J. Phys. Chem. 1994,98, 1933-1938
1933
FT-IRSpectroscopic Studies of Methane Adsorption on Magnesium Oxide Can Li,' Guoqiang Li, and Qm Xin State Key Laboratory of Catalysis, Dalian Institute of Chemical Physics, Chinese Academy of Sciences, Dalian 1 1 6023, China Received: November 15, 1993'
-
Adsorption of CH4 and coadsorption of CH4 and C O on differently treated magnesium oxide have been studied by FT-IR spectroscopy a t 173-273 K. Five IR bands at 3008,3000,2900,2890,and 1306 cm-1 were observed when well-outgassed MgO was exposed to CH4 at 173 K. The bands at 3008 and 3000 cm-l are assigned to a degenerate stretch vibration of CH4, and those at 2900 and 2890 cm-1 originate from an infrared-forbidden mode at 2917 cm-' of free CH4. These bands became noticeably stronger with the outgassing temperature of MgO but could be sufficiently reduced as MgO was pretreated in air or under water vapor. Adsorbed CO on MgO gives two IR bands at 2161 and 21 55 cm-' which exhibit a parallel variation with those bands of adsorbed CH4 with the different pretreatments of MgO. Coadsorption of CH4 and CO indicated that adsorbed CO can deplete part of the adsorbed methane, particularly that with IR bands at 2900 and 2890 cm-1. It is proposed that the adsorbed methane can be substituted by adsorbed C O coordinated to a Lewis acid-base pair site, Mg2+LC02-LC, and the methane unaffected by coadsorbed C O interacts mainly with surface oxygen anions alone. The observed IR bands can be attributed to two types of adsorbed methane on MgO: one with IR bands at 3008 and 1305 cm-l due to weakly adsorbed methane on surface oxygen anions and another with bands at 3000,2900,2890,1305 cm-l due to methane strongly interacting with the Lewis acid-base pair sites of MgO. A possible mechanism of methane interaction with surface of MgO is discussed.
Introduction Methane adsorption and activation have received much attention in recent years.l-ll Methane is chemically the most stable molecule among hydrocarbons, and the activation of methane has presented a formidable challenge in catalysis. The conversion of methane into ethylene, liquid fuels, and chemicals has been intensely explored worldwide during recent decades due to its commercial potential.12-17 Activation of methane on metal oxide surfaces in the presence of oxygen species attracted more interest because methane activation assisted by oxygen is a most promissing way to achieve the tough goal of utilizing methane. Therefore study on the adsorption and activation of methane on metal oxides became a subject of current interest. MgO has been demonstrated to be an active catalyst for activating methane even at low temperatures18.19and extensively studied in methane oxidative coupling.20.21 A common consensus about methane activation on MgO and many other basic metal oxdies is that the surface active oxygen species are majorly responsible for the first hydrogen abstraction of methane. However, how the methane interacts with surface oxygen species and what the role played by other surface species such as surface metal cations, hydroxyl, and carbonate are still far from understood. The studies on the initial interaction of methane with surface active species or sites are necessary to gain an insight into the mechanism of methane activation. Previous worklOJ1found that at least two types of adsorbed methane were formed on cerium oxide, one interacts with normal surface oxygen anion and other interacts with @-like species. The structure of the latter was notably distorted, and accordingly an infrared-forbidden vibration at 2917 cm-1 was not only made infrared detectable but alsodown shifted to 2875 cm-l indicating that the adsorbed methane was activated and most possibly the species of predissociation. The present work is to investigate methane adsorption on differently treated MgO surfaces by FTIR spectroscopywith a hope to know the details of the interaction and the activation of methane on the surface of a typical basic
* To whom correspondence should be addressed. *Abstract published in Advance ACS Abstracts, January 1, 1994.
oxide. A coadsorption of CH4 and CO was performed to clarify the role of magnesium cation played in the methane adsorption and activation.
