Langmuir 1999, 15, 5913-5920
5913
Temperature Programmed Desorption Characterization of Oxidized Uranium Surfaces: Relation to Some Gas-Uranium Reactions† A. Danon,‡ J. E. Koresh,‡ and M. H. Mintz*,‡,§ Nuclear Research CentersNegev, P.O. Box 9001, Beer-Sheva, Israel, and Department of Nuclear Engineering, The Ben-Gurion University of the Negev, P.O. Box 653, Beer-Sheva, Israel Received September 9, 1998. In Final Form: February 3, 1999 The chemisorption characteristics and surface composition of oxidation overlayers developing on metals when exposed to oxidizing atmospheres are important in determining the protective ability of these layers against certain gas-phase reactions (e.g., corrosion and hydriding). In the present study, a special setup of supersonic molecular beam-temperature-programmed desorption was utilized to determine the different chemisorbed species present on oxidized uranium surfaces. The main identified species included water (in different binding forms) and hydrogen. The latter hydrogen originates from the water-uranium oxidation reaction, which produces uranium dioxide and two types of hydrogen: a near surface hydride and a surfacechemisorbed form that desorbs at a lower temperature than that of the hydride. Assignments of the different water desorption peaks to different binding sites were proposed. In general, four water desorption features were identified (labeled W0, W1, W2, and W3, respectively, in the order of increasing desorption temperatures). These features correspond to a reversibly chemisorbed molecular form (W0), a more tightly bound water (chemisorbed on different type of oxide sites) or hydroxyl clusters (W1), and strongly bounded (possibly isolated) hydroxyl groups (W2). The highest temperature peak (W3) is related to the formation of complex water-carbo-oxy compounds and is present only on oxidation overlayers, which contain proper chemisorbed carbo-oxy species. The relation of the water and hydrogen thermal release behavior to some problems addressed to certain effects observed in hydrogen-uranium and water-uranium reactions is discussed. For the latter, a microscopic mechanism is proposed.
1. Introduction Metallic uranium is commonly used as a nuclear fuel component in low-temperature reactors. Its chemical reactivity toward corroding gases (especially water, oxygen, and hydrogen) is of great concern regarding its storage conditions and its in-pile performance. Extensive studies on the hydriding reaction of uranium1 and its oxidation behavior in air or under water vapor exposures2 have been reported in the literature. It has been realized that chemisorption of water on the oxide layer which coats the metal surface plays an important role in determining the kinetics and morphological characteristics of those gas-phase corrosion reactions. Such water chemisorption may significantly affect the coadsorption of other species, e.g., that of hydrogen, as well as modify the transport properties of the reacting species through the oxide layer. For example, it has been pointed out3,4 that vacuum heat pretreatments (performed in the range of about 100-250 °C) drastically increased the reactivity of uranium toward hydrogen, shortening the induction period for the initiation of the reaction, and producing a denser hydride nucleation on the reacting † Presented at the Third International Symposium on Effects of Surface Heterogeneity in Adsorption and Catalysis on Solids, Poland; August 9-16, 1998. ‡ Nuclear Research CentersNegev. § The Ben-Gurion University of the Negev.
(1) For a review, see: Bloch, J.; Mintz, M. H. in Actinide Processing: Methods and Materials; Proceedings of the International Symposium; Mishra, B., Averill, W. A.; TMS: San Francisco, 1994; pp 121-131. (2) For reviews, see: Colmenares, C. A. Prog. Solid State Chem. 1975, 9, 129-239; 1984, 15, 257. (3) Bloch, J.; Swissa, E.; Mintz, M. H. Z. Phys. Chem. NF 1989, 164, 1193. (4) Bloch, J.; Mintz, M. H. J. Less-Common Met. 1990, 166, 241.
