摘要:

J. Chem. SOC., Furaduy Trans. I , 1989, 85(4), 855-867 Effect of Form on the Surface Reactivity of Differently Prepared Zinc Oxides Vera Bolis, Bice Fubini" and Elio Giamello Dipartimento di Chimica Inorganica, Chimica Fisica e Chimica dei Materiali, Universita' di Torino, Via P. Giuria 7, 10125 Torino, Italy Armin Reller Anorganisch-chemisches Institut, Uniuersitat Ziirich, Winterthurerstrasse 190, 8057 Zurich, Switzerland The relationships between the form of crystallites and their adsorption properties have been investigated on ZnO samples of various origins by means of electron microscopy, X-ray diffraction and adsorption calorimetry. Three polycrystalline ZnO samples have been studied: one obtained by ignition of zinc metal (Kadox) and the others obtained by decomposition of zinc carbonate and oxalate.The higher surface reactivity found on Kadox ZnO in comparison with that found on ex-salt ZnO is not due to the preferential development of crystal planes at the surface but to the presence of better defined single microcrystals with sharp edges. In all the cases examined the preparation route leads to a difference in morphology and consequently a difference in reactivity. Carbon monoxide and hydrogen have been used as surface probes for active sites : CO is coordinated onto cations exposed at the edges between the (0001) and (1010) planes; H, is adsorbed in different forms, one of which, type I, occurs on sites located at the same edge as CO. Prereduction of ZnO reduces the adsorption activity towards both gases; the extent of this reduction also depends upon the actual morphology.The reactivity of a solid in general, including surface reactivity in catalysis, depends decisively on its form, which comprises compositional, structural and morphological features. This is particularly relevant in the surface chemistry of zinc oxide. On the one hand, studies of single crystals have shown a dependence of its adsorption properties upon the exposed surface on the other hand, differences in adsorption properties have been found on polycrystalline materials of various origins.*-' In the latter case differences in adsorption capacity and interaction energy have been ascribed by some authors, including ourselves, t o different distributions of the exposed crystal faces on the samples, resulting from the preparation route Since zinc oxide is an active component in mixed catalysts, such as Cu/ZnO and Cr,O,ZnO,lO, l1 for various reactions, e.g.the synthesis of methanol and higher alcohols, the influence of the preparation route on the final state of the ZnO surface is of essential importance in catalyst preparation. Zinc oxide crystallizes in a wurtzite-type structure and is usually made up of hexagonal prisms, where the hexagonal planes, (0001) and (OOOT), are located perpendicular to the c axis and the prismatic planes, (1010) and (1 120), are parallel to it. These low-index planes are the most commonly exposed faces in polycrystalline material^.^ On the polar faces, (0001) and (OOOT), Zn2+ cations and 02- anions, respectively, are the more outwardly exposed.On the non-polar prismatic planes, (1010) and (1 120), both zinc and oxygen ions are located in the same plane. The polar faces, where surface reconstruction produces particular anion/cation vacancies, have been 855856 Surface Reactivity of Zinc Oxides regarded by several authors as the active sites in catalytic reaction^.^. ' 9 8~ l1 In fact, at room temperature both CO and H, are selectively adsorbed on this type of face.2*12-15 A clear-cut relationship between the extension of polar faces, the adsorption capacity towards CO and H, and catalytic activity has thus been but never proved, on the basis of parallel investigations on morphology and surface properties. Lavalley et al.,' however, recently reported that kinks and edges (at the intersection of polar and non-polar planes) should be taken into account as possible active sites instead of the polar planes.Moreover, the actual shape of the microcrystalline particle (i.e. the microstructure of the crystal faces present and thus of the edges formed) can be an essential cause of the existence of specific sites. Therefore, in addition to the crystal morphology (i.e. the type and relative abundance of the various exposed faces), the abundance of ' morphological ' defects such as surface steps, edges, kinks and corners, which cause the presence of ions that are in particularly low coordination at the surface and not available on extended faces, must be taken into account. In fact, the number of these defects (negligible in the case of single crystals) can be very high in finely divided polycrystalline materials and thus can influence the reactivity of the solid by introducing particularly active sites.Conditions used in the preparation route could play an important role in this respect. The aim of this paper is to compare directly adsorption features (the number of active sites and the binding energy) and morphological aspects [as elucidated using X-ray diffraction (X.r.d.), scanning electron microscopy (SEM), transmission electron microscopy (TEM) and high-resolution electron microscopy (HREM)] of ZnO prepared by different procedures : the widely studied Kadox ZnO, obtained by the ignition of Zn metal, and samples obtained by the decomposition of salts (carbonate and oxalate), a procedure which is close to that adopted in the preparation of mixed catalysts. Carbon monoxide and hydrogen have been chosen as test molecules in adsorption work for three main reasons.(i) The adsorption of these gases, at least on Kadox ZnO, has been studied widely by infrared spectroscopy,', ' 9 1 3 7 1 5 9 l6 adsorption microcalor- imetry4. 6 * and temperature-programmed de~orpti0n.l~. (ii) Both of these gases are selectively adsorbed at room temperature on only one (CO) and two (H,) kinds of sites, and assignments of the adsorbed form to a well defined surface arrangement have been proposed. Carbon monoxide is coordinated onto the coordinatively unsaturated Zn2+ cations (c.u.~),,' l2 whereas hydrogen is dissociatively adsorbed onto the so-called type I sitesl3'l7 yielding surface hydroxyls and hydrides.These sites are proposed to be located at the (0001) face14 or at the edge between the (0001) and (1010) faces." Hydrogen is also adsorbed onto the so-called type I1 sites, for which a bridged form requiring suitably oriented pairs of Zn2+ and 0,- ions has been pr~posed.~ In addition, previous calorimetric investigations by some of us indicated the possibility that some adsorbed hydrogen could slowly diffuse into the bulk.4 (iii) Hydrogen and carbon monoxide are the starting materials for syngas reactions. In addition to the morphological aspects, the semiconducting nature of ZnO (n-type) can also play a crucial role in its reactivity. Zinc donor centres in excess, and consequently the presence of a relevant electronic population in the conduction band, will obviously influence any interaction involving electron exchanges between the solid and the adsorbed m01ecule.'~ The effect of reducing pretreatments on adsorption has thus been investigated.A drastic change in reactivity towards CO upon reduction was found by some of us in a previous study of one ZnO sample.6 The study of the influence of reduction upon adsorption of both CO and H, has been extended to all the samples examined in order to verify whether reduction affects samples obtained by different preparation routes in the same way.V. Bolis, B. Fubini, E. Giamello and A . Reller 857 Experimental Three different ZnO samples have been investigated: zinc oxide obtained from the ignition of the metal in an oxidising atmosphere (Kadox 25, New Jersey Zinc Co.), zinc oxide obtained by decomposition of zinc carbonate and zinc oxide obtained by the slow decomposition of zinc oxalate.The B.E.T. surface area of the samples (measured at 77 K by means of a Carlo Erba Sorptomatic apparatus, N, iunit) were 10, 24 and 18 m2 g-l, respectively. The heat of adsorption of CO and H, and related quantities were measured at room temperature by means of a Tian Calvet microcalorimeter connected to a volumetric apparatus as described previously.20 The samples were pretreated following three different procedures. (i) The sample was slowly heated at 673 K under vacuum and then 75 Torr of 0, (1 Torr z 0.133 kPa) was admitted and removed three times at the same temperature. The sample was then cooled and transferred into the calorimeter under oxygen.(ii) The sample was slowly heated up to 673 K in vacuo and then transferred, in vacuo, into the calorimeter. (iii) The sample was slowly heated to 673 K, then cooled to 473 K. At this temperature 75 Torr of H, was admitted and removed three times. The sample was then evacuated at the same temperature (473 K) in order to ensure the elimination of any adsorbed hydrogen irreversibly held at the surface at lower temperatures. The origin of the ZnO sample will be designated as k-, c- or 0-, indicating, respectively, Kadox, ex-carbonate or ex-oxalate. The type of pretreatment undergone by the sample will be indicated by the terms -ox (oxidised), -un (untreated) or -red (reduced), following the ZnO forrnula, e.g. k-ZnO-red indicates Kadox zinc oxide submitted to the reducing pretreatment.Adsorption was performed by admission of successive small doses of the adsorptive (CO or H,) up to an equilibrium pressure of 75 Torr in the case of CO and 25-30 Torr in the case of H, (first run). The calorimetric cell was completely evacuated after the first adsorption run in order to evaluate the reversibility of the process. Furthermore, in some cases a second adsorption run has been performed. Owing to the small size of the calorimetric vessel (containing at maximum 3-4 g of powder) and the relatively low surface area of the samples examined, it was difficult to perform accurate volumetric measurements, particularly in the case of the ex-oxalate zinc oxide, which has a poor adsorption capacity. For this reason data relative to the adsorption of H, onto the reduced ex-oxalate zinc oxide are not available. The three different zinc oxides were characterized with respect to their structural, compositional and morphological features.X-Ray diffractometry was carried out using a Guinier-IV camera as well as a Siemens Kristalloflex diffractometer, both with Cu Ka radiation. In order to confirm the purity of the samples, i.e. the absence of zinc carbonate and, respectively, zinc oxalate, thermogravimetric measurements were performed using a Perkin-Elmer TGS-2 thermomicrobalance. For a detailed analysis of both the morphology and microstructure, including the identification of the cry- stallographic faces constituting the crystallite surfaces, scanning electron microscopy (SEM) using a Cambridge Instruments Stereoscan mk 111 instrument as well as high- resolution electron microscopy (HREM) and selected-area electron diffraction (SAED) using a Jeol 200CX instrument equipped with top-entry stage were applied.For the HREM studies the samples were dispersed in hexane and disposed on copper grids coated with a ‘holey’ carbon film. In order to obtain diffraction patterns of isolated ZnO crystallites, optical diffraction studies were performed by laser-beam diffraction of high- resolution electron micrographs (fringe patterns) on an optical bench with standardized geometry.858 Surface Reactivity of Zinc Oxides Results Surface Morphology The initial thermogravimetric measurements of all three samples gave no evidence for a weight change of > 0.3% within the temperature range 300-900 K.Therefore the presence of traces of zinc carbonate or zinc oxalate can be excluded. In X-ray powder diffractograms, as well as in the patterns obtained from the Guinier IV camera, one observes slight differences between the samples. Whereas for c-ZnO and o-ZnO a slight broadening of the reflections is observed, sharp reflections are registered for k-ZnO. These results can be explained by the presence of smaller crystallites making up c-ZnO and o-ZnO, compared with k-ZnO, or by the presence of poorly crystalline materials. In addition, no indication can be derived from X-ray studies for a possible prevalence of a particular type of morphology. Thus no unambiguous identification of crystallographic faces making up the surfaces of the crystallites could be obtained.A more informative answer, however, could only be achieved from electron-microscopy studies. As is shown in the scanning electron micrographs [plate 1 (a)-(c)] the micromorphology of the three materials can be characterized as follows : for c-ZnO, agglomerates of crystallites with dimensions 4 1 pm are observed. In the case of o-ZnO rod-like agglomerates made up of densely packed crystallites are observed. The characteristic shape of these agglomerates can be described as pseudomorphs of the precursor material, i.e. the zinc oxalate crystals. For k-ZnO the crystallites appear to be randomly agglomerated. As a consequence this material can be easily dispersed. In summary, the scanning electron micrographs give evidence for the presence of very small ZnO crystallites constituting the three samples.Owing to the limited resolution of SEM, however, the exact shape of these crystallites cannot be defined. Therefore, the high-resolution micrographs presented in plate 2(a)-(c) must be examined. First one should note that high-resolution electron microscopy cannot yield quantitative distributions of particle sizes or the exact numbers of crystallographic faces constituting the surfaces of the differently prepared samples. The analysis of the micrographs, however, yields information on the shape and dimensions, as well as on the crystallographic faces actually making up the surfaces of isolated particles. In plate 2 the characteristic features of the three ZnO samples are summarised.Plates 2(a) and (b) present an agglomerate of c-ZnO microcrystals and, as an inset, the high-resolution image of few isolated crystals. The dimension of these particles lies in the range 30-60 nm. The formation of partial large arrays of single- crystal material can be explained by the fact that, during the decomposition of the initial zinc carbonate, relatively high local partial pressures of CO, exist. As has been shown in the decomposition of various carbonates,2' such conditions allow sintering of the product phase, i.e. the zinc oxide. This result is supported by the fact that no sharp edges confining the crystallites are observed. The evaluation of the fringe patterns in the high- resolution micrographs obtained gives evidence that prismatic and hexagonal faces contribute to the surfaces.A predominance of prismatic faces can be derived from the inspection of larger arrays of this sample. The micrographs of o-ZnO [plates 2(c) and ( d ) ] support the abovementioned role of product formation governed by the course of the decomposition of the initial zinc oxalate: agglomerates made up of crystallites with equal dimensions in the range 30-60 nm exhibit the shape of the precursor crystals, i.e. they represent well defined pseudomorphs. Again the isolated crystallites of ZnO exhibit no well defined edges, but rather disc-like shapes. An analysis of the fringe patterns indicates that prismatic and hexagonal faces are present. The k-ZnO sample consists of perfectly formed single crystals with dimensions in the range 2G300 nm [plates 2(e)-(g)].This feature shows the most distinct difference between the samples c-ZnO and o-ZnO. The presence of such well defined single crystals with sharp edges must be a consequence of the basically different method of preparation,J . Chem. SOC., Faradaj) Trans. 1, Vol. 85, part 4 Plate 1 Plate 1. Scanning electron micrographs of microcrystalline ZnO, obtained from the decomposition of (a) zinc carbonate (c-ZnO) (b) zinc oxalate (0-ZnO) and by (c) the ignition of metallic zinc in an oxidising atmosphere (k-ZnO). V. Bolis et cil. (Fucing p . 8 5 8 )J . Chem. Soc., Faraday Trans. 1, Vol. 85, part 4 10 nm Plate 2(a)(b). For caption see overleaf. Plate 2 (a) (b) 100 nm V. Bolis et al.J . Chem. SOC., Furaduy Trans. 1 , Vol. 85, part 4 (C) 500 nrn - 30nm - Plate 2(c)(d).For caption see overleaf. Plate 2(c)(d) V. Bolis et al.J. Chem. SOC., Faraday Trans. 1, Vol. 85, part 4 (e 1 100 nm- 10 nm Plate 2. Electron micrographs of c-ZnO [(a) and (b)], o-ZnO [(c) and (d)] and k-ZnO [(eE(g)]. As insets, lattice images as well as respective optical diffraction patterns are presented for c-ZnO and k-ZnO in order to confirm the crystallographic faces making up the surfaces of crystallites. V. Bolis et al.V. Bolis, B. Fubini, E. Giamello and A . Reller 859 i.e. the combustion of finely dispersed metal. As one may see from the micrographs, the surfaces of these single-crystal particles are made up of prismatic and hexagonal faces. The pronounced predominance of hexagonal faces, which was regarded as a decisive feature in earlier 7 - 8 - l5 cannot be supported by these results.Adsorption of Carbon Monoxide Two runs of adsorption of CO were performed on the ZnO samples. During the first run a small fraction of CO ( < 10 YO) underwent oxidation to carbonate-like species and was irreversibly held at the surface after evacuation. The extent of this process was found to vary in the order o > c > k, and, obviously, ox > un > red. In the second run the adsorption of CO was entirely reversible with respect to simple 0-coordination of the molecule on the acidic cationic centres (Znf:,). Since the adsorption of CO was performed in order to evaluate the population of this type of site, it seems more convenient to compare the second runs obtained for the different samples.In fig. l(a>-(c) the volumetric isotherms of adsorption (adsorbed amounts, n,, us. equilibrium pressure) are reported for the three k-, c- and o-samples. For each sample three curves relative to the standard pretreatment (ox, un, red) are also shown. Inspection of the figures indicates that the number of surface sites interacting with CO depends on both the origin of the sample and the reducing treatment it has undergone. The amounts adsorbed per unit surface area decrease in the order k > c > o and, for each sample, in the order ox > un > red. In fig. 2 the integral heat (Qint) evolved during the adsorption is plotted as a function of adsorbed amount (n,) for all the samples. In all cases the integral heat depends linearly on the adsorbed amount, as all the experimental points fall on straight lines passing through the origin.Under such circumstances the adsorption enthalpy values, AadsH, are given directly by the slope of each line. In the case of the k- and o-samples all points fall on the same line independent of the redox treatment. Instead, three separate lines are found for the c-sample. However, the -AadsH value found for samples pretreated in an oxidising atmosphere is the same (46 kJ mol-') in all three cases. The adsorption data (adsorption enthalpy and adsorbed amounts) for all the samples examined are summarized in table 1. The adsorptive capacity of the k-ZnO-ox sample is two and four times higher than that of c-ZnO-ox and o-ZnO-ox, respectively. Thermal treatments in uacuo and in a reducing atmosphere cause a decrease in the number of active sites in all samples: on k- and o-ZnO the thermal treatment in vacuo eliminates 10 YO of the sites, the reducing one 40 Oh.c-ZnO is more sensitive to thermal treatments as it loses 60 and 70 % of the total sites, respectively. k-ZnO, even in the less-active form (red), exhibits a number of active sites that is higher than that of the other samples in whatever form. In the case of k- and o-ZnO, the energy of interaction is not affected by the reducing treatments (fig. 2), whereas in the case of c-ZnO (as previously reported by some of US)^ the energy of interaction decreases with progressive reduction. The value of - Aads H falls to 29 kJ mo1-' in the case of the untreated sample and to 8 kJ mol-' in the case of the reduced one [fig.2(b)]. Adsorption of Hydrogen Fig. 3 reports volumetric (n, vs. p ) (a) and calorimetric (Qint us. p ) (b) isotherms for the adsorption of H, onto k-, c- and o-ZnO in the oxidised form. The data refer to the first adsorption run. In all cases the process is pressure-dependent in the whole range examined. On k-ZnO, which also in this case shows the highest activity, at ca. 20 Torr the process860 1 . o 0.75 N 'E 4 \ g0.5 a? 0.25 Surface Reactivity of Zinc Oxides ~ I I I ( a ) 0 2'0 6'0 0 20 40 60 8 0.5 0.25 0 20 LO 60 plTorr Fig. 1. Volumetric isotherms of adsorption of CO on (a) k-ZnO (0, ox; @, un; 0, red), (6) c-ZnO (A, ox; A, un; A, red) and (c) o-ZnO (0, ox, 0, un; W, red). is not yet accomplished, whereas both c- and o-ZnO at the same pressure are nearly saturated.Adsorption values deduced from the volumetric isotherms [fig. 3 (a)] indicate that the amounts adsorbed per unit surface area decrease in the same order as that found for CO adsorption, i.e. k > c > 0. Table 2 reports the quantitative data of hydrogen adsorption ('p = 20 Torr) for the whole set of samples. As in some cases the heat of interaction decreases with coverage, the two limiting values of the variation range for the enthalpy are reported [columns 3(a)-(c)]. Two main facts can be observed by inspection of table 2. (i) Similar to the findings for carbon monoxide, the adsorption capacity of k-ZnO-ox is two and fourV . Bolis, B. Fubini, E. Giamello and A . Reller 86 1 3 N ‘E 5 c, - 2 \ E d 01 1 3 N ‘E h 1 0 0.2 0.4 0.6 Fig.2. Integral heat of adsorption us. adsorbed amounts of CO on ( a ) k-ZnO (0, ox; @, un; e, red), (6) c-ZnO (A, ox; A, un; A, red); and (c) o-ZnO (0, ox; 0, un; ., red). times higher than that of c- and o-ZnO, respectively (n,). Furthermore, the numbers of sites on the different samples decrease upon undergoing the reducing treatment. The relative decrease of active sites upon reduction (n,) is the same for the k-, c- and o-ZnO samples. (ii) The adsorption enthalpy decreases with coverage on k- and o-ZnO-ox samples (see also fig. 4), whereas in the case of c-ZnO-ox the differential heat is constant. Furthermore, in this latter case the heat of adsorption is unaffected by the reducing treatments [table 2, column 3(6)]. The same does not apply for k- and o-ZnO, for which the enthalpy range depends upon the treatment [table 2, columns 3 ( a ) and (h), respectively].More detailed information about the variation of the heat of adsorption862 Surface Reactivity of Zinc Oxides Table 1. Adsorption enthalpy and amount adsorbed measured a t p = 20 Torr for CO adsorbed on ZnO samples (a) k-ZnO (b) c-ZnO (c) o-ZnO _ _ _ ~ co/ co/ co/ molecule - AadsH/ molecule - AadsH/ molecule -‘adsH/ nm-, nlU/nzb kJ mol-1 nm-2 nla/nZb kJ mol-l nm-2 n,a/n2b kJ mol-1 ox 0.22 1/1 46 0.11 0.5/1 46 0.05 0.2/1 46 un 0.19 1/0.9 46 0.05 0.3/0.4 29 0.04 0.2/0.9 46 red 0.16 1/0.7 46 0.04 0.3/0.3 8 0.03 0.2/0.6 46 a n, illustrates the variation in adsorptive capacity from sample to sample in the same oxidation state (number of CO molecules referred to those adsorbed on k-ZnO taken as unity).b n z illustrates the variation in adsorptive capacity due to reduction for the same sample (number of CO molecules referred to those adsorbed on the ‘ox’ form taken as unity). I I I I 1.5 (61 6 * - N N ’& 1.0 ‘E 1 h 3 \ s 0.5 0 10 20 30 40 0 10 20 30 40 plTorr p/Torr Fig. 3. (a) Volumetric (n, us. p ) and (b) calorimetric (Qint us. p) isotherms of adsorption of H, on 0, k-ZnO-ox; 0, c-ZnO-ox; and 0, o-ZnO-ox. Table 2. Adsorption enthalpy and amount adsorbed measured at p = 20 Torr for H, adsorbed on ZnO samples (a) k-ZnO (b) c-ZnO (c) o-ZnO H2/ - AadsH/ molecule - ‘ads H,/ - Aads H / molecule H2/ molecule nm-, nla/n2b kJ mol-’ nm-, nla/nzb kJ mol-’ nm-, nlU/nzb kJ mol-l ox 0.64 1 / 1 60-14 0.33 0.5/1 ca.35 0.11 0.2/1 75-50 const. un 0.46 1/0.7 55-30 0.23 0.5/0.7 ca. 35 0.08 0.2/0.7 ca. 59 const. const. red 0.30 1/0.6 ca. 55 0.24 0.6/0.7 ca. 35 - - - const. const. ~~ ~ a n, illustrates the variation in adsorptive capacity from sample to sample in the same oxidation state (number of H, molecules referred to those adsorbed on k-ZnO taken as unity). n4 illustrates the variation in adsorptive capacity due to reduction for the same sample (number of H, molecules referred to those adsorbed on the ‘ox’ form taken as unity).V. Bolis, B. Fubini, E. Giamello and A . Reller 863 Fig. 4. Differential heat of adsorption (qdiff/kJ mol-l) of H, on 0, k-ZnO-ox; a, c-ZnO-ox; and 0, o-ZnO-ox us. amount adsorbed (n,/pmol m-,). with coverage is given in the differential-heat plot reported in fig.4 for the oxidised samples. Arrows in the figure indicate the adsorbed amount corresponding to 20 Torr equilibrium pressure. On k-ZnO the differential heat of adsorption decreases from an initial value of 60 kJ mol-1 to 14 kJ mol-'. A process associated with a very low heat of interaction occurs and becomes more important than the initial process at increasing coverage. Although in the case of o-ZnO the total adsorption is much lower than in the case of k-ZnO, the differential-heat curve lies above that of k-ZnO at low coverage (high-energy) sites. On c-ZnO the heat of adsorption is 35 kJ mol-l over the whole range examined. The thermokinetics of H, adsorption are typical of an activated process, in agreement with the fact that hydrogen is dissociatively chemisorbed on ZnO.'3 a, 1 3 7 l5 Under our experimental conditions the evolution of heat corresponding to the adsorption of a dose of hydrogen lasted typically between 1 and 2 h.Only in the case of k-ZnO-ox is the presence of much slower processes (lasting ca. 6-7 h) observed, as already r e p ~ r t e d . ~ Chemisorption of hydrogen at room temperature consists of a reversible and an irreversible part. In this aspect too the behaviour of the three oxidised samples is different. A large fraction of adsorbed H, (corresponding to ca. 50% of the total heat evolved) is not removed from k-ZnO-ox upon evacuation at room temperature. Irreversible adsorption is lower in the case of o-ZnO (20 % of the total heat evolved) and is practically absent on c-ZnO.Progressive prereduction of the surface somehow inhibits the irreversible phenomena, in parallel with the disappearance of the low-energy process : all the hydrogen on k-ZnO- red is reversibly adsorbed, and the heat of adsorption, 55 kJ mol-l, does not vary with coverage. Discussion The reactivity of zinc oxide surfaces towards CO and H, strongly depends upon form and reduced state of the material. The origin of the sample determines both the form of the powder particles (as evidenced by electron-microscopy results) and the Zn/O ratio, which can be more or less near to the stoichiometry according to the preparation route, as reported by Parravano et aZ.,, In order to evaluate the role of form in determining the surface activity of the sample, it is convenient to compare adsorption results obtained for the different samples864 Surface Reactivity of Zinc Oxides pretreated in oxygen (ox).Stoichiometry differences in pristine samples, caused by the different preparation routes, should be lessened by the standard oxidising thermal treatment. The effect of the reduced state of the sample upon its surface reactivity will be discussed separately. Influence of Form on Adsorption Properties Zinc oxide obtained by direct oxidation of the metal is a material that exhibits unique features if compared with zinc oxides obtained by the slow decomposition of compounds such as zinc carbonate or oxalate. The most relevant differences in form evidenced by X-ray diffraction and electron-microscopy measurements can be ascribed to the two distinct preparation routes.The ex-metal is the most crystalline sample, made up of well separated and defined particles, whereas the ex-salt powders are mainly constituted of densely packed aggregates built up of compact small particles (the specific surface areas of these samples are higher than that of Kadox). The well defined crystallites of Kadox are caused by the growth of ZnO in an oxidising atmosphere at a high temperature. Thus the kinetics of formation of the oxide are very fast in comparison with the kinetics of decomposition of a salt. In the latter case the oxide starts to grow very slowly on a salt particle that is still intact at T < 800 K. Therefore materials with less-defined particle shapes but relatively high surface areas are formed. Similar results have been found in the case of a-Al,O,.23 The main result from a comparison of data on the three samples is that the differences in form of the particles arising from the different preparation routes lead to remarkable differences in adsorption capacity. This feature appears to be important in determining the surface reactivity of k-ZnO : in spite of the fact that the ex-metal zinc oxide bears the lowest surface area among the three samples it exhibits the highest specific surface activity, as shown by the adsorption measurements. The enthalpy values reported for CO adsorption are consistent with a simple coordination process,12 and the constancy of the adsorption enthalpy (46 kJ mol-I) with coverage indicates that CO adsorption involves the same type of equivalent and non- interacting sites on all the samples.In fact, the TEM and HREM observations reported indicate that, contrary to what is generally accepted, k-ZnO does not exhibit a preferential development of polar faces with respect to c- and o-ZnO but shows a better definition of the border of the crystallites, i.e. the edges between polar and apolar planes. The differences in adsorptive capacity between k-ZnO and the two ex-salt samples can thus be explained not in terms of a different distribution of exposed faces but by taking into account the high crystallinity of the ex-metal oxide and thus the presence of well defined intersections of crystallographic planes, i.e. well defined crystal edges. Preferential sites for the reversible adsorption of CO are likely to be highly unsaturated Zn2+ cations exposed at the intersection of polar and apolar planes.This is in agreement with the assignment proposed by Lavalley et al.' In the case of c- and o-ZnO, the low adsorptive capacities observed suggest that, in spite of the poor definition of the microcrystals, which should imply a high abundance of structural defects, the cationic centres are less coordinatively unsaturated than those in k-ZnO. The correspondence between the number of active sites observed on different samples for CO and H,, respectively (the n, values reported in the first rows of tables 1 and 2), suggests that the edge sites also play an important role in hydrogen adsorption, in agreement with results reported in ref. (8).However, interaction of H, with ZnO is more complicated than the simple coordination observed in the case of CO, as indicated by the variation of adsorption enthalpy and thermokinetics with coverage observed in some cases.V. Bolis. B. Fubini, E. Giamello and A . Reller 865 The different phenomena occurring when H, is contacted with ZnO are as follows. (i) A reversible but slightly activated adsorption related to the dissociation of hydrogen molecules with corresponding formation of Zn---H and 0---H species, as previously reported.4 This kind of adsorption, usually indicated as type 1,13 is the only process observed on c-ZnO. It also takes place on k-ZnO, although it is overlapped by other phenomena:4 it is practically absent on o-ZnO. As mentioned above and suggested by the present results, the sites responsible for type I adsorption are the Zn2+ and 0,- ion pairs that are probably present at the intersection of polar and apolar planes.8 (ii) The highest energetic adsorption, slow and irreversible at room temperature, observed at low coverage on k- and o-ZnO, (fig.4) is ascribed to the dissociative homolytic adsorption onto the so-called type I1 sites proposed by Ghiotti et aL7 on the basis of infrared data. This adsorption leads to the formation of Zn---H--- Zn and 0---H---0 bridged species located at the (1010) prismatic faces, where a suitable orientation of Zn” and 02- ion pairs is available.’ (iii) The slow, irreversible and low-energy process (not detectable by infrared spectroscopy) observed in the case of k-ZnO (fig.4, high coverage) has been assigned previously by some of us to the diffusion of molecular hydrogen into the open structure of Zn0.4 This fact could also be related to the presence of structural defects such as steps, kinks and edges that would favour the entrance of the small hydrogen molecules into the structural channels. However, the diffusion of hydrogen is more likely to depend upon the surface concentration’ of H, and not upon the equilibrium pressure. Two neighbouring adsorbed hydrogen atoms must recombine in order to penetrate as a molecule into the bulk. The coverage obtained at the same equilibrium pressure is much greater on k-ZnO than on the other samples. This justifies the occurrence of diffusion only on k-ZnO and the progressive disappearance of the diffusion process itself observed on k-ZnO-un and k-ZnO-red, where the total amounts adsorbed are lower than in the case of k-ZnO-ox (table 2).Finally, note that the two products obtained by the decomposition of the parent zinc compounds exhibit different reactivities. The different adsorption capacities shown by the two ex-salt zinc oxides suggest that several factors have to be taken into account in order to prepare an active ZnO catalyst via precursor salt decomposition, i.e. the reaction conditions, such as temperature and atmosphere, the morphology of the precursor and the nature of the salt. In the case of the decomposition of zinc oxalate carbon monoxide is produced, and the zinc oxide is formed in a more reducing atmosphere than that obtained by decomposition of the carbonate.Influence of the Reducing Treatments on the Activities of the Samples The data reported in tables 1 and 2 clearly indicate that drastic changes in adsorption features occur when the solids are submitted to partially reducing treatments. The variation observed upon ZnO reduction depends, in principle, on both the form of the individual crystal particles and the electronic population of the bands : a clear distinction between these two effects is beyond the aim of the present paper and would require a detailed analysis of the conductivity of each sample. The discussion is limited, therefore, to the most relevant quantitative effects of the reduction on the chemisorption activity of ZnO, which can be summarized as follows: (i) a decrease in the total number of active sites and (ii) a selective inhibition of particular mechanisms of adsorption.The coordination of carbon monoxide involves the donation of electrons to the coordinatively unsaturated surface cations : a surface with an increased electron density is obviously less active towards this kind of adsorption. The dissociative adsorption of hydrogen requires sites involving pairs of zinc and oxygen species : the progressive elimination of oxygen by reducing treatments thus lowers the potential number of active sites. In the case of c-ZnO the decrease in the number of sites coordinating CO also866 Surface Reactivity of Zinc Oxides involves a progressive decrease in the interaction energy (- A,,,H decreases from 46 to 29 and 8 kJ mol-1 for the ‘ un ’ and ‘ red ’ samples, respectively).Since the adsorption enthalpy is a measure of the acid-base interaction, not only the number of acid sites but also their acid strength is affected by the reducing treatments. A similar decrease in the acid strength of the surface sites has not been observed in the k- and o-ZnO samples. The reducing treatments in these cases lower the number of acid sites, although to a lesser extent than in c-ZnO, but leave their acid strengths unchanged. As far as H, adsorption is concerned, the reduction involves not only a decrease in the total amount adsorbed, but also remarkable modifications to the adsorption features. Calorimetric data indicate a different distribution of the various processes. In particular, fewer sites for type I adsorption are available on reduced k- and c-ZnO samples: this decrease parallels that of CO adsorption on these two samples, in agreement with the assignment of the same sites of adsorption for CO and H2(I).Furthermore (as mentioned above), the slow diffusion of H, into the bulk observed on k-ZnO-ox is inhibited upon reduction. Conclusions The insights obtained from the experimental studies presented permit a discussion emphasizing the interdependence between the individual form of a solid and its effect on surface chemistry on processes involving heterogeneous catalysis. The precursor material and the preparation route leading to solid products with characteristic morphological features (and therefore with localised and defined numbers of active sites) are among the important factors influencing specific catalytic activity.In the case of ZnO the predominance of hexagonal faces was regarded as a decisive feature explaining the comparatively high activity of Kadox [see e.g. ref. (7) and (S)]. The experimental findings presented here do not support this assumption, which therefore must be discarded or at least revised. The decisive feature is instead the high perfection of the k-ZnO crystallites, and therefore the presence of well defined edges between the hexagonal and prismatic faces. In comparison with the ex-salt samples, these edges establish a relatively high density of coordination sites for CO. As the energy of interaction of CO is the same in all cases, the variations of the total amounts adsorbed observed for samples of different origins are ascribed to the presence of a different number of equal sites, essentially depending on ‘ form ’ factors.Furthermore, in contrast to the results indicated by Natta in the earliest work on ZnOlO and to what is generally accepted,18 a high surface reactivity of a zinc oxide sample is not necessarily related to a high surface area. An abundance of specific sites, which strongly depends upon the preparation route, is the most important feature for an active zinc oxide surface. In addition, the surface reactivity is influenced by the redox pretreatments undergone by the sample. This work has been supported by the Italian ‘Minister0 della Pubblica Istruzione’ (Progetto Nazionale ‘Struttura e Reattivita’ delle Superfici’). We are indebted to Dr J. C . Lavalley for kindly supplying the ex-oxalate zinc oxide sample. References 1 W. Hotan, W. Gopel and R. Haul, SurJ Sci., 1979, 83, 162. 2 H. R. Gay, M. H. Nodine, E. I. Solomon, V. H. Henrich and H. J. Zeiger, J . Am. Chem. Soc., 1980, 3 S. Akther, K. Lui and H. H. Kung, J . Phys. Chem., 1985, 89, 1958. 4 B. Fubini, E. Giamello, G. Della Gatta and G. Venturello, J . Chem. SOC., Furuduy Truns. 1, 1982, 78, 102, 6752. 53.V. Bolis, B. Fubini, E. Giamello and A . Reller 867 5 M. Bowker, H. Houghton, K. C. Waugh, T. Giddings and M. Green, J. Catal., 1983, 84, 252. 6 E. Giamello and B. Fubini, J . Chem. Soc., Faraday Trans. I , 1983, 79, 1995. 7 G. Ghiotti, F. Boccuzzi and R. Scala, J . Catal., 1985, 92, 79. 8 J. C. Lavalley, J. Saussey and T. Rais, J . Mol. Catal., 1982, 17, 289; C. Chauvin, J. Saussey, J. C. 9 B. Fubini, V. Bolis and E. Giamello, Thermochim. Acta, 1985, 17, 283. Lavalley and G. Djega-Mariadassou, Appl. Catal., 1986, 25, 59. 10 G. Natta, in Catalysis, ed. P. H. Emmet (Reinhold, New York, 1955), vol. 3. 11 K. Klier, Adu. Catal., 1982, 31, 243. 12 C. H. Amberg and D. A. Seanor, in Proc. Third Znt. Congr. Catal. (North Holland, Amsterdam, 1964), 13 R. J. Kokes, A. L. Dent, C. C. Chang and L. T. Dixon, J . Am. Chem. SOC., 1972, 94, 4429. 14 F. Boccuzzi, E. Garrone, A. Zecchina, A. Bossi and M. Camia, J. Catal., 1978, 51, 160. 15 G. L. Griffin and J. T. Yates Jr, J. Chem. Phys., 1982, 77, 3744; 3751. 16 L. A. Denisenko, A. A. Tsyganenko and V. N. Filimonov, React. Kinet. Catal. Lett., 1984, 25, 23. 17 G. L. Griffin and J. T. Yates Jr, J . Catul., 1982, 73, 396. 18 M. Bowker, H. Houghton and K. C. Waugh, J . Chem. Soc., Faraday Trans. I , 1981,77,3023; 1982,78, 19 G. Heiland and H. Luth, in The Chemical Physics of Solid Surfaces and Heterogeneous Catalysis, ed. 20 B. Fubini, RPc. Gen. ThPrmique, 1979, 18, 297. 21 M. Maciejewski and H. R. Oswald Thermochim. Acta, 1985, 85, 39. 22 G. Carnisio, F. Garbassi, G. Petrini and G. Parravano, J. Catal., 1978, 54, 66. 23 D. Scarano, A. Zecchina and A. Reller, work in preparation. paper I. 22. 2573. D. A. King and D. P. Woodruff (Elsevier, Amsterdam, 1984), vol. 3, part B. Paper 8/01620H; Received 25th April, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500855