Experimental Section MgO was obtained from a commercial source (Emerck Co.), and the BET surface area was measured to be -40 m2/g. MgO was pressed into a self-supporting disc with a weight of 20 mg for infrared study. An IR cell used in the study is made of quartz, and the sample disc in the cell can be treated in various ways such as outgassing, oxidation, and reduction at a wide temperature rangefrom 100to 1OOOK. ChandCOadsorptionwasperformed at 173 K, which is 60K higher than the condensation temperature of methane, in order to avoid heavy absorption of physically adsorbed methane. There are large amounts of adsorbed water, carbonate, and hydroxyl specieson the commercial MgO surface, and this sample is so-called C02- and HzO-covered MgO. A carbonate-free MgO surface was obtained by treating MgO in vacuo at 973 K, and most of the surface hydroxyls were also removed after the high-temperature outgassing. An air-contaminated MgO sample was prepared by exposing the welloutgassed MgO to air at 573 K. A hydroxylated MgO sample was made by treating the well-outgassed MgO under water vapor (18 Torr of H20) at 773 K. The CH4 (>99.99%) and CO (>99.99%) were further purified by a liquid nitrogen trap before use. The adsorption pressure of methane was about 1 1 Torr and was maintained by a liquid nitrogen trap. Coadsorption of CH4 and CO was realized by adding about 10 Torr of CO to the methane prefilled system, and the CO pressure at the surface of MgO was controlled by slowly diffusing CO onto the sample through a glass pipe with a diameter of 8 mm. By this method, the coadsorption variation of CH4 and CO with gradually increasing CO pressure could be followed. IR spectra were recorded on a Perkin-Elmer 1800 doublebeam FT-IR spectrometer equipped with a liquid nitrogen cooled mercury-calcium-telluride (MCT) detector, with four scans, and at a nominal resolution of 4 cm-l. Integrated absorbances were calculated from the spectra referenced to the background spectra
0022-3654/94/2098-1933904.50/0 0 1994 American Chemical Society
Li et ai.
1934 The Journal of Physical Chemistry, Vol. 98, No. 7, 1994
TABLE 1: Vibrational Modes of CI& and Observed IR Bands of C& Adsorbed on Well-Outeassed MpO (in cm-1) vibrational mode gas phase adsorbed freq shift 01, sym stretch 29 17" 2900 17 u2, degen deform. u3, degen stretch v4,
degen deform.
1533" 3019 1306
u1+ 04
2890
27
b
b
3008 3000 1305 2828
11 19 1
Infrared inactive. Not observed.
I
I
3100
3300
I
I
2900
2700
Wavenumberkm-' Figure 1. IR spectra of adsorbed CH4 at 173 K on MgO after different pretreatments: (a) well outgassed, (b) hydroxylated, and (c) air
contaminated.
3400
3000
2700
2000
1400
1000
Wavenumberkm-' Figure 2. IR spectra of coadsorbed CHI and CO on well-outgassedMgO at 173 K (a) in CH4 for 30 min, (b) after admission of CO for 5 min, and (c) 60 min after (b). which were taken prior to the adsorption at corresponding temperatures. Results
CI& Adsorbed on Differently Treated MgO Surfaces. When highly outgassed MgO was exposed to CH.,at 173 K there are at least six distinct IR bands that were observed at 3008, 3000, 2900,2890,2828, and 1305 cm-l as shown in Figures l a and 2a. These bands became weaker upon warming the sample and disappeared at around 253 K. The variations of these band intensities are not in parallel, and the bands at 2900 and 2890 cm-1 seem more persistent during the temperature elevation.These bands are attributed to adsorbed methane on the MgO surface, and the IR spectrumof the adsorbedmethane is somewhat similar to that of methaneadsorbedon cerium oxide.lOJ1The assignment of the observed IR bands is collected in Table 1. There are four vibrational modes, v I - v ~ , for free methane whose frequenciesare 2917, 1533, 3019, and 1306 cm-l, respectively. Of these modes u1 and v2 are infrared-forbiddenmodes, and v3 and vqare infraredactive modes. For the adsorbed methane on MgO, the bands at
