摘要:

Contents 4259 4269 4277 4287 4295 431 1 4321 4335 Protonation Constant of Caffeine in Aqueous Solution M. Spiro, D. M. Grandoso and W. E. Price Ionic Equilibria in Acetonitrile Solutions of 2-, 3- and 4-Picoline N-oxide Perchlorates, studied by Potentiometry and Conductometry L. Chmurzynski, A. Wawrzyn6w and Z. Pawlak Liquid-phase Adsorption of Binary Ethanol-Water Mixtures on NaZSM-5 Zeolites with Different Silicon/Aluminium Ratios W-D. Einicke, M. Heuchel, M. v.Szombathely, P. Brauer, R. Schollner and 0. Rademacher Influence of Oxidation/Reduction Pretreatment on Hydrogen Adsorption on Rh/TiO, Catalysts. An lH Nuclear Magnetic Resonance Study J. P. Belzunegui, J. M. Rojo and J. Sanz Volumetric Properties of Mixtures of Simple Molecular Fluids A. C. Colin, E. G. Lezcano, A.Compostizo, R. G. Rubio and M. D. Peiia Study of Ultramicroporous Carbons by High-pressure Sorption. Part 4.-Iso- thems and Kinetic Transport in Activated Carbons J. E. Koresh, T. H. Kim, D. R. B. Walker and W. J. Koros Kinetic and Equilibrium Studies associated with the Solubilisation of n- Pentanol in Micellar Surfactants G. Kelly, N. Takisawa, D. M. Bloor, D. G. Hall and E. Wyn-Jones The effect of Carboxylic Acids on the Dissolution of Calcite in Aqueous Solution. Part 1 .-Maleic and Fumaric Acids R. G. Compton, K. L. Pritchard, P. R. Unwin, G. Grigg, P. Silvester, M. Lees and W. A. House 130-2Contents 4259 4269 4277 4287 4295 431 1 4321 4335 Protonation Constant of Caffeine in Aqueous Solution M. Spiro, D. M. Grandoso and W. E. Price Ionic Equilibria in Acetonitrile Solutions of 2-, 3- and 4-Picoline N-oxide Perchlorates, studied by Potentiometry and Conductometry L.Chmurzynski, A. Wawrzyn6w and Z. Pawlak Liquid-phase Adsorption of Binary Ethanol-Water Mixtures on NaZSM-5 Zeolites with Different Silicon/Aluminium Ratios W-D. Einicke, M. Heuchel, M. v.Szombathely, P. Brauer, R. Schollner and 0. Rademacher Influence of Oxidation/Reduction Pretreatment on Hydrogen Adsorption on Rh/TiO, Catalysts. An lH Nuclear Magnetic Resonance Study J. P. Belzunegui, J. M. Rojo and J. Sanz Volumetric Properties of Mixtures of Simple Molecular Fluids A. C. Colin, E. G. Lezcano, A. Compostizo, R. G. Rubio and M. D. Peiia Study of Ultramicroporous Carbons by High-pressure Sorption. Part 4.-Iso- thems and Kinetic Transport in Activated Carbons J. E. Koresh, T. H. Kim, D. R. B. Walker and W. J. Koros Kinetic and Equilibrium Studies associated with the Solubilisation of n- Pentanol in Micellar Surfactants G. Kelly, N. Takisawa, D. M. Bloor, D. G. Hall and E. Wyn-Jones The effect of Carboxylic Acids on the Dissolution of Calcite in Aqueous Solution. Part 1 .-Maleic and Fumaric Acids R. G. Compton, K. L. Pritchard, P. R. Unwin, G. Grigg, P. Silvester, M. Lees and W. A. House 130-2

ISSN:0300-9599    

DOI:10.1039/F198985FX001

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY ASSOCIAZIONE ITALIANA DI CHIMICA FlSlCA DEUTSCHE BUNSEN-GESELLSCHAFT FUR PHYSIKALISCHE CHEMIE KONINKLIJKE NEDERLANDS CHEMISCHE VERElNlGlNG SOCIETE FRANGAISE DE CHIMIE, DIVISION DE CHlMlE PHYSIQUE FARADAY DIVISION GENERAL DISCUSSION No. 90 Colloidal Dispersions University of Bristol, 10-12 September 1990 Orga nising Com mitte e Professor R. H. Ottewill (Chairman) Professor P. Botherol Professor E. Ferroni Or J. W. Goodwin Professor H. Hoff mann Professor A.L. Smith Professor P. Stenius Dr Th. F. Tadros Professor A. Vrij Dr D. A. Young The joint meeting of the Societies will be directed towards examining current understanding of the behaviour of colloidal dispersions. In particular, stability and instability, short range interactions, dynamic effects, non-equilibrium interaction, concentrated dispersions and order-disorder phenomena will form topics for discussion.The preliminary programme is now availablemay be obtained from: Mrs Y. A. Fish, The Royal Society of Chemistry, Burlington House, London W1V OBN. THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY SYMPOSIUM No. 26 Molecular Transport in Confined Regions and Membranes Oxford, 17-18 December 1990 Experimental, theoretical and simulation studies which address fundamental aspects of molecular transport will be discussed in the following main areas: a) Transport of atoms and molecules in pores, zeolite networks and other cavities; time-dependent statistical mechanics of small systems in confined geometries b) Molecular transport through synthetic membranes, biological membranes, smectic liquid crystalline phases and Langmuir Blodgett films; the dynamics of the molecules forming the membrane c) Diffusion, reorientation, conformational dynamics, viscosity and conductivity of polymer melts, to include papers dealing with bulk systems since the segments of the polymer will move in the anisotropic environment of the complete chain d) Applications of Brownian dynamics to the study of diffusion in porous media and across membranes including the transport of flexible aggregates such as microemulsions e ) The growth of crystals, colloidal aggregates and droplets on irregular surfaces and in pores Contributions for consideration by the Organising Committee are invited and abstracts of about 300 words should be sent by 31 December 1989 to: Dr D.J. Tildesley, Department of Chemistry, The University, Southampton SO9 SNH. Full papers for publication in the Symposium Volume will be required by August 1990.THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY ASSOCIAZIONE ITALIANA DI CHIMICA FlSlCA DEUTSCHE BUNSEN-GESELLSCHAFT FUR PHYSIKALISCHE CHEMIE KONINKLIJKE NEDERLANDS CHEMISCHE VERElNlGlNG SOCIETE FRANGAISE DE CHIMIE, DIVISION DE CHlMlE PHYSIQUE FARADAY DIVISION GENERAL DISCUSSION No. 90 Colloidal Dispersions University of Bristol, 10-12 September 1990 Orga nising Com mitte e Professor R. H. Ottewill (Chairman) Professor P. Botherol Professor E. Ferroni Or J. W. Goodwin Professor H. Hoff mann Professor A.L. Smith Professor P. Stenius Dr Th.F. Tadros Professor A. Vrij Dr D. A. Young The joint meeting of the Societies will be directed towards examining current understanding of the behaviour of colloidal dispersions. In particular, stability and instability, short range interactions, dynamic effects, non-equilibrium interaction, concentrated dispersions and order-disorder phenomena will form topics for discussion. The preliminary programme is now availablemay be obtained from: Mrs Y. A. Fish, The Royal Society of Chemistry, Burlington House, London W1V OBN. THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY SYMPOSIUM No. 26 Molecular Transport in Confined Regions and Membranes Oxford, 17-18 December 1990 Experimental, theoretical and simulation studies which address fundamental aspects of molecular transport will be discussed in the following main areas: a) Transport of atoms and molecules in pores, zeolite networks and other cavities; time-dependent statistical mechanics of small systems in confined geometries b) Molecular transport through synthetic membranes, biological membranes, smectic liquid crystalline phases and Langmuir Blodgett films; the dynamics of the molecules forming the membrane c) Diffusion, reorientation, conformational dynamics, viscosity and conductivity of polymer melts, to include papers dealing with bulk systems since the segments of the polymer will move in the anisotropic environment of the complete chain d) Applications of Brownian dynamics to the study of diffusion in porous media and across membranes including the transport of flexible aggregates such as microemulsions e ) The growth of crystals, colloidal aggregates and droplets on irregular surfaces and in pores Contributions for consideration by the Organising Committee are invited and abstracts of about 300 words should be sent by 31 December 1989 to: Dr D.J. Tildesley, Department of Chemistry, The University, Southampton SO9 SNH. Full papers for publication in the Symposium Volume will be required by August 1990.

ISSN:0300-9599    

DOI:10.1039/F198985BX003

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. I, 1989, 85(1), 11-21 Raman and Infrared Spectroscopy of the A1C13-SOC1, System Pamela A. Mosier-Boss, Roger D. Boss, Cedric J. Gabriel and Stanislaw Szpak Naval Ocean Systems Center, San Diego, CA 92152-5000, U.S.A. Jerry J . Smith Naval Weapons Center, China Lake, CA 93555-6001, U.S.A. Robert J . Nowak Ofice of Naval Research, Arlington, VA 2221 7-5000, U.S.A. The structural aspects of the AlC1,-SOCI, system have been examined by vibrational spectroscopy. The SOCI, molecule exhibits amphoteric charac- ter, i.e. it can act simultaneously as a donor through the oxygen atom and as an acceptor through the sulphur atom. The liquid state contains loosely bound, open-chain dimers/oligomers, (Cl,SO), with n 2 2. The dissolution of AI,Cl, in SOCI, occurs dissociatively with the formation of C1,SO -+ AlCl, adducts.At higher AICI, concentrations, an increase in solution electrical conductivity is attributed to the reaction : 2 C1,SO -+ AlCl, e [Cl,Al( +- OSCl,),]' + AlCl, Hecht,' followed by Spandau and Brunneck, proposed the existence of 1 : 1 and 2: 1 AlC1,-SOC1, adducts. In the early 1960s, Long and Bailey3 examined the structural features of these adducts and concluded that complexation occurred through the oxygen atom of SOCl,. A few years later, Auborn and co-workers4 employed this system in the construction of Li galvanic cells of highest practical energy densities. Their work prompted further inquiry into the nature, structure and reactivity of the various SOCl, complexes associated with the Li-SOCl, cell Here, we discuss the structure of neat SOCl, and is reactivity through interaction with a set of selected miscible liquids.We include an analysis of the structural changes and species produced by the addition of Al,Cl, to SOCl,.* Experimental Chemicals Thionyl chloride was refluxed under an argon or helium atmosphere to remove dissolved HCl and SO, with the end point indicated by pH paper located in the gas exit stream. The SOCl, was then distilled and the middle fraction was collected and stored under argon. Aluminium chloride, benzene, carbon tetrachloride and hexanes, all of spectroscopic grade (Fluka), and toluene and methylene chloride (Aldrich, gold label) were used as received. Solutions All solutions were prepared in a glove bag by diluting known volumes of SOCl, with CCl,, CH,Cl,, C,H,, or C,H,CH3.The AlC1,-SOC1, solutions were prepared from known amounts of AI,Cl, and filtered through a glass-fibre filter prior to use. 1112 Raman and I.R. Spectroscopy of AlC1,-SOCl, Instrumentation Infrared spectra were obtained on a Nicolet 5DXB FT-IR spectrometer with a resolution of 2cm-'. Standard demountable cells with NaCl windows were used throughout. The pathlength was maintained at 2.5 x m; spectra were subtractively normalized to compensate for solvent absorption. The Raman spectra were recorded on a system comprising a Lexel model 85 Ar-ion laser emitting at 488 nm, which was chopped at 20 Hz, a sample chamber with coupling and collecting optics, a Spex model 1400-1 I scanning double monochromator with photon-counting preamplifier and line driver, a photon-counting integratorg and a digital synchronous detector.The samples were contained in a Helma model 162 F stoppered quartz cell. The cell chamber was flushed with dry nitrogen to prevent condensation on the cell windows when operating below room temperature. Tem- peratures were controlled to within kO.1 "C. Most of the spectra were recorded at a laser power of 50 mW. Resolution of Spectral Bands The recorded infrared absorption and Raman-scat tering bands were computer-analysed and decomposed into their component Voigt profiles using a previously described procedure.1° In a dynamic system, comprising a number of species, N , with overlapping spectral lines, the experimentally observed spectral band intensity, I( v), is a superposition of at least N lines, and can be approximated by eqn (1) N I(v) = C,(v-a)+ c CiL( ...) i- 1 where a, C, and Ci are adjustable parameters and f( ...) is defined by eqn (2) where and v,, Av, and AvG are, respectively, the centre frequency, and the full widths at half maximum of the Lorentzian and Gaussian distributions that characterize the line.The linear term Co(v-a) has been added to account for a sloping baseline. a = Av, d ( l n 2)/Av,, cu = 2(v- vo) d(In 2)/Av,, Molecular-orbital Calculations Calculations were performed using AMPAC, a general purpose, semi-empirical molecular- orbital package developed at the University of Texas (Austin, TX). These calculations yield information on electronic and core-core repulsion energies, heats of formation and vibrational frequencies. Results and Discussion The molecular structure of a liquid arises from interactions between neighbouring molecules.This premise allows us to follow structural changes as the molecules adapt themselves to the changing environment, e.g. to changes in solution composition or temperature. For example, SOCl,, being an amphoteric molecule, may form dimers/ oligomers. In the more complex AlC1,-SOCI, system, such interactions can lead to the formation of molecular adducts and ionic species.8P. A . Mosier-Boss et al. 1.0 13 V , ( A ' ) 6,(0SCI 1 - - I ICl1 I 200 400 600 800 1000 1200 1400 wavenumber/cm-' Fig. 1. Raman spectrum of neat SOCI,: vibrational modes and assignments. Method of Analysis The SOCI, molecule is of the Z X Y , type and a member of the C, symmetry point group." It has six normal modes of vibration, all Raman and infrared active.The Raman spectrum of neat SOCI, and the assignments are shown in fig. 1. The structural characteristics of SOCI, in the liquid state as well as the AlCl,-SOCl, system are ascertained by examining the position and lineshape of v,(A'), the symmetric S-0 stretching vibration at 1231 cm-l and the v2(A') and v,(A') symmetric and asymmetric S-Cl stretching vibrations at 492 and 455 cm-', respectively ; these vibrational modes are sensitive to changes in the charge distribution arising from molecular interactions. The S=O bond of SOCI, has partial double bond character which results from the superposition of pn - dn back-bonding from 0 to S upon the S + 0 bond.', According to the valence-shell electron-pair repulsion (VSEPR) model, bonding through the oxygen atom should lessen the pn --+ dn back-bonding and, hence, lower the S=O bond- order and stretching frequency.Conversely, bonding through the sulphur atom increases the pn -, dn back-bonding, thus raising the S-0 stretching frequency. Withdrawal of electron density from the S=O bond will decrease the repulsion between the lone pair of electrons on the sulphur and the chlorine atoms of SOCI,, which, in turn, strengthens the S-Cl bonds and shifts the symmetric and asymmetric S-Cl stretching vibrations to higher frequency. Moreover, coordination through the sulphur atom will reduce the repulsion between the sulphur lone pair of electrons and the chlorines, also shifting the symmetric and asymmetric S-Cl stretching vibrations to higher frequencies.Examples of this kind of behaviour have been reported for the metallic complexes of (CH,),SO, in which complexation with Zn2+, A13+, Ni2+, Co2+, Fe2+ and Fe3+ ions occurs through the oxygen atom while complexation with Pd2+ and Pt2+ occurs through the sulphur atom.', Thionyl Chloride in the Liquid State In the liquid state, the S-0 stretching frequency is at 1231 cm-'; whereas, in the gaseous state it occurs at 1251 cm-'.14 This rather small shift to lower frequency upon14 Raman and I.R. Spectroscopy of AlCl,-SOCl, 1 .O A 0 B 1150 1175 1200 1225 1250 1275 1300 w avenumber im- ' Fig. 2. Decomposition of S-0 stretching spectral band of neat SOCl, into Voigt profiles: A, 23.5 "C; B, -20.0 "C.condensation of SOC1, implies that the intermolecular interactions are weak. At 23.5 "C the S-0 stretch of neat SOC1, at 1231 cm-l, shown in fig. 2, is actually a composite band which has been resolved into two Voigt profiles with peaks at 1242.6 and 1230.5 cm-l, profiles (a) and (b) respectively. However, at -20 "C the S-0 stretch of neat SOCl, can be resolved into three Voigt profiles with peaks at 1243.5, 1230.8 and 1221.5 cm-', profiles (a), (b) and (c), respectively, in fig. 2. This band structure suggests that pure SOCl, is a weakly associated liquid, which, in view of the rather low latent heat of vaporization and the numerical value of the slope of the fluidity as a function of specific The existence of the association implies that the SOCl, molecule is amphoteric, i.e.it acts simultaneously as a donor through the oxygen atom, and as an acceptor through the sulphur. Thionyl chloride may self-associate in two ways ; via sulphur-to-oxygen bonding in either a cyclic-dimer form, or an open chain form, which may contain more than two members, and via sulphur-sulphur association. In longer chains, a cooperative effect may increase the donor character of the S=O bond. Dilution experiments were performed to determine the structure of the associated indicates the presence of small, interacting molecular clusters.P . A . Mosier-Boss et al. 15 species of SOC1,. Infrared spectroscopy was used to examine the changes in the lineshape and position of the S-0 stretching vibrational band upon dilution in three types of solvents : (a) poor acceptors (inert solvents), e.g.C6H14 and eel,; (b) n-interacting solvents, e.g. C,H,CH,; and (c) a polar liquid, e.g. CH,Cl, (with p = 1.6 Dt).', On the basis of the VSEPR model, as well as more general donor-acceptor considerations," the following characteristics of the S-0 stretch of the monomer and associated species of SOCl, are expected. For a monomer, the S-0 stretching vibrational peak should be fairly narrow and its peak position should be solvent dependent. The cyclic dimer, should also exhibit a fairly narrow peak but its position should be relatively solvent independent. Furthermore, because of the decrease in repulsion between the lone pair of electrons on the sulphur atom and the negative charge on the chlorine atoms, the separation between the S-Cl symmetric and asymmetric stretches, v, and v,, respectively, should be less than that for the monomer.For an open chain dimer/ oligomer, the peak position of the S-0 stretching frequency is expected to be solvent dependent and the peak should be broader than that of the monomer. The open-chain dimer probably has two different S=O bonds. The extent to which these differ will depend on the strength of the interaction within the dimer. If the strength of the interaction were on the order of an ionic or covalent bond then both these S=O vibrational bands would be observed. However, as the strength of the interaction decreases, the separation between these two bands would also decrease eventually resulting in one broad band. This would be the expected result for dipole4ipole and dipole-induced-dipole interactions.Finally, the separation between v, and v, is expected to be greater for the open dimer than for the monomer due to increased repulsion between the chlorines and lone pair of electrons on the sulphur atom. Results of dilution experiments are as follows. In all solvents, within the SOC1, concentration range studied, the S-0 Voigt profile at 1221.5 cm-', v(c) in fig. 2, is not observed. For this reason and the fact that it is observed only in neat SOCl, at low temperatures, this band is attributed to higher aggregates, i.e. to trimers, tetramers etc. Furthermore, it is seen that the peak position of the other two Voigt profiles is solvent dependent. In CCI, and hexanes, which are considered to be 'inert' solvents, the high- frequency peak, v(a), is narrower than the low-frequency peak, v(b).In these solvents, the association of SOC1, is expected to be the dominant reaction. Therefore the high frequency peak, v(a) is assigned to the monomer and the low frequency peak, v(b), to dimer. The ratio of the areas of these peaks is consistent with these assignments. Since v(b) is broader than v(a), we conclude that the dimer is one of the open type as shown in fig. 3(a). Further evidence for an open structure can be found in the separation between the symmetric and asymmetric S-C1 stretches, v, and v,. In the gaseous state this separation is 37 cm-l, whereas in the liquid state, it is 47 cm-'.14 In methylene chloride, both v(a) and v(b) of the S-0 stretching bands are very broad.With a dipole moment of 1.60 D, CH,C1, is a weakly polar solvent; thus one expects dipole4ipole interactions to occur between CH,C1, and SOCl,. Such interactions have been observed for (CH,),SO and CHCl,.15*16 The behaviour of the S-0 stretching composite band of SOCl, solutions in C,H,CH, is similar to that observed in CCl, and C6H14; i.e. v(a) is fairly narrow and v(b) is broader. From this we conclude that any interactions between SOC1, and C,H,CH, are very weak and that self-association of SOC1, is the dominant reaction. The results of the MO calculations provide additional insights into the possible structures of neat SOCI, in the liquid state. In particular, calculations have been performed for dimeric SOCI, using the structures S-0-S-0 in cyclic form and in open chain, as well as the 0-S-S-0 type of association.The calculations showed that the oxygen-sulphur associated dimers are favoured, fig. 3. The sulphur-sulphur t 1 D = 3.33564 x C m.16 Raman and I.R. Spectroscopy of AlC1,-SOCl, C l C l 0 ' C l Fig. 3. MO-optimized structures: (a) open dimer; (6) cyclic dimer. associated species gives an S-S bond length exceeding 6 (6 x lo-'" m) thereby suggesting little or no association uia this bonding arrangement. Furthermore, the MO calculations indicate that ordered dimeric structures, as well as longer oligomeric structures, can exist in the liquid state with essentially the same stability as the molecular SOCl, itself, and, yet, are not of significantly different stability so as to dominate the structure of the liquid SOCl,.The dilution experiments supported by molecular-orbital calculations indicate the presence of dimers with the open-chain spatial arrangement illustrated in fig. 3 (a) rather than the cyclic form, fig. 3(b). In fact, calculations of cyclic dimer structures invariably led to optimized geometries in which the cyclic dimer opened to form an open-chain dimer. These MO calculations yield structures whereby the lone pairs of the sulphurs are directed away from one another. This implies that the repulsion between the lone pair of electrons on the sulphur atom is principally responsible for molecular structure of SOC1, dimers. Similar arguments for the formation of the cyclic dimers results in a highly improbable structure; the repulsion between the lone pair of electrons with the simultaneous requirement for the charge transfer for 0, -, S, and 0, + S,, creates a distorted arrangement.In summary, SOCl, in the liquid state consists of weakly associated species having an open structure. The presence of cyclic dimers, which could be possible, is excluded on a basis of MO calculations.P. A . Mosier-Boss et al. 17 0 100 200 300 400 500 600 1000 1050 1100 1150 1200 1250 Raman shift/m-' Fig. 4. Evolution of Raman Spectra as a function of AICI, concentration (mol drn-,). The AlC1,-SOCI, System With the addition of Al,Cl,, a covalent compound containing halogen bridges, a new set of Raman bands appears and indicates the formation of adduct(s). In the course of 1 : 1 adduct formation, the Al-Cl-A1 bonds of the acceptor molecular must first be broken.The improvement in coordination makes the formation of the 1 : 1 adduct energetically favourable. The progressive changes in the Raman spectra are shown in fig. 4. Upon the addition of Al,Cl, the Raman spectra become more complex. In addition to bands at 1108, 523, 383, 217, 167 and 114 cm-l, a band at ca. 1055 cm-l emerges, and gains in intensity with increasing Al,Cl, concentration. It is noteworthy that the appearance of this band is accompanied by an increase in the solution conductivity.*18 Raman and I.R. Spectroscopy of AlC1,-SOCl, I I 1 I I I I 200 400 600 800 1000 1200 1400 w avenumber/cm-' Fig. 5. Polarized (a) and depolarized (b) Raman Spectra of the equimolar AlC1,-SOCl, solution.The polarized and depolarized Raman spectra of the equimolar A1Cl3-SOC1, solution are shown in fig. 5. The most significant changes are the disappearance of the band at 123 1 cm-' the S-0 stretching vibration of neat SOCl,, and an emergence of a new band at ca. 1108 cm-'. Such a change in the S-0 stretching frequency accompanied by the shift observed in the symmetric and asymmetric S-Cl stretches to higher frequencies (i.e. from 492 to 523 and from 455 to 500cm-', respectively) is consistent with complexation of AlCl, through the oxygen atom of SOC1,. The observed vibrational modes of the AlC1,-SOC1, complex(es) and their assignments are shown in table 1. These assignments correspond with those made for the AlCl, complexes with tetrahydrofuran, C,H,O l9 and nitromethane, CH,N0,.20 Species and Equilibria An examination of the 1000-1 300 cm-' spectral region, fig.4, indicates the formation of, at least, two distinct species. The first species, with the S-0 stretching vibrational frequency v,(A') at 1108 cm-', has been attributed to the 1 : 1 adduct, C1,SO -+ AlC1,. The identification of the species exhibiting a broad band at 1055 cm-' is less certain. For example, the broadness of this peak as well as the shift of v,(A') to lower frequency indicates further electron withdrawal from the oxygen atom perhaps from the formation of a 1 : 2 complex, Cl,SO(AlCI,), by the reaction (1) Such a conclusion is consistent with the existence of 1 : 2 complex in the solid state., Spectroscopically, this equilibrium requires that the ratio of the area of the band at 1055 cm-l to the area of the 1108 cm-' band should increase with increasing AlCl, concentration.However, it is observed that this ratio is nearly constant at 0.3 0.1 for all concentrations of AlC1, in which the two bands can be resolved. Fig. 6 shows the decomposition of the S-0 stretch of the AlC1, adducts. Furthermore, an increase in the solution electrical conductivity which coincides with the appearance of this band offers an alternative interpretation, namely that of formation of ionic species. Molecular adducts are neutral molecules. Formation of ions from the C1,SO -+ AlCl, 2 C1,SO + AlC1, + C1,SO -+ (AlCl,), + SOC1,.P. A . Mosier-Boss et al. 19 Table 1. Vibrational assignments and calculated frequencies for the Raman spectrum of the 1 : 1 AlC1,-SOCl, complex and other observed vibrations for the 1 : 1 solution vibrational mode vobs.d/cm-' Vcalcd/cm-' observed calculated frequency frequency 0-Al-Cl bending of complex AlCl, rocking of complex AlCl, bending of complex AICl, bending of complex Cl-Al-Cl bending of AlCl; deformation of complex SCl, bending of free SOCl, SCI, bending of complex Al-0-S bending of complex SOCI, torsion of free SOCl, SOCl, torsion of complex SOCl, deformation of free SOCl, SOCl, deformation of complex, Al-CI A1-0 stretching of complex A1-0 stretchingb A1-0 stretchingb S-Cl asymmetric stretching of free SOCl, AI-CI asymmetric stretching of complex S-Cl asymmetric stretching of free SOCl, S-Cl asymmetric stretching of complex S-Cl symmetric stretching of complex AlCl, degenerate stretching of complexc S-0 stretchingb S-0 stretching of complex S-0 stretching of free SOCl, symmetric stretching of AlCl, and AICl, < 110 114 - 167 181 (w) - 194" (P) 217 274 282 318 355 344 (P> 383 419 (w) 428 (w) 455" 492" (PI 500 523 560 (w) 1055 1 108 (P) 1231 (P) 35,101 116 121 157 184 167 192 283 - - 340 335 503 - - 566 395,566 615 608 645 588 1330 1374 (w) weak.(p) polarized. a not observed in the equimolar solution. band due to either Cl,(SO) --+ (AlCl,), or Cl,Al( + OSCl,);. compare with 525 cm-' for the 1 : 1 AlCl, : tetra- hydrofuran complex, 532 cm-' for the 1 : 1 AlCI,: nitromethane complex, and both bands are weak in the Raman. adduct can occur either via the halide ion transfer or the internal exchange route, eqn (11) and (111), respectively.C1,SO -, AlCl, SOCl+ + AlC1; (11) (111) 2 C1,SO -, AlC1, + Cl,Al( t OSC1,); + AlCl,. Of these ions, the presence of the AlC1, ions is documented by a weak band at 181 cm-', and attributed to the C1-Al-C1 bending mode, fig. 5. The stronger bands associated with a AlCl,, complex ion, i.e. the bands at 122 and 350 cm-', are obscured by vibrational modes of the free and complexed SOC1,. The AlC1, ion also has a band at 495 cm-l which is strong in the i.r.,, An attempt was made to see this band by adding LiCl to 4.0 mol dm-, AlCl, in SOCl,. Such a solution would contain AlCl; as well as Li(SOC1,)~. Owing to the overlapping S-Cl vibrational bands of neat SOC1, as well as SOC1, complexed by AlC1, and Li+, it was not possible to discern the 495 cm-l band of AlCl,.There is no spectroscopic evidence for the positively charged SOCl+ species, as required by eqn (11) and suggested by others.'-, On the other hand, the evidence for the20 Raman and I.R. Spectroscopy of AlC1,-SOCI, 0.2 - 950 1000 1050 1100 1150 1200 Fig. 6. Decomposition of the S-0 stretching spectral band associated with the aluminium complex of 4.0 mol dm-I AlC1, in SOCl, at -20.0 "C into Voigt profiles. presence of the [CI,Al(+ OSCl,),]' ions, eqn (111), is as follows. The S=O bond in the onium ion must be weakened due to the reduced pn -+ dn back-bonding, thus lowering the vibrational frequency us. that of a neutral adduct. The appearance of a new band at 1055 cm-l supports its presence. Furthermore, because the [CI,AI( + OSCI,),]' would have a positive charge on the A1 atom, its A1-0 bond would be stronger than that for a neutral complex.This, in turn, would shift the Al-0 stretching vibrations to higher frequencies. Indeed, in the neutral adduct the A1-0 stretching mode occurs at 383.3 cm-l while in the conductive, equimolar AlC1,-SOC1, solution, new weak bands attributed to the ionic complex, were found at 419 and 428 cm-'. In addition, eqn (111) requires that the concentration of the AICl, ion must be equal to that of the Cl,AI(t OSCI,),+ ion. The equilibrium constant expression may be rearranged to give eqn (IV) : (IV) which requires that the areas of the bands at 1055 and 1108 cm-' be in constant proportion, i.e. independent of the AICl, concentration, as observed. The results of the MO calculations of relative stabilities of the mono- and di-solvated onium complex favour the latter as the dominant species responsible for the increase in solution conductivity. The suggested structure is a tetrahedral arrangement about the Al with SOCl, coordination through the oxygen, i.e.consistent with eqn (IV). On the basis of the argument presented, the 1055cm-' band is assigned to the [CI,AI( + OSCI,),]' ion. [CI,Al( + OSCl,),]' = K[Cl,SO + AICl,] Conclusions (a) Neat SOCl, is an associated liquid which forms open chain dimers/oligomers (CI,SO), with n = 2,3,4 . . . ; (b) aluminium chloride dissolves dissociatively in SOCI, i.e. as AlCl, not Al,Cl,; (c) complexation of SOCl, with AIC1, occurs through the oxygen to form the adduct, CI,SO -+ AICl,; ( d ) the C1,SO -+ AICl, adduct dissociates to yield the ionic species [CI,Al(+ OSCl,),]' and AICl,.This work was in part supported by the Office of Naval Research.P. A. hfosier- Boss et al. 21 References I H. Hecht. Z. Anorg. Client.. 1947, 254. 44. 2 H. Spandau and E. Brunneck. Z. Atiorg. Cltem., 1952. 270, 201 ; 2. 3 D. A. Long and R. T. Bailey. Trans. F~irutlq. Soc., 1963. 59. 594. 4 J. J. Auborn. K. W. French. S. I. Lieberman. V. K. Shah and A. Heller. J. Elecrrocheni. Soc.. 1973, 5 K. C. Tsaur and R. Pollard. J. Electroc*iwnt. Soc-., 1984. 131. 975; 984; 1986. 133. 2296. 6 M. J. Madou, J. J. Smith and S. Szpak. J. El~~c~trocherii. Sue.. 1987, 134. 2794. 7 J. R. Driscoll, R. Pollard, J. J. Smith and S. Szpak. in Powrr Sources I ! (Academic Press. 1987). 8 S. Szpak and H. V. Vcnkatasetty, J. E1t~ctrocli;wt. Sot.. 1984. 131. 961. 9 S. A. Miller. Rer. Sci. Insrr.. 1968. 39. 192-3. 120. 1613. 10 P. A. Mosier-Boss. C. J. Gabriel and S. S~pak. Spec.rrodtirit. Acru, Purr A. 1987. 43. 1293. 1 I R. A. Suthers and T. Henshall. Z. z4n0rg. ,-ffl,q, Chent.. 1972. 388, 269. 12 F. A. Cotton and R. Francis. J. Ant. Citcwi. Soc.. 1960. 82, 2986. 13 J. Selbin, W. E. Bull and L. H. Holmes Jr. J. Inorg. .\'id. Ciiem.. 1961. 16, 219. 14 D. E. Martz and R. T. Lagemann, J. Ciicw. P/i.r.s.. 1954. 22. 1193. I5 W. Lin and S. Tsay. J. Pirjx C i i m . , 1970. 74. 1037. 16 A. L. McChellan. S. W. Nicksic and J. C. Guffy. J. itlo/. Spec-trusc.. 1963. 11. 340. 17 M. Dalibart, J. Derouault. P. Granger and S. Chapellr. Inorg. Clreni.. 1982. 21, 1040. 18 J. Dcrouauit and M. T. Forel. Inorg. Clirtn.. 1977, 16. 3207. 19 .I. Dcrouauit, P. Granger and M. T. Ford. h r g . Chenr., 1977, 16, 3214. 20 M. Dalibart. J. Deroault and P. Granger. I t t o y . Clte~ir.. 1982. 21, 2241. 2 1 1. Lindquist. Inor,gunic* At1ciuc.r Molecu1c.s qt' O . ~ o - c . o ~ ~ i p o u ~ ~ ~ ~ . ~ (Academic Press, New Y ork. 1963). 22 V. Gutmann. Tiw Donor- Acceptor Approcidt to Molt~cdur Interactions (Plenum Press, New York, 1978). 23 J. E. H. Jones and J. L. Wood. Specsrrodtinr. .-le.tti. f l i r t A . 1967. 23. 2695. Pciper 712000; Receired 10th .Yorenther, 1987