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1989, 85(4), 869-881 Characterization of Iron Oxide-dispersed Activated Carbon Fibres with Fe K-Edge XANES and EXAFS and with Water Adsorption Katsumi Kaneko" Department of Chemistry, Faculty of Science, Chiba University, Yayoi, Chiba 260, Japan Nobuhiro Kosugi and Haruo Kuroda Department of Chemistry, Faculty of Science, The University of Tokyo, Hongo, Tokyo 113, Japan The Fe K-edge X-ray absorbtion near-edge structure (XANES) and extended X-ray absorption fine structure (EXAFS) and the water adsorptivity of iron oxide-dispersed activated carbon fibres (ACF) have been investigated. Also, Fe K-edge XANES and EXAFS of various kinds of powdered iron oxides have been measured in order to characterize the species dispersed on the ACF. The XANES and EXAFS indicated that a- Fe00H-like ultrafine particles transform into a-Fe,O,-like particles on the ACF through heating in UCICUO.The adsorption isotherms of water on the a-Fe00H-dispersed ACF were of type V. The water adsorption isotherms were analysed by the Dubinin-Serpinsky equation; the number of polar sites on the surface was estimated and compared with the data from an analysis of the evolved gas. The relationship between the water adsorptivity and the states of the iron oxides on the ACF is discussed. A comparison of nitrogen and water adsorption shows that the layer of water adsorbed on iron oxides dispersed on the ACF has a more sparse structure than normal water, and that the iron oxide-dispersed and preoxidized ACF samples exhibit an excess of water adsorptivity owing to both molecular-sieve effects and surface fractal structure. The surface-chemical properties and the characterization of activated carbon fibres (ACF) have been actively The ACF are highly microporous, with small external surface areas and very little mesoporosi ty, having characteristic adsorption properties.Consequently, an ACF can be a good model system in the study of micropore filling. We are interested in the micropore filling of supercritical NO on metal oxide- dispersed ACF. As the two-dimensional critical temperature of NO is estimated to be 90 K by de Boer's equation,' and the micropore filling is a limited case of physical adsorption which is effective for vapours,8 normal activated carbons cannot adsorb abundant NO by micropore filling. The surface modification of an ACF with iron oxides enhances the NO micropore filling; these sites adsorb large amounts of NO up to 320 mg g-' at 303 K by chemisorption-assisted micropore filling.'^ lo Inter-NO molecular interactions are enhanced in the micropores.l1 Also, surface modification with iron oxides increases the adsorptivity of the ACF for CO, and NH,.12 Characterization of the iron oxides dispersed on the ACF is necessary, therefore, to unveil the mechanism of c hemisorp tion- assis ted micropore filling. Carbons are almost transparent to hard X-rays, and metal K-edge absorption spectroscopy (XANES and EXAFS) is a very useful method for the characterization of metal oxides on carbons ; X-ray absorption near-edge structure (XANES) and extended 869870 Surface Properties of Active Carbon Fibres X-ray absorption fine structure (EXAFS) give important information on the local environment of a specific atom in finely dispersed metal oxides on the ACF.Thus we have applied X-ray absorption spectroscopy to characterize the ultrafine metal oxides on the ACF of interest. The preliminary EXAFS results on Fe,O,-dispersed ACF have already been reported, together with NO adsorption data,', and an EXAFS study of Cu(OH),-dispersed ACF appeared in the previous paper. l4 A number of studies of water adsorption on carbon^^^-^' have shown that the interaction energy of water with the carbon surface is unusually small and that the adsorption isotherms of water on carbons, which change sensitively with the surface oxidized state, are unusual in comparison with nitrogen isotherms.Dubinin and Serpinsky,' assumed that water molecules adsorbed on polar sites act as sites for further water adsorption ; the number of active sites for water can be roughly estimated from the Dubinin-Serpinsky equation.lg a-FeOOH and a-Fe,O, are active to water adsorption through hydrogen-bond formation with surface h y d r o ~ y l s . ~ l - ~ ~ The crystallite size of a- FeOOH can be controlled by changing the time for hydrolysis of the Fe3+ solution. Thus iron oxides dispersed on the ACF are expected to be associated with the water adsorptivity of the modified ACF. Furthermore, a water molecule is smaller than a nitrogen molecule : hence the water adsorption of iron oxide-dispersed ACF can lead to information on the microporosity different from that obtained from nitrogen adsorption.In this work Fe K-edge XANES and EXAFS of iron oxides dispersed on ACF and various kinds of iron oxides and iron oxide hydroxides were measured, together with the corresponding water adsorptivity, so as to characterize the micropore structure of the iron oxide-dispersed ACF systems. Experiment a1 Materials Cellulose-based ACF (Toyobo KF- 1500) and ACF samples preoxidized with 6 mol dm-, HNO, at 373 K, denoted ox-ACF, were used. a-FeOOH particles of different crystallite sizes were dispersed on the ACF by hydrolysing 0.6 mol dm-, Fe,(SO,), solution at 303 K and pH 13 for various hydrolysis times, t, of 0-10 days. These treated ACF samples are designated a-t-ACF. a-FeOOH was also dispersed on the ox-ACF by a 6 h hydrolysis of the Fe3+ solution (a-ox-ACF).a-FeOOH was dispersed on the ACF and ox-ACF after outgassing at 383 K and a reduced pressure of 1 mPa for 15 h. a-ox- ACF-T were obtained by heating a-ox-ACF samples at T K (573 and 773 K) under 1 mPa for 15 h. The Fe content in the ACF and ox-ACF samples was determined by titrating Fe3+ ions eluted in HCl solution with a 0.005 mol dmP3 K,Cr,O, solution. a-Fe00H powders of different crystallite sizes, which are denoted by a-FeOOH-t where t is the hydrolysis time, were prepared for comparison with the ACF-supported samples;24 the (110) crystallite sizes of these a-FeOOH samples are as follows: a-FeOOH-0, 3 nm; a-FeOOH-3h, 7 nm; a-FeOOH-6h, 8 nm; a-FeOOH-2d, 9 nm and a-Fe00H-lOd, 11 nm.Also, a-FeOOH synthesized according to the method of Atkinson rt uZ.~~(~-F~OOH-AT) was examined. Synthetic P-FeOOH,23 y-FeOOH2:3 and 6-Fe00H26 samples were prepared in the same way as described elsewhere. Here we used two kinds of y-FeOOH (y-FeOOH-2 and y-Fe00H-25) having different (002) crystallite sizes, 4.5 and 12 nm, respectively. Four kinds of a-Fe,O, were obtained by decomposition of (i) the amorphous hydroxides from an Fe(NO,), solution at 873 K (a-Fe,O,-I), (ii) a-FeOOH from an FeSO, solution at 673 K (a-Fe,O,-IT), and a-Fe00H from an Fe,(SO,), solution at (iii) 673 K (a-Fe,O,-III) and (iv) 573 K (cc-Fe,0,-IV).27Commercial Fe,O, (Wako Chemicals) and y-Fe,O, (Toda Industry) powders were used for comparison.J . Chem. Soc., Favaday Trans. I , vol.85, part 4 Plate 1 Plate 1. Crystal structures of iron oxides and iron oxide hydroxides: (A) cz-FeOOH, (B) /3- FeOOH, (C) y-FeOOH, (D) Fe(OH),, (E) a-Fe,O, and (F) y-Fe,O, and Fe,O,. K. Kameko, N. Kosugi and H. Kuroda ( Facing p . 87 1 )K. Kaneko, N . Kosugi and H. Kurodu 87 I Water Adsorption and Nitrogen Adsorptions The adsorption isotherms of water on the ACF samples were measured gravimetrically at 303 +O. I K. The entire adsorption apparatus was held at 303 f 0.5 K. The adsorption isotherms of nitrogen on the ACF samples at 77 K were measured in a similar way to the water adsorption. As almost all data on nitrogen adsorption have been reported previously,1° only a brief description of the nitrogen adsorption is given here. The samples were evacuated at 373 K and 1 mPa for 15 h prior to the adsorption experiments. Characterization with XANES and EXAFS and Evolved Gas Analysis The Fe K-edge XANES and EXAFS spectra of a-FeOOH, p-FeOOH, y-FeOOH, 6- FeOOH, a-Fe,O,, y-Fe,O, and Fe,O, powders and a-6h-ACF, a-ox-ACF, a-ox-ACF- 573 and a-ox-ACF-773, were measured using the EXAFS apparatus at BL-7C of the Photon Factory in the National Laboratory for High Energy Physics (Tsukuba, Japan).,* The phase-uncorrected Fourier transforms of k 3 ~ ( k ) were determined from the EXAFS oscillation ~ ( k ) .~ ' The evolved-gas analysis (EGA) spectra of ACF, a-6h-ACF, a-ox-ACF and a-ox-ACF-773 up to 773 K were obtained at a heating rate of 10 K min-' with the aid of a mass filter (ULVAC, MSQ-15OA). Results and Discussion Species Produced oir ACF with XANES and EXAFS Various kinds of iroii oxides and iron oxide hydroxides are known.a-Fe,O, has a corundum structure and y-Fe,O, a spinel-like structure with a deficiency of metal ions. Fe,O, has an inverse spinel structure. The representative four polymorphs of FeOOH are a-FeOOH, P-FeOOH, y-FeOOH and 6-FeOOH. These FeOOH crystals have characteristic structures. Plate 1 shows the crystal structures of these iron oxides and iron oxide hydroxides. The structural unit of these compounds is an iron-centred octahedron with oxygens or hydroxyls at the corners; the octahedra are linked together by sharing corners, edges, faces and hydrogen bonds, giving rise to each characteristic crystal structure. Fe,O, and y-Fe,O, have tetrahedra embracing a central iron atom rather than the octahedra.It is said that the fundamental structure of 6-FeOOH is the same as that of Fe(OH),, with S-FeOOH more defective, having a few tetrahedral Fe sites. These iron oxides have different Fe-0 and Fe-Fe distances, as collected in table 1 ; each compound has an inherent local structure. The interatomic distances are cited or calculated from the literature from the X-ray diffraction studies. The range of data for each bond arises from the existence of different sorts of bonds identified by accurate X-ray examination. Fig. 1 shows the XANES of the Fe IS edge of a-FeOOH, p-FeOOH, y-FeOOH, a- Fe,O,, a-6h-ACF, a-ox-ACF and a-ox-ACF-773. All the XANES spectra are similar; each has three main peaks (A), (B) and (C). A weak pre-edge peak (A) arises from the 1s-3d transition which is dipole-forbidden in the case of the 0, symmetry around an Fe3+ ion.The (A) peak of a-FeOOH-0 is the strongest in these samples; the octahedral unit around Fe3+ of a-FeOOH-0 is therefore much distorted. The intensities of the (A) peaks of a-6h-ACF, a-ox-ACF and a-ox-ACF-773 are so weak that the 0, symmetry of Fe3+ is retained irrespective of the state of dispersion on the ACF. Peaks (B) and (C) are caused by the ls-4pa transition and multiple scattering, respectively. Each sample shows slightly different features in (B) and (C). Only a-Fe,O, has a (B) peak with a shoulder and a relatively high (C) peak. Although these XANES spectra can give no definite identification of the species dispersed on the ACF, XANES spectra of a-6h-872 Surface Properties of Actitle Carbon Fibres Table 1.The nearest Fe-O and Fe-Fe distances (nm) of iron oxides obtained from X-ray dif- fraction studies in the literature Fe-O/nm Fe-Fe/nm ref. a-Fe,O, 0.199 0.206 y-Fe,O, 0.186" 0.205b Fe,O, 0.187" a-FeOOH 0.197 0.202 0.214 B-FeOOH 0.1 89 0.21 1 y-FeOOH 0.189 0.199 0.206 0.207 6-FeOOH 0.20 0.288 0.295 0.297 0.286 0.302 0.301 0.288 - - - - - - - 0.23 30, 31 30, 32 30, 33, 34 30, 35 - - - - 36, 37 30, 38 - - - - 39, 40 a Tetrahedral. Octahedral 7090 7130 7170 photon energylev Fig. 1. XANES of the Fe K edge of iron oxides and iron oxides dispersed on the ACF: (a) a-ox- ACF-773, (b) a-ox-ACF, (c) a-6h-ACF7 ( d ) a-Fe,O,-111, (e) y-Fe00H-25, (f) /I-FeOOH, ( g ) a- FeOOH-2d, (h) a-FeOOH-6h and (i) a-FeOOH-0.K. Kaneko, N.Kosugi and H . Kuroda 873 0 0.2 0.4 0.6 0 0.2 0.4 0.6 distancelnm Fig. 2. Fourier transforms of the EXAFS oscillation for (a) a-Fe,O,-I, (6) a-Fe,O,-11, (c) a-Fe,O,-111, ( d ) a-Fe,O,-IV (e) y-Fe,O, and (f) Fe,O,. ACF, a-ox-ACF and a-FeOOH are similar to each other, with the (C) peak of the XANES spectrum of a-ox-ACF-773 resembling that of a-Fe,O,. We can obtain more information on the local structure of these iron oxides by EXAFS than by XANES, but we should note that the nearest-neighbour bond distance from the EXAFS data without phase correction does not coincide with that obtained from X-ray diffraction. Fig. 2 shows phase-uncorrected Fourier transforms of k 3 ~ ( k ) from the EXAFS oscillation ~ ( k ) for four kinds of a-Fe,O,, y-Fe,O, and Fe,O,.There are two main structures, [A] (0.1-0.2 nm) and [B] (0.24.35 nm), which are ascribed to Fe-0 and Fe-Fe distances, respectively. The low-intensity structure below 0.1 nm comes from errors in finite Fourier transformation, background subtraction etc. The four spectra of the obvious a-Fe,O, samples have similar features, although their intensities are different from each other. On the other hand, y-Fe,O, and Fe,O, exhibit distinctly different EXAFS spectra. Fig. 3 shows Fourier transforms of k 3 ~ ( k ) for a-FeOOH, D- FeOOH, y-FeOOH and 6-FeOOH. These FeOOH modifications have the [A] and [B] structures ;imilar to those of a-Fe,O,. The peak positions of [A] and [B] are similar, but the intensides and the intensity ratio of [A] and [B] are characteristic of each structural modification, Accordingly, we can obtain useful information on the polymorphism from the EXAFS data.The Fourier transforms of k 3 ~ ( k ) for a-FeOOH of different crystallite sizes are shown in fig. 4. The peak positions of [A] and [B] change little with the874 Surface Properties of Active Carbon Fibres 0 0.2 0.4 0.6 - A B 0 0.2 0.4 0.6 distance/nm Fig. 3. Fourier transforms of the EXAFS oscillation for (a) a-FeOOH- 10d, (b) a-FeOOH-AT, (c) a-FeOOH, ( d ) y-FeOOH-2, (e) y-Fe00H-25 and (f) 6-FeOOH. -2 3 (I .r( a 0 d g 9 6 - 3 - t A A B (4 gt A 0 0.2 0.4 0.6 0 02 0.4 0.6 distancehm Fig. 4. Fourier transforms of the EXAFS oscillation for a-Fe00H aged for various periods : (a) none, (b) 6 h, (c) 2 days and (d) 10 days.K. Kaneko.N . Komgi and H. Kuroda 9 - 6 - 3 - A (4 B 0 0.2 0.L 0.6 0 0.2 0.4 0.6 875 distancefnm Fig. 5. Fourier transforms of the EXAFS oscillation for (a) a-FeOOH-fih, (b) a-6h-ACF7 (c) a-ox- ACF, (6) or-ox-ACF-573, (e) a-ox-ACF-773 and (f) a-Fe,O,-111. crystallite size, but we can see the relative change of [A] and [B] peak intensities, which reflects the particle size; the smaller the intensity ratio of [B]/[A], the smaller the particle size. In particular, the [B] peak of a-FeOOH-0 is rather weak, indicating that the a- FeOOH-0 lattice has a low coordination number of nearest-neighbour Fe ions, i.e. a- FeOOH-0 is an ultrafine a-FeOOH particle. Fig. 5 shows the Fourier transforms of k 3 ~ ( k ) for a-FeOOH, a-Fe,O, and iron oxides dispersed on ACF. The EXAFS spectrum of a-ox-ACF is comparable with that of a-FeOOH; the [B] peak of a-ox-ACF is much weaker than that of a-FeOOH, indicating that a-Fe00H-like species, with a lower coordination number of nearest Fe ions, are dispersed on ox-ACF.The EXAFS of a- 6h-ACF leads to a similar conclusion to that for a-ox-ACF, although the [B] peak of a- 6h-ACF is more intense than that of a-ox-ACF. Heating a-ox-ACF brings about a significant change in the EXAFS spectra, as shown in fig. 5(4 and (e). The features of a-ox-ACF-573 and a-ox-ACF-773 resemble those of both a-FeOOH and a-Fe,O,; it is difficult to identify the species. However, the [B] peak grows on heating. It is probable that ultrafine a-Fe00H-like species on the ACF decompose into a-Fe,O,-like particles through heating; these then aggregate.Water Adsorption The adsorption isotherms of water on ACF and a-ACF are of type V, having hystereses (fig. 6); the amount of adsorption is slight until relative pressures of 0.4 are reached. The starting relative pressure associated with steep uptake moves to higher values with the growth of a-FeOOH. As we obtained readings which were invariant over 1 h, there is a possibility that the hystereses disappear after much longer adsorption times, as indicated by McBain.’’ Fig. 7 shows the adsorption isotherms of water on ox-ACF, a-ox-ACF 30 FAR 1876 Surface Properties of Active Carbon Fibres 0 0.5 PlPo 1.0 Fig. 6. Adsorption isotherms of water on ACF and a-ACF: 0, ACF; 0, a-6h-ACF; A, a-2d-ACF. - , O O t 0 0.5 1.0 PIP0 Fig. 7. Adsorption isotherms of water on the preoxidized ACF samples: 0, ox-ACF; 0, CZ-OX-ACF; A, a-OX-ACF-773.and a-ox-ACF-773. The adsorption isotherm of ox-ACF is like that of a-ox-ACF, whereas the adsorption isotherm of a-ox-ACF-773 is clearly different from the latter. The amount of water adsorption of ox-ACF and a-ox-ACF increases monotonically with the relative pressure even below 0.4; the preoxidized ACF interacts more strongly with water than do ACF and a-ACF. We determined the micropore volume for water, W,(H,O), from extrapolation to p / p o = 1. W,(H,O) data are collected in table 2. The type V isotherm can be approximated by the quadratic form of the Dubinin-Serpinsky equation :19* 2o Here w represents the amount of water adsorption at p/p,, a, is the amount of adsorption on the polar sites and c is a constant.Eqn (1) is expressed by a parabola; aK. Kaneko, N . Kosugi and H. Kuroda 877 Table 2. Comparison of water and nitrogen adsorption data and the amount of dispersed Fe ACF a-0-ACF a-3 h-ACF a-6h-ACF a-2d-ACF a- 1 Od-ACF OX-ACF ~z-ox-ACF a-OX-ACF-573 a-OX-ACF-773 W,(H,O) /cm3 g-I 0.62 0.63 0.62 0.55 0.58 0.52 0.48 0.47 0.47 0.48 a0 /cm3 g-I 0.12 0.18 0.18 0.14 0.04 0.08 0.13 0.08 0.10 0.09 4)" /cm3 g-I 0.61 0.6 1 0.62 0.62 0.62 0.59 0.57 0.36 0.37 0.44 Fe (wt Yo) 1 .o I .o 1 .o 0.90 0.94 0.88 0.84 1.3 1.3 1.1 - - 11 8.6 3.5 4.5 2.6 3.0 3.1 - n 0 200 400 adsorbed HzO/mg g-* 500 500 600 300 200 Fig. 8. Dubinin-Serpinsky plots for the adsorption isotherms of water on ox-ACF (0) and a-OX-ACF-773 (0). plot of w(p,/p) us. w should be an inverted parabola if the adsorption isotherms obey eqn (1).Fig. 8 shows Dubinin-Serpinsky plots for the adsorption isotherms of water on ox-ACF and 3-ox-ACF-773. The relationship for ox-ACF is a good parabola, but the Dubinin-Serpinsky plot of a-ox-ACF-773 deviates. We estimated the number of polar sites from the coordinates of the top of the approximate parabola; the a, data are collected in table 2. These a, values decrease with the hydrolysis time of a-FeOOH in the a-ACF system, whereas a, is almost constant irrespective of the heating temperature in the case of the a-ox-ACF system. The surface functional groups on a-ACF must be 30-2878 Surface Properties of Active Carbon Fibres 300 400 500 600 700 T / K Fig. 9. Temperature profile of evolved CO and CO,.(a) open symbols, a-ox-ACF; solid symbols, a-OX-ACF-773; (b) a-6h-ACF; (c) ACF. covered with a-Fe00H-like species that grow as fine particles. As for the a-ox-ACF system, heating removes the surface functional groups and hydroxyls of a-FeOOH on a-ox-ACF, producing hydrophilic a-Fe,O,-like species on the surface. Although a-ox- ACF has a more hydrophilic nature than ACF because of the uptake of water in the low p / p o region, the a, value does not necessarily reflect the water adsorption behaviour. The value of a, is not absolute, but leads to only a qualitative tendency. The EGA spectra give more definitive information on the surface polar groups present on ACF. Fig. 9 shows temperature profiles of the evolved CO and CO, from ACF, a- ACF, a-ox-ACF and a-ox-ACF-773. All samples start to evolve both gases near 400 K.a-ox-ACF has the largest CO and CO, peaks, with a CO, evolution three times larger than that of CO. In the case of a-6h-ACF and ACF, similar amounts of CO and CO, are evolved in the 40&700 K range. However, ACF evolves more CO than CO, above 700 K, whereas a-ACF evolves more CO, than CO. CO, is derived from COOH and lactone groups on the carbon surface and CO is caused by phenolic OH and quinone- type oxygen^.^^ Consequently, a-ox-ACF has more COOH and lactone groups than phenolic OH and quinone-type oxygen on the surface. The amounts of COOH and lactone groups are comparable to those of phenolic OH and quinone-type oxygens on the surface of ACF and a-ACF. a-ox-ACF-773 has no gas evolution peak below 700 K ; hence only a few surface functional groups remain on the micropore surface of a-ox- ACF-773.However, an accurate determination of the surface polar groups requires EGA up to 1300 K. Micro porosity A water molecule (0.108 nm2) is smaller than a nitrogen molecule (0.162 nm2). The difference in saturated adsorption amounts of water and nitrogen should be related to two factors, i.e. a molecular-sieve effect and the surface fractal structure of the micropore walls. The molecular-sieve effect represents the presence of narrower micropores, which are not filled with nitrogen but are filled with water. The molecular fractal nature of solid surfaces developed by Avnir and Pfeifer42.4'3 is also important forK. Kaneko, N . Kosugi and H. Kuroda 879 0.02 cm3 g-' I 3 4 5 ~n [u/(Io-' nm3>1 Fig.10. Separation of the molecular-sieve effect and fractal structure for the difference of water and N, adsorptions on a-ox-ACF; m, observed point of water adsorption; bold line, W cc t j - D r J 3 (Dp = 2.5)." the micropores, the widths of which are greater than the diameter of a nitrogen molecule. One of the present authors determined the exponent, D,, for micropores of the ACF systems by organic vapour a d s ~ r p t i o n . ~ ~ When the ratio of Wo(H20) to Wo(N,) is > 1, the above two factors must be taken into account. The Wo(H2)/Wo(N2) values for a-ACF and ox-ACF are one or less, and are ascribed to causes other than the above two factors. The value Wo(H20)/Wo(N,) < 1 for a-ACF and ox-ACF could be due to the decrease of the density of water in the micropores, as Gregg and Sing suggested;45 the presence of fine a-Fe00H-like particles should lead to a sparser structure for the adsorbed water than that of ordinary liquid water or ice in the micro pore^.^^ On the other hand, the Wo(H20)/Wo(N2) ratios for a-ox-ACF and a- ox-ACF-T series are 1.1-1.3. Accordingly, a-ox-ACF and a-ox-ACF-T probably have narrower micropores and/or a fractal nature.We can distinguish the two factors using the molecular resolution relationship of a-ox-ACF reported earlier.44 Here we extend this fractal relationship to the size of the water molecule, as shown in fig. 10. The ordinate and abscissa of fig. 10 are the saturated amount of adsorption, W; in mmol g-l for various vapours and the molecular volume u, respectively.Here we obtained the u for water from the molecular area of 0.108 nm,. The observed pore volume for water deviates from the fractal linear plot; the amount of deviation (0.09 cm3 g-') probably corresponds to the increment of water adsorption due to the molecular-sieve effect. The difference (0.02 cm3 g-I) between water and N, adsorptions calculated from the molecular resolution relationship in fig. 10 is the adsorption increment due to the fractal nature of the micropore wall. We assume that a-ox-ACF-537 and a-ox-ACF-773 have similar narrow micropores and fractal structures to a-ox-ACF, although we did not obtain such relationships for a-ox-ACF- T. In situ X-ray diffraction and small-angle scattering studies are necessary to elucidate the structural character of the molecular- sieve effect.880 Surface Properties of Active Carbon Fibres Conclusions The Fe K-edge XANES and EXAFS of various kinds of powdered iron oxides of known crystal structures show characteristic features, assisting the identification of species of iron oxides dispersed on activated carbon fibres.Ultrafine a-FeOOH dispersed on activated carbon fibres is probably transformed into fine a-Fe,O,-like particles upon heating at 573 and 773 K in vacuo. All the adsorption isotherms of water on ACF samples have hystereses. ACF surfaces without preoxidation with HNO, are hydrophobic, whereas ACF surfaces preoxidized with HNO, are considerably hydrophilic. The Dubinin-Serpinsky analyses and EGA examinations are available for the assessment of surface properties.In the case of water adsorption of ox-ACF and a-ACF, the density of water in the micropores is lower than that of ordinary water. A comparison of nitrogen and water adsorption on a-ox-ACF and a-ox-ACF- T strongly indicates the presence of molecular- sieve effects and the defective nature of the micropore walls. This work was partly supported by a Grant in Aid for Fundamental Scientific Research from the Ministry of Education of Japan. Special thanks are due to Prof. Y. Iwasawa for suggestions on the EGA experiments and to Mr H. Kuwabara for assistance in preparing plate 1. References 1 2 3 4 5 6 7 8 9 10 1 1 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 J. N. Bohra and K. S. Sing, Acisorption Sci. Tech. 1985, 2, 89.S. J. Hichchock, B. McEnaney and S. J. Waltling, J. Chem. Tech. Biotech., 1983, 33A, 157. J. Koresh and A. J. Soffer, J. Chem. SOC., Faraday Trans. 1 , 1980, 76, 2472. J. J. Freeman, F. G. R. Gimblett, R. A. Roberts and K. S. W. Sing, Carbon, 1987, 25, 559. G. G. Jayson, J. A. Sangster, G. Thompson and M. C. Wilkinson, Carbon, 1987, 25, 523. K. Kaneko, Y. Nakahigashi and K. Nagata, Carbon, 1988, 26, 327. J. H. de Boer, The Dynamical Character of Adsorption (Clarendon Press, Oxford, 1968), p. 147. M. M. Dubinin, Carbon, 1983, 21, 359. K. Kaneko and K. Inouye, Carbon, 1986, 24, 772. K. Kaneko, Langmuir, 1987, 3, 357. K. Kaneko, N. Fukuzaki and S. Ozeki, J. Chem. Phys., 1987, 87, 776. K. Kaneko, S. Ozeki and K. Inouye, Colloid Polym. Sci., 1987, 265, 1018. K. Kaneko, Characterization of Porous Solids, Proc.IUPAC Symp. (Elsevier, Amsterdam, 1988), p. 183. K. Kaneko, A. Kobayashi, T. Suzuki, S. Ozeki, K. Kakei, N. Kosugi and H. Kuroda, J. Chem. SOC. Faraday Trans. 1, 1988, 84, 1795. J. W. McBain, J. L. Porter and R. F. Sessions, J. Am. Chem. Soc., 1933, 55, 2294. P. H. Emmett and R. B. Anderson, J. Am. Chem. Suc., 1945,67, 1492. P. L. J. Walker and J. Janov, J. Colloid Interface Sci., 1968, 28, 449. S. J. Gregg and K. S. W. Sing, Adsorption, Surface Area and Porosity (Academic Press, London, 2nd edn, 1982), p. 262. M. J. B. Evans, Curbon, 1987, 25, 81. M. M. Dubinin and V. V. Serpinsky, Carbon, 1981, 19, 402. E. McCafferty, V. Pravidic and A. C. Zettlemoyer, Trans. Faraday Soc., 1970, 66, 1720. K. Kaneko, M, Serizawa, T. Ihikawa and K. Inouye, Bull. Chem. SOC. Jpn, 1975, 48, 1764. K. Kaneko and K. Inouye, Bull. Chem. SOC. Jpn, 1979, 52, 315. K. Kaneko and K. Inouye, J. Chem. Tech. Biotech., 1987, 37, 11. R. J. Atkinson, A. M. Posner and J. P. Quirk, J. Inorg. Nucl. Chem., 1968, 30, 2371. K. Kaneko and K. Inouye, Acisorption Sci. Tech., 1986, 3, 1 1 . T. Ishikawa, M. Okamoto, Y. Ito and K. Inouye, Nippon Kagaku Kaishi, 1972, 1751. M. Nomura, KEK Report 87-1, National Laboratory for High Energy Physics (1987). N. Kosugi and H. Kuroda, Program EXAFSl/V04 and EXAFS/V03, Research Center for Spectrochemistry, The University of Tokyo (1987). E. J. Fasiska, Corrosion Sci., 1967, 7, 833. L. Pauling and S. B. Hendricks, J. Am. Chem. Soc., 1925, 47, 781. A. K. Nikumbh, K. S. Rane and A. J. Mukhedhar, J. Mater. Sci., 1982, 17, 2503.K. Kaneko, N . Kosugi and H . Kuroda 88 1 33 E. J. Verway, P. W. Haayman and F. C. Romeijn, J. Chem. Phys., 1947, 15, 181. 34 E. Z. Basta, Miner. Mag., 1957, 31, 431. 35 F. J. Ewing, J. Chem. Phys., 1935, 3, 203. 36 A. L. Mackay, Miner. Mag., 1960, 32, 545. 37 A. Szytula, M. Balanda and Z. Dimitrijevic, Phys. Stat. Sol. (a), 1970, 3, 1033. 38 F. J. Ewing, J. Chem. Phys., 1935, 3, 420. 39 M. H. Francombe and H. P. Rooksby, Clay Miner., 1959, 4, 1. 40 0. Muller, R. Wilson and W. Krakow, J. Mater. Sci., 1979, 14, 2929. 41 D. Rivin, Rubber Chem. Technol., 1971, 44, 307. 42 P. Pfeifer and D. Avnir, J. Chem. Phys., 1983, 79, 3558. 43 D. Avnir, D. Farin and P. Pheifer, Nature (London), 1984, 308, 261. 44 K. Kaneko, Langmuir, unpublished work. 45 S. J. Gregg and K. S. W. Sing, Adorption, Surface Area and Porosity (Academic Press, London, 2nd edn, 1982), p. 267. Paper 8/01 7 15H; Received 3rd May, I988

