摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(2), 345-347 Photothermal Imaging of Electrochemical Reaction Dynamics Roger S. Hutton and David E. Williams Department of Chemistry, University College London, 20 Gordon Street, London, UK WCIH OAJ The thermal effects associated with a focused light source have been utilised in conjunction with a confocal scanning laser microscope to investigate localised reaction dynamics. This new technique termed scanning laser photothermal electrochemical microscopy (SLAPEM) was used to study hexacyanoferrate(1i) oxidation at a gold electrode. The SLAPEM response is shown to be sensitive to electron transfer kinetics, enabling local variations in mass transport and kinetic control to be determined. Consequently, the important influence of the electrode edge and surface relief on electrochemical reactions at a disc electrode is shown.The origin of image features such as shadowing and hot spots are described and examined by investigating the influence of elec-trode potential, hexacyanoferrate(ii) concentration and laser-spot velocity and size. Electrochemical techniques can provide much information on a variety of systems. However, usually only an average of all responses across an electrode surface is obtained. Recently an assortment of in situ spatially resolved techniques have been developed which enable localised interfacial phenomena to be investigated. '-' Notably, scanning laser have been developed which enable the photocurrent and photovol- tage generated at semiconducting electrodes to be mapped across a surface.This provides information on crystal imper- fections and non-uniformity in electronic properties of the semiconductor. In this paper the small heating effect caused by the absorption of light at the electrode/electrolyte inter- face has been employed at a metallic electrode with an adapted confocal scanning laser microscope to monitor reac- tion dynamics spatially. This new technique, termed scanning laser photothermal electrochemical microscopy (SLAPEM) is demonstrated for the first time and used to investigate localised reaction dynamics of hexacyanoferrate(I1) oxidation at a gold elec- trode. Results are presented and the observed features are described theoretically. Analysis of images indicates that sub- micrometre surface relief has a strong influence on diffusion to the electrode.Experimental The principles of a scanning laser microscope have been described in detail elsewhere.' In this study a Bio-Rad Micro- science Ltd. MRC 600 confocal laser microscope was used, modified to obtain variable linescan speeds. A focused light spot from an argon ion laser (wavelength 514 and 488 nm, 10 mW) is rastered over the surface of an electrode using mirrors. Light returning from the specimen as a result of spe- cular reflection is passed back through the microscope objec- tive lens and monitored with a photomultiplier tube, to obtain an optical image of the electrode surface. Localised current perturbations for each beam position are accumu- lated and saved in a framestore to obtain a grey scale image.(0-256). If the background current is stationary then the per- turbations are directly related to the influence of the imping- ing laser beam. Experiments were performed in a conventional three-electrode cell under potentiostatic control, the output of which was dc-coupled through a pre-amplifier and offset to match the input range of the framestore. The image contrast therefore represents variations in current due to the move- ment of the laser spot across the surface. For each image the laser was scanned from left to right with variable velocity, 0.2 to 60 cm s-', the beam position was held for 0.06 s and then the next line was scanned. Since the current induced by the laser is small (ca.1 nA), measurements were signal-averaged over 50 frames: the total time to record one image was ca. 4 min. Working electrodes were made by encasing gold wire in epoxy resin, which was then polished to a mirror finish with 30, 10, 3, and 1 pm alumina-powder-water slurries. Analyti- cal grade chemicals were used throughout with triply distilled water. Potentials were measured relative to a saturated calomel electrode (SCE). Theory The interaction of an intense focused light spot at an elec- trode surface may, in principle, cause a variety of effects such as enhancing film growth,'.' changing film ~toichiometry~ and photoemission of electrons." However, under the present conditions of study the overriding influence is degra- dation of incident optical energy to heat.Temperature is an important factor in electrochemical pro- cesses, and several thermal effects may occur. Thermal effects associated with intense illumination on electrode processes at the mercury solution interface have been noted by Benderskii12.13 and Barker14-16 and have been used to study relaxation in the structure of the electrical double-layer. Small changes in standard electrode potentials with tem-perature (related to entropy changes) have been observed for a variety of electrode reaction^.'^.'^ The influence of tem- perature changes on diffusion (Soret effect) has been exten- sively investigated and utilised by Valdes and in thermal-modulation voltammetry.The temperature depen- dence of electron transfer rates (related to activation energies) has also been investigated e~perimentally~~ and postulated' as a possible basis for scanning laser imaging. The expected temperature change from a scanning laser beam has been derived previously7 following the formalism of Benderskii and Vehlichko.' Assuming each image point is illuminated by a pulse, the temperature is allowed to relax between points and the size of the electrode is much greater than the illuminated area, it can be shown that the tem- perature increase produced by a laser pulse is given by AT = where K, C, p and xl, C,, p1 are the heat conductivity, spe- cific heat and density of the metal substrate and solution, respectively, L is the light intensity, R the reflection coeff- cient andf(t) describes the laser pulse time-dependent shape.Thus for a pulse of duration to the maximum rise in tem- perature AT,,, is given by With insertion of the known constants for gold and water, assuming a reflection coefficient of 0.6 and a beam power of 10 mW focused to a spot of diameter 1 pm, the temperature rise is estimated to be of the order of 2 K. If only thermally induced changes in the rate of electron transfer kinetics are considered, it may be shown’ that the laser-induced current perturbation is given by 61 = 6TiAE,a/(RT2) (3) where 6T is the local temperature change, i is the total current density flowing, AEa is the activation energy, a is the area of the focused spot and T is the ambient temperature.Insertion of typical experimental values of the total current density, laser-spot diameter and an activation energy of 57 kJ mol-’ 23 indicate a photothermal current of the order of nA is to be expected. Eqn. (3) indicates that image contrast could arise as a con- sequence of variation over the surface of the temperature excursion, the activation energy or the dark current density. The size of the signal may also be influenced by thermal properties of the entire electrode assembly and the electrolyte. No response should be observed where the reaction is con- trolled by mass transport (i.e. negligible thermal diffusion effects), therefore images should reflect spatial variations in kinetic control.Results and Discussion Cyclic voltammograms of 5 and 10 mmol dmb3 K,Fe(CN), (0.1 mol dm-3 K2HP0, + 0.1 mol dm-3 KH2P0,) at a gold electrode (radius 250 pm) are shown in Fig. 1 highlighting regions of interest. At low potentials (region A) no current flows since this is below the equilibrium potential for the Fe(CN)64-’3-redox couple. As the potential is increased current flows and the reaction rate is controlled by the kinetics of electron transfer (region B). In region C the reac- tion rate is jointly controlled by electron transfer kinetics and the transport of Fe(CN),,- to the electrode surface. Finally I 11 I I I I-1 ‘ 0.0 0.2 0.4 0.6 potentialp vs. SCE Fig. 1 Cyclic voltammogram of 5 mmol dm-3 K,Fe(CN), (0.1 mol dm-3 K,HPO, + 0.1 mol dmP3 KH,PO,) at a gold electrode (radius 250 pm).Scan rate 20 mV s-’. A, no reaction; B, kinetically controlled current; C, mixed control; D, diffusion control. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 in region D the reaction rate is dominated by the diffusion of Fe(CN),4-to the electrode (D). Plate l(a) and (b) shows two images, recorded simulta- neously, of the reflected light and photothermal response, at +200 mV vs. SCE. The optical image indicates that the elec- trode surface is rough, however the height of the observed features was estimated, by closing down the confocal aperture of the microscope, to be less than 1 pm. It is noted that the reflected light intensity falls in several regions near the elec- trode edge.This small height variation is caused by mechani- cal polishing of the electrode surface. Plate l(b) shows the corresponding SLAPEM image, in which the grey scale corresponds to a change in current of 10 nA. Since the image contains a large quantity of information it is convenient initially to consider only the origin of image contrast associated with localised changes in the rate of elec- tron transfer kinetics. It is important to consider the diffusion field. At disc elec- trodes the rate of diffusion is a function of position. This is emphasised by Fig. 2 which indicates that at the centre of the electrode the rate of diffusion (semi-infinite planar) is much lower than at the edge (radial). Therefore, close to the equi- librium potential, reactions at a disc electrode are kinetically controlled at the edges but diffusion controlled in the centre.Inspection of Plate l(b) shows that this is indeed observed experimentally, in that larger photothermal currents are observed at the edges. However the response does not match the border of the electrode perfectly: the image is distorted by local regions where a relatively large response is obtained (hot spots) and transient effects (shadowing on the right of the image). Comparison of Plate 2(a) and (b)indicates that the hot spots are due to surface roughness. One explanation for the effect might be that the photothermal current is enhanced as a consequence of multiple reflections of the laser beam.9 A more likely explanation is that at surface asperities the current is kinetically controlled to a greater extent due to enhanced radial diffusion.This interpretation is supported by the observation that, upon rotation of the electrode, the hot spots moved around with elevated surface features. There are two sorts of transient effect visible: the first is a blurring of the image as the spot scan velocity increases, the second is a shadow (i.e. a decrease in current below the average) following a hot spot. The obvious transient pheno- mena which might be considered to affect the image are tran- sient relaxation of the temperature and of the local diffusion field. As the laser beam travels across a region where a large photothermal effect is observed, the increase in current density depletes the solution of Fe(CN)64- near the electrode, driving the reaction into diffusion control, such that a smaller current flows until the steady-state concentration of Fe(CN),,-is attained.The relaxation time associated with this concentration-depletion transient process is therefore related to the diffusion coefficient of hexacyanoferrate(I1). Alternatively the transient effects may be related to the dissi- pation of heat from the electrode surface (thermal-dissipation A n insulation electrodeI Fig. 2 Diffusion profile to a disc electrode showing semi-infinite planar diffusion at the centre (B) and radial diffusion at the edge of the electrode (C) and at raised points (A) J. CHEM. SOC.FARADAY TRANS., 1994, VOL. 90 Plate 1 (a) Optical and (b) SLAPEM image of 5 mmol dm-3 K,Fe(CN), (0.1 mol dm-3 K,HPO, + 0.1 mol dm-3 KH,P03 at a gold electrode. Pixel size 1.5 pm. Image 767 x 512 pixels. Grey scale 0-256 corresponds to 6i w 10 nA. Beam velocity 5.8 cm s-'. Beam diameter 4 pm.Average of 50 frames. Potential +200 mV us. SCE. Plate 2 (a) Optical and (b) SLAPEM images of 5 mmol dm-3 K,Fe(CN), (0.1 mol dm-3 K2HP0, + 0.1 mol dm-3 KH,POJ at a gold electrode. Potential 200 mV us. SCE. Pixel size 0.35 pm. Image 768 x 512 pixels. Grey scale 0-256 corresponds to 6i z 10 nA. Beam velocity 1.3 cm s-'. Beam diameter 2 pm. Average of 50 frames. R. S. Hutton and D. E. Williams J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Plate 3 SLAPEM images of 5 mmol dm-3 K,Fe(CN), (0.1 mol dm-3 K,HPO, + 0.1 mol dm-3 KH,PO,) at a gold electrode at various potentials: (a) + 100, (b) + 150, (c) +200, (d) +250, (e) +275, (f)+ 350, (9)+375 and (h) +550 mV us.SCE. Pixel size 1.5 pm. Image 767 x 512 pixels. Grey scale 0-256 corresponds to 6i x 10 nA. Beam velocity 11.7 cm s-'. Beam diameter 4 pm. Average of 50 frames. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Plate 4 SLAPEM images of 5 mmol dm-3 K,Fe(CN), (0.1 mol dm-3 K,HPO, + 0.1 mol dm-3 KH,PO,) at a gold electrode at various laser spot velocities: (a) 0.8, (b)1, (c) 2, (d)4, (e)12 and (f)58 cm s-'. Pixel size 1.5 pm. Image 767 x 512 pixels. Grey scale 0-256 corresponds to 6i z 10 nA. Potential +200 mV us. SCE. Beam diameter 4 pm. Average of 50 frames.J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Plate 5 SLAPEM images of 10 mmol dm-’ K,Fe(CN), (0.1 mol dmP3 K2HP0, + 0.1 mol dmP3 KH,PO,) at a gold electrode at: (a) +50, (b) + 100, (c) + 150, (d) +200 and (e) +250 mV us. SCE. Pixel size 1.5 pm. Image 767 x 512 pixels. Grey scale 0-256 corresponds to Si x 100 nA. Beam velocity 11.7 cm s-’. Beam diameter 4 pm. Average of 50 frames. Plate 6 SLAPEM image of 5 mmol dm-’ K,Fe(CN), (0.1 mol dm-’ K2HP0, + 0.1 mol dm-’ KH,PO,) at a gold electrode with a laser spot of diameter 2 pm (cf:Plate 1).Pixel size 1.5 pm. Image 767 x 512 pixels. Grey scale 0-256 corresponds to Si x 10 nA. Beam velocity 5.8 cm s-Average of 50 frames. Potential + 200 mV us. SCE. J. CHEM. SOC. FARADAY TRANS..1994. VOL 90 transient). The laser has the effect of locally heating the elec-trode. this thernial energy will be transferred across the entire surface, increasing the rate of the electron transfer reaction. As the laser is removed from the electrode surface the current decreases as the electrode cools. The characteristic time ior the thermal dissipation is therefore related to the thermal dif-fusion coefficients of the gold substrate. surrounding epc.xy resin and the solution. The shadowing observed in Plate 3b) suggests that the dominant effect is depletion of hexac>ano-ferrate(i1) since shadowing is observed at the top left of the image whilst the laser is still on the surface. Apart from the relatively large thermal effects at the elec-trode edge, a photothermal current was observed at the centre of the electrode.where the current is 0.2 nA above the background level. Within the central region no contrast in SLAPEM images was observed associated with changes in local reflectivity (Plate 2). This is presumably because of' the lack of sensitivity to the photothermal effect (mainly diffusion controlled) and masking of features by transient effects. The origin of variations in the image contrast were investi- gated further by considering the influence of potential. hexacyanoferrate(r1) concentration and characteristics of the laser spot. Plate 3(a)-(h)shows SLAPEM images recorded at various dc potentials. At + 100 mV t's. SCE [Plate 3(u)].the observed image is faint, only a small photothermal current is induced. comparable with the background current.Thc response Is dominated by one hot spot with a small contribution froni the electrode edge. As the potential is increased the total thcr- mally induced current increases, clearly showing the influence of radial diffusion near the electrode edge and at hot spot>. At higher potentials [Plate 3(c)-(e)] diffusion control becomes dominant and the total photothermal response decreases. However. even at high potentials a diffuse ring i.; observed around the electrode edge. indicating that the reac-tion is kinetically controlled at the electrode edge. Above 375 mV L'S. SCE the total photothermal response increases in intensity once more. this is surprising, as in this region the reaction is expected to be dominated by diffusitm control.This observation has been attributed to the occurrence of gold oxidation in parallel with hexacyano-fer ra te(I I) oxidation . Plate 4(u)-(y) shows SLAPEM images recorded at -L 33) mV 1's. SCE with various laser-spot velocities. At low belo-cities, features appear sharp. a response is obtained across the entire electrode surface. and hot spots are clearly defined. At a spot velocity of 0.8 cm s-' only a slight shadowing effect ii observed to the right of the image. As the spot velocit!. increases the major effect is that the images become blurred and the extent of shadowing increases. Plate 5(a)-(e) shows a series of SLAPEM images obtained at various dc potentials with 10 mmol dmh3 Fe(CNIo4 .Increasing the hexacyanoferrate(i1) concentration drives the reaction into greater kinetic control. Plate 5 indicates that increasing the hexacyanoferrate(i1) concentrations enhances the image resolution. in a similar manner to lowering the laser-spot velocity. A larger photothermal effect is observed but with reduced shadowing, supporting the concentration- depletion interpretation. Plate 0 shows a SLAPEM image obtained of 5 mmol dti Fc(CN),~-'showing the influence of changing the ni merical aperture and hence laser-spot size. Comparison w,th Plate 1 shows that decreasing the spot size increases the pt otothermal response because of the enhanced heating effect of a smaller spot. Conclusions PI lotothermal imaging of electrochemical reaction dynamics hcs been demonstrated for the first time with a new tech-ni lue.SLAPEM. SLAPEM images are shown to be sensitive to electron transfer kinetics. allowing spatial variations in ki ietic control at a disc electrode to be investigated. The th:oretical concepts that give rise to image contrast have be 3n discussed and were experimentally investigated. TI is work was funded by the SERC. The authors are grateful to Bio-Rad Microscience Ltd. for technical assistance. Rvferences 1 L. Varquez, J. M. Ci. Rodriguez. J. G.Herrero. A. M. Baro. N. Garciii. J. C. Canullo and A. J. Arvia. Suyf: Sci.. 1987. 181.98. 1-.A. J. Hard, G. Denault. C. M. Lee. D. Mandler and D. 0.Wipf. 4c.c. ('hem.Rex. 1990. 23. 357. 3 R. S. Lillard. P. J. Moran and H. S. Isaacs. .J. Elec~rroc~hem.Sot,.. 1991, 139, 1007. 4 K. F. Cohn, J. W. Wagner and J. Kruger. J. Electrochem. Soc.. 19x8. 135. 1033. 5 M. Kozlowski, W. H. Smyrl. 1.J. Atanasoska and R. Ata-nasoska. Electrochim. Actu. 1989. 34. 1763. h R. Peat. A. R. J. Kucernak. D. E. Williams and L.. M. Peter. Scmicoiid. Sci. Techno/.. 1990, 5. 914. 7 D. E. Williams. A. R. J. Kucernak and R. Peat. Electrochim. ictu. 1993, 38, 57. 8 R. Peat. A. Riley and D. E. Williams. J. Electrochem. Soc... 1989. 1-36. 3352. 9 .I. W. Schultze and J. lhietke. Electrochim. Acru. 1989. 34, 1769. 10 K. Leitner and J. W. Schultze. Ber. Bunsmqes. Phys. Chem.. 1988. 92. 181. I1 K. Mclntyre. D. K. Roe. J. K.Sass and H. Gerischer. Ber. Burt-\rnye.\. Phy. Chem.. 1987. 91. 488. I' V. A. Benderskii and G. I. Vehlichko. J. Electrocinal. C'hem.. 19x2. 140. I. 13 V. ,A. Benderskii, G. I. Vehlichko and I. V. Kreitus. J. Electro-mil. ('hem., 1984. 181, 1. 13 G. C. Barker. P. Fowles and B. Stringer. 7run.s. Ftiraduj So(,.. !970.66. 1509. IS G. C. Barker. Electrochim. .4cta. 1968, 13, 1221. lh G. C. Barker, Ber. Bunseriyes. Ph!x Chem., 1971, 75. 128. 17 K. S. Hutton. A. M. Bond. R. Colton and J. Harvey. in prep- 'i rat io11. 18 M.J. Weaver. J. Phys. Chem.. 1979. 83. 1748. 19 13. Miller, J. Electrochem. Soc.. 1983. 130. 1639. '0 I. L. C'aldes and B. Miller. J. Phjx Chem.. 1988. 92.4483. 'I J. L. C'aldes and B. Miller. J. Electrochem. Soc., 1988. 135. 2223. 11--I. L. C'aldes and B. Miller. J. Ph!:;. Chum.. 1988. 92. 525. $3 I-. A. C'urtiss. J. W. Halley. J. Hautman. N. C. Haung, Z. Nag>. Y. J. Rhee and R. M. Yonco. -1. Elrctrochem. Soc.. 1991. 138. 203'. 24 K. S. Hutton and D. E. Williams. in preparation. Paper 3 02723F: Rewired 13rh .May. 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000345

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 9012), 349-354 Stepwise Growth of Size-confined CdS in the Two-Dimensional Hydrophilic lnterlayers of Langmuir-Blodgett Films by the Repeated Sulfidation-Intercalation Technique lsamu Moriguichi, Katsuhiko Hosoi, Hidenori Nagaoka, lchiro Tanaka, Yasutake Teraoka and Shuichi Kagawa" Department of Applied Chemistry, Faculty of Engineering, Nagasaki University, Nagasaki 852,Japan When a CdS-bearing stearic acid multilayer produced by exposing a cadmium stearate Langmuir-Blodgett (LB) film to H,S gas is immersed in aqueous CdCI,, Cd ions are intercalated to regenerate the cadmium stearate multilayer without the escape of CdS. The repetition of sulfidation-intercalation cycles allowed the size-quantized CdS to grow in a stepwise fashion in the hydrophilic interlayers.The layered structure was main- tained throughout the repeated cycles, although the multilayer-constructing stearate molecules became less oriented by the formation and growth of CdS. The formation of a two-dimensional CdS plane was suggested. Nanosized inorganic semiconductor materials have been attracting much attention owing to their structural, chemical and physical properties which are somewhat different from those of the corresponding bulk materials. The most promi- nent feature of nanosized semiconductors as compared with the bulk materials is revealed in the electronic structure, that is, the blue-shifted energy gap and the discrete electronic levels.*-3 Therefore, the photophysical, photochemical and photocatalytic applications of nanosized semiconductors have been studied e~tensively.~-~ Since nanosized particles are inherently liable to aggregate or grow in order to reduce the surface energy, careful and controlled synthetic methods are required.The methodology applied so far can be conveniently classified into the follow- ing three categories: (1) Arrested precipitation in solutions by controlling solvents, concentration and temperat~re'~~ or by using stabilizers and growth-terminating reagents."?' ' (2) Stabilization of small particles in or by aggregation-preventing matrices such as polymers,'2,'3 gla~ses,'~.' sur- faces of monolayer thickness at the air/water interfa~e'~.'' and bilayer lipid membranes.I8 (3) In situ synthesis in the q2,confined spaces of zeolites,' '3,' clays,21 organized sur-factant aggregates (reverse mi~elles,~~-~' vesicle^^^.^^ and Experimental Materials Stearic acid, cadmium chloride, sodium hydrogen carbonate, benzene (Kishida Chemical Co., Ltd.), cadmium stearate (Shimakyu's Pure Chemicals Co., Ltd.) and hydrogen sulfide gas (>99.9%, Sumitomo Seika Co., Ltd.) were used as received.The water used for subphase solutions was purified by a Milli-Q system (Millipore Corp., resistivity > 14 MR cm). CaF, plates [20 mm (diameter) x 2 mm, Japan Spectro- scopic Co., Ltd.], quartz plates (10 mm x 45 mm x 1.25 mm, Fujiwara Co., Ltd.), borosilicate glass plates (76 mn x 26 mm x 1.5 mm, Matsunami Glass Ind., Ltd.) and gold plates (30 mm x 30 mm x 0.2 mm, Nilaco Corp., purity >99.95%) were used as substrates on which the LB film was deposited.The surfaces of gold plates were polished with an alumina abrasive (0.05 pm, Buehler) and were washed with pure water and then methanol (Kishida Chemical Co., Ltd.). Surfaces of other plates were washed with a dilute aqueous solution of HF (Wako Pure Chemical Industries, Ltd.), pure water and then methanol. Four substrates were used properly to meet the condition of instrumental analysis. bilayer membranes28), protein cages29 and LB filrn~.