ISSN:0300-9599    

DOI:10.1039/F19898500011

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. SOC., Furaday Trans. I , 1989, 85(1). 23--32 Ionic Contributions to the Viscosity B Coefficients of the Jones-Dole Equation Part 5.-Acetonitrile Kenneth G. Lawrence* Department of Chemistry, Birkbeck College, Malet Street, London WCl E 7HX Antonio Sacco and Angelo De Giglio Department of Chemistry, University of Bari, Via Amendola 173, 70126 Bari, Italy Angelo Dell'Atti Department of Physics, Uniuersitj, of Lecce, Via Arnesano, 73100 Lecce, Italy The B coefficients of the Jones-Dole viscosity equation are a measure of the size of the ions and of the interaction between the ions and the solvent. The B coefficients have been determined for the electrolytes Bu,NBu,B, Bu,NBr, Bu,NI, Ph,PBr, Ph,PI, NaI and NaPh,B, and for the homologous series from Et,NBr to Hept,NBr in acetonitrile at 25 and 35 "C.Ionic B values for the bromide and iodide ions have been calculated from the B coefficients for the tetra-alkylammonium solutions and are compared with those obtained from the tetraphenylphosphonium solutions. The transition-state treatment has been applied to the results, and the thermodynamic activation parameters for viscous flow have been calculated. These are compared with those found previously for solutions of dimethyl sulphoxide, hexa- methylphosphoric triamide and N,N-dimethylformamide, and are dis- cussed in terms of the new theory of B coefficients proposed by Feakins. For a number of years we have been studying the viscometric properties of non-aqueous solutions of electrolytes. Viscosity B coefficients of the Jones-Dole equation q r = I+AC;+BC (1) have been determined for each system studied, because these coefficients throw some light on the solute-solvent interactions.It is, however, more useful to be able to divide the B coefficients for the salts into ionic contributions, and our most recent work has focused attention on two of the methods of subdivision. One of these assumes that for salts like tetrabutylammonium tetrabutylborate (Bu,NBu,B) and the corresponding tetraphenylphosphonium salt (Ph,PPh,B) similar cation-solvent and anion-solvent interactions occur for each reference salt, and the contributions made to the viscosity of the solutions by the Bu,N+, Bu,B- and the Ph,P+, Ph,B- ions may be considered to be proportional to their van der Waals volumes, Vw. For example, the cationic B value from the tetrabutyl salt is obtained from the following: V,(Bu,B-) B( Bu,N+) = B( Bu,NBu,B) We have applied this method using both the tetrabutyl and tetraphenyl salts to a few solvent systems with remarkable suc~ess.l-~ In this paper we report viscosity and density measurements of Bu,NBu,B, Bu,NBr, 2324 Ionic Contributions to the Viscosity B Coeficients Bu,NI, Ph,PBr, Ph,PI, NaI and NaPh,B in ACN at 25 and 35 "C.Ph,PPh,B was insufficiently soluble, so a B value for this salt was obtained from (3) Although we would normally prefer to use bromide salts, which in our experience give more consistent results, we found that the Ph,PI was more soluble in ACN, and the B coefficients could then be determined with greater precision from measurements made over a wider concentration range.The tetrabutylammonium salts did not pose the same solubility problem. Additional measurements were made with Bu,NI and Ph,PBr for purposes of halide-ion comparison. In a previous paper we examined another method of determining ionic B values that had been reported in the literature. This involved plotting B coefficients of a series of tetra-alkylammonium halides in dimethyl-sulphoxide and hexamethylphophoric tri- amide as some function of the cation, and extrapolating the curve to obtain the B value of the halide ion, This was not very successful, but we thought it worthwhile to repeat this method in acetonitrile and so measurements were also made with tetraethyl- to tetraheptyl-ammonium bromides. B(Ph,PPh,B) = B(Ph,PI) + B(NaPh,B) - B(Na1). Experiment a1 Details of the purification of the salts and of the apparatus used have been given previo~sly.l-~ Bu,NI (Ega-Chemie), purity > 99 %, m.p. 147-148 "C, NaI (Merck Suprapur) and NaPh,B (Gold Label) purity > 9970, were vacuum-dried and used without further purification.Ph,PI was purified as reported in the literature. ACN (Fluka), purity > 99.8 YO, water content determined according to Karl Fischer < 0.008%0, was used without further purification. All the solutions were prepared in a dry box. Results and Calculations In table 1 are reported the experimental results for the salts examined at 25 and 35 "C. Table 2 shows calculated and experimental viscosity A coefficients. The constants used in the Falkenhagen-Vernon equation to calculate the A values at 25 "C were E = 35.95,4 and 21 = 0.3406 CP (this work), and limiting ionic conductances were taken from the literat~re.~ The experimental A values were obtained from a linear-regression least- squares fitting.The viscosity B coefficients were calculated by the method of orthogonal polynomials with theoretical A coefficients substituted in the statistical analysis. The values of the A coefficients used in the calculations at 25 "C were also used at 35 "C, a procedure we have used and discussed previously.' The B value for Bu,N+ was calculated using eqn (2). The value for Ph,P+ was obtained using eqn (3) and the appropriate form of eqn (2). The van der Waals volumes were taken from table 4 of ref. (1). Apparent molar volumes, #", of the solutions were calculated from the equation 1000(do-d) M CdO do +- q5v = (4) where d and do are the densities of the solution and solvent, respectively, A4 is the formula weight of the solute and c is the concentration.The partial molar volumes, #:, of the salts at infinite dilution were obtained by a least-squares fitting of the results to the Masson equation where S is an empirical constant. The values of q5: are reported in table 4. q5v = #:+sci ( 5 )K . G. Lawrence, A. Scicco, .4. Dc, Giglio and A . DelrAtti 25 Table 1. Data for the relative viscosity (q,.) and density (p,) at 25 and 35 "C 25 "C 35 "C Et,NBr 73 104 145 I63 227 305 382 486 107 157 215 265 314 383 437 505 1 I4 151 250 328 394 470 539 606 143 245 324 41 8 516 599 650 742 71 107 148 20 I 277 328 41 6 499 76 106 179 215 278 350 399 486 Pr ,N BR Bu,NBr Bu,NI Pe,NBr Hex,N Br 454 667 916 980 1354 I788 2270 2772 56 1 798 1126 1280 1772 227 1 3012 3 743 418 63 1 762 909 1303 I685 2310 2714 75 1 I3 150 179 237 305 402 503 569 810 1142 1298 1797 2303 3054 3795 814 1227 1705 2121 2525 31 I5 2429 3879 733 1071 I439 1750 2058 2494 2727 308 1 803 1210 1681 2092 2490 3073 3382 3816 792 1056 1452 1730 2027 2464 2667 3005 142 162 245 287 337 403 45 I 500 121 161 26 1 335 408 483 578 65 1 825 1100 1880 2426 3008 3568 4076 4645 855 1 1 1 1 1817 232 I 2825 3348 3803 430 1 813 1084 1854 2392 2966 3518 4020 4582 836 1090 I778 2259 2745 3252 3730 4212 728 1322 774 1290 1687 2123 3640 2992 328 1 3764 142 249 330 42 1 55 1 608 676 766 738 1341 181 1 231 1 2895 3314 3652 4209 778 1316 1734 2170 2675 3062 3344 3833 1786 2278 2856 3268 3602 4151 514 723 1074 1507 2017 23 74 3070 3708 608 755 1 I95 I669 2172 2610 3285 3877 83 99 151 209 280 3 30 434 526 52 I 733 I089 1528 2046 2408 31 12 3760 600 803 1247 1675 2376 2750 3404 4019 719 968 1349 1857 243 1 301 I 3485 4229 77 107 159 229 306 386 428 51 1 539 76 1 1111 1830 2027 2469 290 1 349 1 739 1004 1412 1905 2498 308 I 3520 4208 53 1 750 1095 1509 1999 2435 286 1 344326 Ionic Contributions to the Viscosity B Coeficients Table 1 (cont.) 25 "C 35 "C c / ~ O - ~ mol dm-3 105(qr- 1) lo5@,- 1) c / ~ O - ~ mol dmP3 105(qr- 1) lo5@,- 1) - 532 780 1113 1534 1993 2432 2973 800 1306 1814 2377 2900 3436 3974 4407 423 596 750 942 1171 1342 1499 1678 359 524 732 974 1123 1204 1372 1583 788 1146 1513 1960 2325 2705 301 1 342 1 804 1130 1583 2000 2273 28 14 3089 820 1235 1742 2265 2885 3494 423 1 995 1573 21 17 2739 331 1 3919 4434 5000 67 1 914 1109 1368 1686 1918 21 19 2377 353 51 1 664 876 989 1066 1197 1366 1034 1472 1909 2425 2860 3292 3658 4128 1031 1428 1952 2443 2756 3379 3726 Hept,NBr 76 525 109 769 206 1097 212 1513 280 1965 345 2398 42 1 2932 Bu,NBu,B 36 68 78 101 119 143 160 179 76 116 141 170 214 244 270 298 68 99 130 179 20 1 223 25 1 287 203 293 406 505 60 1 689 774 868 242 344 478 607 68 1 84 1 937 NaPh,B NaI Ph,PBr Ph,PI 789 1288 1789 2344 2860 3389 3920 4346 417 588 740 929 1154 1324 1478 1655 354 538 72 1 960 1107 1188 1360 1521 777 1130 1492 1933 2293 2668 2969 3374 793 1115 1561 1972 2242 2775 3047 817 1630 171 1 2203 2819 3498 4206 965 1536 2067 2667 3215 3782 4303 4840 667 900 1105 1325 1667 1903 2077 2275 337 479 657 830 949 01 1 189 335 006 438 1828 2344 2794 3207 3545 4024 1022 1388 1901 2374 2699 3279 3630 94 112 184 223 297 360 438 38 82 85 123 138 155 180 20 1 79 121 173 183 232 320 288 313 61 98 113 159 190 204 205 274 189 275 368 486 60 1 687 76 1 8 70 243 339 475 602 686 848 948K .G. Lawrence, A . Sacco, A . De Giglio and A . Dell'Atti 27 Table 2. Viscosity A/dmg mol-; coefficients at 25 "C Et,NBr Pr,NBr Bu,NBr Pe,NBr Hex,NBr Hept,NBr Bu,NBu,B Bu,NI Ph,PI NaPh,B NaI Ph,PBr theoretical 0.0 158 0.0 172 0.0 I84 0.0191 0.0 199 0.0205 0.0236 0.0 182 0.0 187 0.02 18 0.0 164 0.0 190 ~~ ~ experimental" 0.0108 & 0.0030 0.01 89 & 0.0009 0.0 I 70 & 0.0007 0.0070 0.0059 0.0 18 1 & 0.0022 0.0234 f 0.0052 0.01 83 f 0.0022 0.0203 & 0.0009 0.01 53 & 0.001 1 0.021 8 & 0.0016 0.0150+ 0.0016 0.01 87 f 0.0007 a With standard error.Table 3. B/dm3 mol-' coefficient differences for salts with a common ion __ -~ ~ ~~ _ ~ _ _ _ ~ ~ ~~ T/ "C Bu,NBr - Bu,NI Ph,PBr - Ph,PI Ph,PBr - Bu,NBr Ph,PI - Bu,NI 25 0.018 0.0 I6 0.265 0.267 35 0.014 0.013 0.256 0.257 Discussion A Coefficients It is not uncommon to find that, for non-aqueous systems possessing large B coefficients, the agreement between the experimental and theoretical A coefficients is not as close as that found for aqueous systems. If ion pairing is occurring to any great extent experimental A values are found to be consistently higher than the theoretical values,' but this was not the case here.Since the theoretical A values were used in the orthogonal polynomials program from which the B coefficients were obtained, the program was re- run with the experimental values substituted, but the resulting B values were not sufficiently changed to warrant reporting them. In any case, the purpose of this work is to investigate our ideas about splitting the B coefficients into ionic contributions, and for this we require internally consistent B values (see below). Confidence in the wider interpretation of the B values will then depend on the relative magnitudes of the experimental errors, which are reported in tables 4 and 5. B Coefficients The internal consistency of the results may be checked by calculating the differences between the B coefficients of two salts with a common ion.Satisfactory agreement for pairs of bromide-iodide and tetraphenyl-tetrabutyl salts at both temperatures can be seen in table 3. Direct comparisons of our results with those available in the literature can be made in table 4. We have shown' that Bu,NPh,B should not be used as a reference salt for28 Ionic Contributions to the Viscosity B Coejicients Table 4. Viscosity B/dm3 mol-' coefficients and partial molar volumes q5:/cm3 mol-' B K 25 "C 35 "C 25 "C 25°C 35°C Et,NBr Pr,NBr Bu,NBr Pe,NBr Hex,NBr Hept,NBr Bu,NBu,B Bu,NI NaI Ph,PBr Ph,PI NaPh,B Bu,NPh,B 0.650 f 0.005 0.706 f 0.001 0.839 f 0.001 0.984 f 0.008 1.100 f 0.004 1.309 f 0.008 1.002 f 0.003 0.821 f 0.001 0.73 1 f 0.003 1.104 f 0.00 1 1.088 f 0.001 1.240 f 0.002 0.640 f 0.014 0.704 f 0.005 0.829 f 0.002 0.952 f 0.007 1.091 f 0.008 1.313f0.008 0.982 f 0.002 0.8 1 5 f 0.00 1 0.699 f 0.005 1.085 f 0.002 1.072 f 0.002 1.237 f 0.006 0.69" 0.71" 0.93" 276f 10 268f 11 574f4 578f4 0.87" 283f 10 327f 18 293f12 303f4 1.26 f 0.02b 262 f 8 254 f 27 1.35" and 1.32 f 0.02b a Ref.(8). and 95 % confidence limits on the fitted apparent molar volumes. Ref. (7). The error limits for this work are standard errors on the fitted B coefficients, Table 5. Ionic viscosity B/dm3 mol-' values in acetonitrile at 25 "C this work Criss" Krumgalzb Gill" Na+ Et,N+ Pr,N+ Bu,N+ Pen,N+ Hex,N+ Hept,N+ Ph,P+ Br- I- Bu,B- Ph,B- 0.453 f 0.004 0.44 0.33 0.32 0.39 0.39 0.37 0.48 0.500+ 0.002 0.56 0.62 0.59 f 0.06 0.67 0.78 0.99 0.8 10 f 0.003 0.32 f 0.03 0.37 0.31 0.30 f 0.03 0.34 0.25 0.502 f 0.002 0.787 f 0.003 0.87 0.735 0.73 f 0.06 a Ref.(14). Ref. (13). " Ref. (7). splitting viscosity B coefficients into ionic contributions, so we did not include it in our measurements, but a value may be calculated assuming the principle of additivity of B coefficients from our values for Bu,NI, NaPh,B and NaI. From these a value of 1.33 is obtained, which is in close agreement with that of Gill.' With the exception of Pr,NBr, the coefficients of FUOSS~ appear to be rather high. When eqn (2) is used to calculate ionic B values from Bu,NBu,B, Bu,NBr and Bu,NI, values of 0.339 for Br- and 0.321 for I- are obtained, whereas the corresponding bromide and iodide values obtained from the tetraphenyl salts are 0.294 and 0.278, each with a standard error of k0.003 dm3 mol-l.This is the first aprotic solvent that we haveK . G. Lawrence, A . Sacco, A. De Giglio and A . DelVAtti 29 studied for which there are notable differences between the halide ion B values from the two reference salt systems. These differences may arise from specific interactions between solvent molecules and the Ph,P+ and Ph,B- ions as proposed by Coetzee and Sharpeg for a number of nonaqueous solvents including ACN. Their spectroscopic results led them to suggest that anions interact directly with the methyl hydrogens of ACN, and that the anion bonding was weaker than for the cations because of the nature of the charge distribution on the solvent molecule.'o However, no firm conclusion was reached other than that ACN discriminated between Ph,P+ and Ph,B- and that the tetra-alkyl- substituted reference salt should be preferable to its tetraphenyl counterpart.Gas-phase studies of the solvation of alkali-metal and halide ions by acetonitrile also indicated unequal solvation of cations and anions.ll Our division into ionic values assumes the validity of eqn (2) and an analogous equation for the tetraphenyl salt system. When we calculate B(Ph,P+) it is clear that if we allow an additional contribution to V,(Ph,P+) on the grounds of unequal cation-solvent and anion-solvent interactions, the resulting B values for the halide ions derived from the tetraphenyl salt will be even smaller than those reported above. Examination of the single-ion conductances in ACN obtained from transference- number measurements by Kay et a[.'' shows that there would be no improvement if we chose volume ratios based on Stokes radii for the division into ionic B values, as used by Krurngal~'~ and Gill.' Thus although one might look more favourably upon the values derived from the tetra-alkyl salt, there does not seem to be any convincing reason for choosing one set of values in preference to the other, so we have reported the mean values for the halide ions from both reference salts in table 5.The error limits shown for the reference ions are calculated from the standard errors on the fitted