Alkane Oxidation on Rh(111) Single-Crystal Surfaces under High

Mar 10, 2010 - Services, 52 Aintree Ave, Mangere, Manukau 2022, Auckland, New Zealand. This article is part of the C: D. Wayne Goodman Festschrift spe...
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J. Phys. Chem. C 2010, 114, 16946–16954

Alkane Oxidation on Rh(111) Single-Crystal Surfaces under High-Temperature, Short-Contact-Time Conditions: A Molecular Beam Kinetic Study† Jarod N. Wilson‡ and Francisco Zaera* Department of Chemistry, UniVersity of California, RiVerside, California 92521 ReceiVed: January 22, 2010; ReVised Manuscript ReceiVed: February 27, 2010

The kinetics of the partial oxidation of alkanes with molecular oxygen on Rh(111) single-crystal surfaces were studied by using a collimated effusive molecular beam under ultrahigh vacuum conditions. These experiments were conceived to probe the details of the mechanism of these processes under high-temperature, short-contact-time conditions. It was determined that the primary products in these reactions are H2 and CO, the components of syngas, as it has been suggested by the so-called direct mechanism. The production of water is also detected, but at a slower rate because of the slow formation of the surface hydroxo intermediate involved. Carbon dioxide, on the other hand, is never produced, because carbon monoxide desorbs before having the opportunity to react further with adsorbed oxygen atoms. Both surface and subsurface atomic oxygen species form during the reaction, but only the more labile surface intermediate is relevant, blocking Rh sites for alkane activation and reacting to form CO and H2O. 1. Introduction Oxidation reactions are typically quite favorable from a thermodynamic point of view, and usually go to completion. In the case of organic reactants, the final products often are carbon dioxide and water. Such combustion is commonly used to generate heat, but it may not be the best way to spend fossil fuels. In many instances, it is better to convert the raw material into chemicals such as olefins so they can be used as feedstocks in manufacturing and petrochemical industries. It may also be desirable to control the hydrocarbon oxidation to produce a more desirable carbon monoxide + hydrogen (syngas) mixture. To this end, it has been shown that, by using a metal catalysts and relatively high (∼1000 K) temperatures, it may be possible to oxidize alkanes to either olefins1,2 or syngas.3-5 These processes have been shown to require short contact times between the catalyst and the reaction mixture, something that may be achieved by using high space velocities, short catalyst beds, and/or new reactor designs.4,6-12 However, optimization of the selectivity of these processes has been hampered by the limited understanding presently available on the kinetics of the reactions involved. Two main mechanisms have been proposed to explain the high-temperature, short-contact-time conversion of alkanes with oxygen on metal catalysts: (1) a direct partial oxidation to H2 and CO and (2) total combustion to H2O and CO2 followed by reformation back to the desired syngas mixture.5,13-17 The resolution between these two mechanisms has proven quite difficult and is still ongoing. One central problem is that, because oxidations are fast exothermic reactions, it is difficult to extract elementary step information from kinetic experiments using regular catalytic setups, where the true rates may be masked by contributions from heat and mass transport.5,18 Here we report on studies using molecular beams on single crystals in order to †

Part of the “D. Wayne Goodman Festschrift”. * Corresponding author. E-mail: [email protected]. ‡ Present address: Watercare Lab. Services, 52 Aintree Ave, Mangere, Manukau 2022, Auckland, New Zealand.

obtain direct information on the kinetics of the alkane oxidation surface reactions and isolate them from other interfering factors. In this manuscript, we expand on our initial report where we showed that both hydrogen and water are primary products from the oxidation of isobutane with oxygen on rhodium substrates and that the selectivity between the two depends on the conditions used to carry out the reaction.19,20 Also, it was determined that the only carbon-containing product is carbon monoxide; no carbon dioxide is ever produced as a primary product in these systems. Here we show that those initial conclusions are general, applying to other alkanes, and that the activation of the hydrocarbons on the surface is the rate limiting step of the overall conversion. We also discuss the role that oxygen plays in defining the kinetics of the surface reactions. It was determined that the oxygen uptake occurs in two stages, to produce surface and subsurface atomic oxygen species, respectively, and that it is the first, more labile species the one involved in the alkane oxidations. The atomic oxygen adsorbed on the surface significantly inhibits the uptake of the alkane and its subsequent conversion, and recombines with surface carbon and hydrogen to produce CO and H2O. 2. Experimental Section All experiments were carried out in a molecular beam system described in detail previously.21-23 In brief, all instrumentation is placed inside a small (6 L) ultrahigh vacuum (UHV) chamber turbopumped to a base pressure of the order of 2 × 10-10 Torr. The solid sample used as the catalyst, a rhodium single crystal cut and polished to expose a (111) facet, is mounted on a manipulator capable of x-y-z-θ motion and of cooling to liquid-nitrogen temperatures and of resistive heating to temperatures of up to approximately 1250 K. The temperature of the crystal is followed by using a chromel-alumel thermocouple spot-welded to the side of the crystal, and monitored and controlled by using homemade electronics. The front surface is cleaned by a combination of ion sputtering and chemical treatments with O2 until being able to reproduce the temperatureprogrammed desorption (TPD) spectra for CO24 and H225