metal surface. These heat pretreatment effects were attributed to the desorption of chemisorbed water or surface hydroxyl groups from the oxide surface layer, producing active sites for hydrogen dissociative chemisorption on this dehydrated surface. The deleterious effect of chemisorbed hydroxyls on the low-temperature chemisorption of hydrogen has also been referred to in conjunction to the catalytic activity of some rare-earth oxides.5 As for the oxidation reactions of uranium, a complex behavior which depends on a large set of parameters is reported in the literature.2 These parameters include the partial pressure of water, oxygen presence (oxygen has a strong inhibition effect on the water-uranium reaction), temperature, and sample surface characteristics. No global model which accounts for all these effects has been proposed so far. However, it has been postulated that the chemisorption characteristics of water on the formed surface oxide layers, as well as the transport properties of that oxide, are the dominant factors in controlling the oxidation behavior.2,6 An attempt has been made to characterize some oxidation layers formed on uranium under different preparation procedures. In that study,6 nitrogen sorption analysis, scanning electron microscopy (SEM) observations, and diffuse reflectance infrared spectroscopy were utilized. Most of the analysis however was focused on the high surface area UO2 powder produced by exposing a highly dispersed uranium powder (formed by hydriding-dehydriding cycles of uranium) to water vapor at room temperature. (5) Netzer, F. P.; Bertel, E. Adsorption and Catalysis on Rare Earth Surfaces. In Handbook on the Physics and Chemistry of Rare Earths; Gschneidner, K. A., Jr.; Eyring, L. North-Holland Publ. Co.: Amsterdam, 1982; pp 217-320. (6) Fuller, E. L., Jr.; Smyrl, N. R.; Condon, J. B.; Eager, M. H. J. Nucl. Mater. 1984, 120, 174.
10.1021/la981210g CCC: $18.00 © 1999 American Chemical Society Published on Web 05/13/1999
5914 Langmuir, Vol. 15, No. 18, 1999
Danon et al.
Figure 1. SEM photographs of a heavily oxidized sample (LO1). (A) Top view, lower magnification. (B) Top view, higher magnification. (C) Cross-section-view, lower magnification. (D) Cross-sectional view, higher magnification.
In the present work, atmospheric-pressure temperatureprogrammed desorption (TPD) was applied to characterize the oxidation layers formed on air- and water-exposed uranium surfaces. The thermal release of both molecular water and hydrogen was monitored, utilizing a special setup of a supersonic molecular beam and TPD (SMB-TPD), as described recently.7 Also, X-ray diffraction (XRD) line shape analysis was applied to determine the types of oxides and estimate the crystallite size of the formed products. Morphological characterization of the oxide layers was performed by SEM observations. 2. Experimental Section The ability to probe small quantities of molecular H2O desorbing from oxidized metallic surfaces during a programmed temperature scan is not a straightforward task, especially when ambient-conditions characterization is required. Thus, for example, in conventional high-vacuum TPD setups, the weakly bound water existing initially on the surface under the ambient atmosphere may desorb completely during the time interval required to pump the vacuum chamber down to the desired working conditions. In addition, condensation of water makes such measurements quite complex. A recent TPD setup, which incorporates a high-pressure cell connected to a differentially pumped supersonic molecular beam arrangement probed by a mass spectrometer (SMB-TPD), was proposed to provide such characterization capability.7 The details of this SMB-TPD equipment were presented in Danon et al.7 and will not be discussed further in this section. The oxidized uranium samples were of two types: (i) heavily oxidized samples which stayed under air for several years, attaining a black, sometimes flaking oxide layer (about 10-15 µm thick), henceforth denoted as LO (long oxidation) and (ii) polished samples which were slightly oxidized either in air or under various vapor pressures at room temperature, denoted as SO (short oxidation). The above samples of pure metallic uranium (99.9%) were taken from two sources: an extruded wire (2 mm in diameter) (7) Danon, A.; Avraham, I.; Koresh, J. E. Rev. Sci. Instrum. 1997, 68, 4359.
denoted as source 1 and parallelepipeds 3 × 3 × 30 mm, denoted as source 2. The main difference between these two sources (besides their geometrical shape) was in their hydrogen impurity content (of the bulk metal), as described in section 3. Hence, the matrix of all types of examined samples contained LO1, LO2, SO1, and SO2 (where the notation LO1 for example indicates a heavily oxidized sample of source 1, etc.). Figure 1 illustrates the morphological characteristics of the heavily oxidized samples (in this figure, the LO1 sample is presented). The SEM top-view photographs (parts A and B) reveal a nonadherent, flaking oxide (Figure 1A, lower magnification) with a porous microstructure (Figure 1B, higher magnification). The oxide is composed of small (ca. 50-70 nm) agglomerated spherical particles with a typical powderlike structure. These observations closely agree with the reported specific surface areas of uranium oxidation products6 (ca. 6-10 m2/g) with an average particle sizes of about 50-90 nm. It is worthwhile to mention that XRD line shape analysis of these oxide layers indicated small UO2 crystallites of about 7-9 nm. Hence, it seems that each of the small agglomerated particles is composed of a few adherent crystallites. The cross-section view of the oxide is presented in Figure 1 parts C (lower magnification) and D (higher magnification). These