8858
J. Phys. Chem. B 1999, 103, 8858-8863
Redox Properties of Nanocrystalline Cu-Doped Cerium Oxide Studied by Isothermal Gravimetric Analysis and X-ray Photoelectron Spectroscopy Andreas Tscho1 pe,*,† Michel L. Trudeau,‡ and Jackie Y. Ying Department of Chemical Engineering, Massachusetts Institute of Technology, Cambridge, Massachusetts 02139 ReceiVed: May 14, 1999; In Final Form: August 10, 1999
Nanostructured CeO2-x and Cu0.15Ce0.85O2-x were synthesized by high-pressure magnetron sputtering. X-ray diffraction was used for phase analysis and grain size characterization. The structural analysis indicated that Cu was highly dispersed, which could be explained by surface segregation on cerium oxide nanocrystals. The reduction and oxidation of dispersed copper in nanostructured Cu0.15Ce0.85O2-x was investigated by isothermal gravimetric analysis (IGA) and X-ray photoelectron spectroscopy (XPS). In IGA, the sample weight was measured at constant temperature during in situ reduction and oxidation in 2% CO/He or 5% H2/He and 15% O2/He, respectively. The weight changes of microcrystalline CeO2 and nanocrystalline CeO2-x and Cu0.15Ce0.85O2-x samples were measured at temperatures between 200 and 500 °C. The analysis revealed that more than one oxygen atom was extracted per Cu atom in Cu0.15Ce0.85O2-x during reduction, which was reversibly restored during oxidation. The Cu valence state in Cu0.15Ce0.85O2-x after oxidation or reduction was independently determined by XPS. The Cu 2p core-level spectra and LMM Auger spectra of Cu0.15Ce0.85O2-x were compared with those of a polycrystalline copper reference. The analysis showed that all three oxidation states of copper (i.e., 0, 1+, and 2+) could be present depending on the reduction/oxidation specifications. Furthermore, Cu0 and Cu1+ were found to coexist in a stable state under certain conditions.
Introduction Nanostructured metal oxides are of great interest in heterogeneous catalysis due to their high dispersion, i.e., large fraction of atoms that are exposed on the surface of the solid crystallites. Due to their structural and chemical stability at high temperatures, such metal oxides may be employed as supports for catalytically active species, such as precious metals. The main purposes of a support material are to provide the catalyst with a mechanically stable platform, and to prevent sintering and coarsening of the expensive active phase. Various metal oxides may also be catalytically active themselves due to inherent acid/ base or redox properties. Cerium oxide has attracted great interest as a catalytic material due to its distinct defect chemistry. The propensity to release lattice oxygen to the gas phase, accompanied by the formation of oxygen vacancies, gives rise to significant deviation from stoichiometry in cerium oxide even under mildly reducing gas compositions.1-3 The ease of oxygen exchange with the gas phase provides cerium oxide with unique catalytic redox activity and gas sensor properties. Cerium oxide has been employed for the oxidation of CO by O2, 4-7 selective catalytic reduction of SO2 by CO, 6-11 and methane oxidation.6,7 Detailed studies on the catalytic activity in the first reaction revealed that surface reduction is the rate-limiting step.12 Apart from being an active catalytic material, CeO2 has also been reported as an excellent support for precious metal13 and base metal8 catalysts. For example, Cu/CeO2 has been examined for CO oxidation and SO2 reduction.6-8,10-12 In both reactions, the catalytic activity of Cu/CeO2 was found to be significantly enhanced compared † Present address: Universita ¨ t des Saarlandes, FB 10 Physik, 66041 Saarbru¨cken, Germany. ‡ Emerging Technologies, Hydro-Que ´ bec Research Institute, Varennes, Que´bec J3 1S1, Canada.