10.1021/jp1006507  2010 American Chemical Society Published on Web 03/10/2010

Alkane Oxidation on Rh(111)

Figure 1. Isothermal kinetic data for the conversion of a 1:1 isobutane + O2 mixture on a Rh(111) single-crystal surface at 1165 K, acquired using the effusive molecular beam setup and experimental procedure described in the text. The evolution of the net flow of the reactants (oxygen and isobutane) and products (hydrogen, carbon monoxide, water, and carbon dioxide) are shown as a function of time. H2, CO, and H2O, but not CO2, are produced, in the expected stoichiometric ratios; the first two with rapid kinetics following the activation of the alkane and water in a more slow fashion, via the formation of hydroxo surface intermediates.

reported in the literature. The gases, argon (Liquid Carbonic 99.999%), oxygen (Nellcor Purtain Bernett, 99.5%), methane (Matheson, 99.999%), propane (Matheson, 99.993%), and isobutane (Matheson, 99.985%), were all used as provided. n-Hexane (Aldrich, g99%) was subjected to several freezepump-thaw cycles for purification and vaporized under vacuum in the gas manifold to make the appropriate reaction mixtures. The kinetic experiments are performed by using an effusive molecular beam source and an out-of-line-of-sight quadrupole mass spectrometer. The molecular beams are produced by dosing the gases through a 1.2-cm-diameter multichannel microcapillary array. The gas flux is set by presetting the position of an intermediate UHV leak valve in previous calibration experiments, and by fixing the pressure of the gas mixture in the gas manifold, where the hydrocarbon and oxygen gases are premixed to the desired ratios.21 The solid sample is placed at a distance of ∼2 cm from the doser in order to ensure interception of the central portion of the beam, which has been shown previously to exhibit an approximate flat flux profile.21,26 The system is equipped with a stainless steel flag to intercept and reinstate the molecular beam at will. The partial pressure of the different gases is monitored by simultaneously following the time evolution of the signals of up to 15 different masses using a personal computer interfaced to the mass spectrometer. The procedure used for the kinetic experiments has been reported before.21,27,28 It is also illustrated by the data reported in Figure 1, which shows the evolution of the net flow rates of the relevant gases versus time for the case of the oxidation of an isobutane + oxygen mixture on Rh(111) at a surface temperature of 1165 K. The experiments are initiated by turning on the molecular beam on while keeping the flag in the blocking position (t ) 0 s). At that point the gases hit the flag and scatter into the UHV chamber, a process that leads to an increase in the net flow rates of both reactants. Once those pressures reach their new steady-state values, the flag is removed from the path of the beam to allow the gases to directly impinge on the Rh(111) surface (t ) 50 s). As a result, the net flow rates of the reactants decrease, reflecting their consumption by the chemical reaction that takes place on the catalyst, and, accordingly, those

J. Phys. Chem. C, Vol. 114, No. 40, 2010 16947 of the products (H2, H2O, and CO) increase. An initial transient is observed in all traces as the surface coverages of the different adsorbed species adjust to the new conditions, but a new steadystate condition is reached after a few seconds. Following an appropriate time delay, the beam is intercepted again in order to better estimate the contributions of the reactions on the surface to the flux changes measured (t ) 70 s). Once again, a transient is observed, after which the system returns to the steady-state condition seen before the removal of the flag. The unblocking (t ) 90, 130 s) and blocking (t ) 110, 150 s) of the beam is repeated several times to ensure that the kinetic data are reproducible, and eventually the gases are pumped away from behind the capillary array to terminate the isothermal kinetic measurements. If further analysis of the sample is desired, via chemical titrations, the crystal is cooled down once the vacuum returns to its baseline value, the surface dosed with the appropriate exposure of the titrant (CO in the case of the experiments reported in Figure 6), and the temperature ramped linearly as the partial pressure of the different possible products is recorded in typical fashion for TPD experiments. The raw data from these experiments are in the form of voltages versus time, one set for each mass recorded. Those voltages are the signals detected by the mass spectrometer, and are proportional to the partial pressures of the corresponding compounds. Moreover, because of the high pumping speed conditions used in the UHV chamber, they are also proportional to the net fluxes of the gases, a direct reflection of reaction rates. Conversion of the raw signals into either partial pressures or flow rates requires multiple calibration steps.21,29,30 The calibration of the different partial pressures was performed against the readings of a nude ion gauge, adjusted by their reported sensitivities,31,32 as each pure compound was leaked individually into the UHV chamber. In the cases of CO,24,33 H2,25,34 and O2,35 the partial pressures were converted into fluxes, in units of mML/s (1 mML ) 1 molecule per 1000 surface Rh atom) by equating the integrated signals from individual uptake experiments on clean Rh(111) to the corresponding reported saturation coverages. For all other gases, the pressure-to-flux conversion factors were estimated by mass balance arguments using specific molecular beam experiments designed to yield a minimal number of products. 3. Results The main purpose of our studies has been to identify the mechanistic details of partial oxidation reactions for alkanes on rhodium surfaces under high-temperature, short-contact-time conditions. The data in Figure 1 can already be used to identify some of those:36 1. Steady-state alkane conversion can be clearly reached for these systems even under UHV conditions, an observation that attests to the high probability of these reactions. 2. The primary products from these reactions are carbon monoxide and hydrogen. Large signals are detected for those products, and the reaction stoichiometry is close to that expected for such primary reaction: i-C4H10 + 2O2 f 4CO + 5H2. 3. A competing channel is also indicated by the production of a small amount of water. The slow transient seen in the water trace highlights the secondary nature of this surface reaction.19 4. No CO2 production is observed, either as a primary or as a secondary product. No evidence is seen for a water-shift reaction.37 5. The formation of small amounts of other transient species on the surface is hinted by the mirroring spikes in the O2 and H2 traces seen immediately after blocking and unblocking of the beam.