observations were made by mounting the oxidized sample in a plastic matrix and polishing it over the cross-section side. It is seen that the polished oxide surfaces lost their microstructural details described above due to the “smearing” action of the polish. However, these structure details can still be observed in the oxide particles located beneath the polished plane (Figure 1D). The oxide layer coating the heavily oxidized samples seem to be nonuniform with thicknesses ranging between 5 and 20 µm. The average thickness is estimated to be about 10 µm. The oxide is cracked and particulated and no continuous adherent zone is apparent within the resolution limit of these observations (ca. 100 nm at the metal-oxide boundary). Yet, as discussed later (section 4), a thin protective oxide layer does exist at the oxidemetal interface region and controls the kinetics of some gasuranium reactions. Nevertheless, the most significant contribution to the TPD desorption spectra (presented in the next section) originates from the outer porous and cracked (∼10 µm thick) oxide zone, which has a typical powderlike microstructure. This accounts for the broad TPD line widths described later on and for the estimated amounts of total water content. It should
TPD Characterization of Uranium Surfaces
Langmuir, Vol. 15, No. 18, 1999 5915
be realized that a ∼10 µm thick oxide with the above-stated 50-70 nm particles size, produces a surface area of (0.5-1) × 103 cm2/cm2 geometrical area of the sample. Hence, a water content on the order of 1018 molecules/cm2 geometrical area is in fact related to a surface coverage (of the oxide particles) of only a few monolayers. The working procedure of the TPD experiments was basically as described in Danon et al.7 The sample was placed in a 1/4 in. o.d. silicosteel tube. He at a pressure about 800 Torr and a flow rate of 15 cm3/min passed through the sample tube during sample introduction and after the insertion of the sample for exactly 10 min prior to the heating. It is important to note that the He flow removed some of the water from the sample and thereby influenced the TPD signal, especially the W0 peak. The helium removal effect can be seen in the water TPD spectra by the high water signal at the beginning of the TPD run. Accurate measurement of the time between the insertion of the sample to TPD cell and the startup of the heating was thus important in obtaining reproducible results, especially those of the W0 peak. The sample temperature was controlled and measured with a chromel-alumel thermocouple connected on the outside of the silicosteel tube. The sample temperature was about 4-7 °C lower than the measured oven temperature. During the experiment, the column and the nozzle, used to direct the gas flow and produce the molecular beam, were heated to 180 °C to avoid adsorption of water. The 10 min He flow prior the heating was done in order to evacuate the entrance of the water vapors during the sample introduction. The water containment was calibrated by decomposing weighted amounts of CuSO4‚5H2O, and it included the water desorbed prior to the heating.
3. TPD Results and Peaks Assignment 3.1. Heavily Oxidized LO Samples. Figure 2A presents typical TPD spectra of the water (mass 18) and of the concurrent hydrogen (mass 2) desorption, obtained for the LO1 type samples. Figure 2B presents the same spectra obtained for the LO2 type specimens. It is evident that several differences exist between the two types of heavily oxidized samples. 3.1.1. Water Desorption Peaks. Concerning the water desorption spectra (solid lines), even though the temperature range where desorption takes place is similar for the two types of samples (e.g., for the heating rate applied in Figure 2, it extends between about 30 and 400 °C), the TPD line shapes seem to be somewhat different. A relatively resolved spectrum is apparent for the LO1 case (Figure 2A) with at least three (sometimes four) peaks (some of which overlap) labeled in the figure as W0, W1, W2, and W3. For the LO2 samples, the W1 peak is broader, overlapping the W2 peak. This W2 peak seems to be more intense for the LO2 samples than for the LO1 samples and appears in Figure 2B as a broad shoulder on the hightemperature side of the W1 peak. It should be noted that in some of the LO1 samples this W2 peak appeared more distinctly than it did in Figure 2A but still with an intensity lower than that of the LO2 samples. In addition to the difference between the W2 peaks, the most prominent difference between these samples (LO1 and LO2) is in the W3 peak, which appears as an intense feature for only the LO1 type specimens. When a scan of higher mass numbers released during the TPD measurements was performed, some interesting results emerged, shedding light on the origin of the W3 peak in Figure 2. For the LO1 type samples, an intense release of mass 44 (CO2) and mass 28 (CO) concurrently occurred with the W3 water desorption peak. This is illustrated in Figure 3A. Additional small CO emission appeared as a higher-temperature shoulder on the intense CO peak. This feature is less distinct for the heating rate applied in Figure 3; however, it appears clearly for lower heating rates. The CO2 and CO release pattern is
Figure 2. TPD spectra of water desorption (solid lines) and hydrogen release (broken lines) for the heavily oxidized uranium of types (A) LO1 and (B) LO2 (see text). The water desorption peaks are labeled by W0, W1, W2, and W3, and those corresponding to hydrogen release are labeled by Hs, Hh, and Hb. Heating rate, 30°/min. Totals of water containment (of all W peaks integrated area) are (A) 1.7 × 1018 and (B) 1.2 × 1018 molecules/cm2 (where the normalized area is the geometric area of the samples).