3008 and 3000 cm-* are attributed to the v3 mode which shows a red-shift, and the band at 1305 cm-l is due to theu4 mode whose position hardly shifts for adsorption. The weak band at 2828 cm-1 can be assigned to the combination mode of u2 and ~ 4 . 2 2The bands at 2900 and 2890 cm-1 originatefrom the infrared-forbidden modeat 2917 cm-l offreemethane.11.22The fact that theinfraredforbidden mode not only becomes infrared observable but also shifts to lower frequencies implies that the adsorbed methane interacts so strongly with the MgO surface that it leads to a breakdown of the infrared selection rule. Figure 1 gives the infrared spectra of adsorbed methane on three differently treated MgO surfaces (the 1305-cm-1 band is not included in Figure 1 since its position is almost the same for the three surfaces). For the well-outgassed MgO (pretreated at 973 K invacuo) adsorbed methane shows very strong bands at 3008 and 3000 cm-' while these bands are relatively weak for methane adsorbed on either the air-contaminatedor the hydroxylated MgO. Another striking contrast is that the band at 2890 cm-' appeared for the highly outgassed MgO (Figure la) but disappearedfor the hydroxylated and air-contaminated MgO (Figure l b and IC). Although the bands at 2900 and 2890 cm-1 seem weak, they represent large amounts of adsorbed methane species because the infrared absorption coefficients of these modes are very small. Therefore the disappearance of the band at 2890 cm-l indicates that a considerable amount of surface sites available for methane adsorptionis destroyed by the air contaminationor hydroxylation. It is clear that the methane adsorption is very sensitive to the surface states. The IR bands in Figure 1cannot simply be ascribed to the physically adsorbed methane but are ascribed to the adsorbed species involved with certain chemical interaction with surface sites. Coadsorption of CH, and CO on the MgO Surface. Most of the studies of methane oxidative coupling over metal oxides concluded that the activationof methane is majorly due tosurface active oxygen species, i.e., the first hydrogen abstraction of methane is realized on the surface Lewis basic sites. Less attention has been paid to the surface cations of the metal oxide, Le., the Lewis acid sites. In order todistinguish the surfacesites available for the formation of adsorbed methane, coadsorptionof CH4 and CO on MgO was performed at 173 K. Figure 2 shows the IR spectra recorded after introducing CO onto the MgO surface with preadsorbed methane. CO was admitted onto the MgO surface through a long pipe which can control the CO from zero to a desired pressure by a slow diffuse process. Immediately following CO admission, the spectrum contains only the IR bands of adsorbed methane and the IR bands of adsorbed CO are not observed since CO does has not yet diffused onto the MgO surface (see Figure 2a). Two IR bands at 216 1 and 2 155 cm-' due to adsorbed CO appeared after 5 min as shown in Figure 2b. Correspondingly,all the IR bands of adsorbed methane are attenuated, especially the bands at 2900 and 2890 cm-1, which are noticeably reduced. After a long standing, the IR bands of adsorbed CO grow markedly and those of adsorbed methane decline very much, in particular, the bands at 2900 and 2890 cm-1 nearly vanish (see Figure 2c). It is straightforward that part of the adsorbed methane could be replaced by adsorbed CO. In particular, the adsorbed methane
The Journal of Physical Chemistry, Vol. 98, No. 7, 1994 1935
Methane Adsorption on Magnesium Oxide
0
‘-CO 0
in 20
I
l
60 80 Timelmin. 40
l
100
Figure 3. Variations of IR intensities of adsorbed CH4 and adsorbed CO on well-outgassed MgO at 173 K with standing time.
with IR bands at 2900 and 2890 cm-1 can be replaced completely by the adsorbed CO. The methane species with the remaining bands at 3008 and 1305 cm-l possibly adsorb on the sites which are unfavorable for CO adsorption. This result clarifies that at least two types of adsorbed methane species are derived on the well-outgassed MgO surface, one can be depleted by adsorbed CO and the other is intact with the coadsorbed CO. These differently adsorbed methane species should be grouped into two types according to their IR bands, Le., 3008 and 1305 cm-l and 3000, 2900(2890), and 1305 cm-1. The band at 1305 cm-l is common for the two species. The species giving the IR bands at 3008 and 1305 cm-1 is ascribed to the methane interacting with surface oxygen anions denoted as species 11, similar to that for cerium oxide.ll The species with the IR bands at 2900 and 2890 cm-1 may be the adsorbed methane that interacts with both surface magnesium cation and oxygen anion. Figure 3 illustrates the detail variations of the IR band intensities (measured as integrated absorbance) of adsorbed CHI and CO on the MgO surface. The integrated