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Wilson and Zaera TABLE 1: Experimental Parameters and Reaction Rates and Probabilities for the Data Reported in Figure 2a F(HC)/ F(O2)/ R(H2)init/ R(H2)ss/ R(HC)ss/ mML · s-1 mML · s-1 mML · s-1 mML · s-1 mML · s-1 Preact,ss methane propane isobutane n-hexane

258 117 44 11

0 0 0 0

Alkane Alone 19.2 22.1 30.9 14.0

0.5 9.7 22.1 13.1

0.6 2.2 2.6 1.1

0.001 0.021 0.100 0.169

methane propane isobutane n-hexane

250 92 36 13

23 12 91 115

with Oxygen 24.4 14.9 17.5 17.6

0.3 2.0 3.8 4.3

0.0 0.5 1.0 0.7

0.002 0.009 0.033 0.128

a

Figure 2. H2 production probability versus time on Rh(111) at 1130 K under steady-state conditions, normalized to account for the stoichiometries of the reactions. Data are shown for molecular beams of methane, propane, isobutane, and n-hexane containing either the hydrocarbon alone (left panel) or a mixture of the alkane and O2. The compositions of the beams are discussed in the text. The reaction probabilities are in general quite high, higher for the heavier hydrocarbons and for beams without oxygen.

The formation of the primary products, H2 and CO, is believed to occur via the recombination of the constituent atoms on the surface, which is known to be quite fast.25,38 Therefore, the kinetics of alkane oxidation is likely to be controlled by the dissociative adsorption of the reactants, of the alkanes in particular. Those kinetics were investigated independently here by performing molecular beam isothermal kinetic experiments with pure alkanes. The left panel of Figure 2 shows typical data from experiments with four alkanes, methane, propane, isobutane, and hexane, at T ) 1130 K. Their fluxes were fixed in all cases to correspond to a steady-state pressure of 1 × 10-8 Torr in the UHV chamber. The actual fluxes are somewhat different for each compound because of differences in ion gauge and mass spectrometer sensitivities, but such effects were normalized in this figure by reporting reaction probabilities instead, by ratioing the reaction rates against the flow rate of the molecular beam. The details of the parameters used and the rates and probabilities calculated from these experiments are provided in Table 1. The conversion rates in this figure were followed by measuring the evolution of the net fluxes of H2, which were also normalized to account for the appropriate stoichiometry of each compound, but rates are also reported for the consumption of the alkane in Table 1 to illustrate the level of error involved in the measurements. Reaction probabilities could only be determined reliably down to values of about 0.2%. One immediate observation from the data in the left panel of Figure 2 is the high probabilities observed for the activation of alkanes on the clean metallic surface of rhodium at these high temperatures. In the case of n-hexane, the activation probability at 1130 K is almost 20%. Also, this probability increases following the sequence methane < propane < isobutane < n-hexane. A combination of factors can be identified as responsible for this trend. First, the larger the molecules, the larger the number of C-H bonds that can be broken. This, however, does not alone explain our observations: methane in particular is quite unreactive, much less reactive than what would be predicted by this argument. Also to note is the relatively large reaction probability seen for isobutane. The reaction probabilities reported here can be roughly correlated with the energy of the weakest C-H bond in each of the molecules considered: 439 kJ/mol for methane, 409 kJ/mol for

Tabulated here are the beam fluxes of the alkane and oxygen (F(HC) and F(O2)), the initial and steady-state reaction rates for hydrogen production (R(H2)init and R(H2)ss), and the steady-state rate of alkane consumption (R(HC)ss), all in units of mML/s. The last column provides the reaction probability under steady-state conditions, Preact,ss, estimated by dividing the steady-state rate of hydrogen + water production by the flux of the alkane (F(HC)) and normalizing to account for the stoichiometry of the reaction.

Figure 3. Kinetics of n-hexane decomposition on Rh(111) as a function of temperature. Left panel: rate of evolution of H2 as a function of time at four different temperatures showing the slow decay due to poisoning of the surface by the surface carbon byproduct of alkane decomposition. Right: Summary of kinetic parameters extracted from the data on the right, reflecting the activated nature of the uptake and the poisoning of the surface over time.

