4 q

(16)

Assuming that the forward adsorption process of Eq. (20) is a rate-determining step, part of the molecular oxygen adsorption lies in a pseudo-equilibrium state, and all of surface reactions is also initiated by the interactions between ethylene and uOO, the rads and rrel are represented by Eqs. (21) and (22), respectively. KO 2

O2 þ u
!uOO rads ¼ a2 pO2 u ¼ rrel ¼ k2 ¼

(20) a2 pO 2 1 þ KO2 pO2

k2 pO2 1 þ KO2 pO2

1 þ KO2 pO O2 pO O2

(21) (22)

ðconstantÞ;

(23)

where KO2 is pseudo(partial)-equilibrium constant, a2 adsorption rate constant, and k2 is the relative rate constant. 2.5. One-step dissociative oxygen adsorption equilibrium: Model 3 When the dissociative oxygen adsorption of Eq. (24) lies in an complete equilibrium state with equilibrium constant K1e and the surface reaction consisting of ethylene and oxygen adatoms is rate-determining, K1e

O2 þ 2u Ð 2uO

(24)

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

the following rate equations are derived. rred ¼ a3 pC2 H4 uO rrel ¼

k3 ð pO2 Þ

(25)

1=2

1 þ ðK1e pO2 Þ

1=2

and k3 ¼

1 þ ðK1e pO O2 Þ 1=2 ð pO O2 Þ

1=2

(26)

where a3 is rate constant and k3 is the relative rate constant. 3. Experiment 3.1. Catalysts Two kind of silver-supported catalysts were used. One of the catalyst had a bulk composition of Ag/K2SO4/a-Al2O3 = 1/ 0.01/3.78 by weight. Detail of the preparation was described in the earlier reports [65,66]. Another catalyst was Cs and Recopromoted Ag catalyst supported on a-alumina crystal carrier (a-Al2O3(cr)). The contents of Cs and Re were 1000 ppm in each element referred to Ag, and the silver content supported was 12 wt.%. The details were also described in the previous reports [67,68]. 3.2. Reaction systems 3.2.1. Pulse reaction The K2SO4–Ag/a-Al2O3 catalyst showed activity decrease and increase in selectivity to ethylene oxide for 30–50 h [15,16]. If feed gas was switched to new one having different pC2 H4 = pO2 ratio, the activity and selectivity changed over 20 h. So, using the feed gas mixture of pC2 H4 = pO2 = pN2 ¼ 0:032=0:202=0:766, the catalyst was at first stabilized for 50 h at 523 K and at 1750 g-cat h (g-mol-C2H4)
1. After subsequent oxidation by oxygen and then reduction by hydrogen at 553 K each for 2 h, the catalyst indicated repeatedly a constant level of steady activity state within 4 h after restarting the

97

reaction [65]. Nevertheless, the steady activity state sometimes varied after the reactions with feeds of different gas compositions. In the present work, a pulse reaction technique was employed to obtain many sets of kinetic data on the K2SO4– Ag/a-Al2O3 catalyst having a definitely steady activity state. In addition, it was very important to see whether the kinetic data obtained were really measured or not on the catalyst surface with an identical activity state. The confirmation was made by comparing X and S values determined with a standard feed gas of pC2 H4 = pO2 = pN2 ¼ 0:0099=0:0661=0:924, before and after the pulse reactions with various feed gas mixtures. Through all the experiments, no variation in the catalytic activity states was observed. Fig. 2 shows the pulse reaction system comprised of a fixed bed reactor, a pulse injection glass loop (3.4 ml), six glass ball flasks for reactant gas mixture storage, a vacuum system, a feed gas mixing and controlling system, and a TCD gas chromatography equipped with DOP column (6 mm o.d.  1200 mm, 353 K) and Porapack Q column (4 mm o.d.  3800 mm, 353 K). The reactor consisted of a Pyrexglass U-tube of 10 mm i.d., at the center of which a glass thermocouple shield of 4 mm o.d. was attached. The catalyst weight used was 3.0 g and reaction temperature was controlled within 1 K. Flow rate of He carrier was 27.4 ml min
1, pulse size 1.12 ml (STP), and pulse interval 14 min. Kinetic data were measured by changing pO2 of reactant mixtures, pC2 H4 of which was kept constant. The gas compositions and reaction temperatures adopted are listed in Table 1. 3.2.2. Flow reaction The CsRe–Ag/a-Al2O3(cr) catalyst indicated very stable activity and selectivity for a long time at 433–493 K under the standard feed gas of pC2 H4 = pO2 = pN2 ¼ 0:30=0:08=0:62 [68]. Even if the feed gas mixture was altered, new steady activity

Fig. 2. Schematic drawing of the experimental apparatus.

0.66–15.0

506 489

E-4-1 E-4-2

S (%)

q

Catalyst was K2SO4–Ag/a-Al2O3 (3.0 g), flow rate of He carrier 27.4 ml min
1 (STP), pulse size 1.12 ml (STP), and pulse interval 14 min.

X (%)

S (%)

q

E-3-3

X (%)

states were obtained within 2 h. Therefore, kinetic data on the CsRe–Ag/a-Al2O3(cr) catalyst (9.0 g) were measured by the fixed bed flow reaction at 453–493 K and total flow rate 26.4 ml min
1 (GHSV 250 h
1). The apparatus shown in Fig. 2 was also used. Pyrex-glass tubing of 12 mm i.d. with a glass thermocouple shield of 4 mm o.d. at the center was used as reactor. Feed gas mixtures comprised of pC2 H4 ¼ 0:30 or 0.50 atm (constant), pO2 = 0.014–0.088 atm (variable), and pHe = balanced pressure. Two sets of kinetic data, when monochloroethylene of 2.0 ppmv was added, were also measured at 473 and 493 K. The addition of monochloroethylene to the standard feed gas caused activity decrease and selectivity elevation for 10–20 h. The steady activity and selectivity reached were almost unchanged even after measurements of the pO2 dependences of X and S under the conditions of constant pC2 H4 and monochloroethylene concentration. 4. Results

pO2 ðatmÞ X (%)

68.7 74.4 80.6 83.2 84.7 85.2 87.0

pO2 ðatmÞ

0.0100 0.0201 0.0441 0.0657 0.0987 0.131 0.150

S (%)

q

E-2-2 E-1-2

Table 2 shows typical kinetic data for the K2SO4–Ag/aAl2O3 catalyst obtained using the pulse reaction. All of the data was given by an average value of 3rd–5th pulse reactions. Fig. 3 indicates, as an example, dependencies of X, S, and q on pO2 for run-no. E-3-2. The X, S, and q varied smoothly with increasing pO2 . Suitabilities of the linear forms of Eqs. (18), (22) and (26) to the kinetic data E-3-2 were examined (Fig. 4). Eq. (18) resulted in a good straight line, while the plots for Eqs. (22) and (26) indicated an upward (convex) and downward (concave) curve, respectively. The constants of k1 and K1 were determined from the ð pO2 =rrel Þ1=2 versus ð pO2 Þ1=2 plot by the least square method. The k2 and KO2 or the k3 and K1e were also estimated from the pO2 =rrel versus pO2 plot or the ð pO2 Þ1=2 =rrel versus ð pO2 Þ1=2 plot approximated to be linear line, respectively. These constants for all the sets of kinetic data presented in