absorbance of adsorbedCO centered at 21 57 cm-1 was calculated for both bands at 2161 and 2155 cm-I. The amount of adsorbed methane was measured by the infrared integrated absorbance of the band at 1305 cm-1. With the increase of the band absorbanceof adsorbed CO, that of adsorbed methane decreased simultaneously. Finally, the absorbance of adsorbed CO reaches more than 7 arbitrary units but the band absorbance of 1305 cm-1 falls from 5 to 3 arbitrary units indicatingthat about 40%of the adsorbedmethane is replaced by the coadsorbed CO. The adsorbed methane remaining on the MgO surface is primarily species I1 as noted above. It is generally believed that CO adsorbed on a metal oxide interacts with surface Lewis acid sites, Le., surface Mg2+ cations for MgO.23-25 The two bands at 2161 and 2155 cm-* of the adsorbed CO reflect that there are two kinds of Mg2+sites with slightly different chemical environments. It is explicit that the methane replaced by adsorbed CO interacts not only with surface oxygen species but also with Mg2+ cations; in other words, the methane adsorbs on a Lewis acid-base pair site. The interaction of CH4 with surface Mg2+ cations is not stronger than that of CO, which can ocupy the Lewis acid sites irrespective of the presence of adsorbed methane. Adsorption of CH4 and CO on MgO Outgassed at Various Temperatures. Both CH4 and CO adsorbed on well-outgassed MgO producedverystrong IR bands due to their adsorbed species whereas these bands were reduced significantly when the welloutgassed MgO was hydroxylated or air contaminated. This strongly suggests that the adsorption of CH4 and CO on MgO is fairly sensitive to the surface state or is highly selective to the surface sites. With an intention to understand the surface sites available for the formation of adsorbed methane species the adsorption of CH4 and CO on MgO was carried out for the
TemperatureIK
Figure4. IR intensities of adsorbed CHq and CO at 173 K on hydroxylated
MgO which was outgassed at various temperatures.
’L
0273
473
673
873
1(
’3
Temperature / K
Figure 5. IR intensities of adsorbed CH4 and CO at 173 K on aircontaminated MgO which was outgassed at various temperatures.
hydroxylatedand air-contaminated MgO that was pre-outgassed at different temperatures from 298 to 973 K. Figure 4 presents the integrated absorbance of 1305 cm-l (adsorbed CH4) and 2157 cm-1 (adsorbed CO) measured for their adsorption on hydroxylated MgO that was pre-outgassed at 373,573, 773,873,and 973 K,respectively. By comparison with the data in Figure 3, the amount of the adsorbed CH4 and CO on hydroxylated MgO (outgassed at 373 K) is decreased to 60% and 15% of that on well-outgassed MgO, respectively. The results definitely show that the hydroxylation heavily destroys the surface sites for the adsorption of CH4 and CO. With the elevationofoutgassingtemperature from 373 to973 K, theamount of both adsorbed CH4 and CO is enhanced greatly and eventually reaches the level of that prior to the hydroxylation as shown in Figure 3. The reproducible data corroborate that most of the surface sites of adsorbed CH4 and CO can be regenerated after a dehydroxylation treatment. The parallel enhancement of the adsorbed CHI and CO with outgassing temperature manifests that some surface sites may be common for the adsorption of CH4 and CO. In a way similar to the adsorption of CH4 and CO on hydroxylated MgO, the adsorption of CH4 and CO on aircontaminatedMgO was also investigated and the results are shown in Figure 5. After the treatment of MgO in air, the amount of adsorbed CH4 and CO obviously decreased, and by comparing with that on well-outgassed MgO only 40% of the adsorbed CH4 and 25% of the adsorbed CO remained for the air-contaminated MgO. The air contamination reduces the surface sites for the adsorption of CH4 and CO in much the same way as the hydroxylation does. In contrast with the hydroxylation treatment,
1936 The Journal of Physical Chemistry, Vol. 98, No. 7, 1994
Li et al.
TABLE 2: Amounts of Adsorbed CH.I (1305 cm-l) and CO (2157 cm-l) on Differently Pretreated MgO (in Integrated Absorbance) well outgassedo air contaminatedb hydroxylatedC CHI
5.0 7.4
co
2.1 1.9
3.1 1.1
a Outgassed at 973 K. b Well-outgassed MgO was exposed to air at 773 K and then evacuated at 298 K. Well-outgassed MgO was treated under H20 vapor at 773 K and then evacuated at 373 K.