propane (central carbon), and 404 kJ/mol for isobutane (the value for n-hexane was not available).39 This reflects the relative strengths of the C-H bonds involving primary versus secondary versus tertiary carbons. The temperature dependence of the activation of n-hexane on clean Rh(111) is reported in Figure 3. The left panel shows the time dependence of the rate of hydrogen evolution over time, for a total period of 80 s, at several surface temperatures, whereas the right panel summarizes some of the key kinetic data derived from the isothermal molecular beam runs. The dissociative adsorption of n-hexane shows a mild temperature dependence, the reactivity increasing with increasing temperature. This conclusion is general, and was seen with the other alkanes as well (data not shown). The trend observed is not entirely expected, since sticking coefficients in adsorption phenomena tend to display only a weak temperature dependence if any, and often show a slight decrease in value with increasing temperature (because of a decrease in the ability of the molecules to dissipate their excess kinetic energy into a hotter surface).40 On the other hand, adsorption in these systems is activated, and is likely to involve a weakly adsorbed precursor state. Interest-

Alkane Oxidation on Rh(111) ingly, a slight cooling of the surface was observed in these experiments upon exposure to the molecular beam, more at higher surface temperatures (∆T ∼ 1 K at T ∼ 1200 K, data not shown). In any case, an Arrhenius fit of the data for the initial H2 production rate yields an apparent activation energy value of approximately 22 kJ/mol for the case of n-hexane on Rh(111). The other thing to note in Figure 3 is a slow decrease in the rate of hydrogen evolution over time. This is because the surface becomes poisoned by the deposition of carbon from total dehydrogenation of the alkane. In the case of n-hexane, the reaction rate is reduced by a factor of between a quarter and half after t ) 80 s (Figure 3, right). Similar, albeit slower, poisoning was seen with the other alkanes as well. Finally, more hydrogen is produced, and the reaction is poisoned at a slower rate, at 1225 K (compared with 1020 and 900 K), perhaps because at such higher temperatures the carbon atoms produced by the alkane dehydrogenation may diffuse into the bulk of the crystal, freeing some surface site for further reaction. Support for the idea of the excess carbon diffusing into the bulk comes from the fact that after these molecular beam experiments the surface exhibits adsorption behavior in TPD experiments with CO and other probe molecules identical to that expected of clean Rh(111) (data not shown). In any case, according to the H2 yields reported in the right panel of Figure 3, the total alkane uptake over the entire 80 s of the experiments is quite small, never surpassing 10% of a monolayer (it amounts to 3% at 900 K, for instance). Much lower carbon coverages must be reached under steady-state reaction conditions, under which most of the carbon and hydrogen produced by the activated adsorption of the alkane is consumed to yield CO, H2, and H2O. The sticking coefficient and dissociative adsorption rate of alkanes on Rh(111) are significantly reduced when oxygen is added to the molecular beam mixture. This is clearly shown by the data in the right panel of Figure 2. In these experiments, the flux of the alkanes was kept the same in the experiments with versus without oxygen. The oxygen flux used varied in each case but proved to not be critical, because, after a threshold value, further increases in O2 flux did not led to any measurable changes in reaction probabilities (data not shown). In general, it was determined from the data in Figure 2 that not only the initial values for the hydrogen production from alkane decomposition on Rh(111) are lower than with beams consisting of the pure alkane but also that they decrease further over a period of a few seconds, until reaching a low final steady-state value. The decreases in reaction probabilities observed over time after the initial transient (∆t ∼ 1-2 s) are mostly accounted for by an accompanying increase in the production of water (data not shown),19 but the absolute total conversion of the alkanes is nevertheless inhibited by oxygen. Several additional experiments were carried out to further test the effect of oxygen on the activity and selectivity of these alkane partial oxidation reactions. Figure 4 shows the molecular beam data obtained for the conversion of propane on Rh(111) at 1125 K using three different gas mixtures, one consisting of pure propane and two with varying amounts of oxygen added. As in the example discussed above, the initial hydrogen yields are also reduced significantly here once oxygen is added to the mixture. This is true even at the beginning of the exposure of the surface to the beam, and is a phenomenon independent of the further decrease in hydrogen yield observed over a period of ∆t ∼ 20 s, which is due to a switch in selectivity toward the production of water. Another important observation in this figure is the steady production of carbon monoxide when oxygen is

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Figure 4. Raw kinetic data for the steady-state partial oxidation of propane over Rh(111) at 1125 K as a function of the flux of O2 added to a molecular beam of the hydrocarbon. The left panel shows that without any oxygen in the beam the only reactions possible are the uptake of the propane and the evolution of hydrogen. The other two panels, which correspond to experiments where oxygen was added to the gas mixture, displays additional changes in rates for the consumption of O2 and the production of both CO and water. Those two cases display quite similar kinetics in spite of the doubling of the flux of oxygen between the data in the center and right panels; only the initial stages of the uptake of oxygen show a significant variation. Also to note is the slow evolution of the rate of water production toward its steadystate value.