Table 2 Typical experimental data obtained in the pulse reactions

4.1. Pulse reaction data

47.0 49.8 57.6 58.4 62.2 63.3 63.5

3.29

12.9 17.2 26.1 31.5 41.0 44.2 45.0

E-3-1 E-3-2 E-3-3 E-3-4

0.00661 0.0105 0.0209 0.0396 0.0989 0.131 0.150

529 513 491 471

0.325 0.389 0.476 0.540 0.559 0.566

2.2–32.9

48.5 55.4 57.7 60.0 61.8 62.0

1.70

18.2 24.1 30.5 36.0 38.4 39.0

E-2-1 E-2-2 E-2-3

0.0219 0.0518 0.0992 0.204 0.301 0.329

527 506 488

0.514 0.594 0.738 1.032 1.06 1.085

0.68–15.0

54.3 57.4 59.1 62.6 62.7 62.9

0.981

31.3 38.0 48.5 71.9 74.1 76.0

E-1-1 E-1-2 E-1-3

0.00677 0.00983 0.0197 0.100 0.132 0.150

528 506 486

0.897 0.984 1.048 1.080 1.108 1.123 1.129

1.0–15.0

67.8 67.1 68.0 68.1 67.7 67.3 68.1

0.329

S (%)

Run-number

X (%)

T (K)

pO2 ðatmÞ

pO2  102 ðatmÞ

E-4-1

pC2 H4  102 ðatmÞ

q

Table 1 Reaction conditions used in the pulse reactions

0.236 0.302 0.407 0.485 0.592 0.627 0.636

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

pO2 ðatmÞ

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A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

99

Fig. 3. Dependences of (5) total ethylene conversion X, (~) selectivity to ethylene oxide S, and (*) the amount of oxygen consumption q on oxygen partial pressure for the run-no. E-3-2. Fig. 4. Suitabilities of the linear forms of (*) Eq. (18), (*) Eq. (22), and (5) Eq. (26) to the experimental data for the run-no. E-3-2.

P Fig. 3 and Table 2 are summarized in Table 3. The ðdi Þ2 is the sum of square of the residue (SSR), where di = (qi)ex
(qi)cal; (qi)ex is the experimental value and (qi)cal the value calculated using the relative rate constant (ki) and pseudo-equilibrium constant (Ki) for each oxygen adsorption model. The SSR values for Eq. (18) were smaller than those for Eq. (22) by a factor of 10 or 100. The SSR for Eq. (26) was reliably smaller than Eq. (22) but larger than those for Eq. (18), though the difference between the SSR values for Eqs. (26) and (18) were small. Table 4 indicates characteristic features of the plots of ð pO2 =rrel Þ1=2 versus ð pO2 Þ1=2 , pO2 =rrel versus pO2 , and ð pO2 Þ1=2 =rrel versus ð pO2 Þ1=2 , including the SSR values, for all the sets of kinetic data. Fig. 5 compares the reproducibilities of Eqs. (18), (22) and (26) with respect to the experimental q
pO2 curve for run-no. E-3-2. Eq. (18) reproduced the distinctive features of the q
pO2 curve with better agreement to the experimental curve than Eqs. (22) and (26). From these

results mentioned above, it became clear that the Model 1, based on the assumption that the two-step dissociative oxygen adsorptions were rate-determining, was superior to those for the molecular oxygen adsorption (Model 2) and the one-step dissociative oxygen adsorption equilibrium (Model 3). In order to estimate the coverages of adsorbed oxygen species, it was necessary to determine K2 and b in Eqs. (10)– (13) for the Model 1. These parameters calculated using K1 are shown in Table 5 with k1 and K1 values. The b was very small (10
3 order) and K2 very large value of 106–109 order. Here, the magnitudes of uO, uOO, and u were estimated, assuming as s = 1. Table 6 indicates these oxygen coverages for four typical instances. The uO was 10
3 to 10
4 order, being very small compared to the uOO and u. With increasing pC2 H4 the uOO decreased and the u increased. Furthermore, from the relations of K2  K1 and uOO > u  uO, it seems likely to be clarified that, on the silver catalysts during ethylene oxidation, the first

Table 3 The rate constants, the pseudo-equilibrium constants, and the SSR values for Model 1, Model 2, and Model 3 Run-number

Model 1 k1

Model 2 K1

P

2

ðdi Þ

k2

Model 3 KO2

P ðdi Þ2

k3

K1e

P

ðdi Þ2

E-1-2 3750 3440 4.04  10
5 245 238 4.03  10
3 29.5 726 5.60  10
5 E-2-2 452 349 5.01  10
5 98.4 91.8 6.86  10
3 9.19 43.7 5.95  10
4 E-3-2 317 258 4.42  10
4 49.3 46.2 3.26  10
3 8.18 41.4 7.07  10
4 E-3-3 225 176 5.17  10
4 40.7 37.7 2.67  10
3 6.80 25.6 5.27  10
4 E-4-1 271 193 3.96  10
4 71.2 64.5 1.64  10
3 6.93 18.9 1.54  10
3 P ðdi Þ2 is the sum of square of residue (SSR) of di = (qi)ex
(qi)cal, where (qi)ex is the experimental value and (qi)cal is the value calculated using Eqs. (18), (22) or (26).

100

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

Table 4 Tests of kinetic models Run-number

Temperature (K)

Linearity

Table 5 Rate constant (k1) and oxygen adsorption equilibrium constants (K1 and K2) for Model 1

SSR

M1

M2

M3

M1

M2

M3

E-1-1 E-1-2 E-1-3

528 506 486

glla gll gll

uc b uc uc

dc c dc dc

~ * ~

5 5 5

* ~ *

E-2-1 E-2-2 E-2-3

527 506 488

gll gll gll

uc uc uc

dc dc dc

~ * ~

5 5 5

* ~ *

E-3-1 E-3-2 E-3-3 E-3-4

529 513 491 471

gll gll gll gll

uc uc uc uc

dc dc dc dc

* * * *

~ 5 5 5

5 ~ ~ ~

E-4-1 E-4-2

506 489

gll gll

uc uc

dc dc

* *

5 5

~ ~

M1: Model 1 (two-step dissociative oxygen adsorptions are rate-determining; Eq. (18)). M2: Model 2 (molecular adsorption of oxygen is rate-determining; Eq. (22)). M3: Model 3 (dissociative adsorption equilibrium model of oxygen; Eq. (26)). SSR: order of the magnitude of the SSR values; * < ~ < 5. a Good linear line. b Upward (convex) curve. c Downward (concave) curve.

dissociative adsorption of oxygen (Eq. (6)) is slow step and the second dissociative adsorption step (Eq. (7)) shifts very large to the right hand side; that is, the overall rate of surface reactions of ethylene and adsorbed oxygen is very near to the rate of the second dissociative oxygen adsorption on uO, which functions as active center.

Fig. 5. Comparison of the q vs. pO2 curves for Eq. (18), (22), and (26) calculated using the constants given in Table 3 with the experiment curve for the run-no. E-3-2. (*) Experimental values. Calculated curves: (—) Eq. (18), (- - -) Eq. (22) and (– - –) Eq. (26).