ISSN:0300-9599    

DOI:10.1039/F19898500869

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chern. SOC., Faradaj. Trans. I , 1989, 85(4), 883-893 Formate Oxidation induced by a Copper Peroxo Complex produced in Fenton-like Reactions Harry C . Sutton Institute of Nuclear Sciences, DSIR, Lower Hutt, New Zealand Yields of CO, resulting from the oxidation of formate induced by the reaction of copper(1) 1 : 10 phenanthroline [(OP),Cu+] with H,O, have been studied (a) by radiolysis of (OP)2Cu2+ - H,O, -- formate mixtures in which (OP),Cu+ is produced at low steady-state concentrations and (b) by mixing initially high concentrations of (OP),Cu+ with H,O, and formate. The CO, yields per mole of (OP),Cu+ were independent of [H,02], but increased with increasing [HCO;] and with decreasing [(OP),Cu+] to a limiting value of 1.75, which was independent of these concentrations, and hence of dose rate during radiolysis.It is concluded that Fenton-like reactions of (OP),Cu+ produce both (OP),Cu3+ and a peroxo complex, (OP),CuH,O;, but not OH'. The peroxo complex oxidises formate with a rate constant in the region of 3 x 10' dm3 mo1-l s-l, but (OP),Cu3+ does not. Instead it is consumed by reacting quantitatively with (OP),Cu+, thus terminating what would otherwise be an efficient chain mechanism. The peroxo complex decomposes to (OP),Cu3+. and differs from OH' in its reactivity with formate, (OP),Cu'+, and phenanthroline. Fenton-like reactions of Cu' are generally assumed to produce OH': Cu' + H,O, + OH' + OH- + Cu2+ (1) thus initiating the oxidation of organic' and biochemica12q3 substrates by a chain mechanism, as exemplified below for methanol : (2) (3) Evidence for a different mechanism has been found by Johnson and Nazhat,*v5 who showed that the chain yield of formaldehyde resulting from the radiolysis of methanol - H,O, - Cu+ - N,O solutions is much smaller than would result from this mechanism, assuming termination by combination of OH'. They suggested that the product of reaction ( 1 ) is not OH', but Cu3+ (also produced by OH'-induced oxidation of CU")~ reacting as follows: OH' + CH30H + CH,OH' + H,O CH,OH' + Cu2' + HCHO + CU+.Cu++ H,O, + Cu3+ + 20H- CU' + cu3+ + 2cu2+. (4) ( 5 ) (6) If k , is much less than the known high value of k , (8 x lo8 dm3 mol- s-'),' then reactions (3)-(6) account for the small ?hain yields of HCHO, and for their observed proportionality to [CH,OH]/(dose rate)z.Later studies have employed 1 : 10 phenanthroline complexes (OP),Cu2' and (OP),Cu+ to prevent the dismutation of non-complexed Cu+ in oxygen-free solution : 2cu+ -+ c u + CU". Cu3+ + CH,OH -+ CH,OH' + H' + Cu2+ 883884 Fenton-like Reactions of Copper Kinetic studies of the decay of [(OP),Cu+] produced by pulse radiolysis of methanol or formate solutions containing (OP),Cu2+ and H,O, also led to the conclusion that OH' is not formed,8 although the oxidation mechanism has not been fully established, and might involve a peroxo complex [(OP),CuH,O~], rather than (OP),Cu3+. Stopped-flow studies of the decay of [(OP),Cu+] in mixtures of (OP),Cu+ with H,O, and alcohols or formate have led to broadly similar conclusion^.^ The mechanism of formate oxidation is simpler than that of alcohols because the C0;- radical produced in reactions analogous to (2) or ( 5 ) rapidly reduces Cu2+: c0;- + (OP),Cu2+ -+ (OP),Cu+ + co, (7) with k , > 1.5 x lo8 dm8 mol-1 s-l for non-complexed Cu2+,l0 and > lolo for (OP),Cu2+ as shown below.Further, it does not oxidise CU+.~, In contrast, CH,OH' derived from methanol is reported to oxidise (OP),Cu+ 80 times faster than it reduces (OP),Cu2+ in one study,s and more than 30 times slower in a n ~ t h e r . ~ Studies of the yields of formate oxidation to CO, in these systems should therefore be enlightening. Goldstein and Czapski reported CO, yields in radiolysis experiments,8 but these were inferred from measurements of H,O, consumption and are therefore subject to considerable uncertain ties.This paper reports CO, yields for (OP),Cu+ - H,O, - formate solutions in which (OP),Cu+ is produced either by radiolysis at steady-state concentrations estimated to be < loe8 mol dmP3, or is mixed with the other reagents at initial concentrations up to 8 x mol dm-3. The results display unexpected features quite unlike those noted above for methanol solutions, and provide new information on the reactivity of trivalent copper. Experimental Solutions of (OP),CuSO, were prepared from CuSO, - 5H,O (May and Baker) and 1 : 10 phenanthroline hydrate (B.D.H.) with water from a Milli-Q purification system. For y-radiolysis experiments these were mixed with HCOONa (Merck), 6 mmol dmP3 phosphate buffer salts (B.D.H.) to give pH 7, and H,O,. These were bubbled with nitrogen, collected in glass syringes and irradiated with y-rays from a 6oCo source, usually for 2 min at 33 Gy m-', determined by Fricke dosimetry.Prior irradiation of the water did not affect the results. For all experiments involving copper salts the results were identical with H,O, made from (a) B.D.H. AnalaR solutions containing inhibitor, (b) Merck Selectipur solutions, unstabilised, and (c) oxygen-saturated water, pre- irradiated to yield 360 pmol dmP3 H,O,. Radiolysis of formate - H,O, solutions free of copper and buffer salts gave reproducible results with (c) but not (a). All other reagents were of AnalaR grade. For experiments in which performed (OP),Cu+ was mixed with formate - H,O,, [(OP),Cu],SO, was prepared by irradiating glass syringes containing nitrogen-bubbled solutions of (OP),CuSO, (usually 0.26 mmol dm-3) in buffered HC0,Na (usually 26 mmol dm-3) to ca.260 Gy, thus producing typically 115 pmol dm-3 (OP),Cu+. The product was bubbled with nitrogen for > 1 h to remove radiolytic CO,, and analysed for (OP),Cu+ from its optical absorption at 410 nm measured in a flow-through cell, adopting the extinction coefficients reported below. 5 cm3 portions were collected in 10 cm3 glass syringes and rapidly mixed with 8 cm3 nitrogen-bubbled H,O, solution (sometimes containing formate) in a second syringe, connected by plastic tubing. Mixing was complete in ca. 0.1 s. In some cases a screw device was used to mix the two solutions slowly over 2 min (with magnetic stirring), to simulate continuous generation of (OP),Cu+ as in radiolysis experiments.In such cases formate was included in both solutions. Spectra of (OP),Cu+ were obtained by reducing oxygen-free solutions of (OP),CuSO, with excess ascorbate. These altered with addition of formate, givingH. C. Sutton 885 E,,, = 5000 dm3 cm-l mol-1 at 410 nm for 26 mmol dm-3 formate, 4800 at 407 nm for 112 mmol dme3, and 4500 dm3 cm-l mol-1 at 402 nm for 400 mmol dm-3. Johnson and Nazhatg report similar behaviour for less concentrated (OP),Cu+, with extinction coefficients ca. 2 '10 smaller, where conditions are comparable. Solutions were analysed for CO, by gas chromatography as described previously, l1 with reproducibility of 4 1 YO standard deviation for calibrations, and f 1 pmol dm-3 sensitivity for 10 cm3 samples at low [CO,].Nitrous oxide interfered with this analysis and therefore could not be used in radiolytic experiments. Hydrogen peroxide was analysed colourimetrically with a previously calibrated titanium reagent. Results and Discussion Radiolysis of Formate - H,O, Solutions without Copper Radiolysis of these air-free solutions provides a necessary comparison with later experiments. Since the results are likely to be influenced by traces of metals occurring in buffer salts etc., they were performed without buffer and with minimal sodium formate (2 mmol dm-3, ca. pH 6), with H,O, made by preirradiating oxygen-saturated water. The [H,O,] in such solutions declined exponentially with increasing dose, giving [CO,] equal to that of [H,O,] consumed, in yields considerably exceeding those lof radiolytic Cog-, and which increased for a given dose in proportion to (dose rate)-l.Radiolysis of dilute aqueous solutions produces OH' (G = 2.75), H' (G = 0.65) and e;; (G = 2.65) with the total radical yield G, = 6.1.12 These radicals are rapidly and quantitatively converted into CO, as follows : OH'(H') + HCOi- -+ C0;- + H,O(H,) (8) (9) The following subsequent reactions [in addition to reaction (8)] account for the (10) 2CO;- -+ products. (1 1) (12) e& + H,O, --+ OH' +OH-. following observations : Cog- + H,O, -+ CO, + OH' + OH- This leads to F = [H,O,I,/[H,O,I, = 1 - [CO,I,/~H,O,I, = exp [ - k,, t(G, I/2kl,):] for t s irradiation at dose rate I, producing [Cog-] at the rate G, I = 0.38 pmol dm-3 s-' for the high-dose-rate data reported in fig.1. The two linear plots of log F us. t vary 3.5 fold in slope for a 12.5 fold decrease of dose rate, in accordance with eqn (12). The lag or induction period at small doses up to 3 Gy is prpbably due to initial fonsumption of impurities. From these Cata we obtain k10/(2k11)5 = 17 (mol drnp3)t s-5, in reasonable agreement with k,,/(k,,)~ = 22 reported by Buxton and Wilmarth13 from the photolysis of carbon monoxide - H,O, solutions. Since k,, = 7.6 x 10' dm3 mol-1 s-',14 we conclude that k,, x 5 x lo5 dm3 mol-1 s-l. Approximately the same variation of F with dose at high dose rate was found in later experiments with B.D.H. AnalaR H,O, at initial concentrations up to 2 mmol dm-3, but yields at low dose rate were less than expected, and not reproducible. Addition of 10 mmol dmP3 buffer salts (pH 7) or of 2 pmol dm-3 FeSO,, did not greatly influence the results, but traces of CuSo, or of (OP),CuSo, certainly did.Radiolysis of Oxygen-free Formate - (OP),Cu*+ Solutions The spectrum of (OP),Cu+ produced by radiolysis was identical with that resulting from ascorbate reduction of (OP),Cu2+, provided [(OP),Cu2+]/[HCO;] < 0.05. At higher886 Fenton-like Reactions of Copper 0.01 I 1 I 1 300 600 900 irradiation time/s Fig. 1. Radiolysis of oxygen-free solutions of H,O, (0.36 mmol dm-3, made by irradiating oxygenated water) and HC0,Na (2 mmol dm-3): 0, 36 and 0, 2.88 Gy min-'. values there was a noticeable change in the spectrum, attributed to OH' attack on the phenanthroline species. (OP),Cu' is produced during radiolysis by reduction of (OP),Cu2+ with radiolytic CO,, and by the reaction (13) which competes with reaction (9) for eiq, giving G[(OP)2C~+]ca,~.(, = G,.= 6. I . Yields measured at small doses from the emax values given above were slightly less, namely 5.6 & 0.3 for 0.3 mmol dmP3 (OP),Cu2+ in 10 mmol dm-3 HCO,. Consumption of (OP),Cu+ by a fraction of the radiolytic H,O, (G = 0.7) accounts for this. This fraction is estimated to be ca. 0.3 from the mechanism and data given later. Yields observed with 0.1 mol dm+ formate were ca. 5 % higher, probably because of a small increase of G,. at this solute concentration. l2 Copper(1) ions are reported to form a 1 : 1 complex with phenanthroline,'j but irradiated HCO;(OP),CuSO, solutions were unstable after radiolysis, slowly pre- cipitating metallic copper.e;; + (OP),Cu2 -+ (OP),Cu+ Radiolysis of (0P),Cu2+ - HCO, - H202 Solutions This produced CO, in yields proportional to radiation dose which were independent of [H,O,] in the range 0.3-3 mmol dm-3, independent of phosphate buffer concentration in the range 3-20 mmol dm-3, and nearly independent of dose rate, increasing by 3 2 % for a 12.5-fold decrease of dose rate. Variations of CO, yields with [HCO,] at 10 pmol dm-3 (OP),Cu2+, and with [(OP),Cu2'] at low [HCO;], are shown in fig. 2 and 3(a), respectively. Values of G(CO,)/G, inferred by Goldstein and Czapski8 from H,O, consumption measurements in the radiolysis of similar solutions saturated with N,O ranged from 2.1 to 4.65 over a comparable range of [HCO;]. These tended to increase with [HCO;] and with [H,O,], but not in a well defined manner because of a considerableH.C. Sutton 887 3 2 n N 1 0 Fig. 2. Effect of [HCO,Na] on CO, yields in radiolysis of oxygen-free H,O, (0.4mmol dmP3)- copper solutions (dose = 66 Gy, producing 42 pmol dmd3 radicals at G, = 6.1): 0, a, 10 pmol dm-3 (OP),CuSO,, phosphate buffered at pH 7, at 0 , 3 3 and a, 2.64 Gy min-' (----) Calculated from eqn (25) in text, (. - . - . -) calculated including reaction (26) (see text); 0, 2 pmol dmP3 CuSO,, buffer-free, at 33 Gy min-'. 0 , 1000 1 10 100 [(OP), cu2 '1 pml dm3 I I 1000 10 100 [phenanthroline] pmol dm3 Fig. 3. Effect of [(OP),Cu+] and of [phenanthroline] on CO, yields in radiolysis of oxygen-free H,O, (0.4 mmol dm-3)-HC0,Na~OP),CuS0, solutions at pH 7.(66 Gy at 33 Gy min-'). (a) [HCO,Na] = 10 mmol dm-3 0 and 1 mmol dm-3 a; (b) [(OP),CuSO,] = 2 pmol dm-", [HCO,Na] = 10 mmol dm-3 (0) and 1 mmol dm-3 (W).888 Fenton-like Reactions of Copper spread of data. My measurements of H,O, consumption (data not shown) slightly exceeded those calculated from the CO, yields; a better check on stiochiometry is reported later for mixed reagents of known concentration. These CO, yields exceed those arising from primary radiolytic radicals by reactions (7)<9) owing to a chain mechanism initiated by reaction of (OP),Cu+ with H,O,. If all the radiolytic OH' and H' reacts with formate to produce COi-, and all the e& reacts via reaction (9) rather than reaction (13), then the primary radiolytic yields of COi-, and of CO, and (OP),Cu+ derived from it in reaction (7), will all be G, = 6.1 (equivalent to 42 pmol dmP3 for 2 min irradiation).Literature values of k, and k137 indicate that this situation should be approached when [H,0,]/[(OP),Cu2+] > 20. CO, yields arising from the chain process are then given by G(CO,),,,,, = G(CO,),,,,,-G,, and are seen in fig. 2 to increase with [HCO;] to a limiting value of G(C02)ch,i,/G, = G(C02)ch,i,/G{(OP)2C~+) = 2.75 - 1 = 1.75. The chain mechanism will be discussed later; the following section considers complicating factors which arise at high [(OP),Cu2+]/[HCO;], and with excess phenanthroline. Effect of (OP)2Cu2+ and Excess Phenanthroline At high [(OP),Cu2+]/[HCO;], radiolytic OH' attacks the phenanthroline moiety in (OP),Cu2+, as noted above.This decreases the primary radiolytic yield of CO, and presumably also that of (OP),Cu+, which initiates the chain mechanism. Fig. 3(a) illustrates how the total CO, yield varies with [(OP),Cu2+] at two formate concentrations. Fig. 3 (b) shows the corresponding variation with excess [phenanthroline] at 2 pmol dm-3 [(OP),Cu2+]. The unexpected increase of CO, yield as [(OP),Cu2+] is raised from 2 to 10 pmol dmW3 will be considered later. Subsequent decreases at higher [(OP),Cu2+] are attributed to reactions (14) and (15) in competition with reaction (8): (14) (15) A simple treatment of this competition is possible if it assumed that all primary radicals are converted into OH', and the effect of the corresponding competition for the chain carrying oxidant between reactions analogous to reactions (14) and (1 5) is ignored.This leads to OH' + (OP),Cu2+ -+ products, consuming OH' OH' + phenanthroline + oxidation products, consuming OH'. where G(C0,) in the absence of reaction (14) is denoted as G'(C0,). The data in fig. 3 (a) conform approximately with this relation, giving k14/k8 z 4, and hence k,, z 1 x 1 O ' O dm3 mol-1 s-' from literature values7 of k, (2.5 x lo9 dm3 mol-' s-l). Similar treatment of fig. 3 (b) indicates k15/k, z 10, and hence k15 = 2.5 x 1O1O dm3 mol-1 s-l, as compared with a literature value of 8.4 x lo9 dm3 mol-' s-l.16 Clearly the treatment is very approximate, but it shows k15/k14 = 2.5, and demonstrates that [(OP),Cu2+] has negligible influence on G(C0,) when [(OP),Cu2] < 0.001 [HCO,], which applies to all the data in fig. 2 except at the lowest [HCO,].At this point, where [HCO;] = 1 mmol dm-3 and [(OP),Cu2+] = 10 pmol dm-3, G(C0,) is estimated from eqn (16) to increase by 4 YO (equivalent to 3 pmol dm-3 CO,) when reaction (14) is suppressed. A further effect occurs with 2 pmol dmV3 (OP),Cu2+, causing unexpectedly small yields of CO, which are restored to the values expected in the absence of reaction (14) by adding 5 pmol dm-3 phenanthroline [cf. fig. 3 (a) and (b)]. For formate concentrations of 10 (but not quite for 1) mmol dm-3, the same restoration is achieved by increasing [(OP),Cu2+] to 10 pmol dm-3. A similar effect was found in studies of the chain reactionH . C. Sutton 889 1.5 1 .o R 0 . 5 / I I 10 100 [Hc02-]/nnnol d d Fig.4. Effect of [HCO,Na] and initial [(OP),Cu+] on R (= [CO,],/[(OP),Cu+],) for rapidly mixed oxygen-free solutions of H,O, (0.5 mmol dm-3, HCO,Na, (OP),Cu),SO, and (OP),CuSO, at pH 7. (----) Calculated from mechanism in text, ( - - . -. -) calculated including reaction (26) (see text). Main graph: initial [(OP),Cu+] = 45 pmol dmW3, initial [(OP),Cu2+] = 55 pmol dm-3. For point ., solutions were slowly mixed over 2 m. Inset: [HCO,Na] = 10 mmol dm-3, initial [(OP),Cu+]/mmol dmP3 = 0, 80; 0, 5 M O and A, 120. between H,O, and formate induced by paraquat radicals, catalysed by traces of iron salts chelated with an excess of EDTA." If [EDTA] < 5 pmol dmP3 then this catalysis was very inefficient, even though the iron concentration was < 0.1 pmol dm-3.Adding > 10 pmol dm-3 EDTA restored full catalysis. It was shown that ca. 2 % of the Fenton reactions which initiate this chain consume EDTA in a 'cage'-type process, whilst the remainder produce chain-carrying radicals. A similar process involving consumption of phenanthroline might account for the low yield with 2 pmol dm-3 (OP),Cu2+. It may be concluded that G(C0,),,,,,/G[(OP)2Cu+] obtained from fig. 2 gives a true measure of the effect of [HCO,] on chain yields of CO,, except for the 4% effect at 1 mmol dm-3 CO,. Further, when reaction (14) is suppressed, then G(C0,) is independent of [(OP),Cu2+] in the range 2-100 pmol dmF3. This implies that reaction (7) wholly suppresses the OH'-producing reaction (10) over this range, and hence (from the value of k,, inferred above) that k, probably exceeds 10" dm3 mol-l s-'.Data for solutions containing 2 pmol dmP3 CuSO, instead of (OP),CuSO, are shown for comparison in fig. 2, and are seen to rise with increasing [HCO;] above a value which approximates G, at [HCO;] = 1 mmol dm-3. Yields under these conditions increased slightly with [H,O,], and were influenced by phosphate buffer salts. They have not been investigated further. Mixtures of (OP),Cu+ and Formate with H,O, These produced CO, in yields which were independent of initial [H,O,] in the range 0.2-2 mmol dm-3 and independent of phosphate buffer concentration in the range 3-20 mmol dm-3. Similar but less reproducible yields were obtained without buffer at ca. pH 6. Yields varied with [HCO,], initial [(OP),Cu+] and rate of mixing as shown in fig.4. These are expressed as yield ratios R = final[CO,]/initial[(OP),Cu+], which tend at high [HCO,] to a limiting value of 1.74 0.09, equal to G(CO,),,,i,/G {(OP),Cu+) in the radiolysis experiments, as inferred above from fig. 2. Slow mixing over 2 min increased R to 1.43 -10.08 for 10 rnmol dm-3 formate, which agrees with that found in radiolysis890 Fenton-like Reactions of Copper experiments for G(CO,),,,,,/G[(OP),Cu+] at the same [HCO;], namely 1.45 0.03 R values were independent of [(OP),Cu2+] within 5 % experimental error over a threefold range, provided [(OP),Cu2+]/[HCOJ < 0.01, which applies to all the data in fig. 4. Yields were determined by the total [HCO;] in the mixed solution, regardless of whether this was wholly in the copper solution before mixing, or mostly in the H,O,.The reaction with 45 pmol drn+ (OP),Cu+ and 10 mmol dm-3 HCO, was complete after rapid mixing in ca. 5 s for 0.2 mmol dm-3 H,O, (as shown by decay of colour), and in less time at higher [H,O,], yielding CO, in concentrations which were measured 1 min after mixing and were usually constant for 30 min. Corrections were made for slight time-dependent increases sometimes found at high [H,O,] and [(OP),Cu2+], and also for CO, in the premixed reagents (usually < 5% of the reaction yield). H,O, was consumed during the reaction in yields given within experimental error - A[H,O,] = 0.5 initial [(OP),Cu+] +final [CO,]. (fig. 2). by These correspond with overall reactions expressed stoichiometrically as H,O, + 2(OP),Cu' --+ 2(OP),Cu2+ + 20H- and H,O, + HCO; -+ CO, + H,O + OH-.R values were decreased by large [(OP),Cu2+] and by excess phenanthroline, indicating that the chain carrying oxidant is consumed in reactions analogous to (14) and ( I 5). For 2.3 mmol dm-3 formate, R was decreased from 0.65 f0.03 to 0.42k0.02 when ([(OP),Cu2+] + [(OP),Cu+]} was increased from 0.1 to 1.02 mmol dmP3. For 10 mmol dm-3 HCO;, R decreased by 6 + 5 O/O when the total copper concentration was raised from 0.1 to 0.52 mmol dm-3. It is estimated from these data that the chain oxidant reacts with (OP),Cu2+ ca. 1.3 times faster than with formate, as compared with a corresponding factor of 4 estimated above for OH'. Increasing the total [phenanthroline] to 3 times the total [copper] [thus forming (OP),Cu2+, and possibly also (OP),Cu'] did not significantly alter the R values in fig.4. An excess of phenanthroline at 0.31 and 0.71 mmol dm-" decreased observed R values by 15+5 and 35+6%, respectively, for 3 mmol dm-3 formate, and by 2 f 5 and 12 + 5 O/O for 30 mmol dm-3 formate. These data indicate that the chain-carrying oxidant reacts with phenanthroline ca. 2 times faster than with formate and hence ca. 1.5 times faster with phenanthroline than with (OP),Cu2+. The corresponding factors estimated above by very approximate methods for OH' were 10 and 2.5, respectively. Reaction Mechanism The large CO, yields found in the absence of copper show that OH' formation propagates an efficient chain reaction in the system studied. If the reaction of (OP),Cu+ with HzOz p-roduces OH' with k z 2.5 x lo3 dm3 mol-' S-' (half the value found for Cu+ consumption in absence of organic substrate^),^, l5 then the fast reactions (7) and (8) would propagate a chain sequence terminated in mixed solutions by reaction of (OP),Cu+ with OH', and in radiolysis experiments at high [HCO;] by reaction (1 1).In the former case this would give CO, yields nearly equal to the initial [H 0 ] in < 1 min, whilst radiolysis would lead to G(CO,)/G, = k,[(OP),Cu2+]/(2kllGr~~ ihich exceeds 1000 for 10 pmol dmP3 (OP),Cu2+. In contrast, the observed CO, yields were independent of [H,O,], whilst G(CO,)/G, was independent of dose rate and < 3 in all cases. Further, no chain reaction occurred [G(CO,) = G,] in CuSO, solutions at [HCO;] high enough (1 mmol dm-3) to react quantitatively with OH'.This evidence strongly supports earlier conclusions4~ 5 3 '* that Fenton-like reactions of copper(r) ions do not produce OH'.H. C. Sutton 89 1 If, instead, the Fenton product is (OP),Cu3+ reacting as suggested for methanol'*' (17) (18) (7) (19) d[CO,]/ -d[(OP),Cu+] = kl,[HCO~]/k19[(OP)2C~+] (20) in the sequence (OP),Cu+ + H,O, -+ (OP),Cu3+ + 20H- (OP),Cu3+ + HCO, -+ CO, + H+ + (OP),Cu2+ (OP),CU2+ + C0;- -+ co, + (OP),Cu+ (OP),CU+ + (OP),CU3+ + 2(OP),Cu2+ then this mechanism requires at each value of [(OP),Cu+] as the reaction proceeds in mixed solutions. It predicts CO, yields which are directly proportional to [HCO,], and which increase by a constant concentration for each tenfold decrease of [(OP),Cu+], and for each tenfold increase of reaction time, until all H,O, is ultimately consumed.In fact both [HCO,] and [(OP),Cu+] have only a limited range of influence, the first saturating at 0.1 mol dm-3, and the second showing that chain termination occurs by a process independent of [(OP),Cu+] at the low concentrations in slow mixing experiments, and also in radiolytic experiments, with yields independent of dose rate (fig. 2 and 4). These observations can be explained if the chain-carrying oxidant is not (OP),Cu3+ but a peroxo complex of (OP),Cu+, produced and reacting as follows: (OP),Cu+ + H,O, -+ (OP),CuH,O; (21) (22) (23) (24) (OP),CuH,Ol -+ (OP),Cu3' + 20H- (OP),CuH,Ol + (OP),Cu+ + 2(OP),Cu2' + H,O,. (OP),CuH,Oi + HCO, -+ COh-+ H+ + 2 0 W + (OP),Cu2+ The overall chain mechanism in mixed solutions then consists of reactions (17), (21H24), (7) and (19), but excludes reaction (18).Its essential feature is that (OP),Cu3+ formed in reactions (17) and (22) is quantitatively removed in reaction (19), independently of [(OP),Cu+]. Thus the chain is propagated by reactions (21) and (23), and terminated by reactions (1 7), (22), (1 9) and (24). In radiolysis experiments reaction (24) is negligible, and C0;- is generated in a yield of G,. Ignoring reaction (14) for simplicity (since it contributes only partially at 1 mmol dm-3 formate), the mechanism then leads to G(CO2) = G, + 0.5G,{k,i/k21 + k22/(k,3[HC0,1) + k,, k~3/(~27 k23 [Hco,l>>-l. (25) The interrupted curve in fig. 2 has been calculated from this relation, adopting k1i/k21 = 0.288 and k22/k23 = 3.69 x mol dmW3.Agreement with the data is satisfactory. The tendency for observed yields to exceed those calculated by ca. 5 % at 0.1 mol dm-3 formate is attributed to a corresponding increase of G, under these conditions, as indicated by increased values of G{ (OP),Cu+). Reaction (24) occurs at the high values of [(OP),Cu+] in mixed solutions, thus decreasing R below G(CO,),,,,,/G{(OP),Cu+}. A computerised approach has been used to evaluate concentrations of reactants and products as a function of time under these conditions, and thus calculate R values from the mechanism. Rate constants adopted for this purpose are listed in table 1. Provided k,,/kz1 and k22/k23 are maintained at the values given above, and the additional parameter required to fit fig.4 is assigned the value k,,/k,, = 2.71 x then calculated values of R are independent of the absolute values of the seven rate constants listed over a wide range. The value k,, = 1 x lo9 dm3 mol-1 s-' appears reasonable, in892 Fenton-like Reactions of Copper Table 1. Reaction rate constants adopted for calculations rate constant rate constant reaction dm3 mo1-' s-' reaction dm3 mol-1 s-' 7" 1 x 1O'O 21" 5 x 103 1 76 1.44 x 103 22' I x 104 (s-1) 18" < 1x103 23d 2.17 x 107 19' 3~ 109 24f I x 109 * Estimated approximately in the text. Based on k17/k21 = 0.288, Based on k,,/k,, = 2.96 x mol dmP3. Based on k,,/k,, = 2.71 x lo-,. Assumed, but unimportant since eqn (19) is quantitative. Assumed. which case k,, = 2.7 x lo7 dm3 mol-1 s-'.The value k,, = 5000 dm3 mol-' s-' has been assigned to agree with the only measurement of overall reaction rate possible, namely the very approximate estimate of total reaction time (5 s) for mixed solutions at 0.22 mmol dmP3 H,O,. Calculated values of R are shown as interrupted curves in fig. 4 and agree with the data, except at the lowest [HCO;] studied. Rates of decay of [(OP),Cu+] in the mixed solutions may be calculated by computerised methods from the mechanism above. For [(OP),Cu+] < 3 pmol dm-, these are proportional to [H,O,] [(OP),Cu+] with apparent rate constants decreasing from 5.7 to 2.7 dm3 mol-1 s-l as [HCOJ is increased from 1 to 100 mmol drn-,. Decay kinetics of this form were found in pulse radiolysis studies of similar solutions saturated with N,O,, but with constants of ca.1.3 dm3 mol-1 s-l which increased only slightly with increasing [HCO;]. Good absolute agreement is not expected because of my very approximate estimate of k21, but an increase of reaction time with increasing [HCO;] was clearly evident. Evidently the proposed mechanism accounts reasonably for the observations, with three adjustable parameters. No doubt additional reactions also occur to a minor extent, such as (OP),CuH,Oi -+ OH' + OH- + (0P),Cu2+ suggested by Goldstein and Czapski,' followed by reaction (8). If reaction (26) is included with k,, = 1.5 x lo4 dm3 mol-1 s-l, then a slightly better fit to the data in fig. 2 and 4 is obtained, as shown by the dotted curves. However, the exclusion of reaction (1 8) is the paramount requirement; agreement with the data is unsatisfactory if k,, > lo3 dm3 mo1-I s-'. It is recognised that two-electron transfer reactions such as reaction (1 7) are rare, and that formation of the (OP),CuH,Oi complex from (OP),Cu+ may pose coordination problems..f' Nevertheless, the evidence now available shows that reaction (1) does not proceed by a simple one-electron transfer.Its products react much less rapidly with formate and methanol than does OH', and in the case of (OP),Cu+ they include two species, one of which decomposes to the second of very low reactivity, which is almost certainly (OP),Cu3+. A report published after submission of this paper arrives at similar conclusions by thermodynamic arguments, indicating that the primary copper product is analogous to (OP),CUH,O~.~' The biological significance of these and analogous FeIV species has been noted.18 I gratefully acknowledge Dr M. Manning's assistance in providing the computer program which made these calculations possible. I thank Miss D. Chambers for technical assistance and Dr P. €3. Roberts for helpful discussions. t I am indebted to a referee for pointing this out.H. C. Sutton 893 References 1 G. V. Buxton and J. C. Green, J . Chern. Soc., Faruduy Trans. I , 1978. 74, 697. 2 D. A. Rowley and B. Halliwell, Arch. Biochern. Biophys., 1983, 225, 279. 3 B. G. Que, K. M. Doweny and A. G. So, Biochemistry, 1980, 19, 5987. 4 G. R. Alistair Johnson, N. B. Nazhat and R. A. Saadalla-Nazhat, J. Chern. Soc., Chern. Cornmun., 5 G. R. Alistair Johnson, N. B. Nazhat and R. A. Saadalla-Nazhat. J . Chem. Soc., Faraday Trans. I , 6 D. Meyerstein, Inorg. Chem., 1971, 10, 638. 7 Farhataziz and A. B. Ross, Natl. Stand. Ref. Data Ser. 1977, 59. 8 S. Goldstein and G. Czapski, J . Free Radicals in Biology and Medicine, 1985, 1, 373. 9 G. R. Alastair Johnson and N. B. Nazhat, J. Am. Chern. SOC., 1987, 109, 1990. 10 Y. Ilan, Y . A. Ilan and G. Czapski, Biochim. Biophys. Acta, 1978, 503, 599. 11 H. C. Sutton and C. C. Winterbourn, Arch. Biochem. Biophys., 1984, 235, 106. 12 H. A. Schwarz, J . Chem. Ed., 1981, 58, 101. 13 G. Buxton and W. K. Wilmarth, J. Phys. Chem., 1963, 67, 2835. 14 G. V. Buxton and R. M. Sellers, J . Chem. SOC., Faraday Trans. I , 1973, 69, 555. 15 K. V. Ponganis, M. A. de Araujo and H. Leslie Hodges, Inorg. Chem., 1980, 19, 2704. 16 E. Siekierska-Floryan, Nucelonica, 1979, 24, 951. 17 M. Masarwa, H. Cohen, D. Meyerstein, D. L. Hickman, A. Bakac and J. H. Espenson, J . Am. Chem. 18 H. C. Sutton and C. C. Winterbourn, Free Radical Biology and Medicine, 1988, in press. 1985, 407. 1988, 84, 501. SOC., 1988, 110, 4293. Paper 8101789A: Received 6th May, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500883