somewhat different than those for the LO2 type samples (Figure 3B). In this case, the intensities of these highermass species are much attenuated, and their positions are not correlated with any water desorption peak (only the end tail of the W2 feature extends to this temperature range). In addition, the CO peak of the LO2 samples is shifted to a higher temperature as compared with the main CO peak in Figure 3A and seems to correspond to the previously mentioned shoulder in that peak. An infrared spectroscopy study of the interactions between gas-phase CO2 and UO2+x surfaces,8 as well as adsorption studies of this system,9 pointed to the formation of strongly chemisorbed carboxylate (CO2-) or carbonate (CO3-2) species on the oxide surface. The chemisorption characteristics of these forms strongly depended on the stoichiometry of the oxide (i.e., on x in UO2+x). Hence, it is possible that the heavily oxidized samples, which stayed in air for very long periods of time, contained some carboxylate- or carbonate-chemisorbed moieties, which decompose on heating, producing CO2 and CO. The concomitant desorption of water in the LO1 samples may be due to two possible reasons: (i) formation of complex surface groups (H2O-COx) where water or hydroxyls are bonded to the carboxylate or carbonate species (hence, the thermal decomposition of these complex moieties produces the desorbed mixture of water, CO2 and CO) (8) Colmenares, C. J. Phys. Chem. 1974, 78, 2117. (9) Colmenares, C. A.; Terada, K. J. Nucl. Mater. 1975, 58, 336.
5916 Langmuir, Vol. 15, No. 18, 1999
Danon et al.
Figure 4. TPD spectra of water desorption during the first heating cycle (solid lines) and after repeated desorption (at 400 °C) rehydration cycle (broken lines) of LO1 samples. Rehydration was performed by 4 h exposure of the desorbed samples to air at room temperature under controlled RH of 33%. Heating rate, 30°/min.
Figure 3. TPD spectra of water (solid lines) and concurrent desorption peaks of CO (dashed lines) and CO2 (dotted lines) for the heavily oxidized samples (A) LO1 and (B) LO2 presented in Figure 1. Heating rate, 30°/min.
and (ii) a surface reaction which occurs between the decomposed CO2 molecules and hydrogen (which, as will be discussed later, is contained in the LO1 samples) to form water and CO. In this case, the W3 water peak is the product of that surface reaction and is not actually associated with a chemisorbed form on the oxide surface. In any case, it seems that the major contribution to the W3 peak in the LO1 type samples is related to the presence of surface, carbo-oxy species produced by the long period exposure to CO2 (in air). In addition to the above waterrelated decomposition of the carbo-oxy forms, the shifted CO shoulder in Figure 3A might indicate that an additional carboxylate decomposition route, not associated with water release, takes place. This route seems to be the dominant one concerning the LO2 samples, where the main CO peak is located at a temperature corresponding to the abovementioned shoulder. The different behavior of the LO1 and LO2 samples, associated with the W3 peak may be related to the two above-mentioned possible explanations of the W3 formation: (i) A different stoichiometry of the product oxide layer may occur for the two types of specimens which may lead to the different amounts of the chemisorbed carbooxy groups.8,9 These groups in turn may occupy two types of sites, those that can bind water molecules to form H2OCOx complexes and those that form a nonhydrated carbooxy form. The latter seem to exist on both LO1 and LO2 type samples, whereas the former occurs only on the LO1 type specimens. (ii) If the hydrogen-CO2 reaction produces the W3 water peak, the higher concentration of the carbooxy species on the LO1 samples and the presence of a relatively high concentration of dissolved bulk hydrogen in this uranium (unlike in the LO2, as discussed later) lead to the formation of this feature only in the LO1 case.