~'-~~Preparation of LB FilmsThese matrices play an aggregation-preventing role as well. Recently, some groups (including ourselves) reported the synthesis of metal chalcogenides, especially sulfides, in the hydrophilic interlayers of LB films of fatty acids and their Since the interlayers are the two-dimensional reac- tion field and the arrangement and the amount of precursor metal ions are well controlled by metal ion-binding head groups which are highly ordered and assembled, the inter- layers serve as the well restricted reaction field for the in situ synthesis of nanosized materials.The present invention relates to a new method for increas- ing in a stepwise manner the dimensions of size-quantized semiconductor particles or films in the interlayer of LB films. The method has a potential for wide applications to the prep- aration of size-quantized semiconductors with arbitrary dimensions. We have found and reported the in situ growth of size-constrained CdS in the LB film by the repeated sulfidation-intercalation te~hnique.,~In the present paper, a detailed investigation will be described on the synthesis and growth of CdS in the hydrophilic interlayers of LB films and the accompanied structure change of the LB matrix.A benzene solution of stearic acid (1 g dm-') was spread at 20°C on the surface of 3 x mol dmP3 aqueous CdCl, which was adjusted to pH 5.8 by adding aqueous NaHCO,. A separate experiment confirmed that under these conditions cadmium stearate was formed at the air/water interface after reaction between Cd2+ and stearic acid. The LB deposition was performed at a surface pressure of 30 mN m- ',at which the monolayer was in the solid condensed state, in the verti- cal mode with a combination of a film balance (Sanesu Keisoku Co.Ltd., Model FSD-20) and a lifter (Sanesu Keisoku Co. Ltd., Model FSD-23). Monolayers of cadmium stearate were transferred at a deposition rate of 10 mm min-', and the Y-type LB film of cadmium stearate was suc- cessfully built up with a transfer ratio of unity on well cleaned borosilicate glass, quartz, CaF, and gold plates. Production of CdS and Intercalation of Cd Ions The production of CdS was performed by exposure of a cadmium stearate LB film to a flow of H,S gas (105 cm3 3 50 min-') at room temperature, which is referred to as the sul- fidation (S) process. The sulfidized LB film was then immersed in 3 x lo-, mol dm-3 aqueous CdCl, which was adjusted to pH 5.8 with aqueous NaHCO,, followed by rinsing with pure water for 5 min [intercalation (I) process].Thereafter, the film was subjected alternately to sulfidation and intercalation processes. In this paper, the samples after the sulfidation and intercalation processes are denoted as S(n) and I(n) films, respectively, where n is the number of specified processes undergone by the film. It is natural that the S(n) film underwent the sulfidation process n times and the inter- calation process n -1 times, and the I(n) film underwent n sets of S-I cycles. Instrumental Analysis UV-VIS spectra of quartz-supported films were recorded on a Shimazu UV-3 100 spectrometer in the transmission mode, and the absorption onset was determined by the second derivative of the spectrum.IR spectra of CaF2-supported films were measured using a Nihon Bunko-IR-180 instrument or a Perkin-Elmer 1650 FTIR spectrometer in the transmis- sion mode. The reflection-absorption (RA) FTIR spectra were taken for Au-supported films with the FTIR spectrom- eter equipped with a specular reflectance accessory (Spectra- Tech Inc. Model 501, 85" incident angle) using p-polarized light. The sample chamber of the FTIR spectrometer was purged with a flow of dry air so as to minimize background H,O. X-Ray photoelectron spectra (XPS) of quartz-supported films were recorded on a Shimadzu ESCA-850M instrument with an Mg-Ka source (1253.6 eV). The binding energies (E,,) were calibrated with reference to the C Is line of the aliphatic carbon of stearate molecules (285.0 eV).The atomic ratios, Cd : C, S : C and Cd :S, were determined by using integrated areas, photoelectron cross-sections, and inelastic mean free paths of the C Is, S 2p and Cd 3d,,2 photoelectron lines. X-Ray diffraction (XRD) patterns of borosilicate glass-supported films were taken with a Rigaku 2034 diffractometer using Cu-Ka radiation. Results and Discussion Chemical Change during the Sulfidation and Intercalation Processes Changes of the carboxylate groups during the sulfidation and intercalation processes were followed by monitoring the IR absorption band of the C=O stretching vibration, and the formation of CdS was measured using UV-VIS spectroscopy.As shown in Fig. l(a), the original LB film of cadmium stearate (19 layers) showed an antisymmetric CO, stretching band of carboxylate ions (RCO;) at 1548 cm-'. Upon contact with H,S gas for 5 min, the band of RCO; was totally replaced by that of protonated carboxylic acids (RC0,H) at 1702 cm-', and the IR spectrum did not change on prolonged exposure to H,S [Fig. l(b)]. The change in the carboxylate groups from RCO, to RCO,H suggests that the cadmium stearate multilayer reacts with H2S to yield the stearic acid multilayer and CdS (acid-form composite film), [(RCOy),Cd]L, + H2S +[2RCO2H]L, + CdS (1) where the subscript LB refers to constituent molecules of the mu1 tilayer. The exposure to H,S gas gave rise to the appearance of an optical absorption due to CdS.The wavelength of the absorption onset (Aos) of the 19-layer film increased on increasing the exposure time to H2S (sulfidation time) and then reached a constant value after 20 rnin (Fig. 2). This J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 wavenumber/cm-' Fig. 1 Transmission IR spectra of a cadmium stearate LB film (19 layers) after exposure to H,S gas and immersion in aqueous CdCI, : (a) original LB film, (6) after sulfidation, (c) after immersion for 1 h, (d)after immersion for 3 h means that an exposure to H,S for at least 20 rnin is required to form the CdS in an equilibrium state under the present experimental conditions. As stated above, reaction (1) was completed within 5 rnin in view of the change of carboxylate groups.Accordingly, the observed spectral changes between 5 and 20 rnin are due to the growth of CdS rather than concen- tration effects; the CdS produced in the initial stages of the reaction may be in the form of small clusters (or even molecules), which grow with time. This result indicates that the dimensions of the CdS particles can be controlled by the sulfidation conditions. We reported previously34 that when the LB film of cadmium stearate was sulfidized in a flow of H2S gas (25 cm3 min-') for 15 min, the CdS formed had an absorption onset at 370 nm. In the present study, exposure to a flow of H,S gas (105 cm3rnin -') for 20 min was adopted as the sulfidation process in order to examine the equilibrium state.Note that the production and growth of CdS described below is realized even when sulfidation is stopped before reaching the equilibrium state.34 Fig. 3 shows UV-VIS spectra of 19-, 29-and 39-layer films sulfidized for 20 min. The fact that the absorption onset, that is, the dimension of CdS formed, was identical irrespective of 450 440E 278 430 .-cE $ 9 420 410 0 10 20 30 sulfidation time/rnin Fig. 2 Onset of absorption in the cadmium stearate LB film (19 layers) as a function of sulfidation time J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0.6, number of layers 1 0.1 J O'350. ' ' ' 400' ' ' ' ' 450' ' ' ' "500 wavelength/nrn Fig. 3 UV-VIS spectra of (a) 19-, (b)29-and (c) 39-layer cadmium stearate LB films after the sulfidation process.Inset : neat absorption of CdS (AA, see text) as a function of the number of layers. These spectra were obtained by using a quartz plate as a reference on each size of which a five-layer cadmium stearate LB film was deposited. the number of layers demonstrates that controlled formation of CdS takes place within each hydrophilic interlayer and the CdS formed remains in the interlayers. Because the hydro- philic interlayers of a cadmium stearate LB film are spatially confined and well separated from each other by the hydro- phobic organic layers and the amount of Cd2+ ions therein is regulated exclusively by the complexation with carboxylate groups, the hydrophilic interlayers would serve as a restricted reaction field for the in situ synthesis of CdS.In the present case the size of the CdS particles is constant and therefore the size-dependent absorption coefficient at a given wavelength should not be taken into account. The absorption spectra exhibited Beer-Lambert behaviour, as shown in the inset in Fig. 3. In order to eliminate from the UV-VIS spectra the influence of an increase in baseline absorbance (possibly due to scattering or reflection of light), the difference in absorb- ance at 370 and 600 nm (AA) was taken as the neat absorp- tion of CdS. The subsequent immersion of the sulfidized film into aqueous CdCl, caused the transformation of C0,H back to COY, and immersion for 3 h was necessary to complete the transformation, as shown in Fig.l(c) and (4. During the treatment, the intensities of the C-H stretching IR bands (Fig. 2) and UV-VIS absorption remained unchanged. Accordingly, it can be concluded that during the immersion process Cd ions are stoichiometrically intercalated into the acid-form composite film to give the salt-form composite film of CdS and cadmium stearate without the escape of CdS and stearate molecules from the film. Hereafter, the intercalation of Cd2+ ions was carried out by immersion into aqueous CdCl, for 3 h. Growth of CdS CdS coexists with its precursor ions (Cd2') in the hydrophilic interlayers of the salt-form composite films obtained in this way, and it is conceivable that CdS can be grown in a well controlled manner by the repetition of the sulfidation (S) and intercalation (I) processes.Changes in the IR and UV spectra of the 19-layer film caused by the S and I treatments are shown in Fig. 4 and 5, respectively. Stoichiometric and reversible transformation of carboxylate groups between the ionized and protonated forms was observed by the repetition of the S-I cycle while keeping the C-H stretching bands intact. This transformation was confirmed in up to six cycles of I-S treatments, though the spectra before the third S treat-ment are shown in Fig. 4. Repetition of the I-S cycle gave rise to an increase in A,,, which provides direct evidence for the growth of CdS in an LB matrix. Note that CdS formed after 35 1 IIII I LBISIiSI S1 1800 1600 1400 wavenumber/crn-' Fig.4 IR spectra of the C-0 stretching vibration of (a) a cadmium stearate LB film, and (b) S(1), (c) I(l), (6)S(2), (e) I(2) and cf) S(3) films; 19-layer films deposited on each side of a CaF, substrate the sixth sulfidation process even has A,, at 486 nm which is blue-shifted from that of bulk CdS (520 nm) and therefore is still size-quantized. As shown in the inset in Fig. 5, an increase in Aos became moderate after the fourth sulfidation process. This is because Aos gradually approaches the limiting bulk value with the growth of CdS. The optical absorption due to CdS (AA) increased with repetition of I-S cycles. This suggests an increasing amount of CdS, though the change in the absorption coefficient with size should be taken into account.19-Layer films were analysed by XPS. C 1s Spectra were composed of a main peak due to aliphatic carbon (285.0 eV) and a small additional peak due to carbonyl carbon (288.5 eV); the binding energy (EJ of the latter was the same in both the acid- and salt-form multilayers. The Cd 3d,,, signal was observed at 405.6 eV in the original LB film, the CdS- 0.6 490 480 .. 0.0 ' ' ' ' I ' ' ' " " 400 " ' 500' ' 600 wavelength/nrn Fig. 5 UV-VIS spectra of (a) a cadmium stearate LB film and (b)-(f)S(n) composite films of stearic acid multilayers and CdS. (b)n = 1, (c)n = 2, (d) n = 3, (e)n = 4, (f) n = 6. The onsets of absorption (A,) of S(n)films are plotted in the inset.These spectra were recorded with 19-layer films deposited on each side of a quartz substrate using uncoated quartz as a reference Table 1 Quantitative XPS analysis of the 19-layer film atomic ratio sample Cd : C s:c Cd : S S( 1) film S(5) film S( 10)film 0.069 : 1 0.23 : 1 0.33 :1 0.073 : 1 0.23 :1 0.29 :1 0.95 :1 1.0 : 1 1.0 :1 bearing multilayers in both the acid- and salt-form composite films and the reference materials (cadmium stearate and CdS powder), indicating that the divalent Cd ions in cadmium stearate and CdS cannot be distinguished. The S 2p signal was observed at 162.0 eV in the composite films after the first S treatment. The results of quantitative analysis are shown in Table 1. The amounts of cadmium and sulfur relative to carbon (Cd : C, S : C) increased steadily with repeating I-S cycle, and the Cd : S atomic ratio is nearly unity after each sulfidation process.These results confirm the increase in the amount of CdS in the LB film by repeating the I-S cycle. Layered Structure The layered structures of the cadmium stearate LB film and the CdS-bearing composite films were investigated by XRD (Fig. 6). The original LB film, which gave distinct X-ray 001 Bragg peaks [Fig. qa)], had a well organized layered struc- ture with a basal-plane spacing (d) of 50.2 A as reported by Matsuda et After the sulfidation process [Fig. 6(b)],the XRD peaks weakened considerably. However, the XRD pattern characteristic of the layered structure reappeared again after the intercalation process [Fig.qc)], though the peak intensity was weaker than that of the original LB film. Through further repetition of S-I cycles, a weakening in the intensity or almost complete disappearance of the peaks was observed during the S treatment, while reappearance of the peaks occurred during the I treatment [Fig. 6(d)-(f)]. The full width at half maximum of the (002), (003) or (004) peaks of the CdS-bearing composite films was 0.19",which is com- (f)Ik5 10 15 5 10 15 : 2O/degrees 20fdegrees Fig. 6 XRD patterns of (a) a cadmium stearate LB film (lo00 counts s-l) and (b)-cf) CdS-bearing composite films (400 counts s-l): (b)W),(4W),(4I(2), (4 I(3),(S)I(6) J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 parable to that of the original cadmium stearate LB film (0.20'). These results indicate that the layered structure is maintained through repetition of the S-I cycle. Particulate CdS, the strongest XRD peak of which appears at 28 = 26.5", was not detected even in the S(6)film. Orientation of the Stearate Molecules The change in orientation of the stearate molecules was investigated by RA and transmission IR spectroscopies. It is well known that the RA and transmission spectra are selec- tive to bands having transition moments perpendicular and parallel to the substrate surface, respectively. In addition, comparison of the RA and transmission intensities reportedly gave quantitative information about the molecular orienta- ti~n.~~-~'Fig. 7 shows RA spectra of nine-layer films sup- ported on a gold substrate. For the original cadmium stearate LB film, the symmetric COY stretching band at 1433 cm-',the transition moment of which is perpendicular to the surface, was exclusively observed in the RA spectrum (Fig.7). In the transmission spectrum (Fig. 4), on the other hand, the following strong bands with transition moments parallel to the surface appear: the antisymmetric and symmetric CH, stretching bands at 2916 and 2849 cm-', respectively, and the antisymmetric CO, stretching band at 1542 cm-l. These results mean that the stearate molecules in the original LB film are nearly perpendicular to the substrate surface. The RA intensities of the symmetric and antisymmetric CH, stretching bands of the CdS-bearing composite films increased gradually with repetition of the S-I cycles.More- over, the antisymmetric CO, band was observed for the 1(1) and 1(2) films in the RA spectra. These results give clear evi- dence that stearate molecules in CdS-bearing composite films become inclined gradually with repeated S-I cycles. The pres- ervation of the all-trans configuration and the uniaxial orien- tation of the alkyl chains through the spectra in Fig. 7(a)-(f) is suggested by the identical wavenumber of the C-H stretching vibration band^^'-^^ and the appearance of the band progression due to the CH, wagging mode between 1400 and 1200 cm- 1,44respectively. The molecular orientation of the nine-layer film is quanti- tatively evaluated by comparison of the RA absorbance (AR) of the Au-supported film (Fig.7) and the transmission absorbance (AT) of the CaF,-supported film (not shown). The A, value is naturally one half of the observed absorbance, because the CaF, plate carried nine layers on each side LB 3000 2800 Fig. 7 Change in the RA FTIR spectra on repeating the sulfidation-intercalation cycle; nine-layer films deposited on a gold plate J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 353 (totally 18 layers). When stearate molecules are uniaxially orientated and the in-plane contribution of the electric field to the RA spectra can be neglected, the orientation angles (0, and 0,) between the surface normal and transition moments of the symmetric and asymmetric CH, stretching modes are given by ei = tan-'J[rnXA,/A,),]; i = s or a (1) where m is the enhancement factor for the RA intensity on the Au substrate with respect to the transmission intensity on the CaF, substrate. Because the rn value of Au is the same as that for Ag, values of rn, = 9.16 and rn, = 9.05 are used in this study, which were reported for the nine-layer LB film of cadmium stearate on Ag and CaF, substrate^.^' The tilt angle (y) of the alkyl chain axis from the surface normal can be obtained by the orthogonality relation among 8,, 8, and y: cos2 e, + cos2 e, + cos2 = 1 (11) For the original cadmium stearate film, y is calculated to be 10".This value agrees well with the reported tilt angle of cadmium salts of long-chain fatty and the angle estimated from the X-ray diffraction res~lt.~~,~' The calcu- lated tilt angles of the original, sulfidized [S(n)] and inter- calation [I(n)] films are shown in Fig.8, from which two characteristic features are observed. The tilt angles of the salt- form films (open circles) increase on going from the original to 1(2) films and are almost constant from 1(2) to 1(4) films. The same tendency is also observed for the acid-form films (closed circles). This suggests that the formation and growth of CdS up to the second cycle cause the stearate molecules to be less oriented. Another characteristic observed from Fig. 8 is that the stearate molecules become less oriented during the S treatment and rearrange to recover the orientation during the I treatment.This may be due to the weaker molecular interaction in the acid-form films than in the salt-form films.50*51As described above, the acid-form, S(n) films showed no distinct XRD peaks. This is presumably due to the poor orientation of stearate molecules in the acid-form films and to Cd not forming a well arranged plane in the hydrophilic interlayers. Consideration of the Form of CIS The basal-plane spacings (d)of the salt-form films are calcu- lated from XRD results and plotted in Fig. 9. d increased steadily, but gently, with growth of CdS. It is clear that the variation of d includes changes due to both the formation of CdS and the molecular orientation. Thus the net change due to the formation of CdS (Ad,) is estimated by the following 54*01 n in I(n)film Fig.9 Basal-plane spacing (6)and estimated thickness of the CdS layers (Ad, see text) of I(n) films equation which takes the change in molecular orientation into account: d = d, -do[sin(90 -yi)/sin(90-yo)] (111) Here, do and yo are the basal-plane spacing and tilt angle of the original film, and d, and y, are those of the I(n) films. As shown in Fig. 9, Ad increases more steeply than d. Note that the Ad value of the 1(3) film (3.5 A) is close to the ionic diam- eter of sulfide ion (3.4 A). The area per stearate molecule in the original LB film was estimated to be ca. 21 81' from the surface pressure (n)-area (A) isotherm of a calcium stearate monolayer at the air/water interface.Since Cd2+ ions in Y-type cadmium stearate LB films are reportedly arranged in the same plane,37 the area per Cd2+ ion is also 21 A'. If we assume a close-packed planar arrangement of sulfide ions in the hydrophilic layers, then the area per sulfide ion is 10 A'. Therefore two or three repetitions of the I-S cycles are required to form a continuous CdS sheet at a monolayer level. The coincidence between Ad of the 1(3) film and the ionic diameter of the sulfide ions may imply the formation of such a monolayer sheet of CdS. If the reported correlation between band gap and particle size of spherical CdS'*3 is applied to the present cases, the diameters of CdS particles in S(n) and I(n) films are estimated to be 48 (n = l), 66 (n= 2), 70 (n = 3) and >100 A (n = 4).It seems that the spherical CdS particles, whose diameters are comparable to or larger than the basal-plane spacing, are too large to be accommo- dated in the hydrophilic interlayers without destroying the layered structure. Judging from the too large diameter of the imaginary spherical CdS particle and the observed magnitude of Ad, we consider that the CdS material formed is two- dimensional in nature, although a more detailed study will be .~~necessary to confirm this. Smotkin et ~1 and Grieser et .~~~1reported the formation of disc-shaped CdS in LB films. Conclusion301 " o S(l) i(1) S(2) l(2) S(3) i(3) S(4) l(4) Fig. 8 Tilt angles of the alkyl chain axis in original, sulfidized [S(n)] and intercalated [I@)] films The exposure of a cadmium stearate LB film to H,S gas yielded size-quantized CdS in the hydrophilic interlayers with concomitant conversion of the cadmium stearate multilayer to a stearic acid one.By immersion into aqueous CdCl, ,Cd ions are intercalated into the hydrophilic interlayers of the CdS-bearing stearic acid multilayer without the escape of CdS to form the CdS-bearing cadmium stearate multilayer. Further repetition of the sulfidation-intercalation cycle makes it possible to grow CdS in the hydrophilic interlayers. The layered structure is maintained during the repetition of sulfidation-intercalation cycles, though stearate molecules tend to become less orientated with the formation and growth of CdS.The onset of absorption of CdS and thickness 354 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 of the CdS layer, coupled with simple geometric consider- ations, suggest that the CdS produced in the hydrophilic interlayers might be in the form of a two-dimensional plane. In this study, we have developed a novel and simple tech- nique by which size-quantized CdS is formed and grown in the hydrophilic interlayers of LB films without destroying the 21 22 23 24 25 R. D. Stramel, T. Nakamura and J. K. Thomas, J. Chem. SOC., Faraday Trans. 1, 1988,84, 1287. H.Miyoshi, H. Mori and H. Yoneyama, Langmuir, 1991,7,503. M. Meyer, C. Wallberg, K. Kurihara and J. H. Fendler, J. Chem. SOC.Chem. Commun., 1984,90. C. Petit and M. P. Pileni, J. Phys.Chem., 1988,92,2282. M. L. Steigerwald, A. P. Alivisatos, J. M. Gibson, T. D. Harris, layered structure. The present results also suggest that the hydrophilic interlayers of metal-containing LB films serve as a restricted reaction field for the well controlled, in situ synthesis of inorganic materials, CdS in the present case. This is because the interlayers are spatially constrained and the amount of precursor metal cations in the interlayers is regu- 26 27 28 R. Kortan, A. J. Muller, A. M. Thayer, T. M. Duncan, D. C. Douglass and L. E. Brus, J. Am. Chem. SOC., 1988,110,3046. Y. M. Tricot and J. H. Fendler, J. Am. Chem. SOC., 1984, 106, 2475. R. Rafaeloff, Y. M. Tricot, F. Nome and J. H. Fendler, J. Phys. Chem., 1985,89,533. N. Kimizuka, T. Miyoshi, I.Ichinose and T. Kunitake, Chem. lated exclusively by complexation with the hydrophilic moiety of the multilayer-constructing molecules. 29 Lett., 1991, 2039. F. C. Meldrum, V. J. 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ISSN:0956-5000    

DOI:10.1039/FT9949000349

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(2), 355-362 Thermogravimetry-FTIR Study of the Surface Formate Decomposition on Cu, CuCI, Cu,O and CuO Correlations between Reaction Selectivity and Structural Properties Jianyi Lin* Department of Physics, National University of Singapore, Singapore 0511 Koon Gee Neoh and Wah Koon Teo Department of Chemical Engineering, National University of Singapore, Singapore 0511 In this work three types of the surface formate, ionic, monodentate and bidentate species, are identified to be the main products of formic acid adsorption on Cu, CuCI, Cu,O and CuO at room temperature. Combined thermal analysis [differential thermal analysis (DTA), thermogravimetry (TG) and evolved gas analysis (EGA)] and FTIR studies have shown that the decomposition selectivity and activity are highly dependent on the electronic and structural properties of the surface formate.The ionic formate which gives the characteristic infrared bands at 1600 and 1320 cm-' is found to exist on all the surfaces studied in the lower temperature range (<490 K) and is easily (with an activation energy of 5-19 kcal mol-') decomposed at elevated temperatures, forming water as the main product. Bidentate formate with infrared bands at 1630 and 1350 cm-' is observed on the Cu, Cu,O and CuO surfaces and is decomposed at high temperatures (around 500 K) with an activation energy of 1539 kcal mol-' , producing mainly carbon dioxide. The adsorbed formate on CuCl at high temperatures demonstrates only an asymmetric OCO band at 1610 cm-' and is considered to be the monodentate formate, whose pyrolysis favours the production of water, with an observed activation energy of 19 kcal mol-'.The reaction scheme and the origin of the structural dependence of the surfaceformate decomposition are discussed. Formic acid has been shown to be adsorbed mainly as the surface formate on a variety of metals and oxides at room temperature. The decomposition of the adsorbed for-mates has been often used as a model system of catalytic selectivity studies since only one reactant is involved and the products of the reactions (both dehydration and dehydrogenation), CO,, H,, CO and H,O, can be easily monitored. Although due attention has not been paid to studies on copper oxides, a study on copper metal has re- cently been considered particularly attractive as the surface formate species is often thought to be a pivotal intermediate in several catalytic processes including methanol synthesis and water-gas shift.' 7-44 The correlation between the catalytic selectivity and physi- cal properties of the solids which adsorb formic acid has been the subject of systematic studies.'-' Our recent secondary ion mass spectrometry (S1MS)-X-ray photoelectron spectroscopy (XPS) study of formic acid adsorption on copper metal and copper(1) chloride surfaces led to the conclusion that the cata- lytic selectivity of the decomposition of formic acid is closely related to the structure of the surface formate species.45 This paper intends to verify the above assumption by using FTIR and thermal analysis techniques. FTIR of the adsorption layers can provide information on the structure of the surface formate, while thermal analysis allows for a kinetic study of the surface formate pyrolysis by a combination of simulta- neous DTA, TG and EGA.As it will be shown later, the catalytic selectivity and kinetic activity of the surface formate decomposition are indeed dependent on the structure of the surface formate over Cu, CuCl, Cu,O and CuO, with the ionic and monodentate formates favouring dehydration, whereas the bidentate covalent species favour dehydroge- nation. Experimental Preparation and Characterization of the Surface Formates The catalyst samples used in this experiment were all in the powder form as received (Cu, H&W, GR; CuCl, Merck, GR; Cu,O, Merck, GR; CuO, Fluka, AG).Formic acid adsorp- tion was carried out in a gas uptake+volution system where the sample could be evacuated by a rotary pump and the adsorption-desorption could be monitored by a mercury barometer. Normally the sample maintained at a low vacuum was first heated to 320-370 K for 1 h in order to pump off the contaminant gases. Then, after the sample was cooled to room temperature, formic acid vapour (ca. 40 mmHg), which was in equilibrium with liquid formic acid (Merck, GR) at room temperature, was introduced into the system and allowed to come into contact with the sample at room tem- perature for 1 h.The pressure of the system with the vapour was observed to drop, indicating that adsorption had occurred. The adsorbed layers thus prepared were characterized by FTIR on a Shimadzu FTIR-8101, using the pellet technique which involved mixing the finely ground sample and pot- assium bromide powder, and pressing the mixture in an evac- uable die at 8 tonne rn-, of pressure to produce a transparent disk. Strong bands associated with the OCO stretching were observed in the FTIR spectra obtained from these samples, indicating that at room temperature the formic acid was adsorbed mainly as surface formate on all the samples studied. These observations agreed well with those obtained from the earlier vibrational studies under ultrahigh vacuum conditions,".' demonstrating that the surface '~'~9~' formates were fairly stable in air at room temperature.In fact, it was found that samples after being exposed to air for a short period of time and then stored in capped sample bottles for several weeks still produced infrared spectra identical to that of the freshly prepared sample. Therefore, it was possible to study the changes of the surface formate layers during the course of pyrolysis by preparing samples in the gas uptake- evolution system at room temperature, heating to different elevated temperatures in the STA 409 instrument (vide infra) and testing on the FTIR-8101 spectrometer after the samples were cooled to room temperature. 