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1. 1989, 85(4), 895-906 Structural Investigation on a Spinel-related Zn/Cr = 1 Mixed-oxide SystemT Cinzia Cristiani and Pi0 Forzatti" Dipartimento di Chimica Industriule ed Ingegneria Chimica ' G. Natta' del Politecnico - Piazza Leonurdo da Vinci, 32, 20133 Milano, Itall? Maurizio Bellotto Dipartirnento di Ingegneria Meccunicu, Universita di Brescia - Via Vulotti, 9, 25060 Brescia, Italy A Zn-Cr mixed oxide of 1 / 1 composition has been characterized by powder X-ray diffraction, infrared spectroscopy, ultraviolet-visible diffuse reflec- tance spectroscopy, differential thermal analysis-thermal gravimetry, surface-area determination and chemical analysis. The behaviour of the system has been studied over a large temperature range and under different activation atmospheres.In addition to the model involving a single non- stoichiometric spinel phase, already proposed in the literature, a model involving a shell rich in ZnCrO, and a core constituted by ZnCr,O, and ZnO is discussed and shown to be consistent with the experimental evidence. Commercial high-temperature catalysts for the synthesis of methanol and mixtures of methanol and higher alcohols from CO and H, are usually based on pure and alkali- promoted mixed oxides of Zn and Cr, Most of the commercial catalysts present Zn/Cr ratios > 0.5. The catalytic activity has generally been attributed to a promoting effect of Cr because of physical effects related to the formation of ZnCr,O,, which either prevents the sintering of ZnO (regarded as the active phase)3 or acts as a support for ZnO.*- Alternatively the catalytic activity has been attributed to the ZnCr,O, spinel phase.' Recently a number of papers7-13 have appeared dealing with the structure and reactivity of Zn and Cr mixed oxides.In these catalysts the active species was reported to be a spinel-like phase, non-stoichiometric because of an excess of Zn. This excess of Zn was reported to be octahedrally coordinated, and to provide the catalytically active ~ i t e s . ~ * * * l ~ For an understanding of the nature of such a system the most interesting composition was found to be Zn/Cr = 1.' In fact this composition is reported to exhibit typical non-stoichiometry at the maximum degree, being monophasic, while at higher Zn loadings free ZnO is detected.However, the behaviour of the system, particularly at high calcination temperatures and under different activation atmospheres, was not fully clarified. The aim of this paper is to investigate the nature of a Zn/Cr = 1 mixed oxide. A number of techniques were used to obtain information on the valence state of the elements and their local coordination and on the structure of the present phases. The evolution of the sample has been followed in air, N, and H, up to temperatures higher than those reported by other authors. The target was to provide a better understanding of the phase transformations occurring in the system, during both activation and t This is Paper XV in the series Synthesis of Alcohols from Carbon Oxides and Hydrogen. 895896 Structural Investigation of a Zn-Cr Mixed Oxide catalytic reaction.A comparison with literature data is provided. A new model is discussed which is capable of interpreting the experimental evidence collected. Experiment a1 Preparation and Activation A Zn-Cr mixed-oxide sample with an atomic ratio Zn/Cr = 1/1 was prepared by dissolving Zn(NO,), * 6H,O and Cr(NO,), * 9H,O (Carlo Erba RPE). Zn and Cr were precipitated by addition of a 3 mol dm-, solution of (NH,),CO, at 333 K. Evolution of CO, was observed. The final pH of the supernatant liquor was 7.8. The resulting slurry was aged for 12 h at 353 K and then filtered and washed to remove NO, ions from the precipitate. The presence of NO; in the eluate was ascertained by qualitative analysis. After drying at 383 K a precursor was obtained, quoted as Z,C1-4A.Indeed all the Zn and Cr was precipitated, as confirmed by chemical analysis performed on the sample after calcination at 673 K. Throughout the paper samples are identified with a notation indicating the activation temperature and the corresponding reaction atmosphere : e.g. Z,C,-llA indicates sample Z,C, heated at 1073 K in air. Sample ZlC,-4A was activated in air at temperatures extending from 673 to 1073 K, and in N, and H, atmospheres up to 673 and 773 K. Details of the temperature cycles of activation are summarized in table 1. Characterization The samples were characterized by means of powder X-ray diffraction, infrared spectroscopy, ultraviolet-visible diffuse reflectance spectroscopy, differential thermal analysis-thermal gravimetry, surface-area measurements and chemical analysis.X-Ray powder diffraction patterns were collected using a Philips PW 1050/70 vertical goniometer with nickel-filtered Cu Ka radiation. Crystallite dimensions were evaluated using the Sherrer equation. Lattice parameters were calculated with a least-squares fitting routine, taking into account systematic errors produced by the goniometer. Quantitative estimates of multiphase samples were performed according to Klug and A1e~ander.I~ Infrared spectra were recorded on a Perkin-Elmer 147 spectrometer with the KBr pressed-disc technique. Ultraviolet-visible diffuse reflectance spectra were recorded by means of a Jasco UVIDEC-6 10 double-beam spectrophotometer equipped with integrating sphere (bandwidth 2 nm).Surface areas were determined by using a Carlo Erba Sorptomatic 1800 series instrument with a B.E.T. dynamic system. Differential thermal analysis-thermal gravimetry data were obtained with Mettler 2000 instrument. The chemical analysis was performed in order to determine the fraction of Cr present as Crv' in Z,C,-7A. 0.5 mol dm-, H,SO, was added at a quoted portion of the sample, under stirring and boiling for ca. 6 h (without evaporation of the acid solution) in order to dissolve the CrV1.l5 A yellow solution and a black residue (quoted as ZlC,7AR-4A) were separated by filtration. The amount of Crvl in the solution was determined by manganometric titration. A portion of ZlC,7AR-4A was etched and dissolved with a mixture of aqua regia and H,SO, by boiling for 24 h.The chromium present in the solution was again determined as Crv' by manganometric titration. Z1C,7AR-4A was also characterized as usual. Results Precursor The sample dried at 383 K ZlC,-4A is the matrix from which all subsequent samples have been derived. The X.r.d. powder pattern shows [fig. l(a)] only two very broad reflections owing to its microcrystallinity. In fact this sample has a very high surfaceC. Cristiani, P. Forzutti and M. Bellotto 897 Table 1. Activation-temperature cycles sample temperature cycle ZlC,-7A Z,C,-9A Z,C,-11A ZICl-7N ZlC,-8N Z,C1-7H Z,Cl-8H from r.t. to 673 K, 120 K h-l; at 673 K for 4 11 from r.t. to 673 K, 240 K h-l; from 673 to 873 K, 120 K h-l; at 873 K for 2 h from r.t. to 873 K, 240 K h-'; from 873 to 1073 K, 120 K h-l; at 1073 K for 1 h from r.t.to 673 K, 120 K h-l; at 673 K for 4 h from r.t. to 673 K, 240 K h-l; from 673 to 773 K, 60 K h-l; at 773 K for 3 h from r.t. to 673 K, 120 K h-l; at 673 K for 2.5 h from r.t. to 673 K, 240 K h-'; from 673 to 773 K, 60 K h-'; at 773 K for 2 h 0 0 A . 65 60 55 50 45 40 35 30 25 20 281" Fig. 1. Experimental X-ray powder diffraction patterns of (a) Z,C,-4A, (6) Z,Cl-7A, (c) Z,C,-9A and ( d ) Z,C,-llA, together with calculated patterns of a mixture of ZnCr,O, and ZnO: (i) f.w.h.m. = 2.0, (ii) f.w.h.m. = 0.6 and (iii) f.w.h.m. = 0.2". a, ZnO; 0, ZnCr,O,. area, i.e. = 252 m' g-l, which accounts for its small crystallite dimensions, in line with the broadness of the X.r.d. reflections. The i.r. spectrum [fig. 2(a)] gives evidence of bands characteristic of COi- (840, 1020, 1090, 1370 and 1480 cm-l) and possibly of NH; (1390 cm-').16 The diffuse reflectance spectrum [fig.3(a)] shows bands at 260, 370, 440, 580 and 700 nm. From literature data1' we can attribute the bands at ca. 440, 580 and 700 nm to CrII' d 4 transitions [fig. 4(b)], the band at 370 nm to 02- -+ Zn2+ charge transfer, as in ZnO [fig. 4(c)], and the bands at 370 and 270 nm to 02- + Cr6+ charge transfer. The oxidation of some CrIIl to Crv' is probably caused by aging the898 Structural Investigation of R Zn-Cr Mixed Oxide I 0 I I 1 1 I 1600 1200 MK) 400 wavenumber/cm-' slurry at 353 K in alkaline ~olution'~ and/or drying it at 383 K in air. The d.t.a.-t.g. analysis in N, and air up to 673 K [fig. 5(a) and (b)] shows two peaks (the first, endothermic, at 453-553 K and the second, exothermic, at 553-653 K) attributed to the decomposition and subsequent combustion of an ammoniacal compound [possibly (NH,),CrO,, qecomp = 453 K].The first endothermic peak at 453-553 K also has a possible contribution from the decomposition of Zn,(OH),(CO,), ( &ecomp z 538 K).lS The endothermic peak at ca. 373 K may be associated with dehydration of the sample. From the results so far it follows that the precursor may consist of a mixture of ammonium compounds and carbonates. Activation in Air Upon calcination up to 673 K sample Z1C,-7A is obtained. Its X.r.d. powder pattern shows a spinel-like phase [fig. l(b)], related to ZnCr,O, (JCPDS 22-1107). No other phase, namely ZnO (JCPDS 5-0664), is detectable. However, the broadness and asymmetry of the reflections prevents us from excluding the presence of ZnO.Fig. 1 (i) shows the calculated spectrum of a mixture of ZnO and ZnCr,O,. The Zn/Cr ratio of this mixture, resulting from ZnO and ZnCr,O,, is 1 / 1 and corresponds to that measured for sample Z,C,-1lA. The full width at half maximum (f.w.h.m.) is set to 2.0" in orderC. Cristiani, P. Forzatti and M . Bellotto 899 r 1 '-- I I I 1 1 1 1 ' 0 . m 250 350 450 550 650 750 850 wavelength /n m Fig. 3. Ultraviolet-visible diffuse reflectance spectra of (a) Z,C,-4A, (b) ZlC,-7A, (c) Z1C,-9A and (d) Z,C,-IIA. to simulate the plot in fig. l(6). From a rough visual comparison of fig. 1 (b) and (i) it is evident that the hypothesis that sample ZlC,-7A consists in the bulk of a mixture of ZnO apd ZnCr20, cannot be discarded.Particle-size determination gave a mean value of 50 A. This corresponds to a surface area of 126 m2 g-' (calculated assuming spheres with radius equal to the mean particle size), in line with the measured value of 104 m2 g-'. The lattice parameter of this spinel-like phase has been refined using all the resolved measured reflections, i.e. (1 1 l), (2 2 0), (3 1 l), (4 0 0), ( 5 1 1)/(3 3 3) and (4 4 0), and under the hypothesis that the sample is monophasic, i.e. the assumption was made that the line position is due only to the parameters of the spinel phase a n t is not affected by the overlapping lines of other phase!. The refined value was 8.40(2) A, higher than that reported for ZnCr,O, (a, = 8.328 A).However, in this case, the refined a, value is affected by the poor crystallinity of the sample and by the underlying hypothesis that the sample is monophasic. Because of this we cannot speculate further on this assumption. The i.r. spectrum [fig. 2 (b)] shows two absorptions at 500 and 600 cm-' attributed to the bending and stretching of Cr"'-O octahedra in ZnCr20,1g and a composite absorption in the range 85&930 cm-' attributed to the bending and stretching of CrV'-0 tetrahedra in ZnCrO, (a-phase).,' The diffuse reflectance spectrum [fig. 3 (b)] gives evidence of 02- -+ Zn2+ charge transfer at 370 nm as in ZnO and at 300 nm as in ZnCr,O, [fig. 4(a)]. 02- + Cr6+ charge transfer at 370 and 260 nm are also present while no bands due to CrIII d 4 transitions are evident [fig. 4(a) and (b)].CrI" d-d transitions are not evident in the u.v.-visible diffuse reflectance spectrum, possibly because they are broad and thus smeared out, and/or because they are covered by the low-energy tail of the more intense charge-transfer bands.21 The larger half-width may indicate a greater antibonding character of the excited state compared to the stoichiometric spinel, i.e. a larger internuclear equilibrium distance. The band at 300 nm is not so evident as in the spectrum of ZnCr,O, [fig. 4(a)] because of the presence in sample Z,C,-7A of the 02- --+ Cr6+ charge-transfer absorptions. In contrast, the band at 300 nm is well evident in sample ZlC,-8H, as will be discussed later, owing to the absence of Crvl. On annealing this sample in air at 873 and 1073 K samples ZlC,-9A and Z,C,-1 1A are obtained.Their X.r.d. powder patterns evidence a progressive crystallization of ZnO (reflections at 28 = 31.88, 34.54, 36.62, 47.61 and 56.64") and of ZnCr,O, [fig. 1 ( c ) and (41. Fig. (ii) and (iii) show the calculated spectra equivalent to that already discussed inStructural Investigation of a Zn-Cr Mixed Oxide O.OO0 250 350 450 550 650 750 850 w avelengthhm Fig. 4. Ultraviolet-visible diffuse reflectance spectra of (a) ZnCr,O,, (b) Cr,O, and (c) ZnO. fig. l(i). Direct comparison of calculated and experimental X.r.d. patterns is in agreement with the hypothesis that samples Z,C,-9A and Z,C,-1 1A are constituted by a mixture of ZnO and ZnCr204. The f.w.h.m. are set to 0.6 and 0.2, respectively. Quantitative analysis of the pattern shows that all Zn in excess with respect to Zncr,04 is present as ZnO in sampleo Z,C,-l1A.The mean particle size increases to 100 A for sample Z,C,-9A and 259 A for sample Z,C,-llA. The refinedo lattice parameter decreases to a, = 8.35(1) A for sample Z,C,-9A and a, = 8.327(3) A for sample Zlcl- 11 A, thus approaching the value found for the stoichiometric ZnCr204 (a, = 8.328 A). The refinement was made using all measured reflections in these cases also. The surface- area value, calculated starting from crystallite dimensions, is 66 m2 g-' for Z,C,-9A and 27 m2 g-' for Z,C,-11 A, in line with the measured values of 47 and 19 m2 g-', respectively. The i.r. spectra [fig. 2(c) and (d)] indicate a sharpening of the bands attributed to the bending and stretching of CrII' octahedra and the appearance of a band attributed to ZnO (as a shoulder at 450 cm-').A progressive decrease, on increasing the calcination temperature, of the composite absorption in the range 850-930 cm-' associated with Crv' in a tetrahedral coordination is also evident. The diffuse reflectance spectra [fig. 3(c) and (d)] show an increase in the relative intensities of the absorptions corresponding to CrIII d-d transitions on increasing the calcination temperature. In order to quantify the amount of Crvl in Z,C,-7A a chemical analysis has been carried out. The etching conditions prevent oxidation of the CrIII eventually dissolved along with Crvl. Then manganometric titration detects only Crvl species.15 The Crvl dissolved by H,S04 etching is the 14.5% of the total Cr. This amount is in line with the one reported for a similar sample after etching with diluted CH,COOH.' The X.r.d.powder pattern of the residue (sample Z,C17AR-4A) shows a spinel-like phase whose crystallinity is comparable to that of sample Z,C,-7A. The i.r. spectrum of the residue shows no evidence of the composite absorption extending from 850 to 930cm-' associated with Crvl [fig. 2(e)]. The absorptions at 1150, 1050 and 870 cm-', which are not present in the i.r. spectrum of sample Z,C,-7A, are attributed to SO:- ions. The pair of bands at 500 and 600 cm-' do not differ significantly from those detected in sample Z,C,-7A. Chemical analysis confirms that the insoluble portion contains 85% of the total Cr. Thus the etching of sample Z,C,-7A occurs only on the surface and does not induce any structural modification of the bulk.It dissolves 14.5% of total Cr as Crvl;C. Cristiani, P. Forzatti and M . Bellotto 90 1 z 81 W 8j t 0,1 W 4 673 573 473 373 T/K < 673 573 473 373 T / K Fig. 5. Differential thermal analysis-thermal gravimetry up to 673 K of (a) Z1C,-4A in air and (b) Z1C,-4A in N,. (i) D.t.a., (ii) t.g. and (iii) d.t.g. the residual amount of Cr is present as Crl'* in the insoluble portion. No Crvl is detected in the residue. From the evidence collected so far, the evolution in air of the system under study can be summarized as follows. (1) The decomposition of the ammonium compounds and of the carbonates, constituents of sample Z1C,-4A, gives rise to sample Z,C,-7A.This sample may be constituted by a shell rich in Cr", formed by a phase related to ZnCrO, which contains 14.5 O/O of the total Cr, and by an inner core, formed by a spinel- like phase related to ZnCr,O,, and possibly ZnO. (2) When calcining to higher temperatures, we observe a decrease in the content of Cr", together with the crystallization of the present phases. The refined lattice parameter of the well crystallized spinel phase in sample Z,C,-llA is close to that reported in literature for ZnCr,O,. Activation in N, The samples ZlC,-7N and ZlC,-8N were obtained upon heating Z,C,-4A in N, at 673 and 773 K. The outcome of d.t.a.-t.g. analysis on sample ZlC1-4A in N, [fig. 5(a)] is similar to that previously discussed for sample ZlC,-4A in air. The X.r.d.powder patterns [fig. 6(a) and (b)] are consistent with the reflections of ZnCr,O,. The asymmetries and shoulders at 28 z 30 and 36" may be attributed to the presence o[ZnO, as already discussed. The mean crystallite dimension of sample ZlC,-8N is ca. 25 A. The Cr"'-O stretching at 600 cm-l and bending at 500 cm-l, typical of 11-111 type spinel phases, are detected in the i.r. spectra of the two samples ZlC,-7N and Z1C,-8N [fig. 7 (a) and (b)]. The composite absorption typical of tetrahedrally coordinated Crv' in the range 850-930 cm-' decreases on increasing the temperature from 673 to 773 K. Diffuse reflectance measurements on the sample Z1C,-8N [fig. 8(a)] are consistent with the presence of characteristics absorptions due to 0,- + Zn2+ and 02- -+ Cr6+ charge transfer (260, 340 and 360 nm) and due to Cr"' d-d transitions (450, 600 and 700 nm).Indeed, these absorptions are barely evident in the spectrum, but are seen by taking the second derivative.902 Structural Investigation of a Zn-Cr Mixed Oxide I 1 1 1 I I 1 I I I 65 60 55 50 45 40 35 30 25 20 291' Fig. 6. Experimental X-ray powder diffraction patterns of (a) ZlC,-7N, (b) Z,Cl-8N, (c) Z,C1-7H and ( d ) ZlC,-8H. a, ZnO; O,ZnCr,O,. 0 Fig. 7. Infrared spectra of (a) ZICl-7N, (b) ZlC,-8N, (c) Z,Cl-7H and ( d ) ZlC1-8H. 0, ZnCrzO,; ., ZnCrO,. The activating atmosphere is less oxidizing than air; thus there is a smaller amount of Crv' evidenced by both the i.r. and u.v.-visible diffuse reflectance spectra. In this case also, on increasing the activation temperature, we observed a decrease in the i.r.and u.v.-visible diffuse reflectance absorptions related to Crv*. Moreover, as compared to the samples calcined in air, the samples activated in N, are less crystallized.C. Cristiani, P. Forzatti and M. Bellotto (4 903 C c - 0.500 -8 4 \ - - d 1 I I I 1 0.000 Fig. 8. Ultraviolet-visible diffuse reflectance spectra of (a) Z1C,-8N and (b) Z,C,-8H. I I -1 TIK - 573 473 373 Fig. 9. Differential thermal analysis-thermal gravimetry up to 673 K of sample Z,C,-7A in H,. (a) D.t.a., (b) t.g. and (c) d.t.g. Activation in H, The samples ZlC,-7H and ZlC,-8H are obtained by the activation in H, of ZlC,-4A at 673 and 773 K. The X.r.d. patterns of sample Z,C1-7H [fig. 6(c)] indicate a very microcrystalline sample. Reflections consistent with those of a spinel-like phase are detectable.On increasing the temperature of annealing we observe [fig. 6 ( d ) ] the crystallization of the spinel-like phase and, to a much greater extent, of the ZnO phasz. The mean crystallite dimensions of the spinel-like phase in sample ZlC,-8H is ca. 25 A. The infrared spectra [fig. 7(c) and (d)] show bands at 500 and 600 cm.-l typical of octahedrally coordinated CrIII, but there is no evidence that Crv' is present. The diffuse reflectance spectrum of sample ZlC,-8H [fig. 8(b)] shows absorptions at 300 and 360 nm due to 02- -+ Zn2+ charge transfer, and absorptions at 450, 600 and 700 nm attributed to C P d-d transitions.904 Structural Investigation of a Zn-Cr Mixed Oxide Because of the reducing atmosphere in which the activation is performed, no Crvl is present in these samples.In fact, absorptions typical of Crvl in the i.r. and u.v.-visible diffuse reflectance spectra are not evident. In order to investigate further the CrII' + CrIV transformation the sample Z,C,-7A was reduced under H, on a thermogravimetric balance up to 673 K. An exothermic peak is observed at 523-573 K with an associated weight loss of ca. 1.2% (fig. 9). The weight loss is entirely associated with the CrV1 -+ Cr'" reduction. This is in agreement with the calculated value of 1.44 Oh, assuming that the amount of Crvl is 14.5 O/O of the total Cr, as was derived from the chemical analysis. Accordingly, the i.r. spectrum obtained following the d. t.a.-t.g. experiment shows no evidence of absorptions characteristic of Crvl coordinated tetrahedrally.Discussion The data previously reported show evidence of a spinel-like phase for all the samples investigated. ZnO is detected as a separate phase in some cases. Other authors previously reported the spinel-like phase to be non-stoichiometric for similar samples prepared in a slightly different way. 7-13 The evidence of this non-stoichiometry was found primarily in the X.r.d. powder patterns. The non-stoichiometry was explained in terms of octahedrally coordinated ZnI' ions. The structure was regarded as intermediate between a spinel and a rock-salt structure. This hypothesis was supported by a profile-fitting analysis of the X-ray powder pattern,' and occupancy factors were given for octahedral and tetrahedral Zn" and for octahedral CrIII ions.', Studies of H, heterolytic dissociation and CO absorptionlo* l2 provide evidence of surface sites different from pure ZnO and stoichiometric ZnCr,O,. In addition, the high catalytic activity of the Zn/Cr = 1/1 system was related to this non-stoichiometric spinel phase.The decrease in activity with time was attributed to the metastable behaviour of the system, which evolves to stoichiometric ZnCr,O, with segregation of Zn0.7y89 l1 Moreover, the width of e.s.r. signals indicated a distribution of CrI'I-CrIII distances suggesting a non-ordered system.12 Even if the model proposed, considering a non-stoichiometric spinel phase, is in a good agreement with experimental data, other models can be taken into account. In fact profile-fitting analysis of such a microcrystalline pattern is no definite evidence, and the authors themselves could not exclude the presence of small amounts of amorphous ZnO,ll even if this was not present in the profile-fitting model.Moreover, the absorption of probe molecules and catalytic activity can be related to surface properties and are not direct evidence of a bulk non-stoichiometric spinel phase. Furthermore, the presence of Crvl was not accounted for. From the data reported, no definite evidence of non-stoichiometry or stoichiometry of the spinel phase emerges for our sample. The data may suggest an alternative model for the investigated Zn/Cr = 1 mixed oxide, in which the sample is thought to be constituted by a shell rich in ZnCrO, and a core of ZnCr,O, and ZnO.In this model surface Zn", C P and Crvl ions may be located in octahedral and tetrahedral sites in a non-ordered fashion. Note that ZnCrO, has a crystal structure closely related to the spinel structure : in fact it is constituted by the same oxygen lattice, within which Zn is located in the octahedral and Cr in the tetrahedral holes. Thus the presence of octahedrally coordinated Zn" is justified by the simultaneous presence of tetrahedrally coordinated Cr". In the bulk, Zn" ions not accommodated in the stoichiometric spinel phase are present as ZnO. The sample is never monophasic, and ZnO is not detected at 673 K in air by X.r.d. analysis because of the poor crystallinity of the phases and the large width of the peaks. On increasing the calcination temperature the amount of Crvl decreases, probably through the decomposition of ZnCrO, to ZnCr,O, and ZnO.The appearance of ZnOC. Cristiani, P . Forzatti and M. Bellotto 905 in the X.r.d. spectra is due to a better crystallization of all the phases, with a consequently smaller f.w.h.m. and better resolution of the peaks. This hypothesis is supported by the calculated spectra of fig. 1, in which the different crystallization of the phases is simulated by varying the f.w.h.m. Fig. 1 (i), like fig. 1 (b), does not show clearly the presence of ZnO. The samples activated in N, show the same behaviour as those calcined in air. There is a trend towards improved crystallization on increasing the temperature. At the same temperatures, however, crystallite dimensions are smaller compared with the samples calcined in air.The inert atmosphere leads to a smaller CrV' content in these samples. With regard to the samples activated in H,, the reducing atmosphere causes the complete absence of Crv' even at 673 K. Increasing the annealing temperature to 773 K, rapid crystallization of ZnO is detected, while the spinel-like phase behaves as it does in the samples activated in N,. In this case the reduction of Crv' to Cr"' may explain the observed phase transformations, and be the cause of the different crystallization rates of ZnO and ZnCr,O,. The reduction of the surface layers of ZnCrO, may result in a structural situation that is different from a mixture of ZnO and ZnCr,O,. The interaction of the two phases to form a solid solution and/or to give a defective phase may justify the surface behaviour reported by other authors and discussed above.Conclusions The basic features of the Z,C, oxide system under study can be summarized as follows. (1) At 373 K the sample is microcrystalline and may be formed of a mixture of ammoniacal compounds and carbonates. (2) On heating in air at 673 K a spinel-like phase is obtained. Crvl is also present in the surface layers. Calcination at higher temperatures brings about a decrease in the amount of Crv' and the crystallization of all the phases present. (3) Activation in N, shows the same trend on increasing the temperature. A smaller amount of Crv' is present in the structure because of the annealing atmosphere. (4) The activation in H, brings about a more rapid crystallization of ZnO, formed through the reduction of ZnCrO, to ZnCr,O,.The behaviour of the system may be explained hypothesizing non-stoichiometry of the spinel phase in line with previous suggestions. '-13 An alternative model is considered here, which is also consistent with the experimental data. In this model the sample is constituted by a shell rich in ZnCrO, and a core of ZnCr,O, and ZnO. This work was performed under a contract from Progetto Finalizzato Energetica 11. Thanks are due to Dr N. Ferlazzo and to Prof. M. Zocchi for stimulating discussion and criticism. References 1 G. Natta in Catalysis, ed. P. H . Emmet (Reinhold, New York, 1953), vol. 111, chap. 8. 2 G. Natta, U. Colombo and I. Pasquon, in Catalysis, ed. P. H. Emmet (Reinhold, New York, 1953), 3 G.Natta and P. Corradini, Proc. Znt. Symp. React. Solids, Gothenburg, 1952, p. 619. 4 J. E. Germain, J. Bigourd, J. P. Beaufils, B. Gras and L. Pousolle, Bull. SOC. Chim. Fr., 1961, 1777. 5 M. Krans, 2. Zitny, D. Mihajlova and A. A. Andreev, Coll. Czech. Chem. Commun., 1971, 41, 3563. 6 F. Runge and K. Zepf, Brennstof-Chem., 1954, 35, 167. 7 G. Del Piero, F. Trifiro and A. Vaccari, J. Chem. Soc., Chem. Commun., 1984, 656. 8 M. Di Conca, A. Riva, F. Trifiro, A. Vaccari, G. Del Piero, V. Fattore and F. Pincolini, Proc. 8th Znt. 9 G. Del Piero, M. Di Conca, F. Trifirb and A. Vaccari, Reactivity of Solids, 1985, 1029. vol. V, chap. 3. Congr. Catal. (Dechema, Frankfurt am Main, 1984), vol. 2, p. 173. 10 G. Busca and A. Vaccari, J . Catal., 1987, 108, 491.906 Structural Invesitigation of a Zn-Cr Mixed Oxide 11 A. Riva, F. Trifiro, A. Vaccari, G. Busca, L. Mintchev, D. Sanfilippo and W. Manzatti, J. Chem. SOC., 12 M. Bertoldi, G. Busca, B. Fubini, E. Giamello, F. Trifiro and A. Vaccari, J. Chem. Soc., Faraday 13 F. Trifiro, L. Mintchev, G. Busca, A. Vaccari and A. Riva, J. Chem. SOC., Faraday Trans. I, 1988,84, 14 H. P. Klug and L. E. Alexander, in X-Ray Diflraction Procedures (Wiley, New York, 1974), p. 531. 15 I. M. Kolthoff and P. J. Elving, in Treatise on Analytical Chemistry (Interscience, New York, 1959). 16 R. A. Nyquist and R. 0. Kogel, in IR Spectra of Inorganic Compounds (Academic Press, New York, 17 C. K. Jnrgensen, in Absorption Spectra and Chemical Bonding in Complexes (Pergamon Press, Oxford, 18 P. Forzatti, C. Cristiani, N. Ferlazzo, L. Lietti, E. Tronconi, P. L. Villa and I. Pasquon, J. Catal., 1988, 19 J. Preudhomme and P. Tarte, Spectrochim. Acta, 1971, 27, 1817. 20 P. P. Cord, P. Courtine and G. Pannetier, Spectrochim. Acta, 1972, 28, 1601. 21 F. A. Cotton and G. Wilkinson, in Adoanced Inorganic Chemistry (Interscience, New York, 1966), Faraday Trans. 1, 1987, 83, 2213. Trans. I , 1988, 84, 1405. 1423. 1971). 1962), chap. 13. 111, 120. p. 713. Paper 8/01823E; Received 9th May, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500895