added to the gas mixture. This reaction channel provides a way to clean the residual carbon deposited on the surface by the dehydrogenation of the alkane, and allows for the reaction to run indefinitely. It is also interesting to notice that the kinetics of reactions are quite similar for all the products of the alkane oxidation reaction in the two experiments with the two different fluxes of O2. A significant difference is seen only in the initial transient of the uptake of oxygen, where the initial adsorption rate doubles with a doubling of the O2 flux (the initial sticking probability staying approximately the same). However, even for that compound, the steady-state rate settles at approximately the same value regardless. The details of the behavior of the oxygen consumption kinetics will be addressed in more detail below. The reduction in reaction rate in these partial oxidations with O2 addition into the beams can be ascribed to the buildup of a layer of chemisorbed atomic oxygen on the surface. Figure 5 displays data from experiments designed to confirm this hypothesis. The left panel shows the poisoning effect of presaturating the Rh(111) surface with oxygen on the rate of hydrogen production from isobutane. It is clearly seen that, at the relatively low temperature of this experiment (T ) 700 K), a monolayer of adsorbed atomic oxygen completely inhibits the activation of the alkane. Because at this temperature the desorption of oxygen from the surface is quite slow, the poisoning effect lasts a long time. At higher temperatures, by contrast, oxygen desorption does occur, and that cleans up the surface and activates it toward alkane adsorption unless fresh oxygen molecules are fed from the gas phase. The kinetic behavior in such high-temperature regime is illustrated by the data in the right panel of Figure 5, for T ) 1030 K and propane in this example. The hydrocarbon uptake and decomposition kinetics reported before were reproduced here by using a beam of pure propane during the first 50 s of the experiment. Significant hydrogen production was observed, modulated by the blocking and unblocking of the beam, as expected. However, when O2 was added to the mixture (at t )

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Figure 5. Data from experiments designed to test the effect of surface oxygen on the conversion of alkanes on Rh(111). Left panel: Hydrogen evolution rate from the conversion of isobutane on clean (top, blue, trace) and oxygen presaturated (bottom, red, trace) surfaces at 700 K. Right panel: Time dependence of the rates of evolution of hydrogen, carbon monoxide, water, and carbon dioxide in an experiment where a beam of pure propane is first directed at the Rh(111) crystal held at 1030 K until steady state is achieved, and O2 is then briefly added to the mixture. These data show the immediate poisoning of the surface as oxygen is added to the gas, and the slow recovery of the original activity once the O2 is removed.

50 s), the production of hydrogen was shut down almost entirely. A significant amount of CO was also detected at this point (together with a much smaller yield of CO2), the result of a recombination step of the carbon deposited previously from propane decomposition with the newly added oxygen. A slow evolution of water is seen as well, possibly the result of a stepwise hydrogenation of adsorbed atomic oxygen to hydroxo groups first and to water afterward.19 The latter part of the experiment, starting at t ) 70 s, shows how all these changes can be reversed once the oxygen supply is turned back off. CO production stops immediately, and CO2 and H2O stop more slowly. Hydrogen production is restored as well, albeit after a long induction period (∆t ∼ 50 s), during which all remaining O(ads) and OH(ads) groups are removed in the form of water. An alternative way to probe the evolution of the surface species deposited on the Rh(111) surface during these alkane partial oxidation reactions is by performing postreaction titration experiments. An example of the data obtained from such studies is provided in Figure 6, in this case for a 1:2 mixture of isobutane and oxygen directed at a Rh(111) surface heated to 825 K. A low reaction temperature was chosen to better control the buildup of the surface species, and three experiments were carried out for three different reaction times. In each case the crystal was then cooled to room temperature and saturated with CO, a titrant used to estimate the nature and coverage of the oxygen deposited by the reaction mixture, and TPD traces were obtained for CO, CO2, and O2. As expected, the total yields estimated from the TPD signals increase with increasing reaction times, attesting to the buildup of oxygen on the surface. In addition, the titration experiments also suggest the existence of more than one type of chemically different adsorbed oxygen species. Specifically, oxygen removal in these TPD experiments occurs in two ways, as CO2 at low temperatures, and as O2 above 900 K. The CO2 trace, which is produced by recombination of CO with the more reactive adsorbed atomic oxygen, grows significantly between the “low” and “medium” reaction times (t ∼ 2 and 10 s, respectively) but is similar for the “medium” and “long” reaction times (t ∼ 10 and 40 s, respectively). By contrast, the O2 TPD signal, the product of

Wilson and Zaera

Figure 6. CO, CO2, and O2 temperature programmed desorption (TPD) data from postmortem titration experiments performed after the conversion of isobutane + O2 mixtures on Rh(111) at 825 K, by saturating the surface after reaction with CO at room temperature. The three panels correspond to experiments where the reaction was stopped after ∼2, 10, and 40 s (left, center, and right panels, respectively). The increases in the yields of CO2 (first) and O2 (in the last panel) indicate an increase in the total amount of adsorbed atomic oxygen and also the split of that oxygen into two types: one that easily reacts with CO, which populates first, and a second that only desorbs as O2 above 900 K.