Run-number

k1 (atm
1 min
1)

K1 (atm
1)

b  103

K2  10
7 (atm
1)

E-1-1 E-1-2 E-1-3

8640 3750 1610

7990 3440 1410

2.30 5.60 8.78

151 10.9 1.82

E-2-1 E-2-2 E-2-3

822 452 293

681 349 211

1.75 2.65 6.93

22.2 5.00 0.437

E-3-1 E-3-2 E-3-3 E-3-4

473 317 225 139

400 258 176 101

5.43 6.22 7.73 7.30

1.35 0.666 0.293 0.188

E-4-1 E-4-2

271 200

193 133

2.64 4.19

2.77 0.758

The k1 and K1 were calculated using Eqs. (18) and (19). The K2 and b were obtained using the K1 and Eq. (12) with the least square method. The K1, K2, and b values satisfied the relationship of Eq. (13).

4.2. Flow reaction data The kinetic data for the CsRe–Ag/a-Al2O3(cr) catalyst determined using the flow reaction with the feeds of pC2 H4 3 0:30 atm are shown in Table 7 and Fig. 6. The numerals before and after the letter A of the run-number represent pC2 H4  102 atm and reaction temperature in Kelvin, respectively. Two sets of flow reactions of 3A473 and 3A493 were also performed to make sure whether similar results to those derived from the pulse reaction data were available or not from the flow reaction kinetic data. The analysis of the 3A473 and 3A493 data fortunately brought out the same results in the examination of linearlity and in the comparison of SSR values and reproducibility of q
pO2 curve as those in the pulse reaction (Table 8). Fig. 7 indicates suitabilities of the linear forms of Eqs. (18), (22) and (26) to the 50A453 data. A good linear line was obtained only for Eq. (18). Eq. (22) showed a convex curve and Eq. (26) a concave curve. Especially, the ð pO2 Þ1=2 =rrel versus ð pO2 Þ1=2 plot of Eq. (26) indicated a negative slope when it was approximated by a linear line. As the slope, ðK1e Þ1=2 =k3 , should be essentially positive in physical meaning, the negative slope was never accepted; that is, Model 3 was at the present stage rejected. The characteristics in suitabilities of the three equations to experimental data, including reproducibilities for the q
pO2 curve, were quite unchangeable for all the sets of flow reaction kinetic data. Table 8 shows rate constants, pseudo-equilibrium constants of oxygen adsorption, and SSR values. The SSRs for Eqs. (22) and (26) were larger than that for Eq. (18). We concluded that Model 1 was much superior to Model 2 and Model 3, as can be seen from Fig. 7 and Table 8. The same results as those in the pulse reaction data on the K2SO4–Ag/a-Al2O3 catalyst were accessible from the flow reaction kinetic data on the CsRe–Ag/ a-Al2O3(cr) catalyst at pC2 H4 3 0:30 atm. These facts were quite unchanged even if monochloroethylene was added into the feed gas mixture (Tables 7 and 8).

Table 6 The typical values of uO, uOO, and u estimated using the constants for Model 1, which are shown in Table 5 E-1-2

E-2-2
2

506 K, pC2 H4 ¼ 0:329  10

E-3-2
2

506 K, pC2 H4 ¼ 0:981  10

atm

E-4-1
2

513 K, pC2 H4 ¼ 1:70  10

atm

506 K, pC2 H4 ¼ 3:29  10
2 atm

atm

u

uO  103

uOO

pO2 ðatmÞ

u

uO  103

uOO

pO2 ðatmÞ

u

uO  103

uOO

pO2 ðatmÞ

u

uO  10 3

uOO

0.0100 0.0201 0.0441 0.0657 0.0987 0.1313 0.1500

0.146 0.107 0.075 0.062 0.052 0.045 0.042

0.82 0.60 0.42 0.35 0.29 0.25 0.24

0.852 0.890 0.922 0.935 0.946 0.952 0.955

0.00677 0.00983 0.0197 0.100 0.132 0.150

0.394 0.351 0.276 0.145 0.128 0.121

1.04 0.93 0.73 0.38 0.34 0.32

0.606 0.649 0.724 0.855 0.872 0.879

0.0219 0.0518 0.0992 0.204 0.301 0.329

0.296 0.215 0.165 0.121 0.102 0.098

1.84 1.34 1.03 0.75 0.63 0.61

0.703 0.785 0.834 0.878 0.898 0.902

0.00661 0.0105 0.0209 0.0396 0.0989 0.1310 0.1500

0.470 0.413 0.332 0.266 0.186 0.166 0.157

1.24 1.09 0.88 0.70 0.49 0.44 0.41

0.530 0.587 0.668 0.734 0.814 0.834 0.843

Table 7 Typical kinetic data obtained in the flow reactions on the CsRe–Ag/a-Al2O3(cr) catalyst at GHSV 250 h
1 3A473

30A473 S (%)

q  102

ðatmÞ

X (%)

0.038 0.066 0.093 0.179 0.198

26.0 31.5 34.0 38.9 39.5

66.1 65.0 64.9 63.1 63.0

35.0 43.3 46.8 55.3 56.3

pO

2

50A473 S (%)

q  10 2

ðatmÞ

X (%)

0.016 0.023 0.034 0.052 0.057

2.70 3.66 5.00 6.92 7.36

61.1 63.2 64.8 65.1 65.0

3.98 5.20 6.90 9.50 10.12

pO

2

30AE473 S (%)

q  10 2

ðatmÞ

X (%)

0.014 0.017 0.024 0.038 0.059 0.065

1.47 1.79 2.40 3.50 4.83 5.17

63.0 64.2 64.9 66.2 67.3 67.8

2.10 2.50 3.31 4.71 6.36 6.75

pO

2

3A493 S (%)

q  10 2

ðatmÞ

X (%)

0.022 0.036 0.056 0.074 0.081

2.30 3.22 4.82 6.33 7.00

76.0 77.8 77.0 77.0 78.0

2.53 3.40 5.18 6.80 7.35

pO

2

30A493 S (%)

q  10 2

ðatmÞ

X (%)

0.038 0.068 0.095 0.120 0.132

46.9 54.8 58.5 60.2 60.5

61.2 61.2 61.2 61.5 61.4

69.0 80.5 86.0 88.0 88.6

pO

2

50A493 S (%)

q  102

ðatmÞ

X (%)

0.017 0.027 0.038 0.048 0.053

3.30 4.97 6.72 8.20 8.95

56.5 60.0 62.1 63.0 63.7

5.24 7.45 9.72 11.69 12.60

pO

2

30AE493 S (%)

q  102

ðatmÞ

X (%)

0.012 0.018 0.026 0.043 0.047

1.80 2.53 3.41 4.98 5.37

58.4 59.9 61.3 63.4 64.5

2.77 3.80 5.00 7.05 7.45

pO

2

X (%)

S (%)

q  10 2

ðatmÞ 0.028 0.046 0.065 0.081 0.089

2.78 4.38 6.24 7.79 8.10

66.7 67.0 71.9 77.4 78.5

3.70 5.80 7.50 8.30 8.40

pO

2

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

pO2 ðatmÞ

The number before the letter A of the run-number shows ethene partial pressure pC2H4  102 atm, the value after the A indicates reaction temperature, and the letter of E implies the addition of monochloroethene of 2 ppmv into the feed gas.

101

102

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

Fig. 6. Dependences of total ethylene conversion (X), selectivity to ethylene oxide (S), and the amount of oxygen consumption (q) on pO2 for the run-no. 30A453 and 50A453. 30A453: (5) X, (~) S, (*) q. 50A453: (!) X, (~) S, (*) q.