40 scans were taken for each spectrum to obtain a good signal-to-noise ratio.Kinetic Studies of the Surface Formate Pyrolysis The kinetic study of the pyrolysis of the adsorption layers was performed in a simultaneous thermal analysis apparatus, Netzsch STA 409. Ca. 100-200 mg of the sample prepared as described above was placed in an open crucible and was heated from room temperature to CQ. 700 K at a linear rate of 5°C min-' in nitrogen which was used as the carrier gas through the furnace at a flow rate of 100 ml min-'. The change in weight during the experiment and the differential temperature, AT, which was the temperature difference between the sample and the reference material, were contin- uously recorded as a function of the sample temperature, giving TG and DTA curves.EGA by FTIR was performed on the Shimadzu FTIR-8201. The FTIR gas cell was directly connected to the outlet of the STA 409 so that a stable plug- flow state was reached, allowing the outlet gas to be contin- uously tested. As will be shown below, the combination of simultaneous DTA-TG-EGA allows kinetic and catalytic studies of the surface formate decomposition, providing infor- mation on the reaction selectivity, activation energy, calorim- etry and so on. Blank experiments using the solids without the adsorbed overlayer were carried out in parallel to the above described experiments, including both the thermal analysis and FTIR measurements, to make sure that the information obtained was truly from the adsorbed formate overlayers.Results and Analysis Decomposition of the Surface Formate on Cu Thermal Analysis: DTA, TG and EGA Results Fig. 1 displays the DTA-TG results obtained from pyrolysis of the formate-covered copper. At temperatures around 360 K, DTA evidently shows an endothermic heat flow while the TG data exhibit a weight loss. Further increasing the sample temperature to 440 K results in the second weight loss, accompanied by exothermic heat flow. Since the dehydration of the formic acid is known to be endothermic while the dehydrogenation is exothermic,' the TG-DTA data in Fig. 1 clearly demonstrate that the pyrolysis of the surface formate species on Cu favours dehydration at temperatures around 360 K, whereas dehydrogenation becomes predominant at higher temperatures (above 440 K).This conclusion can be also deduced from EGA by FTIR. Fig. 2 displays the infrared spectra of the evolved gas at various temperatures. The infra- red bands at 3500-4000 and 1250-2000 cm-' are known to 0.5 1.5 1.o 0 0.5 7$ -0.5 & 0% i I'd -DTA , -0.5 -1 .o -1 .o -1.5 -1.5 300 320 340360380400420440460480 500 TIK Fig. 1 DTA-TG data for the formate-covered Cu in N, at 5°C min-'. W stands for the total mass loss due to formate decomposi- tion at elevated temperatures and AW/W is the fraction of the formate decomposed. AT is the temperature difference between the sample and reference in mV (thermocouple reading) and represents the heat flow due to formate pyrolysis. J. CHEM.SOC. FARADAY TRANS., 1994, VOL.90 4000 2000 1500 wavenurnber/crn-' Fig. 2 Infrared spectra of the evolved gas from the formate-covered Cu at 310 (a),340 (b),360(c),410 (4,450 (e)and 480K cf) be due to the water vapo~r,~~ whereas the band at 2340-2361 cm-' is due to CO, . Hydrogen, which is a homopolar mol- ecule, is not infrared active. CO is observed normally as a few lines between 2000 and 2250 cm-'. Though the molar absorbability of CO is found to be lower than those of H,O and CO,, note that the amount of CO produced from decomposition of surface formate is lower than that expected based on the stoichiometry of the dehydration of formic acid, which should give rise to equimolar amounts of CO and H,O. Since the ability for the instrument to detect CO has been proven by injecting a trace of CO into the outlet of the STA 409 and obtaining a signal in the corresponding wave- number region, this fact may suggest a reaction scheme for the surface formate decomposition where some of the CO produced from the dehydration may react further with adsorbed OH, adsorbed oxygen or lattice oxygen, as will be shown below. It is obvious from Fig.2 that H,O is the major component in the evolved gas at low temperatures (300-380 K). The water signal becomes weaker at 410 K and then retains high intensity up to 480 K. The CO, signal increases more rapidly at higher temperatures and becomes very pro- nounced at temperatures between 440 and 460 K. At 480 K, the CO, intensity decreases, indicating that the decomposi- tion of the surface formate is almost completed.The conser- vation of the water signal at temperatures above 410 K may result from some side reactions including the proportionation reaction of the surface OH group, which may take place at high temperatures. It is possible to estimate the activation energy of the surface formate pyrolysis using the TG data obtained. Decomposition of the adsorbed formate is known to be a first-order reaction., It has been reported47 that for a first-order reaction the fraction of the reactant decomposed at temperature T, AW/W,and the activation energy of the reac- tion, E, ,have the following relationship: log[-log(l -AW/W)/T2]= log[(AR/aE,)(l -2RT/E,)] -EJ2.3RT where A is the frequency factor of the reaction, a the linear heating rate and R the molar gas constant. In Fig.3, Y = log[ -log(l -AW/W)/T2], calculated using the TG data in Fig. 1, is plotted against 1/T. The data points are best fitted with straight lines, L1 and L2, yielding an activation J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 -5.2 - \ 1 -5.4 -5.6 .1 'rcL2 I -2% I -5.8 Y -6.0 -6.2 -6.4 -6.6 -6.8 0.0021 0.0023 0.0025 0.0027 0.0029 KIT Fig. 3 Y us. 1/T plot for the formate-covered Cu: Y = log[-log(1 -AW/W)/T2],where AW/W is the fraction of the formate decomposed and T the pyrolysis temperature. Data (H)are obtained based on the TG results in Fig. 1 ;the data are best-fitted by two straight lines (L1 and L2), giving the activation energy E = 7 kcal mol-' at low temperatures and E = 15 kcal mol-' at high temperatures.energy of 7 kcal mol-' for the formate pyrolysis over Cu at 340-410 K and 15 kcal mol-' at 410-460 K. Note that the activation energy obtained as described above is based on the loss in sample weight (i.e. the consumption of reactant) and is therefore the apparent activation energy for the overall reac- ti~n.~* FTIR Spectra of the Adsorbed Layers Fig. 4 shows the infrared spectra for the formate-covered Cu surfaces at different temperatures. Included at the top of the figure, Fig. 4(a) is obtained from the bulk copper@ formate in the wavenumber region between lo00 and 2000 cm-'. Similar to the known spectrum for the bulk sodium formate which has IR bands at 1567, 1377 and 1366 cm-', the bands at 1600, 1380 and 1320 cm-' may be assigned to v,,(OCO), 6(CH) and vs(OCO), respecti~ely.'~*~~Fig. 4(b) is for the formate-covered Cu surface at room temperature. The most c 4000 2000 1500 1000 wavenumber/cm-' Fig.4 Infrared spectra of the formate adsorbed on Cu at 300 (b), 380 (c), 480 (d) and 720 K (e).Spectrum (a) is obtained from bulk copper(1) formate and included for comparison. 357 prominent bands in this spectrum are the asymmetric OCO stretching vibration, va,(OCO),in the 1600-1640 cm-' region and the symmetric vibration, v,(OCO), at 1320-1380 cm-'. Note also that the v,,(OCO) band is composed of two well resolved peaks at 1600 and 1630 cm-' while the v,(OCO) band, which is peaked at 1350 cm-',is asymmetric in shape, having a shoulder a 1320 cm-'.Besides these peaks there is also a signal observable at 1380 cm-'.This appears to imply that there may exist more than one chemisorption state for surface formate on polycrystalline copper surfaces. The bands at 1600, 1380 and 1320 cm-' can be ascribed to the ionic complex according to Fig. 4(a)for the bulk copper@) formate. On the other hand, the bands at 1630 and 1350 cm-'appear to be due to the bidentate surface formate since previous vibrational studies have observed bands at 1650 and 1348 cm-' for formate-Cu( 110) surface13 and at 1640 and 1330 cm-' for formate-Cu(lOO),'O respectively, and the geometric structure for the formate on these two well defined surfaces is either a chelating bidentate on Cu(ll0) or a bridging biden- tate on Cu(100),as known from studies using extended X-ray absorption fine structure (EXAFS),'-'*'' see Fig.5(a) and (b). Further support may be found from the study of the formate- covered GaAs(ll0) surface, where the bands at 1640 and 1350 cm-' are also attributed to the bidentate formate.26 When the sample temperature is increased to 380 K, the OCO infra-red band at 1600 cm-', which is assigned to the ionic species, evidently decays while the 1635 cm-' band remains fairly strong [see Fig. 4(c)J.This appears to indicate that the ionic surface formate is less stable and decomposed in the lower temperature range than the bidentate formate which is mostly decomposed at temperatures higher than 380 K.At ca. 500 K, almost all of the signals due to the formate dimin- ish, which is indicative of the completion of the formate pyrolysis and is in good agreement with that reported from temperature-programmed desorption (TPD)' and ultraviolet photoelectron spectroscopy (UPS)4 studies. The bands at 3453,2920, 2360, 1070, 760,670 and 620 cm-', which can be assigned to v(OH), V(CH), v(CO,), n(CH), s(OCO), 74C0,) and v(Cu'-0), respectively, according to the literat~re,'~*~~ may be observed at some temperatures but will not be dis-cussed in detail. three-electron H one-electron donor 0 R I cu-0' monatomic chelate monodentate with additional bridging (9) (h) Fig.5 Structural models proposed in the literature' 7,4531-61 for the bidentate bridging formate on Cu(100) (a), bidentate chelate formate on Cu( 1 10) (b),hydrogen-bonded monodentate formate on CuCI(11 1) (c), and for the carboxylato copper complexes, (d)-(i). Surface Formate Decomposition on CuCl DTA, TG and EGA Results The TG profile of the formate-CuC1 in Fig. 6 shows two weight losses at the temperatures around 440 and 510 K, respectively;while DTA displays two endothermic heat flows in the corresponding temperature ranges. In agreement with the DTA data, EGA by FTIR shows that water vapour is the main product from the pyrolysis of the formate adsorbed on CuCl in the whole temperature range studied, with some CO, being detected mainly at temperatures between 460 and 500 K (see Fig. 7).The activation energy, estimated using the same procedures as described above, is 5 kcal mol-' between 380 and 490 K, and 19 kcal mol-' at 490-560 K. FTIR Spectra of the Adsorbed Layers Fig. 8 illustrates the FTIR spectra of the formate-covered CuCl at a number of different temperatures. Between room temperature and 390 K, the asymmetric OCO stretching vibration band is peaked at 1600 crr-', while in the sym- metric stretching vibration region there are two peaks at 1380 and 1320 cm-', respectively [see Fig. 8(a)]. Following the assignment for the formate-Cu system, these infrared bands 1 10.5oi II 1 -1.5 I --1.57 300 350 400 450 500 550 600 T/K Fig. 6 DTA-TG profiles for the formate-covered CuCl 4000 3000 2000 1750 1500 1250 wavenumber/cm-' Infrared spectra of the evolved gas from the formate-covered Ckl at 310 (a),340 (b),370 (c),470 (4,500 (e)and 520 K cf) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 4000 2000 1500 wavenumber/cm-' Fig. 8 Infrared spectra for the formate adsorbed on CuCl at 300 an2 390 K (a),510-K(b) are due to the ionic surface formate. After heating to 510 K the infrared bands at 1320 and 1380 cm-', which are due to the OCO stretching and the CH in-plane deformation, disap- pear and the peak at 1600 cm-', which is assigned to the asymmetric OCO stretching, is now shifted to higher wave- number, located at 1610 cm-'. These facts suggest that the adsorbed formate on the CuCl surface at high temperatures may exist mainly in the monodentate configuration [Fig.5(c)], as previously suggested in our SIMS-XPS study.45 Note that the OH band in Fig. 8 appears to have a tail at its lower wavenumber side. The tail seems to be indicative of the presence of hydrogen bonding since the stretching vibration of the Cl-H bond is located in this energy range (ca. 2991 cm-l ).46 Close values for the asymmetric OCO vibration of the monodentate surface formate have been reported on Cu/SiO, (1640 ~m-'),~~ Cu(100) at 100 K (1640 cm-'),lo Ru(OO1) (1600 cm-1),22*24 A120, (1621 and Pt(ll1) (1620 cm")." Note that there is no band in the region between 1320 and 1380 cm-' as expected for the symmetric OCO vibration. The absence of the symmetric OCO vibration (1320-1380 cm-') has been reported on the formate-covered Cu-deposited KBr surface.It is attributed to the roughness of the surface and explained in terms of the selection rule for dipolar excitations. l4 Alternatively, the absence of the asym- metric OCO vibration in the region between 1560 and 1680 cm-' for a bidentate formate was reported on Cu(l10),'3 Cu(100)" and R~(00l),~~ and was interpreted on the basis of the selection rule. It appears that the change of the orienta- tion of the dipole moment with respect to the electric field may also be the most probable consideration for the absence of the symmetric OCO band in the case of formate-CuC1. Nevertheless, since the CuCl used in the present experiment was in powder form and the infrared light was unpolarized, the selection rule may not be applicable and the reason for the absence of the symmetric vibration is still not very clear.Surface Formate Decomposition on Cu,O DTA, TG and EGA Results The TG data in Fig. 9 show a weight loss occurring at tem- peratures between 380 and 450 K. This is an endothermic process as shown by the DTA profile in the same figure. EGA by FTIR indicates that the main product of the formate pyrolysis in this temperature range is water vapour. A signifi- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 $ -0.5-d -1.5 300 350 400 450 500 -0.5 550 T/K Fig. 9 DTA-TG data of the formate-covered Cu,O cant weight loss follows immediately at temperatures between 450 and 500 K, with an evident heat release.EGA shows that in this temperature range carbon dioxide is the main product from the pyrolysis of the formate adsorbed on Cu,O. The activation energy is estimated to be 9 kcal mol-I at tem- peratures between 380 and 450 K and 24 kcal mol- ' at 450- 520 K. FTIR Spectra of the Adsorbed Layers The FTIR spectra of the formate-covered Cu,O are displayed in Fig. 10. The Cul-0 stretching vibration is observed at 620 cm-' as a huge band. It does not change during pyrol- ysis and is therefore not presented in the figure. Similar to the spectra of the formate-Cu, the v,,(OCO)band consists of two rather weak and can be observed only as a small bump. When the temperature is higher than 510 K all the bands due to the formate ion disappear, indicating that the decomposi- tion is complete.Surface Formate Decomposition on CuO DTA, TG and EGA Results The DTA-TG data of the formate-covered CuO are illus- trated in Fig. 11 where a small weight loss occurs around 350 K. This is an endothermic process as shown by the DTA profile in the same figure. EGA by FTIR indicates that the main product of the formate pyrolysis in this temperature range is water vapour. A significant weight loss, with an evident exothermic heat flow, is observed at temperatures between 460 and 530 K. EGA shows that carbon dioxide is the main component in the evolved gas at high temperatures. The activation energy estimated based on the TG data is 14 kcal mol-' at 320-360 K and 39 kcal mol-' at 460-530 K.FTIR Spectra of the Adsorbed Layers Fig. 12 illustrates the FTIR spectra of the formate-covered CuO. At room temperature, the formate-CuO surface gives a va,(OCO)band centred at 1600 cm-'and an overall v,(OCO) band with three peaks at 1380, 1360 and 1330 cm-'.Increas-ing the temperature to 420 K results in the decay of the t3 peaks at 1600 and 1635 cm-' while the v,(OCO)band is composed of the peaks at 1350 and 1380 cm-',respectively. The bands at 1600 and 1380 cm-' may be attributed to the -0.5 QO::DTAionic complex, and those at 1635 and 1350 cm-' to the bidentate formate. With increasing temperature, the 1600 cm-' band decreases faster than the 1635 cm-' band as -1 .o shown in Fig.lqb). At 470 K, only the peak at 1630 cm-' is observed [see Fig. lqc)]. The 1350 cm-' band becomes -1.5 300 350 400 450 500 550 600 TIK Fig. 11 DTA-TG profiles of the formate-covered CuO 4000 2000 1500 4000 2000 1500 wavenumber/cm-wavenumber/cm-Fig. 10 Infrared spectra of the formate adsorbed on Cu,O at 300 Fig. 12 Infrared spectra of the formate adsorbed on CuO at 300 (a), (a),400 (b) and 470 K (c) 420 (b) and 520 K (c) 360 formate bands, with the absorption band at 1600 cm-' decreasing more than that at 1630 cm-' [see Fig. ll(b)]. Based on the assignments for the formate-Cu surface, the band at 1600 cm-' may be due to the ionic surface complex while the 1630 cm-' band is attributed to the bidentate formate.The ionic complex is decomposed at relatively low temperatures, while the bidentatic formate may be decom- posed mostly at higher temperatures. The Cu"-O stretching vibration is observed at 530 cm-' but is not presented in this figure. Discussion As summarized in Table 1, three types of adsorbed formate species, ionic, monodentate and bidentate formate, have been identified by the infrared study as the main products of formic acid adsorption on Cu, CuC1, Cu,O and CuO surfaces at room temperature. The ionic formate, which gives the characteristic infrared bands at 1600 and 1320 cm-',is found to exist on all the surfaces studied in the low-temperature range (<490 K) and is easily decomposed at elevated tem- peratures, forming water as the main product.The bidentate formate, with infrared bands at 1630 and 1350 cm-I, is the most abundant species observed on Cu, Cu,O and CuO sur- faces, based on the TG data, and is decomposed at high tem- peratures (around 500 K), producing mainly carbon dioxide. The adsorbed species observed on the CuCl surface at high temperatures exhibits only an asymmetric OCO band at 1610 cm-' with little infrared absorption in the energy region for the symmetric OCO vibrations and is considered as the monodentate formate whose pyrolysis favours the production of water. These observations suggest that the selectivity of the surface formate decomposition may not really depend on the type of catalyst (i.e. whether metal, transition metal oxide or other metal oxide) as often proposed,' but, may actually depend on the structural properties of the formate formed on these materials, i.e.whether the surface formate is ionic or covalently bonded, and whether it is monodentate or biden- tate in geometric configuration. Note that in Table 1 the acti- vation energy of formate pyrolysis is lowest (ca. 5-14 kcal mol-') for the ionic formate, intermediate (19 kcal mol-') for the monodentate and highest (15-39 kcal mol-') for the bidentate surface formate. In particular, for the bidentate formate, the pyrolysis activation energy is shown to be the highest on CuO (39 kcal mol-') and lower on Cu,O and Cu (24 and 15 kcal mol-', respectively). Since the electron charge transfer from the formate ligand to the copper atom/ion takes place to an increasing extent in the fol- lowing sequence: ionic < monodentate (one-electron donor)" < bidentate (three-electron donor)' ' formate and Cuo < Cu' < Cu", it appears that the kinetic activity of the pyrolysis decreases (i.e.the activation energy increases), with increasing electron charge transfer from the formate to the metal atom/ion. This agrees well with our previous SIMS- XPS study4' where the greater electron charge transfer is J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 considered to be responsible for the more stable OCO bonding on Cu compared with that on CuCl. UPS studies4** and theoretical molecular orbital calc~lations~~ indicated that the highest occupied molecular orbitals of the formate ion are la,, 6a,, 4b,.These orbitals are mainly due to the oxygen atoms and are most strongly involved in bonding to metal, donating electron charge to the surface. In particular, the la, orbital, which has .n symmetry, is essentially the anti- symmetric combination of the 0 2p, orbitals and is thus antibonding with respect to the 0-C-0 bond. Therefore, the removal of electron density from this orbital would result in the relative strengthening of the OCO bonding.l3 In con- trast, less electron charge transfer is involved in the ionic bonding interaction or in the monodentate configuration, leading to a less stable 0-C-0 bond. This conclusion can find support both from the previous SIMS results4' and from the present FTIR studies.In the SIMS studies the OCO adspecies is observed to be more stable on Cu than on CuCl, while the FTIR results show that the 0-C-0 stretching frequencies (and thus the OCO bonding) are reduced from 1630/1350 for the bidentate formate to 1610/-for the mono- dentate species and 1600/1320 cm-' for the ionic complex. According to the above conclusion, the decomposition of a bidentate formate would favour dehydrogenation owing to its relatively stable OCO bonding, while a further splitting of the OCO bond may occur with smaller activation energy for the ionic or monodentate formates (because of the less stable OCO bonding), resulting in the predominant dehydration. This can be best described in the following reaction scheme, Scheme 1, which is modified from a reaction mechanism sug- gested by.1glesia and Boudart :'3 /H.* CO, -+CO,(g) + H(ads) HCO,(ads) He .(CO).. .O + CO(ads) + OH(ads) CO(ads) CO(g) CO(ads) + O(ads or lattice) eCO,(g) OH(ads) + CO(ads)e CO,(g) + H(ads) OH(ads) + OH(ads)e H,O(ads) + O(ads) H(ads) + OH(ads) eH,O(ads) H,O(ads) eH,O(g) H(ads) + H(ads) ~t H,(g) Scheme 1 The proposed reaction scheme involves an activated complex as the reaction intermediate, in which the C-H bond is weakened through coordination to the catalyst Table 1 Kinetic activities for surface formate decomposition on Cu, CuCI, Cu,O and CuO surfaces" ionic (1600, 1 320)b monodentate (1610, -)* bidentate (1630, 1350)b AH E.3 AH Ea AH Ea surface T/K /kcal mol-' /kcal mol-' T/K /kcal mol-' /kcal mol-' T/K /kcal mol-' /kcal mol-' cu 340-430 endo 7 ---430-480 exo 15 CuCl 380-490 endo 5 490-560 endo 19 ---CU,O 380-460 endo 9 ---460-520 exo 24 CUO 320-360 endo 14 ---460-530 exo 39 ~~ T,decomposition temperature; AH, heat flow; E,, activation energy; endo, endothermic; exo, exothermic.* IR bands in cm- '. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 361 surface. If the HCO,(ads) adopts a bidentate configuration, the OCO bonding is relatively stable and the decomposition proceeds mainly through reaction (1) which leads to the for- mation of gaseous CO,. On the other hand, the ionic/ monodentate HCO,(ads) in which the OCO bonding is relatively unstable would follow reaction (2) forming CO(ads) 6 7 8 J.Stohr, J. L. Gland, W. E. Eberhard, D. Outka, R. J. Madix, F. Sette, R. J. Koestner and U. Dobler, Phys. Rev. Lett., 1983, 51, 2414. A. Puschmann, J. Haase, M. D. Crapper, C. E. Riley and D. P. Woodruff, Phys. Rev. Lett., 1985,54, 2250. Th. Lindner, J. Somers, A. M. Bradshaw and G. P. Williams, Surf: Sci., 1987, 185, 75, and references therein. and OH(ads) which then further react giving water as the main product. In this scheme the C-H cleavage is viewed as the rate-determining Since the O-H bond of the bidentate HCO,(ads) is normally thought to be in a molecu- lar plane perpendicular to the surface and is away from the ~urface,~~,~",~~a tilting of the HCO,(ads) species may be 9 10 11 P. Hofmann, C. Mariani, K. Horn and A. M. Bradshaw, in Proc. 4th Int.Conf: on Solid Surfaces and 3rd European Cont on Surface Science, ed. D. A. Degras and M. Costa, Paris, 1980, p. 541. B. A. Sexton, Surj Sci., 1979,88, 319. L. H. Dubois, T. H. Ellis, B. R. Zegarski and S. D. Kevan, Surf: Sci., 1986, 172, 385. required in order for the C-H bond to approach the surface. It appears that less energy may be needed for the decomposi- tion of a monodentate formate, which has only one oxygen atom bound to the surface, allowing for the free rotation of the C-H bond toward the surface, than of a bidentate formate with two metal-oxygen Thus, it is under- 12 13 14 15 P. Sen and C. N. R. Rao, Surf. Sci., 1986,172,269. B. E. Hayden, K. Prince, D. P. Woodroof and A. M. Bradshaw, Surf: Sci., 1983, 133, 589. M. Ito and W. Suetaka, J. Phys.Chem., 1975,79, 1190. R. J. Madix, in Vibrations at Surfaces, ed. C. R. Brundle and H. Morawitz, Elsevier, Amsterdam, 1983, p. 25; Surf: Sci., 1979, 89, 540, and references therein. standable that a higher pyrolysis activation energy may be found for the surface formate in the bidentate configuration than for the ionic/monodentate form. Cu, Cu,O and CuO are known to have very different crystal structures, but the surface formates on these catalysts are all dominated by the bidentate species as inferred from 16 17 18 P. Hofmann, S. R. Bare, N. V. Richardson and D. A. King, Surf: Sci., 1983, 133, L459. D. A. Outka and R. J. Madix, in Chemistry and Physics of Solid Surfaces, ed. R. Vanselow and R. Howe, Springer Verlag, Berlin, 1986, vol. 6, p. 133, and references therein.B. A. Sexton, A. E. Hughes and N. R. Avery, Surf. Sci., 1985, 155, 366. infrared studies. This similarity may be derived from the pres- ence of many possible structural configurations for the biden- tate formate ligand. On well defined copper surface^'^,'^ the formates are known to have either chelating or bridging bidentate configurations [Fig. 5(a) and (b)].In most of the 19 20 21 F. C. Henn, J. A. Rodriquez and C. T. Campbell, Surf: Sci., 1990, 236,282. C. Egawa, I. Doi, S. Naito and K. Tamaru, Surf: Sci., 1986, 176, 491. S. W. Jorgensen and R. J. Madix, J. Am. Chem. SOC., 1988, 110, 397. carboxylato Cu"-Cu' complexes, whose crystal structures 22 F. Solymosi, J. Kiss and I. Kovacs, J. Phys. Chem., 1988,92, 796; have been determined by X-ray or neutron diffraction, the carboxylate ions may exist as chelating bidentate ligands, i.e.as a bridging bidentate ligand in a syn-syn, anti-anti, anti-syn configuration [see Fig. 5(d)-(l1)].