Figure 7. Isothermal kinetic data for the uptake of oxygen from CH4 + O2 beams on Rh(111) at four different temperatures. The left panel displays the raw data, whereas the right summarizes some key parameters extracted from the original traces. Two uptake regimes are identified: a rapid adsorption of a labile oxygen species within the first secondary carbon > tertiary carbon. A simple mechanism can be used to interpret the kinetic data obtained here based on the initial dissociative adsorption of both O2 and the alkane and the recombination of the appropriate surface species to produce O2 (the reverse step of its adsorption), H2, and CO.19 In this mechanism, water formation occurs in a stepwise fashion, via the formation of a hydroxo surface intermediate. Modeling of the raw data using such a simple model can not reproduce all of the data, in particular some of the transient behavior, but was already used to highlight a few of the key features of the alkane partial oxidation reactions, including the sharp transient in the uptake of O2 and the slow initial production of water.19 The rate-limiting step in these reactions is most likely the activated adsorption of the alkane, since the subsequent recombination of hydrogen atoms to H225 and of carbon and oxygen atoms to CO38 are both steps known to happen with high probability at much lower temperatures. Therefore, the data in Figures 2 and 3, as well as most of the other raw kinetic data from our isothermal kinetic studies, are likely to directly reflect the kinetics of the dissociative adsorption of the alkane, which is believed to be initiated by the activation of one (the weakest) C-H bond. The data in Figure 3 shows that the dissociative adsorption of alkanes is an activated process. Indeed, the rate of dissociative adsorption of n-hexane (and of the other alkanes tested here) increases with temperature. The temperature dependence is somewhat weak, amounting to the equivalent activation energy of about 22 kJ/mol (for n-hexane), but surprising nevertheless because it does not follow the decrease in probability with temperature typical of adsorption processes. This observation suggests that the incoming molecules may be trapped initially into a weak physisorbed state on the surface, allowing for most of the kinetic energy to be dissipated into other energy modes within the solid. In this picture, the weakly physisorbed state then undergoes C-H bond activation in its way to a more strongly bonded dissociative chemisorbed state. Two things are worth mentioning at this stage, though: (1) it is surprising that the overall process displays such high reaction probabilities, given that the weak physisorbed state would be expected to be mostly unpopulated at the high temperatures of the reaction, the equilibrium shifted toward the gas phase, and (2) the

Wilson and Zaera dissociation of the alkane from the physisorbed state must display a very early transition state, because the measured activation energy represents only a small fraction of the total energy required for the C-H bond dissociation (although that may be compensated by the simultaneous formation of Rh-H and Rh-C bonds). Two more observations are worth highlighting from the data in Figures 2 and 3: (1) the rate of the uptake of the alkane on the clean Rh(111) surface decreases with time, presumably because of the buildup of a layer of adsorbed carbon that poisons the surface, and (2) the reaction probability also decreases substantially when oxygen is added to the gas mixture. This latter observation was studied in more detail in this work, and additional data are reported in Figures 4-7. One observation deriving from the data in Figure 4 is the existence of three time regimes in the kinetic data. Indeed, a pseudosteady-state rate is reached for H2 desorption in the second regime, right after the first fast hydrogen transient seen within the first 1-2 s of the unblocking of the molecular beam. This rate is already lower than that seen with the pure alkane beam. Moreover, the H2 evolution rate decreases further, in a more slowly fashion, over the following 10-20 s of reaction. The first reduction in pseudosteady-state rate is a direct manifestation of the reduced reactive sticking coefficient of the alkane on the oxygen-covered surface, a fact that suggests that activation occurs on rhodium (not oxide) sites. The second reduction, on the other hand, reflects a slow switch in selectivity from H2 to H2O production as the surface coverage of OH(s) species builds up on the surface. The inhibition of the rate of reaction via oxygen adsorption on the surface is further evidenced by the data in Figures 5 and 6. In Figure 5, it is clear that the rate of alkane oxidation can be significantly reduced either by preadsorbing oxygen (left panel) or by adding oxygen to the gas mixture afterward (right panel). This indicates that oxygen adsorption is typically faster than alkane activation, and also suggests that alkane activation requires the presence of Rh metal atoms. The buildup of oxygen on the surface may be slowed down by choosing the right reaction conditions, in particular a low surface temperature (as it was done in the experiment in Figure 6), but at low temperatures the buildup of surface oxygen is essentially irreversible, since oxygen recombination and desorption is negligible, and the poisoning of the surface is therefore permanent once a certain O(ads) coverage is reached. In contrast, at higher temperatures, above ∼ 900 K, the O2 desorption step is activated and can remove the O(ads) species from the surface, but oxygen adsorption is also fast. Therefore, a steady-state coverage of atomic oxygen is reached rapidly on the surface with beams containing O2. This surface oxygen can virtually stop any alkane conversion, as in the case of the experiment reported in the right panel of Figure 5. Interestingly, though, the activity can be (slowly) restituted once all O2 is removed from the gas phase, after which all remaining O(ads) reacts to form mostly water but also a small amount of CO2. The buildup of oxygen on the surface follows a complex kinetics behavior indicative of the deposition of at least two different types of atomic oxygen. It is in fact well known that oxygen atoms can diffuse into the subsurface region of the Rh metal.35,44,47 In our studies this is manifested by the two time regimes for the uptake of O2 reported in Figure 7. An initial coverage of labile O(ads) builds up within the first few seconds of exposure of the surface to the O2 beam, more at higher temperatures. This is followed by a much slower, almost steadystate, uptake of additional oxygen over a period of tens of