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chern. Sac., Faraday Trans. I , 1989, 85(4), 907-916 Characterization of the Mixed Perovskite BaSn,_,Sb,O, by Elect 1-01 y te Elec troreflec t ance, Diffuse Reflectance, and X-Ray Photoelectron Spectroscopy Gerardo Larramona and Claudio GutiCrrez" Instituto de Quimica Fisica ' Rocasolano ', C.S.I.C., Serrano 119, 28006-Madrid, Spain Isabel Pereira CECUL (Centro de Electroquimica e Cine'tica da Universidade de Lisboa j , Faculdade de Ciencias, R. da Escola Polite'cnica 58, 1200 Lisbon, Portugal Manuel Rosa Nunes and Fernanda M. A. da Costa I.N.I.C., Centro de Qulinica Fisica e Radioquimica da Universidade de Lisboa, Faculdade de Ciencias, R. da Escola Polite'cnica 58, 1200 Lisbon, Portugal The band gap of the mixed perovskite BaSn,-,Sb,O, has been reported to be 134 kJ mol-1 (J-M.Hermann, M. R. Nunes and F. M. A. da Costa, J. Chern. Soc., Faraday Trans. I, 1982, 78, 1983), corresponding to 1.4eV, which is the optimum value for solar-energy conversion. In order to confirm this value, pressed discs of the mixed perovskite were sintered at 1673 K and made into electrodes, and their electrolyte electroreflectance (e.e.r.) spectra measured. The band gap was determined by fitting the e.e.r. spectra to Aspnes' formula for the low-field regime. The band gap of the pure, insulating perovskite BaSnO, was determined from its diffuse reflectance spectrum, after application of the Kubelka-Munk formula. The band gap of the mixed perovskite was constant and equal to 330kJmol-' for Sb concentrations in the range x = &lo%, in agreement with the observed constancy of the lattice parameter for x = &17 %, the range over which the perovskite exists as a single phase.The high band-gap energy of this material makes it unsuitable for solar-energy conversion. X-Ray photo- electron spectroscopy measurements showed that all the Sb in the mixed perovskite was present as Sb', as would be expected from the high conductivity of the samples. From the atomic ratios it was concluded that probably the surface of the pure BaSnO, perovskite is enriched in SnO,, and that of the mixed perovskite in Sb,O,. The mixed perovskite BaSn,-,Sb,O, is a semiconductor that exists as a single-phase solid solution over the interval 0 < x < 0.17, a second phase appearing at higher Sb Concentrations.' This material has good conductivity, and its activation energy of conduction is reported to be 134 kJ mol-',' which would make it an ideal candidate for solar-energy conversion, since this is the value of the band gap (1.4 eV,t corresponding to 900 nm) for which conversion efficiency is at a maximum.In this work we have characterized several samples of the mixed perovskite in the solid-solution range, in order to determine its band gap and the chemical composition of its surface. Experimental The mixed perovskites were synthesized by a solid-state reaction. The calculated amounts of BaCO,, SnO, and Sb,O, (Johnson Matthey, Specpure) were ground t 1 eV z 1.602 x lO-"J. 31 907 FAR 1908 Characterization of the Mixed Perovskite BaSn,-,Sb,O, together in an agate mortar and the powder was heated under air in an alumina boat up to a final reaction temperature of 1613 K.Then the powder was immediately made into discs, 6 mm in diameter and 2 mm thick, using a die at a pressure of 20 MN m-2, and the discs were sintered at 1673 K. The lattice parameter of the perovskites was determined with a Philips X-ray powder diffractometer. The electrical resistivity of the discs was determined by the four-point met hod. The perovskite electrodes for the electrolyte electroreflectance (e.e.r.) measurements were prepared as follows. Electrical contact to the discs was made by rubbing their reverse side with an In-Ga eutectic, and then glueing a silver disc to their backs with silver epoxy adhesive, a copper wire having been previously welded to the silver disc.The sample was then mounted in a glass tube with Araldite epoxy resin, so that the electrolyte could only make contact with the perovskite. Ohmic contact to the samples was verified by glueing silver discs to the two sides of each disc. Linear current us. potential plots were obtained for potentials from -2 to 2 v. The e.e.r. setup was conventional, and has been described already.2 Essentially, the (usually minute) variations (AR) in the amount of light (R) reflected from the electrode surface, produced by a modulation of the electrode potential, are detected by a lock-in amplifier, an e.e.r. spectrum consisting of a plot of AR/R us. wavelength. The potential was modulated with a sine wave of frequency 170 Hz. The actual amplitude of the modulation at the electrode-electrolyte interface was a fraction of the nominal amplitude applied at the control input of the potentiostat, since a large fraction of the modulation dropped in the ohmic resistance of the electrolyte layer extending from the electrode surface to the tip of the Luggin capillary.The fraction of the amplitude modulation actually dropping at the electrode-electrolyte interface was calculated by vectorial subtraction of the parasitic electrolyte resistance (measured as the total impedance at high frequency) from the total impedance at 170 Hz between the working and reference electrodes, whose magnitude and phase were measured with a lock-in amplifier. The diffuse reflectance spectra of the undoped BaSnO, perovskite powder was measured with a Lambda 9 u.v.-visible-near-i.r.spectrophotometer from Perkin- Elmer. X-Ray photoelectron spectra were obtained in a Leybold Heraeus LHS 10 spectrometer with Mg K, radiation (1253.6 eV) at a pressure < N mP2, assuming the binding energy (Eb) of C 1s to be 284.6 eV. Results and Discussion Lattice Parameter and Resistivity The lattice parameter, resistivity, thermal activation energy of the resistivity in the range 323-453 K, and colour of the six perovskites are given in hble 1. The lattice parameter remains constant and equal to the literature value of 4.1 17 A for pure BaSnO,, over the whole Sb concentration range (0-1 7 O h ) where a single phase exists. The resistivity is only 0.15 Q cm, low enough for application as an electrode, for Sb concentrations > 5 YO ; this resistivity is several orders of magnitude lower than that of compressed powders of the same composition reported in ref.(l), owing to the elimination of the intergranular resistance brought about by sintering. The thennal activation energy of the resistivity is negligible (< 0.07 eV) for all but the sample with 5 O/O Sb; again, this value is one order of magnitude lower than that of the corresponding powdered samples,' also for the same reason. The colour changes progressively with Sb concentration, from slightly yellowish (which points to a band gap > 3.1 eV) for pure BaSnO,, to grey, and to black for samples with > 10 % Sb. The colour of the powdered samples was lighter.'909 G. Larramona et al. Table 1. Several physical characteristics of the samples Sb mol lattice resistivity disc colour (%I /A ( T = 298 K) (453-323 K)/eV 1673 K) fraction paraqeter P m cm activation energy ( T i n t = 0.0 4.1 14 - 35 2.50 - 0.15 5.0 - 10.0 4.117 0.15 0.15 15.0 - 17.0 4.1 18 - 1.25 4.114 2.103 - yellowish 0.03 dark grey 0.28 dark blue 0.0 1 black 0.07 black - black 0.07 grey -1.1 0 +1.2 EIVvs.SCE Fig. 1. Cyclic voltammograms of the mixed perovskites with an Sb mole fraction of (a) 2.5, (b) 10 and (c) 15 % in 1 mol dm-3 Na,SO,. Sweep speed was 100 mV s-l. 31-2910 Characterization of the Mixed Perovskite BaSn,-,Sb,O, L I I I 300 400 500 Unm Fig. 2. Normalized e.e.r. spectrum of the mixed BaSno,9Sbo,,03-perovskite in 1 mol dm-3 Na,SO, at 0.4 V us. SCE, showing the linearity between the e.e.r. signal and the actual modulation amplitude, which was: 0, 29, 0, 87 and X, 174 mV r.m.s.Fig. 3. Dependence of the e.e.r. signal on the polarization potential at (a) 320 and (b) 360 nm for the mixed BaSn,,Sb,,03 perovskite in 1 mol dm-3 Na,SO,. Actual amplitude modulation, 176 mV r.m.s. Sweep speed of the polarization potential, 5 mV s-l.G. Larramona et al. 91 1 s Q Q I I 1 300 400 500 A/nm Fig. 4. E.e.r. spectra of five mixed perovskites in 1 mol dm-3 Na,SO, at 0.4 V us. SCE. The composition of the perovskite and the actual modulation amplitude, 4,, are as follows : (a) 1.25 YO Sb, 4, = 245 mV; (b) 2.5% Sb, 4, = 160mV; ( c ) 5.0% Sb, 4, = 160mV; (6) 10% Sb, 4, = 176 mV; (e) 15% Sb, 4, = 180 mV.912 3 - 4 - 8 -12 Characterization of the Mixed Perovskite BaSn,-,Sb,O, - - - I - 4 0 3 Q 4 I 9.-. f.. i ,=*.. t . I . I I \ I \ I I \ I I \ d\ I ! I \ I \ I \ I \ t i, \a ,b 0 1 1 1 I I I 2.6 3.0 3 . 4 3 . 8 4 . 2 EIeV Fig. 5. E.e.r. spectrum of the mixed BaSn,,Sb,~,O, perovskite in 1 mol dm-3 Na,SO, at 0.4V us. SCE with an actual modulation amplitude of 176 mV r.m.s. e, Experimental points; (....) best fit to Aspnes' formula, eqn (1). Electrolyte Electroreflectance Owing to its experimental simplicity, the e.e.r. technique, derived by Cardona et aZ.,4 made popular the study of the band structure of semiconductors by means of their electroreflectance. Aspnes5 provided the theoretical basis for the analysis of the e.e.r. results under low-field conditions, i.e. when the amplitude of the modulation of the electric field at the electrode4ectrolyte interface is low.The electrolyte used here was 1 mol dmV3 Na,S04. The potentials were within the stability range of the perovskites, clearly apparent in the voltammograms (fig. 1). Owing to the low reflectivity of the samples, the e.e.r. measurements were effected initially at near-normal incidence and with the perovskite disc very near the (single) silica window, in order to maximize the amount of light scattered by the electrode reaching the photocathode of the photomultiplier tube. However, later it was found that the same results were obtained in a conventional two-window cell with an angle of incidence of 45". Also, the same spectra were obtained without any treatment of the discs and after 'polishing' them with filter paper. The e.e.r. spectra increased linearly with the modulation amplitude, as can be seen in fig.2 for the BaSno,9Sbo,,0, perovskite. All the other samples presented the same behaviour. This is a characteristic of low-field e.e.r. spectra of semiconductors. In fig. 3 we show the dependence of the e.e.r. signal on the polarization potential for two wavelengths, 320 and 360 nm. The signal remains constant over the potential range - 0.1 to 0.7 V vs. SCE, which corresponds to the plateau in the voltammograms of fig. 1, i.e. the potential region without faradaic reaction, in which the current is consumed only in charging the electrical double layer. The independence of the e.e.r. signal from the polarisation potential is another characteristic of semiconductors in the low-field regime. The e.e.r.spectra of five mixed perovskites at a potential of 0.4 V vs. SCE, well within the region where the e.e.r. signal is independent of the polarization potential, are given in fig. 4, All of them are fairly similar, pointing to a lack of influence of the Sb contentG. Larramona et al. 913 Table 2. Parameters obtained from fitting Aspnes' eqn (1) to the experimental e.e.r. spectra Sb mol fraction (1 Yo) E,/eV wo T/eV C 1.25 3.43 270 2.5 3.35 320 5.0 3.45 360 10.0 3.46 0 15.0 3.65 270 20.0 3.70 270 0.43 0.42 0.44 0.47 0.45 0.42 0.70 0.90 0.45 0.20 0.75 0.55 on the band structure of the perovskite. The spectra were fitted to Aspnes' formula for the low-field regime :5 (1) where n = for a three-dimensional simple parabolic critical-point model, Eg is the width of the band gap, C and 9 are amplitude and phase factors, respectively, and r is a broadening parameter related to the lifetime.Fitting was effected with an HP 9816 computer by trial and error, without constraining any of the parameters, and was very sensitive to the value of E,, increments as low as 0.05 eV producing readily appreciable changes in the quality of the fit. Aspnes' formula, eqn (l), fitted the experimental points fairly well, as can be seen in fig. 5 for the perovskite with 10% Sb. The parameters obtained from fitting the e.e.r. spectra of the six perovskites to Aspnes' formula are given in table 2. The band gap remains constant, and equal to 3.4eV, for Sb concentrations up to lo%, increasing to 3.7 eV for higher Sb concentrations. This is in perfect agreement with the constancy of the lattice parameter for Sb concentrations up to 17% (table 1).It becomes clear that the band-gap value of 1.4 eV reported earlier' corresponds to excitation to some energy level within the band gap or, more probably, is simply an artifact due to the intergranular resistance in the compressed powder. The fact that the activation energy of the resistivity for these sintered samples (table 1) was 3-6 times smaller than that for the compressed powders' lends support to this last hypothesis. AR/R = Re [C exp (i0) (E- Eg + iT)-"] Diffuse Reflectance The diffuse reflectance (r) spectrum of the pure BaSnO, perovskite is given in fig. 6(a), and the logarithm of the corresponding Kubelka-Munk function, (1 - r)2/2r, is given in fig. 6(b).The band gap is also 3.4eV, in perfect agreement with that of the perovskites with up to 10% Sb. X-Ray Photoelectron Spectra Only two samples were measured, the pure BaSnO, perovskite and the mixed BaSn,,,,Sb,,,, perovskite. The binding energies (Eb) of the different atoms are given in table 3. The binding energies of the Ba 3d and Sn 3d multiplets are the same for both samples, and correspond to Ba2+ and Sn4+, respectively.6 The spectrum of the level 0 1s is more complex, showing two maxima at 529.7 eV (lattice 02- ions) and 531.7 eV (probably from chemisorbed water). We have also found these two 0 1s peaks in many compounds, including metals, metal oxides, perovskite oxides of the type LnMO, (Ln = La, Pr or Nd; M = Cr, Mn, Fe, Co and Ni), and superconductors of the type LnBaCuO (Ln = Y, Nd).In principle the 0 1s signal at higher binding energies could be due, not914 Characterization of the Mixed Perovskite BaSn,-,Sb,O, 200 400 600 800 1000 1200 Alnm 2.5 3.0 3.5 4.0 h vleV Fig. 6. (a) Diffuse reflectance spectrum of the pure BaSnO, perovskite. (b) Its corresponding Kubelka-Munk function (full line). The band gap energy is ca. 3.4eV. Table 3. Binding energies of several atomic levels (in eV) sample 0 1s Sn 34,; Ba 3dg.t Sb 3d; BaSnO, 02-: 529.7; 494.5-485.7 795.6780.3 - 53 1.7" BaSno.,,Sbo., 7 0 3 530.8 494.5-48 5.8 79 5.7-780.3 539.8 a The 0 1s signal at higher binding energies could be due to 0-, OH-, or chemisorbed water. We attribute it to the latter for reasons stated in the text. only to chemisorbed water, but also to OH- or 0-.This latter species can be discounted, since its high reactivity makes it very unstable. The very high value of the O/Sn ratio in both perovskites (see below) leads us to think that the peak at 531.7 eV is due to strongly chemisorbed water. In the mixed perovskite the levels 0 1s and Sb 3dg overlap. The maximum appears at & = 530.8 eV, intermediate between the two 0 1s maxima for the pure perovskite. Sb is present as Sb5+, since the binding energy of the Sb 3d; level appears at 539.8 eV; the maximum corresponding to the Sb 3d; level, shown as dotted line in fig. 7, has an area $ times larger, and a binding energy 9.40 eV lower, than those of the Sb 3d; maximum. The fact that Sb is in the pentavalent state was to be expected from the high conductivityG.Larramona et al. 915 545 540 535 530 525 Eb /eV Fig. 7. X-Ray photoelectron spectra of (a) the pure BaSnO, perovskite and (b) the mixed BaSn,~,,Sb,,,70, perovskite. The dotted curve corresponds to the Sb 3dg level, and was calculated from the area and binding energy of the Sb 3d; level. Table 4. Surface atomic ratios as determined by X.P.S. sample O/Sn Ba/Sn Sn/Sb Sb/O BaSnO, measured 02-: 2.60; 0.77 - - expected 3 1 2.67" - - BaSn0.83Sb0.1703 measured 5.58 1.03 2.56 14.3 expected 3.61 1.21 4.88 17.7 a The 0 1s signal a t higher binding energies could be due to 0-, OH-, or chemisorbed water. We attribute it to the latter for reasons stated in the text. of the samples, since each Sb5+ ion contributes one electron to the conduction band of the perovskite.It could be thought that these conduction-band electrons should be detected (which they are not) in the X.p. spectra at energies near the Fermi level, but in the very similar case of SnO, doped with 3% Sb they are only seen in the ultraviolet photoelectron spectra, and at a very expanded scale.7916 Characterization of the Mixed Perovskite BaSn,-,Sb,O, The electrons whose energy is analysed in X.p. spectra proceed from a layer < 3 nm deep. The atomic composition in this surface layer has been calculated using the sensitivity factors reported by Wagner et aZ.8 and is given in table 4. For the peaks a Gaussian-Lorentzian fraction of 0.4 was assumed ; for each peak the base line was taken as a straight line tangent to the peak wings. As a check of the method we have effected quantitative measurements of several stoichiometric oxides (CuO, Fe,O, and V,O,), the agreement being within 6-10 Oh.In the pure perovskite both the 02-/Sn and the Ba/Sn ratios are lower than expected for BaSnO,: 3 and 1, respectively. Therefore, we think that the surface of the pure perovskite is enriched in SnO,, although in an amount certainly undetectable by X-ray diffraction. In the mixed perovskite the ratio O/Sn is higher, and the Ba/Sn, Sn/Sb and Sb/O ratios lower, than those expected. The lower value of the Sb/O ratio could be due to the presence on the perovskite surface of a small amount of Sb,O,. Conclusions (1) The band gap of the pure BaSnO, perovskite, as estimated from its diffuse reflectance spectrum, is 3.4eV. (2) True e.e.r.spectra were obtained from all the mixed perovskite samples. (The e.e.r. signal was independent of the polarization potential, and increased linearly with the modulation amplitude.) The band gap obtained by fitting the e.e.r. spectrum to Aspnes’ formula was 3.4 eV for Sb concentrations up to 10 %, and 3.7 eV for the samples with 15 and 20% Sb. (3) X-Ray photoelectron clearly showed that all the Sb in the mixed perovskites was present as Sb5+, as expected from the high conductivity of the samples. From the atomic ratios it was concluded that probably the surface of the pure perovskite is enriched in SnO,, and that of the mixed perovskite in Sb,O,. The financial help of Comitk Conjunto Hispano-Norteamericano para la Cooperacion Cientifica y Tecnologica (project CCA 8309038), the Spanish Comision Interministerial de Ciencia y Tecnologia (project ID 87094), Portuguese Junta Nacional de Investigacao Cientifica e Tecnologica (research contract no. 912.86.183) and NATO (collaborative research grant no. 0168/85) is gratefully acknowledged. We thank Dr J. L. Garcia- Fierro, Instituto de Catalisis y Petroleoquimica, C.S.I.C., for the XPS spectra, and Dr A. Duran, Instituto de Ceramica y Vidrio, C.S.I.C., for the diffuse reflectance spectrum. References 1 J-M. Herrmann, M. R. Nunes and F. M. A. da Costa, J. Chem. Soc., Faraday Trans. I , 1982, 78, 2 C. Gutierrez and M. A. Martinez, J. Electrochem. Sac., 1986, 133, 1873. 3 A. J. Smith and J. E. Welch, Acta Crystallogr., 1960, 13, 653. 4 K. L. Shaklee, F. H. Pollak and M. Cardona, Phys. Rev. LRtt., 1965, 15, 883. 5 D. E. Aspnes, SurJ Sci., 1973, 37, 418. 6 C. D. Wagner, W. M. Riggs, L. E. Davis, J. F. Moulder and G. E. Muilenberg, Handbook of X-Ray 7 P. A. Cox, R. G. Egdell, C. Harding, W. R. Patterson and P. J. Tavener, Surf. Sci., 1982, 123, 179. 8 C. D. Wagner, L. E. Davis, J. A. Taylor, R. H. Raymond and L. H. Gale, SurJ Interfuc. Anal., 1981, 1983. Photoelectron Spectroscopy (Perkin-Elmer Corp., Minnesota, 1979). 3, 211. Paper 8/01838C; Received 10th May, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500907

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. Soc., Furaday Trans. I, 1989, 85(4), 917-928 3 4 5 6 7 8 3' 4' 5' 6' 7' 8' IH and 13C Longitudinal and Transverse Relaxation in Aerosol OT in Methanol Solution and Inverted Microemulsions in Benzene Frank Heatley Department of Chemistry, University of Manchester, Manchester MI3 9PL 'H and 13C longitudinal and transverse relaxation in inverted micro- emulsions of Aerosol OT (AOT) in deuterated benzene has been studied as a function of composition and resonance frequency. Except at high AOT concentrations (> 2 mol AOT per kg C6D6) and low water content [AOT/ H,O mole ratio ( R ) < 21 the frequency dependence of T,, T, and the nuclear Overhauser effects can be satisfactorily interpreted in terms of the two-step correlation function of Wennerstrom et al. (H. Wennerstrom, B.Lindman, 0. Soderman, T. Drakenberg and J. B. Rosenholm, J. Am. Chem. Soc., 1979, 101, 6860). The slow correlation time decreases as the water content increases up to R x 3 and thereafter is constant. This behaviour suggests that the slow process is loss of AOT from the aggregate into the hydrocarbon phase. Relaxation of AOT in dilute methanol solution has been used to determine the relaxation mechanism. 917918 Relaxation in Aerosol OT Inverted microemulsions of Aerosol OT [AOT, sodium bis(2-ethylhexy1)sulpho- succinate, 11 in p-xylene have been studied using 13C & and NOE data.8 For the head- group nuclei C-1 and C-1', z, and 2 , were of the order of 10-l' and s, respectively, while S varied from ca. 0.8 at an H,O/AOT mole ratio (R) of 3.8, to ca.0.6 at an R value of 34. However, the C-1' linewidth (i.e. T,) was not consistent with the motional parameters extracted from the and NOE data, and a third process slower than z, was postulated. In view of the uncertainties associated with determining from the linewidths of lH decoupled spectra, it was thought appropriate to test the validity of eqn (1) more reliably by combining 13C and NOE data with the 13C T, data determined using the spin-echo technique. This paper reports such an investigation focussing on the AOT head-group in C,D, microemulsions. In addition 'H & and T, relaxation has been investigated for comparison with 13C relaxation, and relaxation in unaggregated AOT in methanol has been studied in order to obtain a comparison with unrestricted motion. Experiment a1 'H Data at 80 (TJ, 300 (q, T,) and 500 MHz (q) were obtained using Bruker Spectrospin WP-80, AC-300E and AM-500 spectrometers, respectively. The WP-80 and AC-300E instruments also provided 13C data at 20.1 (q, NOE) and 75.5 MHz (&, T,, NOE), respectively.T, measurements were performed using the (~-2-n/2) inversion recovery sequence, and T, measurements using the Carr-Purcell-Meiboom-Gill (CPMG) spin-echo sequence with an interval of 1 ms between z pulses. 'H decoupling was not applied during the 13C experiments at 75.5 MHz in order to avoid heating problems with the large decoupling fields necessary. l4 Nuclear Overhauser enhancements were measured using the gated decoupling method with relaxation intervals of at least 5 T,, for the enhanced and 8 T,, for the unenhanced spectrum.AOT was obtained from Sigma Chemical Company, St. Louis, U.S.A. and was used without further purification. [This procedure is justified in ref. (1 5).] Deuteriated solvents were obtained from CEA, France, and were used as received. Solutions were not degassed; trial degassing of a few samples showed that the AOT relaxation was essentially unaffected. Relaxation Theory T,, & and the NOE depend on a spectral density function J(w), which is related to the correlation function (1) by (1 -s")z, 9 2 , G(z)exp(imz)dz = +- 1 +m22,2 1 +m2g I3C Relaxation For a 13CH, group where the 13C is relaxed entirely by dipole4ipole interaction with the attached protons, the 'H-decoupled and the NOE are expressed in terms of the spectral density for motion of the C-H bond, JCH(w), by where Q = (3/10) o10/4~)2y,2yc2~2/rcH6.yx is the magnetogyric ratio and w, is the resonance frequency of nucleus X. rCH is the C-H bond length, taken to be 109 pm. TheF. Heatley 919 NOE is the ratio of the enhanced and unenhanced l3C intensities. For the CH, groups, eqn (3) and (4) do not take into account cross-correlation spectral densities which appear in rigorous density-matrix relaxation theory.16 However, eqn (3) does apply to the initial relaxation rate, while cross-correlation effects on the NOE are small compared to experimental errors and the limits of fitting data to eqn (2). Transverse 13C relaxation measurements in this work were performed without ‘H decoupling. In CH, groups, cross-correlation terms manifest themselves in different relaxation rates for the outer and centre lines of the 13C triplet.” However, they can be eliminated by taking the average of the two rates. For a 13CH, group, the average transverse relaxation time i d 4 ‘H Relaxation The H-1 proton (X) and H-1’ protons (A and B) form a coupled relaxation system. Although in C,D, microemulsions A and B are strongly spin-spin coupled,15 conventional first-order coupled longitudinal relaxation theory1* has been applied. This is felt to be acceptable because the AB multiplet is so broadened that only the total AB intensity was observable, and relaxation of this signal is dominated by the strong AB geminal interaction.The time dependence of the intensity of proton i, Si, is given by dS, - = - (& - sg) (p* + 2 Pij) - 2 Oij(s, - Si”) d t j i where The superscript O indicates equilibrium values, and the parameter p* represents relaxation by other mechanisms such as intra- and inter-molecular interactions with other protons.For the purpose of calculating relaxation times, since only the total AB intensity, designated SAB, is measured, it is convenient to combine the three equations governing SA, % and ,!& into two equations expressing the evolution of SAR and J;(. Eqn (7a) and (7b) give biexponential relaxation for SAB and &. In practice, the recovery of SAB was effectively exponential, although the recovery of S;E was in some systems distinctly not so. Further details are described in the Results section below. To calculate NOE values correctly, the three equations (6) were retained. The resulting expressions are AX + 0nx XA+PXB+P* NOEx(AB) = 1 +920 Relaxation in Aerosol OT Table 1.'H and 13C relaxation data at 22 "C for AOT in CD30Da concentrationb frequency parameter nucleus /MHz 0.24 2.1 T,, c- 1 75.5 c-1' 75.5 NOE C-1 75.5 c-1' 75.5 T H H-1' 300 H- 1 300 T , H H-1' 80 T2H H- 1 ' 300 0.77 0.22 0.44 0.13 2.94 1.96 2.93 2.06 0.55 2.0 0.52 0.52 a Relaxation times in s, uncertainty+_ 5 %. mol (kg CD,OD)-l. NOEJj} is the ratio of the i intensity with j saturated, relative to the unperturbed intensity. For lH transverse relaxation, only the AB peak has been studied. Because of the dominance of the geminal AB interaction TfB is given to good approximation by the expression for a pair of protons with additional contributions from other protons: where T,* represents external contributions.Results CD,OD Solutions The purpose in studying relaxation in methanol solution is to examine a system of unassociated molecules with all correlation times meeting the extreme-narrowing limit (W;Z: 6 1). In this regime is independent of resonance frequency and equal to q, and l3C(lH)NOE values for entirely dipolar relaxation with protons reach their maximum of 2.988. By these criteria, solutions of AOT in CD,OD containing 0.24 mol (kg CD, OD)-' met the desired objective (table 1). At a concentration of 2.1 mol (kg CD,OD)-l, however, the 13C values are much lower and the 13C NOE values are significantly less than maximal. Evidently, aggregation can occur in methanol as well as in water and hydrocarbons, although at a much higher concentration.Data are analysed only for the former solution. All relaxation curves were exponential within experimental error over a decade recovery. values for C-1 and C-1' is not quite 2, as expected for isotropic rotation. From eqn (2) and (3), setting S2 = 0, we obtain effective correlation times of 4.8 x lo-" and 5.5 x s for C-1 and C-l', respectively. It has been shown15 that AOT exists predominantly in conformation (111) in fig. 1 . Assuming tetrahedral bond angles and exact staggering, and defining A to be the H-1' trans to H-1, we calculate r,, = 178 pm, rA, = 306 pm and rBx = 249 pm. Eqn (7a) and The ratio of the 13CF. Heatley 92 1 SO” s 0, CO, R Fig. 1. Head-group conformations of AOT.(7b) with p* = 0 then yield TAB = 0.66 s and T,, = 4.74 s, compared with experimental values of 0.55 f 0.02 s and 2.0 0.1 s, respectively. The discrepancy between the calculated and experimental values, particularly Tx, is outside reasonable uncertainties in the internuclear distances and correlation times, and indicates a significant contribution to relaxation from sources other than interaction between H-1 and H-1’. These ‘ external ’ interactions are either intermolecular (including dissolved oxygen) or intramolecular with protons in the alkyl chains. In both cases, the H-1 and H-1’ signals are probably essentially equally affected, but the relative effect on H-1’ is less than on H-1 because of the comparatively short geminal AB distance. If p* is set equal to (pAX +psx + oAx + o,,), the calculated relaxation times are TAB = 0.60 s and Tx = 2.2 S, in good agreement with experiment.To support this analysis, the NOE of H-1 on irradiation of H-1’ was examined. The experimental NOE was 0.23f0.03, in good agreement with the value 0.2 calculated using this value of p*. Thus it appears that H-1 is relaxed as effectively by ‘external’ interactions as by interactions with H-1’. In the analysis of lH relaxation in microemulsions, p* was therefore incorporated using the same relationship p* = pAX +pBx + oAX + oBx. To obtain the same correction for relaxation of the H- 1’ protons, 1/T: in eqn (9) was set equal to ((IAx KAX + Q, KBx). The effect of these ‘external’ contributions on H-1’ relaxation is small. ‘’C Relaxation in Microemuisions Relaxation of C-1 and C-1’ in microemulsions of varying composition was studied at 307 K.Results are given in table 2. It is found that T, and the NOE depend on the resonance frequency, that T, is considerably less than and that the NOE is less than maximum. The participation of a slow motion (w:zt 2 1) in the relaxation process is therefore indicated. Note that if C-1 and C-1’ experienced identical motions, the relaxation times of C-1 should be twice those of C-l’, In fact the difference is less than a factor of 2. Qualitatively, the increase in and the NOE with increasing water content922 Relaxation in Aerosol OT Table 2. 13C relaxation data at 307 K in AOT microemulsions in C,D,* frequency/MHz 75.5 20.1 concentrationb Rc nucleus Tl T, NOE NOE 2.19 0.37 C-1 c-1’ c-1’ 2.1 1 2.9 C-1 c-1’ 2.1 1 11.0 c-1 c-1’ 0.222 0.98 C-1 c-1’ 0.2 14 12.1 c-1 c-1’ 2.16 0.78 C-1 0.220 0.129 0.187 0.093 0.198 0.120 0.306 0.169 0.164 0.102 0.288 0.163 0.043 0.027 0.07 1 0.042 0.120 0.077 0.182 0.122 0.112 0.071 0.23 1 0.141 1.27 1.31 1.38 1.42 1.73 1.80 2.3 1 2.37 1 S O 1.62 2.37 2.43 0.046 1.58 0.030 1.69 0.047 1.77 0.031 1.79 0.073 2.25 0.043 2.31 0.147 2.46 0.087 2.59 0.065 2.42 0.043 2.53 0.198 2.49 0.125 2.51 a Relaxation times in s, uncertainty f 5 %.mol (kg C6D6)-I. H,O/AOT molar ratio. and decreasing concentration indicates an increasing contribution from the faster motions. In order to test the validity of eqn (2), data were first analysed for those solutions where values were available at two resonance frequencies.The data set thus comprised five items: at 20.1 and 75.5 MHz, T, at 75.5 MHz and the NOE at 20.1 and 75.5 MHz. A computer search was performed for the values of S2, z, and z, giving the best fit to this data set, the best fit being defined as the minimum value of the quantity 2, given T,exptl is the experimental and the calculated value of parameter y. Data for carbons C-1 and C-l’-were analysed independently, with the results shown in table 3. With the exception of sample 1, the data are represented within the experimental error of ca. 5 % by eqn (2). It is not always necessary to invoke8 a slower process to reconcile and T,. It is difficult to quote an uncertainty in the motional parameters because they are highly covariant, as was found in a study of relaxation in micellar sodium 0ctan0ate.l~ Taking as an example C-1 of the system containing 2.16 mol AOT (kg C,D,)-’ and R = 0.78, if two of the three motional parameters are kept fixed at their best-fit values and only the third allowed to vary, the ranges giving an r.m.s.difference of 5 % or less are S2 = 0.66-0.76, z, = 3.1 to 3.9 ns and z, = 36 to 170 ps. However, if for example z, and z, are readjusted to give the best-fit for S2 = 0.66, their best-fit values are z, = 3.7 ns and z, = 110 ps with an r.m.s. difference of 4.1 YO. For S2 = 0.76 the best-fit parameters are z, = 3.3 ns and z, = 58 ps with an r.m.s. difference of 4.5%. In general, the trend on increasing water content is a large decrease in the order parameter, together with a smaller decrease in 2,.z, appears to increase slightly, but the effect is marginal. For both C-1 and C-1’ in the sample with R = 0.37, the r.m.s. deviation for the five-F. Heatley 923 Table 3. Best-fit motional parameters [eqn (2)] from the 13C relaxation data in table 2 concentra- tion‘ R‘ nucleus S2 zJns q/ps dev.O 2.19 0.37 C-1 c-1’ 2.16 0.78 C-1 c-1’ 2.11 2.9 C-1 c-1’ 2.11 11.2 c - 1 c-1’ 0.222 0.98 C-1 c-1’ 0.214 12.1 C-1 c-1’ 0.79 5.6 0.63 5.4 0.70 3.5 0.61 3.1 0.46 2.3 0.39 2.1 0.15 3.4 0.14 2.2 0.66 1.6 0.50 1.5 0.09 2.6 0.065 2.4 85 8.9 54 11.6 90 1.8 87 1.7 130 2.7 93 4.4 140 6.4 110 4.4 84 2.0 78 3.4 150 4.3 130 5.1 a See table 2. and experimental parameters. R.m.s. percentage difference between calculated item data set is significantly larger than the experimental error of ca.5 %, suggesting that eqn (2) is inaccurate in this case. In order to examine the form of deficiency, a four-item data set of and NOE values only was fitted, yielding the following parameters for C-1 (C-1’): S2 = 0.70(0.55), z, = 4.2 (3.7) ns, z, = 33 (23) ps, r.m.s. percentage deviation = 0.8 (1.3). With these motional parameters, the expected T, values are 65 ms for C-1 and 46 ms for C-1’. Both of these are some 50 O h larger than the experimental values. The same discrepancy was noted by Carnali et aL8 for a more concentrated [6.08 mol (kg solvent)-’] solution of AOT in p-xylene using & estimated from linewidths. ‘H Relaxation in AOT in Microemulsions The most reliable ‘H relaxation data are those for the H- 1’ protons, since their relaxation is dominated by their mutual geminal interaction. This leads to two important advantages in interpreting the data.First the magnitude of the dipolar interactions is known more accurately than for H-1, since it depends less on such variables as the dihedral angle in the head-group conformation and the contribution from intermolecular interactions. Secondly H-1’ relaxation is exponential over more than a decade and can therefore be characterised by a time constant. In contrast, for H-1 it was found that the spin-lattice relaxation recovery curves were frequently non-exponential, particularly in those systems of high-order parameter. The illustration of this behaviour in fig. 2 shows that the instantaneous relaxation ‘ time’ decreases as the relaxation proceeds.This can be understood in terms of the coupled equations (7a) and (7b). Immediately after the n pulse, which inverts both H-1 and H-1’, the relaxation ‘time’ of H-1 (i.e. &) is (pXA +pXB +p* +oAX +oBx). Because the H-1’ protons relax some three times more quickly than H-1, in the later stages of H-1 relaxation, the condition SAB z SABo holds, and the relaxation ‘time’ of H-1 is (pXA+pXH+p*). In the less mobile systems the correlation times are such that +Jj(wi - w j ) > 24(wi + mi). Hence oAX and oBX are negative, and the initial relaxation ‘time’ is less than the final relaxation ‘time’. The same phenomenon leads to ‘H NOE values of less than unity. Table 4 gives data for H- 1’ relaxation in a number of microemulsion systems. Because924 Relaxation in Aerosol OT “r--- 2 P 0 0.5 1.0 tlS 5 Fig.2. Inversion recovery curve of H-1 in a microemulsion containing 2.16mol (AOT) (kg C6D6)-l, R = 0.78, at 296 K and 300 MHz. The line is for eye guidance only. Go - &(t) is in arbitrary units. Table 4. H-1’ relaxation data in AOT microemulsions in C6D6a frequency/ MHz 500 300 80 concentra- tionb R b T/K T T T, T, 2.19 2.16 2.1 1 2.11 0.206 0.222 0.2 12 0.214 0.025 0.021 0.37 296 307 0.78 296 307 2.9 296 307 11.2 296 307 0.18 296 0.98 296 2.5 296 12.1 296 1.0 296 10.9 296 1.37 0.777 0.736 1.10 0.655 0.544 0.682 0.404 0.320 0.480 0.275 0.297 0.512 0.907 0.490 0.310 0.511 0.323 0.420 0.493 0.300 0.019 0.025 0.036 0.037 0.051 0.085 0.086 0.133 0.040 0.054 0.070 0.118 0.061 0.116 ~~ 0.140 0.109 0.107 0.093 0.083 0.087 0.105 0.140 0.076 0.077 0.079 0.122 0.074 0.1 14 a Relaxation times are in s, uncertainty & 5 YO.See table 2.F. Heatley 925 Table 5. Best-fit correlation parameters [eqn (2)] from the H-1' relaxation data in table 4" concen tra- tion' R b T/K S 2 t,/ns t,/ps dev.c 2.19 2.16 2.1 1 2.1 1 0.206 0.222 0.212 0.2 14 0.025 0.021 0.37 0.78 2.9 11.2 0.18 0.98 2.5 12.1 1 .o 10.9 296 0.76 6.6 307 0.80 4.8 296 0.69 3.7 307 0.71 3.5 296 0.61 2.6 307 0.48 1.6 296 0.34 2.2 307 0.20 1.8 296 0.79 2.8 296 0.71 2.3 296 0.59 1.8 296 0.31 1.6 296 0.68 1.9 296 0.32 1.5 190 1.3 190 98 1.5 160 180 2.2 150 180 3.6 130 220 110 1.3 150 110 4.4 130 120 4.6 a Where an r.m.s. deviation is entered, four relaxation times were fitted. For the other systems, three relaxation times were fitted exactly.See table 2. See table 3. of the much greater sensitivity of 'H n.m.r. compared to 13C, it was possible to investigate a wider range of systems, especially those of low AOT concentration. Qualitatively the 'H relaxation behaviour parallels that of 13C in that the AOT motion becomes less restricted as the water/AOT ratio increases. The 'H results also show that for a given water/AOT ratio the AOT motion becomes slightly less restricted as the concentration decreases. For some of the systems in table 3, four independent relaxation times were available, and an extended test of eqn (2) was possible using the same best-fit criterion as for the 13C analysis. The derived correlation parameters are given in table 5, with the r.m.s. deviation, The quality of the simulation is within experimental error.For the other systems, only three relaxation times were available which can be fitted exactly by a unique set of the three correlation parameters. The good quality of the overdetermined four-item fits justifies accepting these unchecked three-item fit parameters as reasonably accurate. This analysis was restricted to H-1' relaxation data because of the problems associated with H- 1 mentioned above. The numerical results confirm the qualitative conclusions described above. If the molecular dynamics can be represented as ' fast ' and ' slow ' processes, the values of zf and z, from lH and 13C relaxation should be reasonably consistent. This is so for those systems appearing in both tables 3 and 5, within the limits of the fitting procedure.Some variation in S 2 may be expected because S 2 depends on the range of orientations allowed to an internuclear vector by the fast process. S 2 for the H-1' geminal interaction tends to be larger than for the C-H interactions, particularly at the higher values of R. Comparing systems in table 5 with the same R value but different AOT concentrations, it appears that when R > 2, z, is independent of concentration, but at lower values of R, z, increases with increasing AOT concentration. z, is effectively independent of composition.926 Relaxation in Aerosol OT Table 6. Comparison of experimental and simulated (using parameters in table 5) values of relaxation parameters involving HXa T"ff NOE,.{W NOEX{ AB} 1x concentra- tionb Rb exptl calcd exptl calcd exptl calcd 2.19 0.37 0.69 0.67 0.43 0.22 1 .o 1.36 2.16 0.78 0.78 0.82 0.51 0.48 0.95 1.59 2.1 1 2.9 0.93 0.931 0.67 0.64 0.85 1.22 2.11 11.2 0.997 0.986 0.913 0.903 0.95 1.04 0.222 0.98 0.69 0.58 a All values are at 296 K and 300 MHz.See table 2. As a further check on the validity of this interpretation of the 'H relaxation, calculated values of relaxation parameters involving the H-1 proton have been compared with experiment. These parameters were NOE,,{X}, NOE,{AB} and an X longitudinal 'relaxation time'. Because of the non-exponential nature of the X longitudinal relaxation described above, a single relaxation time cannot be defined, nor is it easy to measure the initial or final slopes. For the purposes of this exercise, an effective 'relaxation time', Tf! was defined as the time constant obtained from the slope of a linear least-squares fit to a log(time) plot for the X recovery over a period of time equal to the AB relaxation time.This quantity is easily obtained experimentally and theoretically. The comparison is presented in table 6. There is excellent agreement between the calculated and experimental values of NOE,,{X), a parameter which depends principally on the relative A-B, A-X and B-X distances. It therefore appears that the geometrical assumptions are reasonable and that all three interactions have essentially the same correlation function. The agreement for the other parameters, NOE,{AB} and T;;, is less good, but these are more susceptible to uncertainties in the external relaxation contribution.Discussion We consider first the relatively poor quality of fit given by eqn (2) to the 13C data for the microemulsions with R = 0.37. There are two possible reasons for this deficiency. The first is that eqn (2) does not adequately represent the spectral density for both and q. The simplest modification, suggested by Carnali et al.,' is to add a third term to give a three-stage loss of correlation. If this third process is very slow (w:z," 9 1) it could make a significant contribution to T, without materially affecting T,, even if its contribution to averaging the dipolar interaction is very small. However, it is difficult to find an acceptable physical explanation for such a slow process. Those processes contributing to z, have been proposed as overall aggregate tumbling and molecular diffusion around the aggregate surface.' A further possibility is exchange either with free surfactant or with surfactant in other aggregates.All these processes would lead to complete loss of correlation, and any slower process would not be separately detectable. It is perhaps more likely that the form of the local motions is such that eqn (1) no longer represents the correlation function satisfactorily. It would be necessary to obtain further relaxation data over a greatly extended frequency range to test this possibility. The second explanation is that an additional relaxation mechanism contributes to q, such as exchange between states of different chemical shift. An example relevant to micelles is proton exchange between a carboxylic acid and carboxylate anion which has been shown to contribute to the T, relaxation of the methylene group adjacent to theF.Heatley 927 Table 7. Order parameters for rotation about an axis axis c,-c,, interaction C,-SO, bisector C,-H, 0.1 1 0.25 C,,-H, 0.1 1 0.25 C,,-H, 0.1 1 0.25 H,-H, 0.25 0.0 16 carbonyl in sodium octanoate micelles if the acid form is not suppressed by an excess of alkali.’* The contribution to q from such a process is given by19 where PA and P, are the fractional populations of the two sites, 6v is the frequency difference between the sites, and k,, is the (first-order) rate constant for A+B. If such a process is responsible for the deviations in T,, the exchange contributions to T, may be obtained from the difference between the experimental q and that calculated using the motional parameters fitting the and NOE only.The exchange contributions were found to be 127 and 65 ms for C-1 and C-l’, respectively. That the exchange contribution is twice as great for C-1’ as for C-1 is difficult to interpret, since such a process is most likely to involve the SO, group, hence producing a larger 6v for C-1 than for C-1’. The fact that the deviations in T, are in the same ratio as the number of protons is perhaps not fortuitous, and supports the idea that eqn (2) inadequately represents the dipolar correlation function under certain conditions. and NOE data at 300 K for p-xylene microemulsions containing 6.08 mol AOT (kg p-xylene)-l and R = 3.7G33.8. Comparing the R values closest to those used here, for C-1 and C-1’ at R = 3.76 their analysis gave S 2 z 0.8, z, = 2.1 kO.25 ns and z, = 37+ 13 ps, while at R = 8.46 their analysis gave S2 x 0.8, z, = 0.9k0.1 ns and zf = 60+40 ps.The values of z, are comparable to those reported here, confirming the lack of any significant concentration dependence when R 2 2. However, the values of S 2 are rather larger and the values of zf rather smaller. Carnali et pointed out that their values of z, were not entirely compatible with expectations based on a model of aggregate tumbling plus monomer diffusion. The present results are also inconsistent with this interpretation. Note particularly that the decrease in z, with increasing R in table 3 runs counter to evidence from viscosity that the droplet radius increases with increasing R in toluene.20 It is easier to understand the behaviour of z, if the slow process originates in an ejection of a surfactant molecule from the aggregate into the hydro- carbon phase.In the free state the molecule undergoes several rotations, more or less isotropically, before rejoining an aggregate. The rate-determining step for the average loss of correlation will be the slower ejection step, which could well be independent of concentration or R as long as the surfactant molecules are not tightly bound in the aggregate. The increase in z, with decreasing R and the high values of S 2 when R is low are both consistent with a more rigid droplet structure. Even at high R, the values of S2 in these microemulsions are very much higher than those in normal m i c e l l e ~ ~ - ~ ” - ~ ~ (typically ca.0.05 at the head-group) indicating that the AOT local motions are much more restricted in amplitude. A simple model for the local Carnali et a1.8 have derived motional parameters from 13C928 Relaxation in Aerosol OT motion is diffusion within a cone of semi-angle a, for which S2 is given by21 S 2 = [cos a (cos a + 1)/212. (12) The order parameters in table 3 correspond to angles a ranging from ca. 25 to 60" according to water content. An alternative model is rotation about an axis, for which S 2 is given by21 where 9 is the angle between the internuclear vector and the rotation axis. On examination of the predominant conformation of AOT (I11 in fig. I), it appears that the most likely axes for such a rotation are either the C-SO, bond or the bisector of the C-2 and C-2' dihedral angle.The former would minimise disturbance of polar interactions while the latter would minimise resistance to motion of the alkyl chains. The order parameters calculated for rotation about these two axes are given in table 7, from which it can be seen that the order parameters for the C,-SO, axis are more consistent with the results for the systems of high R. S 2 = g3 COS2 8- (13) I am grateful to Prof. G. C. K. Roberts and Dr Ly-yun Lian of the University of Leicester Biological N.M.R. Centre for making the AM-500 spectrometer available, and to the S.E.R.C. for grants towards the other spectrometers, References 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 H. Wennerstrom, B. Lindman, 0. Soderman, T. Drakenberg and J. B. Rosenholm, J. Am. Chem. SOC., 1979, 101, 6860. T. Ahlnas, 0. Soderman, C. Hjelm and B. Lindman, J. Phys. Chem., 1983, 87, 822. H. Walderhaug, 0. Soderman and P. Stilbs, J. Phys. Chem., 1984, 88, 1655. 0. Soderman, H. Walderhaug, U. Henriksson and P. Stilbs, J. Phys. Chem., 1985,89, 3693. P. Stilbs, 0. Soderman and H. Walderhaug, J . Magn. Reson., 1986, 69, 411. U. Olsson, 0. Soderman and P. Gutring, J . Phys. Chem., 1986, 90, 5223. H. Nery, 0. Soderman, D. Canet, H. Walderhaug and B. Lindman, J . Phys. Chem., 1986, 90, 5802. J. Carnali, B. Lindman, 0. Soderman and H. Walderhaug, Langmuir, 1986, 2, 51. J. E. Ellena, R. N. Dominey and D. S . Cafiso, J . Phys. Chem., 1987, 91, 131. 0. Soderman, U. Henriksson and U. Olsson, J . Phys. Chem., 1987, 91, 116. T. Ahlnas, G. Karlstrom and B. Lindman, J . Phys. Chem., 1987, 91, 4030. 0. Soderman and U. Henriksson, J . Chem. SOC., Faraday Trans. I , 1987, 83, 1515. F. Heatley, Hoon Hong Teo and C. Booth, J. Chem. SOC., Faraday Trans. I , 1984, 80, 981. F. Heatley, J . Chem. SOC., Faraday Trans. 1, 1987, 83, 2593. F. Heatley, J . Chem. SOC., Faraday Trans. I , 1987, 83, 517. L. G. Werbelow and D. M. Grant, J. Chem. Phys., 1975, 63, 4742. R. R. Vold and R. L. Vold, J. Chem. Phys., 1976, 64, 320. J. H. Noggle and R. E. Schirmer, The Nuclear Overhauser Effect (Academic Press, New York, 1971). A. Allerhand and H. S . Gutowsky, J . Chem. Phys., 1965, 42, 1587. R. A. Day, B. H. Robinson, J. H. R. Clarke and J. V. Doherty, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 132. 0. W. Howarth, J. Chem. SOC., Faraday Trans. I , 1979, 75, 863. Paper 8/01847G; Received 16th May, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500917