8
C
31
3iOO
4600
IO
Wavenumberlcm-'
Figure 7. IR spectra of hydroxylated MgO outgassed at elevated temperatures: (a) 373, (b) 573, (c) 773, (d) 873, and (e) 973 K. 4000
3600
2600 1500 Wavenumberkm-'
1600
Figure 6. IR spectra of commercial MgO outgassed at elevated temperatures: (a) 373, (b) 473, (c) 573, (d) 673, (e) 773, and (f') 873 K a spectrum at 973 K was taken as a reference background. the air-contamination is more effective for poisoning the surface sites of methane adsorption. The amount of adsorbed CH4 increases monotonously with the elevation of outgassing temperature, and that of adsorbed CO increases steeply above 573 K. The same amounts of adsorbed CH4 and CO as that on a well-outgassed MgO can be achieved when the air-contaminated MgO was outgassed at 973 K. The results substantiate that the surfacesites poisoned by air contaminationcan be recovered again through outgassing at 973 K. Table 2 lists the integrated absorbance of adsorbed CH4 and CO on well-outgassed, hydroxylated, and air-contaminated MgO. Hydroxyls and Carbonate Species on MgO Surface. Figure 6 shows IR spectra of a commercial MgO sample (Emerck Co.) outgassed in a stepwise manner from 373 to 973 K. The spectra recorded at different temperatures are compared with a background spectrum which was taken after a prolonged evacuation at 973 K. The background spectrum of MgO (not given here) exhibits almost no bands due to surface carbonate species and only very weak bands due to surface hydroxyls. The spectrum of MgO outgassed at room temperature shows extremely strong bands in 3800-2500-cm-1 and 1700-1 300-cm-1 regions. The bands in the two regions are attributed to surface hydroxyls together with adsorbed water and surface carbonate species, respectively. After an outgassing at 373 K, physically adsorbed water was removed and therefore the spectrum (Figure 6a) displays distinguishable bands at 3675,3475, and 3100 (broad) cm-1 due to surface hydroxyls but still very strong bands of surface carbonate species. The relatively weak bands that appeared at 1115, 885, and 855 cm-' are attributed to v,,(OCO) and 6(COJ)of surface carbonate species.26 The split bands at 885 and 855 cm-I indicate that at least two kinds of surface carbonate species exist on MgO. The very strong bands of the surface carbonate species are indicative of the large amount of C 0 2 adsorbed on surface basic sites of MgO. The carbonate and hydroxyl species are difficult to remove at 673 K even in vacuo (Figure 6a-c). Obviously, most of the surface hydroxyls and carbonate species can be removed only after the outgassing at 873 K (Figure6f). The stronglyadsorbedcarbonate species imply that MgO is a strong base oxide whose surface base sites could be derived after an outgassing at high temperatures. The weak
-
0-c
I I -Mg-O-
//O
.,,
.
I
,
I
.. ,
-Mg-OC.U.S.
sites
Figure 8. Schematic description of the formation of surface cus sites. band at 3735 cm-l in Figure 6f may represent the sparselylocated structural hydroxyls on the highly outgassed MgO. When MgO is outgassed at 973 K, the carbonate species and most of the hydroxyls can be removed and, accordingly,the surface is highly decarbonated and dehydroxylated. The dehydroxylated MgO was treated under the vapor of H2O (18 Torr) at 773 K in order to regenerate the surface hydroxyls. Figure 7a is an IR spectrum of hydroxylated MgO. The distinct bands at 3764, 3750,3706,3686,3493, and 3600 (broad) cm-1are representative of five individual hydroxyls reproduced by the hydroxylation. From the band intensities in Figures 6 and 7, however, only a small part of the surface hydroxyls can be recovered. These hydroxyls were gradually removed in vacuo at elevated temperatures (Figure 7a-e), and only a weak band at 3750 cm-' can be maintained after an outgassing at 973 K as seen in Figure 7e. When a well-outgassed MgO samplewas treated in air at room temperature or higher temperatures, surface carbonate species and a small amount of hydroxyls can be reproduced as evidenced by the IR spectrum. The surface carbonate speciesand hydroxyls on theair-contaminated MgOcan beremoved invacuo at elevated temperatures in a similar fashion to that in Figure 6. Figures 6 and 7 demonstrated that the outgassing at high temperature can effectively remove the surface carbonate and hydroxyls on MgO. The dehydroxylation and decarbonation result in surface coordinatively unsaturated (cus) sites that are surfacemagnesium cation and oxygen anion with low coordination, commonly denoted as Mg2+LC and 0 2 - ~ c . The formation of surface cus sites can be schematically illustrated as in Figure 8. The surface cus sites are naturally the Lewis acid and base centers which may be generated conjunctly. Figure 8 also explains how the surface cus sites could be destroyed by the reverse reaction, carbonation and hydroxylation. The IR bands at 2161 and 2155 cm-l in Figure 2 are readily assigned to the CO adsorbed on surface Mg2+LCsites. The twin bands indicate that there are different Mg2+LCsites derived on the MgO surface after a good outgassing. The IR band intensities