5. Discussion

Fig. 7. Suitabilities of the linear forms of (*) Eq. (18), (*) Eq. (22), and (5) Eq. (26) to the experimental data for the run-no. 50A453.

The consecutively two-step dissociative oxygen adsorption model (Model 1) was superior to the molecular oxygen adsorption model (Model 2), K2 was over 104 times K1, being uO extremely small compared to uOO and u. These results suggested that the first dissociative adsorption step of oxygen on vacant active center (Eq. (6)) was apparently very slower, compared to the second dissociative adsorption step represented by Eq. (7). The couplings of adsorbed oxygen atoms, i.e., the reverse reactions of Eqs. (6) and (7) seemed to occur hardly, because K1 3 102 and K2 3 106 on the K2SO4–Ag/aAl2O3 catalyst and, in the case of the CsRe–Ag/a-Al2O3(cr) catalyst, the reaction atmosphere was strongly reducible ( pC2 H4 3 0:30 atom) and K2 3 107 though K1 = 0.16–196.

This agrees with the facts that, during ethylene oxidation, no oxygen coupling reaction between adatoms or adatom and admolecule occurred [62,63]. Models 1 and 2 assume that oxygen adsorption is ratedetermining. If Eq. (6) lies in a complete equilibrium state (Eq. (7) is absent) and all of the surface reactions is ratedetermining, the Model 3 is presented by Eq. (26). For the pulse reaction kinetic data, this model showed concave curves, but the SSR frequently resulted in the least values in the run-no. E-1 and E-2 series at pC2 H4 2 0:01 atom (Fig. 4 and Table 4). This indicates that, only under such oxygen-rich conditions, the onestep dissociative oxygen adsorption of Eq. (6) almost reaches an

Table 8 The rate constants, the pseudo-equilibrium constants, and the SSR values of Model 1, Model 2, and Model 3 for the flow reaction data Run-number

k1 3A473 3A493 30A453 30A473 30A493 50A453 50A473 50A493 30AE473 30AE493 a

Model 2 a

Model 1

119.7 286.3 35.0 36.3 38.5 22.6 32.9 49.2 13.6 24.7

K1 75.3 196.3 3.17 3.70 4.20 1.34 3.13 5.57 0.16 1.85

b  10 20.5 20.7 4.0 8.0 9.0 2.0 2.0 2.5 1.25 1.06

4


7

K2  10 1.79 4.57 19.5 42.8 53.6 33.6 78.2 83.6 1.07 1.65

k2 33.8 63.3 27.0 28.4 44.6 18.8 25.4 36.2 13.7 27.5

Model 3a,b KO2 29.0 55.1 8.05 11.4 27.3 6.07 10.0 15.1 0.35 4.16

k3 4.89 7.96

K1e 6.72 25.4

SSR Model 1

Model 2
5

6.27  10 6.97  10
5 1.82  10
6 4.43  10
6 3.25  10
6 1.80  10
7 6.00  10
7 3.80  10
7 1.86  10
6 2.17  10
5

Model 3
4

1.61  10 2.59  10
4 2.95  10
4 1.22  10
5 4.72  10
5 8.30  10
7 7.00  10
7 1.09  10
6 3.25  10
5 6.53  10
3

1.15  10
4 3.40  10
4

Model 2 showed convex curves and Model 3 did concave curves. The ki and Ki were estimated from the approximated linear lines. The rate constant and adsorption equilibrium constant should be positive, but the slope ðK1e Þ1=2 =k3 of the Model 3 was nevertheless negative in all the runs except for the run-number. 3A473 and 3A493. As these results unsatisfied the physical meanings, Model 3 was rejected. b

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

equilibrium state (Eq. (24)). These situations are however very rare cases. In practice, the pC2 H4 adopted in the air epoxidation processes was 0.02–0.07 atm [69]. Also, in the recent commercial processes using pure oxygen as oxidant, the feed mixture of pC2 H4 = pO2 ¼ 20=7 [70] or 30/8 [71,72] has been employed. On the other hand, for the flow reaction kinetic data with feeds of pC2 H4 3 0.30 atm on the CsRe–Ag/a-Al2O3(cr) catalyst, the slope of approximated line, ðK1e Þ1=2 =k3 , was surprisingly negative, as shown in Fig. 7. Because rate constant and equilibrium constant have to be fundamentally positive value, the Model 3 could be not accepted. Assuming that Eq. (20) reaches a complete equilibrium state and the surface reactions between ethylene and oxygen admolecule are rate-determining, the relative rate equation is the same as Eqs. (22) and (23). Yong and Cant investigated the oxygen isotope exchange effect when the mixture of (C2H4 + 16O2 + 18O2) was reacted through the Ag-sponge catalyst bed, in addition to the reaction of (C2H4 + N216O + 18O2) mentioned above [63]. In this case, the isotope exchange also did not take place at all. The results indicate strongly that the stripping of adsorbed oxygen species by ethylene is much faster than the oxygen isotope exchange. They claimed that the oxygen molecular adsorption step of Eq. (27) lay in a complete equilibrium state (KOe 2 ), the route to the oxygen pool (Eq. (28)) was very slow step, and ethylene oxide and carbon dioxide were derived from the oxygen pool represented by uO. KOe

2

O2 þ u Ð uOO

(27)

uOO þ u ! 2uO

(28)

The reaction rate of Eq. (28) is represented by r4 ¼ c4 uOO u + c4 KOe 2 pO2 u2 , where uOO ¼ KOe 2 pO2 u
uO , uO is very small, and u ¼ s=ð1 þ KOe 2 pO2 Þ. So, the following equation is obtained: k4 pO2 (29) rrel ¼ ð1 þ KOe 2 pO2 Þ2 However, the plots of ð pO2 =rrel Þ1=2 versus pO2 resulted in convex curves for the kinetic data set obtained in the present work. The radius of the curvatures was near to or smaller than that for the pO2 =rrel: versus pO2 plot of Model 2. Consequently, the model of Eqs. (27) and (28) was undoubtedly inferior to Model 1. Furthermore, if Eq. (6) is rate-determining and Eq. (7) is disappeared, Eq. (18) can be also derived. This model implies that only oxygen adatom is active species and oxygen admolecule is quite absent on silver surface. As the situation was clearly in conflict with the existence of many spectroscopic evidences measured under conditions near to practical processes of ethylene oxidation [17,36,40–44], the model was ruled out in the present work. About the presence of superoxide ion (O2
) on silver catalyst surface, many investigators have reported: Gerei et al. [36] and Kilty et al. [17] observed IR band of n(O–O
) of  CH2–CH2–O–O–Ag at 870 cm
1, Seo and Sato indicated that the superoxide ion had lower entropy than (O
)ads thermo-