to that of pure CeO2. It was comparable even with the activity of commercial precious metal catalysts for CO oxidation.12 The high activity of Cu-doped CeO2 raises questions with regard to the atomic, chemical, and electronic structure of this material. X-ray diffraction (XRD)11 and scanning transmission electron microscopy (STEM)14 studies showed that Cu0.15Ce0.85O2-x generated by magnetron sputtering consisted of nanostructured CeO2 matrix with highly dispersed copper, and that the high copper dispersion was stable at temperatures as high as 500 °C. There is no phase diagram available for the Ce-Cu-O system to our knowledge. However, a rather low solubility of Cun+ in CeO2 might be expected due to the different ionic radii of copper and cerium. Therefore, a high driving force for segregation of Cu to the CeO2 surface exists so that most of the Cu may be present as small clusters or individual cations at the surface of cerium oxide grains. It is well-known that the oxidation state of supported transition metal species has an important effect on the activity in various catalytic reactions. For instance, the reduced Cu0 and Cu1+ species are catalytically more active than Cu2+ in NO reduction by NH3,15 methanol synthesis,16 and the dehydrogenation of 2-propanol to 2-propanone.17 For the latter reaction, there is also some evidence indicating that the coexistence of Cu0 and Cu1+ represents the most active state.17 Typical dehydrogenation catalysts, such as Cu/Cr2O3, have to be activated by hydrogen reduction. However, activation has to be performed at temperatures below 670 K to prevent sintering of Cu. Given the intimate contact between the surface-segregated Cun+ and the CeO2 support, it would be of interest to investigate the oxidation state of copper as a function of pretreatment conditions. It is conceivable that the strong chemical interaction between Cu and CeO2, which provides for the high thermal stability of Cu dispersion, may also strongly affect the redox properties of the copper catalyst.
10.1021/jp9915681 CCC: $18.00 © 1999 American Chemical Society Published on Web 09/28/1999
Nanocrystalline Cu-Doped Cerium Oxide The objective of this study is to investigate the thermochemical stability of surface-segregated copper in Cu-doped CeO2 catalysts. For this purpose, nanostructured Cu0.15Ce0.85O2-x samples were subjected to various oxidation and reduction treatments. Two complementary analytical techniques were employed: isothermal gravimetric analysis (IGA) and X-ray photoelectron spectroscopy (XPS). This combination of techniques follows two aspects of the reduction of a metal oxide, i.e., weight loss and change in cation valency due to oxygen release. Whereas isothermal gravimetric analysis could be performed in situ, XPS measurements had to be carried out separately after heat treatment of the samples in various atmospheres. These characterizations could lead to a better understanding of the redox properties of supported base metal catalysts, as well as provide guidelines for optimizing their activation conditions. Experimental Section Nanocrystalline pure CeO2-x and Cu0.15Ce0.85O2-x were synthesized by magnetron sputtering of cerium and ceriumcopper alloy targets, respectively, in an ultrahigh vacuum chamber at an argon atmosphere of 30 Pa.18 The nanometersized metallic clusters generated through gas condensation were collected on a liquid nitrogen cooled substrate. After sputtering for about 30 min, the ultrahigh vacuum chamber was evacuated and then slowly backfilled with oxygen up to a pressure of 1 kPa. This controlled oxidation allowed the generation of nonstoichiometric CeO2-x and Cu0.15Ce0.85O2-x samples. The nanocrystalline oxide particles were scrapped off the substrate and collected as a powder. Microcrystalline CeO2 and Cu reference powders were obtained from commercial sources (Alfa). All samples were compacted with a uniaxial pressure of 0.5 GPa into porous self-supporting pellets of 6 mm diameter and 1-2 mm thickness. The N2 adsorption BET surface area was determined to be ∼60 m2/g for the nanocrystalline samples, and less than 1 m2/g for the microcrystalline samples. X-ray diffraction patterns were obtained using a Rigaku rotating anode X-ray generator with Cu radiation and a Ni filter. Diffraction patterns were taken after sample compaction and after annealing at various temperatures in 15% O2/Ar atmosphere for 10 h. The volume-weighted average grain size of nanocrystalline samples was calculated from the peak broadening in the diffraction patterns. Instrumental peak broadening was taken into account by subtracting the peak width (fwhm) of a coarse-grained sample from that of the nanocrystalline sample. Strain contribution to CeO2 peak broadening was separated by Kochendorfer analysis19 of the (111) and (222) peaks. Isothermal gravimetric analysis was performed using a PerkinElmer TGA 7. The nanocrystalline and microcrystalline pure cerium oxide samples were first annealed at 550 °C in 15% O2/He for 20 h before further analysis. This annealing step was omitted in the case of Cu-doped CeO2 to avoid irreversible structural changes. All TGA measurements involved sequential reduction and reoxidation under isothermal conditions. The samples were first equilibrated at a constant temperature in 15%O2/He for several hours. Then, the purge gas was changed to either 2% CO/He or 5% H2/He. The weight change was measured as a function of time for a period of 500 min. The purge gas was then switched back to 15% O2/He, and the sample weight was monitored for another 500 min. The oxidation state of Cu was determined by X-ray photoelectron spectroscopy using a Perkin-Elmer PHI-5500 ESCA system with monochromatic Al KR radiation. High-resolution spectra were obtained using a pass energy of 11.75 eV with a
J. Phys. Chem. B, Vol. 103, No. 42, 1999 8859
Figure 1. X-ray diffraction patterns of nanocrystalline Cu0.15Ce0.85O2-x after annealing for 10 h in 15% O2/Ar at the temperatures indicated.