~"-~"The Cu-Cu distance within those carboxylato copper complexes ranges from 2.64 to 5.77 A. The above features found for the carboxylato copper complexes appear to provide flexibility, allowing a surface formate to adopt a bidentate configuration among several possible coordination geometries and to fit the surface structures over the Cu, Cu20 or CuO catalysts. Among the carboxylato Cu"-Cu' complexes with known crystal struc- tures there are a few exceptions where the carboxylato group 23 24 25 26 27 28 29 30 Surf. Sci., 1987, 192, 47. F. Solymosi and I. Kovacs, Surf: Sci., 1991,259,95.M. Chtats, P. A. Thiry, J. P. Delrue, J. J. Pireaux and R. Caudano, J. Electron Spectrosc. Relat. Phenom., 1983, 29, 293. B. H. Toby, N. R. Avery, A. B. Anton and W. H. Weinberg, J. Electron Spectrosc. Relat. Phenom., 1983, 29, 3 17. R. Matz and H. Luth, Surf. Sci., 1982,117, 362. M. A. Karolewski and R. G. Cavell, Surf: Sci., 1989, 210, 175; 219, 261. B. F. Lewis, M. Mosesman and W. H. Weinberg, Surf. Sci., 1974, 41, 142. J. Klein, A. Leger, M. Berlin, D. Defourneau and M. J. L. Sang- ster, Phys. Rev. B, 1972, 7, 2336. K. Yamashita, S. Naito and K. Tamaru, J. Catal., 1985,94,353. is bound in the monodentate configuration [Fig. 5(i)].65 The fact that the surface formate adsorbed on CuCl is observed to function as a monodentate ligand [Fig.5(c)], different from those over Cu, Cu,O and CuO, may be related to the rela- tively high ionic character of the bonding in CuCl and the 31 32 33 K. Lui, M. Vest, P. Berlowitz, S. Akhter and H. H. Kung, J. Phys. Chem., 1986,90,3183. J. M. Vohs and M. A. Barteau, Surf: Sci., 1986, 176, 91; 1988, 197, 109. Y. Noto, K. Fukuda, T. Onishi and K. Tamaru, Trans. Faraday SOC., 1967,63,2300; 3072. presence of the highly electronegative anion in the vicinity of 34 A. Ueno, T. Onishi and K. Tamaru, Trans. Faraday SOC., 1970, the copper ions, which may lead to the formation of hydro- gen bonding, favouring the monodentate configuration. 35 66,756. M. Bowker, H. Houghton and K. C. Waugh, J. Chem. SOC., Faraday Trans., 198 1,77,3023. We thank Prof. E.I. Solomon of Stanford University and Dr. J. Y. Lee for critical reading and valuable comments prior to the submission of this paper. Experimental assistance from Mr. K. W. Wong is gratefully acknowledged. 36 37 38 D. G. Walmsley, in Vibrational Spectroscopy of Adsorbates, ed. R.F. Willis, Springer-Verlag, Berlin, 1980, p. 67. A. Deluzarche, J. P. Hindermann, R. Kieffer and A. Kienne- mann, Rev. Chem. Int., 1985,6, 255, and references therein. P. A. Taylor, P. B. Rasmussen and I. Chorkendorff, J. Phys.: Condens. Matter, 1991,3, s59. 39 K. Waugh, Catal. Lett., 1990, 7, 345. 40 G. C. Chinchen, K. C. Waugh and D. A. Whan, Appl. Catal., 1986, 25, 101. References P. Mars, J. J. F. Scholten and P. Zwietering, in Advances in Catalysis, ed. D. D. Eley, H. Pines and P.B. Weisz, Academic Press, New York, 1963, vol. 14, p. 35, and references therein. J. Fahrenfort, L. L. Van Reyen and W. M. H. Sachtler, in Proc. Symp. Mechanism of Heterogeneous Catalysis, ed. J. H. Boer, 41 42 43 44 R. G. Herman, in New Trends in CO Activation, ed. L. Guczi, Elsevier, Amsterdam, 1991, p. 265, and references therein. V. Ponec, in ref. 41, p. 117, and references therein. K. Klier, Adv. Caral., 1982,31,243. G. W. Bridger and M. S. Spencer, in Catalyst Handbook, ed. M. V. Twigg, Wolf, London, 2nd edn., 1989, p. 446, and refer- ences therein. Elsevier, Amsterdam, 1960, p. 23. 45 J. Lin, A. T. S. Wee, A. C. H. Huan and K. L. Tan, Surf: Sci., P. Mars, in ref. 2, p. 49. R. W. Joyner and M. W. Roberts, Proc. R. SOC.London, A, 1976, 350, 107. D.A. Outka, R. J. Madix and J. Stohr, Surf. Sci., 1985,164,235. 46 47 1993, 285, 31. K. Nakamoto, in Infrared and Raman Spectra of Inorganic and Coordination Compounds, 4th edn., and references therein. A. W. Coats and J. P. Redfern, Nature (London), 1964,201,68. 362 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 48 M. A. Newton, S. M. Francis and M.Bowker, in Catalysis and Surface Characterisation, ed. T. J. Dines, C. H. Rochester and J. 58 59 D. A. Langs and C. R. Hare, Chem. Commun., 1967,890. B. W.Skelton and T. N. Water, J. Chem. SOC., Dalton Trans., 49 Thomson, Royal Society of Chemistry, Cambridge, 1992, p. 165. G. J. Miller, C. H. Rochester and K. C. Waugh, J. Chem. SOC., Faraday Trans., 1991,87, 1491. 60 61 1972,2133. G. A. Barclay and C. H.L. Kennard, J. Chem. SOC., 1961,3289. G. B. Drew, D. A. Edwards and R. Richards, J. Chem. SOC., 50 51 N. R. Avery, Appl. Surf: Sci., 1983,14,149. F. A. Cotton and G. Wilkinson, in Advanced Inorganic Chem- 62 Chem. Commun., 1973,124. D. A. Edwards and R. Richards, J. Chem. SOC., Dalton Trans., 52 istry, Wiley, New York, 4th edn.,1980, p. 63. S. D. Peyerimhoff and R. J. Buenker, J. Chem. Phys., 1969, 50, 1846. 63 1973,2463. K. Okada, M. I. Kay, D. T. Cromer and I. Almodovar, J. Chem. Phys., 1966,44, 1648. 53 E. Iglesia and M. Boudart, J. Phys. Chem., 1991,9!5,7011. 64 A. G. Massey, in The Chemistry of Cu, Ag and Au, Pergamon 54 J. Catterick and P. Thornton, Adv. Inorg. Chem. Rudiochem., 1977,20, 291. 65 Press, Oxford, 1973, p. 1. K. Anzenhofer and L. N. A. Ten Rouwelaar, Rev. Trav. Chim., 55 C. Oldman, Progr. Inorg. Chem., 1968,10,223. 1967,86,801. 56 57 G. B. Deacon and R. J. Philips, Coord. Chem. Rev., 1980,33,227. J. N. Van Niekerk and F. R. L. Schoening, Acta Crystallogr., 1953,6, 227. Paper 3/02662K; Receioed 1lth May, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000355

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(2), 363-367 Heat of Water Chemisorption on a-Al,O, at 200-400 OC Pier Francesco Rossi,* Giovanni Oliveri and Marta Bassoli lstituto di Chimica (CNR),Facolta di lngegneria , Universita di Genova , Fiera del Mare, P.1e.J.F. Kennedy, Pad. D, 16129 Genova, Italy The differential heats and the adsorption isotherms of water on a-Al,O, have been measured at five different temperatures in the range 200-400 "C using a microcalorimetric-volumetric apparatus. The amounts of water adsorbed and integral heats decreased with increasing temperature. At all temperatures the differential heats decreased with increasing amount of chemisorbed water, suggesting an ordinary surface heterogeneity. At very low coverages, the heats of adsorption at all temperatures are >200 kJ mol-', at higher coverages they are < 200 kJ mol-'.Two types of adsorption were considered: dissociative adsorption and coordination of molecular water. We also carried out some thermokinetic investigations on heat emission as a function of time at increas- ing equilibrium water pressures at 200°C and 400"C,and at different adsorption temperatures at 0.05 Torr equilibrium water pressure. The complex interactions of water vapour with the surface of different oxides have been studied extensively by many methods, such as gravimetric,' infrared2*, and calorimetric The chemisorption of water vapour on dehy- droxylated polymorphous aluminas has been studied at dif- ferent The adsorption process, at room temperature, consisted of both physisorption and chemisorp- tion.However, at higher temperature, the chemisorption of water vapour prevailed and it occurred by two different mechanisms : irreversible dissociative and reversible coordi- native adsorption, with different chemisorbed species of molecular water. The literature data for the heats of adsorp- tion of water on aluminas at temperatures >2OO"C appear really poor. For this reason, we carried out adsorption mea- surements of water on a-A120, in the temperature range 200-400 "C. The purpose of the study is to probe the distribution of different strength sites on a-Al,O, surfaces as a function of temperature. Experimental Materials The sample of metal oxide (ca. 2 g) used was a-Al,O, (99.999%) supplied by Aldrich (USA).The sample was immersed in the liquid water at room temperature for 48 h to give maximum hydration and then dried at 110°C. The spe- cific surface area of the a-Al,O, ,degassed at 500"C and in a vacuum of lo-' Torr, was found to be 15.37 m2 g-' on the basis of N, adsorption data. Microcalorimetric Apparatus The microcalorimetric assembly consisted of a heat-flow microcalorimeter (Calvet type, measuring temperature with an accuracy of fO.l "C) connected to a volumetric apparatus to measure adsorption. The measurements were carried out at 200, 250, 300, 350 and 400°C. The microcalorimeter Cali- bration was performed using a standard cell of alumina with Pt resistance, supplied by Setaram (Lyon, France).The reference cell and the laboratory cell, supplied by Glass-Emery (Genoa, Italy) were made of quartz, equipped with inner walls to reduce the geometrical volume. The water was purified by double distillation after passage through an ion-exchange resin. The microcalorimetric measurements were carried out by means of an IBM-AT computer, interfaced to the micro- calorimeter with A/D converter. The apparatus scheme has been shown elsewhere.' Prior to adsorption, the sample of ct-A120,, was degassed under a vacuum of low5Torr at 500°C for 5 h (to eliminate most of the chemisorbed water without possible sintering of the alumina surface)." After this pretreatment, the calorimetric cell, containing a-A120, under vacuum, was inserted into the microcalorimeter at 200°C and was connected to the volumetric Pyrex glass line for vapour adsorption.A first water adsorption up to surface saturation of the sample was performed (run 1). The alumina sample was then degassed again inside the micro- calorimeter and a second water adsorption was performed at 200°C to measure the reversible fraction of the adsorbate at the same temperature (run 2). The procedure was repeated on other samples at 250, 300, 350 and 400°C to give the reversible fractions of the adsorb- ate at these temperatures. Processing of the microcalorimetric data with a computer provides a good characterization of the water adsorption process. The peak areas and volumetric isotherms give the total heat and total adsorption, respectively, as a function of pressure.The integral heats for the successive increments are analytically fitted and the differential heats, qdiff, are then obtained by differentiation of the integral function with respect to the amount adsorbed n,. A study of the thermo- kinetics of heat evolution provides additional insight into the adsorption mechanisms. Results and Discussion Fig. l(a) shows the primary calorimetric isotherms (run 1) and Fig. l(b) the secondary calorimetric isotherms (run 2) at the five adsorption temperatures. Fig. 2 shows the corre-sponding volumetric isotherms. An increase in experimental temperature always involved a decrease in both heat released and amounts adsorbed. The first run appears to consist of an irreversible chemisorption followed by a pressure-dependent branch, because, at very low pressure, the water vapour was irreversibly chemisorbed on the strongest (Lewis acid) sites, while at higher pressure, the water was adsorbed reversibly on intermediate sites.In Fig. 2(c) we show the irreversibly adsorbed amounts of water (run 1 -run 2) as a function of experimental tem- perature. We note that this amount decreases slightly with increasing temperature (i.e.from 0.47 pmol m-2 at 200°C to 0.22 pmol m-at 400 "C).The volumetric isotherms of run 1 364 0.50 N 0.25 E OF--- 0.251 II Ih \ \" I I I 0 1 2 3 4 5 PP-orr Fig. 1 Calorimetric isotherms: (a) run 1, empty symbols; (b) run 2, filled symbols.0,200; 0,250; A, 300; 0,350 and V, 400°C. seem to show a Langmuir-like trend and none of the iso- therms has the BET type I1 behaviour as we observed, for example, for the physisorption of H20 on high-energy sur- faces." We observed that a part of the irreversibly bound water at 200°C becomes reversibly bound at higher tem- peratures. The initial adsorption decreases with increasing temperatures and at all temperatures, the heat released and amounts adsorbed from run 1 are higher than those for the reversible adsorption (run 2). Such data confirm the complexity of the chemisorption of water vapour on a-A120, at high temperatures. In particular, we think that in the temperature range 200-400"C, the simultaneous existence of different adsorption mechanisms is present.In fact when the sample of a-Al,O, is degassed under vacuum at different temperatures, the surface hydroxy groups react to form oxygen bridges and water according to scheme (1):l2 0 1 2 3 4 5 p/rorr Fig. 2 Volumetric isotherms: (a) run 1, empty symbols; (b) run 2, filled symbols; (c) run 1 -run 2. Symbols as in Fig. 1. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 140 100 EE--. .-0 c .-> -Q 60 \ 20 II I I I 15 75 135 195 255 time/mi n Fig. 3 Heat emission peaks as a function of increasing water vapour equilibrium pressures at 200°C: (a) 0.001, (b) 0.1, (c) 0.2 and (d) 1.6 Torr. (y axis refers to the galvanometer deviation of the micro- calorimeter assembly.) On contact with water vapour, the oxygen bridges react to re-form surface OH groups (dissociative chemisorption), according to reaction (2): HH The coordinative chemisorption of water vapour onto unsaturated aluminium ions can be presented by reaction (3)9 ,*I, I Such adsorption mechanisms, lead to the irreversible and reversible phases, and their balance depends on both cover- age and temperature.We think that at very low coverages, irreversible chemisorption due to the dissociative mechanism of water adsorption is present and at relatively higher cover- ages reversible chemisorption due to the coordinative mechanism is prevalent. Thermokinetic Study Semiquantitative information may be obtained from the peak shapes of heat emission from water adsorbed on alumina.In Fig. 3 we show some significant heat emission peaks with increasing coverage (increasing equilibrium pressure) of J. CHEM. SOC. FARADAY TRANS., 1993, VOL. 90 water vapour at 200°C as a function of time.8 Each peak corresponds to admission of a single vapour dose, and there- fore to a point on the calorimetric isotherms [Fig. l(a)]. Moreover Fig. 3(a)shows that the initial adsorption appears to be activated. The heat evolution rate varies with coverage in a complex manner: the peak shape does not show a regular increase in heat emission rate with coverage. Two factors appear to be in competition: the general trend towards instantaneous adsorption and a significant increase in activation energy with coverage [Fig.3(b)and (c)].At high coverages, the peak shape becomes close to that typical of reversible phenomena [Fig. 3(d)]. In conclusion (a)represents a slow, clearly activated, irreversible adsorption which is also apparent from the slow (many minutes) decrease in the gas pressure on the adsorbent whereas (d), represents a typical reversible phenomena of fast heat evolution, (b) and (c) show the presence of both phenomena. In Fig. 4 we show the rate of heat emission at 400°C for different coverages. The conclusions are similar to those of Fig. 3, but the emission peak shapes vary more slowly from (4to (4.In Fig. 5, we show the rate of heat emission at the same equilibrium pressure (0.05 Torr) (i.e. same coverage) at five different temperatures.The peak shape varies from (a) to (e): that of (a), 200°C is rounded (activated process) and they become sharper with increasing temperature (fast heat evolution). The increase of temperature, at the same equi- librium pressure, seems to oppose the activated irreversible process and to promote fast reversible adsorption. Surface Rehydroxylation In Fig. 6 the differential heats of adsorption on the outgassed surface (run 1) are reported as a function of coverage for the five temperatures and in the inset the variation of the integral heat of adsorption (Qi,J with coverage is shown. Such trends (at very low coverage the adsorption mainly represents the irreversible fraction) are characterized by five curves whose convexity decreases as the temperature in-creases.In fact, the distribution of the strengths of the active sites changes over the temperature range 200-400°C. On introduction of the first dose of water vapour to a-Al,O,, the average value of the differential heat is 270 kJ mol- ' for the five temperatures. This value decreases very steeply to near the heat of water vaporization (43.8 kJ mol-','3 or 10.51 kcal loo i 365 14C 100 E E 1 0.-c .-m > U 60 20 15 75 135 195 255 time/min Fig. 5 Heat emission peaks as a function of adsorption temperature at the same water vapour equilibrium pressure (0.05 Torr): (a)200, (b) 250, (c) 300,(d)350 and (e) 400 "C mol-').'4 At very low coverage, the value of qdiff (270 kJ mol-'), can be assigned to irreversible dissociative chemi- sorption.At higher coverage, &iff = 200-90 kJ mol- ', reversible coordination of molecular water is present. This is a region of gradually decreasing heats, representing adsorption on sites of intermediate strength. Finally a region of low differential heat ( <90 kJ mol-') follows, corresponding to H-bonded water. Our initial values of qdiff of water adsorption on a-Al,O, are in fairly good agreement with those measured at 150°C by Della Gatta (251 kJ mol-')' and are higher than those measured at 23 "C by Yung-Fang Yu Yao (>125 kJ mol-'), but these measurements were carried out on single crys- tals. However, our measurements of differential heats are 300 I I 1.5 01 I I I 15 75 135 195 255 0 1 2 t ime/min n,/pmol m-2 Fig.4 Heat emission peaks as a function of increasing water Fig. 6 Differential heats of adsorption (run 1). In the inset, integral vapour equilibrium pressure at 400°C: (a) 0.2, (b)0.4, (c) 0.7 and (d) heats of adsorption (run 1). (0) 250, (A) 300, (0)350 and 200, (O), 2.0 Torr (V)400 "C. 366 higher than those measured by Venable (84kJ rnol-')l6 and those measured by Hendriksen (ca. 109 kJ mo1-1),'4 but they both used immersion calorimetry at 25 "C. We can try to explain the reason why the differential heat curves show a rapid decrease with increasing coverage and temperature. There are two possible causes for this drastic decrease in differential heats: the extreme heterogeneity of the surface and repulsive interactions between adsorbate mol- ecules.As the temperature increases, these two factors have an increasing effect on the differential heats.I6 In Fig. 7, we show a theoretical study of the differential heats for derivation of the integral heats of adsorption. We carried out the following mathematical operations on an IBM Computer AT:'7 (a) Interpolation of the curves of the integral heats by means of a polynomial; (b)calculations of the polynomial coefficients by means of a mean-square method for a best fit to the experimental data; (c) polynomial derivation; (d)determination of the curve of differential heats. The polynomial type that best approaches the experimental differential heat curves is a polynomial of the fourth degree.We can verify that, at all experimental temperatures, the curves show an exponential decrease, apart from the varia- tion of the convexity of 350°C and at 400"C,towards the surface saturation. In Fig. 7, we can see that the increasing temperature affects the trend of the qdiff(and hence the trend in the site strength distribution) with coverage. In particular, it seems that the trend at 200°C and at 250°C is very similar; whilst at 300"C, the slope of the curve clearly changes and the trends at 350°C and 400°C are again different but very similar to each other and fall very rapidly towards the value of the heat of water vaporization. We can attempt to explain the trend of the differential heats at 350°C and at 400°C.Water molecules adsorbed on stronger sites must be immobile especially at lower tem- perature, while those on the less active sites of hydrogen- bonded hydroxy groups will have a good possibility of moving. As the temperature increases, the alumina surface becomes more porous to water vapour and so more water molecules penetrate into the bulk, producing a higher inter- action. Reversible Phases Differential heats of adsorption for run 2 are shown in Fig. 8 for all five temperatures. Such reversible phases seem to 300 -200-E 3 1-p" 100 0 0.5 1.o 1.5 2.0 na/pmol m-2 Fig. 7 Differential heats of adsorption obtained from derivation of a fourth-degree polynomial. (a) 200, (b) 250, (c) 300, (6)350 and (e) 400 "C.J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 300 E na/pmol m-2 I I 01 II I I I I I 0 0.5 1.o 1.5 2.0 n./pmol m-Fig. 8 Differential heats of adsorption (run 2). Inset: integral heats of adsorption (run 2). (m) 200, (0)250, (A)300, (+) 350 and (V) 400 "C. consist of molecular water chemisorbed on the dehydroxy- lated surface. Comparison of the differential heats of reversible adsorp- tion at the different temperatures, gives the following infor- mation on molecular water chemisorbed on the dehydroxylated surface: (a) At very low coverage, the same initial average value is always obtained (190 kJ mol-'); (b) the adsorption temperature does not seem to have much effect on the curve shape (exponential trend).We note that the number of available sites for the first adsorptions (run 1) are not yet useful for the second adsorp- tion (run 2). Furthermore the differential heats of run 2, at all five temperatures, fall in a very narrow range of values. As the temperature increases, the surface sites of a-Al,O, are previously saturated with a lower amount of adsorbed water and the differential heats decrease steeply from 190 kJ mol-' to a value near to that of the heat of water vaporization (43.8 kJ mol- '). Conclusions Surface calorimetry is a powerful technique for studying the chemistry of a surface and its interactions with different mol- ecules. We have shown that the active sites of or-Al,O, react again with water vapour at 400°C and the distribution of different strength sites changes on increasing the temperature from 200 to 400°C.Moreover, all our experimental data confirm the complex- ity of the adsorption process of water vapour on alumina at high temperature. References 1 R. B. Gammage, E. L. Fuller Jr. and H. F. Holmes, J. Phys. Chem., 1970,74,4276. 2 E. Borello, G. Della Gatta, B. Fubini, C. Morterra and G. Ven- turello, J. Card., 1974,35, 1. 3 J. B. Peri, J. Phys. Chem., 1965,69, 21 1. 4 E. McCafferty and A. C. Zettlemoyer, Discuss. Faraday Soc., 1972,52, 239. 5 G. Della Gatta, B. Fubini and L. Stradella, J. Chem. SOC., Faraday Trans. 2, 1977,73, 1040. 6 M. Nagao, K. Yunoki, H. Muraishi and T. Morimoto, J. Phys. Chem., 1978,82, 1032. 7 T. Morimoto, M. Kiriki, S. Kittaka, T. Kadoka and M. Nagao, J. Phys. Chem., 1979,89, 2768. 8 B. Fubini, G. Della Gatta and C.Venturello, J. Colloid Interface Sci., 1978,64,470. J. CHEM. SOC. FARADAY TRANS., 1993, VOL. 90 367 9 10 11 12 P. F. Rossi, G. Milana and A. Vettor, AFCAT. Calorimdrie et Analyse Thermique, Clermont-Ferrand, 1990, vol. XX-XXI, p. 137. C. Morterra, G. Magnacca and N. Del Favero, Langmuir, 1993, 9, 642. J. Barto, J. L. Durham, V. F. Baston and W. H. Wade, J. Colloid Interface Sci., 1966, 22,49 1. T. Morimoto, N. Katayama, H. Naono and M. Nagao, Bull. 14 15 16 17 B. A. Hendriksen, D. R. Pearce and R. Rudham, J. Card., 1972, 24,82. Yung-Fang Yu Yao, J. Phys. Chem., 1965,69,3930. R. L. Venable, W. H. Wade and N. Hackerman, J. Phys. Chem., 1965,69, 317. G. Della Gatta, B. Fubini and C. Antonione, J. Chim. Phys., 1975,12, 66. 13 Chem. Soc. Jpn., 1969,42, 1490. J. G. Dawber, L. B. Guest and R. B. Lambourne, Thermochim. Acta, 1972,4,471. Paper 3/041711; Received 16th JuIy, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000363

