8692
J. Phys. Chem. B 2000, 104, 8692-8703
Decomposition of Saturated Hydrocarbons Adsorbed on Ni(755): Comparison of Decomposition Starting Temperatures among Cyclic and Straight-Chain Hydrocarbons Hideo Orita,* Hiroshi Kondoh, and Hisakazu Nozoye National Institute of Materials and Chemical Research, Tsukuba, Ibaraki 305-8565, Japan ReceiVed: March 27, 2000; In Final Form: June 30, 2000
The surface chemistry of saturated hydrocarbons (cyclic and straight-chain ones between C5 and C8) adsorbed on a stepped Ni(755) {Ni(S)[6(111) × (100)]} has been investigated mainly with temperature-programmed desorption (TPD). Coadsorbed CO shows several significant effects on decomposition of hydrocarbons (e.g., promoting effect on decomposition of “low-reactivity” hydrocarbons). By using these effects of CO, we have determined desorption energy, decomposition fraction, ratio of desorption-limited H2 peak area to reactionlimited one, and decomposition starting temperature of hydrocarbons. These quantities are very dependent on molecular structure. The decomposition starting temperature of straight-chain hydrocarbons increases only slightly with increase of carbon atom numbers, suggesting that activation energy for decomposition is similar for all the straight-chain hydrocarbons and only part of the chain reorients regardless of chain length by rotation about C-C bond in a transition state for decomposition. Cyclohexane shows a much higher decomposition starting temperature and lower decomposition fraction than other cyclic hydrocarbons, which can be explained on the basis of the difference in conformational energy to attain eclipsed C-H bonds on adjacent carbons in a transition state. Relating decomposition starting temperature with desorption energy, we have discussed the decomposition mechanism of hydrocarbons and proposed some candidates for transition state of decomposition.
1. Introduction A fundamental study of the interactions of hydrocarbons with metal surfaces is not only interesting but also important because it may provide insight into the mechanism of hydrocarbon decomposition reaction (e.g., dehydrogenation, hydrogenolysis) in practical catalytic systems.1 Several recent investigations revealing the interactions of saturated hydrocarbons with singlecrystal surfaces under ultrahigh vacuum condition have been carried out by using many experimental techniques such as temperature-programmed desorption (TPD),2-6 high-resolution electron energy loss spectroscopy (HREELS),7,8 and infrared reflection absorption spectroscopy (IRAS).9,10 The structure and reactivity of adsorbed hydrocarbons have been studied in detail, and some reactivity adsorbed intermediates have now been wellcharacterized in some cases. Among these studies, Campbell et al. proposed the utility of bismuth postdosing thermal desorption spectroscopy (BPTDS), which used the characteristics of bismuth adatoms (e.g., strong site-blocking effect without pronounced electronic effect).4 They identified adsorbed benzene as an intermediate in the dehydrogenation of cyclohexane on Pt(111). The subsequent works8,11,12 of Campbell et al. displayed the further application of BPTDS, combined with HREELS, to other hydrocarbons. They reported that cyclohexane decomposition began at ca. 195 K and was completed by ca. 230 K. They also identified π-allyl adsorbed c-C6H9 as a dominant intermediate in the early stage of decomposition of cyclohexane and cyclohexene, and found that the c-C6H9 species converted to adsorbed benzene at ca. 290-340 K. Recently, we have investigated coadsorption of saturated hydrocarbons and CO on Ni(755), which is denoted [6(111) × * To whom correspondence should be addressed. Tel: +81-298-61-4527. Fax: +81-298-61-4504. E-mail:
[email protected].
(100)] in step notation, changing the coverages of adsorbates carefully. For the purpose of understanding, we briefly outline the behavior of hydrocarbons/CO on Ni(755) reported in our previous studies.13,14 We have found a novel promoting effect of coadsorbed CO on decomposition of “low-reactivity” hydrocarbons, which decompose little during TPD run (e.g., cyclohexane, n-pentane).13 For “high-reactivity” hydrocarbons, which decompose completely at their low coverages without desorbing parent molecules during a TPD run, coadsorbed CO at a coverage below ca. 0.2 monolayer (ML) does not suppress the high decomposition probability of the hydrocarbons.14 When CO coverage increases, coadsorbed CO displaces a hydrocarbon from the first layer on the surface to the second layer on CO and the decomposition of low-reactivity as well as highreactivity hydrocarbons is suppressed. We call this the inhibiting effect of coadsorbed CO. We have also shown that the coverage and decomposition fraction (ratio of decomposed molecules to adsorbed ones) can be quantitatively determined by using the effect of coadsorbed CO. We have also presented a method to determine the decomposition starting temperature (the temperature where a hydrocarbon starts to decompose) of a high-reactivity hydrocarbon by making use of the inhibiting effect of coadsorbed CO.14 In short, the method consists of the following procedures. The surface is predosed with a hydrocarbon at ca. 120 K, subsequently flashed to a predetermined temperature, quenched back to the adsorption temperature, and postdosed with a saturated amount of CO. After this pretreatment, TPD is measured for H2, CO, and the hydrocarbon. The desorbed amount of H2 corresponds to the decomposed amount of the hydrocarbon. When the preflash temperature is not high enough, the hydrocarbon is not decomposed and no decomposition intermediate is produced. The hydrocarbon that is not decomposed on the
10.1021/jp001137e CCC: $19.00 © 2000 American Chemical Society Published on Web 08/19/2000
Decomposition of Hydrocarbons Adsorbed on Ni(755) surface is displaced from the first layer to the physisorbed second layer by the postdosed CO and decomposition of the hydrocarbon is inhibited. On the other hand, when some intermediates are produced in the preflash step, they bind to the surface strongly and are not displaced by the postdosed CO. Decomposition of intermediates is not inhibited by the postdosed CO, and the amount of produced hydrogen increases as the amount of intermediates does in the preflash step. Therefore, by changing the preflash temperature, we can determine the temperature above which decomposition of a hydrocarbon becomes noticeable. We have defined the term “decomposition starting temperature” as the preflash temperature where the hydrogen peak area begins to increase. In the present work, we have applied this method to various adsorbed hydrocarbons (i.e., cyclic and straight-chain hydrocarbons between C5 and C8). Even if a hydrocarbon does not decompose at all without coadsorbed CO, we can determine its decomposition starting temperature by using the promoting effect of CO on decomposition and modifying the method reported previously. We have obtained a relationship between decomposition starting temperature of the hydrocarbons and their structures and desorption energy and have discussed their decomposition mechanism. 2. Experimental Section All the experiments were conducted in a stainless steel ultrahigh vacuum system equipped with a single-pass CMA for Auger electron spectroscopy (AES), four-grid low-energy electron diffraction (LEED) optics, a quadrupole mass spectrometer for TPD, and an ion gun for cleaning. The base pressure was less than 1 × 10-10 Torr (1 Torr ) 133.3 Pa). A diskshaped Ni(755) crystal (ca. φ 8 × 1 mm) was heated resistively and could be cooled below 90 K. The sample temperature was measured with a chromel-alumel thermocouple spot-welded to the edge of the crystal. The Ni(755) crystal was cleaned by Ar ion sputtering followed by annealing to 1080 K. The cleanliness and ordering of the surface were checked by AES and LEED. Adsorption of gases on the surface was performed using a gas doser that was composed of a glass capillary array. The exposure was controlled by varying dose time and the backpressure of the doser (usually 1 × 10-4 Torr, which is not corrected for ion gauge sensitivity to various gases). TPD experiments were carried out with linear heating of 10 K/s controlled by a personal computer. Four masses were monitored simultaneously in a single experiment, and the data were stored in the computer. The coverage of adsorbed CO relative to the exposed surface nickel atoms was determined from the integrated mass intensity of the TPD peak, assuming that the saturation coverage around 300 K is 0.5 ML (1 ML ) one molecule or atom per exposed surface nickel atom).13 Under the present experimental conditions, the saturation coverage of CO around 300 K (i.e., 0.5 ML) was accomplished with a dose time of 50 s. We can determine the mass sensitivity factor for hydrogen from TPD for H2 adsorption, assuming the saturation coverage of hydrogen atom is 1 ML. We can calibrate also the mass sensitivity factor for hydrocarbons, assuming the mass balance between hydrogen and hydrocarbons (no other products than hydrogen were observed and AES showed that carbon remained on the surface after TPD measurements). Thus, the amounts of desorbed and decomposed hydrocarbons are calculated by using these sensitivity factors. Although all the TPD spectra for hydrogen were not corrected for the hydrogen adsorption from