Methane Adsorption on Magnesium Oxide of the adsorbed CO could be approximately proportional to the surface concentration of Mgz+Lc sites. As a consequence, the amount of adsorbed CO increases with the elevation of outgassing temperature (Figures 4 and 5) since the higher the outgassing temperature the more surface cus sites are produced. The data in Table 2 explicitly demonstrate that most of the surface cus sites can be saturated by either hydroxyls or carbonate species. Discussion Details of the interaction of methane with the surface of a metal oxide are much less known since methane is normally difficult to capture by the surfaces. It may be assumed that the initial interaction between methane and the surface of a catalyst is naturally related to the activation as well as the reaction of methane, e.g., C-H bond rupture to produce CH3 radical as proposed for methane oxidative c o ~ p l i n g . An ~ ~ early * ~ ~ study of methane adsorption on the Si02 surface was made by Sheppard and Yates2Zwho observed an IR band a t 2900 cm-l for adsorbed methane. They have assigned the band to the infrared-forbidden mode, u1, which became infrared active due to the interaction by surface force. In previous work,lOJ1we observed the infraredforbidden band for adsorbed methane on well-outgassed CeO2. This band was not only intense but shifted from 2917 to 2875 cm-I. The large red-shift of the u1 mode may be an indication of the preactivation of adsorbed methane. The breakdown of the infrared selection rule implies that methane does not retain its Td symmetry on the surface. The I R bands at 2900 and 2890 cm-1 of adsorbed CH4 on MgO permit us to conclude that the interaction between methane and the surface leads to symmetry or even structure deformation and preactivation of methane. The fact that I R spectra of adsorbed methane strongly depend on the pretreatment of MgO confirms that the chemical interaction is involved in the methane adsorption. The chemical interaction, although weak, may be as motivating a force to cleave the C-H bond at high temperatures as that for methane oxidative coupling. The activation of methane could be also interpreted in the formalism of quantum chemistry. Methane is difficult to activate because the level of its LUMO (lowest unoccupied molecular orbital) is too high and that of its HOMO (highest occupied molecular orbital) is too low, and this is due partly to its high symmetry. A strong interaction with the surface can induce methane from the high-symmetry structure, Td symmetry, to a structure with lower symmetry, e.g., C3,, Cb,or C3 symmetry as shown by the appearance of the infrared band of the u1 mode. The surface-induced structure distortion may also change the relative position of the LUMO and HOMO of CH4 and possibly make it easy to move an electron from the HOMO or to donate an electron to the LUMO as a result of methane activation or reaction. Surface sites for the activation of methane on a metal oxide are still open to discussion in the literature. An acceptable conclusion is that surface active oxygen species are responsible for the first hydrogen abstraction of CH4. However, it is not well understood if the methane is activated solely by the active oxygen species and what role is played by Lewis acid sites. Recently Choudhary et ~ 1 . 3 2 ~ 3 reported 3 that the catalytic activity and selectivity of methane oxidative coupling are dependent on both the surface acidity and basicity, and they assumed that there is a possibility of an involvement of an acid-base pair in methane oxidative coupling. From the IR spectra of coadsorbed CH4 and CO (Figures 2 and 3), part of the adsorbed methane interacts with both surface Mg2+LCand 0 2 - ~sites, c Le., a Lewis acid-base pair. When CO was introduced onto the methane-preadsorbed MgO surface, CO selectively adsorbed on a MgZ+LCsite as a result of depletion of the adsorbed methane because CO more strongly interacting with the MgZ+LCsite than methane does. Figure 2 shows that invaded CO cannot completely deplete the preadsorbed methane even after a prolonged equilibrium. For the sake of a systematic study,” the adsorbed methane that
The Journal of Physical Chemistry, Vol. 98, No. 7, 1994 1937 H H+H
I
H
.