step size of 0.1 eV. The insulating nature of CeO2 resulted in charging of the sample during photoemission, which was compensated by an electron flood gun. The flood gun current was set to minimize the peak width in the photoelectron spectra. The energy scale was calibrated using the C-(C,H) components of the C 1s band at 284.8 eV and the highest binding energy peak for Ce4+ d3/2 at 916.5 eV, giving an uncertainty in the peak positions of about 0.2 eV. The oxidation or reduction of the samples prior to XPS measurements was performed in an ultrahigh vacuum chamber directly connected to the XPS system. In this reaction chamber, reduction in 8% H2/He or oxidation in 1% CO2/He was performed under a total pressure of 30 Pa. After heat treatment for 15 h in the controlled atmospheres, the samples were rapidly cooled to room temperature and immediately analyzed in the XPS chamber under an ultrahigh vacuum (10-7 Pa base pressure). The CuO reference spectrum was obtained from a microcrystalline copper sample that had been oxidized at 200 °C in air in an external tube furnace. The oxidation states of Cu in the various samples were determined from the Cu 2p core-level signals and the Cu LMM Auger signals. The XPS spectra were curve-fitted with the Perkin-Elmer software assuming mixed Gaussian-Lorentzian functions for the peak profiles. The ratio of different copper oxidation states was determined from the relative peak intensities after profile fitting. Results X-ray Diffraction. X-ray diffraction patterns of a nanocrystalline Cu0.15Ce0.85O2-x sample annealed in 15% O2/Ar at various temperatures are shown in Figure 1. The diffraction peaks found after sample annealing at 500 °C were all assigned to the fcc cerium sublattice of the CeO2 fluorite structure. The volumeweighted average crystal size was calculated to be 7 nm. CuO emerged as a second phase with increasing peak intensity in samples annealed at temperatures above 500 °C. The diffraction peaks corresponding to CeO2 became sharper after sample annealing above 500 °C due to grain growth of the cerium oxide phase. Isothermal Gravimetric Analysis. The relative weight changes in nanocrystalline CeO2-x and Cu0.15Ce0.85O2-x during reduction in 2% CO/He and reoxidation in 15% O2/He at 200 °C are shown in Figure 2. The measurements were performed with the same sample up to five times and were found to be completely reversible. When CO was introduced, two processes with different time constants were observed: (i) rapid CO adsorption (small weight gain) and (ii) slow reduction of the sample (weight loss). Sample oxidation and CO2 desorption (as
8860 J. Phys. Chem. B, Vol. 103, No. 42, 1999
Figure 2. Relative weight of (a) nanocrystalline CeO2-x and (b) nanocrystalline Cu0.15Ce0.85O2-x in an isothermal thermogravimetric study at 200 °C. The gas composition was 15% O2/He in regions I and III, and 2% CO/He in region II.20
Figure 3. Relative weight change (∆, as defined in Figure 2) as a function of temperature for (a) microcrystalline CeO2, (b) nanocrystalline CeO2-x, and (c) nanocrystalline Cu0.15Ce0.85O2-x.20
confirmed by mass spectroscopy) occurred when oxygen was introduced to the system. The signal of the adsorption/desorption process was observed at low temperatures only and was absent at temperatures g400 °C. This relative weight change ∆ is plotted as a function of redox reaction temperature for different samples in Figure 3. For the microcrystalline CeO2 sample, detectable weight changes were only noted at temperatures g450 °C (Figure 3a), indicating that bulk reduction in cerium oxide required such high temperatures. After being preannealed at 550 °C, nanocrystalline CeO2-x with a specific surface area of > 40 m2/g exhibited greater weight change (Figure 3b) than the microcrystalline CeO2 with low surface area. Nanocrystalline Cu0.15Ce0.85O2-x (Figure 3c) showed a weight change that was an order of magnitude larger than pure nanocrystalline CeO2-x at temperatures below 500 °C. To investigate whether the reduction process at 200 °C was specific for CO as the reducing gas, similar measurements were performed using 5% H2/He. The relative weight changes for nanocrystalline Cu0.15Ce0.85O2-x at 200 °C were 1.83 and 2.14 wt % in 2% CO/He and 5% H2/He, respectively. X-ray Photoelectron Spectroscopy. X-ray photoelectron reference spectra were obtained from microcrystalline Cu after various oxidation/reduction treatments. As described in the Experimental Section, oxidation in air was performed in an external furnace. However, reduction in 8% H2/He could be achieved in the reaction chamber connected to the XPS system, so that the reduced samples could be analyzed by XPS without exposure to ambient atmosphere. The Cu 2p spectra and Cu L3M4,5M4,5 Auger spectra obtained are shown in Figures 4 and 5, respectively. The Cu2+ oxidation state can be easily identified
Tscho¨pe et al.