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(2), 369-376 Frequency Response Analysis for Multicomponent Diffusion in Adsorbents L. M. Sun, G. M. Zhong, P. G. Gray and F. Meunier LIMSI-CNRS, B.P.133 91403, Orsay France Frequency response behaviour of diffusion of multiple sorbates in adsorbents is theoretically investigated. A mathematical model describing diffusion of n components with both equilibrium and diffusional interference is developed and analytical solutions for frequency response and moments are derived. It is shown that, for the model system CO, and C,H, in 4A zeolite, the fastest diffusing component can be strongly affected by other slower diffusing components and its partial-pressure frequency response curves exhibit a roll-up phenomenon induced exclusively by the diffusional interference.The frequency response curves of the total pressure gener- ally have multi-modal forms as a result of multicomponent diffusion ; the peak positions and heights within these responses, which are closely related to the diffusivities, are found to be sensitive to the extent of equilibrium and diffusional interference. Comparison of these resonance frequencies obtained from purecomponent and multicomponent experiments can therefore give a good indication of the interference between the components. The effects of film resistance and heat dissipation, which may be important for fast diffusing systems, are also discussed. The modelling of diffusion of multiple adsorbates in a micro- porous adsorbent in general requires knowledge of both the main-term and cross-term diffusivities, as the effect of the cross-term diffusivities may be important for strongly inter- acting systems.As a consequence, the number of the diffusion coefficients to be determined may increase rapidly with the number of components; n(n + 1)/2 diffusion coefficients needed for a system of n components, by application of the Onsager reciprocal relation. Simultaneous experimental mea- surement of all these coefficients is very difficult, even for a binary system. This has necessitated the development of theo- retical models capable of predicting multicomponent diffusi- vities from pure component diffusivities. The advantage of predicting multicomponent data from single component data by using theoretical predictive models has long been proved in the correlation of multicomponent adsorption equilibria.For certain systems, the diffusional interference of different components mainly arises from the coupling through multi- component equilibria. The cross-terms of the phenomenologi- cal transfer coefficients, based on the irreversible thermodynamics formulations, can therefore be neglected and the diffusion flux of a given species depends only on its own chemical potential. '-3 The resulting Fickian diffusivity matrix has non-zero cross-terms which are concentration dependent. This model can be referred to as the ideal multicomponent diffusion model. Krishna has developed a model based on a generalized Maxwell-Stefan formulation in which the adsorb- ent is treated as an additional c~mponent.~ The transfer coef- ficients of this model were composed of the pure component coefficients and those for the counter-exchange between the adsorbates which were related to the pure component diffu- sion Coefficients by using empirical correlations.Fully predic- tive models have recently been developed by Wei and co-workers using a stochastic Markovian formalism, and by Chen and Yang using a kinetic In these models, the cross-diffusivities are explicitly related to the diffusivities of pure components, and to the occupancy, through a number of parameters characterizing interactions between adsorption sites and diffusing molecules. The model of Chen and Yang has a simple form, which reduces to the ideal multicomponent diffusion model with an extended Langmuir isotherm when all the interaction parameters are equal to zero.Their model was able to predict experimental data accurately for the diffusion of C0, and C,H6 in 4A zeolite and of benzene and toluene in ZSM-5 zeolite. In order to check the applicability of predictive multi- component diffusion models, experimental multicomponent diffusion data for different representative systems are needed. However, due to the experimental difficulties in measuring the uptake of each adsorbate, very few experimental investi- gations have been reported so far in the literature and avail- able experimental techniques are quite limited : the microscopic NMR technique,' * the differential adsorption bed (DAB) technique,2*' and the constant-volume technique.' 3-' Recently, the frequency response technique has been found to be very useful for the measurement of diffusion rates of single species in adsorbents. 16v1 'Detailed theoretical studies have revealed a high sensitivity of the technique to the nature of the governing transport equations.'8-21 For example, a surface barrier effect causes the so-called in-phase and out-of- phase functions to intersect, while intrusion of heat effects can lead to bimodal frequency response curves.The intersection and bimodal form of the frequency response curves constitute an unusual pattern of behaviour which cannot be described by a pure diffusion model.It therefore provides, in principle, a good way of distinguishing between intracrystalline diffu- sional resistance, surface resistance to mass transfer and the effect of heat dissipation. The high sensitivity of the frequency response technique may also be useful for studying multi- component diffusion in adsorbents. The first frequency response experiments for diffusion of binary systems have recently been reported in the literat~re.~~-,~ However, firm conclusions concerning the applicability of the technique cannot be drawn from these papers owing to the use of overly simplified theoretical models. In this paper a detailed theoretical analysis of frequency response for the diffusion of multiple adsorbates in mono- dispersed adsorbent particles (e.g.zeolite crystals) will be pre-sented. The model to be developed will include both equilibrium and diffusional interference and is, therefore, general enough to allow a reliable investigation of the sensi- tivity of the frequency response to main-term and cross-term diffusivities. The theoretical study, using the theory of Chen and Yang for predicting multicomponent diffusivities, will enable us to determine whether measurement of the total pressure alone is sufficient to yield a reliable estimation of both main-term and cross-term diffusivities. Moreover, effects of heat dissipation and surface barrier are also included in the model and discussed. Mathematical Model Consider an adsorption chamber in which are placed a number of well separated monodispersed adsorbent particles occupying a volume V,.The adsorbent particles are assumed to be of uniform size; a significantly non-uniform particle size distribution would lead to distortion or smoothing of the fre- quency response curves, thus resolution of characteristic peaks in the frequency spectrum would be very difficult. Hence, in any experimental study, it would be necessary to work with a rather uniform particle size distribution; an assumption which is made in this mathematical model. The chamber contains an ideal gaseous mixture of n components in a volume V which can be varied in a controlled manner. The volume variation causes the total pressure and partial pressures of the mixture to change, which in turn causes the gases to diffuse into or out of the adsorbent particles.The sorption of the adsorbates produces a change in the adsorb- ent temperature and a heat exchange occurs between the uniform-temperature adsorbent and the constant-temperature surroundings. The system can be considered as linear since volume changes are, in general, quite small (< 5%). The mass balance in the particles is given by the following multicomponent diffusion equation : with the reduced time t = t/R: and a = 0, 1, 2 for slab, cylin- drical and spherical particle geometry, respectively. q is a vector of dimension n denoting the incremental quantity adsorbed by the particles. [D] is a square matrix of diffusion coefficients in which the off-diagonal terms are generally non- zero.(See glossary for definition of other symbols). The corresponding initial and boundary conditions are : d$, 0)= 0 (2) (3) where the parameter k, is introduced to account for a surface barrier effect, which can be due to surface heterogeneity and/or film resistance. q* is related to the gas pressures p and temperature T through the linearized multicomponent adsorption equilibrium : q*=[Klp-kTT (5) The heat balance on the particles is given by the following scalar equation : where t, = R,CJ[(o + 1)h) is the time constant for heat exchange between the adsorbent and surroundings. The pressures of the adsorbates in the gas phase in the chamber are assumed to be uniform throughout the chamber J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 and are determined by the following mass conservation con- dition :20 P + diag(B,M = ; for a constant-volume condition for a volume modulation (7) i"iu exp(iot)p, ; where pi is the initial pressure in the gas phase, u and o are the relative amplitude and angular frequency of a sinusoidal volume modulation, respectively. The parameter B, is defined as KRk q/K.The zero-order moments for the above system under the constant-volume condition can be obtained from the Laplace solutions using Van der Laan's theorem.25 The solution pro- cedure is similar to that used for deriving the frequency response solution which will be presented later. The moments of the temperature T and quantity adsorbed q are: thM(T)= J':.dt = --(AH -CS R: [D]-'(a + 1)(a + 3) where [A] = [I] + [K]diag(/3,) represents the ratio of the masses contained in the adsorbent and gas phase, respec- tively.t,, = R,/[(a + l)k,] is the time constant for the surface barrier effect. The tensor product of k, and AH leads to a square matrix whose elements (i, j) are kTiAHj.q, is the amount adsorbed at the equilibrium state: 4, = [KIWI + diag(B,)[KlI -'Pi (10) The temperature moment is linearly proportional to the heat exchange time constant and the moments for the adsorbed masses are given by a linear addition of the resist- ances caused by surface barrier, diffusion and heat exchange, respectively. For one-component systems (n= l),the moment solutions [(eqn.(8) and (9)] reduce to those given in a pre- vious study.lg From a practical point of view, the zero-order moments are useful for determining both the heat transfer coefficient, h, and the total resistance to mass transfer. Analytical Frequency Response Solution The frequency response solution of the above system can be analytically obtained by using the matrix analogue.26-28 The key point is to diagonalize the matrix [D]in terms of the eigenvalues and eigenvectors : [XI -[D] [XI = diag(L,) (11) where 1, are the eigenvalues of [DJ and [XI is formed from columns of the corresponding eigenvectors. For a binary system, the eigenvalues are : and the eigenvector matrix is : 1 CXl = i D2 1 11 -D22 4,) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 With the diagonalization of the diffusivity matrix, the n coupled diffusion equations [eqn. (l)] can be separated into n individual equations which can be easily solved as in the case of pure component systems. The periodic steady solution of eqn. (1)-(4) is: q = [Xldiag [B]([K]p -kr T)(;&::;IThe function f(+) is defined by: cosh(b,+) ; for 0 = 0 ; for 0 = 1 fk($) = I (k = 1, 2, ..., n) -+ sinh(b, +) ; for 0 = 2 where b, = (1 + i),/wR,2/(2Ak). The matrix [B] is derived from the particle surface bound- ary condition [eqn. (4)] : [BI -= [XI + diag(wtSk)[X]diag(6,,+ idck) with a+l6,k -id,, = -bk tan h(b,) ; for 0 = 0 'l(bk)/zO(bk) ; for B = 1 (k = 1, 2, ..., n) (bk coth(b,) -l)/bk; for B = 2 where d,, and 6, are in-phase and out-of-phase functions for the intrinsic diffusion mechanism only.The volumetic average for qis: 4 = CAI"Klp -k, T) (13) with [A] = [X]diag(G,, -iS,,)[B]. From the heat balance equation [eqn. (6)],the following equation is obtained: where aT= iwtd(1 + iwt,). This yields : 4 = CClCAlCKIP = CHIP (15) where [C]-' = [I] -(cl~/C,)[A][kr x hw. Eqn. (15) is introduced into eqn. (7) to obtain solutions for the gas pressures p: p = u exp(iwt)[Elp, (16) where [El-' = [I] + diag(B,)[q. This allows one to deter- mine the amplitudes and phase lags of the partial and total pressures ak,cp, and A,0,respectively. The in-phase and out-of-phase functions for the partial and total pressures are defined as: Pek ' Bk 0,6in.k -iaout,k= -eXp(-icpk) -1 = -(k = 1, 2, .. . , n)a& Yk (17) S!tot)In -i'(tot) out = APv exp(-ia) -1 = c;=1 Pk ek A C;=l Yk (18) where N 37 1 When the interference between the different components is negligible, both for the equilibrium and kinetics (Kkj = Dkj = 0, k #j), the following simple frequency response solutions are obtained: where the heat effect has been neglected and These solutions are identical to those obtained for one-component systems. ' The above frequency response functions can be normalized with respect to their asymptotes of the in-phase functions at zero frequency, which is done by replacing [HI and [El -by[K] and [I] + diag(B,)[K), respectively in eqn. (17) and (18).In the case of negligible equilibrium interference, i.e. [K] x diag(K,,), the asymptotes are given by: din,klo=O= jkKkk (20) This simplified solution has been given by Yasuda for a binary system.,, Results and Discussion The adsorption of CO, and C,H, in 4A zeolite was used as a reference system for theoretically investigating multi-component frequency response behaviour. The equilibrium and diffusion parameters of this system have been given by Chen and Yang7 and are summarized in Table 1. The main characteristic of the chosen system is that the diffusion of CO, in 4A zeolite is much faster than that of C,H, [D,(CO,)/Do(C,H,) x 501. Values of the other parameters are: V, = m3, V,= 3 x m3, 0 = 2 (spherical crystals), R, = 1.7 pm, u = 0.02, C, = 10, J m-3 K, t, = 0.1 s and t, = 0 (no surface barrier effect).With these values, the heat effect proved to be negligible owing to the very low rates of diffusion of CO, and C,H, in 4A zeolite. The equilibria of the binary mixture C0,-C,H, were cal- culated using the following LRC isotherm: where the constants b and n were assumed to be the same as for the pure components. The prediction of the binary diffusi- vities from pure-component diffusivities was made using Yang's m~del~.~ where Do, is the zero coverage Fickian diffusivity of com- ponent k, and wkjrepresents the ratio of the sticking prob- ability of molecule k on adsorbed molecule j to that on a Table 1 Equilibrium and diffusion parameters of CO, and C,H, in 4A zeolite at 25 "C taken from Chen and Yang.' no.sorbate qJkgm-' bfPa-' n -AH/J kg-' D,/m's-' UJ~ 1 CO, 290 3.7 x lo-' 0.530 0.75 x lo6 4.2 x 10 Is 0.0936 2 C,H, 121 7.0 x 0.982 0.93 x lo6 8.2 x lo-" 0.0117 vacant site, which can be obtained from pure-component data (AHk and okk).When all o are equal to zero, the above model leads to the same expressions obtained from irrevers- ible thermodynamics for ideal multicomponent diffusion with an extended Langmuir i~otherm."~ When all o are equal to unity, the model predicts independent diffusion behaviour with constant main-term diffusivities and zero off-diagonal diffusivities(Dkk= Do,and Dkj= 0). All the simulations were made with P, = 15 kPa and x(C0,) = 0.33, which gave the following values for [K] and CDl : 0.0067 -0.0023 w1=( -0.0010 0.0012 1.4 10-l~ 1.1 x 10-l~) ['I = (3.7 10-17 1.2 x 10-14 Effect of Equilibrium and Diffusional Interference on the Partial-pressure Response The frequency response of the binary mixture, with both dif- fusional and equilibrium interference, was calculated using the reference equilibrium and diffusion values.The normal- ized in-phase and out-of-phase functions for both partial and total pressures are shown in Fig. 1. It can be seen that the frequency response curves of the slower diffusing component (C,H4)behave monotonically (as in a single component case) and are little affected by the presence of CO,. On the other hand, the faster diffusing component (CO,) is strongly influ- enced by C,H4 and has a negatively valued out-of-phase function near the resonance frequency of the slower com- ponent, while its in-phase function becomes greater than unity.This much stronger influence of the slower component on the faster component had previously been pointed out by Chen and Yang.' Effects of the diffusional interference can be better under- stood by comparing Fig. 1 with Fig. 2, wherein are shown frequency response curves obtained by cancelling the cross- term diffusivities (ill,= D,,= 0). In this case, the behaviour of the frequency response of CO, becomes 'normal', in the sense that the in-phase function does not exceed unity and the out-of-phase function is always positive. Owing to the equilibrium interference, however, the frequency response of CO, exhibits an additional maximum of smaller magnitude at lower frequency, which disappears completely when the cross-terms of the equilibrium are also set to zero (K12 = K,, = 0, Fig.3). The comparison of these figures suggests that the 'roll-up' (in-phase greater than unity and out-of- ...1.2 -f in-phase \\ I\ total0.6 0.4 0.2 0.0 --.-.--I out-of-p hase -0.21 '"'~1'1 111111111 ' 1I''I111 1 IIIII 1 TIIlITT 10-~ lo4 10-3 lo9 lo-' 1 freq uency/H z Fig. 1 Normalized in-phase and out-of-phase functions of the partial and total pressures for the reference case. Short dashed line: CO, ,long dashed line: C,H, and solid line: total pressure. J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 co2 f req uency/Hz Fig. 2 Same as in Fig. 1 except that the cross-term diffusivities (D,, and D,,)have been set to zero frequency/Hz Fig. 3 Frequency response of two independent components with zero equilibrium and diffusion cross terms (KI2= K,, = D12= D,,= 0), i.e. the independent single component case. Reference values for the other parameters. phase less than zero) of the CO, partial-pressure frequency response curves is caused by diffusional interference. This can be further confirmed by Fig. 4, wherein the frequency response is obtained by swapping the zero-coverage diffusi- vities of CO, and C,H, . The roll-up now occurs for the light component C,H, which has a larger diffusivity, but the bimodal pattern of the total-pressure frequency response -.1.2-1 , 10" lo4 10" lo-* lo-' freq u ency/H z Fig.4 Frequency response obtained by exchanging the zero-coverage diffusivities of CO, and C,H,. Reference values for the other parameters. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 curves has largely disappeared owing to the smaller adsorp- tion capacity of the fast-diffusing component. The change in behaviour of the CO, frequency response curves in Fig. 1-3 can be explained by investigating the two components, phase shift (qco2)and amplitude (aco2),of the in-phase (din) and out-of-phase (dout)functions, where (25) The phase shift and amplitude behaviour of CO, corre-sponding to Fig. 1-3 is shown in Fig. 5-6. Curves (a)in Fig. 5 and 6 show the behaviour of the amplitude (aco2)and phase shift (qco,)representative of a pure component with neither equilibrium nor diffusional interference (Kl2 = K,, = D,, = D,,= 0 corresponding to Fig.3). When equilibrium inter- ference is introduced uia the non-zero cross-terms, K,, and KZ1(curves (b)in Fig. 5 and 6, corresponding to Fig. 2), one observes that the amplitude increases in two stages with increasing frequency, while there are now two peaks in the phase shift spectrum. The additional phenomena in the amplitude and phase shift spectra at low frequency are induced by the behaviour of C,H, in the region of its own resonant frequency (ca. lop4Hz). As the frequency increases and enters the resonant frequency region of C2H6, the resist- I -____-----__---.____----250-200 $ 150 lo4 io4 lo-' 1 10 frequency/ Hz Fig. 5 Gas-phase CO, partial pressure amplitude us.frequency. Curve (a):no equilibrium or diffusional interference (K,,= KZ1= D1,= D,, = 0),corresponding to Fig. 3. Curve (b):no diffusional interference (D,,= D,,= 0, K,,, K,,# 0),corresponding to Fig. 2. Curve (c): both diffusional and equilibrium interference, reference case (K,,, K,, ,D,,,D,, # 0),corresponding to Fig. 1. 0.20 0.1 5 v) m.-x 0.10 s-'c r 0.05 v) c 0 0.00 -0.05 lo4 lo-' I 10 frequency/ Hz Fig. 6 Phase shift as a function of frequency. Parameters for curves (a),(b)and (c)as in Fig. 5. ance to diffusion inhibits the movement of C& into and out of the adsorbent, resulting in a concentration profile; conse- quently the average adsorbed-phase amplitude decreases and the gas-phase amplitude increases.Expanding the multi- component equilibrium [eqn. (5)], for constant temperature and no surface barrier, the particle surface concentration of CO, is given by: qco2I*= 1 = Kl lPCO2 + K12PC2H6 (26) As K,, is negative, the increasing C,H6 gas-phase amplitude causes a decrease in the (amplitude of) CO, surface concen- tration oia eqn. (26),and consequently a quasi-instantaneous decrease in the amplitude of the CO, concentration within the adsorbent and therefore an increase in the gas-phase CO, amplitude [the first stage increase of curve (b)in Fig. 51. This CO, response is entirely induced by the equilibrium inter- ference of C,H6 [oia eqn.(26)],giving a positive phase shift peak at the same frequency as that for C,H6 (ca. Hz, Fig. 6).The second stage increase of the CO, amplitude, and the second CO, phase shift peak (at CQ. lo-, Hz) are caused by the retarded transfer of CO, resulting from the adsorbent diffusional resistance (at a timescale characterized by D,l). Curves (c) in Fig. 5 and 6 (corresponding to Fig. 1) show the CO, behaviour with both equilibrium and diffusional interference. The CO, in-phase function greater than unity in Fig. 1 is mainly due to the decrease in amplitude (in the region 10-5-10-2 Hz) in Fig. 5, while the negative CO, out-of-phase function (region 10-5-10-3 Hz in Fig. 1) is mainly a result of the negative phase shift in Fig.6. This unusual behaviour is due to a combination of C,H6 diffu- sional and equilibrium interference and can be explained by inspecting the concentration profiles within the adsorbent particle. The particle concentration profiles of CO, and C2H6, for the three cases of equilibrium and diffusional inter- ference (presented in Fig. 1-3), were generated by numerical solution (finite difference) of model eqn. (1)-(4). These profiles (in Fig. 7) are given for a frequency of lop3Hz (diffusional resistance is very large for CzH6, but negligible for CO,), at 1/8 of a period (i.e. volume decreasing, total pressure increas- ing, adsorption in progress). It is evident from the similarity of all three ethane curves in Fig.7 that the C2H6 profiles are little affected by either equi- librium or diffusional interference with CO,; the C2H6 behaviour is dominated by diffusional resistance within the particle at a timescale characterized by D,,.Unlike the pro- files of C,H6, Fig. 7 shows that the CO, profiles are strongly influenced by equilbrium and diffusional interference. The CO, profile with no interference (corresponding to Fig. 3, i.e. 0.25 I P" 0.154 -0.05 !0.0 I0.2 I0.4 10.6 I0.8 1.o radial coordinate (+) Fig. 7 Adsorbed phase concentration profiles after 1/8 period (volume decreasing) at Hz. Parameters for curves (a), (b)and (c) as in Fig. 5. independent single component behaviour) is given by curve (a),which has a very weak gradient at the surface as a result of negligible diffusional resistance (at 10-Hz).As explained previously, the addition of equilibrium interference via the cross-terms, K,, and K,, (corresponding to Fig. 2), results in a decrease in CO, concentration within the particle due to the surface CO, concentration being reduced by C,H6 inter- ference. This is shown by curve (b) in Fig. 7 where, again, the profile is relatively flat due to minimal diffusional resistance. Profile (c) in Fig. 7 results from both equilibrium and diffu- sional interference (Kl,, K21, D,,, D,, non-zero), corre-sponding to the frequency response curves presented in Fig. I. Expanding eqn. (1) for CO, gives (27) showing that the C2H6 gradient within the particle can also contribute to the accumulation of CO, (as D,,is positive and the same order of magnitude as Ill,).Hence, the strong posi- tive C,H6 gradient near the surface induces an additional flux of CO, into the particle which increase the mass of CO, adsorbed in the particle, raising profile (c) above profile (b); the adsorbed mass amplitude is therefore greater, and the gas-phase CO, amplitude is smaller, than for case (b) (see Fig. 5). The additional CO, flux induced by the C,H6 gradient is, however, not the only phenomenon affecting the shape of the CO, profile. As in case (b),C,H6 equilibrium interference (via the K,, term), and also reduced CO, in the gas phase, reduces the CO, surface concentration much below that inside the particle (independently induced by the C,H6 gradient).This results in a steep negative CO, gradient near the adsorbent surface [curve (c), Fig. 71 and, consequently, desorption from the particle in the region of the surface, Although it appears paradoxical that both co-diffusion and counter-diffusion can occur simultaneously, this is in fact per- missible as the two CO, fluxes, due to the CO, and C,H6 gradients [eqn. (27)], are completely independent of each other; the result of these two processes being the negative CO, phase shift (see Fig. 6) in the region of lo-’ to Hz. Conversely, if the sign of the diffusion cross-terms (D,,, D2J is made negative, the CO, out-of-phase peak (and hence the phase shift) at lop4Hz becomes positive and the CO, in-phase is always less than unity (Fig.8), and the ‘roll-up’ 0.0 , 10-5 lo4 lo3 lo-’ 1 frequency/Hz Fig. 8 Normalized in-phase and out-of-phase functions of the CO, (short dashed line) and C,H, (long dashed line) partial pressures and the total pressure (solid line). Parameters same as reference case except with D,, and D,, negative. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 behaviour of Fig. 1 completely disappears. The very large C,H, gradient at the particle surface (starting to appear before ca. Hz), combined with the negative value of largely blocks the net movement of CO, into and out of the adsorbent [via eqn. (27)]. Thus the CO, out-of-phase peak at ca. Hz is more pronounced than that at the CO, reson- ant frequency (ca.Hz); the opposite behaviour to that observed for diffusional interference with positive cross-terms (Fig. 1). Comparing Fig. 1, 2, 3 and 8, one can clearly see that the roll-up of the CO, partial pressure frequency response curves is uniquely a consequence of diffusional interference with positive cross-terms ; no other type of interference can cause this behaviour. Sensitivity of the Total-pressure Response The out-of-phase curves of the total pressure shown in Fig. I, 2, 3 and 8 exhibit a bimodal form due to the difference in diffusion rates of the two components: the peak at the lower frequency is induced by diffusion of C,H6 and the peak at the higher frequency by diffusion of CO,. While the frequency response curves of the partial pressures have qualitatively different forms due to equilibrium and diffusional interference between the components (as discussed previously), it is evident from Fig.1, 2, 3 and 8 that the corresponding responses of the total pressure have similar forms and there- fore are only quantitatively different. This is summarized in Fig. 9, wherein the total in-phase and out-of-phase curves from Fig. 1,2, 3 and 8 are assembled. It is clear from Fig. 1, 2, 3 and 8 that one could easily discriminate between the different types of interference if the CO, and C,H6 partial-pressure responses were experimen- tally measured. However, current frequency response tech- niques can only give the total-pressure response, so it is important to know whether this is sufficient to identify accu- rately the type of interference involved.The difference between the curves in Fig. 9 indicates that, in principle, it should be possible to identify the type of interference based on the experimental total-pressure response alone; noting that, although curves (a) and (b)(no interference and equi- librium interference, respectively) are virtually identical when normalized, they are quite different in the unnormalized form. In practice, however, one would require very ‘clean’ data (i.e. with very little experimental scatter); which has not always been the case, i.e. some frequency response data appearing in the literature has quite a lot of scatter of the experimental 1.o 0.8 0.6 0.4 0.2 0.010-5 10-4 10-3 10-2 lo-’ frequency/Hz Fig.9 Comparison of total-pressure responses (from Fig. 1, 2, 3 and 8) for different types of interference. Curves (a),(b) and (c) same as in Fig. 5. Curve (4:both diffusional and equilibrium interference, but with negative cross-term diffusion coefficients (K ,,, K,, # 0; D,,, D,,< 0),corresponding to Fig. 8. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 points. Without 'clean' data, one could not discriminate between the different theoretical curves on a statistically sig- nificant basis. The separation between the total-pressure out-of-phase peaks is dependent not only on the type of interference, but also on the magnitude of the diffusivities of each component and the thermodynamic properties of the system. Thus, for an experimental study, one should choose a binary system in which the diffusivities of each component differ by more than an order of magnitude, otherwise, as shown in Fig.10, the out-of-phase peaks will not be sufficiently distinct. Further- more, the influence of total pressure and composition of the gas mixture should not be ignored as, particularly when the isotherms for each component are quite different (as in the case of CO,/C,H, on 4A zeolite, see ref. 7), an inappropriate choice of these experimental conditions may render the out-of-phase peaks inseparable or cause one peak to disap- pear. Effect of Surface Resistance For multicomponent diffusion in zeolite crystals, the surface barrier effect can be either of an intrinsic nature and/or caused by film resistance in the gas phase.The former can result, for example, from obstructions of the pore entrances or the inte- rior windows by hydrothermal pretreatment3 and can be quantified by pure-component experiments. The latter is due to the concentration gradient in the gas phase created by the difference in diffusion rates of the components inside the crys- tals and does not exist in a pure-component experiment. As can be seen from Fig. 11 (reference case with the addition of a surface barrier of time constant is), the existence of any surface resistance may cause the in-phase and out-of-phase functions of the total pressure to intersect, as in the case of pure systems.In principle, the presence of a gas-phase con- centration gradient induced surface resistance may be detected by a quasi-single-component experiment in which the adsorbable component is mixed with a second, non-adsorbable, component (e.g., helium), hence artificially cre- ating a surface barrier. Intrusion of film resistance in multicomponent frequency response experiments has been observed by Shen and Rees when studying diffusion of p-xylene and benzene in silicalite-I:24 the experimental data of the binary mixture revealed the presence of a surface resist- ance, while the frequency response curves obtained from pure-component experiments excluded the existence of any structural surface barrier. 375 I in-phase4 0.8 % 1 out-of-Dhase I 0-5 lo4 10" lo-* 10-1 frequency/Hz Fig. 11 Effect of surface resistance on the frequency response of the total pressure. Solid line: t, = 0 (no surface barrier effect); long dashed line: t, = 6 s and short dashed line: t, = 60 s.Reference values for the other parameters. Effect of Heat Dissipation For diffusion of the mixture CO, and C2H, in 4A zeolite, the rate of diffusion is so slow that the effect of heat dissipation is negligible under usual experimental conditions. However, for fast diffusing systems, the heat effect may become important and significantly alter the frequency response behaviour. This is illustrated by the curves in Fig. 12 where, by artificially increasing the diffusivities more than lOO-fold, a thermal resistance is induced ; the frequency response curves were obtained using Do, = 5 x lo-'' m2 s-', Do,2= m2 s-', R, = 10 pm, t, = 7 s for the non-isothermal case, P, = 10 kPa and x(C0,) = 0.5 (reference values for the other parameters).It can be seen that the out-of-phase curves, which are bimodal in the isothermal case, exhibit an addi- tional peak in the non-isothermal case, as a result of a rate- limiting heat exchange. The three peaks remain, however, sufficiently distinct that the corresponding mechanisms could be discriminated with good experimental data. The fact that the resonance frequencies of each of the different transport mechanisms are little affected by the other mechanisms con- stitutes one of the most attractive advantages of the fre- quency response technique.1.o 0.8 -\'vphasea-, 10 f req uency/H z frequency/Hz Fig. 10 Out-of-phase curves of the total pressure for three ratios of Fig. 12 Effect of heat dissipation on the frequency response of the diffusion rates: short dashed line (Do, = 4.2 x lo-"), long dashed total pressure. Solid line: isothermal case (th= 0 s) and dashed line: line (Do, = 4.2 x 10-16) and solid line (Do,,= 4.2 x 10-15) non-isothermal case (th= 7 s). Curves obtained with Do, = 5 (reference values for the other Parameters). The ratios of Do, to Do,, x mz s-l, Do,z = m2 s-l, R, = 10 pm,P, = 10 kPa are given by the numbers on the curves. and x(C0,) + = 0.5 and reference values for the other parameters. J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 Conclusions Owing to its high sensitivity to the nature of rate-limiting transport processes, the frequency response technique offers the possibility of studying multicomponent adsorption systems. The present work has theoretically investigated the sensitivity of the partial-and total-pressure frequency responses to the nature of the interference affecting the multi- component diffusion. It has been shown, in the case of a binary system, that partial-pressure frequency response curves of the fast diffusing component can be qualitatively different, depending on the presence or absence of equi-librium and diffusional interference, while the resulting total- pressure frequency response curves are only quantitatively different. Nevertheless, with good experimental data, it should be possible to identify the type of interference from the total-pressure frequency response, particularly if pure and multicomponent experimental results are compared. Because of the complexity of multicomponent diffusion, however, it is preferable to make experiments over a range of conditions (compositions, pressures and temperatures), guided by a pre- dictive diffusivity model, in order to have a precise and unique determination of the multicomponent diffusivities. Furthermore, it has been shown that, for the case of equi- librium plus diffusional interference with positive cross-diffusivities, both counter-diffusion and co-diffusion of the faster diffusing component can coexist.Frequency response experiments can be disturbed by the presence of surface barrier and heat effects. The former can result from surface heterogeneity and/or film resistance in the gas phase (not present in pure-component experiments, but may arise in multicomponent experiments) and lead to inter- sections of the in-phase and out-of-phase functions. The effect of the intrinsic surface barrier can, in principle, be determined from pure-component experiments while the effect of the film resistance has to be determined from multicomponent experi- ments. In principle, the latter effect could be estimated using a binary mixture composed of an adsorbate and an inert gas which does not diffuse in adsorbents. For fast diffusing systems, effect of heat dissipation can be important and lead to an additional peak for the frequency response of the total pressure.In order to avoid attributing this thermal peak to a diffusion process, it is highly necessary to quantify the heat effect, which once again can be estimated from pure-component experiments. Glossary cs volumetric heat capacity of crystals/J K-' Do diffusion coefficient at zero coverage/m2 s-' D Fickian diffusion coeficient/m2 s-' h heat transfer coefficient/W m-* K-AH heat of adsorption/J kg- ' mass transfer coefficient for surface resistance/m s-' kS derivative of isotherm with respect to the temperature/ kT kg m-3 K-' K adsorption isotherm equilibrium constant/kg m-Pa-' n number of components P incremental pressure of adsorbates/Pa Pi initial gas pressures in the chamber/Pa absolute partial pressures at equilibrium/Pa Pe P incremental total pressure/Pa Pe absolute total pressures at equilibrium/Pa incremental amount adsorbed by crystals/kg m- 4s saturated amount adsorbed by crystals/kg m-3 r spatial coordinate/m R gas constant/J kg- ' K-' RC radius of crystals/m time/s time constant of heat exchange/s time constant of surface resistance to mass transfers/s incremental temperature/K amplitude of volume modulation volume of the chamber excluding the adsorbent sample/m3 volume occupied by the adsorbent sample/m3 molar fraction surface coverage shape parameter of particles reduced time (= t/R:)/s m2 normalized spatial coordinate ( =r/Rc) angular frequency of the volume modulation/s -',ratio of sticking probabilities in Yang's model [eqn.(23) and (2411 Subscripts m equilibrium values after a pressure or volume step e reference state j,k component index Superscripts -volume average values variables in the Laplace domain Matrix and Vector Notation boldface letters vectors of dimension n [ 3 square matrix of dimension n x n References 1 J. Karger and M. Bulow, Chem. Eng. Sci., 1975,30,893. 2 X. Hu and D. D. Do, Chem. Eng. Sci., 1992,47,1715. 3 J. Karger and D. M. Ruthven, Digusion in Zeolites and Other Microporous Solids, 1992, Wiely Interscience, New York. 4 R. Krishna, Chem. Eng. Sci., 1990,45, 1779. 5 J. G. Tsikoyiannis and J. Wei, Chem.Eng. Sci., 1991,46, 233. 6 J. G. Tsikoyiannis and J. Wei, Chem. Eng. Sci., 1991,46, 255. 7 Y. D. Chen and R. T. Yang, Chem. Eng. Sci., 1992,47,3895. 8 Y. D. Chen, R. T. Yang and L. M. Sun, Chem. Eng. Sci., 1993, 48,2815. 9 Y. D. Chen and R. T. Yang, Chem. Eng. Sci., 1993,48,2815. 0 J. Karger, M. Bulow and P. Lorenz, J. Colloid Interface Sci., 1978,65, 181. 11 N. W. Carlson and J. S. Dranoff, Fundamentals ofAdsorption, ed. A. L. Liapis, Engineering Foundation, New York, p. 129. 12 R. T. Yang, Y. D. Chen and Y. T. Yeh, Chem. Eng. Sci., 1991,46, 3089. 13 Y. H. Ma and T. Y. Lee, Znd. Eng. Chem. Fund., 1977,16,44. 14 V. R. Choudhary, D. B. Akolekar and A. P. Singh, Chem. Eng. Sci., 1989,44, 1047. 15 W. R. Qureshi and J. Wei, J. Catal., 1990, 126, 126. 16 Y. Yasuda, Y. Suzuki and H. Fukada, J. Phys. Chem., 1991, 95, 2486. 17 D. M. Shen and L. V. C. Rees, Zeolites, 1991, 11,684. 18 R. G. Jordi and D. D. Do, J. Chem. SOC.,Faraday Trans., 1992, 88,241 1. 19 L. M. Sun, F. Meunier and J. Karger, Chem. Eng. Sci., 1993, 48, 715. 20 L. M. Sun, F. Meunier, Ph. Grenier and D. M. Ruthven, Chem. Eng. Sci., 1993, in the press. 21 L. M. Sun and V. Bourdin, Chem. Eng. Sci., 1993,48,3783. 22 Y. Yasuda and S. Shinbo, Bull. Chem. SOC.Jpn., 1988,61, 745. 23 Y. Yasuda and K. Matsumoto, J. Phys. Chem., 1989,93,3195. 24 D. M. Shen and L. V. C. Rees, 9th International Zeolite Con- ference, Montreal, 1992. 25 D. M. Ruthven, Principles of Adsorption and Adsorption Pro- cesses, Wiley Interscience, New York, 1984. 26 R. M. Marutovsky and M. Bulow, Zeolites, 1987,7, 11 1. 27 S. V. Sotirchos and V. N. Burganos, AZChEJ., 1988,34,1106. 28 W. R. Qureshi and J. Wei, J. Catal., 1990,126, 147. Paper 3/04725C; Received 5th August 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000369

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(2), 377-381 Photocatalytic Decomposition of Dinitrogen Oxide on Cu-containing ZSM4 Catalyst Kohki Ebitani, Munehiro Morokuma, Jong-Ho Kim and Akira Morikawa* Department of Chemical Engineering, Faculty of Engineering, Tokyo Institute of Technology,2-12-1 Ookayama, Meguro-ku, Tokyo 152,Japan Photodecomposition of dinitrogen oxide (N,O) into N, and 0, on Cu-exchanged ZSM-5 zeolite (ion-exchange degree, 145%), degassed at 723 K, proceeds catalytically at 278 K. The dependence of the decomposition rate on the wavelength of the irradiated light reveals the importance of the excitation of monovalent Cu ions (<300 nm) in the photocatalytic decomposition of N20. The rate of decomposition can be expressed by the rate equa- tion : where a and b are constants, and ais proportional to the light intensity.The desorption of 0, from the surface is not a rate-determining step in the photodecomposition or the thermal decomposition of N,O on the zeolite which occurs with an apparent activation energy of 37 & 3 kJ mol-'. The decomposition of nitrogen oxides (NO,) is important for reducing global air pollution. While NO decomposition has been intensively investigated, relatively little attention has been paid to the decomposition of dinitrogen oxide (N,O) over various metal and oxide catalysts,'-13 in order to reduce its emission into the atmosphere. The thermal decomposition of N,O requires temperatures >573 K,9 because molecular oxygen, a product of the reac- tion, is not evolved below 573 K, as it is tightly bound to the catalyst surface at these temperature^.^ Photodecomposition of N,O has also been reported on semiconductor catalysts such as ZnO,"." TiO, anchored onto porous Vycor glass (PVG),', and Pt/TiO,, in the presence of water vapour.13 On these photocatalysts,''-' trapping of the photoformed electron by the N,O molecule to form the N,O-ion is believed to be the key process in N,O photodecomposition. The N,O-ion has been directly detected by the EPR tech- nique and its decomposition into N, and 0-ions has been reported by Anpo et ~1.'~However, no evolution of gaseous oxygen during N,O photodecomposition has been reported except for decomposition on ZnO" and Pt/TiO, in the pres- ence of water. ' It seems likely that Cu-containing ZSM-5 zeolite (Cu/ ZSM-5) would be active for the photocatalytic decomposition of N,O, since it is well known that the zeolite is highly active for the thermal dec~mposition'~~' of nitrogen monoxide (NO) into nitrogen and oxygen.'6*'7 It has also been reported that Cu/ZSM-5 is also active in the photocatalytic decompo- sition of NO," in which electron transfer from the excited copper(1) ion (Cu') to the antibonding orbital of the NO molecule plays an important role.' 8, ' Here we report the photocatalytic decomposition of N 2O into N, and 0, at 278 K over a Cu-exchanged ZSM-5 cata-lyst. The mechanism of photocatalytic decomposition of N,O will be discussed through kinetic investigation of the reaction rate and examination of the diffuse reflectance (DR) spectrum of the catalysts. Experimental Na-ZSM-5 (Si : A1 = 41 : I), prepared by the procedure pro- posed by Yoshimura et was ion exchanged at 353 K for 12 h in a 0.05 mol 1-' aqueous solution of Cu(NO,),, washed with distilled water several times, and dried at 383 K for 24 h, followed by calcination at 773 K for 5 h. The Cu content of the Cu-exchanged ZSM-5 zeolite, thus prepared, was 1.74 wt.% (145%, ion exchanged) by atomic absorption spectrometry.The XRD pattern of the Cu/ZSM-5 was iden- tical to that of the parent Na-ZSM-5, and no pattern due to CuO was detected. These observations imply that the Cu ions are incorporated inside the ZSM-5 cavities without formation of a large CuO crystal either in the pore structures or on the external surface of the zeolite, which is consistent with the report of Cu-ion dispersion of a Cu-exchanged ZSM-5 zeolite by means of X-ray absorption (XANES/EXAFS) spectro- scopy., ' The Cu/ZSM-5 (180 mg) was degassed in a vacuum of < lop4Torr (1 Torr = 133.3 Pa) at 723 K for 2 h prior to being used in the reaction.The products were analysed by on-line gas chromatography. Reaction of N,O (99.8% purity), diluted with He so that the total pressure was 6.7 kPa, was carried out in the conventional closed gas-circulating system with a volume of 167 ml to which a quartz reactor with a flat, transparent quartz plate bottom was installed (Fig. The powdered Cu/ZSM-5 catalyst was spread on the bottom of the reactor.The catalyst was irradi- ated with a medium-pressure mercury lamp (H-400P, Toshiba Co., Ltd.) through a water filter. Oxygen gas was purified by passing it through molecular sieves at 77 K. Various UV cut-off filters (Toshiba, Co., Ltd.) were used to examine the effect of wavelength on the reaction. The light intensity was controlled by changing the number of quartz plates placed between the light source and the water filter. The thermal decomposition of N,O on Cu/ZSM-5, degassed at 723 K, was also performed in the same apparatus and reactor that were used for the photocatalytic decomposi- tion. The diffuse reflectance (DR) spectra (240-800 nm) were recorded with a Shimadzu spectrophotometer (MPS-2000) at room temperature, using BaSO, as the reference material.For the DR measurements, the ground sample was placed in a quartz cell with the two planar windows 2 mm apart. Results Photodecompositionof N*O Fig. 2 shows the change in the nitrogen yield as a function of time, for N,O decomposition at 278 K on the Cu/ZSM-5 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 G C Fig. 1 (a) Schematic diagram of the apparatus. V, vacuum line; S, gas supply; G, to gas chromatograph; C, pressure sensor; P, circulation pump; R, reactor; W, water filter; M, front surface mirror; L, light source. (V, + V2= 167 ml). (b) Detail of the reactor. powder degassed at 723 K (the N,O pressure was 1.2 kPa). During the reaction in the dark, little nitrogen was formed.When the Cu/ZSM-5 was irradiated by an Hg lamp through a water filter, N, formation was initiated immediately and increased during the irradiation. When the irradiation ceased, N, formation stopped immediately. The initial N,O decom-position activity was reproduced when the catalyst was irra- diated again. This strongly indicates that photoinduced N,O decomposition occurs on the Cu/ZSM-5 powder. The number of nitrogen molecules formed per Cu atom of the catalyst after 350 min was about two, indicating that pho- toinduced N,O decomposition on the Cu/ZSM-5 occurs catalytically. The N, : 0, ratio in the photocatalytic decomposition pro- ducts was in the range (3.5-4) : 1 at the initial stage of the reaction (<200 min), and gradually decreased to 3 : 1 with increase of the reaction time.The observed N, : 0, ratio was >2 : 1 throughout the reaction, which is the predicted value for the reaction :N,O -,N, + 1/20,. The dependence of the rate of N, formation under irradia- tion as a function of N,O partial pressure is shown in Fig. 3 (the N,O pressure was varied under a constant total pressure adjusted with He diluent). The rate of formation of N, was 15 1 Y N z 100 200 300 400 time/min Fig. 2 Change of N, yield for the decomposition of N20 at 278 K on the degassed Cu/ZSM-5 catalyst determined from the change in N, yield between the irradia- tion times of 30 and 120 min. At any N,O pressure, the response of the N,O decomposition activity to the irradia- tion was the same as shown in Fig.2, suggesting that the same N,O decomposition mechanism is operative under irra- diation at any N,O pressure. Fig. 3 shows that the rate of N, formation is independent of N,O pressure. The N, :0, product ratios were also independent of the N,O pressure “2 0,= (3-4) 11. Fig. 4 shows the change in rate of N, formation during irradiation at 278 K as a function of molecular oxygen pres- sure on the Cu/ZSM-5 powder degassed at 723 K, when the N,O pressure was 0.68 kPa. The degree of retardation of the rate of N, formation by 0, is not drastic but is still appre- ciable for the higher 0, pressures (2.5-4 kPa). When Cu/ZSM-5 was irradiated through UV cut-off filters, UV-34, UV-32 and UV-30,t no N,O decomposition occurred.With a UV-28 filter, however, the N,O decomposi-tion proceeded with a rate 15% of that without the UV 0.06 -I .-C 0.04 f -E @.--. P CI2 C .-c;0.02 5+ N z 0 0 1 2 3 N,O pressure/kPa Fig. 3 Nitrogen monoxide pressure dependence of the rate of N, formation under irradiation at 278 K t The filter notation is after Toshiba Co. The filter number denotes one tenth of the wavelength (nm) at 50% reduction of the light inten- sity transmitted. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0 1 2 3 4 0, pressure/kPa Fig. 4 Effect of the 0, pressure on the rate of N, formation under irradiation at 278 K cut-off filter. Therefore, light of wavelength <300 nm was effective for photocatalytic decomposition of N,O on Cu/ ZSM-5.When the light intensity was reduced by placing quartz plates in the light path, the rate of N, formation, at an initial N,O pressure of 0.7 kPa, decreased, as shown in Fig. 5. The relative light intensity shown in the figure is calculated from the attenuation of the light intensity per quartz plate for the wavelength of 250 nm. Note that the rate of N, formation is proportional to the irradiated light intensity. Thermal Decompositionof N,O Fig. 6 shows the change in nitrogen yield for the thermal decomposition of N,O at 513 and 613 K on the Cu/ZSM-5 powder degassed at 723 K (the N,O initial pressure was 1.31 kPa) as a function of time. At 513 K, the N, yield was satu-rated at ca.6% after 70 min. The N, : 0, ratio at 513 K was 100 h $ v $ 80 B c 60 E 40 z Q,.> 20 c. E 0 0 20 40 60 80 100 relative light intensity (%) Fig. 5 Effect of the irradiated light intensity of the rate of N, forma-tion at 278 K 50 n 40 ,\" U I! 30 .- >. z" 20 10 0 0 50 100 150 200 tirne/rnin .: I I I I 0 1 2 3 4 N,O pressure/kPa Fig. 7 Dependence of the rate of N, formation on dinitrogen oxide pressure at (a)513, (6)613 K (5--6):1, which is larger than that obtained for the photo-catalytic decomposition of N,O at 278 K. At 613 K, the N, yield increased with the reaction time, and the N, : 0, ratio of (2-3) : 1 was close to the expected value of 2 : 1.This corre-sponds to the result obtained on a Cu-containing ZSM-5 catalyst prepared by a cation-exchange method with copper(I1)a~etate.~ The dependence of the rate of N, formation, determined from the initial slopes of the N, formation curves in Fig. 6, obtained at 513 and 613 K, on N,O pressure is shown in Fig. 7. The N, formation rates increased with increase of N,O pressure and approached a constant value at high N,O pres-sures, suggesting the strong interaction of N,O with the reac-tion sites (Cu ions), differing from previous reports in which the N,O decomposition rate is proportional to the N,O pres-~ure.~,~ The effect of 0, pressure on the rate of N, formation in the thermal decomposition of N20 was also examined at 513 and 613 K, with an initial N,O pressure 0.7 kPa, as sum-marized in Fig.8. For both reaction temperatures, coexisting molecular oxygen did not affect the rate of N, formation, which is different from the results obtained under irradiation at 278 K (Fig. 4), The lack of 0, inhibition in the thermal decomposition of N,O has also been reported for Fe-containing ZSM-5 catalyst4 and Fe-exchanged m~rdenite.~ The rate of N, formation was measured at the initial stage of the thermal decomposition of N,O at various reaction temperatures between 473 and 608 K, and the apparent acti-vation energy for N, formation was estimated by changing the reaction temperature between 473 and 608 K, to be 37 3 kJ mol-'. This is comparable to the activation energy observed for the thermal decomposition of N,O on Fe-containing ZSM-5 zeolite, reported by Panov et al.(42 f8 kJ mol-1),4 but is smaller than that obtained on ZnO (146 kJ mol-')'' and on Fe-exchanged mordenite (134 kJ mol-I).' -" 0 1 2 3 4 5 0, pressure/kPa Fig. 6 Change of N, yield for the thermal decomposition of N,O at Fig. 8 Effect of the 0, pressure on the rate of N, formation at (a) (a) 513, (6) 613 K on the degassed Cu/ZSM-5 catalyst 513, (6) 613 K h v)+-..-C 3 rj Y 0, C e D 300 500 700 wavelength/nm Fig. 9 Diffuse reflectance spectra of Cu/ZSM-5 catalyst: after degassing H/ZSM-5 at 723 K (a),423 K (b)and 723 K (c) Diffuse Reflectance Spectroscopy Study The DR spectra of Cu/ZSM-5 are shown in Fig. 9.The Cu/ ZSM-5 degassed at 723 K showed a strong absorption band at around 240-400 nm with an absorption maximum at 270- 280 nm on its spectrum [Fig. 9(c)]. This absorption band cor- responds to the (3d)",'S, +(3d)9(4s)','D, electronic transition of Cu' ions in zeolite cavities,23 with an absorp- tion band at 322-227 nm, and is included in the region of the light of wavelength <300 nm which is effective for N,O decomposition. Therefore, it is suggested that the excitation of the Cu 'ions causes N20 photodecomposition. The absorption around 700 nm [Fig. 9(b)], assigned to the d4 transition of octahedrally coordinated Cu2 ' disappeared after degassing of the Cu/ZSM-5 [Fig. 9(c)] at 723 K, which is attributed to the partial reduction of the Cu2' ions to low-valence copper such as Cu' and/or Cuo after the desorption of water (or OH groups and/or lattice 0,-)from the Cu2+ sites.The Cu/ZSM-5 treated with hydrogen (6.7 kPa) at 603 K followed by degassing at the same temperature gave a diffuse reflectance spectrum charac- teristic of Cuo, different from that obtained for Cu/ZSM-5 after degassing alone. Therefore, it is plausible that the oxida- tion state of the copper in ZSM-5 after degassing at 723 K is not Cuo but Cu'. A separate experiment on adsorption (at 273 K) of CO, which is well known to be adsorbed irrevers- ibly on monovalent copper ions,26 showed that an irrevers- ible adsorption of CO did, in fact, occur on the Cu/ZSM-5 degassed at 723 K.The number of irreversibly adsorbed CO molecules per copper species in the catalyst (CO : Cu ratio) was 0.73: 1. This confirms the existence of monovalent copper species (Cu') on Cu/ZSM-5. The Cu" ions in the ZSM-5 should be reduced to Cu' ions during evacuation by releasing molecular oxygen and H,O according to the mechanism suggested by Jacobs et ~21.