Alkane Oxidation on Rh(111) seconds. Most of the initial labile O(ads) desorbs once the O2 beam is removed, but some remain on the surface, a coverage of approximately 10 mML. The total oxygen coverage in similar molecular-beam experiments was estimated to be much larger, somewhere between 50 and 130 mML,19 which indicates that the rest correspond to the more strongly bonded state, most likely oxygen in the subsurface region. Nevertheless, it seems that the surface oxygen is the one that limits the reactivity of the Rh(111) surface toward alkane activation. In particular, the overall rate of alkane conversion is fairly independent of the flux of O2 in the beam, even though the total uptake of O2 in the (sub) surface scales with it.19 Both the total oxygen uptake and the rate of alkane activation increase with temperature, but the effect on the latter is milder and also seen on the clean Rh(111), with no O2 in the beam (Figure 3). Finally, in the experiment reported in the left panel of Figure 5 alkane activation is totally inhibited on a surface containing only surface oxygen. We suggest that alkane activation occurs on Rh metal atoms, that those are blocked by O(ads) when O2 is added to the gas mixture, and that neither subsurface oxygen nor any form of oxidized Rh play a role in this chemistry. Finally, it is worth reiterating here some additional mechanistic conclusions reached from our initial experiments on this system.19 With respect to the formation of water, its kinetics do not exhibit a strong dependence on gas composition or total flux; the main changes are seen as a function of reaction temperature. Indeed, both the selectivity for water production and its rate of formation decrease with increasing temperature. It was proposed in our initial report that the first observation may reflect a higher activation energy for hydrogen recombination to H2 relative to that for OH(ads) formation whereas the second may be a consequence of a faster rate for the reverse OH(ads) dissociation step.19 Alternatively, both trends may be related to the slight decrease in surface atomic oxygen coverage with increasing temperature suggested by the data in Figure 7. It is likely that the OH(ads) intermediate is formed only from surface (not subsurface) oxygen. It was also noted in the initial report that under no circumstances carbon dioxide production could be detected in these molecular beam experiments. This is because of the fast rate of carbon monoxide desorption from Rh(111), which does not leave sufficient time for it to react with the surface oxygen.27,28 It is worth noticing that a reverse in relative rates is seen on Pt(111), where CO2 formation has been detected under similar molecular-beam conditions.48 5. Conclusions Our molecular beam experiments have provided some insight into the mechanism of alkane partial oxidation reactions on rhodium surfaces under high-temperature and short-contact-time conditions. Some of the general conclusions are as follows: 1. Alkane partial oxidation can be sustained catalytically with high reaction probabilities even under vacuum conditions. 2. The reaction probability, which is controlled by the dissociative adsorption of the alkane, increases with molecular weight, and inversely correlates with the strength of the weakest C-H bond in the molecule. 3. The reaction rate of these alkane partial oxidation reactions increases with increasing surface temperature. 4. In the initial stages, H2 and CO, the desired syngas components, are the only products of the reaction. 5. Water is also produced, but at a slower rate. 6. No carbon dioxide is detected as a primary product in these reactions, because carbon monoxide desorbs before having the chance to react with any O(ads).

J. Phys. Chem. C, Vol. 114, No. 40, 2010 16953 In addition, the work reported here focused on determining the role of oxygen in these reactions. In that respect, the following was learned: 1. Two types of atomic oxygen species develop under the reaction conditions: a labile adsorbed surface intermediate and a more strongly bonded subsurface oxygen. 2. The surface oxygen is the relevant species for the partial oxidation of alkanes. 3. This surface oxygen blocks sites for alkane dissociative adsorption, reducing its effective sticking coefficient and with that the overall rate of reaction. Alkane activation must occur on Rh metal sites and not on oxygen or oxidized Rh atoms. 4. The surface oxygen also participates in the formation of water, which occurs in a stepwise manner, via the formation of an OH(ads) intermediate. 5. The steady state coverages of both oxygen species decreases with increasing surface temperature because of an increase in their recombination rate. This effect is more marked with the subsurface oxygen, subtler with the surface species. 6. Because of the latter, the rate of water formation in alkane partial oxidation reactions decreases with temperature, and the selectivity between H2 and H2O production shifts toward the former. Acknowledgment. Financial assistance for this research was provided by the U.S. National Science Foundation. References and Notes (1) Iordanoglou, D. I.; Bodke, A. S.; Schmidt, L. D. J. Catal. 1999, 187, 400. (2) Cavani, F.; Ballarini, N.; Cericola, A. Catal. Today 2007, 127, 113. (3) Choudhary, V. R.; Mammon, A. S.; Sansare, S. D. Angew. Chem., Int. Ed. 1992, 31, 1189. (4) Hickman, D. A.; Schmidt, L. D. Science 1993, 259, 343. (5) Enger, B. C.; Lødeng, R.; Holmen, A. Appl. Catal., A 2008, 346, 1. (6) Fathi, M.; Hofstad, K. H.; Sperle, T.; Rokstad, O. A.; Holmen, A. Catal. Today 1998, 42, 205. (7) Williams, K. A.; Leclerc, C. A.; Schmidt, L. D. AIChE J. 2005, 51, 247. (8) Roychoudhury, S.; Castaldi, M.; Lyubovsky, M.; LaPierre, R.; Ahmed, S. J. Power Sources 2005, 152, 75. (9) Neumann, D.; Veser, G. AIChE J. 2005, 51, 210. (10) Burke, N. R.; Trimm, D. L. Catal. Today 2006, 117, 248. (11) Vlachos, D. G.; Mhadeshwar, A. B.; Kaisare, N. S. Comput. Chem. Eng. 2006, 30, 1712. (12) Choudhary, T. V.; Choudhary, V. R. Angew. Chem., Int. Ed. 2008, 47, 1828. (13) Horn, R.; Williams, K. A.; Degenstein, N. J.; Schmidt, L. D. J. Catal. 2006, 242, 92. (14) York, A. P. E.; Xiao, T. C.; Green, M. L. H.; Claridge, J. B. Catal. ReV.-Sci. Eng. 2007, 49, 511. (15) Liu, T.; Snyder, C.; Veser, G. Ind. Eng. Chem. Res. 2007, 46, 9045. (16) Donazzi, A.; Beretta, A.; Groppi, G.; Forzatti, P. J. Catal. 2008, 255, 241. (17) Maestri, M.; Vlachos, D.; Beretta, A.; Forzatti, P.; Groppi, G.; Tronconi, E. Top. Catal. 2009, 52, 1983. (18) Bitsch-Larsen, A.; Horn, R.; Schmidt, L. D. Appl. Catal., A 2008, 348, 165. (19) Wilson, J. N.; Pedigo, R. A.; Zaera, F. J. Am. Chem. Soc. 2008, 130, 15796. (20) The numerical values of the temperatures in the reference above were mistakenly reported in °C, even though the units indicate K. (21) Liu, J.; Xu, M.; Nordmeyer, T.; Zaera, F. J. Phys. Chem. 1995, 99, 6167. (22) Zaera, F. Int. ReV. Phys. Chem. 2002, 21, 433. (23) Un˜ac, R. O.; Bustos, V.; Wilson, J.; Zgrablich, G.; Zaera, F. J. Chem. Phys. 2006, 125, 074705ASAP. (24) Thiel, P. A.; Williams, E. D.; Yates, J. T., Jr.; Weinberg, W. H. Surf. Sci. 1979, 84, 54. (25) Yates, J. T., Jr.; Thiel, P. A.; Weinberg, W. H. Surf. Sci. 1979, 84, 427. (26) Guevremont, J. M.; Sheldon, S.; Zaera, F. ReV. Sci. Instrum. 2000, 71, 3869. (27) Gopinath, C. S.; Zaera, F. J. Catal. 1999, 186, 387.