J. Phys. Chem. B, Vol. 104, No. 36, 2000 8693 residual gas, the amount of hydrogen adsorbed from residual gas was usually much smaller than that from decomposition of hydrocarbons and could be subtracted easily from the total amount of desorbed hydrogen except when mentioned in the text. All the following hydrocarbons used in the experiments were purchased and purified by at least five freeze-pump-thaw cycles: cyclopentane, methylcyclopentane, cyclohexane, cyclohexane-d12, cycloheptane, cyclooctane, n-pentane, n-hexane, n-heptane, and n-octane. The hydrogen and CO were the research grade purity and used without further purification. The purity of all the gases was verified by mass spectroscopy after admission into the ultrahigh vacuum system. 3. Results and Discussion We have investigated 10 various hydrocarbons (cyclopentane, methylcyclopentane, cyclohexane, cyclohexane-d12, cycloheptane, cyclooctane, n-pentane, n-hexane, n-heptane, and n-octane). Among these hydrocarbons, cyclohexane and n-pentane are classified as “low-reactivity” hydrocarbons, and the other hydrocarbons are classified as “high-reactivity” hydrocarbons. In this report, results for cyclohexane and n-pentane are presented below in detail as representative cases of lowreactivity cyclic and straight-chain hydrocarbons, respectively. Results for high-reactivity cyclic and straight-chain hydrocarbons (adsorption, decomposition, effects of coadsorbed CO, and so on) are similar to those of cycloheptane, which were reported before.14 We do not describe the detailed procedures for these high-reactivity hydrocarbons but only show the results as desorption temperatures etc. in Tables 1 and 2. 3.1. Cyclic Hydrocarbons. 3.1.1. Cyclohexane (C6H12 and C6D12). TPD results of cyclohexane on clean Ni(755) show three peaks at 137, 194, and 219 K, depending on cyclohexane coverage (cf. Table 1).13 The peak at 137 K is assigned to desorption from the multilayer and the other two peaks are due to monolayer adsorption (we use the word “monolayer” to indicate the first adsorbed layer, which is in direct contact with the surface, and designate these peaks as monolayer desorption). Zebisch et al. have investigated the adsorption and desorption of cyclohexane on Ni(111).15 Cyclohexane desorbs from Ni(111) without dissociation, and TPD spectra of cyclohexane show three peaks around 140, 145, and 190 K that are assigned to desorption from the multilayer, the second layer, and the first layer, respectively. The peak position of the first layer on Ni(111) is in good agreement with that of the lower temperature peak of monolayer desorption from Ni(755), so this peak is assigned to the desorption from the (111) terrace. For Ni(111), no desorption of cyclohexane above 200 K is observed. Therefore, the desorption peak from Ni(755) observed above 200 K is due to the desorption of cyclohexane, which interacts with the step sites of Ni(755). This peak above 200 K is saturated at a coverage of 0.03 ( 0.01 ML, and all of the monolayer desorption peaks are saturated at a coverage of 0.08 ( 0.03 ML. We carried out LEED measurements after adsorption of cyclohexane on Ni(755), but we did not observe any additional spots due to an ordered overlayer structure. For Ni(111), however, the monolayer cyclohexane exhibits a well-ordered (x7 × x7)R19.1° LEED structure for a coverage range from 0.04 ML up to the saturation coverage of 0.143 ML, indicative of island formation.15 The difference in saturation coverage of monolayer between Ni(111) and Ni(755) probably comes from the small terrace width of Ni(755). The saturation coverage of a possible ordered adsorption model on Ni(755) (the unit cell
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TABLE 1: Comparison of TPD Peak Temperature, Desorption Energy, Initial Decomposition Fraction, Ratio of Desorption-Limited H2 Peak Area to Reaction-Limited One, Decomposition Starting Temperature, and Activation Energy for Decomposition among Cyclic Hydrocarbons Adsorbed on Ni(755) moleculea c-C5 Me-c-C5 c-C6 c-C6D12 c-C7 c-C8
TPD peak temp (K)
desorp energyb (kJ/mol)
170 (mono)e 141 (multi)f 191 (mono)e 137 (multi)f 194 (mono)e 219 (mono)g 138 (multi)f 192 (mono)e 214 (mono)g 153 (multi)f 207 (mono)e 163 (multi)f 226 (mono)e
41.5 34.3 46.8 33.3 47.6 53.9 33.5 47.1 52.6 37.3 50.9 39.7 55.7
initial decomp fraction
peak ratioc (X)/(Y)
decomp starting temp (K)
activation energyd (kJ/mol)
1.0h,i 1.0h,i
5/5 5/7
128 148
36 42
0.10h 0.96j
5/7
187k
52
0.0h 0.89j
5/7
206k
58
1.0h,i
7/7
162
45
1.0h,i
8/8
177
50
a c-C , cyclopentane; Me-c-C , methylcyclopentane; c-C , cyclohexane; c-C D , cyclohexane-d ; c-C , cycloheptane; c-C , cyclooctane. b Estimated 5 5 6 6 12 12 7 8 from the TPD peak temperature in column 2 by using Redhead’s method with a preexponential factor of 1013 s-1; the surface heating rate is 10 K/s. c Ratio of desorption-limited H2 peak area (X) to reaction-limited one (Y) (the sum of X and Y equals the total number of hydrogen atoms in a hydrocarbon, and error bars for the value of X and Y are at least (1). d Estimated from decomposition starting temperature with a preexponential factor of 1013 s-1; see section 3.1.1 for a discussion of the method of calculation. Estimation of activation energy is very dependent on the value of the preexponential factor used. A difference in preexponential factor of one order of magnitude leads to an uncertainty of (4 kJ/mol in activation energy. e Peak temperature of monolayer from the terrace at its saturation coverage. f Peak temperature of multilayer for the longest dose time examined. g Peak temperature of monolayer due to the molecule which interacts with the step at its saturation coverage. h Initial decomposition fraction without CO. i Increase of decomposition fraction was not observed in the presence of coadsorbed CO. j Initial decomposition fraction with maximum CO promotion. k Obtained by the extrapolation method for coadsorption with CO.
TABLE 2: Comparison of TPD Peak Temperature, Desorption Energy, Initial Decomposition Fraction, Ratio of Desorption-Limited H2 Peak Area to Reaction-Limited One, Decomposition Starting Temperature, and Activation Energy for Decomposition among Straight-Chain Hydrocarbons Adsorbed on Ni(755) moleculea n-C5 n-C6 n-C7 n-C8
TPD peak temp (K) 185 (mono)e 217 (mono)f 139 (multi)g 211 (mono)e 154 (multi)g 230 (mono)e 170 (multi)g 241 (mono)e
desorp energyb (kJ/mol) 45.3 53.4 33.8 51.9 37.5 56.7 41.5 59.5
peak ratioc (X)/(Y)
decomp starting temp (K)
activation energyd (kJ/mol)
0.50h 1.0i 1.0h,j
4/8
184l
52
5/9
187
52
1.0h,k
6/10
1.0h,k
7/11
197 190l 195
55 53 55
initial decomp fraction
a n-C , n-pentane; n-C , n-hexane; n-C , n-heptane; n-C , n-octane. b Estimated from the TPD peak temperature in column 2 by using Redhead’s 5 6 7 8 method with a preexponential factor of 1013 s-1; the surface heating rate is 10 K/s. c Ratio of desorption-limited H2 peak area (X) to reaction-limited one (Y) (the sum of X and Y equals the total number of hydrogen atoms in a hydrocarbon, and error bars for the value of X and Y are at least (1). d Estimated from decomposition starting temperature with a preexponential factor of 1013 s-1; see section 3.1.1 for a discussion of the method of calculation. Estimation of activation energy is very dependent on the value of the preexponential factor used. A difference in preexponential factor of one order of magnitude leads to uncertainty of (4 kJ/mol in activation energy. e Peak temperature of monolayer from the terrace at its saturation coverage. f Peak temperature of monolayer due to the molecule which interacts with the step at its saturation coverage. g Peak temperature of multilayer for the longest dose time examined. h Initial decomposition fraction without CO. i Initial decomposition fraction with maximum CO promotion. j Complete decomposition of n-hexane only occurred at a very small coverage, below 0.004 ML. The promoting effect of CO was clearly observed for n-hexane when the decomposition fraction became smaller than 0.8 at larger coverage of n-hexane. k Increase of decomposition fraction was not observed in the presence of CO. l Obtained by the extrapolation method for coadsorption with CO.