103

dynamically [73], Clarkson and Cirillo succeeded in the measurement of ESR signal of O2
[25], Tanaka and Yamashina found that the intensity of O2
ESR signal decreased with exposure to ethylene at 373 K [40], and also Joyner and Roberts confirmed directly the existence of oxygen admolecule on Ag catalyst exposed to oxygen gas flow using XPS and UPS [30]. Similarly, many reports for the observation of O2 or O2
species using ESR [26,27], XPS [31,32,60,74], UPS [28,29,74], and HREELS [33,34] have been published. Pettenkoffer et al. reported first the presence of O2 and O2
using SERS [35], and McBreen and Moskovits succeeded in the observation of Raman bands of O2 at 676 cm
1 and O2
at 995 cm
1 [42]. Boghosian et al. [43] and Kondarides et al. [44] observed the SERS band of 815 cm
1 assigned to the n(O–O) of  CH2–CH2–O–O–Ag on Ag heated at 573 K under reaction conditions of the epoxidation. They also confirmed that the band intensity was reduced with increasing the pC2 H4 = pO2 ratio. Nakatsuji et al. calculated theoretically by the MP2 method using the dipped adcluster model (DAM) for the oxidation of ethylene [45–47]. They concluded that, from comparison of the calculated energy barriers and the heats of reaction, the primarily important species for the epoxidation on silver catalyst was the superoxide O2
, which was in the bent end-on geometry. Ethylene molecule attacked the terminal oxygen atom of the superoxide with a barrier of only 47.3 kJ mol
1, and then the subsequent reaction proceeded quite smoothly without energy barriers and led to ethylene oxide. The overall reaction was exothermic by 155.3 kJ mol
1 from the level of (C2H4 + O2 + Ag surface) and by only 42.7 kJ mol
1 from the level of (C2H4 + superoxide on Ag surface). The atomically adsorbed oxygen, which was left on the surface after the epoxidation and/or formed by the dissociative oxygen adsorption on the bare Ag surface, had two reaction channels leading to ethylene oxide and to acetaldehyde (acetaldehyde was a precursor of carbon dioxide and water). The heats of reaction for the former and the latter reaction were exothermic by 90.0 and 212.7 kJ mol
1 from the level of (C2H4 + O2 + Ag surface), respectively. Consequently, in the reactions induced from the oxygen adatom, carbon dioxide and water was preferentially produced. The many experimental evidences described above are considered to justify one of the conclusions derived from the present work; that is, the mainly active oxygen species on silver surface during the ethylene oxidation is admolecule (perhaps, superoxide ion) and ethylene oxide is produced through the following reaction;
C2 H4 þ O
2 ! C2 H4 O þ O

(30)

Using the rate constant (k1) and the pseudo-equilibrium constants (K1 and K2), u, uO, and uOO were estimated (Tables 6 and 9). In both the pulse and flow reactions at pC2 H4 3 0.03 atm, uOO was very larger than u. The uOO in the flow reaction at pC2 H4 3 0.30 atm was only 0.12–0.34, being the u 0.66–0.87. The uO was, however, extremely small compared to the u and uOO. The abnormally small values of uO seem to be concealed behind the extremely large values of uOO, i.e., the uO estimated corresponds to the number of adatom left

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A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

after the formation of uOO and/or the epoxidation. In other words, these results mean that the dissociative oxygen adsorption on the adatom (uO) is much faster than the dissociative adsorption on vacant sites (u); the oxygen adatom seems likely to act as active center in the second dissociative oxygen adsorption process, as have been suggested by Harriott et al. [11,75] and Temkin et al. [76]. When monochloroethylene of 2 ppmv was added (30AE493), the uO values were remarkably reduced compared to those for 30A493, even though they were 10
4 order values (Table 9). The result shows that the chlorine or chloride ion produced from monochloroethylene is adsorbed selectively on such active sites (Ag) as forming the uO or oxygen adatom. This fundamentally agrees with the chlorine additive effect reported by Kilty et al. [17,37]. Force and Bell explained the in situ IR data measured on silver catalyst during the ethylene oxidation by the mechanism that the intermediates were CH2–CH2–O–Ag and CH3–CH– O–Ag and oxygen admolecule was absent [77,78]. The IR spectra were recorded under the catalytic reaction conditions, i.e., in the presence of all reactants and products. So, these spectra provide information on the long-living surface complexes prevailing in the steady state, but short-living species such as the intermediate CH2–CH2–O–O–Ag and –O– O–Ag seem to be not recorded. Backx et al. denoted that both the epoxidation and complete oxidation proceeded through an intermediate formed on the oxygen adatom and the selectivity was determined by difference in the frequencies of the subsequent two reactions [22]. van Santen and de Groot observed that, in the temperature-programmed reaction of (C2H4 + 18O2) mixture on the silver surface which was exposed to 16O2 at 473 K and subsequently quenched to room temperature in UHV, at 300–370 K only C2H416O was produced, while C2H418O began to be formed above 380 K [57]. They concluded that oxygen adatom produced ethylene oxide and the presence of subsurface oxygen promoted the formation of ethylene oxide. Grant and Lambert reported that the ethylene oxide yield was independent of dioxygen coverage but approximately proportional to the coverage of oxygen adatom in the range of 0.03 < uO < 0.09 on Ag(1 1 1) [39]. The investigators mentioned above claimed that the adsorbed oxygen species on the silver surface quenched to lower temperature in UHV were only atomic oxygen and subsurface oxygen while the superoxide ion was absent. But, direct and clear evidences for the absence of superoxide on the silver

surfaces cooled down to room temperature from the temperatures above 473 K, including spectroscopic works under practical reaction conditions, seem to be poor. It has been well known that the oxygen adsorption isobar on silver indicates a distinct maximum in the range of 453–503 K [18–20,79–82]. This suggests that, with cooling of silver samples exposed to gaseous oxygen above 453 K, the equilibriums among adsorbed oxygen species shift to nondissociative oxygen species side, i.e., the desorption side of adsorbed oxygen molecule. Using a dynamic in situ XPS measurement technique, Suzuki and Ayame confirmed that, on pure silver-selfsupporting disk exposed to O2-gas jet of 10
6 Pa in UHV, the oxygen admolecules (O2 and O2
) occupied about 44% and 64% of the total amount of surface oxygens at room temperature and 483 K, respectively [59,60]; on the Re and Csdoped silver disk resulting in high ethylene oxide selectivity, the admolecule species occupied 84% and 51% at room temperature and 483 K, respectively [60,61]. The presence of the oxygen admolecules were also observed on the silver surfaces quenched to room temperature in UHV after exposing to 133 Pa-O2 at 483 K for 2 h in a pretreatment chamber [60,61]. These experimental results are in conflict with such a consideration as no oxygen admolecule species exist on the silver surface quenched to room temperature in UHV after exposing to O2 at 473 K. From these viewpoints, the consideration that ethylene oxide is formed by the reaction between ethylene and oxygen adatom is doubtful. According to the assumption described in the Section 2, the relationships between the selectivity to ethylene oxide and the coverage of oxygen admolecule were examined. Figs. 8 and 9 show the plots of S versus uOO for the pulse and flow reaction kinetic data, respectively. The S elevated almost linearly with increase in uOO. On the other hand, the relations between S and uO of 10
4 order showed negative dependences, as can be seen in comparison of S in Table 2 and uO in Table 6 and those in Tables 7 and 9. These results support strongly the concept that oxygen admolecule is responsible for the production of ethylene oxide. In Fig. 8, the S of E-1-1 (~ mark) deviated from the linear line and that of E-1-2 (~ mark) was almost unchangeable even if the uOO increased. The slight decrease and constancy in S with increase in uOO are due to the increase in the consecutive oxidation of ethylene oxide, which is also promoted with elevating pO2 . At pC2 H4 ¼ 0:329  10
2 atm, the reaction temperatures above 506 K seem to be too high to form ethylene oxide selectively.