B coefficients for these salts, whereas for the halide ions we show the resulting standard deviation from the mean of the pairs. The derived ionic values shown in table 5 are therefore only reported to two decimal places.The B values for Bu,N+ and Bu,B- are considerably smaller than the values for Ph,P+ and Ph,B-, and table 4 shows that the partial molar volume of Bu,NBr is of comparable magnitude to the corresponding tetraphenyl salt. Similar results have been seen in all of our studies with these reference salts, and have been explained in terms of the differing flow patterns around the long-chain tetra-alkyl ions and the propeller-like tetraphenyl ions. Also shown in table 5 are Criss's ionic B values,', which must be regarded as speculative, since he used the experimental results of FUOSS' and chose, without giving a reason, a value of 0.25 for the Me,N+ ion. Krumgalz also used Fuoss's values. We have previously tested another method of obtaining ionic B values that involves measuring the B coefficients for an homologous series of tetra-alkylammonium salts with a common anion, and plotting the B coefficients of these salts as a function of the formula weight or van der Waals volumes of the ~ a t i 0 n s .l ~ Extrapolation of the independent variable, e.g. to V,(R,N+) = 0, should give the B(X-) as intercept. This method assumes that the B coefficients are a linear function of this variable, and the contribution of the chosen cationic variable to the B value should vanish at zero. Fig. 1 shows that the assumption of linearity of B coefficients with V, is true only for tetrapropyl to tetrahexyl ; the circles approximately correspond to 95 % confidence limits.Extrapolation for these four salts produced the value at V,(R,N+) = 0 of B(Br-) = 0.26f0.02 at 25 OC, and this does not compare well with the reference-salt value. A similar situation was found for both DMSO' and HMPT2 solutions published previously. Krumgalzl' also tried extrapolating B coefficients against the cube of Stokes radii for the tetra-alkylammonium ions; the radii were values averaged from literature sources of ionic conductances. Using his radii a value of B(Br-) = 0.48f0.03 was obtained. This further confirms our view that extrapolation methods should be viewed with caution.30 Ionic Contributions to the Viscosity B Coeflcients 0 Et,N+ Pr,N* Bu,N+ Pe4N+ Hex4N+ Hept,N+ 100 150 200 2 50 300 v, /an3 mol- I I I I I Fig. 1. Plot of B(R,NBr) us. Vw(R,N+) at 25 "C.Transition-state Treatment The detail of the viscous flow process as it pertains to B coefficients has recently been re- examined in terms of the transition-state treatment.17 This new theory suggests that the magnitude of the ionic molar contribution to the activation energy, A&*, depends only on differences in ion-solvent interactions between the ground and transition states. The movement of an ion into the transition state may be considered to involve the breaking of ion-solvent bonds to create a cavity ahead of the moving ion, and the re-making of ion-solvent bonds in the transition state. For highly structured solvents the coordination of the ion in the ground state may be incomplete, but the viscous flow process disrupts more of the solvent structure, so in the transition state the coordination of the ion is likely to increase.For less well structured solvents an ion's ground-state coordination may even diminish in the transition state. Consequent changes that may occur in the solvent-solvent interactions are reflected in the enthalpies and entropies of activation only, and do not contribute to Api*. We have calculated the activation parameters for the salts involved in this work and divided them into cationic and anionic contributions using the same technique as that used to divide the B coefficients. These are shown in table 6, along with values calculated from the work presented in previous papers for comparison. Values for the bromide ion in the various solvents are the mean values obtained from Bu,NBr and Ph,PBr.Examination of this table reveals that the results for the tetrabutyl and tetraphenyl ions in ACN are unusual; the TAS,O* values are negative, their AH,"* values are smaller than in the other solvents, and their A&* values are considerably larger than the corresponding AH;* values. Of the three solvents DMSO, HMPT and DMF, DMSO is believed to possess the most structure through dipoledipole interactions to form small chains ; nevertheless the molecular association is probably not strong enough to prevent complete coordination of an ion by DMSO molecules in the ground state, so in the transition state a reduction in the coordination is most likely. Thus ion-solvent and solvent-solvent bond-breaking contributes to the large positive values of the enthalpies and entropies of activation for the reference ions in DMSO.HMPT and DMF are less structured than DMSO, and this results in a similar pattern but smaller numerical values for AH;* and TAS,"* for these ions.Table 6. Ionic thermodynamic activation parameters for viscous flow at 25 "C AH,"+/kJ mol-' DMSO" Bu,N+ Ph,P+ Bu,B- Ph,B- Br- I- Na' solvents 42.5 78.3 43 .O 76.1 44.6 39.1 32.9 13.7 HMPTb DMF' 38.3 38.7 65.5 42.5 38.7 39.2 64.1 41.3 32.5 29.8 - - - - 14.4 8.3 ACNd DMSO HMPT DMF 31.6 29.6" 31.8 28.6" 29.5 32.0 37.9 8.8 11.6 32.2 11.9 31.3 21.2 16.6 14.3 - 0.9 12.5 6.3 12.8 6.6 26.5 - 1.5 14.3 12.0 26.8 - 1.5 -3.6 -4.2 ACN - 8.9 - 25.0' - 8.9 - 24.5" 11.2 14.6 12.5 - 0.6 DMSO 30.9 46.1 31.1 44.8 23.4 22.8 18.5 14.57 Api*/kJ mol-' HMPT DMF 25.8 32.4 38.7 44.0 25.9 32.6 37.6 42.8 18.2 17.8 - ~ - - - 17.99 12.49 ACN 40.5 54.6e 40.7 53.1" 18.3 17.4 24.5 9.44 a , b Table 7, ref. (2). Calculated from ref. (3). This work. " From Ph,PI +NaPh,B-NaI. b b b cp w c)32 Ionic Contributions to the Viscosity B Coeficients ACN molecules have a strong tendency to form dimers by association.l' The dissociation energy of the dimers is estimated to be 16-30 kJ mol-l, with 96 YO of the CN groups associating in pairs at 30 "C. For a given temperature the degree of solvation of an ion in the ground state will depend on whether the ion-solvent interactions are strong enough to open the solvent dimers. In the transition state there will be more molecules freed from dimer association by the viscous flow process allowing an increase in the solvation of the ions.The concomitant bond-making results in negative entropies of activation, and a negative contribution to the enthalpies of activation that lowers their numerical values (see table 6). As mentioned above, changes in the ion-solvent interactions from ground to transition state are effectively measured by A&*, and we see that the largest values occur for the reference ions in ACN. NaI in DMSO and ACN makes an interesting comparison with the reference salts. A strong interaction between Na+ and I- ions and solvent molecules will occur in the ground state, and formation of the transition state will be accompanied by some ion-solvent bond-breaking (?!-'AS: * positive) and a reduction in the solvation of these ions in both solvents.The Influence of Temperature Table 4 can be used to show the change in the B coefficients with rise in temperature, AB/A T. A comparison with results reported for the salts in solvent systems previously studied in this series of papers shows that the values for ABIAT for ACN are unusually small. A full discussion of this aspect of the work will be reported in a subsequent paper. References 1 K. G. Lawrence and A. Sacco, J. Chem. Soc., Faraday Trans. 1, 1983, 79, 615. 2 A. Sacco, M. D. Monica, A. De Giblio and K. G. Lawrence, J. Chem. SOC., Faraday Trans. I , 1983,79, 3 A. Sacco, A. De Giglio, A. Dell'Atti and K. G. Lawrence, 2. Phys. Chem. N.F., 1983, 136, 145. 4 G. P. Cunningham, G. A. Vidulich and R. L. Kay, J. Chem. Eng. Data, 1967, 12, 336. 5 G. J. Janz and R. P. T. Tomkins, Nonaqueous Electrolytes Handbook (Academic Press, New York, 1972); A. K. Covington and T. Dickinson, Physical Chemistry of Organic Solvent Systems (Plenum Press, London, 1973). 6 J. Crudden, G. M. Delaney, D. Feakins, P. J. O'Reilly, W. E. Waghorne and K. G. Lawrence, J. Chem. SOC., Faraday Trans. 1, 1986, 82, 2195. 7 D. S. Gill, M. S. Chauhan and M. B. Sekhri, J. Chem. SOC., Faraday Trans. 1, 1982, 78, 3461. 8 D. F. Tuan and R. M. Fuoss, J. Phys. Chem., 1963, 67, 1347. 9 J. F. Coetzee and W. R. Sharpe, J. Phys. Chem., 1971, 75, 3141. 10 J. A. Pople and M. Gordon, J. Am. Chem. SOC., 1967, 89,4253. 11 W. R. Davidson and P. Kebarle, J. Am. Chem. SOC., 1976, 98, 6125. 12 C. H. Springer, J. F. Coetzee and R. L. Kay, J. Phys. Chem., 1969, 73,471. 13 B. S. Krumgalz, Russ. J. Phys. Chem., 1973,47, 956. 14 C. M. Criss and M. J. Mastroianni, J. Phys. Chem., 1971, 75, 2532. 15 K. G. Lawrence, R. T. M. Bicknell, A. Sacco and A. Dell'Atti, J. Chem. Soc., Faraday Trans. 1, 1985, 16 B. S. Krumgalz, J. Chem. SOC., Faraday Trans. I , 1980, 76, 1275. 17 D. Feakins, W. E. Waghorne and K. G. Lawrence, J. Chem. SOC., Faraday Trans. I , 1986, 82, 563. 18 A. M. Saum, J. Polym. Sci., 1960, 42, 57. 263 1. 81, 1133. Paper 8/0022OG; Received 18th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500023

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Cheni. Soc., Farackij. Truns. I , 1989. 8 3 1). 3 3 4 5 Temperature-programmed Desorption of Ethanol from ZSM-5, ZSM-11 and Theta-1 Chen Li-feng,? Thomas Wacker$ and Lovat V. C. Rees* Department of Chemistry, Impwiirl c'olltyc of Science und Technologj., Lomion SW7 2'4 Y Ethanol was adsorbed on low Si/AI ratio ZSM-5. ZSM-I I and Theta-l at room temperature and relative prcssure p/p(, = 0.5. in situ. in a thermal gravimetric balance. The saturation capacity of ZSM-I 1 was found to be slightly greater than ZSM-5 and cci. twice as large as that of Theta-I. Temperature-programmed desorption profiles of ethanol were determined and analysed by ( u ) a variable heaLing rate and ( h ) single heating rate method using an on-line computer. The effect of cation exchange on the sorption capacities and activation cnergieq of desorption have also been determined. In previous studies the temperature-programmed desorption (t.p.d.1 of non-polar hydrocarbons from low silica-alumina ratio ZSM-5, ZSM-I 1 and Theta-l have been reported.'.' In the present paper the t.p.d.of the polar molecule, ethanol, from the same zeolites is studied. Comparison with the former results shows that the electrostatic interaction between the permanent dipole moment of ethanol and the zeolite channel surface is strong. This electrostatic interaction tends to overcome the steric hindrance of the cations present in these zeolite channels. This hindrance is quite important in the packing of hydrocarbon molecules. '. The t.p.d. of ethanol from the above zeolites in their H' and Na' forms demonstrates that increasing the Na content decreases the amount of ethanol adsorbed slightly and increases the activation energy of desorption. As before,'.' two methods of analysing the t.p.d.profiles have been used, i.e. the variable heating rate (v.h.r.) and the single heating rate (s.h.r.) methods. Experimental Details of the thermogravimetric balance are given in ref. (1) and (3). The unit-cell formulae of the low Si/AI ratio zeolites used in this study are as follows: ZSM-S(§i/Al = 15) : H, ,,Na, 52A16Si900192 ZSM- 1 1 (Si/Al = 15) : H, ,,Na, 64Al,Si,,,01,2 Theta- I (Si/AI = 32) : H,, ,,Na, ,,Al, 91Si93,09019z. The sodium-exchanged forms of these three zeolites were prepared by ion exchange with 0.1 mol dm-3 NaCl solution under a controlled pH of 10.The sodium contents of these samples were determined by neutron activation and an Na/A1 ratio of 1 was found, within experimental error, for all three samples. The hydrogen form of Theta-l was prepared by ion exchange with 0.1 mol dm '' NH,Cl at pH 5 followed by calcination at 500 "C. The unit-cell formula of this sample was found to be Na,,,H,,,Al,,,Si,,,,,O,,,. i- Present address : Chemical Engineering Department. Nanchang Institute of Aeronautical Technology, No. 2 Shanghai Road, Nanchang, Jiangxi. The People's Republic of China. Present address : Weidhoelzliweg 19, 5024 Knettigen. Switzerland. 33 2-234 T.P.D. of Ethanol from Zeolites Table 1. Comparison of the amount of ethanol adsorbed under various saturation conditions experimental conditions ZSM-5 ZSM- 1 1 Theta- 1 p / p , z 0.5; room 13.68 14.12 6.42 temperature ; sorption time, 30 minu temperature ; sorption time, 4 hb 1 h at 50 "C then 3 h at room temperatureb p / p , x 0.5; room 13.62 14.36 6.51 p / p , x 0.5; sorption for 13.73 14.18 6.30 The saturation capacities are given as molecules per unit cell.a Average value from the 10 different heating rate experiments. Determined from 10 K min-' run. The sodium content of this sample was determined after dissolution of the sample by atomic absorption spectroscopy. Absolute alcohol (A. R. quality, > 99.7%) was used. The reproducibility of the t.p.d. profiles was very good. Results and Discussion Various, in situ experimental conditions were tested to ensure complete saturation of the sample with adsorbate prior to starting the t.p.d.runs. The results of these preliminary experiments are given in table 1. Table 1 shows that the amount of ethanol sorbed in ZSM-11 is slightly greater than that in ZSM-5 and much greater than that in Theta-1. The sorption of various hydrocarbons was found previously' to follow the same pattern. From the saturation amounts sorbed and a density of 0.7893 g cm-3 for liquid ethanol, micropore volumes, W,, of 0.138, 0.142 and 0.064 cm3 g-l for ZSM-5, ZSM-11 and Theta-1, respectively, were calculated. These sorption volumes are much lower than the theoretical volumes of the channels of these zeolites. The micropore volumes calculated from the saturation capacities of five different adsorbate molecules are compared in table 2.The W, value of ethanol in ZSM-5 is close to the corresponding value in silicalite. The low saturation sorption capacity of ethanol in silicalite was ascribed to the hydrophobic character of the channel s ~ r f a c e . ~ The surface of the low Si/A1 ratio ZSM-5 sample used in the present study should be less hydrophobic than that of silicalite and the low sorption capacity of ethanol in this sample must reflect, therefore, some steric hindrance of the cations present in the channels. However, the larger sorption capacities of n- hexane and n-octane over ethanol, in table 2 for the ZSM-5 sample indicates that the surface of these low Si/Al ratio zeolites is still dominantly hydrophobic in character. Ethanol T.P.D. Studies The t.p.d. profiles of ethanol from ZSM-5 (Si/Al = 15), ZSM-11 (Si/Al = 15) and Theta-1 (Si/Al = 32) are shown in fig.1-3 along with the differential of these profiles.C/tm LiTfeng, T. Wuckc~ cind L. P'. C. Reps 35 n-hexaneb n-octaneb p-x ylene' benzenec ethanol 10 % > 0.5 E a Table 2. y, (cm3 g') obtained for different adsorbate molecules ~~~ ~~ - ~~ ZSM-5 ZSM- 1 1 Theta- 1 (SI/AI = 15) (SI/AI = 15) (Si/AI = 32) sillcall tea -~ ~ 0. I50 0 156 0.060 0.173 0.143 0.147 0.055 0.149 0.128 0.133 0.042 0 181 0.108 0.1 10 0.042 - 0.138 0.143 0.064 0.136 I' Ref. (3). ', Kcf. ( I ). Ref. (2). I 0 I 3 0 II] U 14.0 10.5 ..a .- c b a 7.0 2 - ; 8 3.5 0 300 350 400 450 500 550 600 650 TIK Fig. 1. T.p.d. profile and differential of profile of ethanol in (NaH)ZSM-5. /3 = 10 K min-'. 