is p(3 × 3) for the (111) terrace and cyclohexane occupies the top site with the C-C bonds parallel to the 〈211〉 directions) is calculated to be about 0.11 ML. This value is a little overestimated, because half of the cyclohexane molecules in this model touch the step. TPD spectra were measured after dosing clean Ni(755) at 118 K to a fixed amount of cyclohexane followed by exposure to various amounts of CO. We used a mass peak of m/e ) 56 to monitor cyclohexane, since it has the highest intensity in the cracking pattern of cyclohexane, and we have checked that its behavior is similar to that of the parent peak (m/e ) 84). Figure 1 shows the typical dependence of the TPD peak areas of CO, hydrogen, and cyclohexane on CO dose time. The peak area of CO increases almost linearly with dose time, indicating a mobile precursor mediated adsorption. The amount of adsorbed CO is almost the same as that for a clean surface, and the reverse
dosing sequence (i.e., CO dose followed by cyclohexane exposure) gives almost the same results as in Figure 1, which indicates that preadsorbed cyclohexane little affects the sticking probability of CO and vice versa. Therefore, it is appropriate that a CO dose time of 50 s corresponds to 0.5 ML in Figure 1. On increasing CO dose time, the peak area of hydrogen becomes larger and reaches its maximum value at about a dose time of 20 s. On further increasing dose time, the peak area of hydrogen decreases considerably. The behavior of cyclohexane is opposite to that of hydrogen. These results clearly indicate that coadsorbed CO promotes the decomposition of cyclohexane.13 The promotion effect of CO during TPD shows maximum efficiency around a CO coverage of 0.2 ML. The decomposition fraction (ratio of decomposed molecules to adsorbed ones) at zero CO dose time and maximum decomposition fraction in Figure 1 are 0.08 and 0.58, respectively.
Decomposition of Hydrocarbons Adsorbed on Ni(755)
J. Phys. Chem. B, Vol. 104, No. 36, 2000 8695
Figure 1. Dependence of the TPD peak area of CO, hydrogen, and cyclohexane (m/e ) 56) on the postdose time of CO. Cyclohexane was predosed onto clean Ni(755) at 118 K for 45 s. The coverage of cyclohexane was determined to be 0.02 ML from the mass balance between hydrogen and cyclohexane (see the text).
It should be noted that the amount of produced hydrogen with the saturation coverage of coadsorbed CO is still larger than that without CO (cf. Figure 1). Therefore, we need to modify the previous method14 for high-reactivity hydrocarbons in order to determine the decomposition starting temperature of cyclohexane. TPD spectra were measured after preflash, quenching, and postdosing of CO on Ni(755) with preadsorbed cyclohexane and predosing of CO. Typical spectra for hydrogen and CO are shown in parts a and b of Figure 2, respectively. The dose time of cyclohexane is 45 s (its coverage is 0.02 ML), and those of CO before and after preflash are 10 and 50 s, respectively. Figure 3 displays the variations in TPD peak area of hydrogen and CO with preflash temperature. (The preflash temperature indicates the maximum temperature achieved at the preflash step; the surface heating is 10 K/s). After switching off the heating current of the sample, the temperature stays around the preflash temperature (within 5 K) for about 5 s and then goes down to the adsorption temperature.) In the preflash temperature range between 150 and 175 K, TPD spectra for hydrogen and CO as well as the peak area remain almost unchanged. The TPD spectrum for CO after preflash to 152 K is quite similar to that for CO adsorbed on Ni(111) reported before by others.16,17 At a preflash temperature around 180 K, the hydrogen peak area begins to increase while the lower temperature desorption feature of CO below 350 K decreases in intensity and the total peak area of CO starts to decrease. By preflash to 200 K, the hydrogen and CO peak areas reach constant values. The small lower temperature desorption peak of CO around 315 K is clearly observed (Figure 2b). The results in Figures 2 and 3 show that (1) at ca. 180 K the cyclohexane coadsorbed with CO begins to decompose to hydrogen atoms and some dehydrogenated intermediates on Ni(755); (2) these intermediates reduce the amount of postadsorbed CO, and a new lower temperature desorption peak of CO appears by interaction between the intermediates and postdosed CO; and (3) decomposition of the coadsorbed cyclohexane to the intermediates is completed already at ca. 200 K. The starting edge of the increase of hydrogen peak area is sharp and reflects the onset of the first step of the decomposition of cyclohexane. Thus, we use the temperature where the hydrogen peak area begins to increase as decomposition starting temperature for the CO and cyclohexane coadsorption system. In Figure 3, the decomposition starting temperature of the cyclohexane coadsorbed with CO at a predose time of 10 s is determined to be 176 K. The
Figure 2. Variation of TPD spectra for (a) hydrogen (m/e ) 2) and (b) CO (m/e ) 28) with preflash temperature. Cyclohexane was predosed for 45 s (the coverage is 0.02 ML) at 118 K followed by exposure to CO for 10 s (0.1 ML). The surface was flashed at 10 K/s to a preflash temperature. After recooling to 118 K, CO was postdosed to the surface for 50 s and then the TPD spectrum was measured.
Figure 3. Variations in TPD peak area of hydrogen and CO with preflash temperature. The experimental procedure is the same as that in Figure 2.
experimental error of the decomposition starting temperature is below (5 K. Similar experiments were repeated with changing CO preand postdose time (the sum of CO pre- and postdose time was fixed to 60 s). Figure 4 shows the dependence of decomposition starting temperature of the coadsorbed cyclohexane on CO predose time. The decomposition starting temperature decreases almost linearly with increasing CO predose time and can be extrapolated to zero CO predose time (we call this method “the extrapolation method”). This extrapolated value is determined to be 187 ( 5 K.
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Figure 4. Dependence of decomposition starting temperature of the cyclohexane coadsorbed with CO on CO predose time. The decomposition starting temperature was measured as a function of pre- and postdose time of CO as shown in Figure 3. The experimental procedure is the same as that in Figure 3, and only the CO dose time has changed. The sum of CO pre- and postdose time is fixed to 60 s.