.* I
5 . :
. 8
~..
-Mg-O-
-Mg-O-
(11)
(1111
Figure 9. Proposedinteractionmodels of CH4 with surfacesites. Species I1 (3008, 1305 cm-*)formed on surface lattice oxygen, and Species 111 (3000, 2900(2890), 1305 cm-l) formed on a Lewis acid-base pair.
interacts mainly with O2-~cis denoted as species 11, and the adsorbed methane interacting with a Lewis acid-base pair is denoted as species I11 (species I is the adsorbed methane that interacts with surface @-like species’l). Figure 9 depicts the two types of adsorbed methane on the MgO surface, where species I1 is formed via a direct interaction of a hydrogen atom with the 0 2 - ~site c alone and species I11 is formed uia the interaction of methane with both O2-~c and Mg2+Lcsites where a carbon atom or a C H u donation interacts with Mgz+Lc site and a hydrogen atom interacts with O2-LCsite. The species I1 is weakly adsorbed; therefore, its spectrum is analogous to that of free methane with two bands at -3008 and 1305 cm-l. The species I11 is strongly adsorbed, and the infrared-forbiddenmode at 29 17 cm-1 becomes infrareddetectableand shifts to 2900 and 2890 cm-I. The species I11 is activated markedly since the surface pair sites are favorable for weakening the C-H bond of methane. The two bands a t 2900 and 2890 cm-1 indicate that the species I11 can be further divided into two species whose chemical environments are slightly different, but the formation of the species with a band a t 2890 cm-1 is more selective to surface sites because this band was only detected for methane adsorption on highly outgassed MgO as seen in Figure 1. The species I1 and I11 are observed a t low temperatures that aremuch lower than that needed for the reaction of methane, reactions such as methane oxidative coupling, but it can be speculated that the two species are possible candidates for C-H bond rupture at reaction conditions. At reaction temperatures, the C-H bond is cleaved as soon as methane collisions occur on surface sites and the activated species are hardly detected on the surface. However, we believed that the initial interaction of methane with surface sites prior to its C-H bond rupture at high temperatures probably just follows the ways of species I1 and 111, especially the latter. The species I11 may be a possible precursor of methane activation because the charge separation of the C-H bond is more easily realized on the Lewis acid-base pair. In fact, MgO, when pretreated a t high temperature in vacuo, can lead to the dissociation of the C-H bond of methane even below room temperat~re.’~,’~ This is only attributed to the high reactivity of surface cus sites which are the active centers for methane activation; particularly at high temperatures, the surface cus sites are more readily derived and populated on Mg0.35336 Conclusions Two types of adsorbed methane (species I1 and 111) can be formed on well-outgassed MgO surfaces at low temperatures. Species I1 solely interacts with an 0 2 - ~ site c while species I11 does so with a Mg2+LC02-pair on the MgO surface. The species I1 is weakly adsorbed, and its I R spectrum is similar to that of free methane. Species I11 interacts strongly with surface Lewis acid-base sites and its Td symmetry is distorted as evidenced by the appearance of the infrared-forbiddenbands at 2900 and 2890 cm-1. The two types of species can be successfully distinguished by the coadsorption of CO. It is assumed that the surface cus sites are responsible for methane activation at reaction temperatures and the two adsorbed species (species I1 and 111), particularly the species 111, may be the possible candidates for precursor prior to the C-H bond rupture of methane at reaction conditions.