1.0 6 2 05 a 0 0 -0 \ 0 0 0 a U T- !3 14.0 I - 10.5 .C c b a 7.0 3 3 8 I 2 3.5 0 300 350 400 450 500 550 600 650 T/K Fig.2. T.p.d. profile and differential of profile of ethanol in (NaH)ZSM- 11. j3 = 10 K min-'.36 T.P.D. of Ethanol from Zeolites 1.0 9 3 .. Q5 0 0 -0 I 0 0 0 0 0 0 n o D U . I I 300 350 400 450 500 550 600 650 T/K Fig. 3. T.p.d. profile and differential of profile of ethanol in (NaH)Theta-1. /? = 10 K min-'. Table 3. Peak temperatures and peak widths (B = 10 K min-') peak temperature/K peak width/# major peak I peak I1 peak peak I peak I1 diffuse 425490 ZSM-5 diffuse 465 peak I1 ZSM- 1 1 320 470 peak I1 300-380 430-490 Theta- 1 320 510 peak I 30&380 48&550 The peak temperatures and peak widths are listed in table 3. Note that the low- temperature peak at 320 K found with ZSM-11 and Theta-1 is absent or very diffuse with ZSM-5.This result suggests that the sorption sites in ZSM-5 from which the first seven molecules of ethanol are desorbed have gradually increasing energies. ZSM- 1 1 and Theta-1 both have sets of low-energy sites from which the first seven and two molecules, respectively, are easily desorbed at low temperature (i.e. ca. 320 K). Comparison of the peak temperature of ZSM-5 of 465 K with that of silicalite3 of 340 K indicates that the electrostatic interaction between the permanent dipole moment of ethanol and the sites in ZSM-5 (Si/Al = 15) with large electric fields is obviously strong. A small high-temperature peak was found above 625 K with ZSM-5 when heating rates of 2 or 4 K min-' were used. This small peak must be the result of some catalytic reaction of ethanol molecules strongly associated with framework aluminiums or acid sites. At higher heating rates the small peak is not observed, probably because the catalytic reaction is spread out over a larger temperature The second peak observed with Theta-1 occurs at a higher temperature than that found with ZSM-5 and ZSM-11, probably due to the higher Na content of Theta-1 and the lack of cavities in the one-dimensional channel structure.The electrostatic interaction between the permanent dipole moment of ethanol and the walls of the channels of Theta-1 must be stronger than that with ZSM-5 and ZSM-11.1LO 120 - 100 d 8 c1 \ 4" 80 60 40 37 0 2 i 6 a 10 12 14 molecules per unit cell Fig. 4. Activation energy of desorption of ethanol as a function of coverage.0. ZSM-5: 0, ZSM-I I : 0. Theta-1 in their (haH)-forms. V.h.r. method of analysis. Analjtsis of T.P. D. Profiles by the V.H. R. Mcthod Heating rates at 2 K intervals between 2 and 20 K min-' were used as before.' ' The activation energies, E(I, and entropies, -AS:, of desorption as a function of coverage derived from the v.h.r. method of analysis:' are shown in fig. 4 and 5 , respectively. The activation energy, averaged over the whole coverage range, Eli, for the three zeolites is given in table 4. All three zeolites show high E,, values at low coverages, consistent with the presence of sites with high electric fields in these low Si/Al samples. The variation of El, with coverage is similar for both ZSM-5 and ZSM- 1 1.as would be expected for two zeolites with very similar channel structures and the same Si/AI ratio. The variation of E,, with coverage for ZSM-5 is very similar in form to that found for the differential molar heats of sorption, (2, of ethanol in silicalite - 1 ;' but Et1 is cu. 30 kJ molpl greater than Q at all corresponding sorbate loadings. In both samples 12 ethanol molecules per unit cell out of the total sorption capacity of 14 molecules are desorbed with near-constant Eli values of ca. 8&90 kJ mol-'. Although ethanol has been reported to be dehydrated in H-forms of ZSM-5 at quite low temperatures, there is no clear evidence in these desorption energy distributions of any catalytic activity. Fig. 1 shows that the desorption of these 12 ethanol molecules from ZSM-5 cover a temperature span of 25-180 "C.i.0. the last of these ethanol molecules is being desorbed at 180 "C where some catalytic activity would be expected. However, Ed remains sensibly constant over this coverage range and similar in form to the variation of Q in silicalite- I , where no catalytic activity would be expected.38 T.P.D. of Ethanol from Zeolites 200 160 - k 120 - 2 h \ t m I a 80 LO 0 0 2 4 6 8 10 12 1L molecules per unit cell Fig. 5. Activation entropy of desorption of ethanol as a function of coverage. 0, ZSM-5; 0, ZSM-11; 0, Theta-1 in their (NaH)-forms. V.h.r. method of analysis. Table 4. Average activation energies of desorption (in kJ mol-') of ethanol by the v.h.r. and s.h.r. methods of analysis ZSM-5 ZSM- 1 1 Theta- 1 Ed (v.h.r.) 89 88 72 Ed (s.h.r.) 85 84 96 It would seem, therefore, that these Ed energies do represent desorption energies and not activation energies for the dehydration reaction.These conclusions should be confirmed by analysis of the effluent gas from the t.g. system but, unfortunately, we do not have a mass-spectrometer attachment on the back-end of the system and could not, therefore, analyse the effluent gases. Although the first seven ethanol molecules to be desorbed from ZSM-5 require slightly higher activation energies of desorption compared with the desorption of these molecules from ZSM- 1 1, there is no obvious reason for these differences observed in the desorption profiles shown in fig. 1 and 2 and table 3 for these two samples in this sorption range. The desorption of the first two ethanol molecules from Theta- 1 involves quite a distinct maximum in Ed at 5 molecules per unit cell which can account for the sharp low temperature peak in fig.3. The smaller E d value for Theta-1 in table 4 and the much lower Ed values in fig. 4 at120 - 100 L 8 2 y" 80 60 Chen Li-feng, T. Wacker and L. V. C. Rees 39 i 0 2 1: 6 8 10 12 14 molecules per unit cell Fig. 6. Activation energy of desorption of ethanol as a function of coverage. 0, ZSM-5; 0, ZSM-11; 0, Theta-1 in their (NaH)-forms. S.h.r. method of analysis. higher coverages compared with the corresponding values for ZSM-5 and ZSM-11 are consistent with the higher Si/A1 ratio of Theta-1 and are comparable to the differential molar heats, Q, found for ethanol in silicalite-1.5 Analysis of T.P.D.Profiles by the S.H.R. Method A heating rate of 6 K min-l was chosen. The variations of Ed and -AS$ with coverage derived from the s.h.r. method of analysis are presented in fig. 6 and 7, respectively. The average activation energies, Ed, for the three zeolites are given in table 4. The Ed values for all three zeolites at low coverages are quite large at ca. 110 kJ mol-1 for ZSM-5 and ZSM-11 and ca. 130 kJ mo1-l for Theta-1. These low-coverage E d values are similar to those in fig. 4. However, the s.h.r. method of analysis produces a h e a r decrease in Ed with increasing coverage to a final value of ca. 60 kJ mo1-l at saturation coverage for all three zeolites. There is no obvious strong interaction of ethanol with the limited number of sites of high electric field in these samples.The larger Ed value for Theta-1 in table 4 than that found for ZSM-5 and ZSM-11 is the opposite of the behaviour found using the v.h.r. method of analysis in table 4. This high E d value for Theta-1 is not consistent with the higher Si/Al ratio of this sample. The high E d value in Theta-1 could arise because of the one-dimensional channel network and the Naf ions acting as barriers in these channels. However, this explanation introduces the concept of a diffusion-controlled desorption process, which is unlikely since the Ed values are much greater than the activation energies of diffusion normally found in 10- ring channels. The average activation energy of desorption, Ed, was calculated using the s.h.r.method of analysis on each of the 10 t.p.d. profiles needed for the v.h.r. method of analysis. Each profile produced very consistent E d values which varied by < f 2 kJ mol-' for ZSM-5 and ZSM-11 and < f 3 kJ mol-' for Theta-1 from the Ed values listed in40 T.P.D. of Ethanol from Zeolites 160 - k - b 120 h \ t w I a 80 Table 5. Comparison of Ed values (kJ mol-') for ZSM-5 and silicalite adsorbate adsorbent n-hexane n-octane p-xylene benzene ethanol ZSM-5 85 92" 85b 8gb 85 silicalite ca. 73" ca. 86" ca. 77" ca. 70d ca. 53" (Si/A1 = 15) a Ref. (1). Ref. (2). " Ref. (3). Ref. (6). table 4 for the 6Kmin-l profile. Thus the s.h.r. method of analysis seems to give consistent Ed values reasonably independent of the heating rate employed at least over a heating range of 2-20 K min-'.The average desorption energies of ethanol from ZSM-5 and silicalite are compared in table 5. The large difference in Ed between these two samples is the result of the additional electrostatic energy introduced by the sites with high electric fields which exist in the ZSM-5 sample. The value of Ed of ca. 53 kJ mol-' for silicalite is.in very good agreement with the differential molar heat of sorption of ethanol in silicalite5 of ca. 57 kJ mo1-l. The difference in & for the non-polar hydrocarbon molecules listed in table 5 for these two samples is much smaller than that for ethanol, demonstrating that the high electric fields in ZSM-5 have a smaller effect on these energies when the desorbing molecule has no dipole moment.The desorption entropies, ASS, in fig. 4 and 7 for both ZSM-5 and ZSM-11 are ca. 80-90 J mol-' K-'. The v.h.r. and s.h.r. methods of analysis, therefore, give similar desorption entropies, but there are subtle differences in these entropies as a function of coverage between the two methods. The corresponding curves for Theta-1 in fig. 4 and 7 are quite different.Chen Li-feng, T. Wacker and L. V. C. Rees 41 1.0 k 3 .? 0.5 0 0 a 300 350 400 450 500 550 600 650 T / K Fig. 8. T.p.d. differential profile of ethanol in NaTheta- 1 . b = 10 K min-’ 1.0 h ? 0.5 3 0 0 0 a a O R 0 0 0 s o n I 0 0 0 300 350 400 450 500 550 600 650 T/K Fig. 9. T.p.d. differential profile of ethanol in HTheta-I. /? = 10 K min-’. T.P. D. of Ethanol from Diflerent Cationic Forms of High-silica Zeolites The t.p.d.differential profiles of ethanol from the sodium and hydrogen forms of Theta- 1 (Si/Al = 32) and the sodium forms of ZSM-5 and ZSM-11 are shown in fig. 8-1 1, respectively. These are the same zeolites which were used in the first part of this study in their mixed (NaH)-forms. The peak temperatures and widths and the maximum rates of desorption of ethanol from these samples are listed in table 6 . The differential dm’/dTcurves in fig. 8, 3 and 9 for ethanol desorbing from the Na-, (NaH)- and H-forms of Theta- 1, respectively, are very similar iq nature. Table 6 shows that the Na- and (NaH)-forms have similar temperatures for peaks I and 11, with the H- form having slightly lower peak temperatures. The maximum rate of desorption from the H-form is greater than the rate of the other two cationic forms. These results suggest that the desorption of ethanol from Theta- 1 involves larger interaction energies when Na ions are present in the one-dimensional channel system.42 T.P.D.of Ethanol from Zeolites Table 6. Peak temperature, peak width and maximum rate of ethanol desorption (,9 = 10 K min-') peak temperature/K ~~ peak I peak I1 larger peak max. desorption rate/ lo-" mg s-' NaTheta- 1 (NaH)Theta- 1 HTheta- I NaZSM-5 (NaH)ZSM-5 NaZSM- 1 1 (NaH)ZSM- 1 1 320 510 320 510 315 500 320 475 diffuse 465 320 475 320 470 I I I I1 I1 I1 I1 1.685 1.809 2.235 2.279 4.13 2.828 3.774 peak width/K peak I peak I1 300-380 465-550 300-380 480-550 _ _ _ _ ~ . . - 300-375 475-525 300400 435-500 diffuse 425490 300-380 430-490 300400 435-540 Table 7.Ethanol saturation capacities and Ed values obtained by the s.h.r. method @? = 10 K min-') Ed amount adsorbed /molecules per unit cell /kJ mol-' NaTheta- 1 (NaH)Theta- I HTheta- 1 NaZSM-5 (NaH)ZSM- 5 NaZSM- I 1 (NaH)ZSM- 1 1 5.98 99 6.25 96 6.50 89 13.21 98 13.60 88 13.93 94 14.27 84 1.0 .$ 0.5 0 a 0 a 0 0 O k 0 U 0 0 a 0 0 a 300 350 400 450 500 550 600 650 TIK Fig. 10. T.p.d. differential profile of ethanol in NaZSM-5. /I = 10 K min-'. The differential curves in fig. 1 and 10 show the effect on the ethanol desorption of additional Na+ ions in the ZSM-5 channel system. The low-temperature peak I is now quite distinct in the pure Na ZSM-5, while peak I1 occurs at a higher temperature with a much broader peak width than that found with the (NaH)-form.Similar behaviour forChen Li-feng, T. Wacker and L. V. C . Rees 1.0 0 8 z OS5 0 300 350 400 450 500 5% 600 650 TIK Fig. 11. T.p.d. differential profile of ethanol in NaZSM-I 1. p = 10 K min-'. 140 120 - L 8 2 100 4" \ 80 60 43 0 1 2 3 4 5 6 7 molecules per unit cell Fig. 12. Activation energy of desorption of ethanol as a function of coverage. 0, (NaH)Theta-1 ; 0, HTheta-1. S.h.r. method of analysis. 0, NaTheta- 1 ; peak I1 is found with the Na- and (NaH)-forms of ZSM-11. Finally, the rate of desorption of ethanol is lower in the pure Na-forms of these two zeolites. The amount of ethanol adsorbed at saturation and the average activation energy of desorption, Ed, calculated using the s.h.r. method of analysis of the profiles in fig. 8-1 1 are listed in table 7.In all of these samples increasing concentrations of Na ions decrease the saturation sorption capacity. The decrease is not very large and may only represent the additional volume of the Na+ ions over that of the replaced H+ ions.44 T.P.D. of Ethanol from Zeolites 140 120 @ k 2 4 100 80 60 0 2 4 6 8 10 12 14 molecules per unit cell Fig. 13. Activation energy of desorption of ethanol as a function of coverage. 0, NaZSM-5; 0, (NaH)ZSM-5. S.h.r. method of analysis. 1 LO 120 @ \ 2 100 4 80 60 0 2 L 6 8 10 12 l h molecules per unit cell Fig. 14. Activation energy of desorption of ethanol as a function of coverage. 0, NaZSM-11; 0, (NaH)ZSM- 1 1. S.h.r. method of analysis.Chen Li-feng, T. Wacker and L. V . C . Rees 45 The average desorption energies also clearly demonstrate the increasing sorb- ate-sorbent interaction energies introduced by the Na+ ions in the channels of these three zeolites.In fig. 12 the activation energy of desorption of ethanol determined by the s.h.r. method of analysis for the three cationic forms of Theta-1 is presented as a function of coverage. The three cationic forms all show very similar behaviour, with the H-form demonstrating slightly smaller E d values than the other two forms throughout the coverage range. In fig. 13 and 14 the corresponding Ed us. coverage curves are given for the two cationic forms of ZSM-5 and ZSM-1 I , respectively. Both of these figures clearly demonstrate the higher interaction energies between ethanol molecules and the channel surfaces at low loadings when additional Na+ ions are present in the channels.The desorption energies, however, are very similar for the respective cationic forms of ZSM- 5 and ZSM-11. At saturation loadings the desorption energy for both zeolites and both cationic forms is ca. 60 kJ mo1-l and not dissimilar from the E d value given in table 5 for silicalite. However, this desorption energy in the case of silicalite occurs over a much wider coverage range.3 These comparable desorption energies must represent the energy of desorption of ethanol from a silica site far removed from sites with large electric fields. Conclusions The t.p.d. of a polar molecule such as ethanol on low Si/Al ratio ZSM-5. ZSM-1 I and Theta-1 and different Na-, (NaH)- and H-forms of these zeolites shows that as the Na cation content of the zeolite increases the amount of ethanol adsorbed at saturation decreases and the activation energy for desorption increases.Both the saturation capacity and desorption energy demonstrate a subtle interplay between the electrostatic interaction of the dipole moment of the polar ethanol molecule and the electric field in the zeolite channels around an Al centre and the steric inhibitions created by the cations associated with such A1 centres. The variation in Ed with coverage for ethanol desorbing from ZSM-5 is very similar in form to the variation of the differential molar heat of adsorption with coverage found with silicalite. However, a difference of ca. 30 kJ mol-1 exists at all coverages between these two curves. These results are consistent with Ed being a measure of the desorption energies for these various systems and there is no clear evidence of any catalytic activity in the H- forms of these samples at the lower desorption temperatures. However, the broader high-temperature peaks in the Na-ZSM-5 and Na-ZSM-11 in fig. 10 and 11, respectively, compared with the corresponding narrower, somewhat lower temperature peaks in fig. 1 and 2 could be due to catalysed reaction increasing the ease of desorption of ethanol molecules from the H-forms of these zeolite samples. The authors are indebted to Dr R. E. Richards and Dr E. A. Dima for the development of the computer programs used in this study. We also wish to thank British Petroleum Research Centre, Sunbury-on-Thames for the three zeolite samples and Mr A. M. McAleer for help in the characterisation of these samples. References 1 L. F. Chen and L. V. C. Rees, Zeolites, 1988, 8, 310. 2 L. F. Chen and L. V. C. Rees, Zeofites, in press. 3 R. E. Richards and L. V. C. Rees, Zeolites, 1986, 6, 17. 4 Y. Tokoro, M. Misono, T. Uchijima and Y. Yoneda, BUN. Chem. Soc. Jpn, 1978. 51, 85. 5 H. Thamm, J . Chem. Soc., Faraday Trans. I , 1989, 85, 1. 6 H. Thamm, J . Phys. Chem., 1987, 91, 8. Paper 8/00230D; Received 20th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500033