The decomposition starting temperature thus determined (187 K) is comparable to that of cyclohexane on Pt(111) (195 K), which is determined from experiments of BPTDS and HREELS by Campbell et al.,8 but it is much higher than that of cycloheptane on Ni(755) (162 K)14 (cf. Table 1). The desorption of cyclohexane starts from ca. 185 K, and maximum desorption peak is observed at 218 K when only cyclohexane is adsorbed on clean Ni(755) at the same dose time in Figures 2 and 3. Therefore, the decomposition of cyclohexane to dehydrogenated intermediates occurs simultaneously as the desorption of molecular cyclohexane does. This fact corresponds well with a low initial decomposition fraction (10%) of cyclohexane.13 The decrease of decomposition starting temperature with increase of coadsorbed CO coverage correlates well with the promoting effect of CO. In the presence of a large amount of coadsorbed CO, most of the cyclohexane is displaced to an overlayer of adsorbed CO and cannot interact directly with the nickel surface.13 Then, the amount of produced hydrogen decreases (see Figure 1) while CO is reducing the decomposition starting temperature of the coadsorbed cyclohexane, which still interacts with the surface. We have also investigated the decomposition property of cyclohexane-d12 (c-C6D12). The promoting effect of CO on decomposition of cyclohexane-d12 is observed similarly as for cyclohexane (cf. Figure 1). As the results tabulated in Table 1 show, cyclohexane-d12 has the same desorption temperature as cyclohexane, whereas the deuterated molecule shows lower initial decomposition fraction (0%) and higher decomposition starting temperature (206 K) than the unlabeled one. These facts demonstrate that the first step of decomposition involves breaking of the carbon-hydrogen bond and that the decomposition starting temperature correlates with the activation energy for decomposition. It is worth noting that by using the extrapolation method we can determine the decomposition starting temperature of cyclohexane-d12 that does not decompose at all without coadsorbed CO. The rate of dehydrogenation can be expressed in Arrhenius form as
rate ) -dθ/dt ) ν exp(-Ea/RT)θ
(1)
where θ is the coverage of cyclohexane, ν is the preexponential factor, Ea is the activation energy, and R is the gas constant. We have evaluated the amount of cyclohexane that has undergone decomposition for any given preflash temperature by the numerical integral of eq 1 between the adsorption
temperature (118 K for cyclohexane) and the preflash temperature with a given Ea and a given preexponential factor. Assuming that 5% of cyclohexane decomposes when the substrate temperature is increased from 118 K to the measured decomposition starting temperature, we can estimate the activation energy for a given preexponential factor. The activation energies of cyclohexane and cyclohexane-d12 thus estimated are 45 and 49.5 kJ/mol for a preexponential factor of 1011 s-1, 52 and 58 kJ/mol for 1013 s-1, and 59.7 and 65.8 kJ/mol for 1015 s-1, respectively. Although the activation energy depends on the preexponential factor, the difference in the activation energy between cyclohexanes does not depend on it so much and is always about 5 kJ/mol. Taking the error bar of (5 K in the measured decomposition starting temperature into account, the difference in the activation energy between cyclohexanes is determined to be 5 ( 2 kJ/mol and compares well with the difference in the zero-point energies for C-H and C-D bonds (about 5 kJ/mol). Campbell et al. have estimated the activation energy for the dehydrogenation of cyclohexane adsorbed on Pt(111) by simulating the BPTDS data.8 Assuming a preexponential factor of 7 × 1012 s-1 based on their prior results,5 they have obtained an activation energy of 56 kJ/mol for dehydrogenation of cyclohexane. The estimated activation energy for decomposition of cyclohexane on Ni(755) is comparable to that on Pt(111). If we treat the decomposition starting temperature as a parameter at the estimation of activation energy, the relationship between the activation energy and the decomposition starting temperature can be fitted to the following equation, Ea ) 0.06 + 0.28T kJ/mol (where T is the decomposition starting temperature) when we use a common preexponential factor of 1013 s-1. The equation becomes Ea ) 0.05 + 0.24T and Ea ) -0.17 + 0.32T kJ/mol for the preexponential factor of 1011 and 1015 s-1, respectively. As there is a linear relationship between the activation energy and decomposition starting temperature and we have no information of preexponential factor for decomposition of hydrocarbons at present, we use the decomposition starting temperature as an index of activation energy in further discussion of this work. The peak area of hydrogen after preflash to ca. 200 K in Figure 2a can be divided approximately at about 380 K into two regions as discussed in detail for cycloheptane.14 In the present case, the separation of peaks is not very good, but the lower temperature peak and the higher one correspond to desorption-limited and reaction-limited peaks, respectively, because postdosing of D2 decreases only the lower temperature peak by isotope exchange reaction between H and D atoms. The ratio of desorption-limited peak area (X) to reaction-limited one (Y) is estimated to be X/Y ) 5/7 (X + Y ) 12; i.e., the sum of X and Y is normalized to the total number of hydrogen atoms in cyclohexane, and error bars for the value of X and Y are at least (1). According to this ratio, the following dehydrogenation pathway can be tentatively considered: C6H12(a) f C6H7(a) + 5H(a) below ca. 200 K, 5H(a) f 5/2H2(gas) between ca. 300 and 380 K, and C6H7(a) f 6C(a) + 7/2H2(gas) between ca. 380 and 523 K. Although we have no information about the structure of intermediates, π-cyclohexadienyl-adsorbed c-C6H7 species is the most probable candidate from the stoichiometry. This intermediate has been detected in the dehydrogenation of 1,3cyclohexadiene on Pt(111) with IRAS by Manner et al.18 In the higher temperature region in Figure 2a, at least three peaks were observed at ca. 390, 425, and 480 K. These peaks probably come from the sequential dehydrogenation of intermediates. The clear assignment of the peaks is difficult because these peaks
Decomposition of Hydrocarbons Adsorbed on Ni(755) are separated only after preflash to higher temperature and postdosing of saturated amount of CO. As the desorption of CO also occurs around 400 K, the adsorbed structure starts to be changed greatly in the temperature region of the reactionlimited peaks. This makes the analysis of the reaction-limited peaks difficult, too. Jiang et al. have studied the effects of potassium on adsorption and the reaction of methylcyclohexane on Pt(111) and also measured the ratio of desorption-limited H2 peak area to reaction-limited one.18 From the peak ratio, they propose a dehydrogenation pathway of methylcyclohexane over K-promoted Pt(111) as follows: C7H14(a) f C6H6(a) (adsorbed benzene) + CH2(a) + 6H(a) below 230 K, C6H6(a) f C6H6(gas) at 230 K, 6H(a) f 3H2(gas) at 450 K, and CH2(a) f C(a) + H2(gas) between 500 and 700 K. For their case, K adatoms promote the desorption of benzene as a product, but we did not observe any desorption of dehydrogenated species in the present work, probably because of the high activity of the stepped surface. 3.1.2. Other Cyclic Hydrocarbons. Other cyclic hydrocarbons studied here (cyclopentane, methylcyclopentane, cycloheptane, and cyclooctane) are classified as high-reactivity hydrocarbons, whose initial decomposition fraction (i.e., decomposition fraction extrapolated to zero coverage of a hydrocarbon) is unity and their behavior of desorption and decomposition was similar to that of cycloheptane14 (no products other than hydrogen were detected in all the TPD measurements). The results (desorption peak temperature, desorption energy (Ed), initial decomposition fraction, ratio of desorption-limited H2 peak area to reactionlimited one, decomposition starting temperature, and activation energy for decomposition) are tabulated in Table 1. Desorption energies of hydrocarbons are estimated approximately by assuming that all the peaks at longer dose time obey first-order desorption kinetics20 with a preexponential factor of 1013 s-1 (peaks of a high-reactivity hydrocarbons such as cycloheptane14 do not obey first-order desorption kinetics very strictly, but their peak temperatures reach almost constant values at longer dose time). The values of desorption energy are probably only accurate to (8 kJ/mol, since the actual prefactor could be different by (2 orders of magnitude. However, relative desorption energies are accurate to within 4 kJ/mol because the desorption processes of cyclic hydrocarbons are likely to have similar preexponential factors. It is worth noting that high-reactivity hydrocarbons show only one peak of monolayer desorption. The desorption energies of highreactivity hydrocarbons and the desorption energy of cyclohexane from the (111) terrace have good correlation with the number of CH2 units in cyclic hydrocarbons (monotonic increase by 4.6 kJ/mol per CH2 unit; Ed ) 19.1 + 4.6n kJ/mol for c-Cn), but the desorption energy of cyclohexane from the step does not (two monolayer desorption peaks of cyclohexane are different in desorption energy by ca. 6 kJ/mol). Therefore, monolayer desorption peaks of high-reactivity hydrocarbons are assigned to the desorption from the (111) terrace. High-reactivity hydrocarbons do not show any desorption peak of the molecule that interacts with the step sites. This fact is consistent with the complete decomposition of these hydrocarbons at their lower coverage. The ratio of desorption-limited H2 peak area (X) to the reaction-limited one (Y) in TPD spectra for hydrogen nearly equals one (i.e., X/Y ≈ 1; the sum of X and Y is normalized to the total number of hydrogen atoms in a hydrocarbon), indicating that each carbon atom of cyclic hydrocarbons releases one hydrogen atom on average when decomposition to dehydro-
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a
b
Figure 5. TPD spectra for (a) n-pentane (m/e ) 43) and (b) hydrogen (m/e ) 2) after adsorption of n-pentane on clean Ni(755) at 118 K as a function of dose time of n-pentane.