Table 9 The typical values of uO, uOO, and u estimated using the constants for Model 1, which are shown in Table 8 for the flow reactions 3A493

30A493

pO2 ðatmÞ

u

uO  10

0.039 0.068 0.095 0.120 0.132

0.266 0.215 0.188 0.171 0.164

5.51 4.45 3.89 3.54 3.39

3

50A493

uOO

pO2 ðatmÞ

u

uO  10

0.734 0.785 0.811 0.829 0.835

0.017 0.027 0.038 0.048 0.053

0.789 0.749 0.716 0.691 0.679

7.10 6.74 6.44 6.22 6.11

3

30AE493

uOO

pO2 ðatmÞ

u

uO  10

0.211 0.251 0.284 0.309 0.321

0.012 0.018 0.026 0.043 0.047

0.795 0.759 0.724 0.671 0.662

1.99 1.90 1.81 1.68 1.66

3

uOO

pO2 ðatmÞ

u

uO  10 3

uOO

0.205 0.241 0.276 0.329 0.338

0.028 0.046 0.065 0.081 0.089

0.814 0.774 0.743 0.721 0.711

0.86 0.82 0.79 0.76 0.75

0.186 0.226 0.257 0.279 0.289

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

Fig. 8. Relationships between uOO and S for the pulse reaction data on the K2SO4–Ag/a-Al2O3 catalyst. (~) E-1-1 (528 K), (~) E-1-2 (506 K), (5) E-13 (486 K), (!) E-2-1 (527 K), (*) E-2-2 (506 K), (*) E-2–3 (488 K), (+) E-31 (529 K), (*) E-3-2 (513 K), (&) E-3-3 (491 K), ()) E-3-4(471 K),  E-41(506 K),  E-4-2(489 K).

105

dioxide formation rate on the CsRe–Ag/a-Al2O3(cr) catalyst at 453 K, as the selectivity S was unchangeable even if pO2 increased (Fig. 6). Using k1 in Tables 5 and 8, activation energies (DEa) for the oxygen adsorption were estimated. The DEa values obtained on the K2SO4–Ag/a-Al2O3 catalyst were 87, 57, 42, and 35 kJ mol
1 at pC2 H4 ¼ 0:0033; 0:0098; 0:0170; and 0:0329 atm, respectively. On the CsRe–Ag/a-Al2O3(cr) catalyst, DEa = 84, 14, and 9 kJ mol
1 were obtained at pC2 H4 ¼ 0:03; 0:30; and 0:50 atm, respectively. Since the silver surface was usually exposed to reducing force of ethylene in the present works, with increasing pC2 H4 the surface reactions between ethylene and adsorbed oxygen species were promoted, the vacant active site (u) became larger, and then the dissociative oxygen adsorptions on the u and uO became easier. Consequently, the energy barriers for the oxygen adsorptions seem most likely to be effectively reduced with increase in pC2 H4 . The DEa values on the K2SO4–Ag/a-Al2O3 catalyst overlap with the activation energies, 53–79 kJ mol
1, for the ethylene oxide formation and 72–88 kJ mol
1 for carbon dioxide formation [6,8,13, 14,16,18]. Furthermore, DEa of 35–57 kJ mol
1 coincides with 47 kJ mol
1 of the activation energy in the route from (C2H4 + Ag–O–O
) to CH2–CH2–O–O–Ag [46,47] and with

In the 30A453 and 50A453 series (Fig. 9), the S was also hardly changed although the uOO increased from 0.12 to 0.28. These seem to suggest that the dependence of ethylene oxide formation rate on uOO almost coincides with that of carbon

Fig. 9. Relationships between uOO and S for the flow reaction data on the CsRe– Ag/a-Al2O3(cr) catalyst. (*) 30A453, (~) 30A473, (5) 30A493, (*) 50A453, (+) 50A473, (*) 50A493, (!) 30AE473, (~) 30AE493.

Fig. 10. Dependences of the activation energies of oxygen adsorption and the heats of oxygen adsorption on ethylene partial pressure. K2SO4–Ag/a-Al2O3 catalyst: (*) DEa; (~) DH1; (5) DH2. CsRe–Ag/a-Al2O3(cr) catalyst: (*) DEa; (~) DH1; (!) DH2.

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44 kJ mol
1 for the partial oxidation of ethylene on the NaCl–Ag catalyst, on which the elimination reaction of ethylene oxide hardly occur and the selectivities near to 6/7 and more are available [83–85]. Czanderna [20] and Kilty and Sachtler [17,37] reported that the activation energies of dissociative adsorption (very fast process) and molecular adsorption were 12 and 33 kJ mol
1, respectively. The 12 kJ mol
1 for the very fast dissociative adsorption agrees with the 9–14 kJ mol
1 observed on the CsRe–Ag/a-Al2O3(cr) catalyst under high ethylene partial pressures. This fact may indicate that the fast dissociative oxygen adsorption corresponds to the diffusion process of oxygen from gaseous bulk fluid to the u and uO, because the 9–14 kJ mol
1 is very near to DEa of physical processes. The extrapolation of the two plots of DEa versus pC2 H4 to pC2 H4 = 0 resulted in an activation energy of 118 kJ mol
1. This is comparable with 105 [13], 121 [19], and 136 [86] kJ mol
1, which have been reported as DEa for oxygen adsorption on silver determined using isostatic adsorption methods. The heats of adsorption (DH1 and DH2) derived from temperature dependences of the K1 and K2 were 43–89 and 91– 242 kJ mol
1 on the K2SO4–Ag/a-Al2O3 catalyst, respectively. In the flow reaction on the CsRe–Ag/a-Al2O3(cr) catalyst, DH1 = 26–90 and DH2 = 46–106 kJ mol
1 were obtained (Fig. 10). Both the DH1 and DH2 decreased remarkedly with increasing pC2 H4 . Comparison of DH1 and DH2 values indicates that the oxygen adatom (uO) is apparently unstable more than oxygen admolecule (uOO) even in the presence of ethylene. The lowering of DH values seems also to show that the stabilities or lifetimes of adsorbed oxygen species are reduced with increasing pC2 H4 . Furthermore, extraporating of DH1 for both the catalysts to pC2 H4 ¼ 0 provided about 107 kJ mol
1. The DH2 versus pC2 H4 plot for the CsRe–Ag/a-Al2O3(cr) catalyst resulted in 113 kJ mol
1 at pC2 H4 ¼ 0, while the plot of DH2 vs. pC2 H4 for the K2SO4–Ag/ a-Al2O3 catalyst reached 260 kJ mol
1. The heat of adsorption of 107–113 kJ mol
1 is comparable with the highest values of 53–119 kJ mol
1 [79], 63–105 kJ mol
1 [19], and 42– 142 kJ mol
1 [17] measured with the isostatic oxygen adsorption method and coincides with 112 kJ mol
1 in the formation of Ag–O–O
from (Ag + O2) [47]. The 260 kJ mol
1 is near to 225 kJ mol
1 for the Ag2O formation determined by means of thermodynamic decomposition equilibria of water vapor [87] and almost agrees with 257 kJ mol
1 estimated for dissociative oxygen adsorption when the adcluster is Ag4O2 [88] and the bond energy 243 kJ mol
1 of Ag and O [89]. In the reactions at pC2 H4 ¼ 0:329  10
2 atm on the K2SO4–Ag/a-Al2O3 catalyst, DH2 = 242 kJ mol
1 was obtained and the uOO was very large value such as 0.79–0.96. So, it is easily stipulated that the silver surface at pC2 H4 ¼ 0 is completely oxidized to silver (I) oxide, which is further covered by oxygen adatoms. 6. Conclusions Oxygen partial pressure-depending kinetic data were measured at constant partial pressure of ethylene, using a pair of the pulse reaction method and the well-stabilized K2SO4–