Figure 4. Cu 2p photoelectron spectra of microcrystalline Cu (a) after oxidation in air at 200 °C, and after reduction in 30 Pa of 8% H2/He at (b) 200 °C and (c) 400 °C for 15 h.
Figure 5. Cu LMM Auger spectra of microcrystalline Cu (a) after oxidation in air at 200 °C, and after reduction in 30 Pa of 8% H2/He at (b) 200 °C and (c) 400 °C for 15 h.
by the shake-up structure at 940-945 eV in Figure 4a. In the most oxidized state, the relative intensity of the satellite with respect to the main line at 934.2 eV was determined to be 0.55, in agreement with literature data.21 Upon reduction at 200 °C, the satellite disappeared completely, and the main line became sharper and was shifted to 932.6 eV (Figure 4b). Further reduction at 400 °C did not change the Cu 2p spectrum significantly (Figure 4c). The Cu LMM Auger spectra also exhibited clear differences for the sample after oxidation (Figure 5a) and reduction (Figure 5b) at 200 °C. The position of the main peak shifted from 917.2 to 915.8 eV. In contrast to Cu 2p spectra, Auger spectra showed further changes upon reduction at 400 °C (Figure 5c). The position of the main line was shifted to 918.6 eV, and a plasmon satellite structure was resolved. The Auger spectra in Figures 5b,c were used as reference spectra for distinguishing the Cu+ and Cu0 oxidation states. The Cu 2p and Ce 3d core-level spectra of nanocrystalline Cu0.15Ce0.85O2-x are shown in Figure 6. The shake-up structure in Cu 2p was found after oxidation at 200 °C (Figure 6a), but disappeared after reduction at 200 °C (Figure 6b). The ratio of integrated intensity of the satellite with respect to the main line was 0.35 for the sample after oxidation at 200 °C, and increased at higher oxidation temperatures. The spectral features at 880920 eV correspond to Ce 3d core levels. The six peaks indicated by solid dots in Figure 6 are associated with the four-valent Ce 4f0 initial state of CeO2.22-26 Additional lines in the cerium core levels are characteristic of the Ce 4f1 initial state and correspond to trivalent cerium.21,27 Further quantitative analysis of the Ce oxidation state was not part of this study, but has been previously investigated for nanocrystalline La-doped CeO2.27 The Auger spectra of the Cu0.15Ce0.85O2-x sample after various redox
Nanocrystalline Cu-Doped Cerium Oxide
J. Phys. Chem. B, Vol. 103, No. 42, 1999 8861
TABLE 1: Oxidation States of Cun+ after Oxidation or Reduction Treatments at Different Temperatures, As Determined from XPS Cu 2p Core-Level Spectra and LMM Auger Spectra 200 °C
300 °C
400 °C
microcrystalline Cu reduced in 30 Pa of 8% H2/He
100% Cu1+
100% Cu0
nanocrystalline Cu0.15Ce0.85O2-x oxidized in 30 Pa of 1% CO2/He
65% Cu2+ 35% Cu1+ 100% Cu1+
45% Cu1+ 55% Cu0 65% Cu2+ 35% Cu1+ 70% Cu1+ 30% Cu0
nanocrystalline Cu0.15Ce0.85O2-x reduced in 30 Pa of 8% H2/He
Figure 6. Cu 2p and Ce 3d photoelectron spectra of nanocrystalline Cu0.15Ce0.85O2-x (a) after oxidation in 30 Pa of 1% CO2/He at 200 °C and (b) after reduction in 30 Pa of 8% H2/He at 200 °C. The solid dots indicate the position of the six core levels of the four-valent Ce 4f0 initial state.