generated intermediates is completed. Although the error bars for X and Y are (1 or 2 due to the poor separation of the peaks, dehydrogenated intermediate species which have stoichiometries similar to those in Table 1 are reported for Pt(111).18,21,22 For the dehydrogenation of the cyclic C5 and C7 hydrocarbons, stable planar intermediates have been identified by IRAS to be η5cyclopentadienyl (C5H5)18 and η7-cycloheptatrienyl (C7H7) species,21 respectively. The dehydrogenation of cyclic C8 hydrocarbons proceeds through sequential C-H bond cleavage steps to transient formation of cyclooctatetraene (C8H8), which is converted via a ring-closure process to the stable bicyclic ring species of stoichiometry C8H6.22 Decomposition starting temperature is dependent on structure of cyclic hydrocarbons (comparison of decomposition starting temperature among cyclic and straight-chain hydrocarbons will be discussed later). For high-reactivity hydrocarbons, their decomposition starting temperature is much lower than their TPD peak temperature of monolayer desorption. The last column of Table 1 shows the estimated value of activation energy for decomposition that is calculated from decomposition starting temperature assuming a common preexponential factor of 1013 s-1 (see section 3.1.1 for a discussion of the method of calculation). The absolute values of activation energy are probably only accurate to (8 kJ/mol since the actual prefactor could be different by (2 orders of magnitude. 3.2. Straight-Chain Hydrocarbons. 3.2.1. n-Pentane (C5H12). A series of TPD spectra for n-pentane adsorbed on Ni(755) is shown in Figure 5. Parts a and b of Figure 5 display the evolution of undecomposed molecular n-pentane and hydrogen for various doses of pentane, respectively. In Figure 5a, we used a mass peak of m/e ) 43 to monitor n-pentane since it has the highest intensity in the cracking pattern of n-pentane and we
8698 J. Phys. Chem. B, Vol. 104, No. 36, 2000
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a
Figure 6. Variations in TPD peak area of hydrogen and n-pentane (m/e ) 43 and 72) with dose time of n-pentane.
checked that its behavior is similar to that of the parent peak (m/e ) 72) (cf. Figure 6). No other products than hydrogen and molecular n-pentane are detected by TPD in all the experiments. The desorption peak of pentane is observed at 218 K at the lowest dose time. This peak grows in intensity and shifts slightly to 210 K by further increase of dose time and starts to saturate at a dose time of 60 s. Another peak below 200 K appears from a dose time of 60 s and grows in intensity. This peak also begins to saturate at a dose time of 150 s (the spectrum is the same as that at a dose time of 180 s) and is located at 185 K at its saturation. These two states can be assigned to desorption of molecularly adsorbed n-pentane on Ni(755) (monolayer desorption) as for cyclohexane described above: the lower temperature peak is due to the desorption of n-pentane from the (111) terrace and the higher one comes from the n-pentane which interacts with step sites of Ni(755). Even for higher dose time (>150 s), any sharp peak due to multilayer adsorption of n-pentane does not appear below 150 K, whereas the desorption from multilayer is observed clearly for cyclohexane13 and cycloheptane.14 This is probably because the adsorption temperature (118 K) used is too high to fill the multilayer adsorption state of n-pentane. The TPD spectra for hydrogen evolution (Figure 5b) are characterized by the presence of a main peak between 380 and 390 K and a shoulder around 450 K, but these peaks are not separated enough to accurately estimate the ratio of two peaks. These spectra are not corrected for the H2 adsorption from residual gas. The amount of hydrogen observed at zero dose time is not negligible, about one-third of that produced at a dose time of 15 s (cf. Figure 6). Figure 6 shows variations in the peak areas of hydrogen (m/e ) 2) and n-pentane (m/e ) 43 and 72) with dose time of n-pentane. The fragment and parent peak areas of desorbed n-pentane at m/e ) 43 and 72 show the same dependence on dose time, although the peak area of m/e ) 72 is quite smaller than that of m/e ) 43. The peak area of desorbed n-pentane is saturated at a dose time of 150 s. As to hydrogen, its peak area increases abruptly at first and then gradually with dose time. The gradual increase of the peak area for dose time >60 s mainly comes from the residual H2 adsorption. These results indicate that decomposition of n-pentane does not occur efficiently. The effect of coadsortion of CO on preadsorbed n-pentane was studied by TPD. TPD spectra were measured after dosing a fixed amount of n-pentane (60 s dose time) to clean Ni(755) at 118 K followed by exposing to various amounts of CO. In Figure 7, the variations in TPD spectra for n-pentane, hydrogen, and CO with CO dose time are indicated. Figure 7a clearly
b
c
Figure 7. Variations in TPD spectra for (a) n-pentane (m/e ) 43), (b) hydrogen, and (c) CO with dose time of CO. n-Pentane was predosed for 60 s onto clean Ni(755) at 118 K.
shows a drastic change in TPD spectra for n-pentane as follows: (1) the higher temperature desorption state at 210 K disappears first in the presence of coadsorbed CO (this state comes from the n-pentane which interacts with the step sites, but cannot be decomposed without CO); (2) the peak shifts to lower temperature with broadening; (3) a sharp peak appears below 140 K and grows in intensity at CO dose time between 30 and 40 s; and (4) this sharp peak shifts down to 130 K with a decrease in intensity on further increase of CO dose time. As mentioned above (cf. Figures 5 and 6), the adsorption of n-pentane clearly saturates at a dose time of 150 s and the multilayer adsorption state cannot be produced under the present experimental conditions. The sharp peak of n-pentane in Figure 7a might come from the n-pentane affected strongly by coadsorbed CO. This peak disappears in the presence of a larger amount of coadsorbed CO, indicating that preadsorbed n-pentane is expelled from the surface during the further dose of CO. As
Decomposition of Hydrocarbons Adsorbed on Ni(755)
Figure 8. Dependence of the TPD peak area of hydrogen, CO, and n-pentane (m/e ) 43) on CO dose time. n-Pentane was predosed for 60 s. The coverage of n-pentane was estimated to be 0.022 ML.
to hydrogen, Figure 7b clearly shows that the peak intensity increases at first and then decreases. The position of the main peak is shifted only slightly from 384 to 378 K in the presence of coadsorbed CO. The peak position of CO is shifted to lower temperature with increase of CO dose time (Figure 7c), as observed for Ni(111).17 Figure 8 shows the dependence of the TPD peak areas of hydrogen, n-pentane, and CO on CO dose time. The peak area of CO in Figure 8 increases linearly with CO dose time, as in Figure 1. A CO dose time of 50 s corresponds to 0.5 ML. On increasing CO dose time up to 20 s, the peak area of hydrogen becomes larger and reaches its maximum value at a CO dose time of 20 s (0.2 ML). By further increase of CO dose time, the peak area of hydrogen decreases considerably. The first increase of peak area clearly indicates that coadsorbed CO promotes the decomposition of n-pentane. As seen in Figure 8, the behavior of the TPD peak area of n-pentane is a little complicated. Its peak area decreases at first, reaches its minimum value at about a dose time of 15 s, increases at dose time between 15 and 30 s, and then decreases again on further increase of dose time. Since the longer dose time of CO expels the preadsorbed n-pentane considerably, the mass balance between hydrogen and n-pentane is not preserved and the amount of adsorbed n-pentane cannot be estimated accurately. However, some estimation can be done by using only the results in the smaller dose time between 0 and 20 s because the mass balance in this region is preserved well. The amount of preadsorbed n-pentane in Figure 8 is estimated to be 0.022 ML. The decomposition fraction at zero CO dose time and maximum decomposition fraction thus estimated are 0.18 and 0.64, respectively. The saturation coverage of n-pentane on Ni(755) in Figure 6 is estimated to be 0.055 ML by the mass sensitivity of n-pentane determined from Figure 8. By changing the amount of adsorbed n-pentane, we can extrapolate the decomposition fraction to zero coverage of n-pentane, as reported previously for cyclohexane.13 The initial decomposition fractions of n-pentane without and with the maximum CO promotion are estimated to be 0.5 and 1.0, respectively. The experiments to determine decomposition starting temperature of n-pentane were carried out by using the extrapolation method described for cyclohexane. Figure 9 shows a typical result of variations in TPD peak area of hydrogen and CO with preflash temperature for coadsorption of n-pentane and CO (the dose time of n-pentane is 60 s and CO predose time is 20 s). The hydrogen peak area clearly begins to increase at ca. 150 K, and the decomposition starting temperature of the coadsorbed
J. Phys. Chem. B, Vol. 104, No. 36, 2000 8699
Figure 9. Variations in TPD peak area of hydrogen and CO with preflash temperature. n-Pentane was predosed at 118 K for 60 s (the coverage is 0.022 ML) followed by exposure to CO for 20 s (0.2 ML). The surface was flashed at 10 K/s to a preflash temperature. After recooling to 118 K, CO was postdosed to the surface for 40 s and then a TPD spectrum was measured.