1938 The Journal of Physical Chemistry, Vol. 98, No. 7, 1994
Acknowledgment. We gratefully acknowledge the Natural Science Foundation of China (NSFC) for support of this research. References and Notes (1) Beekrle, J. D.; Johnson, A. D.; Ceyer, S.T. Phys. Rev. Lett. 1989, 62, 685. (2) Ceyer, S.T. Langmuir 1990, 6, 82. (3) Anderson, A. B.; Maloney, J. J. J . Phys. Chem. 1988, 92, 809. (4) Zaera, F. Catal. Lett. 1991, 1 1 , 95. (5) Mirodatos, C.; Ducarme, V.; Mozzanega, H.; Holmen, A,; SanchezMarcano, J.; Wu, Q.; Martin, G.A. In Natural Gas Conversion; Elsevier Science Publishers B. V.: Amsterdam, 1991; p 41. (6) Swang, 0.;Fagri, K., Jr.; Gropen,Q.;Waklgren, U. In Natural Gas Conversion; Elsevier Science Publishers B. V.: Amsterdam, 1991; p 191. (7) Kaliaguine, S.L.; Shelimov, B. N.; Kazansky, V. B. J . Coral. 1978, 55, 384. (8) Wang, F.; Wan, H.;Tsai, K. R.; Wang, S.;Xu, F. Caral. Lett. 1992,
12, 319. (9) Borve, K. J.; Pettersson, L. G . M. J . Phys. Chem. 1991, 95, 3214; 95, 7401. (10) Li, C.; Xin, Q.J . Chem. SOC.,Chem. Commun. 1992, 782. ( 1 1 ) Li, C.; Xin, Q.J . Phys. Chem. 1992, 96, 7714. (12) Lee, J. S.; Oyama, S.T. Catal. Rev.-Sci. Eng. 1988, 30, 249. (13) Hutchings, G.J.;Scurrell, M. S.; Woodhouse, J. R. Chem.Soc. Rev. 1989, 18, 251. (14) Brown, M. J.; Parkyns, N. D. Catal. Today 1991,8, 305. (15) Kolts, J. H.; Kimble, J. B.; Porter, R. A. In Chemical Industries; Albright, L. F., Crynes, B. L., Nowak, S.,Eds.; Marcel Dekker Inc.: New York, 1992, Vol. 46, p 1. (16) Srivastava, R. D.; Zhou, P.;Stiegel, G.J.; Rao,V. U.S.; Cinquegrane, G. Catalysis 1992, 9, 182.
Li et al. (17) Han, S.;Martenak, D. J.; Palermo, R. E.;Pearson, J. A.; Walsh, D. E. J . Catal. 1992, 136, 578. (18) Ito, T.; Tashiro, T.; Watanabe, T.; Toi, K.; Ikemoto, I. Chem. Lett. 1987, 1723. (19) Aika, K.; Lunsford, J. H. J . Phys. Chem. 1977,81, 1393. (20) Ito, T.; Wang, J.; Lin, C.; Lunsford, J. H. J . Am. Chem. Soc. 1985, 107. 5062. (21) Lunsford,J. H. InNaturalGas Conuersion;ElsevierSciencehblishers B. V.: Amsterdam, 1991; p 3. (22) Sheppard, N.; Yates, D. J. C. Proc. R. SOC.London, A 1956, 238, 69. (23) Kung, M. C.; Kung, H. H. Catal. Rev.-Sci. Eng. 1985, 27, 425. (24) Zaki, M. I.; Knozinger, H. Mater. Chem. Phys. 1987, 17, 201. (25) Bordiga, S.; Platero, E.E.; Arean, C. 0.;Lamberti, C.; Zecchina, A. J . Catal. 1992, 137, 179. (26) Li, C.; Sakata, Y.; Arai, T.; Domen, K.; Maruya, K.; Onishi, T. J . Chem. SOC.,Faraday Trans. I 1989,85, 929. (27) Coluccia, S.;Marchese, L.; Lavagnino, S.;Anpo, M. Spectrochim. Acta 1987.43A. 1573. (28) Kobayashi, M.; Kanno,T.; Konishi,Y. J. Chem.Soc.,Faraday Trans. I 1988,84, 281. (29) Hadiiivanov, K.; Klissurski, D.; Davydov, A. J. Chem.Soc., Faraday Trans. I 19&, 84, 37. (30) Ito, T.; Lunsford, J. H. Nature (London) 1985,314, 721. (31) Driscoll, D. J.; Campbell, K. D.; Lunsford, J. H. Ada Catal. 1987, 35, 139. (32) Choudhary, V. R.; Rane, V. H.; Chaudhari, S.T. Catal. Lett. 1990, 6, 95. (33) Choudhary, V. R.; Rane, V. H. J . Catal. 1991, 130,411. (34) Hamada, E.; Toi, K. Catal. Sci. Technol. 1991, 1 , 437. (35) Klabunde, K. J.; Nieves, I. J . Phys. Chem. 1988, 92, 2521. (36) Peng, X.D.; Burteau, M. A. Catal. Lett. 1992, 12, 245.