Figure 7. Cu LMM Auger spectra of nanocrystalline Cu0.15Ce0.85O2-x (a) after oxidation in 30 Pa of 1% CO2/He at 200 °C, and after reduction in 30 Pa of 8% H2/He at (b) 200 °C and (c) 400 °C.
treatments are shown in Figure 7. The changes in the Auger spectra are similar to those found in the reference material. Reduction of the sample at 200 °C resulted in a shift of the main line from 917.9 to 915.9 eV. After reduction at 400 °C, the main peak was shifted to 918.1 eV (Figure 7c). Determination of the Cu oxidation state was performed on the assumption that Cu2+ and Cu1+ could be distinguished by the presence of the shake-up peak in Cu 2p spectra. The intensity ratio of this shake-up feature to the main line is 0.55 for pure Cu2+ and 0 for pure Cu1+. Cu1+ and Cu0 oxidation states were distinguished by the LMM Auger spectra. The various spectra were normalized, and intensity ratios were determined for semiquantitative analysis. The results of this analysis are summarized in Table 1. Discussion Thermochemical stability of the macroscopic phases CuO, Cu2O, and Cu is well characterized by their Gibbs free energy
80% Cu2+ 20% Cu1+ 100% Cu0
of formation, and the corresponding data can be found in standard tables.28 In contrast, redox properties of Cu in nanostructured Cu0.15Ce0.85O2-x can be expected to be affected by the particular microstructure of this material. Nanocrystalline Cu0.15Ce0.85O2-x exhibited only the diffraction peaks corresponding to CeO2 fluorite structure after being annealed at e500 °C. Annealing at temperatures above 500 °C resulted in the precipitation of CuO as a second phase. CuO phase precipitation at high temperatures supported the assumption of a very low Cun+ solubility in the CeO2 lattice. Hence, the single-phase XRD pattern of nanocrystalline Cu0.15Ce0.85O2-x indicated a high dispersion of Cu that was thermally stable at temperatures e500 °C. Recent measurements of the reversible Nernst voltage of a solid-state electrochemical cell consisting of a coarse-grained CuO/Cu2O reference electrode, CuBr solid electrolyte, and nanocrystalline CuxCe1-xO2-y samples revealed that the chemical activity ratio aCuO/aCu2O was as low as 0.2 for compositions of x < 0.1 and increased to 0.75 for 0.1 < x < 0.3.29 The low solubility of Cu in bulk CeO2 on one hand and the low chemical activity (i.e., apparently high solubility) of CuO as measured in nanostructured CuxCe1-xO2-y on the other hand suggest a metastable segregation of Cu at the surfaces of CeO2 nanocrystals. According to this model, surface-segregated Cun+ in high dispersion is thermodynamically favorable compared to the precipitation of a new phase, provided that the specific surface area of the CeO2 matrix is large enough to ensure submonolayer coverage. Only upon grain growth of the CeO2 matrix with a corresponding loss of specific surface area would the precipitation of CuO become necessary. A high dispersion of surface-segregated copper in CeO2 is also confirmed by electron paramagnetic resonance (EPR) studies.30 Analysis of the EPR spectra of Cu0.01Ce0.99O2-y indicates the presence of single Cu2+ ions and Cu2+ ion pairs at the cerium oxide surface. The purpose of the present study was to investigate the reduction and oxidation properties of Cu, which was segregated at the surface of CeO2 nanocrystals. Isothermal gravimetric analysis of the weight change upon reduction and oxidation of the various samples was used as one of the analytical methods. When pure nanocrystalline CeO2 after equilibration in 15% O2/ He was exposed to carbon monoxide, CO adsorbed on it as surface carbonates.31 If these carbonates desorb as CO2, the cerium oxide surface is left in a reduced state. As demonstrated by temperature-programmed reduction (TPR) and CO oxidation catalytic studies,12 such surface reduction is readily observed at temperatures above 100 °C. The time-dependent weight change in nanocrystalline CeO2-x after exposure to CO (Figure 2a) could be explained by assuming (i) adsorption of CO reaching a steady-state coverage and (ii) surface reduction. When the sample was exposed again to 15% O2/He, the surface was rapidly oxidized and desorption of the surface carbonates occurred. High thermal stability of the bidentate and monodentate carbonates on CeO231 caused desorption to be slow as compared to surface oxidation. Furthermore, the sample weight at the end of region II, i.e., just before switching to 15% O2/ He, included the mass of adsorbed gases, which slowly desorbed