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(28) Gopinath, C. S.; Zaera, F. J. Catal. 2001, 200, 270. (29) Zaera, F.; Liu, J.; Xu, M. J. Chem. Phys. 1997, 106, 4204. ¨ fner, H.; Zaera, F. J. Phys. Chem. 1997, 101, 396. (30) O (31) Summers, R. L. Empirical obserVations on the sensitiVity of hot cathode ionization type Vacuum gages; NASA: Washington, DC, 1969; Report No. NASA-TN-D-5285. (32) Nakayama, K.; Hojo, H. Jpn. J. Appl. Phys. 1974, 2 (1), 113. (33) Beutler, A.; Lundgren, E.; Nyholm, R.; Andersen, J. N.; Setlik, B. J.; Heskett, D. Surf. Sci. 1998, 396, 117. (34) Fukuoka, M.; Okada, M.; Matsumoto, M.; Ogura, S.; Fukutani, K.; Kasai, T. Phys. ReV. B 2007, 75, 235434. (35) Peterlinz, K. A.; Sibener, S. J. J. Phys. Chem. 1995, 99, 2817. (36) Wilson, J.; Guo, H.; Morales, R.; Podgornov, E.; Lee, I.; Zaera, F. Phys. Chem. Chem. Phys. 2007, 9, 3830. (37) Wheeler, C.; Jhalani, A.; Klein, E. J.; Tummala, S.; Schmidt, L. D. J. Catal. 2004, 223, 191. (38) Matolin, V.; Masek, K.; Elyakhloufi, M. H.; Gillet, E. J. Catal. 1993, 143, 492. (39) CRC Handbook of Chemistry and Physics, 85th ed.; Lide, D. R., Ed.; CRC Press: Boca Raton, FL, 2004.

Wilson and Zaera (40) Zangwill, A. Physics at Surfaces; Cambridge University Press: Cambridge, 1988. (41) Thiel, P. A.; Yates, J. T., Jr.; Weinberg, W. H. Surf. Sci. 1979, 82, 22. (42) Root, T. W.; Schmidt, L. D.; Fisher, G. B. Surf. Sci. 1983, 134, 30. (43) Wehner, S.; Paffett, M. T.; Zaera, F. J. Phys. Chem. B 2004, 108, 18683. (44) Wider, J.; Greber, T.; Wetli, E.; Kreutz, T. J.; Schwaller, P.; Osterwalder, J. Surf. Sci. 1998, 417, 301. (45) Lyubovsky, M.; Roychoudhury, S.; LaPierre, R. Catal. Lett. 2005, 99, 113. (46) Dalle Nogare, D.; Degenstein, N. J.; Horn, R.; Canu, P.; Schmidt, L. D. J. Catal. 2008, 258, 131. (47) Ganduglia-Pirovano, M. V.; Reuter, K.; Scheffler, M. Phys. ReV. B 2002, 65, 2454261. (48) Liu, J.; Xu, M.; Zaera, F. Catal. Lett. 1996, 37, 9.

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