n-pentane under the conditions of Figure 9 is determined to be 149 K. By changing CO pre- and postdose time (the sum of CO pre- and postdose time was fixed to 60 s), the decomposition starting temperature of n-pentane with coadsorbed CO is extrapolated to zero CO predose time. The decomposition starting temperature of n-pentane is determined to be 184 K. The desorption of n-pentane starts from ca. 175 K and desorption peaks are observed at 194 and 211 K when n-pentane alone is adsorbed on clean Ni(755) at the same dose time in Figure 9 (see Figure 5a). The decomposition of n-pentane proceeds concomitantly with desorption of n-pentane. This fact is compromised with a low decomposition fraction of n-pentane. 3.2.2. Other Straight-Chain Hydrocarbons. The desorption and decomposition properties of other straight-chain hydrocarbons (n-hexane, n-heptane, and n-octane) were investigated by the same method as cycloheptane,14 and the results including those for n-pentane are tabulated in Table 2. The straight-chain hydrocarbons except n-pentane show only one peak of monolayer desorption which can be assigned to desorption from the (111) terrace as the case for high-reactivity cyclic hydrocarbons (see section 3.1.2). Desorption energy is estimated from the peak temperature of TPD, as described above for cyclic hydrocarbons. Desorption energy of straight-chain hydrocarbons (C5-C8) from the (111) terrace increases by 4.7 kJ/mol per CH2 unit (Ed ) 22.5 + 4.7n kJ/mol for n-Cn), suggesting the flat-lying adsorption configuration of straight-chain hydrocarbons. The increase factor of desorption energy per CH2 unit for straight-chain hydrocarbons is similar to that for cyclic hydrocarbons (4.6 kJ/mol). A similar increase of 4.6 ( 0.4 kJ/mol per CH2 unit has been reported for n-alkyl alcohols on Ag(110); in that case determination of surface coverages by X-ray photoelectron spectroscopy indicates a flat-lying configuration.23 Ratio of desorption-limited H2 peak area to reaction-limited one is estimated also, as described above for cyclohexane by dividing H2 TPD peak area at about 380 K into two regions after preflash and postdosing of CO. The ratio for a straightchain hydrocarbon (n-Cn) can be expressed by (n - 1)/(n + 3), suggesting that each carbon atom except one methyl group releases one hydrogen on average when decomposition to dehydrogenated intermediates is completed. The ratio for straight-chain hydrocarbons is a little different from that for cyclic hydrocarbons (cf. Table 1), which suggests that the decomposition pathway of straight-chain hydrocarbons may be different from that of cyclic hydrocarbons.
8700 J. Phys. Chem. B, Vol. 104, No. 36, 2000
Figure 10. Dependence of decomposition starting temperature of the n-heptane coadsorbed with CO on CO predose time. Decomposition starting temperature was measured as a function of pre- and postdose time of CO as shown for cyclohexane in Figure 3. The experimental procedure is the same as that in Figure 3, only using n-heptane in place of cyclohexane. The sum of CO pre- and postdose time is fixed to 60 s.
For straight-chain hydrocarbons other than n-pentane, their initial decomposition fraction is unity, and their decomposition starting temperature could be determined by using the method for high-reactivity hydrocarbons. Their decomposition starting temperature is much lower than their TPD peak temperature of monolayer desorption, which is also consistent with the fact that these hydrocarbons decompose completely at their lower coverage. The last column of Table 2 shows the estimated value of activation energy for decomposition that is calculated from the decomposition starting temperature assuming a common preexponential factor of 1013 s-1, as described for cyclic hydrocarbons. The decomposition starting temperature of n-heptane was measured also by using the extrapolation method for lowreactivity hydrocarbons to check the validity of the extrapolation method and see whether coadsorbed CO influences the decomposition starting temperature of high-reactivity hydrocarbons, too. Figure 10 clearly shows that the decomposition starting temperature of n-heptane decreases with increase of coadsorbed CO coverage as that of cyclohexane (cf. Figure 4). Although coadsorbed CO reduces the decomposition starting temperature of n-heptane, CO cannot enhance the high decomposition probability of n-heptane any more because n-heptane decomposes efficiently even in the absence of CO. The decomposition starting temperature of n-heptane is determined to be 190 K by the extrapolation method. This value is within experimental error ((5 K) of that determined by the method for high-reactivity hydrocarbons (197 K), which demonstrates the validity of the extrapolation method. 3.3. Relation between Decomposition Starting Temperature and Desorption Energy of Various Hydrocarbons. Figure 11 shows the dependence of decomposition starting temperature of various hydrocarbons on their desorption energy of monolayer desorption. Desorption energy expresses activation energy that is necessary for a molecule to desorb into a vacuum, and it is a good index for the interaction between a molecule and surface. The desorption energy from the (111) terrace is used for low-reactivity hydrocarbons as well as high-reactivity ones. A broken line shows a relationship with desorption energy among straight-chain hydrocarbons, and a solid line shows a relationship among high-reactivity cyclic hydrocarbons. The decomposition starting temperature of cyclohexane is quite different from those of other cyclic hydrocarbons and rather close to those among straight-chain hydrocarbons. Comparing
Orita et al.
Figure 11. Dependence of decomposition starting temperature of various hydrocarbons adsorbed on Ni(755) on their desorption energy from the (111) terrace. Desorption energy of hydrocarbons was estimated by assuming that peaks of monolayer desorption obey firstorder kinetics with a preexponential factor of 1013 s-1. Triangles indicate the values determined from the extrapolation method for coadsorption with CO. Open and filled circles show the results for a hydrocarbon whose decomposition was promoted in the presence of coadsorbed CO and whose decomposition was not promoted, respectively. Broken and solid lines show the relationship with desorption energy among straightchain hydrocarbons and that among high-reactivity cyclic hydrocarbons whose decomposition was not promoted by coadsorbed CO, respectively. For some hydrocarbons, two or three points are plotted to indicate the experimental errors in decomposition starting temperature.