8862 J. Phys. Chem. B, Vol. 103, No. 42, 1999 in region III. Therefore, the rapid weight gain (indicated by ∆) upon exposure to 15% O2/He characterized the oxidation of the sample. The weight gain ∆ may be attributed to the oxidation of the oxide surface and to the oxidation of adsorbed gases. An upper value for the latter contribution can be estimated. Oxidation of the adsorbed CO to CO2 is accompanied by a welldefined weight change reflecting the difference in molecular weight of the two molecules. The amount of oxidized CO2 can be obtained from the slow weight loss during CO2 desorption in region III of Figure 2. Thirty-six percent of this relative weight change would be expected during the oxidation of the corresponding amount of CO. However, the majority of adsorbed CO can be expected to be present as surface carbonates with no further oxidation necessary prior to desorption. Therefore, the given estimate can be regarded as an upper bound. In comparison to nanocrystalline CeO2-x, the nanocrystalline Cu0.15Ce0.85O2-x sample has a much larger weight change. The results were also rather similar for CO (1.83 wt %) or H2 (2.14 wt %) reduction, indicating that the choice of reducing gas did not matter significantly in terms of reaction kinetics. The different order of magnitude of the reduction/oxidation effect on sample weight was also observed at higher temperatures (Figure 3). For microcrystalline CeO2, a noticeable weight change could only be measured at temperatures above 450 °C. This is assigned to bulk reduction of CeO2, as consistent with TPR measurements.12 The nanocrystalline CeO2 sample exhibited much larger weight changes. The effect of high specific surface area on reduction may be two-fold: (i) bulk reduction may be observed at a higher rate due to enhanced diffusion through the interface, and (ii) the cerium oxide surface itself may exhibit much greater propensity to reduction as compared to the bulk phase. Enhanced surface reduction at lower temperatures as compared to bulk reduction was also observed in TPR studies.12 Compared to pure nanocrystalline CeO2-x, the measured relative weight changes for the nanocrystalline Cu0.15Ce0.85O2-x were much larger at all temperatures. For further analysis, it may be useful to estimate the relative weight change of a sample of given Cu/Ce ratio, assuming that all copper was present as CuO in 15% O2/He and reduced to Cu in 2% CO/He, and that reduction of cerium oxide in the mixed oxide was similar to that of pure nanocrystalline CeO2-x. With this simplification, the total relative weight change ∆M/M can be estimated from a rule of mixing
∆M ) M
∑∆mi ) ∑xi‚Wi‚ri ∑mi ∑xi‚Wi
The total absolute weight change ∆M and total weight M are given by the sum of ∆mi and sum of mi, respectively, i.e., the weight change and weight of each contributing oxide. These are related to the molar concentration xi, the molar weight Wi, and the relative weight change ri of each oxide component i. With this assumption, the expected total weight change of nanocrystalline Cu0.15Ce0.85O2-x is calculated to be 1.65 wt %, which is smaller than the measured values of 1.83 and 2.14 wt % with CO and H2 as reducing gases, respectively. On the basis of this comparison, it can be concluded that more than one oxygen atom is removed per copper atom upon reduction and restored upon oxidation at all temperatures between 200 and 500 °C. It should be emphasized that CuO and Cu were used in this model to characterize the oxidation state of Cu in terms of the O/Cu ratio. These oxidation states should not be associated with the presence of CuO and Cu as bulk phases. In the
Tscho¨pe et al. calculation, it was assumed that reduction of the cerium oxide support was not affected by the presence of Cu. However, the fact that the measured weight changes were larger than the calculated value suggests that Cu does promote the reduction of cerium oxide. One may consider various origins of such a promoting effect. The presence of adsorbed metal ions could have an electronic effect, such as a change in surface electrical potential, which would lead to a corresponding change in surface defect formation energies. It is also possible that exposed transition metal ions or clusters represent preferred adsorption sites for CO, which might diffuse onto or “spill over” the cerium oxide surface, thus establishing a new reaction pathway of lower activation energy for cerium oxide reduction. IGA could be used to analyze the reduction and oxidation of metal oxides since these reactions are accompanied by a considerable weight