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1989, 85(4), 929-943 Carbon Monoxide and Carbon Dioxide Adsorption on Cerium Oxide studied by Fourier- transform Infrared Spectroscopy Part 1.-Formation of Carbonate Species on Dehydroxylated CeO, at Room Temperature Can Li,? Yoshihisa Sakata, Toru Arai, Kazunari Domen, Ken-ichi Maruya and Takaharu Onishi" Research Laboratory of Resources Utilization, Tokyo Institute of Technology, 4259 Nagatsuta, Midori-ku, Yokohama 227, Japan The adsorption of CO and CO, on cerium oxide has been studied by Fourier- transform infrared spectroscopy (F.t.i.r.). For CO adsorption at room temperature, in addition to linearly adsorbed CO (21 77 and 21 56 cm-l), two kinds of carbonate (unidentate: 854, 1062, 1348 and 1454 cm-* and bidentate: 854, 1028, 1286 and 1562 cm-') and inorganic carboxylate (1310 and l m cm-l) species were identified spectroscopically.As for CO, adsorption, apart from weak bands at 1728, 1396, 1219 and 11 32 cm-' attributed to bridged carbonate species, bands due to unidentaie carbonate, bidentate carbonate and inorganic carboxylate species, similar to those arising from CO adsorption, were observed. Except for the linearly adsorbed-CO, all species arising from CO and CO, are stable at room temperature in uacuo. The desorption of these species at elevated temperatures shows that the order of thermal stability is bridged carbonate < bidentate carbonate < inorganic carboxylate < unidentate carbonate species, and the residual of unidentate carbonate species can remain on the surface up to 773 K under evacuation.Forming carbonate and inorganic carboxylate species proved that the CeO, surface could be partially reduced by CO even at room temperature. No bands in the region 2300-800 cm-' were detected below 373 K for CO adsorption on hydroxylated CeO,. This indicates that CO adsorption depends on the degree of dehydroxylation of the surface. The mechanism of CO adsorption is also discussed. CO is a useful probe in the characterization of surface properties of both metals' and metal oxides2 via adsorption. In addition to surface investigations, the subject of CO adsorption is closely related to the studies of reactions involving CO, such as CO hydrogenation and CO oxidation. Therefore, studies on CO adsorption are an important field and will be continued in the future.Although CO adsorbed on metal oxides has been extensively studied by infrared spectroscopy,2. infrared-spectroscopic studies of CO adsorption on cerium oxide have not been carried out in detail until now. Some authors4* have studied CO adsorption on ceria-supported metal catalysts ; however, their main attention was focused on the metals instead of on ceria. Recently, Jin et aL6 reported the adsorption of CO and CO, on Pt/Ce02 and CeO, by infrared spectroscopy and other techniques, but they failed to find any peaks corresponding to CO on pure CeO, in the region 1000-3000 cm-'. For the above reasons, i.r. studies of CO adsorption on CeO, have been initiated in 7 Permanent address: Dalian Institute of Chemical Physics, Chinese Academy of Sciences, 129 Street, Dalian, China.929930 CO and CO, Adsorption on CeO, our laboratory. We have found that CO not only is adsorbed in a linear form as on other oxides, but also interacts strongly with the CeO, surface, forming a considerable number of carbonate species even at room temperature. A similar phenomenon on oxides has been discussed previously. '-11 In this paper i.r. spectra of adsorbed CO on cerium oxide are displayed. In order to assign those bands arising from adsorbed CO, the adsorption of 13C0 and CO, on cerium oxide has also been studied by Fourier-transform i.r. spectroscopy. Furthermore, the thermal stability of the species obtained from adsorbed CO and CO, has been investigated by in situ F.t.i.r. Finally, the results of CO adsorption on hydroxylated CeO, and the mechanism of forming carbonate and carboxylate species are discussed.Experiment a1 Materials Cerium oxide (ceria or CeO,) used in this study was obtained by the thermal decomposition of cerium hydroxide gel at 773 K in air for 3 h. The cerium hydroxide was prepared through precipitation from an aqueous solution of cerium(Ir1) nitrate with NH,OH at pH 8-9, and the resulting precipitate was repeatedly washed with deionized water until NH; and NO, ions were eliminated, prior to drying and calcination. H, was purified via a Deoxo arrangement and then through a liquid-nitrogen trap. 0,, CO and 13C0 (99.7%) were refined through a liquid-nitrogen trap in a circulating system for 15 min. CO, was prepared from the decomposition of NaHCO,. Apparatus 1.r.spectra were recorded on a JEOL JIR-100 F.t.i.r. spectrometer with 256 scans at 4 cm-l resolution using a liquid-nitrogen-cooled HgCdTe detector. All the spectra shown in the paper are in absorbance, and their backgrounds were recorded before admitting the adsorbate gas under corresponding experimental conditions of the spectra. The same calcined at 773 K was pressed into a self-supporting disc, weight ca. 150 mg, with a diameter of 20 mm. The disc was placed in a conventional i.r. cell with NaCl windows and a furnace capable of heating the disc in situ up to 1100 K. The i.r. cell, connected to a vacuum system can be kept below lop4 Torr (1 Torr w 133.3 N m-,) by an oil-diffusion pump and a mechanical pump. Procedure In order to eliminate the species contaminants, adsorbed H,O and hydroxyl groups prior to the first experiment, a new disc in the cell was outgassed at 1000 K for 10 min and then treated in 0, at 873 K for 3 days (the time was shortened to 12 h for subsequent experiments).After pretreatment the sample was outgassed again at 1000 K for 30 min. The i.r. spectra showed that no peaks of contaminants, residual water or OH groups remained on the surface after pretreatment and outgassing. Therefore, unless noted otherwise, the sample used in this study is a dehydroxylated CeO,, symbolized as CeO,( 1000 K). Continuing the above processes, the disc was progressively cooled to the given temperatures for recording background spectra and finally to room temperature for CO or CO, adsorption. The thermal-desorption and CO-adsorption experiments at elevated temperatures wee performed in a stepwise way and the temperature rate of an interval was ca.25 K min-l. A given temperature was not changed until the spectra reached a steady state. A hydroxylated CeQ, surface was obtained by treating the sample which had undergone the above pretreatment in water vapour from room temperature to 673 K and at 673 K for 1 h, then outgassing for 1 h at the same temperature. The i.r. spectraC. Li, Y. Sakata, T. Arai, K. Domen, K-i. Maruya and T. Onishi 93 1 2300 2100 1900 2000 1600 1200 800 wavenumberfcm- ' Fig. 1. 1.r. spectra of CO adsorbed on CeO,(l000 K) at room temperature: (A) with decreasing CO pressure, (a) 110, (b) 87 and (c) 21 Torr and ( d ) outgassing for 1 min after (c); (B) after admission of CO for (a) 5, (6) 20 and (c) 124 min.of the hydroxylated sample showed that two sharp peaks at 3664(strong)cm-' and 3629(weak) cm-l, due to different terminal OH groups, and a broad peak centred at ca. 3427 cm-l, attributed to hydrogen-bonding OH groups, were reproduced, indicating that the CeO, surface was truly hydroxylated. Results CO Adsorption at Room Temperature Fig. 1 shows the i.r. spectra of CO adsorbed on CeO, at room temperature, and the results are described in two sections as follows. In the range 230&1900 cm-' in fig. 1 A, two bands at 2177 and 21 56 cm-', attributed to CO linearly adsorbed on Ce", were observed. The intensities of the two bands decrease simultaneously with a decrease in the CO pressure [fig. 1 A(a)-(c)] and both bands disappear quickly on evacuation at room temperature [fig.1 A(d)]. The dependence of the band intensities upon the pressure of CO agrees with a Langmuir relationship, and the CO pressure for saturated adsorption is ca. 100Torr at room temperature. The two bands were gradually weakened at elevated temperatures even in a CO atmosphere (1 10 Torr) and completely vanished at ca. 350 K. These results suggest that the two bands are associated with weakly adsorbed CO and are similar to those on other oxides. In the region 1800-800 cm-', as soon as CO was introduced into the i.r. cell, nine distinct peaks at 854, 1028, 1062, 1286, 1310, 1348, 1454, 1510 and 1562 cm-' were detected as shown in fig. 1 B(a). All the bands grew slowly with time (fig. 1 B) and it took ca.200 min to reach saturated adsorption at room temperature. Various CO pressures in the range 1&100 Torr made no apparent difference both to the spectra and to the time of saturated adsorption (thus hereafter pressures of 10-20 Torr of CO and 13C0 were932 CO and CO, Adsorption on CeO, 4 - A h - A - 0.95 0.85 rD 00 2 -4 3 0.75 0.65 0.5 5 10 60 110 160 tlmin Fig. 2. Variations of i.r. band intensities of CO adsorbed on CeO, (1000 K) at room temperature as a function of prolonged time: A, 1286; A, 1562; 0, 1348; @, 1454 and 0, 1510 cm-I. adopted). All the bands become stronger and more distinguishable with prolonged adsorption time, except that at 1310 cm-l, which was submerged by two intense neighbouring bands at 1286 and 1348 cm-l. Comparing the bands in the two regions, 2300-1900 and 1800-800 cm-', the former are rapidly formed in a CO atmosphere and easily outgassed by pumping; however, the latter are produced slowly and cannot be removed by evacuation at room temperature. This indicates that the surface species related to the bands in the two regions are appreciably different from each other in nature.With a view to distinguishing the species with i.r. bands in the region 1800-800 cm-l, the variations of five main bands with adsorption time after admission of CO are presented in fig. 2, in which the scale is the ratio of absorbance of the indicated band to that of the 1286 cm-l band. The bands at 1286 and 1562 cm-l grow at the same rate; hence their ratios are parallel. The two curves at 1348 and 1454 cm-' are nearly coincident; the two bands keep pace with each other and obviously increase the most quickly among the five bands.The curve of 1510 cm-' is different from those of other bands in slope. On the basis of a knowledge that the bands from the same species will exhibit the same tendency during a variation in the amount of species present, the five bands can be reasonably classified into three groups due to three kinds of surface species, here termed (A) 1286 and 1562 cm-l, (B) 1348 and 1454 cm-' and (C) 1510 cm-l. Isotopic-shift methods are often used to clarify and confirm problems in the interpretation of i.r. spectra. The i.r. spectra of adsorbed 13C0 on CeO, are basically the same as those of adsorbed CO apart from a reasonable isotopic shift.The bands of adsorbed 13C0 correspond to those of adsorbed CO, and the isotopic shifts and ratios, are listed in table 1.C. Li, Y . Sakata, T. Arai, K. Domen, K-i. Maruya and T. Onishi 933 Table 1. Vibrational frequencies (below 1800 cm-l) of l2C0 and 13C0 adsorbed on CeO, (crn-l). isotopic isotopic l2C0 13C0 shifts ratios 854 (m) 1028 (m) 1062(w, b) 1286 (s) 1 3 10 (w) 1348 (s) 1454 (s) 1510(s) 1562 (s) 827 (m) 1024 (m) very weak 1256 (s) 1279 (w) 1319 (s) 1410 (s) 1475 (s) 1518 (s) 27 4 30 31 29 44 35 44 - 1.033 1.004 1.024 1.024 1.022 1.03 1 1.024 1.029 - a w, weak; m, medium; s, strong; b, broad. ; I : I l l I I I I I I I 2000 1600 i200 800 wavenumber/cm-' Fig. 3. 1.r. spectra of CO adsorbed on CeO, (1000 K) at elevated temperatures in the presence of 10 Torr CO: (a) 300 K for 60 min; (b) 373 K for 35 min after (a); (c) 473 K for 20 min after (b) and (d) 573 K for 30 min after (c).CO Adsorption at Elevated Temperatures Fig. 3 displays a series of i.r. spectra of adsorbed CO formed at various temperatures. Spectra were recorded at elevated temperatures stepwise from room temperature to the indicated temperature in the presence of 20 Torr CO. At room temperature bands at 1286 and 1562 em-' appear at the beginning of CO adsorption and are dominant in the spectra. With increasing temperature the bands at 1348, 1454 and 15 10 cm-l become934 CO and CO, Adsorption on CeO, 0.7 0.6 0.5 0.4 8 e ' 0.3 0.2 011 0 I I I 273 373 473 573 673 TIK Fig. 4. Increase in i.r. band intensities of 1286 (O), 1348 (A) and 1510 (0) cm-' in the presence of 10 Torr CO at elevated temperatures (data from fig.3). significantly strong, such that the 1286 are 1562 cm-' bands are almost overwhelmed at 573 K. The weak bands at 854 and 1062 cm-l also grow, apparently at the same time. This implies that the two bands may be ascribed to species B or C instead of species A. The band at 1028 cm-l seems to be in step with the 1286 and 1562 cm-l bands during this process. Fig. 4 quantitatively illustrates the enhancement of the three main bands at 1286, 1348 and 1510 cm-', which are taken as representatives of the three species A, B and C, respectively. The bands at 1348 and 1510 cm-l grow steeply with increasing adsorption temperature, and their intensities vary more than tenfold in the range 300-573 K, while the 1286 cm-l band only doubles in intensity from room temperature to 473 K, and shows no obvious change above 473 K.These results clearly suggest that the activation energies for formation of species B and C are higher than that for species A, and that species A might reach saturated adsorption at 473 K. CO, Adsorption at Room Temperature In an attempt to assign the bands arising from CO adsorption, an i.r. study of CO, adsorption on CeO, was performed under the same conditions as in CO adsorption, and the spectra of CO, adsorbed at room temperature are presented in fig. 5. Adsorbed CO, gives more and stronger bands than adsorbed CO in the region 2000-800 cm-', but the main peaks in the spectra for CO, and CO adsorption are very similar, i.e. bands at 856, 101 1, 1045, 1286, 1354, 1454, 1506 and 1568 cm-l stemming from adsorbed CO, resemble in position and relative intensity corresponding bands at 854, 1028, 1062, 1286,C.Li, Y. Sakata, T. Arai, K. Domen, K-i. Maruya and T. Onishi 93 5 I I I I I 2000 1600 1200 800 wavenumberlcm-' Fig. 5. 1.r. spectra of CO, adsorbed on CeO, (1000 K) at room temperature after admission 1 Torr CO, for (a) 1 , (b) 3, (c) 5 and ( d ) 125 min. 1348, 1510 and 1562 em-' for adsorbed CO. In addition to the abovementioned bands, four weak bands at 1 132, 1219, 1396 and 1728 cm-l appear for CO, adsorption but are absent for CO adsorption. Another apparent feature is that all the bands increase in intensity rapidly with time of exposure of CO,, and reach a maximum within 10 min. Desorption of Adsorbed CO and CO, An experiment on the desorption of adsorbed CO and CO, was undertaken with a view to examining the thermal stability of surface species and to investigate the behaviour of different species during the course of desorption.Fig. 6 shows a series of spectra of adsorbed CO recorded at the temperatures indicated after achieving a steady state of desorption. As shown in fig. 6(a)-(c), bands at 1028, 1286 and 1562 cm-' are reduced remarkably in the same step by heating the sample from room temperature to 473 K, while in contrast bands at 1062, 1348 and 1454 cm-' grow slightly instead of decrease, and the 1510 cm-l band is almost unaffected. Above 473 K, all the bands decrease with further increasing temperature, and the bands at 1062, 1286, 1562 and 15 10 cm-' are almost removed at 673 K, while the bands at 854, 1348 and 1454 cm-' survive a thorough elimination even up to 773 K [fig.6(d)-(f)]. These results confirm the conclusion that the bands below 1800 cm-I are due to three kinds of surface species distinguished by their variations in behaviour when subjected to warming and degassing. The weak bands at 1028 and 1062 cm-l might belong to species A and B, because their behaviour is in accord with bands at 1286, 1562 and 1348 and 1454 cm-I, respectively.936 CO and CO, Adsorption on CeO, I I I I I 2 000 1600 1200 800 w avenumber/cm- ' Fig. 6. 1.r. spectra of CO adsorbed on CeO, (1000 K) in the course of desorption at elevated temperatures in uacuo: (a) 300 K outgassing for 25 min after 100 rnin contact with 20 Torr CO, (b) 373 K for 60 min after (a), (c) 473 K for 30 min after (b), (d) 573 K for 20 min after (c), (e) 673 K for 30 rnin after (d) and (f) 773 K for 20 min after (e).The band at 854 em-' is unique during desorption. In the range 300-473 K it becomes weaker, accompanying the decrease of species A (1028, 1286 and 1562 cm-l); how- ever, in the range 473-573 K its behaviour coincides with species B (1062, 1348 and 1454 cm-l). Therefore it seems reasonable to attribute the 854 cm-l band to both species A and B. The variations of the three species A, B and C marked by the bands at 1286, 1348 and 15 10 cm-l, respectively, are illustrated more clearly in fig. 7. Species B is more stable than species A and C, and the order of their thermal stabilities will be B > C > A.The increase in species B is only due to the conversion of species A, since species C remains the same and there is no further addition of CO to the gas phase during the process. By combining the results shown in fig. 3 and 4, we concluded that the species A may be an intermediate of species B and C. As soon as CO was introduced onto CeO,, only species A was formed on the surface initially; it was then converted into species B and C slowly (fig. 1 B and 3). As a consequence, a high temperature should facilitate the conversion of species A into B and C. This must be the reason why there is an increase of B and no loss of species C on going from 300 to 473 K in vacuu (see fig. 7).C. Li, Y. Sakata, T. Arai, K. Domen, K-i. Maruya and T.Onishi 937 0.6 0.5 0.4 8 5 e 0.3 2 % a2 0.1 2 73 373 473 573 673 TIK Fig. 7. Variations of i.r. band intensities of 1286 (a), 1348 (A) and 1510 (0) cm-' in the course of desorption at elevated temperatures (data from fig. 6). The desorption of adsorbed CO, is very like that of adsorbed CO, except for four weak bands at 1132, 1219, 1396 and 1728 cm-l which are not present in CO adsorption and disappear simultaneously on heating to 373 K. This suggests that, in addition to the three species A, B and C which are the same as formed from CO adsorption, at least one other species is produced from CO, adsorption ; species (named D) may be generated via weakly adsorbed CO,, and differs from species A, B and C in its thermal stability. The same variation in the spectra of adsorbed CO and CO, strongly supports the view that the same species A, B and C could be formed from both CO and CO, adsorption on dehydroxylated CeO,. CO Adsorption on Hydroxylated CeO, As previously described, hydroxylated CeO, was prepared by heating CeO, in the presence of H,O vapour.The hydroxylated surface was exposed to 20Torr CO at 300 K for 120 min and then heated to 373 K for 20 min, while no peaks due to adsorbed CO were observed at the two temperatures, as shown in fig. 8(a) and (b). On progressively heating the sample above 473 K, besides very weak bands at 2939, 2848, 1576, 1558, 1369 and 1307 cm-l arising from formate species produced via the reaction of adsorbed CO and surface OH groups, similar bands to those in fig. 3 appeared, but all the band intensities were much weaker than those on dehydroxylated CeO,.This enabled us to suggest that the surface OH groups hinder the CO from forming adsorbed species, especially at low temperatures.938 CO and CO, Adsorption on CeO, 4000 3000 2000 1600 1200 800 wavenumber/cm-' Fig. 8. 1.r. spectra of CO adsorbed on CeO, (hydrated at 673 K) in the presence of 20 Torr CO at elevated temperatures: (a) 300 K for 120 min, (6) 373 K for 20 min after (a), (c) 473 K for 25 min after (b) ( d ) 573 K for 22 min after (c) and (e) 673 K for 1 min after (d). Discussion Linearly Adsorbed CO On most of the oxides the i.r. bands from CO adsorption appear in the region 230&2000 cm-l, and their wavenumbers are often substantially higher than that of gas- phase CO (2143 cm-').The explanation of this phenomenon is controversial, but two popular corollaries seem reasonable :' (1) these bands are attributed to CO linearly adsorbed on surface-exposed metal ions, e.g. Lewis-acid sites; (2) the more positive the state of the metal ions, the higher is vibrational frequency of adsorbed CO. In the light of these views, we assigned the band at 2177 cm-l to CO linearly adsorbed on Ce4+ and the band at 21 56 cm-l to Ce4+ in a more unsaturated coordination state created by severe degassing at 1000 K. No linearly adsorbed CO was formed on hydroxylated CeO, because the surface was almost covered with OH; consequently no metal ions were exposed to the surface for CO coordination. Assignment of the Bands below 1800 em-' Generally speaking, i.r.bands below 1800 cm-' arising from adsorbed CO, are due to carbonate-like species. Extensive i.r. studies on the assignment of the spectra in this region have been made for CO, and CO adsorption on transition-metal oxides, since the pioneer work on NiO by Eichens and Pliskin.12 Fujita et a1.13 calculated the presence of two kinds of carbonates in a Co"' complex, and their results have been considered as criteria for the assignment of carbonate species3. 14, l5 on the surface of metals and oxides. On CeO,, only Guenin" and Jin' have reported i.r. studies of CO, adsorption. Table 2 lists some assignments of carbonate-like species related to the present work. By comparing our results presented above with those displayed in table 2, one may assignC. Li, Y.Sakata, T. Arai, K. Domen, K-i. Maruya and T. Onishi 939 Table 2. Assignment of carbonate-like species (from the liter at we) a frequency/cm-l assignment ref. CO,/CeO, COJCeO, calculated results of CO(II1) carbonate complex generally 1670-1 695 (6) C 1310-1338 650-970 1590-1630 -'\ 1260-1270 0 1020-1030 ,- do T+ 0 -c : 1560 1410 1470-1530 /.S=O 1040-1080 T- 0 1300-1370 - 0-c: 1580-1290 bidentate carbonate ( 1 6) 1680- 1240 bidentate carbonate 1480- 1 370 uniden tate carbonate 1405 inorganic carboxylate inorganic carboxylate (symmetric) 1 570- 1 360 1483 CO; (as) / O \ , 9 ( 1 3 ) 1373 CO; (sy) co c .eM 1030 C=O \. 0 0 1595 C=O 1282 CO; (as) Co 1038 CO;(sy) \ / (3) -O\ c=o 1780 C=O 1260 CO; (as) 1020 co; (sy) - 0 / -~ bridged carbonate sy, symmetric vibration ; as, asymmetric vibration.species A to a bidentate form and species B to unidentate form, since the bands of species A and B in this study are in excellent agreement with those of bidentate and unidentate carbonates, respectively, in CoI" carbonate complexes. The band at ca. 854 cm-l has been detected in complexes and inorganic compound^,'^^^^ and was attributed to the out- of-plane vibration of the carbonate COi- group, but it is difficult to detect this species in adsorption experiments because of the low transmission of oxides below 1000 cm-'. Our postulation that the band at 854 cm-l is due to both species A and B is confirmed, since both the bidentate and unidentate carbonate species possess a C0:- group. Its isotopic shift of 27 cm-' on substitution by 13C given in table 1 is also in accord with the isotopic shift for Ca13C0, in comparison with Ca12C03.1s The bands at 1510 cm-l (1506 cm-l from adsorbed CO,), previously grouped as 32 F A R I940 CO and CO, Adsorption on CeO, Table 3.Assignment of i.r. bands (below 1800 cm-l) due to CO and CO, adsorbed on CeO, species frequency/cm-l assignment CO (CO,) A 854(856) 0 on CeO, 1028 (1011) Ce ' ' C E O B 854(856) / O \ do \ O / 1286 (1 286) 1562 (1 568) bidentate carbonate 1062 (1045) Ce C ' 1348 (1354) \. 1454 (1454) unidentate carbonate C 1310 (very weak) Ce-C 1510 (1506) +.o inorganic carboxylate CO, on D 1132 CeO, 1219 1396 0 1728 bridged carbonate species C, and 1310 cm-l (weak) may be attributed to inorganic carboxylate rather than carbonate species, since no band in the 1200-1000 cm-l region is in agreement with them.The recognition of inorganic carboxylate species on oxide surfaces is difficult because their bands are usually confused with those of carbonate species. Fortunately, on CeO, (as shown in fig. 1 B and 3) the band at 1510 cm-l(l506 cm-l in fig. 5 ) assigned to an asymmetric vibration is well resolved, and the band at 1310 cm-l ascribed to a symmetric vibration can be clearly identified at the beginning of CO adsorption, although it is very weak and is overwhelmed by neighbouring bands eventually. From the table summarized by Little,3 the inorganic carboxylate species was identified by two bands whose positions vary in the regions 1570-1 510 and 1410-1 310 cm-l, respectively, on different oxides; in addition, the band at 1410-1310 cm-' is weak or absent on some oxides.The two bands at 15 10 and 13 10 cm-' in our study fit well with the above features of inorganic carboxylate species. The assignment concerning carbonate-like species on CeO, may be further verified by the results (in table 1) for isotopic shifts. Among the three bands of species A, i.e. the bidentate carbonate species, the biggest isotopic shift is 44 cm-l for the 15 18 cm-l band. This is close to that for the stretching frequency of the l3C_O molecule, which is estimated to be ca. 48 cm-', so the band arises from the C=O stretching vibration in bidentate carbonate species. The 1024 cm-l band is hardly shifted; this confirms that the carbon atom is unaffected in the vibration. Therefore, it should be ascribed to a symmetric vibration.The band at 1256 cm-', with an isotopic shift of 30 cm-I, is due to an asymmetric vibration, since the carbon atom participates in the vibration but is not involved to the same extent as in the C=O vibration. On the same principle, the isotopic shift of the asymmetric vibration is larger than the symmetric vibration in the carbonate; the two bands at 1410 and 1319 cm-' with isotopic shifts of 44 and 29 cm-l, respectively,C. Li, Y. Sakata, T. Arai, K. Dornen, K-i. Maruya and T. Onishi 94 1 are correspondingly attributed to the asymmetric and symmetric vibrations of unidentate carbonate species. For inorganic carboxylate species the isotopic shifts of 15 18 and 1279 cm-l are closed to each other, not as in carbonate species.As the inorganic carboxylate is a triatomic group the carbon atom is not at the centre as in the carbonate; therefore the carbon atom takes part in both the asymmetric and symmetric vibrations, as is the case for the oxygen atom in a water molecule. This is the reason why the asymmetric and symmetric vibrations show an isotopic shift to the same extent. Species D, with bands at 1728, 1396, 1219 and 1132 cm-', is ascribed to a bridged carbonate species by the characteristic band at 1728 cm-l, which is due to the C=O vibration of an organic carbonate. Furthermore, species D is weakly adsorbed on CeO, and can be eliminated upon heating to 373 K in vacuo; similar results have been reported for Cr203.19 By summarizing the above discussions, the assignment of the bands below 1800 cm-l arising from adsorbed CO and CO, on dehydroxylated CeO, are listed in table 3.1.r. spectroscopic studies of CO adsorption on MgO,' CaO and SrOlO well degassed at high temperatures have shown that adsorbed CO gives rise to a large number of bands in the range 2200-1000 cm-l. These bands were attributed to carbonate species and some unusual CO polymeric species, (C0):- (n = 2,3,4, . . ., x = 2,4, . ..). Although adsorbed CO on CeO, also exhibited rich bands in the same region, they were completely different from those of the CO polymeric species in terms of their band positions and relative intensities. Furthermore, the polymeric species were sensitive to oxygen, being easily destroyed by exposure thereto; however, no apparent change in the spectra of CO adsorbed on CeO, has been found after dosing 0, on the CeO, surface preadsorbed with CO.Apart from these differences, the increase in the intensities of i.r. bands due to CO adsorbed on CeO, at room temperature is a function of CO contact time but is independent of the CO pressure (10-100 Torr); nevertheless, the formation of the CO polymeric species is strongly dependent on the CO pressure. We therefore rule out the possibility of existing CO polymeric species on CeO, in the present study. Mechanism of CO Adsorption on Hydroxylated CeO, The formation of carbonate and inorganic carboxylate species from CO adsorption indicates that CO is oxidized by surface oxygen species; in other words, the surface of CeO, is partially reduced by CO. It is expected that CO reacts with some oxides to form carbonate species at high temperatures, but it seems impossible that the phenomenon occurs at temperature.The results on CeO, are not general for all oxides, and imply that the CeO, surface might possess either special adsorption sites or very active oxygen species, or both. Therefore a short discussion on the mechanism of CO adsorption is indispensable to the conclusions of our paper. CO oxidation on CeO, above 473 K has been investigated by Breysse et aZ.20*2r Their results confirmed the participation of lattice oxide ions during the reaction. A mechanism was proposed involving cyclic reduction/oxidation of the surface to explain CO oxidation on CeO,. Jin2, and his colleagues concluded that the formation of CO, on Pt/CeO, from CO temperature-programmed description was due to lattice oxygen from the interface between Pt and CeO,.From a temperature-programmed reaction study by Yao et aZ.,23 two peaks at 770 and 1020 K were ascribed to the reduction of surface-capping oxygen ions and bulk oxygen ions, respectively; another t.p.r. peak was due to oxygen species which could be converted into the capping and bulk oxygen ions. The oxygen species, i.e. mononuclear and molecular oxygen anions, on CeO, have also been identified by e.s.r. spectro- In the present case, we are inclined to the view that the surface oxygen species and part of the capping oxygen ions (but not the lattice oxide ions) are responsible for the formation of carbonate and carboxylate species on CeO, at room temperature, because 32-2942 CO and CO, Adsorption on CeO, the former are more easily reduced than lattice oxide ions.It is not plausible to extract lattice oxygen, especially at room temperature. It may be possible that the lattice oxide ions migrate to the surface to compensate for the surface oxygen vacancies created by CO reduction with increasing temperature; as a result, the surface carbonate and carboxylate species are enhanced at elevated temperatures, as shown fig. 3 and 4. The surface oxygen species and capping oxygen ions are necessary for forming carbonate and inorganic carboxylate species, but the more important reason for CO oxidation at room temperature should be the presence of surface-active sites that can activate CO to be oxidized easily. The large difference in CO adsorption on dehydroxylated and hydroxylated CeO, (fig.1, 3 and 8) leads us to suggest that the coordinatively unsaturated surface sites may play a key role for activation of CO to form carbonate and inorganic carboxylate species, since the surface unsaturated sites can be generated via dehydroxylation at high temperature. E.s.r. have shown that oxygen defect sites could be formed on the CeO, surface by outgassing at high temperatures, and similar e.s.r. resultsz6 postulated that Ce3+ sites may be produced after pretreatment at 773 K in uacuo. The proposal that the surface-active sites could be produced through dehydration has also been made by Ze~china,~ on Cr,O,. For hydroxylated CeO,, the surface was saturated with OH groups and was hence inactive for both CO adsorption and oxidation at room temperature.Fig. 1 B(a) and fig. 3 ( a ) clearly show that bidentate carbonate species predominate on the surface during the first stage after the admission of CO. It is assumed that CO reacts directly with surface oxygen species and capping oxygen ions to form bidentate carbonate species, and then the bidentate carbonate species as an intermediate are converted into unidentate carbonate and inorganic carboxylate species, as shown in fig. 2, 4 and 7. Conclusions The following summarizations can be made from this work. (1) CO adsorption at room temperature on the dehydroxylated CeO, surface shows the formation of four adsorbed species, linearly adsorbed CO, unidentate carbonate, bidentate carbonate and inorganic carboxylate species.(2) The linearly adsorbed CO can be removed at room temperature by pumping, and its amount is dependent on the CO pressure. The other three species are stable at room temperature and their thermal stability is in the order bidentate < carbonate < unidentate carbonate < inorganic carboxylate species. (3) The bidentate carbonate species can be converted into unidentate carbonate and inorganic carboxylate species, especially at high temperature. (4) Adsorbed CO, produces two carbonate and inorganic carboxylate species which are the same as those from adsorbed CO and bridged carbonate species. (5) Dehydroxylated CeO, can be partially reduced at room temperature by CO. (6) Surface OH groups prohibit CO adsorption. This is interpreted as indicating that there are no active sites for CO coordination. C.L.is grateful to the Ministry of Education, Science and Culture of Japan and the United Nations Educational, Scientific and Cultural Organization (UNESCO) for his acceptance as a research fellow in its 23rd International Postgraduate University Course in Chemistry and Chemical Engineering held at Tokyo Institute of Technology, Tokyo, Japan from October 1987 to September 1988. References 1 J. B. Peri, in Catalysis, Science and Technology, ed. J. R. Anderson and M. Boudart (Springer, New 2 M. C. Kung and H. H. Kung, Catal. Rev.-Sci. Eng., 1985, 27, 425. 3 L. H. Little, Infrared Spectra of Adsorbed Species (Academic Press, New York, 1966), p. 47. York, 1984), vol. 5, p. 171.C. Li, Y. Sakata, T.Arai, K. Domen, K-i. Maruya and T. Onishi 4 J. C. Summers and S. A. Ausen, J . Catal., 1979, 58, 131. 5 A. Kiennemann, R. Breault and J-P. Hindermann, J. Chem. Soc., Faraday Trans. I , 1987, 83, 21 19. 6 T. Jin, Y. Zhou, G. J. Mains and J. M . White. J . Phys. Chem., 1987, 91, 5931. 7 J. W. London and A. T. Bell, J . Catal., 1973, 31, 32. 8 K. Tanaka and J. M. White, J . Phys. Chem., 1982, 88, 4708. 9 E. Guglielminotti, S. Coluccia, E. Garrone, L. Cerruti and A. Zecchina, J . Chem. Soc., Faraday Trans. I , 1979, 75,96. 10 S. Coluccia, E. Garrone, E. Guglielminotti and A. Zecchina, J . Chem. Soc., Faraday Trans. I , 1981,77, 1063. 11 G. Busca and V . Lorenzelli, Mat. Chem., 1982, 7, 89. 12 R. P. Eichens and W. A. Pliskin, in Adtjances in Catalysis, ed. D. D. Eley, W. G. Eley, W. G. Franken- burg, V. I. Komarewsky and P. B. Weisz (Academic Press, New York, 1957), vol. 9, p. 662. 13 J. Fujita, A. A. Martell and K. Nakamoto, J . Chem. Phys., 1962, 36, 339. 14 M. L. Hair, Infrared Spectroscopy irz Surface Chemistry (Marcel Dekker, New York, 1967). 15 K . Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds (Wiley, New 16 M . Guenin, Ann. Chim. (Paris), 1973, 8, 147. 17 B. M. Gatehouse, S. E. Livingstone and R. S . Nyholm, J . Chem. Soc., 1958, 3137. 18 S. Pinchas and I . Laulicht, Infrared Spectra qf Labelled Compounds (Academic Press, New York, 197 l), 19 A. Zecchina, S. Coluccia, E. Guglielminotti and G. Ghiotti, J . Phys. Chem., 1971, 75, 2790. 20 M. Breysse, M. Guenin, B. Claudel, H. Latreille and J. Veron, J . Catal., 1972, 27, 275. 21 M. Breysse, M. Guenin, B. Claudel and J. Veron, J . Catal., 1973, 28, 54. 22 T. Jin, T. Okuhara, G. T. Mains and J. M. White, J . Phys. Chem., 1987, 91, 3310. 23 H. C. Yao and Y. F. Yu Yao, J , Catal., 1984, 86, 254. 24 M. Gideoni and M. Steinberg, J . Solid State Chem., 1972, 4, 370. 25 M. Che, J. F. J. Kibblewhite, A. J. Tench, M. Dufaux and C. Naccache, J . Chem. Soc., Faraday Trans. 26 J. L. G. Fierro and J. Soria, J . Solid State Chem., 1987, 66, 154. 27 A. Zecchina, S. Coluccia, E. Guglielminotti and G. Ghiotti, J . Phys. Chem., 1971, 75, 2774. York, 1978), p. 244. p. 206. 1, 1973, 69, 857. Paper 8/01 990H ; Receiced 19th Muy, I988