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. f, 1989, 85(1), 47-53 The Properties of Boralites studied by Infrared Spectroscopy Jerzy Datka* and Zofia Piwowarska Faculty of Chemistry, Jagiellonian University, Cracow, Poland Because of its small dimensions boron is three-coordinated in boralites and is situated in one face of a tetrahedral hole. The i.r. band of the B-0 asymmetric stretch in BO, units is split into two maxima (at 1380 and 1405 cm-'), thus suggesting that two kinds of boron sites exist in boralites. Boron becomes four-coordinated following adsorption of the electron- donor molecules NH,, H,O and pyridine, but not of benzene. Three kinds of OH groups (at 3680, 3720 and 3460cm-') are formed following the substitution of Na' ions by protons. The 3720cm-' OH groups act as Brernsted-acid sites in the reaction with pyridine.The concentration of Brsnsted sites determined by pyridine sorption is comparable with theoretical values. The low strength of the Bronsted-acid sites (illustrated by pyridine thermodesorption experiments) can be explained by the large separation and weak interaction between boron and the OH groups in the O,B***OH-SiO, units. The electrostatic field of the Lewis-acid sites in boralites is stronger than that in zeolites. There are numerous methods whereby the acidic and catalytic properties of zeolites may be modified. One is the isomorphic substitution of A1 or Si atoms by other elements such as B, Be, Ga, Cr, Fe, V, Ti, Ce, Zr or P. Isomorphically substituted zeolites have been extensively studied,'-14 although as yet only boron-substituted pentasil zeolites have any industrial applications [in the Assoreni (conversion of methyl butyl ether into methanol and isobutene) and Amoco (xylene isomerization and ethylbenzene conversion) processes].Apart from these industrial applications, boralites are also interesting from the chemical view point; there are two features which are characteristic of boralites and do not occur in zeolites: the presence of three-coordinated boron and the low acid strength of the OH groups. T atoms (A1 or Si) in the zeolite lattice are four-coordinated, but boron (being much smaller) is located near three oxygen atoms in a trigonal environment within a tetragonal hole.6. l1 The fourfold coordination of boron is observed only after the adsorption of H,O, CH,OH and NH,.Another interesting feature of boralites is that the acid strength of their OH groups is much lower than that in zeolites, despite the fact that the electronegativity of boron is higher than that of aluminium (2.93 and 2.22 in the Sanderson scale).6'11T12 The aim of our work was to study by i.r. spectroscopy the properties of boron atoms and OH groups in boralites. Na-Boralite, H-boralite and zeolite H-ZSM-5, all containing the same amount of boron and aluminium, were studied. Experimental The boralite samples were synthesized by Dr A. Cichocki (Department of Chemical Technology, Jagiellonian University). Zeolite ZSM-5 (Ultrazet) was synthesized at the Institute of Industrial Chemistry (Warsaw). X-Ray analysis has shown all the samples to be highly crystalline, while an m.a.s.n.m.r.study has shown all the boron atoms to be situated in the boralite 1atti~e.l~ The NH, form of boralite was obtained by ion exchange 4748 I.R. Study of Boralites Table 1. Compositions of the boralites and ZSM-5 zeolite, and the concentration and strength of the Brmsted-acid sites sample composition Brmsted-acid sites per unit cell Na-boralite Na1.0H0.1~B02~1.1~si02~94,9 0.27 - NH,-boralite (NH4)1.0(B02)1.0(si0'2)95 0.67 0.16 NH4-ZSM-5 (NH4,H)0.,Na0.2(A102)1.0(si02)95 0.73 0.65 with an NH,CI solution at room temperature. The compositions of the Na- and NH,- boralites, as well as that of zeolite NH,-ZSM-5, are presented in table 1. Prior to the i.r. experiments, wafers of the boralites and the zeolite were activated at 773 K in situ in the i.r. cell for 1 h.The i.r. spectra were recorded using a Specord 75 i.r. spectrometer working on-line with a KSR 4100 minocomputer. Results BO, Vibrations Spectra of the Na- and H-forms of boralite are presented in fig. 1. A doublet at 1380 and 1405 cm-l is present in the spectra of the activated boralites. Sorption of H,O, NH, or pyridine results in the disappearance of both the 1380 and 1405 cm-' BO, bands [fig. I A and B, spectra (bHd). These effects are reversible. Evacuation at higher temperatures (373-673 K) removes the sorbed molecules and the doublet at 1380 and 1405 cm-' reappears. Sorption of benzene does not influence the BO, vibrations [spectrum (e)]. OH Groups Spectra of OH groups in the Na- and H-forms of boralite, as well as spectra of OH groups in H-ZSM-5, are presented in fig.2A. The OH band at 3740 cm-l, characteristic of silanol Si-OH groups, is present in the spectrum of Na-boralite [fig. 2A, spectrum (a)]. The frequency and the intensity (expressed per 1 g of boralite) of the 3740 cm-' OH band in Na-boralite (3.65 g-l) are practically the same as in the Na-ZSM-5 zeolite (3.35 g-'). The spectrum of H-boralite (obtained by decomposition of the NH,-form) is presented in fig. 2A, spectrum (b). Introduction of protons into the boralite results in the appearance of new OH bands at 3680 and 3720 cm-' as well as a weak and broad band at 3460 cm-'. The band due to Si-OH silanol groups is seen as a shoulder at ca. 3740cm-'. The spectrum of H-ZSM-5 zeolite [fig. 2A, spectrum (c)] contains two distinct OH bands : 3738 cm-' (Si-OH) and 3609 cm-' (Si-OH-AI).Pyridine Adsorption Sorption of pyridine in H-boralite (fig. 2 B) results in a strong decrease (or disappearance) of the 3720 cm-' OH band and in a small diminution of the 3680 cm-' OH band. The 3740 cm-' Si-OH band remains unchanged. The sorption of pyridine results in the formation of pyridinium ions HPy+ (1 545 cm-'), pyridine complexes bonded to Na+ ions PyNa+ (1440 cm-l) as well as pyridine complexes bonded to Lewis-acid sites PyL (1460 cm-'). The same complexes are formed after pyridine sorption in H-ZSM-5. The frequency of the PyL complex formed in H-boralite (1460 cm-') is higher than that in H-ZSM-5 (1450 cm-l).J. Datka and Z . Piwowarska 49 rjoo ubo 1500 l600 wavenumber/cm-' Fig.1. Spectrum of Na-boralite (A) and H-boralite (B): (a) activated boralite, and after sorption of (6) H,O, (c) NH,, ( d ) pyridine and (e) benzene. f ' (b) (4 I 7 l I I l 3400 3600 3800 ' 3bO 3600 38a3 I400 1600 wavenumber/cm- ' w avenumtxr/cm- w avenumtxr/cm-' Fig. 2. (A) OH groups in (a) Na-boralite, (6) H-boralite and (c) H-ZSM-5. (B) OH groups in H-boralite : (a) the activated boralite and (b) after pyridine sorpton. (C) The spectrum of pyridine sorbed in (a) H-boralite and (6) H-ZSM-5.50 I.R. Study of Boralites In order to determine the concentration of Brsnsted-acid sites in boralites and in zeolite, small portions of pyridine were sorbed at 443 K up to a constant intensity of 1545 cm-l PyH+. The concentration of PyH+ (and therefore of Brransted-acid sites reacting with pyridine) was calculated from the maximum intensity of the 1545 cm-l band and from the extinction coefficient of this band determined in a previous study1fi (0.058 cm pmol-2).The calculated values are presented in table 1. The concentrations of Brsnsted-acid sites in Na-boralite, H-boralite and H-ZSM-5 are comparable with the theoretical concentrations of protons calculated from the difference between the contents of B(A1) and Na. In order to study the strength of the Brransted-acid sites, experiments with thermodesorption of pyridine were made. Pyridine was desorbed at 770 K, and the ratio A770/A0 was determined. ( A , and A7,,, are the intensities of the PyH+ 1545 cm-' band before and after desorption.) The values of A770/A0, giving information about the fraction of Brsnsted-acid sites still holding pyridine after desorption at 770 K, are taken as a measure of the acid strength of these sites and are presented in table 1.The strength of the Brsnsted-acid sites in the H-boralites is lower than that in H-ZSM-5. Discussion BO, Vibrations The boron atom is smaller than aluminium (atomic radii 0.82 and 1.18 A, respectively) and occurs preferably in trigonal coordination. In boralites boron is situated between three oxygens at one face of a tetrahedral hole in a planar, or nearly planar, configuration. This conclusion is based on the results of i.r. studies of Coudurier and Vedrine,ll who observed a strong band corresponding to a B-0 vibration in BO, units in dehydrated boralites, and also on the results of Scholle and Veemaq6 who detected a m.a.s.n.m.r.signal for three-coordinated boron. (The quadrupole parameters of this signal were similar to those in planar BO,.) The results obtained in our study are in good agreement with these earlier interpretations. A strong i.r. band of the asymmetric stretching BO, vibration is present in the spectrum of dehydrated Na-boralite and H-boralite [fig. 1 A, spectrum (a)]. This band is split into two maxima at 1380 and 1405 cm-l, thus suggesting that two kinds of BO, units with different B-0 force constants exist in boralites. The configuration of the boron atom in the boralite lattice is presented schematically as (1). The boron atom I 0' B I O \ 0 is situated between three oxygens : the distance to the fourth oxygen is greater, and the interaction is much weaker.The splitting of the BO, band suggests that two different sites exist for the boron atom in boralites. In one such site the distance to this fourth oxygen may be shorter than in the other, and thus the interaction with this oxygen may be stronger. Because the boron atom is shifted closer to this oxygen, its interactions with the three oxygens in the BO, unit are weaker and the B-0 stretching force constant in BO, is lower. Note that 12 crystallographically different 'T' sites have been found in the ZSM-5- type stru~ture.'~ M.a.s.n.m.r. studies1* have shown that the number of different T sites is even higher. The sorption of H20, NH, and pyridine in both Na- and H-boralites results in a reversible decrease in the intensity of the BO, doublet (1380 and 1405 cm-l) (fig.1 A and B). These effects were also observed by Coudurier and Vedrinell following theJ . Datka and Z . Piwowarska 51 adsorption of H20, CH,OH and NH,. Scholle and Veeman' reported the disappearance of the BO, m.a.s.n.m.r. signal and the appearance of a BO, signal upon the adsorption of H20 on H-boralite. The i.r. band of BO, vibrations (at 1150 cm-l) was also observed in our studies of hydrated b0ra1ites.l~ The results obtained in our study (and also the earlier results) prove that boron increases its coordination number (from three to four) after sorption of electron donors such as H,O, NH,, pyridine or CH,OH. An interaction with benzene (a n-electron donor) does not change the coordination number of boron (fig.1). Two possible interpretations of these effects can be considered. (1) The sorption of electron-donor molecules removes a proton or Na+ ion from the framework. Protons form PyH+ or NH; ions, and they can be also dissolved in H20 clusters.6 Na+ ions also form complexes with NH, or pyridine molecules. The abstraction of a proton or Na+ ion from the framework generates four-coordinated boron according to the scheme: O,B...OH(Na)-SiO, - 0,B-0-SO,. The interaction of OH groups with benzene (hydrogen bonding) does not remove the proton from the framework: boron remains three coordinated (fig. 1). (2) The second interpretation (which seems to be less probable) assumes a direct interaction between the boron atom and free electrons of the oxygen atom (in H,O) or nitrogen atom (in NH, or pyridine).This interpretation assumes that the geometry of the tetrahedral hole in which such a boron atom is situated makes such a close contact possible. -H+(Na+) OH Groups An OH maximum at 3740cm-' is present in the spectrum of Na-boralite [fig. 2A, spectrum (a)], thus indicating that silanol Si-OH groups exist in boralites. The frequency and intensity of the Si-OH band in boralite (3740 cm-', A / m = 3.65 g-') approach those in ZSM-5 zeolite (3738 cm-l, A / m = 3.35 g-'), thus suggesting that the nature and properties of the silanol Si-OH groups are the same in both cases. It has been foundlg that silanol Si-OH groups in ZSM-5 zeolites are formed during the composition of TPA ions, and they are situated inside the zeolitic chanels. It may be taken that the same is true of silanol Si-OH groups in boralites.The intoduction of protons into the boralites (by the decomposition of NH; ions) results in the appearance of new OH bands at 3680 and 3720 cm-' as well as a broad, weak band at 3460 cm-l. This last band can be attributed to a hydrogen bond between two adjacent OH groups.20*21 The bands at 3680 and 3720cm-l represent free OH groups. Note that only one OH band (apart from silanol Si-OH) was reported in previous studies of boralites. Chu and Chang' reported an OH band at 3725 cm-l, and Coudurier and Vedrinell reported a band at 3695 cm-l. Both the 3680 and 3720 cm-l OH bands were found in our study. It seems that both kinds of OH groups can be formed in H-boralites, but (depending on the sample composition and on the pretreatment conditions) one becomes dominating and is observed in the spectrum.The frequencies of the 3680 and 3720 cm-' OH bands are markedly higher than the OH frequencies in suggesting that the acid strength of these OH groups is lower. Acidic properties of both kinds of OH groups (3680 and 3720 cm-') in H-boralites were studied by pyridine sorption. The molecules of pyridine react with the 3720 cm-l OH groups, forming PyH+ ions, thus indicating that these OH groups are Brarnsted-acid sites. A decrease of the 3680 cm-l band is also observed after pyridine sorption (fig. 2B); however, it is difficult to decide if this decrease is due to proton transfer or only to a physical interaction. The 3720 cm-l OH groups are the main source of protons in H-52 I.R.Study of Boralites boralites. The numbers of Brransted-acid sites determined by pyridine sorption (presented in table 1) in Na-boralite, H-boralite and zeolite H-ZSM-5 are comparable with the theoretical numbers of protons calculated as the difference between the contents of B (or Al) and Na. A proton deficit in H-boralite and H-ZSM-5 may be due to dehydroxylation, which may occur during the pretreatment (Lewis-acid sites were found in both cases, Experiments involving the desorption of pyridine (table 1) show that the Brarnsted- acid sites in H-boralites are much weaker than those in H-ZSM-5. The same conclusion can also be drawn when comparing the OH vibration frequencies. Low-strength acid sites in boralites were also evidenced by earlier t.p.da6P ''9 l2 and m.a.s.n.m.r.6 studies. This effect is unexpected, since the electronegativity of boron is higher than that of aluminium (2.93 and 2.22, respectively), and (according to the collective model) an increase in average electronegativity is followed by an increase in acid-site strength.The following explanation of the low strength exhibited by the acid sites in boralites can be proposed. Because of its small size, the boron atom is situated between three oxygens in one face of a tetrahedral hole (1). The distance to the fourth oxygen (forming the OH group) is longer and the interaction B--*OH- is much weaker than that in zeolites. Such an OH group behaves more like a silanol Si-OH group than like Al-OH-Si (1). The vibrational frequency vCC(N) of a pyridine molecule bonded to an electron-acceptor site depends on the nature of the site.In a series of cationic forms of zeolite Y Ward23y24 observed an increase in this frequency with the electrostatic field of the cation (an increase of the cation charge and a decrease of the radius). In H-boralite the frequency v ~ ~ ( ~ ) of pyridine bonded to Lewis-acid sites (1460 cm-') is higher than that in zeolite H-ZSM-5 (1450 cm-'), thus indicating that the electrostatic field of Lewis-acid sites in boralite is stronger than in the zeolite. This may be due to the higher electronegativity and and smaller size of boron in comparison with aluminium. fig. 2C). We thank Dr A. Cichocki of the Jagiellonian University for donating samples of boralites. References 1 M. Taramasso, G.Perego and B. Notari, Proc. 5th Znt. Conf. Zeolites, Naples, 1980 (Hayden, London, 2 R. M. Barrer, Hydrothermal Chemistry of Zeolites (Academic Press, London, 1982), p. 251. 