The above proposed relation of the W3 peak to the carbooxy chemisorbed moieties is substantiated by the fact that after a cycle of desorption (by heating to 400 °C) followed by air- or water-exposure, both types of samples, i.e., LO1 and LO2, display a similar TPD pattern, as depicted in Figure 4. In both cases, the W3 peak is missing, and only the W0, W1, and W2 peaks remain. In this context, it should be pointed out that only a fraction of the original water containment is restored after the desorption-rehydration cycle. This restored fraction decreases with the increases in the maximum temperature that is applied in the desorption scan. It is likely that the partial sintering process of the oxide, which reduces its specific surface area, affects the reduced amount of restored chemisorbed water. Such a sintering process has been reported for corrosion-produced UO2 powders.6 In addition, X-ray diffraction of the as-received LO1 samples indicated (as mentioned in section 2) broad UO2 diffraction lines which corresponded, by shape analysis, to a very small crystallite size (about 7-9 nm). Heating the samples to 650 °C significantly narrowed these XRD lines, indicating crystallite growth.10 Additional understanding was gained when isotope exchange experiments were performed, by exposing the LO1 samples to D2O (at room temperature for 24 h). A substantial exchange was apparent in the W0 and W1 peaks, whereas a less efficient exchange occurred for the W3 peak (Figure 5). The fact that some nonnegligible exchange is apparent in this latter peak may suggest that the formation of complex H2O-COx surface groups is more likely to be the origin of the W3 peak of the LO1 samples and not the H2-CO2 reaction. The H2O in this former case, though present on the surface, seems to be strongly bonded to the COx- groups. Hence, it is accessible to the D2O molecules to some extent producing a discernible exchange, but its isotopic conversion rate is much slower than that of the other hydrated forms associated with the W0 and W1 peaks which seem to be less tightly bound. It has been proposed that carboxylic acid derived moieties are chemisorbed on pair sites of neighboring metaloxygen.11 A possible assignment of the lower temperature water desorption peak will now be proposed. This peak, W0, is (10) Danon, A.; Koresh, J. E.; Mintz, M. H.; Kimmel, G., to be published. (11) Barteau, M. A. J. Vac. Sci. Technol. A 1993, 11, 2162.
TPD Characterization of Uranium Surfaces
Figure 5. Isotope exchange experiments performed on the LO1 type samples. TPD spectra after exposure at room temperature to D2O (100% RH) for 20 h. Total integrated water containment (of all isotopic mass numbers of all peaks) is 2.4 × 1018 molecules/cm2. Heating rate, 30°/min.
Figure 6. Effect of room temperature outgassing on the LO2 type samples. The solid line is TPD before outgassing and the broken line is TPD of a sample outgassed at room temperature for 20 h (both normalized to the same W2 peak intensities). Heating rate, 30°/min.
likely to be associated with a reversibly adsorbed molecular form of water. Such a form is anticipated to desorb easily, even at the ambient temperature, when the partial pressure of the gas-phase water vapor is reduced (either by evacuation or by passing a dry inert gas over the samples). Figure 6 illustrates the effect of room temperature outgassing (performed by passing dry helium over the sample for about 20 h) on the TPD spectra of the LO2 samples. It is seen that the only peak affected is the W0 one which is strongly attenuated by that treatment. In fact, the intensity of the W0 peak is altered even for much shorter outgassing periods (tens of minutes), and as stated in section 2, some variations in the time period between the insertion of the sample into the TPD cell and the startup of the heating scan may induce discernible changes in the intensity of that peak. The relatively weak binding of the corresponding water molecules may thus suggest that this W0 peak can be attributed to either a second layer chemisorption (i.e., chemisorption of water over a first layer of chemisorbed hydroxyls or water) or chemisorption on some weakly binding surface sites present on the oxide. It has been argued12 that nondissociative chemisorption (12) Thiel, P. A.; Madey, T. E. Surf. Sci. Rep. 1987, 7, 211.