~~ Discussion From the dependence of the rate of N,O formation on the wavelength of the irradiated light, it is concluded that the N20 photocatalytic decomposition on Cu/ZSM-5 catalyst is initiated by the electronic excitation of Cu' ions, (3d)",'S0 +(3d)9(4s)1,'D,. Charge transfer from the excited Cu' to the antibonding molecular orbital of N,O must trigger the decomposition of N,O into a nitrogen molecule and an adsorbed oxygen atom, Oads.A similar role of the J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 excited Cu' ions has been proposed for the photocatalytic decomposition of NO on Cu/ZSM-5" and Cu/SiO, cata-lysts.l9 Thus, the kinetic scheme for the photocatalytic decomposition of N,O on Cu/ZSM-5, shown in reactions (1)-(4), is suggested: CU' + hv -(Cu+)* (CU')* (CU')* + N,O ki kz -Cu'-Cu' + hv' + N, + 1/20, (1) (2) (3) or (Cu')* + N,O kz. -Cu+-O + N2 (3') 2cu'-0 &t 2Cu' + 0, (4) k-3 where I, is the light intensity absorbed by the Cu' ions and (Cu+)* denotes the excited states of the Cu', whose quantum yield is 4. Reaction (2) represents the radiative decay of the excited states of Cu'.Reaction (4) is the equilibrium between adsorbed oxygen and molecular oxygen in the gas phase. Reaction (3) represents the photosensitized decomposition of N,O, and reaction (3') corresponds to N,O decomposition into the adsorbed 0 atom on Cu' and gaseous N,. It is stressed that the concentration of Cu'-0 species formed by the step shown in reaction (3') is very small under irradiation, since molecular oxygen has been evolved, though the N, :0, ratio under irradiation is larger than the stoichio- metric value of 2 : 1. If the photocatalytic decomposition of N,O is terminated by forming the Cu'-0 species without recovering Cu' ions, 0, formation in the gas phase should not be detected. However, the amounts of N, and 0, evolved after 350 min (shown in Fig.2) corresponded to twice and 1.4 times of the amount of Cu on the catalyst, respectively. Therefore, the amount of 0, evolved was about 70% of the stoichiometric value corresponding to N, formation. In other words, the rate of the forward reaction (k3)is larger than that (k-3)of the backward one in reaction (4). The decrease in the amount of Cu' by interaction between oxygen atoms and Cu' ions, may cause the reduction of 61, and result in the decline of the rate of N,O decomposition as noted in Fig. 2. Application of the steady-state approximation to the con- centration of (Cu')* under irradiation gives the N, forma-tion rate Since enough catalyst is used to achieve complete absorp- tion of the irradiating light, 41, is proportional to the inci- dent light intensity, I,.Eqn. (I) is consistent with the experimental results shown in Fig. 6, i.e. the rate is dependent on the intensity of the irradiating light. The observed independence of the rate of N, formation of N,O pressure (Fig. 3) is explained by assuming that k, 4k, , indicating the higher probability of charge transfer from the excited Cu+ to the N,O molecule than of decay of excited Cu+ to its ground state. If it is assumed that 0, molecules have little effect on the quenching of the Cu' excited state but oxygen atoms formed by the dissociation of 0, on the catalyst are highly efficient for quenching, reactions (5) and (6) should be considered: K 0, 20 (5) J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 Then, the rate of N, formation, r’, in the presence of 0, is given by: where K is the equilibrium constant of reaction (5). When the relation k, $ k, is assumed, the following equation is obtained : This equation predicts that at constant N,O pressure, the reciprocal of the reaction rate will change linearly with square root of 0, pressure. Fig. 10 shows the linear relation- ship between the reciprocal of the rate of N, formation and the square root of 0, pressure under irradiation at 278 K. For thermal decomposition of N,O, it has been proposed that the reaction comprises the following catalyst reduction and oxidation steps: (a)catalyst reduction N2O + (0)--+ N2 + 02 + ( ) (7) and (b)catalyst oxidation (9) 0, + 2( )-+2(0) (10) where ( ) and (0)show an empty site and a site holding an oxygen atom, re~pectively.~If this redox mechanism is applied to the present reaction, reactions (8) and (10)are not important, because the initial rate of N, formation in the thermal decomposition of N,O was independent of the pres- sure of oxygen coexisting with N,O as shown in Fig.9. A similar phenomenon has been observed for the thermal decomposition of N,O at 623 K on Fe-exchanged morde- ~~ite.~Therefore, the effect of oxygen gas on the rate of N20 photodecomposition, observed in the present work, results from the quenching of the photoexcited Cu+ by the oxygen atoms and is not attributed to the thermal oxidation of Cu’ by N,O or 0,.It is, furthermore, stressed that the desorption of oxygen through reaction (5) is not rate-determining. Tanaka and Blyholder have also suggested that the desorp- tion of oxygen is not a rate-determining step in the photo- catalytic decomposition of N,O on ZnO, since an oxygen 0 1 2 (0,pressure)’’2/kPa’’2 Fig.10 Plot of the reciprocal of the rate of N, formation vs. square root of 0, pressure under irradiation at 278 K. The N,O initial pres- sure was 0.7 kPa. 38 1 inhibition effect was not observed in the thermal decomposi- tion of N,O at the same reaction temperature.’ The formation of molecular oxygen in the photocatalytic decomposition of N,O on Cu/ZSM-5 suggests that recombi- nation of the Oads,resulting in the desorption of oxygen mol- ecules, is favoured under irradiation conditions.The recovery of the octahedrally coordinated Cu2+ ion was scarcely detected in the diffuse reflectance spectrum of Cu/ZSM-5 irradiated in the presence of 2.7 kPa of N,O. Therefore, it is concluded that the Cu+ ion is reformed after N,O decompo-sition induced by electronically excited Cu ions. The behav- + iour of N,O thermal decomposition on the Cu/ZSM-5 catalyst is different from that on Fe-containing ZSM-54 and Fe-exchanged mordenite’ with the result that the dependence of N,O pressure on decomposition rate is of the order of 0.6 with respect to N20 partial pressure as reported by Li and Armor.’ The unique behaviour of Cu/ZSM-5 displayed in Fig. 8 shows the strong interaction between N,O and Cu/ ZSM-5.This work is supported by a Grant-in-Aid for Scientific Research from the Ministry of Education, Science and Culture, Japan. References 1 E. S. R. Winter, J. Catal., 1974,34,431. 2 C. Korodulis, H. Latsios, A. Lycourghiotis and P. Pomonis, J. Chem. SOC.,Faraday Trans., 1990,86, 185. 3 J. Christopher and C. S. Swamy, J. Mol. Catal., 1990, 62, 69. 4 G. I. Panov, V. I. Sobolev and A. S. Kharitonov, J. Mol. Catal., 1990,61, 85. 5 S. A. Tan, R. B. Grant and R. M. Lambert, J. Catal., 1987, 104, 156. 6 Y-S. Yong and N. Cant, J. Catal., 1990,122,22. 7 J. Leglise, J. 0.Petunchi and W. K. Hall, J. Catal., 1984,86, 392. 8 L. M. Aparicio, M. A. Ulla, W. S. Millman and J. A. Dumesic, J. Catal., 1988, 110, 330. 9 Y. Li and J.N. Armor, Appl. Catal. B, 1992,1, L21. 10 J. Conningham, J. J. Kelly and A. L. Penny, J. Phys. Chem., 1971, 75, 617. 11 K. Tanaka and G. Blyholder, J.Phys. Chem., 1971, 75, 1037. 12 M. Anpo, N. Aikawa, Y. Kubokawa, M. Che, C. Louis and E. Giamello, J. Phys. Chem., 1985,89, 5017. 13 A. Kudo and T. Sakata, Chem. Lett., 1992,2381. 14 M. Iwamoto, H. Yahiro, Y. Mine and S. Kagawa, Chem. Lett., 1989,213. 15 Y. Li and W. K. Hall, J. Catal., 1991, 129,202. 16 E. Giamello, D. Murphy, G. Magnacca, C. Morterra, Y. Shioya, T. Nomura and M. Anpo, J.CataI., 1992,136,510. 17 W. K. Hall and J. Valyon, Catal. Lett., 1992, 15, 311. 18 M. Anpo, Y. Shioya, T. Nomura, E. Giamello, C. Morterra, G. Centi and M. Che, in Meeting Abstr. Catal. SOC.Jpn., 1992, 34, 65; M. Anpo, T. Nomura, Y. Shioya, M. Che, D. Murphy and E. Giamello, in 10th Znt. Con$ Catalysis, 1992. 19 M. Anpo, T. Nomura, T. Kitao, E. Giamello, D. Murphy, M. Che and M. A. Fox,Res. Chem. Intermed., 1991,15,225. 20 A. Yoshimura, S. Namba and T. Yashima, Shokubai, 1981, 23, 232. 21 H. Hamada, N. Matsubayashi, H. Shimada, Y. Kintaichi, T. Ito and A. Nishijima, Catal. Lett., 1990,5, 189. 22 Y. Wada and A. Morikawa, Bull. Chem. SOC. Jpn., 1987, 60, 3509. 23 J. Texter, D. H. Strome, R. G. Herman and K. Klier, J. Phys. Chem., 1977,81, 333. 24 H. Tominaga, Y. Ono and T. Kei, J. Catal., 1975,40, 197. 25 J. J. Freeman and R. M. Friedman, J. Chem. SOC., Farday Trans. I, 1978, 74, 758. 26 Y-Y. Huang, J. Catal., 1973,30, 187. 27 P. A. Jacobs, W. De Wilde, R. A. Schoonheydt, J. B. Uytterhoe- ven and H. Beyer, J. Chem. SOC., Faraday Trans. I, 1976, 12, 1221. Paper 3/03553K; Received 21st June, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000377

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(2), 383-386 Bransted Acid Strength in US-Y: FTIR Study of CO Adsorption Marina A. Makarova, Khalid M. Al-Ghefaili and John Dwyer* Chemistry Department, UMIST, P.O. Box 88, Manchester: UK M60 IQD Proton affinities of Brsnsted hydroxy groups in H-forms of zeolites and SAPOs are determined by an FTIR study of the low-temperature adsorption of carbon monoxide. This method reveals the range of heterogeneity of Brsnsted hydroxy groups associated with unmodified zeolites where only one band interacting with carbon monoxide is evident in the infrared spectrum. Three types of acid hydroxy groups associated with different IR bands are detected in a modified zeolite (US-Y), (i) strongly acidic [3599 cm-', proton affinity (€,J = 1112 kJ mol-'1, (ii) medium acidic (3627 cm-', E,, = 1142 kJ mol-') and (iii) weakly acidic (ca.3740 and 3675 cm-', ED, = 1192 f29 kJ mol-I). Bransted acidity in zeolites, particularly in modified zeolites such as ultra-stabilised Y zeolite (US-Y), is of considerable interest not least because of the wide application in com-mercial catalysts.' *2 Several methods for measuring acidity in zeolites are available and have been extensively re~iewed.~ Because proton affinities are an important feature of acidity, it is of particular interest to measure these experimen- tally both for comparison with theoretical values and in order to clarify the role of both composition and structure in zeolites described as having 'enhanced activity'. Sites showing enhanced catalytic activity are observed following hydrothermal treatment of zeolite^^.^ where increased rates of hydrocarbon transformation are observed. Enhanced cata-lytic activity is generally attributed either to defects in the framework (=A1 +SiE), which are presumed to be very strong Lewis sites,6 or to Bransted sites where acidity is enhanced by interaction with proximate Lewis sites associ- ated with non-framework al~minium.~ The present paper aims to clarify the experimental pro- cedure for adsorption of carbon monoxide for acidity testing and to apply this technique for characterisation of Bransted acid sites in such a complex system as US-Y.Experimental Carbon monoxide was supplied by BOC (R grade) and the range of materials listed in Table 1 was used.A sample of US-Y was prepared using as-synthesized Na-Y zeolite (Si/A1 = 2.7) which underwent ion exchange (to Na/ A1 = OM), was steamed at 560°C for 30 min and then sub- jected to further ion exchange. After this treatment, total Si/Al = 3.0, Na/Al < 0.01; the crystallinity (as compared with that of the initial as-synthesized sample) was 81% with a unit cell parameter of 24.50 A. The framework Si/Al ratio deter- mined by ,'Si MAS NMR was 5.4. The spectroscopic technique was described in detail else- where.8 Briefly, FTIR studies were carried out using a Cygnus-100 Mattson FTIR spectrometer and a special IR Table 1 Chemical composition of the samples sample Si/Al Na/Al H-ZSM-5 16 <0.01 H-Y 2.7 0.01 H-EMT 4.0 <0.01 Si/(Al + P + Si) S APO-37a 0.08 SAPO-37b 0.22 cell having a thermostatted zone which facilitated high-temperature treatment of samples in situ.The cell was also connected to a vacuum rig and a calibrated volume, fitted with a pressure gauge, for quantitative gas adsorption. The samples were pressed into self-supporting disks (rn = 10-20 mg, p = 5-10 mg ern-,), placed into the cell, heated at 1 "C min-' to 350°C (400°C in the case of US-Y) under vacuum and then held at this temperature overnight (pressure of lop5 Torr). Aluminophosphate samples were additionally heated under oxygen (300 Torr, 550"C, 2 h) to remove organic tem- plates. For spectroscopic/adsorption measurements, the IR cell was lowered into a quartz Dewar situated inside the spec- trometer chamber, filled with liquid nitrogen (temperature of sample, -100 to -120"C), and small increments of carbon monoxide were admitted stepwise into the cell.All the spectra were collected using 100 scans and a resolution of 2 cm-in the wavenumber interval of 2500-4000 cm- '. The shifts were determined as peak-to-peak distances in difference spectra obtained by subtraction of the spectra of the samples before adsorption from their spectra after adsorption. Results and Discussion An acidity scale based on the shift in position of hydroxy groups after perturbation with carbon monoxide at low tem- peratures,' is adopted for the present study. It represents a correlation between an energetic characteristic of a hydroxy group, the proton affinity (Ep,/kJ mol-'), and the shift (AvOH) measured in cm -: A v,,(sample) 'pa, OH = 'pa, SiOH -442*5log AVoH(SiOH) (1) SiO, is used as a standard with Epa,SiOH= 1390 kJ mol-' and Av(Si0H) = 90 cm-I,' which results in the final equa- tion : Epa,,&J mol-' = 2254.8 -442.5 log(Av,&m- ') (2) The stronger the acid site, the less the proton is con-strained which results in a larger shift during the 0-H. -CO interaction and in a smaller value of E,, .Previous studies"." show that even when there is a single hydroxy-group band interacting with CO, for example in H-ZSM-5 or the HF band in H-Y, the observed shift AvOH depends upon surface coverage. A larger shift observed at low surface coverage decreases as further amounts of CO are sorbed, reflecting some heterogeneity in acid site strength.When all the hydroxy groups are covered (0 = 1) the shift in AVO, is observed to increase, corresponding to the formation of the condensed sorbate phase which results in enhanced perturbation of the hydroxy groups. As a result, the pertur- bation of hydroxy groups is minimal at 8 approaching 1 and in consequence the conditions for spectroscopic/adsorption methods require strict specification if comparisons are to be made between materials and between laboratories. Interaction of Carbon Monoxide with Unmodified Zeolites In the first part of the present study, CO was adsorbed onto a number of well known zeolites and aluminophosphates in which only one OH band could interact with sorbate: H-ZSM-5, H-Y, H-EMT, SAPO-37 for reference purposes.Titration curves of Brransted hydroxy groups with carbon monoxide for different samples using the coordinates AvOH us. 8 are shown in Fig. 1. These functions can be approximated by straight lines with negative slopes. Presumably, AvOH(8+0) corresponds to the strongest acid sites whereas AvoH(6+ 1) can be considered as a measure of the mean acidity for a given sample. Analysis of the plots in Fig. 1 shows excellent agreement between the results obtained by different research groups in the case of H-Y."*" It also suggests that the conventional ion-exchanged H-Y zeolite (Si/Al x 2.5, Na/Al < 0.1) can be used as a standard for such measurements giving AvOH(B +0) = 296 cm-' and AvOH(8 --+ 1)= 273 cm-'.In the case of H-ZSM-5, our results also concur with previous reports.' O A comparison of the results for different samples (Fig. 1) shows the following: (1) The average acidity (8-P 1) increases in the order: SAPO-37a < SAPO-37b < H-Y < H-EMT < H-ZSM-5. (2) The distribution of the acid strength, d(AvOH)/dO, widens as follows, SAPO-37a < H-ZSM-5 < SAPO-37b < H-Y < H-EMT. (3) The homogeneity in acid sites demonstrated for SAPO-37a, by the almost con- stant value for AvOH as 8 increases, is consistent with the fact that this material contains mainly a SAP0 phase.12 Increased site heterogeneity in SAPO-37b reflects the increased silica content in this material which results in two c 310I E 2 0 4 290 270 0 0.2 0.4 0.6 0.8 1.o 00 H Fig.1 Shift in the position of the IR band corresponding to Brmsted hydroxy groups us. their fractional coverage with carbon monoxide for different samples: a,H-Y; 0,H-EMT; 0,H-ZSM-5; @, SAPO-37a; @I, SAPO-37b. Literature data: a, H-Y (Si/ A1 = 2.9, Na/Al = 0.05);'' A, H-Y (Si/Al = 2.5, Na/AI = 0.1);" 0, H-ZSM-5 (Si/Al = 13.5)'' J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 compositional regions, and the generation of some stronger acid sites.12 Consequently, both the maximum shift, Avo,(8 +0),and the decrease of AvOH with increase in 8 are greater for SAPO-37b than for SAPO-37a. (4) Comparison of H-Y (Si/AI = 2.7) and H-EMT (Si/AI = 4.0) suggests that H-EMT possesses some sites stronger than those in H-Y and that both zeolites show heterogeneity in acid-site strength.Moreover, the average acid-site strength (8+ 1) in H-EMT is close to that of the strongest sites (0 -P 0) in H-Y. Interaction of Carbon Monoxide with US-Y The identification of Brernsted acid sites in US-Y is more complicated than in the materials examined above. It has been shown previously, using FTIR and ammonia thermode- sorption," that Brernsted hydroxy groups vibrating in the wavenumber range of 3500-3650 cm-' are represented by two pairs of bands, one at 3599 and 3525 cm-', indicating enhanced acidity, and one at 3627 and 3554 cm- ' associated with less acidic sites which are typical of siliceous H-Y zeo- lites.In each pair, the high-frequency peak corresponds to the hydroxy groups vibrating in the supercages and the low-frequency band, to the OH groups pointed into the #l-cages. FTIR spectra reflecting adsorption of the increased amounts of carbon monoxide on US-Y are given in Fig. 2. Although the structure of the spectra is very complex, the difference spectra (Fig. 3, the spectrum of the zeolite with adsorbed portion of CO minus the spectrum of the sample before adsorption) allow for analysis of the changes caused by adsorption. The interaction of carbon monoxide with H-US-Y clearly shows that the hydroxy groups initially at 3627 and 3599 cm-' are accommodated in the large cages. Fig.3(a) demonstrates that a small amount of CO perturbs mainly the strongest OH groups with the band at 3608 cm- '. This band corresponds to that at 3599 cm-', since at low temperatures the positions of the bands shift towards higher wavenumbers than those at room temperature. With further addition of CO, the less acidic band at 3631 cm-' (corresponding to 3627 cm- ' at room temperature) becomes perturbed [Fig. 3(b)] and, finally, the two other bands from -3800 3600 3400 3200 3000 wavenumber/cm -' Fig. 2 FTIR spectra for US-Y zeolite before and after adsorption of different amounts of carbon monoxide: (a) initial, (b)0.5, (c) 3.1, (d) 5.7 pmol CO J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0.1 J 0 /C (De 2 / weak-medium strong 7 3800 3600 3400 3200 3000 wavenumber/cm-' Fig.3 FTIR difference spectra for different amounts of carbon monoxide adsorbed on US-Y: (a)0.5, (b)3.1, (c) 5.7 pmol CO weakly acidic hydroxy groups at ca. 3747 and 3675 cm-' (ca. 3740 and 3675 cm-' at room temperature) disappear [Fig. 3(c)]. These bands were previously assigned to Si-OH and Al-OH species, respectively. l4 All these perturbed bands together give rise to the complex band from hydrogen-bonded hydroxy groups consisting of at least three superimposed components. The present results for the strength of accessible acid sites using CO adsorption concur with previous results on ammonia thermode~orption'~ which show that hydroxy groups vibrating at ca. 3740 cm-' desorb ammonia at tem- peratures below 20O0C, those vibrating at 3627 cm-' do so at temperatures below 250°C and those vibrating at 3599 cm-' do so mainly in the temperature interval 300-350°C.The values of the shifts and corresponding proton affinities discussed above are summarised in Table 2. Since the bands from two types of hydrogen-bonded weak hydroxy groups are not resolved, the shift from their medium position is given. Clear evidence is presented for the enhancement in acidity which is associated with the generation of the band at lower frequency (3599 cm-') in the large cages of H-USY. Moreover, the reduced hydroxy-group stretching frequency, the retention of ammonia at higher temperature^'^ and the larger shift, Av,,, due to CO, are all consistent with the pres- Table 2 Shifts in the position of the hydroxy-group IR bands resulting from perturbation with carbon monoxide and correspond- ing proton affinities (Epa) at 0-0 at041 sample AVO! /cm 'pa/kJ mol-' AVOH /cm-' 'pa/kJ mol-' SAPO-37a 275 1176 268 1180 SAPO-37b 288 1167 275 1175 H-Y 296 1161 273 1177 H-EMT 325 1143 293 1163 H-ZSM-5 318 1147 308 1154 us-Y: strong medium 382 327 1112 1142 weak 255 k38 1192 f29 ence of Brnrnsted sites of enhanced acidity (generated by hydrothermal treatment of zeolite H-Y).Thus, the present work, whilst not excluding a role for very strong framework Lewis sites, shows very clearly the significant enhancement of Br~rnsted acidity in zeolites containing extraframework alu- minium.Additionally, this work provides further evidence for site heterogeneity in zeolites when there are local variations in composition, i.e. it offers further support for a local rather than global view of acid sites in zeolites." Conclusions It is shown using a number of zeolites and alumino-phosphates that the conditions for acidity testing using CO adsorption should be strictly specified. The shift in the posi- tion of the OH band resulting from perturbation with carbon monoxide, AvOH, determined as an extrapolation of the frac- tional coverage of hydroxy groups with sorbate to 0 (0 +0), reflects the acid strength of the strongest Brnrnsted hydroxy groups, whereas the extrapolation to full coverage (0+ 1) gives the average acidity in a given sample.The difference between these two values reflects the range of heterogeneity of the Brsnsted acid sites. In the case of US-Y, three types of hydroxy groups differ- ing in their acid strength have been detected. (1) Hydroxy groups, presumably associated with OH groups interacting with proximate Lewis sites in supercages, which vibrate at 3599 cm- ',adsorb the initial increments of carbon monoxide and produce the largest shift, AvOH, reflecting their stronger acidity. (2) Hydroxy groups which are typical of siliceous H-Y zeolite (supercages), vibrate at 3627 cm-' and adsorb CO after the stronger sites are perturbed and give a smaller shift, AvOH. (3) Hydroxy groups vibrating at ca.3740 and 3675 cm-', ascribed to Si-OH and Al-OH species, which show weak acidity. Using the observed values of AvOH for US-Y and the refer- ence samples, an estimate of proton affinity has been deter- mined. The estimated values of E,, can be used for comparison with theoretical calculations and as a basis for correlation with the relative activity of catalysts in catalytic cracking. We thank Drs. K. Karim, A. F. Ojo, D. Rawlence and W. J. Smith for the samples of microporous materials and the group of Prof. J-L. Guth for the 29Si MAS NMR character- isation. We are grateful to the EC (BRITE EURAM 4633) and the Saudi Arabian Government for financial support for M.A.M. and K.M.Al-G, respectively. We also thank SERC for an equipment grant (GR/D/99768).References 1 C. V. McDaniel and P. K. Maher, in SOC. Chem. Ind. London Monogr., 1968, 186. 2 P. K. Maher, F. D. Hunter and J. Scherzer in, Ado. Chem., 1971, 101, 266. 3 J. Dwyer and P. O'Malley, in Stud. Surf.Sci. Catal., ed. S. Kalia-guine, Elsevier, Amsterdam, 1988, vol. 35, p. 5. 4 A. G. Ashton, S. Batmanian, D. M. Clark, J. Dwyer, F. R. Fitch, A. Hinchliffe and F. J. Machado, in Stud. Surf: Sci. Catal., ed. B. Imelik et al. Elsevier, Amsterdam, 1985, vol. 20, p. 101. 5 R. M. Lago, W. 0. Haag, R. J. Mikowsky, D. H. Olson, S. D. Hellring, K. D. Scmitt and G. T. Kerr, in Proc. 7th Znt. Zeol. Con$, ed. Y. Murakami et al., Kodansha-Elsevier, Tokyo- Amsterdam, 1986, p. 677. 6 V. I. Zolobenko, L. M. Kustov, V. B. Kazansky, E. Loemer, U. Lohse and G. Oehlmann, Zeolites, 1991,11, 132. 7 C. Mirodatos and D. Barthomeuf, J. Chem. Soc., Chem. Commun., 1983, 39. 386 8 N. E. Thompson, PhD Thesis, UMIST, 1991. 9 E. A. Paukshtis and E. N. Yurchenko, Usp. Khim., 1983,53,426. 10 L. Kubelkova, S. Beran and J. A. Lercher, Zeolites, 1989,9, 539. 11 N. Echoufi and P. Gelin, J. Chem. SOC.,Faraday Trans., 1992, 88,1067. 12 M. A. Makarova, A. F. Ojo, K. M. Al-Ghefaili and J. Dwyer, in Proc. 9th IZA Con$, Montreal, 1992 ed. R. von Ballmoos et al., Butterworth-Heinemann, Boston, 1993, vol. 2, p. 259. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 13 M. A. Makarova and J. Dwyer, J. Phys. Chem., 1993,93,6337. 14 P. 0.Fritz and J. H. Lunsford J. Catal., 1989,118,85. 15 J. Dwyer, in Guidelinesfor Mastering the Properties of Molecular Sieves, ed. D. Barthomeuf er al., Plenum Press, New York, 1990, p. 241. Paper 3/03474G; Received 16th June, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000383