Ag/a-Al2O3 catalyst and that of the flow reaction method and the CsRe–Ag/a-Al2O3(cr) catalyst. These kinetic data were used in the analysis of the three kinds of kinetic models, which were derived from such a redox model as the rate of surface reduction with ethylene (i.e., surface reaction rate) was equal to the rate of oxygen adsorption on the silver catalyst surface when the ethylene oxidation proceeds. One is Model 1 based on the assumption that the two-step dissociative oxygen adsorption processes are rate-determining and form both of oxygen adatom and admolecule, second is Model 2 assuming that the molecular oxygen adsorption is rate-determining and forms oxygen admolecule only, and third is Model 3 postulating that the one-step dissociative oxygen adsorption process forming oxygen adatom only lies at a complete equilibrium state and the surface reaction, which is initiated by the interactions between ethylene and oxygen adatom, is ratedetermining. The suitabilities of these models, that is, relative rate equations to the kinetic data were examined using: (i) linearity of the plot of subordinate parameter versus independent parameter in the linear forms of the rate equations, (ii) whether the rate constants and the pseudo-equilibrium constants of oxygen adsorption steps are positive or not, (iii) comparison of the SSR values for the q parameter, and (iv) reproducibility with regard to the q
pO2 experimental curves. The examination and discussion for the (i)–(iv) led to the following conclusions. (a) The suitability of Model 1 to the kinetic data was superior to Model 2 and Model 3. Model 1 and Model 2 represents that ethylene oxide is formed by the Rideal–Eley mechanism between ethylene in gas-phase and adsorbed admolecule or superoxide ion, while Model 3 means that ethylene oxide is produced from gaseous ethylene and oxygen adatom. (b) The surface oxygen coverages were estimated using the rate constant k1 and the pseudo-equilibrium constants (K2 was at least 104 times K1) for Model 1: the uO was of 10
3 to 10
4 order, the uOO was 0.12–0.34 at pC2 H4  0:30 atm and larger than 0.48 at pC2 H4 0:03 atm, and the u balanced value. The abnormally small values of the uO seem likely to correspond to the number of adatom left after the formation of uOO, that is, the second dissociative oxygen adsorption on the adatom (uO) is much faster than the first dissociative adsorption on vacant sites (u) and the oxygen adatom acts as active center in the second oxygen adsorption. (c) Good linear relationships were observed between the selectivity to ethylene oxide and the estimated uOO. (d) Activation energies for the dissociative oxygen adsorption during the course of ethylene oxidation remarkably depended on pC2 H4 ; that is, DEa was effectively reduced with increasing pC2 H4 . The activation energies estimated by extraporation of DEa versus pC2 H4 plots to pC2 H4 ¼ 0 were comparable with literature values obtained by isostatic measurements. (e) Heat of adsorption, extraporated to pC2 H4 ¼ 0, was also comparable with literature values, which are determined by the isostatic oxygen adsorption and the desorption kinetics of adsorbed oxygen, and the values estimated

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theoretically. The heat of adsorption decreased vigorously with the increase in pC2 H4 . This means that the adsorbed oxygen species become very unstable in the presence of ethylene. Acknowledgements The authors would like to express their grateful acknowledgment to Professor Noel W. Cant of Macquarie University, Sydney, Australia and Professor Hiroshi Nakatsuji, Kyoto University, Kyoto, Japan for valuable discussion. References [1] A. Ayame, H. Kanoh, Shokubai 15 (1973) 151. [2] X.E. Verykios, F.P. Stein, R.W. Coughlin, Catal. Rev. Sci. Eng. 22 (1980) 197. [3] R.A. van Santen, H.P.C.E. Kuipers, Adv. Catal. 35 (1987) 265. [4] D. Kondarides, Y. Iwasawa, Hyomen 32 (1994) 295. [5] A. Ayame, Shokubai 38 (1996) 205. [6] A. Orzechowsky, K.E. MacCormack, Can. J. Chem. 32 (1954), 388, 399, 432 and 443. [7] L.G. Nault, O.W. Bolme, L.N. Johanson, I. & E.C., Proc. Des. Dev. 1 (1962) 285. [8] A.I. Kurilenko, N.V. Kul’kova, L.P. Baranova, M.I. Temkin, Kinet. Katal. 3 (1962) 208. [9] T. Seiyama, S. Kagawa, K. Kurama, Kogyo Kagaku Zasshi 70 (1967) 1137. [10] Sh. Kagawa, M. Iwamoto, H. Mori, T. Seiyama, J. Phys. Chem. 85 (1981) 434. [11] P.D. Klugherz, P. Harriott, AIChE J. 17 (1971) 856. [12] C.T. Campbell, M.T. Paffet, Surf. Sci. 139 (1984) 396. [13] G.H. Twigg, Proc. Roy. Soc. (Lond.) A 188 (1946), 92, 105 and 123. [14] O.M. Todes, T.I. Andrianova, Zh. Fiz. Khim. 27 (1953) 1485. [15] H. Kano, A. Ayame, Mem. Muroran Inst. Tech. 4 (1964) 871. [16] A. Ayame, H. Kanoh, T. Kanazuka, H. Baba, Bull. Jpn. Petrol. Inst. 15 (1973) 150. [17] P.A. Kilty, N.C. Rol, W.M.H. Sachtler, in: Proceedings of the Fifth International Congress Catal., Palm Beach, No. 64, (1972), p. 929. [18] I. Yasumori, T. Tazumi, T. Sasamoto, Shokubai 10 (1968) 174P. [19] W.W. Smeltzer, E.L. Tollefson, A. Cambron, Can. J. Chem. 34 (1956) 1046. [20] A.W. Czanderna, J. Phys. Chem. 68 (1964) 2765. [21] D.I. Kondarides, X.E. Verykios, J. Catal. 143 (1993) 481. [22] C. Backx, C.P.M. de Groot, P. Biloen, Surf. Sci. 104 (1981) 300. [23] Yu.Ts. Vol, N.A. Shishakov, Izv. Akad. Nauk SSSR, Otd. Khim. Nauk (1962) 586. [24] Sh. Kagawa, H. Tokunaga, T. Seiyama, Kogyo Kagaku Zasshi 71 (1968) 775. [25] R.B. Clarkson, A.C. Cirillo Jr., J. Catal. 33 (1974) 392. [26] N. Shimizu, K. Shimokoshi, I. Yasumori, Bull. Chem. Soc. Jpn. 46 (1973) 2929. [27] A. Abou-Kais, M. Jarjoui, J.C. Verdrine, P.C. Gravelle, J. Catal. 47 (1977) 399. [28] M.K. Rajumon, K. Prabhakaran, C.N.R. Rao, Surf. Sci. Lett. 233 (1990) L237. [29] K.C. Prince, A.M. Bradshaw, Surf. Sci. 126 (1983) 49. [30] R.W. Joyner, M.W. Roberts, Chem. Phys. Lett. 60 (1979) 459. [31] C.T. Campbell, Surf. Sci. 157 (1985) 43. [32] H. Miura, A. Ayame, H. Kanoh, K. Miyahara, I. Toyoshima, Shinku 25 (1982) 302. [33] B.A. Sexton, R.J. Madix, Chem. Phys. Lett. 76 (1980) 294. [34] K. Prabhakaran, C.N.R. Rao, Surf. Sci. 186 (1987) L575.