3 N. A. Kutz, Proc. 2nd Symp. of Industry-University Cooperative Chemistry Program (Texas A and M University Press, College Station, 1984), p. 121. 4 Z. Gabelica, G. Debras and J. B. Nagy, Catalysis on the Energy Scene (Elsevier, Amsterdam, 1984). 5 R. M. Dessau and G. T. Kerr, Zeolites, 1984, 4, 315. 6 K. F. M. G. J. Scholle, A. P. M. Kentgens, W. S. Veeman, P. Frenken and G. H. P. van der Velden, 7 K. F. M. G. J. Scholle and W. S . Veeman, Zeolites, 1985, 5, 118. 8 M. G. Howden, Zeolites, 1985, 5, 334. 9 G. T. W. Chu and C. D. Chang, J. Phys. Chem., 1985,89, 1569. 10 M. Tielen, M. Geelen and P. A. Jacobs, Proc. Znt. Symp. Zeolite Catal., Siofok, 1985, p. 1. 11 G. Coudurier and J. C. Vedrine, Pure Appl. Chem., 1986, 58, 1389. 12 P. Ratnasamy, S. G. Hedge and A. J. Chandwadker, J. Catal., 1986, 102, 467. 13 Zeolites, ed. Guo Wengui, Lieng Juan, Ying Muliang, Hu Jieban, H. Drzaj, S. Hocevar and S. Pejownik 14 C. T. W. Chu, G. M. Kuehl, R. M. Lago and C. D. Chang, 1985, J. Catal., 93, 451. 15 A. Gichocki, J. Datka, J. Klinowski, M. Michalik, A. Olech and Z. Piwowarska, to be published. 16 J. Datka and E. Tuznik, Zeolites, 5, 230. 17 D. M. Olson, G. T. Kokotailo and S. L. Lawton, J. Phys. Chem., 1981, 85, 2238. 18 C. A. Fyfe, J. H. O'Brien and M. Strobel, Nature (London), 1987, 326, 281. 19 J. Datka and E. Tuznik, J. Catal., 1986, 102, 43. 20 R. S. McDonald, J. Phys. Chem., 1958, 62, I 168. 1980), p. 40. J. Phys. Chem., 1984, 88, 5. (Elsevier, Amsterdam, 1985).J . Datka and Z . Piwowarska 53 21 G. Ghiotti, E. Garrone, C . Morterra and F. Bocuzzi, J. Phys. Chem., 1979, 83, 2863. 22 P. A. Jacobs, Catal. Rev. Sci. Eng., 1982, 24, 415. 23 J. W. Ward, J . Catal., 1968, 10, 34. 24 J. Ward, J . Colloid Interface Sci., 1968, 28, 269. Paper 8/00367J; Received 1st February, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500047

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1989, 85(1), 55-64 Effect of Spinel Oxide Composition on Rate of Carbon Deposition Geoffrey C. Allen* and Josephine A. Jutson Central Electricity Generating Board, Berkeley Nuclear Laboratories, Berkeley, Gloucestershire GL13 9PB The deposition of carbon on fuel cladding and other steels results in a reduction in heat-transfer efficiency. Methane and carbon monoxide are added to the gaseous coolant in power reactors to reduce the radiolytic oxidation of the graphite moderator and this is known to increase the rate of carbon deposition. However, the composition of oxides formed on steel surfaces within the reactor may also influence deposition. In this investigation carefully characterised spinel-type oxides of varying com- position have been subjected to y-irradiation under conditions of tem- perature, pressure and atmosphere similar to those experienced in the reactor.The rate of carbon deposition has been studied using scanning electron microscopy (SEM) and energy-dispersive X-ray analysis (EDX). Carbon deposition has been observed in power reactors on surfaces used in reheater and superheater pipework under operating conditions. Such deposition results in a reduction in heat-transfer efficiency and may require ‘ down-grading ’ of the reactor to prevent overheating of the fuel. Two probable sources of the carbon deposit are methane and carbon monoxide which are added to the carbon dioxide gas coolant to reduce the radiolytic oxidation of the graphite moderator. Much research has concentrated on the identification of an optimum gas coolant composition to prevent corrosion of the graphite moderator while reducing the rate of carbonaceous deposit on fuel cladding and other surfaces.l Recent observations have indicated that the surface morphology of fuel cladding also affects the rate of carbon deposition. If this is the case it is important to know how this occurs and to identify accurately the nature of the surface responsible for producing such surface accumulations. The steels used form ‘protective’ oxide layers composed of two major phases, a rhombohedra1 phase consisting of Cr,O,, Fe,O, or a solid solution of these binary oxides together with a cubic spinel phase.Oxides formed in fuel regions consist mainly of Cr,O, and a manganese-iron-chromium spinel,2 while those formed on reheater and superheater alloys consist mainly of a chromium-nickel-iron spinel with Fe,O, and Cr203., The temperature of these surfaces varies according to their position in the reactor.Since the composition of the oxide formed on a steel also varies with temperature, a range of oxide compositions would be expected to form on each type of steel. It has been suggested that iron and nickel may be active in catalysing the deposition of To verify these suggestions a programme of research has been initiated to investigate the effect of iron, nickel, manganese and chromium on the rate of carbon deposition. A range of standard spinels has been prepared by solid-state reaction6 in atmospheres similar to those of the Advanced Gas Cooled Reactor.Each series of spinels synthesised involved two or more of the component metallic elements of the substrate alloy in varying proportions. Each of the prepared spinels was subjected to radiation in the y cell 5556 Carbon Deposition on Spinels facility at Berkeley Nuclear Laboratories under conditions similar to those experienced by fuel cladding and other steels in reactors. Results from this study are given below. Experimental Spinels in the series Ni,Fe,-,Fe,O,, Mn,Fe,-,Fe,O,, FeFe,-,Cr,O, and Mn,Fe,-,Cr,O, were prepared by solid-state reaction at 950 "C in an atmosphere of C0,/2 O h C0.6 The structure and composition of each spinel was confirmed by X-ray diffraction (XRD) and energy-dispersive X-ray analysis (EDX). ' X-Ray Diffraction Lattice parameters were recorded using a Philips PW 1050 vertical X-ray diffractometer with unfiltered Cu radiation.A graphite crystal monochromator in the diffracted beam removed Cu K,, fluorescent and incoherently scattered radiation. The powdered samples, which consisted of small randomly orientated crystals, were dispersed in methanol and mounted on a non-reflecting silicon crystal wafer to improve the signal- to-noise ratio. Evaporation of the methanol gave an even dispersion of the specimen. Energy-dispersive X-Ray Analysis SEM and EDX analysis was carried out using a Cambridge Instruments Model S150 Mk I1 fitted with a Kevex windowless detector analysis system. Qualitative analysis was obtained by measurement of the peak energy in the characteristic spectrum of a sample.Relative amounts of carbon on each surface were compared by measuring the intensity of the carbon K, line in the spectrum obtained from each of the samples examined. Radiation Experiments Each spinel was pressed into a disc of 8 mm diameter, 1-2 mm thick and loaded into a siiica tube with silica spacers designed to allow unimpeded gas flow over the discs (fig. 1). The silica tube was then loaded into a stainless-steel capsule and placed in the y cell facility at Berkeley Nuclear Laboratories (fig. 2). A 6oCo source is used in this facilitys and the general arrangement of the source within the y cell is indicated in fig. 3 at the end of the guide tube. The capsule was placed in an outer irradiation position (fig. 4) where a dose rate of 8.20-8.34mW g-l was obtained Samples were exposed to two separate temperature regimes within the y cell and examined by SEM and EDX after these exposures.Experiment 1 : y Cell Exposure at 200 "C Spinels in the series FeFe,-,Cr,O, and MnFe,_,Cr,O, were maintained at a temperature of 200 "C for 30 days in the y cell facility. Gas of composition CO,/l O h CO with 800 vpm CH, and 15 vpm C2H6 flowed through the capsule at a rate of 2-3 cm3 min-l and 40 bar pressure in a single-pass experiment. Water was added to the gas mixture to the extent of 200-300 vpm. Experiment 2: y Cell Exposure 550 "C Spinels in the series Ni,Fe,-,Fe,O, and Mn,Fe,-,Fe,O, were maintained at 550 "C for 15 days using the gas composition and flow conditions of experiment 1.J . Chem. SOC., Faraday Trans. I , Vol.85, part 1 Plate 1 Plate 1. Electron micrographs showing Fe,O, spinel surface before (a) and after (b) exposure in the y cell for 30 days at 200 "C: G. C. Allen and J. A. Jutson (Facing p . 56)J . Chem. SOC., Faraday Trans. I , Vol. 85, part 1 Plate 2 Plate 2. Electron micrographs showing Fe2,BCr0,404 spinel surface before (a) and after (b) exposure in the y cell for 30 days at 200 "C. G. C. Allen and J. A. JutsonPlate 3. Electron micrographs showing carbon spheres on the surface of ( 1 1 ) FeCr,O, and ( h ) Mn,,7,Fe,,,,Cr10, after expobure in the ;I cell for 30 days at 200 "C. G. C . .411en and J. A. JutsvnG . C. Allen and J . A . Jutson 57 / f f / 0 \ . \ 1 cam part m en t \ . f f , . . .. . silica silica spinel tube spacer disc gas flow vent - 10 mm Fig.1. Arrangement of spinel discs in y-cell capsule. double doors 1 gamma [cell - rack housing L source B- flask concrete - shielding Plug door Fig. 2. Plan of y cell. Results y Cell Exposure at 200 "C Visual examination of the discs indicated a blue interference film at the surface of the spinels with high iron content in the series FeFe,-,Cr,O,. Subsequent examination by SEM and analysis using EDX showed that this was due to carbon. In fact, all of the spinels in both series had some carbonaceous deposit on the surface. Micrographs of back-scattered images showed what appeared to be a general layer of carbon across the spinel surface. Secondary electron micrographs of this layer are shown in plates 1 and58 Carbon Deposition on Spinels y//A irradiation posit ion 1000 mm t 1 Fig.3. General arrangement of source. irradiation positions W \source carrier 10 cm I I Fig. 4. Arrangement of irradiation position. 2, and a corresponding layer with scattered spherules of carbon is shown in plate 3. EDX analysis was carried out using 20 and 10 kV electron-beam exciting voltages. The intensity of the carbon K, X-ray signal measured in counts (100 s)-' was much higher for analysis of a given sample using a 10 kV rather than 20 kV electron-beam energy and in every case the carbon intensity measured from spherule regions by EDX was higher than that recorded from the general surface (tables 1 and 2). y Cell Exposure at 550 "C No spherules of carbon were identified on the spinel surface following exposure to the same gas conditions but higher temperature.A layer of carbon was apparent only over small areas of the surface of the spinel samples with high iron content. The measured intensity of the carbon K, X-ray line for each spinel is listed in tables 3 and 4.G. C. Allen and J. A. Jutson 59 Table 1. Carbon counts per 100 s for spinel series FeCr,Fe,-,O, maintained in the cell for 30 days at 200 "C spinel general surface sp herule 10 kV 20 kV 10kV 20kV Fe3+(Fe2+Fe3+)0," - 5583 4266 19 989 1776 1159 1506 990 3138 2026 672 205 3386 1467 1025 985 322 1 1059 928 478 15341 - - - Inverse (a) - normal (b) spinel transition, degree of inversion estimated for intermediate spinels. Octahedral site occupancy indicated by brackets. Table 2. Carbon counts per 100 s for spinel series Mn,Fe,-,Cr,O, maintained in the cell for 30 days at 200 "C spinel general surface spherule 10kV 20kV 10 kV 20 v 928 478 15341 - - 447 533 71 1 1430 303 - 4392 - - 506 - Normal-normal spinel transition, octahedral site occupancy indicated by brackets.Table 3. Carbon counts per 100 s for spinel series Mn,Fe,-,Fe,O, maintained in the cell for 15 days at 550 "C carbon counts (general surface) spinel 10 kV 20 kV Fe3+(Fe2+Fe3+)Oda - 2165 /MnOb - 1157 1081 I805 433 263 36 1 526 Inverse (a) - normal (b) spinel transition, degree of inversion estimated for intermediate spinels. Octahedral site occupancy indicated by brackets.60 Carbon Deposition on Spinels 0. Table 4. Carbon counts per 100 s for spinel series Ni,Fe,-,Fe,O, maintained in the cell for 15 days at 550 "C b I I I I spinel carbon counts (general surface) 20 kV Fe3+( Fe2+Fe3+)0, 2165 487 390 473 Fig.inverse-inverse spinel transition, octahedral site occupancy indicated by brackets. 6000 A -- loo0 t - -A - - -- -A- - - -- -A - 0.5 1.0 1.5 2.0 X Fe C r 0 5. Carbon deposition on spinels in the series FeFe,-,Cr,O, after 30 days in the y cell at 200 "C (EDX). A, 10 kV; 0, 20 kV. For both experiments the intensity of the recorded signal for a given sample was plotted against its composition. For a given series of spinels the relative carbon count was considered to be proportional to the relative amount of carbon deposition. Discussion y Cell Exposure at 200 "C The results for the series FeFe,-,Cr,O, showed a decrease in the amount of carbon deposited with decreasing iron(II1) content of the spinel-type oxide and increasing chromium concentration (fig.9, indicating iron to be active in promoting carbon deposition. The increased penetrating power of the 20 kV electron beam compared with that of the 10 kV beam, gave lower intensity carbon K, signals suggesting that this element was present as a surface layer on the spinel rather than incorporated within the oxide matrix. This is in agreement with the visual evidence from SEM micrographsG. C. Alien and J . A . Jutson 61 2 8 I I I 1 0.25 0.50 0.75 1 Fe C r20, X MnCr204 Fig. 6. Carbon deposition on spinels in the series Mn,Fe,-,Cr,O, after 30 days in the y cell at 200 "C (EDX). 20 kV electron beam (plates 1 and 2). A comparison of the disc surface before and after exposure in the y cell showed an almost transparent layer surrounding the oxide particles which was attributed to carbon.The spherule shapes identified at the spinel surface which gave rise to more intense energy-dispersive X-ray signals appeared to consist solely of carbon. By contrast, the intensity of the carbon K , signal recorded from the Mn,Fe,_,Cr,O, series showed little variation with change in iron(I1) or manganese(I1) content (fig. 6). A careful consideration of the results from these two series indicated that while iron@) in the absence of iron(m) had little effect on carbon deposition, a variation in the iron(m) concentration when iron(I1) was present in constant concentration markedly affected the rate of deposition. y Cell Exposure at 550 "C In general the intensity of the carbon K, line recorded for the series Mn,Fe,-,Fe,O, and Ni,Fe,,Fe,O, was lower than that measured at 200 "C.The exposure time was also lower, but previously it had been supposed that, at a higher temperature, deposition would occur at a faster rate since methane destruction and vinyl radical (C&) production, which is thought to be a precursor to deposition, is known to increase with increasing temperat~re.