ISSN:0300-9599    

DOI:10.1039/F19898500929

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. SOC., Furuduy Trans. I , 1989, 85(4), 945-955 Isopiestic Measurement of Salt Imbibition in Zeolites Na-X and Na-Y Barrie M. Lowe Chemistry Department , University of Edinburgh, Scot land Christopher G. Pope* Chemistry Department, University of Otago, Dunedin, New Zealand Factors affecting the reliability of the isopiestic method for the measurement of the imbibition of salts from aqueous solution by zeolites have been examined. It has been found that the sharp intersection of the isopiestic curves, which is essential to the accuracy of the method, does not occur when the zeolite crystals are so small that sorption of the salt on the external surface of the zeolite is substantial. Results for the occlusion of sodium nitrate by zeolite Na-X (ca. 6 pm crystals) and Na-Y (ca.7 pm crystals) at 298.2, 313.2 and 328.2 K are reported. Those for sodium nitrate and sodium chlorate at 298.2 K are used to test the applicability of the Donnan equation. The application of the isopiestic technique to the measurements of the uptake of solutes by zeolites was introduced by Fegan and Lowe.' The method is versatile, as it can be applied to any solution containing a single involatile solute, and the experiments are simple to perform and require little equipment. The results yield the compositions of both the imbibed and external solutions directly, and it is easy to confirm that equilibrium has been established while the experiments are being performed. Fegan2 has demonstrated that accurate data can be obtained at room temperature, but the experiments tend to be rather slow and tedious.This work is concerned with a further examination of the reliability of the method, with simplifying and speeding up the procedure, and with extending the technique to higher temperatures where solvent condensation on the sample containers presents a possible problem. Experiments were carried out on the imbibition of sodium nitrate at 298, 313 and 328 K by samples of Na-X and Na-Y with different Si/Al ratios and particle size. One sample was studied in more detail at 298 K in order to examine whether the salt imbibition increased smoothly as the water activity decreased, as Fegan' has reported a sharp change of slope at the water activity of the saturated solution when sodium chloride was used, although it does not seem to be clear why such an effect should exist.Finally, experiments with sodium chlorate, sodium citrate and sodium fluoride were carried out to investigate the dependence of imbibition on the salt used. Chlorate and nitrate are of interest because of their different structure directing influence during the synthesis of sodalite and ~ancrinite,~ citrate for its charge, size and hydration, and fluoride because its reactivity with aluminium and silicon suggested that it would probably interact strongly with the zeolite lattice. 945946 Salt Imbibition in Zeolites Experimental Materials A set of Na-X and Na-Y samples were made using a common general procedure and starting materials. Only the proportions of the reactants used were changed, and this in turn resulted in different times being required to complete crystallisation.Sodium aluminate solution was prepared by dissolving AR-grade aluminium wire with AR sodium hydroxide solution in a screw-capped polythene bottle. The rest of the water required was mixed with Cabosil M-5 fumed silica in a second bottle. When the aluminium was completely dissolved, and the solution was cold, the aluminate solution was added to the silica and the mixture was blended thoroughly to a creamy consistency. This was aged for five days at 298 K, and then transferred to a water bath at 367 K where it was left unstirred until crystallisation was complete, as judged by the appearance of the suspension and optical microscope examination of the product (1-3 days). The clear supernatant liquid was then rapidly removed and the remaining slurry filtered and washed with hot water.All the preparations carried out by this method produced small irregularly shaped solid particles which appeared to be made up of agglomerates of smaller material, and were pure as judged by X.r.d. Only details of F1 and F5 at the ends of the Si/A1 composition range synthesised are given. Large crystals of essentially the same composition as F5 were made from the same starting materials using the same apparatus and temperatures by a procedure based on that reported by Kacirek and Le~hert.~ A more dilute synthesis mixture was prepared with the molar composition shown in table 1. A seeding gel was formed by following the preparative procedure used before for F5, and after a short time, a percentage of this was thoroughly mixed with the synthesis mixture, which was also at 367 K.In the product used in this work, which was pure as judged by electron microscopy and X.r.d., 1 YO of seeding mix was added after it had been aged for 26 h. Other proportions of seed and aging times also were possible, but these were not examined in detail. When the synthesis mixture was left unseeded crystallisation was much slower, and the product was a mixture of faujasite and zeolite S. Large crystals of high aluminium content were prepared by the method of K ~ h l . ~ The synthesis mixture was prepared from sodium silicate solution (1 2 % Na,O, 30 % SiO, from B.D.H.) and AR-grade aluminium wire, sodium hydroxide and potassium hydroxide. The silicate solution was added to the aluminate solution quickly, with rapid stirring at room temperature in a polythene bottle.The mixed reagents gelled in about 30 s, and were transferred to an oven at 354 K for 1 h and then to a water bath at 369 K for 4 h. The product was washed and filtered as before. Complete exchange to the sodium form was carried out by repeated treatment with sodium nitrate solution (1 mol dm-3) which had been made alkaline with sodium hydroxide (0.01 mol dm-3). Washing was completed with sodium hydroxide (0.01 mol dm-3), and finally with small volumes of water. Electron microscopy showed no impurities but the X.r.d. pattern showed small peaks (ca. 1 YO total intensity) due to zeolite A. This is designated sample K. Details of all the zeolites used in this work are shown in the table.Imbibition Experiment Procedure The apparatus differed from that of Fegan and Lowel in that simple glass specimen bottles were used to hold the zeolite samples, and these were capped with polythene stoppers after the desiccator was opened following equilibration. The desiccators were not rocked in the thermostat bath. A simple filter-paper cover proved effective at preventing condensation falling from the desiccator lid into the bottles. The solutionB. M. Lowe and C. G. Pope 947 Table 1. Synthesis and properties of zeolites composition of synthesis mixtures (mol) Si/Al in sample Na,O K,O A1,0, SiO, H,O producta size and morphology F1 75 0 15 60 4500 1.58 irregular shapes, probably agglomerates ; typical overall size ca. 0.5 pm F5 27 0 3 60 900 2.50 as F1 but larger at ca.2pm F5L 27 0 3 60 3600 2.54 spheroidal, platy surfaces typical size ca. 7 pm K 159 53 28.6 60 3550 0.98 as F5L size ca. 6pm a Estimated from the cubic unit cell parameter. D. W. Breck and E. M. Flanigen, Molecular Sieves (Society of Chemical Industry, London, 1968), p. 47. controlling the water activity was placed as close to the samples as practicable to minimise temperature differences, and wide diffusion paths were provided to further speed vapour transport. When experiments were performed above 298 K the use of an unsaturated control solution was preferred, as otherwise the change in solubility with temperature delayed the isopiestic equilibration. One of the sample bottles contained a weighed amount of control solution.Monitoring the weight of this provided an extra check on the composition of the external solution at equilibrium, Results and Discussion The use of glass bottles and the crude capping technique did not introduce errors in the weights of the equilibrated samples larger than the balance precision of f0.2 mg. Equilibration was not speeded up by allowing the control solution to evaporate from a larger surface area, but decreasing the length of the vapour diffusion paths, and avoiding constrictions were important. Equilibrium was typically established in 1-4 days, depending on the control solution composition. Reproducibility of results was similar at 298.2, 313.2 and 328.2 K, and condensation of water onto the sample containers when the desiccator was removed from the thermostat bath at higher than ambient temperature was not important.The theory of the isopiestic method, as presented by Fegan2 and Fegan and Lowe' predicts that at high salt contents the zeolite particles will be surrounded by a solution film in which the water activity is the same as that of the control solution. If the salt content is reduced, whilst the water activity is held constant, the liquid film becomes thinner but remains at the same concentration, as does the imbibed solution. Thus, the slope of this part of the isopiestic plot is constant as it is determined by the water activity, and can be predicted from literature data on the salt solution used. The observed and predicted slopes have always agreed within experimental error in both this and previously published work.', 2 , 6 The slopes of isopiestic plots at low salt concentrations, where no external film of solution is present, depend on the affinity of the zeolite for the salt and the water.Usually this part of the isopiestic curve is observed to be linear, or nearly SO,^^^^'*' and of different slope to the high salt side, although there is no thermodynamic requirement that this should be so.948 Salt Imbibition in Zeolites 0 0-04 0.08 mass ratio salt/zeolite Fig. 1. Isopiestic curves for NaNO, at 298.2 K on small crystal sample F5. The numbers on the curves are the molalities of the external solutions of NaNO,. The arrows mark the salt content at which the large crystal sample showed sharp intersection points for the same external solutions. At the intersection point the liquid film has just disappeared but the imbibed material is still at equilibrium with the external solution.Determination of the intersection point, which defines the composition of the imbibed solution is thus the basis of the isopiestic measurements of salt uptake which have been reported. Fig. 1 shows a set of isopiestic curves measured at 298.2 K in which sodium nitrate was imbibed by sample F5. The distinct curvature of the lines at intermediate salt concentrations which was also observed with F1 and other small crystal samples, made it uncertain as to how the concentration of the imbibed solution in vapour contact with the external solution should be determined. The curves did represent true equilibria, as judged by the following criteria.(i) The experimental points were unchanged after extended contact times. (ii) The slopes of the curves at high salt content were consistent with the compositions of the control solutions. (iii) Addition of water to the zeolite samples, followed by further equilibration, did not alter the final steady weights achieved. A further check on the reversibility of the data was performed by successively changing the control solution from saturated KCl to saturated KNO, and back again. Only two, reproducible isopiestic curves were obtained, irrespective of the order in which the measurements were made. Fig. 2 shows a set of results obtained with the larger crystal sample, FSL, which had essentially the same chemical composition as F5. Little curvature was observed. When data points very close to the intersection were obtained, they sometimes lay slightly above the curves, but the deviation was much smaller than with the small crystal samples, and only slightly greater than the probable experimental error.Results obtained at the same water activity with the two samples are shown for comparison in fig. 3. The intersection which is sharp, or nearly so, occurs at the same salt loading at which the small crystal data begin to display curvature. This was observedB. M. Lowe and C. G. Pope 949 0.40 0 Y .- .-.I 8 \ 92 3 'a 1 E 0 2 0.35 0.30 b 2.14 0.84 0.08 mass ratio salt/zeolite Fig. 2. Isopiestic curves for NaNO, at 298.2 K on large crystal sample F5L. The numbers on the curves are the molalities of the external solutions of NaNO,. It is possible that a small amount of curvature exists at the intercepts, and that this has led to the slope at low salt content appearing less steep with the two most dilute external solutions.0 at each water activity studied. The corresponding concentrations are marked by arrows on the curves of fig. 1. The different behaviour of the two samples seems most likely to arise from the difference in external surface area of the crystals, or from capillary-condensation effects. However, the latter explanation is inconsistent with the results, as the two zeolite samples have similar water contents and because the phenomenon occurs essentially unchanged at high and low water activities. It is possible to explain the observations if salt adsorption occurs more strongly on the external surface of the crystals than within their zeolitic pore space. If the F5 particles consist of agglomerates of material, size ca.0.1 pm, which is consistent with their appearance under the electron microscope, an adsorption of ca. one NaNO, unit per 0.3 nm2 area of external surface would be required to produce the observed effect. The external surface of crystals larger than ca. 2 pm would, on this basis, produce an almost negligible effect on the isopiestic curves. Whilst these magnitudes seem to be possible, it is surprising that adsorption should be so great, bearing in mind the low concentrations of imbibed salts which have been observed in this work, and by others.1,2i6,8,9 We believe that the local salt concentrations which occur in different parts of porous crystals, and on their external surfaces depend on the net result of ion/water and ion/ zeolite surface interactions, and on the influence of the Donnan exclusion effect.In large950 Salt Imbibition in Zeolites I 1 I 0.04 0.08 mass ratio salt/zeolite Fig. 3. Comparison of large and small crystal imbibition of NaNO, at 298.2 K. The numbers by the curves are the molalities of the external solutions of NaNO,. 0, F5L (large crystals); 0, F5 (small crystals). 0 cavities, where salt ions can only approach one part of the zeolite framework closely, the Donnan effect, and the more effective hydration which can be achieved in the external solution generally results in low imbibition. Adsorption on the outside surface of the crystals should be stronger, as Donnan exclusion does not apply there.Stronger adsorption may also occur in small cavities, where stronger ion framework interaction is possible. This probably accounts for the large concentrations of salts which become locked into sodalite and cancrinite during synthesi~,~! lo and may partly explain the profound effect that salts may exert on synthesis pr0ducts.l' We have observed that perchlorate is trapped in faujasite if it is included in the synthesis mixture, though only to the extent of about one ion per two sodalite cages in our experiments.12 Large post-synthesis imbibition of salts is not expected in aqueous systems, because ions which are small enough to enter the windows of small cages have large solvation energies which makes them favour a bulk solution environment. Our results suggest that salt imbibition experiments, whether carried out by the isopiestic or classical analysis procedures, may well give misleading information if external surface adsorption is significant.The isopiestic technique is to be preferred, because the shape of the characteristic curves shows immediately whether this is the case. Fig. 4 shows the imbibition of NaNO, at 298.2 K on F5L over an extended range of water activity. No abrupt change in imbibition behaviour is shown by this system at a water activity corresponding to a saturated external solution of the salt. Decreasing the water activity down to very low values only causes extra salt to enter the zeolite to a small extent. The isopiestic curves are all linear, with a constant slope of - 0.46 0.02B. M .Lowe and C. G . Pope 95 1 0.5 0.9 water activity 0.1 Fig. 4. Imbibition of NaNO, in F5L at 298.2 K over an extended range of water activity, Water activity in a saturated solution of NaNO, arrowed. 4 0 4 6 8 molality of external NaNQ solution 0 Fig. 5. NaNO, imbibition at 298.2 K (O), 313.2 K (0) and 328.2 K (A) on F5L and K samples. Ordinate scales displaced for clarity.952 Salt Imbibition in Zeolites 0.7 0.5 Y* 0.3 0.1 0 \ \\ I I I 4 8 12 ionic strength Fig. 6. yf values at 298.2 K. External solution values are calculated from the equation of Br0m1ey.l~ (a) external chlorate; (b) external nitrate. NaNO, in F5L (0) and in K (a). NaClO, in F5L (0). at the low salt side of their intersection points.This slope is nearly equal to the density ratio, water/solid NaNO,, (0.44) and suggests that the pore space is always full, so that salt uptake can only occur if an equivalent volume of water is expelled. The regular and simple nature of the results appears to confirm further the accuracy of the data, and that complications due to adsorption or capillary condensation are absent. Fig. 5 compares the imbibition by a low and a high aluminium content sample over a range of temperatures. There is little temperature dependence and the F5L results display the form typical of the operation of a Donnan equilibrium, as has been reported b e f ~ r e . ~ Surprisingly, the high aluminium content sample imbibed similar amounts of NaNO,, though previous work had suggested2,9 that higher salt uptake might be expected.The graphs of internal us. external molality were almost straight lines, and displayed little of the curvature which has been attributed to the operation of a Donnan effect. ** Donnan theory leads to the equation mzyf: = mi(mi+m,)(y~i)z (1) for imbibition of a 1 : 1 electrolyte. The terms on the left-hand side relate to the external solution, and those on the right-hand side to the imbibed solution. m, is the molality of the charge balancing cations. Even though plots of mi(mi + m,) us. mf have been found to be straight lines with some of the systems studied in the past, this does not constitute a sensitive or reliable test of the Donnan equation. There appears to be no reason toB. M. Lowe and C. G. Pope 953 6 0 4 0 - c e W 2 0 0 0 3 0 6 0 9 0 Fig.7. Conventional tests of the Donnan equation, NaNO, in F5L at 298.2 K, calculated with different assumed internal cation molalities. 0, Calculated from correct zeolite composition. , Calculated with pure zeolite cation concentration halved. expect that the activity coefficient ratio should be constant. Indeed, the large, varied, and different ionic concentrations outside and inside the zeolite, together with the different environments in the two locations suggest that we might expect to see the ratio vary a great deal. As yf, values are available from independent sources13 we believe that attention should be directed to the terms relating to the imbibed solution. If the zeolite contains imbibed salt, the cations from this material are not distinguishable from the original charge-balancing cations of the same chemical type, and should be included in the calculation in the same way.m, is therefore properly calculated from the chemical composition of the original zeolite, the amount of imbibed salt, and the water content at imbibition. With the values of mi, m,, me and y f , it is possible to obtain y+ from eqn (1). These values are dependent on the environment within the zeolite pores and cannot be calculated from first principles, as the problem involves all the difficulties associated with the treatment of concentrated electrolyte solutions, together with the extra complication of the influence of the zeolite pore walls. are acceptable. (1) y should change systematically with salt imbibition. (2) y f should be of similar magnitude, but probably somewhat higher than expected for an external solution of the same ionic strength.(3) y should vary systematically with the charge on the zeolite framework for a given salt and zeolite. The basis for proposition (2) is that the zeolite cavities must restrict the ability of the solvent to respond to the demands of the internal electrolyte, which must destabilise the internal salt to some extent compared with the free solution environment. Fig. 6 shows We suggest three criteria which might be used to assess whether the values of y954 Salt Imbibition in Zeolites 0 2 4 6 8 10 external molality Fig. 8. Imbibition of sodium nitrate (O), sodium chlorate (a) and trisodium &rate (0) at 298.2 K.the variation of y k and y k with ionic strength of the salt solutions. Bulk solution activity coefficients are calculated using the equations of Bromley. l4 The data for the high aluminium sample follow the bulk solution values remarkably closely. The more siliceous zeolite also yields yki values which vary smoothly with concentration, but are somewhat higher than y+ at the same ionic strength. The results appear to be consistent with the predictions which can be made from Donnan theory. The lower values obtained with the sample K may reflect some stabilisation of the salt by the charge on the zeolite framework. Fig. 7 shows the conventional test of the Donnan equation. The agreement appears convincing, until the results of using quite incorrect charge balancing cation concentrations are also examined.This illustrates quite clearly that the test is insufficiently sensitive to be very useful. The relation between salt and water uptake into initially salt-free zeolites was examined in experiments with sodium nitrate, sodium chlorate, sodium citrate and sodium fluoride. Nitrate and chlorate gave very similar isopiestic curves, and their different structure directing effect in synthesis3 was not reflected in these results. Citrate was used because of its larger size, higher charge and greater water affinity. It was found that as this salt entered the zeolite, the water content remained almost unchanged over a wide range of water activity. The space required to accommodate the salt, and the influence of the ions on the internal water packing clearly produced compensating effects.The imbibition of the three salts is shown for comparison in fig. 8. The low uptake of citrate is consistent with hydration causing the salt to favour concentration in the bulk solution.B. M. Lowe and C. G. Pope 955 Experiments with sodium fluoride were unsuccessful, as the isopiestic curves showed no sharp changes of slope. Thus no estimate of salt occlusion could be made, even though this probably occurred to some extent. The failure of the method arose as follows. Salt entering the zeolite causes little change in water content, due to the strong solvation and rather small molar volume of the salt. The low solubility of sodium fluoride means that no external solution can form except at high water activity, so that the high salt content side of the isopiestic curves is also horizontal with the control solutions used.The sample showed the same X.r.d. pattern before and after fluoride treatment, so that lattice attack was not significant during these experiments. Experience in this work, and with other large crystal material’ suggests that at high water activities and low salt contents, the slopes of isopiestic curves on the low salt side of the intercept have constant slope, or a slope which changes very slowly with water activity except perhaps for very high water activities. Larger variations may well be an indication of external adsorption being important. At low water activity, behaviour appears to be more varied, and no generalisation can be offered at the present time. C.G.P. thanks the University of Otago for study leave, and the University of Edinburgh, Chemistry Department, for its hospitality during 1987. References 1 S . G. Fegan and B. M. Lowe in Proc. 6th Int. Zeolite Conf., Reno, ed. A. Bisio and D. Olson (Butterworths, London, 1983), p. 288. 2 S . G. Fegan, Ph.D. Thesis (University of Edinburgh, 1985). 3 R. M. Barrer and J. F. Cole, J. Chem. SOC., A., 1970, 1516. 4 H. Kacirek and H. Lechert, J. Phys. Chem., 1975, 79, 1589. 5 G. H. Kuhl, Zeolites, 1987, 7 , 451. 6 K. R. Franklin, B. M. Lowe and G. H. Walters, J. Chem. Res. ( S ) , 1988, 32. 7 B. M. Lowe, C. G. Pope and C. D. Williams, J. Chem., SOC., Chem. Commun., 1988, 1186. 8 R. M. Barrer and W. M. Meier, Trans. Faraday SOC., 1958, 54, 1074. 9 R. M. Barrer and A. J. Walker, Trans. Faraday Soc., 1968, 60, 171. 10 R. M. Barrer, J. F. Cole and H. Villiger, J. Chem. SOC. A . , 1970, 1523. 11 R. M. Barrer, Hydrothermal Chemistry of Zeolites (Academic Press, London, 1982). 12 B. M. Lowe and C. G. Pope, unpublished work. 13 J. F. Zemaitis, Jr, D. M. Clark, M. Rafal, N. C. Scrivner, Handbook of Aqueous Electrolyte 14 L. A. Bromley, AZChE J., 1973, 19, 313. Thermodynamics (D.I.P.P.R., American Institute of Chemical Engineers, New York, 1986). Paper 8/02148A; Received 13th September, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500945