107

[35] C. Pettenkoffer, I. Pockard, A. Otto, Surf. Sci. 135 (1983) 52. [36] S.V. Gerei, K.M. Kholyavenko, M.Ya. Rubanik, Ukrain Khim. Zhur. 31 (1965) 449. [37] P.A. Kilty, W.M.H. Sachtler, Catal. Rev. Sci. Eng. 10 (1974) 1. [38] R.B. Grant, R.M. Lambert, Surf. Sci. 146 (1984) 256. [39] R.B. Grant, R.M. Lambert, J. Catal. 92 (1985) 364. [40] S. Tanaka, T. Yamashina, J. Catal. 40 (1975) 140. [41] D. Hermerschmidt, R. Haul, Ber. Bunsenges. Phys. Chem. 85 (1981) 739. [42] P.H. McBreen, M. Moskovits, J. Catal. 103 (1987) 188. [43] S. Boghosian, S. Bebelis, C.G. Vayenas, G.N. Papatheodorou, J. Catal. 117 (1989) 561. [44] D.I. Kondarides, G.N. Papatheodorou, C.G. Vayenas, X.E. Verykios, Ber. Bunsenges. Phys. Chem. 97 (1993) 709. [45] H. Nakatsuji, H. Nakai, Can. J. Chem. 70 (1992) 404. [46] H. Nakatsuji, H. Nakai, J. Chem. Phys. 98 (1993) 2423. [47] H. Nakatsuji, H. Nakai, K. Ikeda, Y. Yamamoto, Surf. Sci. 384 (1997) 315. [48] C.S.A. Fang, Surf. Sci. Lett. 235 (1990) L291. [49] H.A. Engelhardt, D. Menzel, Surf. Sci. 57 (1976) 591. [50] M. Bowker, M.A. Barteau, R.I. Madix, Surf. Sci. 92 (1980) 528. [51] M.A. Barteau, R.I. Madix, Surf. Sci. 97 (1980) 101. [52] C.T. Campbell, M.T. Paffet, Surf. Sci. 143 (1984) 517. [53] A.I. Boronin, V.I. Bukhtiyarov, A.L. Svishnevskii, G.K. Boreskov, V.I. Savchenko, Surf. Sci. 201 (1988) 195. [54] T. Hashizume, M. Taniguchi, K. Motai, H. Lu, K. Tanaka, T. Sakurai, Surf. Sci. 266 (1992) 282. [55] C. Rehren, G. Isaac, R. Schogl, G. Ertl, Catal. Lett. 11 (1991) 253. [56] X. Bao, J.V. Bath, G. Lehmpfuhl, R. Schuster, Y. Uchida, R. Schogl, G. Ertl, Surf. Sci. 284 (1993) 14. [57] R.A. van Santen, C.P.M. de Groot, J. Catal. 98 (1986) 530. [58] Ch.J. Bertole, Ch.A. Mims, J. Catal. 184 (1999) 224. [59] A. Ayame, K. Yamazaki, A. Miyazaki, Reports of the Asahi Glass Foundation 55 (1989) 335. [60] H. Suzuki, A. Ayame, Nippon Kagaku Kaishi (1992) 930. [61] A. Ayame, H. Suzuki, T. Abiko, Hyomen Kagaku 14 (1993) 558. [62] S. Yokoyama, K. Miyahara, J. Res. Inst. Catal., Hokkaido Univ. 22 (1974) 63. [63] Y.C. Yong, N.W. Cant, Appl. Catal. 48 (1989) 37. [64] O.A. Hougen, K.M. Watson, Chemical Process Princples—Part III (Kinetics and Catalysis), John Wiley & Sons Inc., London, 1947 , p. 902. [65] A. Ayame, Y. Shibuya, T. Yoshida, H. Kanoh, Nippon Kagaku Kaishi (1973) 2063. [66] A. Ayame, A. Numabe, T. Kanazuka, H. Kanoh, Bull. Jpn. Petrol. Ins. 15 (1973) 142. [67] A. Ayame, Y. Uchida, Japan Patent 2000-44331 (2000) to Sumitomo Chemical Co. Ltd. [68] A. Ayame, Y. Uchida, H. Ono, M. Miyamoto, T. Sato, H. Hayasaka, Appl. Catal. A: Gen. 244 (2003) 59. [69] A. Ayame, Shokubai Koza, vol. 7, Kodansha Saientifikku Ltd., Tokyo, 1985, p. 170. [70] Sh. Nagase, H. Tanabe, H. Imai, Japan Patent H5-329368 (1993) to Nippon Shokubai Co. Ltd. [71] J.E. Batsfam, W.H. Gardes, R.M. Kowareski, Japan Patent H5-84440 (1993) to Shell Intern. Res. Maatschappij B.V. [72] K. Hertzauk, K. Beaning, U. Pryucan, V.D. Muros, M. Schvartzman, Japan Patent H4-3177411 (1992) to BASF Aktiengellschaft. [73] M. Seo, N. Sato, Denki Kagaku 40 (1972) 557. [74] H. Miura, A. Ayame, H. Kanoh, K. Miyahara, I. Toyoshima, Shinku 26 (1983) 406. [75] P.L. Metcalf, P. Harriott, Ind. Eng. Chem., Process Des. Develop. 11 (1972) 478. [76] M.I. Temkin, N.V. Kul’kova, Kinet. Katal. 16 (1975) 1211. [77] E.L. Force, A.T. Bell, J. Catal. 38 (1975) 440. [78] E.L. Force, A.T. Bell, J. Catal. 40 (1975) 356. [79] A.F. Benton, L.C. Drake, J. Am. Chem. Soc. 56 (1934) 255. [80] N. Sato, M. Seo, Denki Kagaku 38 (1970) 649.

108

A. Ayame et al. / Applied Catalysis A: General 304 (2006) 93–108

[81] A. Ayame, Y. Ito, H. Kanoh, T. Kanazuka, Mem. Muroran Inst. Tech. 8 (1973) 107. [82] L.M. Akella, H.H. Lee, J. Catal. 86 (1984) 465. [83] A. Ayame, H. Miura, H. Kimura, M. Yamaguchi, H. Kanoh, K. Miyahara, I. Toyoshima, J. Res. Inst. Catal., Hokkaido Univ. 32 (1984) 49. [84] A. Ayame, T. Kimura, M. Yamaguchi, H. Miura, N. Takeno, H. Kanoh, I. Toyoshima, J. Catal. 79 (1983) 233.

[85] A. Ayame, T. Yoshida, M. Yamaguchi, H. Miura, Y. Sakai, N. Nojiri, J. Catal. 100 (1986) 401. [86] Y.L. Sandler, D.D. Durigon, J. Phys. Chem. 69 (1965) 4201. [87] D. Gonzalez, G. Parravano, J. Am. Chem. Soc. 78 (1950) 4533. [88] H. Nakatsuji, H. Nakai, J. Phys. Chem. 98 (1993) 2423. [89] L.Ya. Margolis, Adv. Catal. 14 (1963) 429.