change. However, a more complete picture may be obtained by investigating the valency of the cations using a technique such as XPS. XPS measurements on microcrystalline CuO were used to obtain reference spectra for the various copper oxidation states. With these reference spectra, the oxidation states of Cun+ in nanocrystalline Cu0.15Ce0.85O2-x after various reduction and oxidation treatments were determined (Table 1). Oxidation in 30 Pa of 1% CO2/He resulted in a mixture of Cu2+ and Cu1+, with 80% Cu2+ after oxidation at 400 °C. Upon reduction at 200 °C in 30 Pa of 8% H2/He, only Cu1+ was found in the nanostructured Cu0.15Ce0.85O2-x sample, as for the microcrystalline copper sample. Reduction to Cu0 was incomplete at 300 °C for both nanocrystalline Cu0.15Ce0.85O2-x and microcrystalline copper. This coexistence of Cu1+ and Cu0 was observed even after another 15 h of reduction at this temperature. Complete reduction to zerovalent Cu0 was achieved at a temperature of 400 °C. XPS illustrates that the highly dispersed copper species in nanocrystalline Cu0.15Ce0.85O2-x can adopt any of the three oxidation states, 0, 1+, and 2+, depending on the environment and treatment history of the sample. The presence of pure Cu1+ after reduction at 200 °C also showed that reduction from Cu2+ to Cu0 did not occur in a single step as reported in the literature.32 Furthermore, it is possible for Cu0 and Cu1+ to coexist under certain thermodynamic conditions (i.e., temperature and oxygen partial pressure). In summary, the measured weight change during isothermal gravimetric analysis and the changes in cation valency as determined by X-ray photoelectron spectroscopy indicate that the highly dispersed copper in nanocrystalline Cu0.15Ce0.85O2-x can be readily reduced and oxidized at temperatures as low as 200 °C. The weight change measured in the temperature range of 200-500 °C corresponds to an exchange of more than one oxygen atom per copper atom in the sample during an oxidation/ reduction cycle. In XPS measurements, copper valencies of 0, 1+, and 2+ were found depending on the reduction and oxidation conditions. It should be noted that oxygen exchange and valency change are not necessarily related as they are in the reduction of bulk CuO since (i) oxygen, which is released during reduction, is not necessarily present in a fully ionized state as O2- and (ii) electronic interaction between dispersed copper and the cerium oxide matrix or other adsorbed species may exist. The combined measurement of oxygen exchange and valency change as a function of the relevant thermodynamic parameters (i.e., temperature and oxygen chemical potential) should yield interesting information on the surface chemistry of such mixed metal oxides. Acknowledgment. This work was supported by the National Science Foundation (Grant CTS-9257223) and Hydro-Que´bec
Nanocrystalline Cu-Doped Cerium Oxide Research Institute. A.T. acknowledges fellowship support from the German National Scholarship Foundation (BASF program). References and Notes (1) Bevan, D. J.; Kordis, J. J. Inorg. Nucl. Chem. 1964, 26, 1509. (2) Iwasaki, B.; Katsura, T. Bull. Chem. Soc. Jpn. 1971, 44, 1297. (3) Sørensen, O. T. J. Solid State Chem. 1976, 18, 217. (4) Breysse, M.; Guenin, M.; Claudel, B.; Latreille, H.; Ve´ron, J. J. Catal. 1972, 27, 275. (5) Breysse, M.; Guenin, M.; Claudel, B.; Ve´ron, J. J. Catal. 1973, 28, 54. (6) Liu, W.; Flytzani-Stephanopoulos, M. J. Catal. 1995, 153, 304. (7) Tscho¨pe, A.; Liu, W.; Flytzani-Stephanopoulos, M.; Ying, J. Y. J. Catal. 1995, 157, 42. (8) Liu, W.; Flytzani-Stephanopoulos, M. In EnVironmental Catalysis; Armor, J. N., Ed.; ACS Symposium Series 552; American Chemical Society: Washington, DC, 1994; p 375. (9) Tscho¨pe, A.; Ying, J. Y. In Nanophase Materials: SynthesisProperties-Applications; Hadjipanayis, G. C., Siegel, R. W., Eds.; Kluwer: Dordrecht, The Netherlands, 1994; p 781. (10) Liu, W.; Sarofim, A. F.; Flytzani-Stephanopoulos, M. Appl. Catal. B 1994, 4, 167. (11) Tscho¨pe, A.; Ying, J. Y.; Liu, W.; Flytzani-Stephanopoulos, M. In Materials and Processes for EnVironmental Protection; Voss, K. E., Quick, L. M., Gadgil, P. N., Adkins, C. L. J., Eds.; MRS Symposium Proceedings 344; Materials Research Society: Pittsburgh, 1994; p 133. (12) Tscho¨pe, A.; Schaadt, D.; Birringer, R.; Ying, J. Y. Nanostr. Mater. 1997, 9, 423. (13) Yao, H. C.; Yu Yao, Y. F. J. Catal. 1984, 86, 254. (14) Tscho¨pe, A.; Ying, J. Y.; Chiang, Y.-M. Mater. Sci. Eng. A 1995, 204, 267. (15) Otto, K.; Shelef, M. J. Phys. Chem. 1972, 76, 37.