Langmuir, Vol. 15, No. 18, 1999 5917
of molecular water can occur on oxide surfaces up to relatively high temperatures (ambient and above). Also, the binding strengths of different sites may differ considerably.12 Thus for example, an oxide surface of a given crystal orientation may terminate with oxygen anions [e.g., the UO2 (111) plane13,14], metal cations, or both, anions and cations being in a coplanar arrangement [e.g., the UO2 (110) plane13]. The surface of the very heterogeneous oxide formed by the water corrosion reaction of uranium may compose a variety of different crystal planes, as well as a significant concentration of low-coordination defects and “nonlattice oxygen”.12 Hence, the very broad shape of the W0 peak may be attributed to the occupation of a certain distribution of chemisorption sites rather than to a certain defined site assignment. The water desorption peak W1 which overlaps the preceding W0 one is even broader and more intense. It may probably be attributed to either a class of sites which interact more strongly with chemisorbed water molecules or involvement with dissociated hydroxyl groups. As known in the literature, significant differences may exist for example between the stability of water molecules attached to cation-terminated planes as compared to anion-terminated ones.12,15 Accordingly, the W0 peak may be related to water chemisorbed on some oxygenterminated planes of the UO2, whereas the W1 peak may be related to water chemisorbed on a U4+ array. Certainly, as mentioned above, a nonuniform distribution of different coordination numbers and defects can occur within each class of sites, producing the widening of the corresponding TPD peaks. An alternative or additional source to the W1 feature may arise from decomposition of clustered hydroxyl groups to form water. Evidently, these assignments of the lower-temperature desorption peaks are speculative in the absence of additional supporting information and can be regarded as possible hypotheses only. Finally, the W2 desorption peak may originate from isolated hydroxyl species,12 which as often observed decompose at higher temperatures. Some additional indication to the origin of the W1 and W2 peaks is provided by the rehydration behavior of the desorbed samples, as illustrated in Figure 4. It is seen that the rehydrated water preferentially occupies the lower stability sites associated with the lower-temperature peaks. Usually, the reverse chemisorption trend is observed with the more energetically stable sites being occupied first. Hence, in the present case, a certain kinetic barrier seems to control the re-occupation of the more stable sites associated with the W1 and W2 peaks. The formation of hydroxyl groups provides a straightforward explanation to such a kinetic barrier because the dissociation process of the chemisorbed water molecules may be a slow, kinetically limited process associated with surface reconstruction phenomena. For example, the existence of pairs of neighboring U4+-O2- sites might be required for such dissociation, with such pairs provided by a slow reconstruction process of the surface. 3.1.2. Hydrogen Release Peaks. The release of hydrogen, which was monitored concurrently with that of water, is illustrated by broken lines in Figure 2. For the (13) Ellis, W. P.; Taylor, T. N. Surf. Sci. 1978, 75, 279; 1980, 91, 409; 1981, 107, 249. (14) Castell, M. R.; Muggelberg, C.; Briggs, G. A.; Goddard, D. T. Vac. Sci. Technol. B 1996, 14, 966. Castell, M. R.; Dudarev, S. L.; Muggelberg, C.; Sutton, A. P.; Briggs, G. A. J. Vac. Sci. Technol. A 1998, 16, 1055. (15) Zwicker, G.; Jacob, K. Surf. Sci. 1983, 131, 179.
5918 Langmuir, Vol. 15, No. 18, 1999
Figure 7. TPD spectra of hydrogen release of a polished, oxidefree sample (source 1) hydrided to form a thin surface hydride layer. Heating rate, 30°/min.
LO1 type samples (Figure 2A), three distinct hydrogen release features can be identified and are labeled as Hs, Hh, and Hb, respectively. For the LO2 samples (Figure 2B), the latter high-temperature feature, Hb, is missing and the low-temperature peak, Hs, is less pronounced than that of the LO1 case. The broad high-temperature feature, Hb, has a very smeared and asymmetric shape. Its existence within a temperature range where all of the hydrated forms had already been desorbed may suggest that this feature arises from a bulk hydrogen source. Hydrogen is known to exist as a bulk dissolved (or hydride) impurity in many metals, and its contamination level depends on the details of the manufacturing process of the metal. The above view was checked by performing TPD measurements on oxide-free samples which were polished before being introduced into the TPD cell. In these experiments, the only measurable peak was that of the Hb, which appeared only for the samples of source 1. Its intensity was the same as that obtained for the oxide-coated LO1 samples. Hence, the relation of the Hb feature to a bulk impurity in the metal was confirmed. The contamination level of source 1 uranium seemed to be much higher than that of source 2, possibly because of the different manufacturing processes of these sources. In an attempt to identify the other hydrogen desorption features, an oxide-free, polished sample (from source 1) was partially hydrided by exposure to hydrogen gas at 200 °C. The TPD spectrum of this sample is illustrated in Figure 7. Besides the previously described small bulk peak, the main decomposition peak of the hydride closely corresponds to the labeled Hh peaks of the LO1 and LO2 samples. Hence, it is likely that the main contribution to this Hh feature arises from a near-surface hydride (possibly located at the oxide-metal interface region) produced by the water-uranium oxidation reaction (see section 4). It might be argued that the desorption of water associated with the W3 peak induces a water-metal reaction that produces the Hh release. However, a comparison of parts A and B of Figure 2 indicates that a comparable Hh feature is produced (for the LO2 samples) even in the absence of a concomitant W3-related water decomposition. Hence, the above assignment of the Hh to surface hydride decomposition seems to be more likely. The occurrence of a lower temperature hydrogen release feature, Hs, may probably be associated with the water decomposition related to the W1 peak. The difference in the temperature of these two peaks (about 50-60 °C) may
Danon et al.
Figure 8. Schematic structure of the oxide coating the metal. Most of the oxide (zone 1) is composed of a porous, nonadherent oxide, whereas a thin layer (zone 2) corresponds to an adherent protective layer.