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. SOC., Furaday Trans. I , 1989, 85(4), 957-967 Study of the Conformational Equilibrium between Rotational Isomers using Ultrasonic Relaxation and Raman Spectroscopy Part 3.-1 -Bromo-2-chloroethane Shinobu Koda, Hirohito Matsui and Hiroyasu Nomura" Department of Chemical Engineering, School of Engineering, Nagoya University, Chikusa-ku, Nagoya-shi 464, Japan The rotational isomerism of 1 -bromo-2-chloroethane has been studied by Raman and ultrasonic spectroscopy. Ultrasonic relaxation spectra have been measured in the wide frequency range from 10 MHz to GHz, using the pulse, high-resolution Bragg reflection and Brillouin scattering methods. The enthalpy and volume differences between trans and gauche forms have been obtained from the temperature and pressure dependences of Raman bands, respectively.Combining the ultrasonic and Raman data, the entropy difference between two conformers has been estimated. The enthalpy difference of 1-bromo-2-chloroethane is almost the same as that of 1,2- dichloroethane, but the volume difference and entropy difference are small, in comparison with those of 172-dichloroethane and 1,2-dibromoethane. These results are interpreted in terms of the differences in the intermolecular interaction and packing state in the dihalogenoalkane molecules. In our series of the rotational isomerism of several halogenoalkanes in liquid and dissolved states was investigated by means of the ultrasonic relaxation and Raman spectroscopy. In the previous paper,4 we pointed out that a linear relationship between the ultrasonic relaxation time for the rotational isomerism and the molecular weight holds for most part of chloro- and bromo-homologues, but the relaxation times of 1,2-dichroloethane and 1,2-dibrornethane were abnormally short and deviated largely from the values expected from their respective homologues.However, dihalogenoethanes are still typical samples for conformational equilibrium between trans and gauche forms in a molecule. In order to elucidate more clearly the rotational isomerism of the dihalogenoethanes, we studied the conformational equilibrium of 1 -bromo-2-chloroethane in liquid state by means of ultrasonic relaxation and Raman spectroscopy. Earlier works on dihal~genoethanes~. suggested that the ultrasonic relaxation due to the rotational isomerism of 1-bromo-2-chloroethane was to be in the frequency range above a few hundred MHz.For measurements of the ultrasonic velocities and absorptions in the frequency ranges from one hundred to several hundred MHz, the high-resolution Bragg reflection method proposed and developed by Takagi and co-workers5~ is particularly useful. In this work, an apparatus based on the high-resolution Bragg reflection method was constructed and used. Since the thermodynamic and kinetic parameters related to the rotational isomerism of 1 -bromo-2-chloroethane have been determined completely, these values will be discussed in comparing with those of 1,2-di~hloroethane~ and 1,2-dibr0moethane.~958 Rotational Isomers of 1 -Bromo-2-chloroethane 9 10 11 ,2 n 7 3 1 analyser - e3 Lock- in A+, recorder I I Fig.1. Block diagram of HRB method. 1, He-Ne laser; 2, light chopper; 3, 28 MHz oscillator; 4, diffraction cell; 5, rotary table; 6, sample cell; 7, differential transformer; 8, signal generator; 9, half-mirror ; 10, lens ; 1 1, photodiode ; 12, pre-amplifier ; 13, spectrum analyser ; 14, lock-in amplifier ; 15, X- Y recorder. Experiment a1 Sample The sample, 1 -bromo-2-chloroethane, was purchased from Tokyo Kasei Kogyou Co. Ltd, and was fractionally distilled before use. Raman Spectroscopy The characteristic bands of the C-X (X = halogen) sensitive stretching mode of I-bromo-2-chloroethane are as follows; for trans conformer, 724 and 631 cm-l for C-CI and C-Br stretching modes, respectively, and for the gauche form, 664 and 568 cm-l for C-CI and C-Br stretching modes, respectively.The Raman spectra under high pressure and the temperature dependence of Raman spectra were recorded using the same laser Raman spectrometer, apparatus, and attachments as used in our previous The operating pressure was up to 2500 x Pa. Detailed experimental procedure and data analysis are described el~ewhere,~ The accuracy of the integrated intensities was within 3 O/O. Ultrasonic Measurements ( I ) Below 100 MHz, the ultrasonic absorption coefficients were measured by the usual pulse method. Several different types of measurement cells were used depending on the measuring frequencies. The error in the ultrasonic absorption coefficient was within ca. 3 % on the a/f x lo1’ Np s2 cm-l scale. (2) In order to measure accurately the ultrasonic absorption coefficient and velocity in the frequency range higher than 100 MHz, we constructed an apparatus based on the high-resolution Bragg reflection (HRB) method, in which the optical heterodyne system was set up.* The experimental arrangement is schematically shown in fig.1. The laser beam (He-Ne Laser of NEC, GLG5800) is expanded to 6 mm# in diameter. The parallel light incident on the sample cell is scattered by the continuous sound wave of frequency,S. Koda, H . Matsui and H. Nomura 959 c 28 MHz 1 I I I I I I 1 I Cb 19634' 33' 32' 31' 193' 59' 58' 57' 56' 55' Fig. 2. Typical beat spectra obtained for water at 351 MHz at 25 "C. The beat signals at 28 and 379 MHz correspond to the incident and scattered radiation, respectively.f, under the condition of Bragg reflection. The sound wave comes from a quartz buffer rod, to which a ZnO film transducer with a fundamental frequency of 500 MHz is deposited. Part of the parallel light is modulated by an ultrasonic light modulator at a frequency of 28 MHz and is used as a reference light. In practice, the light reflected under Bragg's condition is superimposed with the reference light on a half-mirror. The intensity of the beat signal is detected with a pinphoto-diode (HP4220) and a Spectrum- Analyzer (Takeda Riken TR4113AL). In order to obtain the intensity spectrum of the beat signal around the Bragg angle, the direction of sound beam is changed by rotating the rotary table. A typical recorder chart obtained for water at f = 35 1 MHz is shown in fig.2. This shows the Lorentzian distribution with an angular width proportional to the ultrasonic absorption (a =q cos @,A@, where q is the wavenumber and a,, the Bragg reflection angle). After the beat signal at f+28 MHz is recorded, the table is rotated until a sharp beat note at 28 MHz is detected. The difference between the two peak points gives directly half of the scattering angle, 0,. The sound velocity can be determined as U = nf/(q sin 0,). Velocities of water determined by HRB method in the frequency range 26MOO MHz are listed in table 1. The temperature of water was controlled within kO.1 "C. The accuracy of velocity measurement was better than kO.05 O/O. Also given in table 1, are the sound velocities that Greenspan and Tschiegg measured at 3 MHz' and the ultrasonic absorptions measured by Mukhopadhyay. lo Agreement is satisfactory.(3) The sound velocity in the GHz region was measured by the Brillouin scattering method. The details of the apparatus and experimental procedures are given elsewhere. l1 The accuracy in the sound velocity was within 1 %. Results and Data Analysis Raman Spectroscopy Fig. 3 shows the typical Raman spectra of C-X sensitive stretching vibrational mode of the sample, together with those of 1,2-dichloroethane and 1,2-dibromoethane (hereafter abbreviated as 1,2-DCE and 1 ,2-DBE, respectively). For assignment, the960 Rotational Isomers of 1 -Bromo-2-chloroethane Table 1. Ultrasonic absorption coefficient and sound velocity of water obtained from the high- resolution Bragg Reflection technique (at 25 "C) f/MHz U/m s-' lO"a/fZNp s2 cm-' 1497.00" 22.2b 26 1 1498.57 21.8 351 1497.45 22.5 457 1497.77 22.9 56 1 1496.09 22.5 " Ref.(9). Ref. (10). I 1 I 1 1 I I 550 600 650 700 750 800 v/m- dibromoethane in the liquid state and 1-bromo-2-chloroethane in the solid state. Fig. 3. Raman spectra of (a) l-bromo-2-chloroethane, (b) 1,2-dichloroethane and (c) 1,2- Raman spectrum of 1-bromo-2-chloroethane (1,2-BCE) in solid state is also shown in fig. 3. The assignment is complete and the frequencies of their characteristic C-X sensitive stretching vibrational modes of each conformer are given in fig. 3. The vibrational frequencies used were independent of temperature and pressure ranges concerned in this work. Assuming that the ratio i2,/SZ, (where SZ is the absolute scattering cross-section of Raman scattering of each mode) is independent of pressure and temperature in the whole range of measurements, the volume and enthalpy differences between the twoS.Koda, H. Matsui and H. Nomura -1.7- -1.8 -1.9- 96 1 0 - a a - - - a - - - I I I I I conformers, that is, trans and gauche forms, can be estimated from the Arrhenius and van’t Hoff equations : where I refers to the Raman intensity. In a previous paper,7 it is confirmed that the depolarization ratio of the C-X stretching mode for dihalogenoethanes little changes with increasing pressure up to 2500 x Pa. Therefore, the above assumption is reasonable for ethane derivatives in the pressure range investigated here, as a first approximation. Fig. 4 and 5 show the relationships between In (ZJI,) us.1 / T and P for the characteristic band of C-Cl and C-Br, respectively. The absolute values of ln(Ig/It) are quite different for C-Cl and C-Br bands but the volume and enthalpy differences, A V and AH, obtained from respective stretching vibrational modes are in excellent agreement with each other. The values thus obtained are summarized in table 2, together with the enthalpy difference in gaseous state found in the literature.12 Ultrasonic Spectroscopy Ultrasonic absorption and sound dispersion spectra of 1,2-BCE are shown in fig. 6 and 7 as a function of temperature. The ultrasonic relaxation is observed in the whole temperature range investigated. The relaxation spectra can be well expressed by the following single relaxation equation wherefrepresents the measuring frequency, f, the relaxation frequency, A the relaxation amplitude and B the absorption from processes other than relaxation.The solid curves in fig. 6 represent the single relaxation curves which are calculated from the linear least mean-squares method minimizing the total relative deviation from the experimental values. The fitting errors are within 3 % in this work. The relaxation parameters thus obtained are summarized in table 3.962 Rotational Isomers of 1 -Bromo-2-chloroethane 1 I I I I I 1 1 400 800 1200 1600 2000 24( p m 5 Pa Fig. 5. Relationships between the logarithmic ratio of integrated intensities and pressure, (0); C-Br stretching mode and (a) C-Cl stretching mode. Table 2. Enthalpy and volume differences obtained from the analysis of Raman intensities AH/kJ mol-' AV/cm3 mol-' AH,,"/kJ mol-1 1,2-DCE 0 - 2.7 6.2 1,2-DBE 3.6 - 5.2 8.4 1,2-BCE 0.4 - 2.0 7.7 "Ref.(12). The sound dispersion curves can also be expressed by the following single dispersion equation :13 where Urn and U, are the sound velocities at the high- and low-frequency limits, respectively, and f" is the dispersion frequency. If eqn (4) can well reproduce the experimental data, the plot between q/(v"- q) and l/f" should be linear. For example, the plot of the data at 25 "C is shown in fig. 8. The straight line was obtained and the dispersion parameters, U,, Uo, f" and the relaxation strength, E (= u",/(u" - E)), were estimated from the slope and intersection of the line. The data thus obtained are summarized in table 4.It is well known that the sound dispersion and sound absorption can be expressed as the real and imaginary parts of the propagating sound wave in a relaxing medium. If the relaxation and sound dispersion curves can be expressed by the single relaxation andS. Koda, H. Matsui and H . Nomura 963 300 200 400F ;OO--.--., loot \ 10 100 1000 fIMHz Fig. 6. Ultrasonic absorption as a function of frequency. Temperatures in "C. 12201 10601 , , I I I , , I I 1 1 , , , , , , , ,,Tf 5 100 1000 10000 f I M H Z Fig. 7. Dispersion of sound velocity. Temperatures in "C.964 Rotational Isomers of 1 -Brorno-2-chloroethane Table 3. Relaxation parameters estimated from the absorption data T/ "C f,/MHz 101'A/Np s2 cm-' 101'B/Np s2 cm-' r - 5 350f10 310& 10 30-60 0.042 5 420f20 260+ 10 30-60 0.041 15 680+30 230+ 10 3&50 0.056 25 780f30 220 + 10 30-50 0.060 10-17 s-2if2 Fig.8. Plot of u",/(UZ-- u",) us. l/f at 15 "C. Table 4. Relaxation parameters estimated from the sound velocity dispersion data T/ "C f"/MHz U,,/ms-l U,/ms-l E -5 350 1160 1202 0.068 5 530 1130 1166 0.061 15 850 1099 1133 0.059 25 950 1067 1098 0.056 dispersion equations, respectively, i.e. eqn (3) and (4), then the following Cole-Cole type equation should hold between the a and V at each measuring frequency.13 where a'3, is the absorption per wavelength. As a typical example, the above ColeCole plot is shown in fig. 9 for 1,2-BCE at 15 "C. In fig. 9, the data are on the half circle within experimental errors in the whole temperature ranges.This indicates that our data analysis is self-consistent. Taking account of the discussion for the ultrasonic data of 1,2-DCE5 and 1,2- DBE,' the ultrasonic relaxation and sound dispersion are ascribed to the rotational isomerism of the 1,2-BCE molecule in liquid state. As pointed out previou~ly,~~~ the free-energy difference AG between the twoS. Koda, H. Matsui and H. Nomura 965 Fig. 9. Plot of -- us. -. : conformers can be evaluated according to the following equation for the rotational isomerism : r = Y'(. RT C, where r is the relaxation strength and values A V and AH are obtained from the Raman spectroscopy. The relaxation strength, r, can be calculated as r = AfrU/2. In eqn (6), g, and g, are, respectively, the degeneracies of the lower and upper energy levels of the two conformers, 8 is the expansion coefficient, C, is the heat capacity at constant pressure and the other symbols have their usual meanings.If the rate constant of the backward reaction is large in comparison with that for the forward reaction, the relaxation frequency, f,, can be described in terms of the following Eyring-type equation :14 f, = V e x p 2nh (- AG:/RT) (7) where AG: is the activation energy and AG: = AH- TAS. Fig. 10 shows the plot of lnf, against 1/T which provides the activation enthalpy, AH and also activation entropy, 9, from the slope and intersection, respectively. The thermodynamic parameters thus estimated are summarized in table 4 together with those of 1 ,2-DCE5 and 1 ,2-DBE6 obtained previously. Discussion AHliquid between two conformers consists of the two parts, that is, intra- and inter- molecular interactions in liquid state.As the AH,,, can be ascribed mainly to be intramolecular interaction, A ( = AH,,, - AHliquid) is a good measure of the inter- molecular interaction in the liquid state. As shown in table 2, the A of 1,2-BCE has966 Rotational Isomers of 1 - Bromo-2-chloroethane 3.8 4.0 4.2 4.4 104/RT Fig. 10. Plot of In (2nf,/T) us. l / R T . Table 5. Summary of the thermodynamic parameters (25 "C) AG/kJ mol-' AS/J K-' mol-1 AH'/kJ mol-1 A$/J K-l mol-1 AG:/kJ mol-1 1,2-DCE 4 - 12 8 - 27 16 1,2-DBE 7 - 10 12 - 17 17 1,2-BCE 2 - 4 16 -4 17 large values similarly as for the cases of 1,2-DCE, and 1,2-DBE, indicating that 1,2-BCE molecules are in fairly strong molecular interaction.However, as shown in tables 2 and 5, the volume difference, AV, and also entropy difference, AS, of 1,2-BCE are small compared with those of 1,2-DCE, and 1,2-DBE. This means that in the case of 1,2-BCE, the local configuration is similar for trans and gauche forms and the differences in packing states are rather small in comparing with those of the 1,2-DCE, and 1,2-DBE. On the other hand, the activation free energy, AGI? of 1,2-BCE is almost equal to those of 1,2-DCE, and 1,2-DBE. In other words, the rate constants of backward reaction of conformational isomerism are in the same order of magnitude for 1,2-DCE, 1,2-DBE and 1,2-BCE. In our previous paper,4 the ultrasonic relaxation frequencies of dihalogenoethanes deviated from those expected from other halogenoalkanes.The relaxation frequencies of 1,2-BCE observed locate in the lower frequency side than that of 1 ,2-DCE and are nearly equal to that of 1,2-DBE at the same temperature, but they still deviate from those of halogenoalkane homologues. The large values of the relaxation frequency, f,, mean that the process of the rotational isomerism for dihalogenoethanes is much faster than those of other halogenoalkanes. If the molecular collision process dominates the rotational isomerism, the short relaxation times of dihalogenoethanes for rotational isomerism suggest that the energy transfer from translation to rotation induced by molecular collision occurs more easily for these dihalogenoalkanes than for other homologues. Combination of the Raman and ultrasonic relaxation spectroscopic data permitsS. Koda, H .Matsui and H. Nomura 967 experimental determination of the free-energy difference between two conformers. The free energy difference, AG, is also related to the equilibrium constant, K , of the conformational equilibrium which can be expressed as : K = exp (- AG/RT) = (n,/n,) = (IgsZt/ZtsZ,). Therefore, we can estimate the ratio of the absolute Raman scattering cross-section between tram and gauche conformers. From the data of AG and Ig/It of 1,2-DCE, we estimated the values of Cl,/n, for C-Cl stretching vibrational mode as 0.21,, while n,/n, for C-Br can be estimated as 0.33, from the data of 1,2-DBE. Using the data for 1,2-BCE, the values of Qt/Q, for C-Cl and C-Br stretching vibrational mode can be obtained simultaneously as 2.87, and 0.48,, respectively. Although the nJQ, values for C-Br obtained from 1,2-DBE and 1,2-BCE are nearly equal, those for C-Cl are quite different.These situations can easily be understood from the Raman spectra of 1,l-DCE, 1,2-DBE and 1,2-BCE, as shown in fig. 3. Unfortunately, no data are available on the absolute Raman scattering cross-section of 1,2-DCE, 1,2-DBE and 1,2-BCE investigated here. Therefore, discussion on this matter is impossible at this stage. However, it is worthwhile pointing out that this method is very useful for estimating the absolute Raman scattering cross-section. We express our deep gratitude to Prof. K. Takagi of Tokyo University for his helpful advice in constructing our HRB apparatus. This work was supported in part by a Grant- in-Aid for Scientific Research from the Ministry of Education, Science and Culture (No. 6 1 134043). References 1 H. Nomura, Y. Udagawa and K. Murasawa, J. Mol. Struct., 1985, 126, 229. 2 H. Nomura and S . Koda, Bull. Chem. SOC. Jpn, 1985, 58, 2917. 3 H. Nomura, S. Koda and K. Hamada, J. Chem. SOC., Faraday Trans. 1, 1987, 83, 527. 4 H. Nomura, S. Koda and K. Hamada, J. Chem Soc., Faraday Trans. I, 1988, 84, in press. 5 W. Seki, P.-K. Choi and K. Takagi, Chem. Phys. Lett., 1983, 98, 518. 6 K. Takagi, P.-K. Choi and W. Seki, J . Chem. Phys., 1983, 79, 964. 7 H. Nomura, K. Murasawa, N. Ito, F. Iida and Y. Udagawa, Bull. Chem. Soc. Jpn, 1984, 57, 3321. 8 K. Takagi and K. Negishi, J. Appl. Phys., 1975, 14, 29. 9 M. Greespan and C. E. Tschiegg, J . Acoustic Soc. Am., 1959, 31, 75. 10 S. K. Mukhopadhyay, Acoustica, 1956, 6, 25. 11 S. Koda, H. Nomura, M. Nakamura and Y. Miyahara, Bull. Chem. SOC. Jpn, 1985, 58, 1484. 12 J. Powling and H. J. Bernstein, J. Am. Chem. SOC., 195 1, 73, 18 15. 13 K. F. Hertzfeld and T. A. Litovitz, Absorption and Dispersion of Ultrasonic Waves (Academic Press, 14 A. J. Matheson, Molecular Acoustics (John Wiley, Chichester, 1971). New York, 1959). Paper 8/02225T; Receiued 3rd June, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500957

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. SOC., Furudajj Trans. 1, 1989, 85(4), 969-976 Electrochemical Behaviour of Polyaniline in Weak Acid Solutions Takayuki Hirai, Susumu Kuwabata and Hiroshi Yoneyama" Department of Applied Chemistry, Faculty of Engineering, Osaka University, Yamada-oka 2-1, Suita, Osaka 565, Japan The electrochemical behaviour of polyaniline (PAn) in weak acid solutions such as pH 4 has been studied. It has been found that PAn gives a stable redox reaction involving electrolyte anions in a limited potential region, whereas it loses its electrochemical activity by being oxidized at more anodic potentials where deprotonation preferentially takes place. Detailed mecha- nisms of the redox and the inactivation reactions of PAn in weak acid solutions are discussed. During the past several years, polyaniline (PAn) films prepared by electrolytic polymerization of aniline have been intensively investigated with regard to their electrochemical properties.Polyaniline exhibits a reversible redox wave due to protonation and deprotonation of PAn in a limited potential region in highly acidic and several applications such as electrochromic devices6-' and electronic devicesg*'O have been attempted using this redox reaction. It has also been found that a redox reaction occurs in non-aqueous solutions with doping and undoping of electrolyte anions if PAn is prepared in highly acidic solutions, and investigations on the utility of this property as an active mass of secondary batteries have been conducted.''~12 However, electrochemical reactions of PAn in relatively weak acid or neutral aqueous electrolyte solutions have not yet been clarified. For example, it is known that PAn loses its electroactivity with increasing pH value of electrolyte solutions, but the detailed reason is not known." Concerning electrochemical reactions in weak acid solutions, Huang et al.suggested from their voltammetric studies that electrolyte anions instead of protons are involved in the redox reaction at pH lH.l3 In this paper, it is shown that a stable redox reaction of PAn occurs even in weak acid solutions such as pH 4 in a limited potential region and that species which are involved in the redox process gradually change from protons to electrolyte anions with an increase in pH. Furthermore, mechanisms on the inactivation process of PAn that occurs in weak acid solutions are proposed based on the finding that the inactivation takes place at potentials a little positive of the anodic limit that gives the stable redox reaction.Experimental PAn films were prepared on Au plates (1 cm2) for electrochemical measurements and on indium-tin-oxide coated glasses (ITO) (2 cm2) for measurements of visible absorption spectra of the films. The films were prepared by anodic polymerization of 1 mol dmP3 aniline dissolved in 2 mol dm-3 HC1 solution at 0.3 mA cm-2 under an N, atmosphere. The quantity of electricity chosen for the film preparation was 45, 90 and 180 mC for measurements of visible absorption spectra, electrochemical measurements, and electron probe X-ray microanalyses, respectively. Electrochemical measurements of the prepared PAn films were carried out using a potentiostat (Hokuto Denko HA- lOl), a function-generator (Hokuto Denko HB- 104) 969970 Electrochemical Behaviour of Polyaniline and an X-Y recorder (Riken Denshi F-3C).A saturated calomel electrode (SCE) served as a reference electrode. Electrolyte anions incorporated into the PAn films were analysed with an energy dispersive electron probe X-ray microanalyser (EPMA) (Horiba EM AX- 1 500E) connected to a scanning electron microscope (SEM) (HITACHI S-450). Visible absorption spectra of the PAn film were measured using a Pyrex cell [20 mm x 10 mm x 30 (height) mm] equipped with an IT0 electrode deposited with PAn film, a Pt counter electrode and an SCE. This cell was set in a u.v.-visible spectrophotometer (Shimadzu MPS-5000), and absorption spectra of PAn films were measured in situ under steady-state polarization at various electrode potentials.Results and Discussion Redox Behaviour of PAn in Weak Acid Solutions Cyclic voltammograms of PAn taken in chloride and benzenesulphonate solutions at pH = 0, 2 and 4 are shown in fig. 1 and 2 , respectively. The scanning range of the potentials were selected in such a manner as to avoid H, evolution at the cathodic limit and to avoid at the anodic limit decomposition or inactivation of PAn that will be described below. All the voltammograms taken under such conditions were quite stable with potential sweep cyclings. It is well established that the redox wave in pH 0 is due to protonationvdeprotonation reaction, but the contribution of this reaction becomes less with increasing P H .~ Both chloride and benzenesulphonate solutions used in the present study did not have buffering actions. It is then highly probable that the pH within and at the surface of the PAn films was varied by potential sweepings. If this were true, interpretations of the changes in voltammograms with pH and/or with potential sweep cyclings should be made carefully. Even so, however, it is recognized that there are marked differences in pH dependences of voltammetric behaviour between the chloride and the benzenesulphonate solutions. Such different behaviour of cyclic voltammograms with solution pH between the chloride and the benzenesulphonate solutions can be reasonably interpreted by assuming that the redox mechanism is changed gradually with decreasing the proton concentration.In the case of NaC1, C1- anions become incorporated with increasing pH in place of deprotonation as will be evidenced below, but benzenesulphonate anions do not owing to the anion-sieving effect of PAn films,14 resulting in a decrease of voltammetric waves. The incorporation of C1- anions into PAn was found by EPMA analyses. Results obtained are shown in fig. 3. Cl signals are seen at PAn film oxidized at 0.3 V us. SCE in NaCl solution of pH 4, but these signals disappeared when the film was reduced at -0.6 V us. SCE. The ratio of Cl to N determined by elemental analyses was 0.34 for the oxidized PAn, but less than 0.05 for the reduced one. These results strongly evidence that at pH 4 the doping and undoping of electrolyte anions take place in the course of the oxidation and reduction of the PAn film, respectively.As described above, the anion- sieving effect is seen for benzenesulphonate solution. Then the incorporation of benzenesulphonate into PAn seems unlikely. However, it seems important to remark here that if the electrode potential was held for about 5 min at 0.3 V us. SCE in benzenesulphonate solution of pH 4, the incorporation of benzenesulphonate anions into the film took place, as observed by EPMA analysis. Therefore, the lack of sharp waves in the cyclic voltammogram taken in the benzenesulphonate solution at pH 4 may be caused by slow rates of insertion of large benzenesulphonate anions into PAn. It is then concluded that the redox mechanism of the PAn film changes from a protonationvdeprotonation mode in the high acidic solutions to a mode of doping- undoping of electrolyte anions in weak acid solutions.T.Hirai, S . Kuwabata and H. Yoneyama 97 1 I I I 1 I I I -0.6 0.0 0-6 EIVos. SCE Fig. 1. Cyclic voltammograms of PAn electrodes in chloride solutions having three different acid concentrations; in 1 rnol dm-3 hydrochloric acid (---), in 0.2 rnol dmP3 sodium chloride/H,O adjusted to pH 2 (. - .) and that adjusted to pH 4 (---). dE/dt = 10 mV s-l. 0.8 -0.6 0-0 0-6 EIVvs. SCE Fig. 2. Cyclic voltammograms of PAn electrodes in benzenesulphonate solutions having three different acid concentrations; in 1 mol dm-3 benzenesulphonic acid (---), in 0.2 mol dm-3 sodium benzenesulphonate/H,O adjusted to pH 2 ( .. . ) and that adjusted to pH 4 (-). dE/dt = 10 mV s-l. 3 3 F A R I972 Electrochemical Behav io ur of Po ly an iline I I I 0.25 0.3 X - ray energylkev Fig. 3. EPMA spectra of C1 for PAn films which were oxidized at 0.3 V (-) and reduced at -0.6 V (---) in 0.2 mol dm-3 NaCI/H,O (pH 4). 1.0 c -0.6 0.0 0-6 EIVvs. SCE Fig. 4. Cyclic voltammograms of PAn electrodes in 0.2 mol dm-3 NaCI/H,O (p 4) taken in a potential region between -0.6 and 0.6 V. The number in the figure denotes the cycle number. dE/dt = 10 mV s-l. Oxidation of PAn at More Anodic Potentials Fig. 4 shows cyclic voltammograms of PAn taken in the NaCl aqueous solution of pH 4 in a potential range between - 0.6 and 0.6 V us. SCE. By expanding the anodic limit in potential sweep cycling from 0.3 to 0.6 V, anodic and cathodic waves were decreased with successive potential sweep cycles, and a vivid colour change of the film observed in early stages in the potential cyclings was gradually weakened finally to give a blue film regardless of the electrode potential. The observation is in qualitative accordance with the results shown in several papers,1.2 that PAn films are electroinactive at pH > 3-4.This inactivation behaviour of the PAn film was then investigated spectrochemically.T. Hirai, S . Kuwabata and H. Yoneyama 973 * 1 I I I I I I (4 0.0' ' I I I I I I r J 0.0' ' I I I I I I I I I I I I 1 0-6 0.0' ' 1 I I I I I I 400 500 600 700 wavelength/nm Fig. 5. Visible absorption spectra of PAn films obtained in situ at 0.2(a), 0.4(b) and 0.6(c) V us.SCE in chloride solutions. The numbers in the figure denote the solution pH. The spectra of the PAn films polarized at 0.2, 0.4, and 0.6 V vs. SCE are shown in fig. 5(a), (b) and (c), respectively, for four kinds of chloride solutions of pH 0, 1.8, 3.0 and 4.0. Almost the same spectra were obtained for PAn films polarized at 0.2 V in the four kinds of solutions, for those at 0.4 V in solutions of pH 0, 1.8 and 3, and for those at 0.6 V at pH 0. The PAn films polarized under these conditions must have the same structure. The spectra obtained under other conditions are different. The spectra obtained at 0.4 V in pH 4 and 0.6 V in pH 3 and 1.8 look eventually the same, though that at 0.6 V in pH 4 is a little different. There are common features in these spectra, however, in comparison with the spectra shown in fig. 5(a); the absorption around 400 nm is relatively weak and a broad absorption which covers wavelengths from ca.450 to 600 nm is intense. Once when PAn films were oxidized at 0.4 V in solutions having pH greater than 4 or at potentials positive of 0.6 V in pHs greater than 3, they were hardly reduced in the same electrolyte solutions. The PAn films were electroinactive, and the spectrum was little changed by cathodic polarization of the oxidized film as a matter of course. The inactive film seemed highly stable, because the blue colour of the film was maintained even when the film was washed with water and/or air-dried. Judging from absorption spectra and voltammetric behaviour of PAn, the inactivation depends on the electrode potential and solution pH.However, since the correct pH value in the PAn film could not be known, the inactivation potential could not be formulated with the pH. Qualitatively, the potential at which the inactivation commences to occur is said to become negative with increasing pH. Indirect evidence to support this view was found in voltammetric behaviour of PAn in a NaCl solution of pH 6, shown in fig. 6. Potential 33-2914 Electrochemical Behaviour of Polyaniline 1 I -0-6 0.0 0.4 EIVvs. SCE Fig. 6. Cyclic voltammograms of PAn electrodes in 0.2 rnol dmP3 NaCI/H,O (pH 6). The number in the figure denotes the cycle number. dE/dt = 10 mV s-l. I' I I 1 I I 1 I 0.6 0.0' I I 1 I I I 1 1 1 400 500 600 700 wavelength /nm Fig.7. Visible absorption spectra of PAn films which were oxidized at 0.4 V us. SCE in 0.2 mol dmP3 NaCl/H,O(pH 4) in advance. (a) Immersed in 0.2 rnol dmP3 NaCI/H,O(pH 4). (h) Immersed in 1 mol dm-3 HCl. ( c ) Successively immersed in 0.1 mol dm-3 NaOH/H,O. sweep cyclings with anodic switching potential at 0.3 V us. SCE caused a steady decrease in voltammetric waves, though in solutions of pH 4 stable voltammograms were obtained with the same polarization conditions (see fig. 1). The unstable voltammetric behaviour observed at pH 6 must be caused by the inactivation reaction. It should be remarked here that in strong acid solutions the imposition of high anodic potentials such as 0.6 V us. SCE causes decomposition of PAn and not the ina~tivati0n.l~ If the electroinactive films were soaked in a high acidic solution such as 1 mol dmP3 HCl, a spectrum characteristic of an electroactive film as given by fig.5 (a) was obtained and indeed the resulting film exhibited electroactivities. Fig. 7 shows changes in absorption spectrum of an electroinactive PAn film by acid-soaking. Spectrum (a),T. Hirai, S. Kuwahata and H. Yoneyama 975 Scheme 1. obtained at an inactivated PAn film that was prepared by oxidation at 0.4 V us. SCE in pH 4, was changed into the spectrum (b) by simply soaking the PAn film in 1 mol dm-3 HC1. The blue colour of the film turned into yellow-green and became electroactive. However, if the film giving spectrum (b) was soaked in 0.1 mol dm-3 NaOH, the spectrum returned to an electroinactive one [spectrum (c)].These results are evidence that electroinactive films become electroactive with protonation.' The electroinactivation must then be caused by deprotonation of PAn. Redox Mechanism and Inactivation Mechanism of PAn in Weak Acid Solutions Scheme 1 gives proposed mechanisms for the redox reaction and the inactivation reaction of PAn in weak acid solutions such as pH 4. In this scheme, A- denotes an electrolyte anion such as C1-. The stable redox reaction of PAn at a potential range between -0.6 and 0.3 V PS. SCE concerns a redox reaction in which structures (I) and (11) are involved. Polyaniline films reduced in a weak acid solution must have a structure like leucoemeraldine base, because electrolyte anions were eventually absent in the reduced polymer.The intermediate oxidized form (TI) is deduced from absorption spectra given by fig. 5(a) which is almost the same as the spectra of an oxidized form in highly acidic solution such as pH 0. A similar redox mechanism has been presumed by Huang et al.13 based on their voltammetric studies at several different pH values. The oxidation of PAn at more anodic potentials, which causes the inactivation of the film seems to progress with the transformation from structure (11) to (IV) given in scheme 1. The oxidation in weak acid solutions does not accompany the decomposition of polymer chains and then the oxidation seems to result in an increase in the quinonedi- imine structure in the polymer. The occupancy of the quinonedi-imine structure is given by suffix ' a ' in structure (111), although the occupancy has not been determined.On the other hand, the inactivation of PAn is caused by the deprotonation reaction as976 Electrochemical Behaviour of Polyaniline mentioned before, and a resulting structure of PAn is deduced as structure (IV). In this scheme, suffixes ' b' and ' c' denote proportions of the quinonedi-imine structures which are deprotonated and protonated, respectively. The relative concentration of Cl to N was determined by elemental analyses to be 0.29 for an inactive PAn film prepared by oxidation at 0.8 V us. SCE in chloride solution (pH 4) giving the occupancy of the protonated quinonedi-imine structure. The rate of the deprotonation reaction of structure (111) to give structure (IV) seems to be higher than that of the protonation reaction of structure (IV) to give back structure (111), because the proton concentration in the PAn film may be relatively low in weak acid solution.Then once PAn is transformed to structure (IV), it may hardly return to structure (111), resulting in the apparent inactivation of PAn. The finding that the electroactivity of the inactive PAn film was recovered by immersing in a highly acidic solution gives support to the hypothesis that the inactivation observed in weak acid solutions is due to the slow rate of the protonation of the deprotonated quinonedi-imine structure. References 1 A. F. Diaz and J. A. Logan, J. Electroanal. Chem., 1980, 111, 1 1 1 . 2 T. Kobayashi, H. Yoneyama and H. Tamura, J. Electroanal. Chem., 1984, 177, 281. 3 A. Kitani, J. Izumi, J. Yano, Y. Hiromoto and K. Sasaki, Bull. Chem. Soc. Jpn., 1984, 57, 2254. 4 E. M. Genies and C. Tsintavis, J. Electroanal. Chem., 1985, 195, 109. 5 E. M. Genies and C. Tsintavis, J. Electroanal. Chem., 1986, 200, 127. 6 T. Kobayashi, H. Yoneyama and H. Tamura, J . Electroanal. Chem., 1984, 161, 419. 7 A. Kitani, J. Yano and K. Sasaki, J. Electroanal. Chem., 1986, 209, 227. 8 M. Kaneko, H. Nakamura and T. Shimomura, Makromol. Chem., Rapid Commun., 1987, 8, 179. 9 E. W. Paul, A. J. Ricco and M. S. Wrighton, J. Phys. Chem., 1985, 89, 1441. 10 S. Chao and M. S. Wrighton, J. Am. Chem. Soc., 1987, 109, 6627. 1 1 K. Okabayashi, F. Goto, K. Abe and T. Yoshida, Synth. Met., 1987, 18, 365. 12 A. G. MacDiarmid, L. S. Yang, W. S. Huang and B. D. Humphrey, Synth. Met., 1987, 18, 393. 13 W. S. Huang, B. D. Humphrey and A. G. MacDiarmid, J. Chem. Soc., Faraday Trans. I , 1986, 82, 14 H. Shinohara, M. Aizawa and H. Shirakawa, J. Chem. Soc., Chem. Commun., 1986, 87. 15 T. Kobayashi, H. Yoneyama and H. Tamura, J. Electroanal. Chem., 1984, 177, 293. 2385. Paper 8102240B ; Received 3rd June, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500969

出版商:RSC    

年代:1989

数据来源: RSC