J. Phys. Chem. B, Vol. 103, No. 42, 1999 8863 (16) Chinchen, G. C.; Denny, P. J.; Jennings, J. R.; Spencer, M. S.; Waugh, K. Appl. Catal. 1988, 36, 1. (17) Cunningham, J.; Al-Sayyed, G. H.; Cronin, J. A.; Fierro, J.; Healy, C.; Hirschwald, W.; Ilyas, M.; Tobin, J. J. Catal. 1986, 102, 160. (18) Tscho¨pe, A.; Ying, J. Y. Nanostr. Mater. 1994, 4, 617. (19) Kochendo¨rfer, A. Z. Kristallogr., Mineral. Petrogr. 1944, 105, 393. (20) Tscho¨pe, A.; Ying, J. Y.; Amonlirdviman, K.; Trudeau, M. L. In Molecularly Designed Ultrafine/Nanostructured Materials; Gonsalves, K. E., Chow, G.-M., Xiao, T. D., Cammarata, R. C., Eds.; MRS Symposium Proceedings 351; Materials Research Society: Pittsburgh, 1994; p 251. (21) Ghijsen, J.; Tjeng, L. H.; van Elp, J.; Eskes, H.; Westerink, J.; Sawatzky, G. A.; Czyzyk, M. T. Phys. ReV. B 1988, 38, 11322. (22) Praline, G.; Koel, B. E.; Hance, R. L.; Lee, H. I.; White, J. M. J. Electron. Spectrosc. Relat. Phenom. 1980, 21, 17. (23) Fujimori, A. Phys. ReV. B 1983, 28, 2281. (24) Jin, T.; Zhou, Y.; Mains, G. J.; White, J. M. J. Phys. Chem. 1987, 91, 5931. (25) Le Normand, F.; Hilaire, L.; Kili, K.; Krill, G.; Maire, G. J. Phys. Chem. 1988, 92, 2561. (26) Le Normand, F.; El Fallah, J.; Hilaire, L.; Le´gare´, P.; Kotani, A.; Parlebas, P. C. Solid State Commun. 1989, 71, 885. (27) Trudeau, M. L.; Tscho¨pe, A.; Ying, J. Y. Surf. Interface Anal. 1995, 23, 219. (28) Lide, D. R. Handbook of Chemistry and Physics, 73rd ed.; CRC: Boca Raton, FL, 1992. (29) Knauth, P.; Schwitzgebel, G.; Tscho¨pe, A.; Villain, S. J. Solid State Chem. 1998, 140, 295. (30) Abou Kaı¨s, A.; Bennani, A.; Aı¨ssi, C. F.; Wrobel, G.; Guelton, M. J. Chem. Soc., Faraday Trans. 1992, 88, 1321. (31) Li, C.; Domen, K.; Maruya, K.; Onishi, T. J. Chem. Soc., Faraday Trans. 1989, 85, 929. (32) Himelfarb, P. B.; Wawner, F. E., Jr.; Bieser, A., Jr.; Vives, S. N. J. Catal. 1983, 83, 469.