Figure 9. TPD spectra of source 1 type uranium samples outgassed at 650 °C, polished, and oxidized again at room temperature with atmospheric air at 33% RH for different periods of time.
be due to the time lag between the release of the water from the oxide surface and its access to the metal interface. As described in section 2, most of the 10-15 µm thickness of the oxide layer coating the LO1 and LO2 samples is actually composed of a powderlike porous nonadhesive structure which does not provide a corrosion barrier to the underlying metal.2,6 Only a very thin layer (less than about 100 nm) is adherent and coherent enough to provide a diffusion barrier, protecting the underlying metal from gas-phase corrosion reactions. This is depicted schematically in Figure 8. Hence, some of the water vapor desorbing from the porous oxide (zone 2 in Figure 8) may reach the surface of zone 1 oxide and react with the metal underneath (by the mechanism discussed in section 4). The time duration required for the mass transfer of water from zone 2 to zone 1 and the water-metal reaction kinetics may then determine the above time lag between the water desorption peak and the associated hydrogen release feature. The fact that hydrogen produced by the above reaction is liberated at temperatures below the Hh peak might indicate that at least part of it is not accommodated as a surface hydride phase. This point will further be discussed in section 4. 3.2. Slightly Oxidized (SO) Samples. Figure 9 presents the TPD spectra of uranium samples outgassed at 650 °C (to release all bulk hydrogen containment producing the Hb feature), polished and washed with 50% HNO3 to obtain an oxide-free surface, and then reacted at room temperature with air at 33% constant RH for
TPD Characterization of Uranium Surfaces
different times (1 week-4 months). It is seen that the trend discussed in section 3.1, the reverse occupation sequence of the different sites (i.e., the less stable sites are occupied first), also prevails in this case. The W0 and part of the W1 sites accommodate the rehydrated water first, then the proportion between the W1 and W0 increases with increasing extent of oxidation, and finally the W2 sites start to evolve. The explanations suggested in the previous section, the kinetically limited processes, may also account for the present trend. It should be pointed out that the significant increase in the total water content observed for increasing exposure time is mainly due to the increase in the thickness of the formed oxide and is much less affected by the changes in the surface coverage of the oxide particles. 4. Discussion The TPD characterization of oxidized uranium surfaces may elucidate some of the phenomena encountered in certain gas-phase-uranium reactions. 4.1. The Hydrogen-Uranium Reaction. First, the so-called “activation” process of uranium toward hydrogen attack is considered: When uranium metal is exposed to hydrogen gas under a given pressure and temperature, a hydriding reaction takes place, converting the metal into UH3. This reaction starts at the gas-surface region, where the nucleation and growth of the hydride phase characterize the incipient stage of the reaction.1 Usually, some induction period precedes the nucleation and growth process. This period is longer for the lower temperatures and pressures applied in the reactor. Yet, for given reaction conditions, it is possible to shorten these induction periods by applying vacuum heat pretreatments (referred to as “activation”).1,3,4 Besides shortening the induction periods, such treatments accelerate the nucleation rates of the hydride, yielding a denser pattern of growing nuclei on the metal surface. The dependence of the induction periods on the heat pretreatment temperature (under given hydriding conditions i.e., temperature and pressure) displays the behavior of a strongly decaying function, with a strong attenuation of the induction periods occurring with increasing activation temperature up to about 200250 °C. A significantly less pronounced effect is produced for higher pretreatment temperatures. It has been proposed3,4 that activation is induced by the desorption of chemisorbed groups (possibly water) present on the protective oxide layer coating the metal, thereby facilitating the dissociative chemisorption of hydrogen on this oxide surface and the penetration of hydrogen through the oxide. From the TPD data presented in section 3, it is evident that the above range of temperatures which is most effective in activation correlates with the W0 + W1 water desorption peaks. Hence, the indirect, speculative assumption proposed in the previous activation studies3,4 is substantiated by the present TPD results which are consistent with the view that the dehydration of chemisorbed water (or clustered hydroxyl groups) provides the required active sites for the dissociative chemisorption of hydrogen on the oxidized uranium surface. 4.2. The Water-Uranium Oxidation Reaction. The kinetics of the water-uranium oxidation reaction (to form UO2 and hydrogen) are dependent on the partial pressure of the vapor gas phase. The exact functional form of the pressure dependence is quite ambiguous and controversial;2,6 however, the qualitative trend is an increase in the reaction rate with increasing water pressure.