straight-chain and cyclic hydrocarbons with the same carbon atom number, straight-chain ones show larger desorption energy and higher decomposition starting temperature than cyclic ones. (We use the same preexponential factor of 1013 s-1 to estimate desorption energy for both straight-chain and cyclic hydrocarbons. When we use the preexponential factor of 1012 s-1 for straight-chain hydrocarbons, the desorption energies of straightchain and cyclic hydrocarbons with the same carbon atom numbers become almost equal.) The difference in adsorbed structures between straight-chain and cyclic hydrocarbons may result in the difference of desorption energy and/or the difference of the preexponential factor for desorption. The decomposition starting temperature of straight-chain hydrocarbons increases only slightly with an increase of desorption energy, indicating that the activation energy for decomposition is similar for all the straight-chain hydrocarbons. The activation energy for decomposition has several influencing factors (e.g., bond strength between carbon and hydrogen, conformational and strain energy in molecule, interaction with active sites for decomposition, character of transition state for decomposition, and structure transformation) that are necessary for decomposition. For all the straight-chain hydrocarbons, the strength of the C-H bond in the methylene groups is almost equal. Since the straight-chain hydrocarbons have similar structure with different chain length, they interact with the surface similarly. In fact, Chesters et al. have reported that straight-chain hydrocarbons (C2-C6) are adsorbed on Pt(111) in a similar form with their carbon chains parallel to the metal surface in an anti conformation.9 Firment and Somorjai have studied surface structures of adsorbed monolayer of straightchain hydrocarbons (C4-C8) on Pt(111) by LEED and found that the molecules lie with their chain axes parallel to the Pt surface and Pt[11h0].24 It is possible that straight-chain hydrocarbons adsorb along the step sites of Ni(755) with their chain axes parallel to [11h0]. We can explain the fact that all the straight-chain hydrocarbons have similar decomposition starting temperatures if only part of the chain reorients, regardless of chain length, by rotation about the C-C bond in a transition
Decomposition of Hydrocarbons Adsorbed on Ni(755) state for decomposition. It is also possible to explain the fact that the straight-chain hydrocarbons have higher decomposition starting temperature than the corresponding cyclic ones if the straight-chain hydrocarbons adsorb along the step sites with larger adsorption energy than the cyclic ones and/or the straightchain ones have a larger activation barrier to the transition state for decomposition than the cyclic ones. The decomposition starting temperature of high-reactivity cyclic hydrocarbons increases linearly as hydrocarbon becomes heavier, indicating that heavier hydrocarbons need larger activation energy for decomposition. Desorption energy increases monotonically with the number of carbon atoms in the molecule, as reported for Ru(001) by Madey and Yates25 and by Hoffman and Upton,26 suggesting that the molecules are adsorbed with their rings more-or-less parallel to the surface to maximize coordination with the surface. Dependent on ring size, cyclic hydrocarbons have different conformations27 and ring strains (e.g., the strain energy in cyclopentane and cycloheptane is about 26 kJ/mol higher than that in cyclohexane, which is virtually strainless28). The equilibrium conformation of cyclopentane is puckered and can be described in terms of 10 envelope conformations (in which there are eclipsed C-H bonds on adjacent carbons) and 10 half-chair ones (in which hydrogens are in a nearly eclipsed geometry), all of equal energy and interconverting through free pseudorotation with an activation barrier less than 0.4 kJ/mol.27 Therefore, the conformers of cyclopentane can interconvert freely by pseudorotation, and little energy is required for two adjacent C-H bonds to attain eclipsed positions. The thermodynamically stable form of cyclohexane is a chair conformation where hydrogens on adjacent carbons are all staggered. A boat conformation, which is about 23 kJ/ mol higher in energy than a chair conformation, is needed to achieve eclipsed C-H bonds on adjacent carbons (i.e., the two C-H bonds on adjacent carbons lie in the same plane). The most stable form of cycloheptane is the twist-chair conformation, and the 14 possible twist-chair conformations can be easily interconverted by a pseudorotational process with energy maxima at the 14 chair conformations which have eclipsed C-H bonds on adjacent carbons.27 The lowest energy conformation of cyclooctane is the boat-chair form, which has nearly eclipsed C-H bonds on adjacent carbons.27 For cyclic hydrocarbons except cyclohexane, it is considered that little energy is required to attain eclipsed C-H bonds. Teplyakov and Bent have investigated β-hydride elimination from alkyl and cycloalkyl groups on a Cu(100) surface by using TPD.29 They suggest that the transition state for decomposition is a planar geometry (i.e., a dihedral angle of 0° between the substituents on R- and β-carbons). They consider the lowest energy conformation of cyclopentane with its near-eclipsed interactions as being about 23 kJ/mol closer to the transition state for a planar elimination reaction relative to the chair conformation of cyclohexane, where all interactions are staggered, and that the rate difference between cyclopentyl and cyclohexyl comes from the difference in conformational energy. They point out that differences in strain energy between the reactants and the products also play some role in accounting for differences in rate. If we consider that the transition state of cyclic hydrocarbons on Ni(755) is similar to that on Cu(100) (i.e., the two C-H bonds on adjacent carbons lie in the same plane), one of the important factors is the structure transformation which makes decomposition proceed more easily (e.g., a partially desorbed form which can interact strongly with active step sites and/or distortion of the ring structure) because of the following
J. Phys. Chem. B, Vol. 104, No. 36, 2000 8701 reasons: (1) heavier cyclic hydrocarbons must need a larger activation energy for structure transformation in proportion to the strength of interaction with the surface because the whole body of the cyclic hydrocarbons probably needs to reorient in a transition state for decomposition due to their planar adsorbed structure; (2) the energy for desorption into a vacuum must correlate with activation energy necessary for structure transformation which is suitable for decomposition (especially with the energy for partially desorbed form). In the present work, cyclohexane shows much higher decomposition starting temperature than other cyclic hydrocarbons. This result can be explained on the basis of the difference in conformational energy if the rate-determining step of decomposition passes through a high energetic transition state that needs a conformational change from the chair conformation to the boat one. A probable intermediate of this reaction pathway is eclipsed 1,2-diadsorbed cyclohexane, which is proposed for an intermediate of catalytic H/D isotopic exchange reaction between cyclohexane and deuterium (the isotopic exchange reaction occurs faster for cyclopentane than for cyclohexane).30 Lamont et al. have investigated the dehydrogenation of cyclohexane and cyclohexene on the (5 × 20) and (1 × 1) surfaces of Pt(100) using vibrational spectroscopy.31 They have proposed that on the (5 × 20) surface cyclohexane and cyclohexene both undergo a similar conformational change from the chair form to the boat one around 200-220 K and that this change seems to be necessary for dehydrogenation to occur. For cyclic hydrocarbons except cyclohexane, it is considered that little energy is required to attain eclipsed C-H bonds. Methylcyclopentane is the most stable in the conformation with the methyl group at the “tip of the flap” of an envelope form (in which there are eclipsed C-H bonds).27 The difference in activation energy between methylcyclopentane and cyclohexane shown in Table 1 might correspond to the activation energy for cyclohexane to attain eclipsed C-H bonds on adjacent carbons because these two molecules show similar desorption energies from the (111) terrace. The difference in activation energy (10 kJ/mol) is much lower than the difference in conformational energy between the chair form and the boat one (23 kJ/mol) and the ring inversion barrier (46 kJ/mol27) in a free molecule. This result suggests that the conformational change can occur much more easily on the surface than in the gas phase, as Lamont et al. have proposed. It is worth noting here that for straight-chain hydrocarbons the eclipsed C-H bonds on adjacent carbons can be attained by rotation about the C-C bond, and the eclipsed conformation is about 15 kJ/mol higher in energy than the anti conformation (the conformational energy is not dependent on chain length).27 This fact might explain the results that the straight-chain hydrocarbons have higher decomposition starting temperature than the corresponding cyclic ones with the same carbon atom numbers. 3.4. Mechanism for Promoting Effect of CO on Decomposition. In the previous paper,13 we have presented three possible mechanisms for the promoting effect of CO: (1) transformation of the adsorbed structure of a hydrocarbon by the interaction with coadsorbed CO, which may cause a reduction in activation energy for the transition state suitable for decomposition; (2) change in the electronic state of the nickel surface by the electron-accepting property of CO; and (3) CO acts as a hydrogen acceptor, which makes the dehydrogenation step proceed more readily. In relation to decomposition starting temperature, we would like to further discuss possible mechanisms for the promoting effect of CO on decomposition.