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(2), 387-393 Influence of Framework Substitution of A13+ by Fe3+on the Sorption Characteristics of p Zeolite Praphulla N. Joshi, Eapen M. Joseph and Vasudeo P. Shiralkar* Catalysis Division, National Chemical Laboratory, Pune 4 10 008, India The sorption characteristics of Al-B (SiO, : AI,O, = 39)and Fe-/I (SiO, : Fe,O, = 39-46) zeolites in both sodium and protonic forms have been compared. Sorption capacities at PIP, = 0.8 and 298 K for water, hexane, cyclo-hexane and butylamine (BA) yielded estimates of the void volume of the fi zeolites. Sorption kinetics with these probe molecules also confirmed the highly crystalline nature of the samples. BA sorption isotherms in both the zeolites in the temperature range 33H83 K have also been measured to characterize the acidic centres in the zeolites. Analysis of the BA sorption data in terms of different sorption models yielded useful information on the nature of sorbent surface and sorption centres. Sorption selectivities were also compared by evaluation of the chemical potential of the sorption at isotherm temperatures.lsosteric heats (QSt)of BA sorption were evalu-ated by the application of a thermodynamic approach to the sorption energetics. Over the entire coverage Q,, values varied within 40-60 kJ mol-' for HIA1-P and within 18-40 kJ mol-' for H/Fe-B zeolites. High-silica zeolites are important and attractive catalysts by virtue of their hydrophobicity and potential, thermal, hydro- thermal and acidic properties with resistance to coke forma- tion.' /? Zeolite is a wide pore, high-silica, crystalline aluminosilicate first synthesised by Wadlinger et al.' Although the synthesis, characterization, structure and cata- lytic properties of B zeolite are already documented,2-" the sorption properties have not been discussed in detail./3 Zeolite is the only high-silica zeolite to have a three-dimensional, 12-ring pore system. It also has a near-random degree of stacking faults whilst maintaining full sorption capacity. Sorption studies are often of primary importance in characterizing zeolite channels and pore openings. The extent of surface heterogeneity and changes in the physico-chemical properties are also studied by sorption measurements with probe molecules such as water, hexane and cyclohexane.The evaluation of thermodynamic parameters from sorption iso- therms of basic molecules such as butylamine (BA) helps to characterize the acidic nature of the zeolite catalysts. Varia- tions in the Si : A1 ratios12 and the extra-framework cationsI3 are reported to influence the sorption properties. In view of these aspects, studies on the sorption properties +of /Izeolite with framework A13 and Fe3 + have been carried out. Experimental Materials Al-B zeolite was synthesised following the method reported earlier3." and Fe-#l zeolite was prepared following a pro- cedure of Kumar et al." Both zeolites were calcined carefully around 773 K for 8 h to drive off templating species. The zeolites in the sodium form were then ion exchanged with 1 mol dm-3 ammonium nitrate (10 cm3 g-' of zeolite) solution at 368 K for 4 h.The ion-exchange procedure was repeated till the resultant solid contained <200 ppm sodium. The ammonium forms were then deammoniated at 723 K for 10 h to give the protonic form of the zeolites. AnalaR Grade (purity > 99.9%) butylamine, cyclohexane and hexane, dried over 3A molecular sieve extrudates, were used for the sorp- tion measurements. Doubly distilled water was used for the sorption measurements and for washing of ion-exchanged products. Methods Powder X-ray diffraction patterns were obtained using high- purity Si powder as an internal standard, on a Rigaku D Max/III VC diffractometer with nickel-filtered Cu-Ka radi- ation.The morphology and crystallite size were examined on a scanning electron microscope (model Cambridge, Stereoscan- 150 UK). The chemical analysis was performed by a combination of wet chemical and atomic absorption (Hitachi-2 8000, Japan) and inductively coupled plasma emission (Jobin Yuon-JY-38 VHR) spectroscopic methods. The zeolites were futher characterized by thermal analysis (NETZSCH, model STA-490), FTIR (Nicolet 60 SXB) and EPR (Brucker E 2000) spectroscopies. Solid-state magic-angle spinning (MAS) NMR for 29Si and 27Al were recorded at 295 K using a Brucker MSL-300 spectrometer. While acquiring 29Si spectra, a recycle time of 3 s was used with MAS at 3.5 KHz. TMS and aqueous AlC1, were used as references for 29Si and 7Al spectra.The magnetic susceptibility measure- ments, in the temperature range 94-297 K were carried out using a Faraday balance (Cahn-Ventron, USA). Sorption studies (both the kinetics and isotherms) were carried out in an all-glass, McBain-Baker type gravimetric vacuum unit described elsewhere. Results and Discussion Characterization Fig. 1 illustrates the powder X-ray diffraction (XRD) patterns of HIA1-P and H/Fe-P zeolites. The identification and the purity of these phases were examined using Si powder as an internal standard. The XRD pattern of the H/Al-#l sample closely matches the reported data.5 Although XRD patterns of both HIA1-B and H/Fe-fi are nearly identical, the intensity of the prominent peaks of H/Fe-P was found to be lower with a small shift of the peak towards lower 28 values.This may be due to the lattice expansion on account of the larger Fe3+ atoms in the framework positions. The absence of impurity peaks and amorphous material indicates the highly crys- talline nature of both the samples. Fig. 2 shows the frame- work IR spectra of H/Al-P and H/Fe-/I samples. The bands due to asymmetric stretching vibrations of Si-0-A1 at 1174 and 1074 cm-' in H/Al-B were found to be shifted to 1168 and 1059 cm- ' in the H/Fe-#l sample. The IR profiles of both the samples and the shift observed in the H/Fe-#l sample J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 9u .* 38 30 22 14 2tl/degrees Fig. 1 X-Ray powder diffraction patterns of (a) HIA1-B and (b) H/Fe-B zeolites (*Si as an internal standard) support the pure and crystalline nature of the A1 and Fe"' silicate /3 phases.Uniform ellipsoid-shaped crystallites of 0.4-0.6 pm and 0.5-0.75 pm is size were observed for HIA1-P and H/Fe-fl respectively. The SEM photographs (Fig. 3) also con- firmed the high purity and crystallinity of the /3 zeolite phases. A signal at g = 4.4in the EPR spectra and the data on magnetic susceptibility measurements support the pres- ence of Fe3+ in a tetrahedral coordination with no inter- action of Fe-0-Fe. Nuclear electron coupling of Si-0-Fe was revealed by signal broadening in the 29Si solid-state MASNMR spectra. On the basis of the above observations, it is clear" that the /3 zeolite samples are of high crystallinity without significant impurity or amorphous matter and that isomorphous replacement of Fe3+ in the framework has been achieved.The unit cell compositions cal- culated from the chemical analysis of the products are tabu- lated in Table 1. Table 1 Unit cell compositions of B zeolites ~ sample unit cell composition uc g-'/1020 n 1 I 1200 1000 800 600 ' wavenumber/cm -Framework IR spectra of (a) HIA1-b and (b)H/Fe-B zeolites Fig. 3 SEM photographs of (a) HIA1-B and (b)H/Fe-B zeolites Sorption Properties Equilibrium sorption capacities at PIP, = 0.8 at 298 K for different probe molecules in both H/Al-/3 and H/Fe-P are summarized in Table 2. The salient features of these results are as follows.Equilibrium sorption capacities for all four probe molecules in the protonic forms of the zeolites are mar- ginally higher than those in the sodium form, possibly because of the smaller size of the proton. However, the increase is larger in case of the Fe-p sample. During ammon- ium exchange, possibly Fe3 + species acting as charge bal- ancing extraframework cations were leached out and hence more void made available for sorption in H/Fe-/3. The enhanced Si : Fe for H/Fe-p (Table 1) supports these findings. Equilibrium sorption capacities usually depend on the size of the probe molecule, the zeolitic void volumes, the geometry of the cages and the probe molecule, the zeolitic void volumes, the geometry of the cages and the packing geometry/eficiency.Polar water molecules, being sufficiently smaller in size (kinetic diameter 2.65 A), penetrate almost all the cages in the zeolitic lattice and assume close packing in an attempt to interact with extra-framework cations which balance the charge on alumina tetrahedra. The equilibrium sorption capacity for water is usually regarded as an indica- Table 2 Sorption properties" of B zeolites Al-/!? Fe-#I sorbate Na form H form Na form H form hexane cyclohexane water BA 8.60 (0.28)b 10.31 (0.28) 59.51 (0.27) 11.09 (0.28) 8.61 (0.28) 10.34 (0.29) 59.83 (0.28) 11.12 (0.29) 8.30 (0.26) 8.83 (0.24) 51.90 (0.24) 9.94 (0.25) 8.82 (0.28) 10.60 (0.29) 61.36 (0.28) 10.33 (0.26) Expressed in molecules uc-',at 298 K, PIPo = 0.8.Values in par- entheses represent pore volume in cm3 g-'. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 ,.-I, 0 2 4 6 8 10 12 14 16 2 4 6 8 10 12 14 16 Pporr Pporr Fig. 4 BA sorption isotherms for A, HIAI-0 and B, H/Fe-P zeolites at (a) 333, (b)363, (c) 393, (6)423, (e)453 and (f)483 K tion of the hydrophilic/hydrophobic character of the zeolite framework. Water uptake is ca. 59-61 molecules uc-' for H/Al-/? and H/Fe-/?. Uptake of BA, the next largest probe molecule, is ca. 10.33-11.12 molecules UC-I in the protonic forms of the zeolites. Cyclohexane (kinetic diameter = 6.0 A) being larger than BA, has a slightly lower uptake (10.3-10.6 molecules uc- ') than BA. The uptake of hexane (8.6-8.8 mol-ecules uc-') is rather lower than that of the BA, which may be due to the difference in the packing efficiency of the cylin- drical hexane (kinetic diameter = 4.3 A) and the almost spherical cyclohexane molecule.Table 2 shows that, with the exception of hexane, the uptake of probe molecules follow the order of the reciprocals of their molecular size in zeolites of comparable pore volumes. BA Sorption Isotherms Fig. 4 shows families of typical sorption isotherms in the tem- perature range 333-483 K for Al-/? and Fe-B zeolites. The shape of the isotherms in both the zeolites was found to be similar to Type I (Langmuir type) according to Ki~elev's'~ classification. It can be seen from the nature of the isotherms, that ca. 80% of the total sorption takes place over a low- pressure range (up to 4 Torr).The sorption isotherms clearly indicate that H/Fe-P has a lower equilibrium sorption capa- bility (molecules uc- ') than H/Al-/? over the entire isotherm temperature range. This may be due to the strong interaction between basic BA molecules with the sorption centres of higher acidic strength (usually hydroxyl) groups leading to solvation and volume-filling phenomena. The different extent of uptake may partly be explained on the basis of dif-ference in the charge on the framework oxygen i.e. by the basic character and by the degree of heterogeneity of the charge distribution. In addition to this, the distortion of the framework structure as indicated by the XRD and the com- paratively less non-specific orientation, may be hindering the rearrangement of BA molecules necessary to achieve satura- tion capacity at a particular temperature.This may also, in part, be responsible for the slower uptake and reduced equi- librium saturation capacities for BA sorption in the H/Fe-fi sample. Therefore, in spite of the nearly equal pore volume of the two samples, the saturation capacity of BA varies con- siderably. Applications of Isotherm Equations Analysis of the sorption data in terms of different isotherm models usually yields useful information about the nature of the sorption centres. It was therefore thought appropriate to investigate the influence of the nature of the framework cations on the applicability of various isotherm equations to the BA sorption data obtained in the present study on HIA1-P and H/Fe-/? zeolites.Langmuir Isotherm Equation The Langmuir isotherm model describes sorption equilibrium in a system wherein all the sorption centres are of the same energy and the sorbate molecule is localized on a sorption centre. Fig. 5 shows typical Langmuir plots for BA sorption in HIA1-B and H/Fe-P zeolites. It can be seen that both the zeolites exhibit excellent linear plots with different intercepts. Thus it is evident that the Langmuir sorption model is applicable to sorption in these samples. Similarly the Lang- muir sorption equation was found to represent satisfacto-rily BA sorption data in EU-1 zeolites. Accordingly all the sorption centres appear to have the same sorption potential. The monolayer capacities obtained from the reciprocals of the slopes of these linear plots are tabulated in Table 3.These values of the saturation capacities are in good agreement Table 3 Comparision of saturation capacities of /3 zeolites saturation capacities (molecules uc- ') temperat ure/K method HIA1-B H1Fe-B 333 experimental Langmuir 9.58 9.61 8.47 8.80 9.60 8.75 363 experimental LangmuiriBET 8.80 9.09 8.98 7.55 7.89 7.85 393 experimental Langmuir 8.07 8.33 8.22 6.66 7.06 6.97 423 experiment a1 Langmuir 7.08 7.40 5.85 6.27 7.30 6.23 453 experimental Langmuir 6.45 6.89 5.20 5.57 6.79 5.54 483 experimental Langmuir 6.04 6.45 6.33 4.68 5.28 5.03 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 I 1 L I 0 5 10 15 PfTorr P/Torr Fig. 5 Langmuir plots for BA sorption in A, H/Al-P and B, H/Fe-B zeolites at (a) 333, (b)363, (c) 393, (6)423, (e)453 and (f) 483 K. a is the amount sorbed in molecules (uc)- '. with those obtained experimentally. These data also show the higher sorption capacity of H/Al-fi than that of H/Fe-fl. It was also suggested16 that BA sorption is localized and a basic molecule such as BA shows strong interaction with acidic protons and the other acidic species. In the present studies, because of localized sorption of BA and interaction with acidic protons in p zeolites, the Langmuir isotherm equation yields linear plots and represents BA sorption data satisfactorily.The lower saturation capacity in H/Fe-p indi- cates the moderate interaction of BA with the acidic protons suggesting the lower acidity of protons bridged to Fe3+ through oxygen. Another salient feature of these Langmuir plots is the extent of decrease in the value of the intercept on the y-axis with temperature. It was observed that the inter- cept on the y-axis of HIAl-fl is lower than that of H/Fe-p for A -120 IIIIIII all temperatures. The intercept on the y-axis is usually related to the strength of sorption and as the intercept decreases, stronger interaction is involved in the sorption process. This is reflected in the higher acidic strength of H/AI-@ than that of HIFe-B. BET Isotherm Equation Linear plots were obtained by applying the BET equation to the sorption data of BA in HIA1-P and H/Fe-P zeolites, Fig.6. BET plots of H/Fe-/l yield higher values of the intercept on the y-axis indicating a lower value of C, i.e. the heat of sorp- tion of the first layer of sorbate on the sorbent surface. Monolayer capacities calculated from the slopes and the intercepts of these plots (shown in Table 3) are in good agree- ment with those obtained experimentally. The linearity of the r B 40 30 c-I pp 20 v (D 10 0 4 8 12 16 20 (pipo)x 103 Fig. 6 BET plots for BA sorption in A, H/Al-8 and B, H/Fe-/3 zeolites at (a) 333, (b)363, (c)393, (d)423, (e) 453 and (f)483 K J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 plots indicates that, the BA sorption data for B zeolite can be described by the BET sorption model.The lower value of heat of BA sorption in H/Fe-/? also suggests moderate inter- action of acid centres with BA molecules. Dubinin Isotherm Equation An attempt is made here to apply Polanyi’s potential theory modified by Dubinin and Radushkevich17 for BA sorption in HIA1-P and H/Fe-B in the temperature range 333-483 K. The Dubinin-Rudushkevich equation is expressed as, B log w = log w,--CT 1% P/P,lZ2.303 /?’ where W is the amount sorbed at equilibrium pressure P, W, is the total sorption capacity, B is the constant independent of temperature and characteristic of sorbent pore structure and fi is the affinity coefficient. When BA sorption data were fitted to the above equation, reasonably linear plots were obtained at all temperatures. Typical Dubinin plots are shown in Fig.7 and indicate that the BA sorption data in P zeolite could be satisfactorily expressed by the Dubinin- Rudushkevich equation. The slopes of these plots were found to increase with increase in temperature. The values of satu- ration capacities and B//?’ obtained from the intercepts on the y-axis and the slopes respectively are summarized in 0 1 2 3 4 5 6 (log 7 1.0 I B I .. 0.5L I 1 I 1 I 0 1 2 3 4 5 6 (log p,lp)2 Fig. 7 Dubinin plots for BA sorption in A, HJAl-8 and B, HJFe-fl zeolites at (a)333, (b)363, (c)393, (6)423, (e)453 and (f)483 K 39 1 Table 4. These saturation capacities are in close agreement with those obtained experimentally.This shows that the BA sorption in fl zeolites nearly follows the Polanyi potential theory of volume filling. Since B is independent of tem-perature, the fact that B/B2 increases with increase in tem- perature and at constant temperature is higher for H/Fe-P than HIA1-B confirms the high affinity coefficient, B, for HIA1-B for a particular set of temperatures. Thus, the varia- tion in the affinity coeficient at constant temperature can be used as a means of confirming the isomorphous framework substitution and strength of acidic sites. Sips Isotherm Equation Sips’8 suggested a new theoretical absorption model to cal- culate the distribution of adsorption energies of the sites of a sorbent surface, when sorption isotherms were known and sorption was localized without sorbate-sorbate interaction. The original Sips eq~ation’~ on linearization takes the form: 0log -= log A + c log P1-8 where A and c are constants and P is the equilibrium pres- sure at coverage 8.For calculating 8, the saturation capacities were obtained from Langmuir plots. In order to check the applicability of Sips equation, the experimental values of 8 and P were substituted in the above equation. The salient feature of the Sips plots obtained in the present study is that BA sorption data only yielded linear plots for H/Fe-B, as shown in Fig. 8. The magnitude of c (obtained from the slopes) deviates from unity in the lower temperature region up to 393 K, suggesting deviation from the Langmuir approach in the same region.However, the value c was nearly constant, 1.0 0.04, in the temperature 1.4 1.2 1.o 0.8 m I 0.6 &-a-0 0.4 0.2 0.0 -0.2 I L I I , I -1.0 -0.8 -0.6 -0.4 4.2 0 0.2 log P Fig. 8 Sips plot for BA sorption in H/Fe-8 zeolite at (a) 333, (b)363, (c)393, (6)423, (e)453 and (f)483 K Table 4 Saturation capacities and BIB2 obtained from Dubinin plots for B-zeolites HjAl-f3 H/Fe-f3 temperature saturation capacity saturation capacity /K B/,P x 107 /molecules uc- BJ~x 107 /molecules uc - 333 1.13 9.60 2.34 9.12 363 1.19 9.22 2.50 8.12 393 1.28 8.17 2.67 7.48 423 1.61 7.20 2.84 6.76 453 2.10 6.64 3.10 6.23 483 3.39 6.20 3.81 6.02 392 J.CHEM. SOC. FARADAY TRANS., 1994,VOL. 90 equation to the BA sorption data. Deviation from linearity was observed at higher pressures and at higher temperatures for both the samples. The extent of deviation from linearity for HIA1-/3 was less than for H/Fe-fl at the same temperature. BA sorption data in Fe3+-exchanged type Y16 and, in a dif- ferent cationic form, LTL' zeolites were also satisfactorily represented by linear plots, but the Freundlich sorption model failed to represent BA sorption in EU-1" and in titan- V-0.4L ' ' ' ' ' ' ' ' ' 1 osilicates with MFI structure. 2o 0.1 0.3 0.5 0.7 0.9 1.1 log P Chemical Affinity and the Selectivity of the Sorbed Phase 1.0 -B A reversible and isothermal transformation of a gas, at a standard pressure Po (760 Torr) into an infinite amount of 7h -*_--u--/-sorbent-sorbate mixture under equilibrium pressure, P, decreases the chemical potential.The chemical affinity, when the non-ideality of the sorbate is neglected may be expressed2' as : Ap = RT ln(P/P,) The value of -Ap may be taken as the quantitative measure of the chemical affinity of the sorbate for the sorbent. The ,/f 1 I 1-range 423-483 K. Therefore, complicating factors such as irreversibility in the BA sorption, sorbate-sorbate inter-actions etc., may be operative in the low temperature region in H/Fe-P. However, Sips equation was not found applicable to the case of HIAl-/3 suggesting stronger sorbate-sorbate interactions through localized sorption.Freundlich Isotherm Equation Analysis of BA sorption data in H/Al-fl and HIFe-8 zeolites in terms of the Freundlich isotherm model yielded linear plots in the higher-pressure region in the temperature range 333-483 K. Typical Freundlich plots for BA sorption in both samples are shown in Fig. 9. The excellent linearity of these plots confirms the applicability of the Freundlich isotherm 6.5 I* BA(rnolecu1es uc-') plots of -Ap against the amount sorbed also serves as useful criteria for the comparison of the sorption affinities of a probe molecule in the lattices of zeolites with different tetra- hedral cations. Typical chemical affinity plots for BA sorp-tion in H/Al-P and H/Fe-#l are shown in Fig.10. It can be clearly seen from the figure, that the decrease in -Ap is sharper with the increase in sorption temperature. H/Al-P shows higher sorption affinity than H/Fe-B and the decrease in the sorption affinity is rather gradual in HIA1-P while it is comparatively rapid in HIFe-P. In other words, the chemical affinity for BA sorption in HIA1-B is higher than that in H/Fe-/3 over the entire coverage over the isotherm tem-perature range 333-483 K. Isosteric Heats of Sorption (Q,,) The isosteric heat of sorption is derived from the sorption isosters by applying the Clausius-Clapeyron equation at con- stant sorbate loading using the equation: If Q,, is independent of temperature, the plots of In (P)vs. 1/T 7.3, 1 IB 6.51 I 5.7 .-I-z 4.9 Y,63 I 4.1 3.3 2 4 6 8 BA(rnolecu1esuc-') Fig.10 Chemical affinity curves for BA sorption in A, HIA1-P and B, H/Fe-fi zeolites at (a)333,(6)363,(c) 393,(d)423,(e)453 and (f)483 K J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 -32 -56 c -I -28-52 3 -'wam 24 -48 -20 -44 -16 -40 123456 BA(molecu1es uc-') Fig. 11 Profiles of isosteric heats of the BA sorption for A, H/Al-P and 0,H/Fe-/3 zeolites should be linear. In the present study, isosters for BA sorp- tion in H/Al-p and H/Fe-P were found to be linear. The iso- steric heat was calculated using the slopes of these isosters. Fig. 11 depicts the isosteric heats of the BA sorption in H/Al-p and H/Fe-P zeolites. H/Al-fl zeolite exhibits Q,, values ranging from 40-63 kJ mol-' which is significantly higher than that (18-32 kJ mol-') for H/Fe-/?.For HIA1-P in the lower-coverage region Q,, increases slowly, passes through a maximum, decreases rather rapidly and then exhibits humps in the higher-coverage region. In the lower-coverage region, sorbate-sorbent interactions are usually predominant fol-lowed by both the sorbate-sorbent and sorbate-sorbate interactions in the mid-coverage region. In the higher-coverage region only sorbate-sorbate interactions are oper- ative. For H/Fe-P Q,, also increases initially and then decreases slowly in the mid-coverage region, passes through a minimum and then increases considerably in the higher-coverage region. Both the zeolitic surfaces display heter-ogeneous character towards BA sorption.In the lower-coverage region, HIA1-P exhibits higher values of Q,, than those of H/Fe-P, suggesting the higher acid strength of H/Al-P zeolite than H/Fe-p zeolite. Conclusions Sorption kinetics and equilibrium sorption capacities were found to be in accordance with the molecular sizes of the probe molecules. The void volume, evaluated from equi- librium sorption capacities of different probe molecules, was found to be almost identical for Al-P (0.27-0.29 cm3 g-') and for Fe-p (0.24-0.26 cm3 g- ') in sodium form. Void volumes increased only marginally for /?-A1 but increased notably for p-Fe on their conversion from the sodium to the protonic form. Of the total Fe3+ in the solid, a small fraction, present as charge-balancing cations, seemed to be leached out during conversion from sodium to protonic form via ammonium exchange.Sorption capacitites also revealed a difference in packing efficiency in zeolite voids of hexane and cyclohexane because of their shape differences. BA sorption isotherm data were satisfactorily represented by BET, Dubinin and Langmuir approaches. The Sips equation could represent BA sorption in H/Fe-P zeolite whereas, the Freundlich sorption model was found to be applicable at higher pressures only. H/Al-P exhibits higher sorption selectivity than H/Fe-b for BA sorption. The isosteric heat ((Is,)variation with the cover- age revealed the heterogeneous character of both the sorbent surfaces.Higher isosteric heat for HIA1-p than H/Fe-P is an indication of the comparatively higher strength of acid centres in the former. References 1 J. Schemer, Catalytic Materials: Relationship between Structure and Reactivity, ed. R. T. K. Baker, E. G. Derouane, R. A. Dalla Betta and T. E. Whyte, jun., American Chemical Society, Wash- ington, DC, 1984, p. 157. 2 R. L. Wadlinger, G. T. Kerr and E. J. Rosinski, US Pat. 3 308 069, 1967. 3 M. M. J. Tracy and J. M. Newsam, Nature (London), 1988, 352, 249. 4 J. B. Higgins, R. B. La-Bierre, J. L. Schlenker, A. C. Rohrman, J. D. Wood, G. T. Kerr and W. J. Rohrbaugh, Zeolites, 1988, 8, 446. 5 J. Perez-Pariente, J. A. Martens and P. A. Jacobs, Zeolites, 1988, 8,46. 6 J. A. Martens, M.Tielen, P. A. Jacobs and J. Weitkamp, Zeo-lites, 1984, 4, 98. 7 J. A. Martens and P. A. Jacobs, Zeolites, 1986,6, 334. 8 J. Perez-Pariente, J. A. Martens and P. A. Jacobs, Appl. Catal., 1987, 31, 35. 9 R. N. Bhat and R. Kumar, J. Chem. Technol. Biotechnol., 1990, 48,453. 10 R. Kumar, A. Thangraj, R. N. Bhat and P. Ratnasamy, Zeolites, 1990, 10, 85. 11 S. G. Hegde, R. Kumar, R. N. Bhat and P. Ratnasamy, Zeolites, 1989,9, 231. 12 V. P. Shiralkar and S. B. Kulkarni, Z. Phys. Chem. (Leipzig), 1984,262,3 13. 13 P. N. Joshi and V. P. Shiralkar, J. Phys. Chem., 1993,97,619. 14 A. V. Kiselev, Discuss.Faraday SOC.,1965, 40,205. 15 G. N. Rao, P. N. Joshi, A. N. Kotasthane and V. P. Shiralkar, J. Phys. Chern., 1990,94,8589. 16 S. J. Kulkarni and S. B. Kulkarni, Indian J. Chern. Sect. A., 1989, 28,6. 17 M. M. Dubinin, L. V. Radushkevich, Proc. Acud. Sci., USSR, 1974, 55, 327. 18 B. Coughlan and P. M. Larkin, Proc. R. Irish Akad. B, Centen-ary Issue, 1977,77, 383. 19 R. Sips, J. Chem. Phys., 1948, 16,491. 20 S. P. Mirajkar, A. Thangraj and V. P. Shiralkar, J. Phys. Chern., 1992,%, 3073. 21 V. P. Shiralkar and S. B. Kulkarni, Zeolites, 1985,5, 37. Paper 3/03635I; Received 24th June, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000387

出版商:RSC    

年代:1994

数据来源: RSC