Langmuir, Vol. 15, No. 18, 1999 5919
From the TPD results presented in section 3, it turns out that the only water desorption peak that is affected by the partial pressure of the water is the W0 peak, which probably corresponds to the reversibly adsorbed molecular form. This form was also observed by FTIR measurements made on UO2 powders.6 It is thus possible that this is the actually active component in the oxidation reaction of uranium. The fact that this form is the dominant one for the early stages of oxidation (e.g., as presented in Figure 9) supports its assignment as the active component in the reaction. The relation proposed in the literature between the oxidation reaction rate and the water containment of the oxide6 might thus refer only to that reversible form of the chemisorbed water and not to the total hydrated amount. As for the microscopic mechanism controlling the oxidation process, some understanding may be gained from the above assignment of the active water component and from the hydrogen release behavior (discussed in section 3.1.2). The appearance of the Hs (the low-temperature feature), which as discussed before seems to be related to the water oxidation reaction, suggests that part of the hydrogen formed by this reaction is not located within a hydride phase (which appears as an additional Hh peak at higher temperature). Evidently, it is not accommodated as a solid solution in the metal, because the peak corresponding to the release of that form, Hb, occurs at even higher temperatures. It is thus likely that this low-temperature release might be associated with either a surface form present on the protective oxide (see Figure 8) or to some near-surface hydrogen contained within that oxide. Combining that with the previously mentioned characterization of the W0-peak-related component, the following sequence of microscopic events can be postulated (see Figure 10). (i) A reversible chemisorption of the H2O molecules on a certain weakly binding site on the surface of the protective oxide layer [henceforth denoted as H2O (ad.)]. (ii) An electron transfer from the metal (located at the metal-oxide interface) through the protective oxide, to the chemisorbed water molecule, producing neutral hydrogen at the surface (Hs). This reduction reaction can occur either for one or for both of the protons of the water molecule:
H2O(ad) + U+int + e- f OH-(ad) + Hs
(1a)
H2O(ad) + U2+int + 2e- f O2-(ad) + 2Hs
(1b)
2Hs f H2(g)
(2)
[The subscript int refers to the oxide-metal interface location.] (iii) As a consequence of step ii, an electric field gradient is produced across this protective oxide (formed by the negatively charged surface groups and the positively charged uranium cations at the oxide-metal interface). This gradient induces a barrier for further electrons transfer (eq 1) and thus impedes the continuation of these reactions. To cancel the electric gradient, diffusion of OH- or O2should occur across the oxide, into the oxide-metal interface. These diffusion processes can be written as: diffusion
OH- (ad) + U+int 98 [U(OH)]int diffusion
O2-(ad) + U2+int 98 [UO]int
(3a) (3b)
5920 Langmuir, Vol. 15, No. 18, 1999
Danon et al.
[U(OH)]int f [UO]int + Hint
(4)
In that case, the neutral hydrogen is located at the metal vicinity, and thus can react and form the hydride:
Uint + 3Hint f [UH3]int
(5)
(v) The oxidation of the UO can further proceed (by a route similar to the above sequence) until the U4+ oxidation state is achieved. The above-proposed mechanism may thus account for the low-temperature hydrogen release peak, Hs, associated with eq 2 and that of the higher temperature hydride peak, Hh, associated with eq 5. Another unresolved effect associated with the water oxidation reaction of uranium is the inhibition of this reaction by presence of small amounts of oxygen.2 Two possible mechanisms were proposed to account for that oxygen effect: (1) Preferred occupation of water chemisorption surface sites by oxygen. (2) Changes induced by minor amounts of chemisorbed oxygen on the transport properties of the oxide. In an attempt to identify the proper mechanism, two heavily oxidized samples (of the type LO1) were outgassed at room temperature and then exposed to two atmospheres containing the same partial pressure of water. One of the two atmospheres contained 100 Torr of oxygen, whereas the other was oxygen-free and contained (in addition to the water) only inert argon. The TPD spectra of the two samples were identical without any change in either of the desorption peaks. This result seems to rule out the first mechanism, because a blocking effect of water chemisorption sites (by oxygen) would produce an attenuation of the corresponding water desorption peaks. Hence, oxygen-induced variations of the mobilities of the O2- and OH- ions (eq 3) are probably controlling the deleterious phenomena induced by oxygen.
Figure 10. Sequence of microscopic events for uranium oxidation with water molecules (see text).
(iv) For eq 3a, further proton reduction can occur, according to
Acknowledgment. This study was partially supported by a grant from the Israel Council of High Education and the Israel Atomic Energy Commission. Thanks are due to Dr. U. Admon, I. Avraham, and D. Weinraub for their valuable assistance. LA981210G