8702 J. Phys. Chem. B, Vol. 104, No. 36, 2000 The decomposition starting temperature of a hydrocarbon decreases with increasing amount of coadsorbed CO (cf. Figures 4 and 10). The desorption peak of a hydrocarbon also shifted to lower temperature by interaction between a hydrocarbon and CO (cf. Figure 7a), indicating that the adsorption state of a hydrocarbon becomes less stable in the presence of coadsorbed CO. Thus, coadsorbed CO probably reduces the activation energy for structure transformation suitable for decomposition due to adsorbate-adsorbate interactions and makes decomposition proceed easily. Another possibility of adsorbate-adsorbate interactions is that a hydrocarbon gets closer to active step sites by the lateral interactions between CO and a hydrocarbon. Weiss et al. have reported multilayer-induced dissociative chemisorption of cyclobutane on Ru(001) and observed that the activation energy leading to decomposition of cyclobutane is reduced from 42.2 to 31.4 kJ/mol by the presence of multilayers.32 They propose a similar mechanism wherein the adsorbed structure of the first layer and the structure/energetics of the transition state are changed due to adsorbate-adsorbate interactions between the first and higher layers. Concerning adsorbateadsorbate interactions of the hydrocarbon multilayer, Nuzzo et al. have recently investigated the transport and structural phase dynamics in multilayer assemblies of hydrocarbons on Pt(111).33 They have found that a mixed monolayer of cyclooctane with n-octane-d18 shows two distinct states for the desorption of the cyclooctane; one is seen at the normal peak desorption temperature for cyclooctane and the other at a higher temperature appropriate for n-octane. They suggest that the attractive lateral interactions in the monolayer lead to the formation of island domains and that the desorption kinetics appear to reflect the underlying rate/structure sensitivity. The coadsorption system of CO and a hydrocarbon probably forms island domains by the lateral interactions, too, and the change of adsorbed structure may lead to the promotion of hydrocarbon decomposition. Cooper et al. have investigated the adsorption of cyclohexane on clean and oxygen-covered Ni(111) surfaces by using IRAS.34 The effect of coadsorbed oxygen is strongly coverage dependent; adsorption of cyclohexane on the (2 × 2) oxygen-covered Ni(111) surface results in further downshift in frequency of the C-H stretching “soft” mode. They consider that this result indicates the importance of charge transfer from the filled CH σ-orbital to the metal in weakening the C-H bond. Adsorption on the (x3 × x3)R30° oxygen-covered surface leads to total suppression of C-H stretching “soft” mode, which is due to steric blocking of bare metal sites by the adsorbed oxygen atoms. Coadsorbed CO might influence the charge-transfer interaction between the C-H bond and the metal surface and weaken the C-H bond as oxygen adatom due to the electron-accepting property of CO. Although it is difficult to determine which mechanism for the promoting effect of CO is preferable within our knowledge of interactions between CO and hydrocarbons, we favor the mechanism of adsorbate-adsorbate interactions since there is much evidence of dissociation and desorption changes due to those interactions on other surfaces. Summary From TPD experiments using the effects of CO coadsorbed with cyclic and straight-chain saturated hydrocarbons (C5-C8), we have determined desorption energy (assuming a first-order preexponential factor of 1013 s-1), decomposition fraction (ratio of decomposed molecules to adsorbed ones), decomposition starting temperature (the temperature where a hydrocarbon starts to decompose), and ratio of desorption-limited H2 peak area to
Orita et al. reaction-limited one. These quantities are very dependent on the molecular structure of hydrocarbons. Comparing straightchain and cyclic hydrocarbons with the same carbon atom number, the straight-chain ones show larger desorption energy and higher decomposition starting temperature than the cyclic ones, which show the difference in the interaction with the surface between straight-chain and cyclic hydrocarbons. The decomposition starting temperature of straight-chain hydrocarbons increases only slightly with the increase of carbon atom number, whereas desorption energy from the (111) terrace increases by 4.7 kJ/mol per CH2 unit. This fact suggests that the activation energy for decomposition is similar for all the straight-chain hydrocarbons and only part of the chain reorients, regardless of chain length, by rotation about the C-C bond in a transition state for decomposition. For cyclic hydrocarbons except cyclohexane, the decomposition starting temperature increases linearly with an increase of desorption energy from the (111) terrace (the desorption energy of cyclic hydrocarbons increases by 4.6 kJ/mol per CH2 unit), suggesting that the whole ring structure needs to reorient in a transition state for decomposition. Cyclohexane shows a much higher decomposition starting temperature and lower decomposition fraction than other cyclic hydrocarbons, which can be explained on the basis of the difference in conformational energy to attain eclipsed C-H bonds on adjacent carbons in a transition state for decomposition. Relating decomposition starting temperature with desorption energy, we have discussed the decomposition mechanism of hydrocarbons and proposed some candidates for transition state (e.g., a partially desorbed form). References and Notes (1) Maire, G. L. C.; Garin F. G. In Catalysis, Science and Technology; Anderson, J. R., Boudart, M., Eds.; Springer-Verlag: Berlin, 1984;Vol. 6, p 161. (2) Wittrig, T. S.; Szuromi, P. D.; Weinberg, W. H. J. Chem. Phys. 1982, 76, 3305; Surf. Sci. 1982, 116, 414. (3) Szuromi, P. D.; Engstrom, J. R.; Weinberg, W. H. J. Phys. Chem. 1985, 89, 2497. (4) Campbell, C. T.; Rodriguez, J. A.; Henn, F. C.; Campbell, J. M.; Dalton, P. J.; Seimanides, S. G. J. Chem. Phys. 1988, 88, 6585. (5) Rodriguez, J. A.; Campbell, C. T. J. Phys. Chem. 1989, 93, 826. (6) Kuramochi, H.; Kunimori, K.; Uchijima, T.; Kondoh, H.; Shindo, H.; Kaise, M.: Nishihara, C.; Nozoye, H. Surf. Sci. 1991, 251/252, 926. (7) Lehwald, S.; Ibach, H. Surf. Sci. 1979, 89, 425. Hoffman, F. M.; Felter, T. E.; Thiel, P. A.; Weinberg, W. H. Surf. Sci. 1983, 130, 173. (8) Bussell, M. E.; Henn, F. C.; Campbell, C. T. J. Phys. Chem. 1992, 96, 5978. (9) Chesters, M. A.; Gardner, P.; McCash, E. M. Surf. Sci. 1989, 209, 89. (10) Raval, R.; Parker, S. F.; Chesters, M. A. Surf. Sci. 1993, 289, 227. (11) Henn, F. C.; Diaz, A. L.; Bussell, M. E.; Hugenschmidt, M. B.; Domagala, M. E.; Campbell, C. T. J. Phys. Chem. 1992, 96, 5965. (12) Hugenschmidt, M. B.; Diaz, A. L.; Campbell, C. T. J. Phys. Chem. 1992, 96, 5974. (13) Orita, H.; Kondoh, H.; Nozoye, H. Chem. Phys. Lett. 1994, 228, 385. (14) Orita, H.; Kondoh, H.; Nozoye, H. J. Phys. Chem. 1995, 99, 3648. (15) Zebisch, P.; Huber, W.; Steinruck, H.-P. Surf. Sci. 1991, 244, 185. (16) Trenary, M.; Uram, K. J.; Bozso, F.; Yates, J. T., Jr. Surf. Sci. 1984, 146, 269. (17) Surnev, L.; Xu, Z.; Yates, J. T., Jr. Surf. Sci. 1988, 201, 1. (18) Manner, W. L.; Girolami, G. S.; Nuzzo, R. G. J. Phys. Chem. B 1998, 102, 10295. (19) Jiang, L. Q.; Avoyan, A.; Koel, B. E.; Falconer, J. L. J. Am. Chem. Soc. 1993, 115, 12106. (20) Rehead, R. A. Trans. Faraday Soc. 1961, 57, 641; Vacuum 1962, 12, 203. (21) Manner, W. L.; Hostetler, M. J.; Girolami, G. S.; Nuzzo, R. G. J. Phys. Chem. B 1999, 103, 6752. (22) Manner, W. L.; Dobois, L. H.; Girolami, G. S.; Nuzzo, R. G. J. Phys. Chem. B 1998, 102, 2391. (23) Zhang, R.; Gellman, A. J. J. Phys. Chem. 1991, 95, 7433. (24) Firment, L. E.; Somorjai, G. A. J. Chem. Phys. 1977, 66, 2901. (25) Madey, T. E.; Yates, J. T., Jr. Surf. Sci. 1978, 76, 397.
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