~ However, it is likely that more than one process leading to carbon deposition was involved. At the lower temperature, deposition may have occurred via the Boudouard reaction : CO,(g) + CO(ads) (1) CO(ads)+ C(ads) + O(ads) (2) (3) while at higher temperatures a deposition mechanism in which the radical products of radiolysis were involved was likely to have been important. O(ads) + CO(g) or CO(ads) - CO,(g) CO;(g) - cow (4) COi(g) ----+ C0,CO; or CO,CO,CO;t polymerisation ( 5 )62 Carbon Deposition on Spinels 0.25 0.5 X 0.75 1 Mn Fe 0, Fig.7. Carbon deposition on spinels in the series Mn,Fe,-,Fe,O, after 15 days in the y cell at 550 "C (EDX). A, 10 kV; .,20 kV. positive ions C2H&3ds) + 3CHj(g)- 2C(ads) + 3CH,(g) . (10) At this stage in the investigation, however, it is not possible to identify either as being the major pathway to deposition. For the spinel series Mn,Fe,-,Fe,O, the intensity of the measured carbon K, signal decreased very rapidly at first and slowly thereafter with decreasing iron(I1) content (fig. 7). The exception to this general observation is MnFe,O,, where the presence of manganese appeared once more to encourage deposition.The preparation of MnFe,O, in C0,/2% CO in fact produced a mixture of a spinel which appeared to be depleted in Mn" and enriched in Fe" with a lattice parameter corresponding to that for Mno,,Fe,,,Fe,O, together with a small amount of MnO. Therefore manganese ferrite was specially prepared in CO, but this product again decomposed to give a manganese depleted spinel and MnO after exposure in the y cell. Manganese oxide is known to promote iron as a catalyst for CO decomposition," and this could have been responsible for the increased carbon count. More than one solid-state property may be responsible for promoting the deposition of carbon, although the activation energy for the thermal outer-sphere Mn"/Mn"' electron transfer between ions in aqueous solution is close to that for the Fe"/Fe"' electron transfer under similar conditions" and both types of intervalence charge- transfer transition have, for example, been identified in the mineral yoderite (Al, Mg, Fe, Mn), Si,01,(OH)2.12 The question here is whether the manganese ion occupies tetrahedral or octahedral positions and whether it has an oxidation state of two or three.Conductivity measurements on manganese-iron-oxygen spinels indicate the following equilibrium between octahedral ions for the temperature range - 180 to 300 "C. Fe3+ + Mn2+' Fez+ + Mn3+; + 0.30 eV.G. C. Allen and J. A . Jutson 0 63 I I I I Fig. 8. Carbon deposition on spinels in the series Ni,Fe,-,Fe,O, after 15 days in the y cell at 550 "C (EDX).20 kV electron beam. This equilibrium lies well to the left and, according to Rieck and Driessens,13 the ground- state formula for manganese ferrite may be written and the equilibrium constant for the above distribution is ca. 36.14 In view of this it is difficult to attribute the enhanced deposition behaviour of our sample to Mn2+ + Fe3+ charge transfer when the Mn2+ occupancy of octahedral sites in the spinel lattice is so small. Moreover, it is highly unlikely that a heteronuclear Mn2+ + Fe3+ intervalence interaction would occur more readily than the homonuclear Fez+ + Fe3+ transfer. A possible explanation is that the presence of manganese and the reducing conditions of the experiment enhance the Fez+ occupancy of octahedral sites, thereby increasing the Fe2+/Fe3+ ratio in these positions and hence the catalytic activity of the oxide substrate.For the series Ni,Fe,-,Fe,O, the recorded carbon Ka signal decreased with increasing nickel content (fig. 8). However, X-ray diffraction showed that exposure in the y cell in the reducing gas atmosphere produced partial decomposition of nickel-containing spinels, the products being nickel and an iron(I1)-enriched spinel. This reduction process has been observed previously during the preparation of nickel spinels.6 It is possible that for these spinels the rate of deposition depends on both the iron and nickel content. In all the experiments carried out so far the greatest rate of deposition has occurred on magnetite (Fe,O,) which, as an inverse spinel, contains equal amounts of Fe" and Fe"' in octahedral sites.Electron exchange between these ions stabilises the spinel structure, offering a catalytic surface through the process of mixed-valence interaction. l5 When Cr'l' replaces Fe"' in the series Fe Cr,Fe,-,O,, Fe" is displaced to the tetrahedral position, thereby reducing the capability for intervalence electron exchange in the octahedral lattice. Similarly, when Fe" is replaced by Ni" or up to 50% Mn" the capability for intervalence electron exchange is reduced and in each experiment this reduction was accompanied by a reduction in carbon deposition. Previous work on the catalytic effect of iron and iron oxides on carbon deposition16 has shown that Fe,O, is a more active catalyst for deposition than Fe,O,, FeO or Fe, giving further support to the view that the Fe"/Fe"' couple may be the most effective of 3 F A R 8564 Carbon Deposition on Spinels the ion species.Rethwisch and Dumesicl' have also noted a high activity of magnetite for water-gas shift relative to other oxides. This they attribute to the variable oxidation state of iron which facilitates surface oxygen transfer. A number of processes involving iron, manganese and nickel may be important in deposition and the present studies indicate that of these Fe"/Fe"' exchange could be the most important. Conclusions Preliminary investigations into the effect of surface oxides on carbon deposition have shown that the iron content of the spinel-type oxide plays a major role? at least in the initial stages. The effect of the presence of nickel and manganese is not yet clear, but future experiments involving exposure of oxides for much longer periods in the y cell should clarify their role in deposition.Future work will also include the quantification of deposits, the deposition behaviour of other spinel surfaces, the effect of exposure to reactor-type gas environment without y radiation and the variation in deposition behaviour with temperature. This work was carried out at the Berkeley Nuclear Laboratories of the Research Division and the paper is published with permission of the Central Electricity Generating Board. The authors thank Mr S. M. Underwood for his work on the scanning electron microscope, also Mr J. V. Best and Mr R. M. Parfitt for their assistance with the irradiation experiments in the y cell. References 1 G. R. Marsh, D. J. Norfolk and R. F. Skinner, CEGB Report TPRD/B/0592/N85 (1985). 2 P. A. Tempest and R. K. Wild, J. Nucl. Mat., 1981, 102, 183. 3 A. F. Smith, Werkstofle Korrosion, 1981, 32, 1. 4 A. M. Brown, A. M. Emsley and M. P. Hill, CEGB Report No. RD/L/M/160/80 (1981). 5 R. T. K. Baker and J. J. Chludzinski, J. Catal., 1980, 64, 464. 6 G. C. Allen, J. A. Jutson and P. A. Tempest, J. Nucl. Mat., in press, 7 J. T. Buswell and G. K. Rickards, CEGB Report RD/B/N3365 (1975). 8 J. V. Best, CEGB Report RD/B/N2431 (1972). 9 D. J. Norfolk, R. F. Skinner and W. J. Williams, Gas Chemistry in Nuclear Reactors and Large Zndustrial Plant, Conference Proeedings, University of Salford, April 1980. 10 K. M. Kreitman, M. Baerns and J. B. Butt, J. Catal., 1987, 105, 319. 1 1 N. S. Hus, Electrochim. Acta, 1968, 13, 1005. 12 R. M. Abu-Eid, K. Langer and F. Siefert, Phys. Chem. Min., 1978, 3, 271. 13 G. D. Rieck and F. C. M. Driessens, Acta. Crystallogr., 1966, 20, 521. 14 A. Miller, J. Appl. Phys., 1960, 31, 2615. 15 G. C. Allen, P. M. Tucker and R. K. Wild, Philos. Mag., 1982, 46, 41 1. 16 R. T. K. Baker, R. J. Alonso, J. A. Dumesic and D. J. C. Yates, J. Catal., 1982, 77, 74. 17 D. G. Rethwisch and J. A. Dumesic, Appl. Catal., 1986, 21, 97. 18 G. C. Allen, J. A. Jutson and P. A. Tempest, CEGB Report TPRD/B/1028/R88 (1988). 19 H. J. Yearian, J. M. Kortvight and R. H. Langenheim, J. Chem. Phys., 1954, 22, 1 196. Paper 8/0047 1 D ; Received 29th January, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500055

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1989, 85(1), 65-70 Effect of Pd-TiO, Interaction on the Enthalpy of Hydrogen Absorption Chung Hsian Foo, Cheng Tsung Hong and Chuin-tih Yeh* Department of Chemistry, National Tsinghua University, Hsinchu, Taiwan, Republic of China The uptake of hydrogen (chemisorption and absorption) by palladium supported on titania has been measured volumetrically at various temperatures. On increasing the hydrogen pressure, a-phase palladium hydride absorbed hydrogen and transformed into p-phase. The transition pressure increased with the measuring temperature. The enthalpy of the transition was decreased by increasing the temperature for reduction pretreatment. These variations are attributed to electronic interactions between the metallic palladium and the reduced support (TiO,).Eight years ago Tauster et a1.l reported suppression of hydrogen and carbon monoxide absorption on supported metallic catalysts by a high-temperature treatment of hydrogen reduction. Many papers have been published on the metal-support interaction since this report., Two major mechanisms, i.e. modification of the chemical properties of the active metal by transferring electrons from the partially reduced support and blockage of adsorption sites on the active metal by a migration of the reduced support, have been suggested as the possible reasons for this suppre~sion.~ However, whether electronic interaction does occur remains a matter of controversy.4 An electronic interaction should cause a variation in the enthalpy of chemisorption.Recently the metal-support interaction has been studied using calorimetric measurement of the enthalpy of hydrogen adsorption by two groups i~~dependently.~?' Distinct conclusions were reached from these two studies. Herman et al. found that the initial enthalpy of hydrogen adsorption on Pt/TiO, decreased (from 92.0 to 79.5 kJ mol-l) owing the the metal-support interaction using differential enthalpy meas~rements.~ An electronic interaction is therefore indicated. On the other hand, Chou et al., by measuring integral enthalpy, studied the average enthalpy of adsorption.' They found negligible change (60 us. 64 kJ mol-') in the heat of hydrogen adsorption on Pd/TiO, for the interaction. They suggested that the electronic properties of metal are not markedly affected.The formation of bulk hydride is a distinctive property of palladium. The uptake is0 them of hydrogen on supported palladium generally includes chemisorption and ab~orption.~ Both the chemisorption and the absorption of hydrogen on palladium catalysts have been found to be suppressed by the metal-support interaction.'-1° Burch and Bray' and Chou et al.' have pointed out that if the metal-support interaction induces an electronic interaction in palladium crystallites, all the atoms in the crystallites are likely to be affected. Therefore, the enthalpy of palladium hydride formation provides a good test for the existence of electronic effects in the metal-support interaction. In this paper we report our results for the variation of the enthalpy of hydrogen absorption on a Pd/TiO, catalyst caued by reduction with hydrogen.65 3-266 Enthalpy of Hydrogen Absorption Table 1. Procedures for the volumetric mea- surements on the Pd/TiO, catalyst step procedure 1 2 reduce the sample with flowing hydro- evacuate the sample to a vacuum gen at 373 K for 2 h (< Torr) for 10 min at 333 K and measure the hydrogen uptake isotherm at 273 K repeat step 2 twice, but increase the temperature of the measurement to room temperature and 315 K, respectively. repeat steps 1-3 three times, but increase the temperature of reduction pretreatment to 473, 573 and 673 K, respectively. 3 4 Experimental Two palladium samples, 2 % Pd/TiO, and palladium black, were used in this study. The Pd/TiO, catalyst was prepared by impregnating TiO, (Degussa P-25) with PdCl, (Merck). The wet suspension was evaporated on a steam bath with constant stirring and then dried in an oven at 383 K overnight.The catalyst was calcined at 673 K and stored in a desiccator for further use. The change in particle size of palladium crystallites on Pd/TiO, with reduction treatment was measured by TEM. The catalyst samples were deposited on a 400 mesh copper screen coated with collodium film for the measurements. A JEM-200 CXTEM (200 kev) was used for the detection. Palladium black was purchased from Hayashi Pure Chemicals and used without further purification. The average particle size of this sample is ca. 1 pm. For hydrogen reduction treatments and hydrogen uptake measurements, samples were placed in a cell connected to a vacuum system equipped with a digital precision gauge from Texas Instruments.The ultimate vacuum of this system is ca. 5 x Torr.? The samples were pre-reduced in flowing hydrogen (30 cm3 min-l at 1 atml) for 2 h at the desired temperature and then evacuated at 333 K for 10 min before all isotherm measurements. The procedures for the volumetric experiments performed on the Pd/ TiO, are listed in table 1. The hydrogen gas used in the isotherm measurement was purified with a Metheson 8361 hydrogen purifier. Results and Discussion Fig. 1 shows uptake isotherms of hydrogen by palladium black measured at different temperatures. At low pressures (P < 1 Torr) palladium was in the a-palladium hydride phase with a stoichiometry of PdH, (a x 0.05). As the pressure of hydrogen was increased, the a-hydride absorbed more hydrogen and transformed into P-hydride, PdH, ( b x 0.60), according to: 2PdH, + ( b - a)H, f 2PdHb.(1) t 1 Torr = 101 325/760 Pa. 1 1 atm = 101 325 PaC. H. Foo, C. T. Hong and C-t. Yeh 67 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 hydrogen uptake (H/Pd) Fig. 1. Hydrogen-uptake isotherms for palladium black measured at various temperatures : 0, 273; U, 297; 0, 315; 0, 333 K. Note that the amount of hydrogen absorption during the phase transition, i.e. (b-a), did not change significantly in the small temperature range of this experiment. The transition pressure (Ptr), on the other hand, is shown in fig. 1 to increase with the sorption temperature. According to the thermodynamics of phase transitions P,, should change with temperature according to :11 (2) where AH is the isosteric enthalpy of the phase transition.Taking P,, as the pressure at which half of the phase transition was accomplished (indicated with a horizontal bar on each isotherm in fig. I), AH = 36.740.6 kJ (mol H2)-' was found for palladium black. This value is in good agreement with reported results [in the range 35.4-39.3 kJ (mol H,)-'] from calorimetric and isosteric measurements on palladium black and Palladium is a noble metal that can be reduced very easily. T.p.r. studies in this laboratory showed that palladium oxides supported on TiO, could be fully reduced to metal at a lower temperature (below 300 K) than those supported on SO, and A1,0, (ca. 400 and 450 K, respectively).'" Palladium crystallites on the Pd/TiO, sarnples were finely dispersed.The size of :he crystallites was found to increase from 45 A for the 373 K reduced sample to 100 A for the 673 K reduced sample according to TEM measurements. The sorption of hydrogen on these fine palladium crystallites therefore included chemisorption as well as absorption. Fig. 2 shows sorption isotherms obtained for the Pd/TiO, catalysts reduced at 373, 473, 573 and 673 K, respectively. These isotherms were measured at 273 K after evacuation at 333 K. Absorbed hydrogen is unstable in the absence of hydrogen gas. It can remain in the palladium bulk only under an equilibrium hydrogen pressure > 1 Torr at room ternperatu~e.~,'~ Consequently, such a mild out- gassing treatment might desorb all the hydrogen absorbed in the palladium bulk.This mild desorption (compared to the 673 K desorption pretreatment normally used in our laboratory for palladium catalysts supported on other oxides) was accepted in present study to prevent the possible desorption of oxygen from TiO, at elevated [d In PJd( 1 / T ) ] = - AH/ R wires.69 11-13 temperatures, l5 i.e. TiO, TiO, + (2 - x)/2 0,. (3)68 Enthalpy of Hydrogen Absorption hydrogen uptake (H/Pd) Fig. 2. Hydrogen-uptake isotherms (measured at 273 K) for 2% Pd/TiO, catalysts reduced at 373 (o), 473 (m), 573 (0) and 673 K (0). The isotherms were obtained after evacuation of the reduced samples below 2 x Torr at 333 K for 10 min. A short horizontal line on each isotherm indicates the P,,. The procedure for determining the P,, is illustrated in the uptake isotherm of the 373 K reduced catalyst.Point a on the abscissa is considered to be the amount of weak chemisorption while (6-a) is the amount of absorption. P,, is assumed to be the uptake pressure under which the absorption is half accomplished. This thermal desorption may induce an additional metal-support interaction. Two uptake steps were found in each isotherm in fig. 2. Chemisorption occured at low hydrogen pressure, while absorption started at pressures around 5 Torr (see fig. 1). The bond strength of chemisorbed hydrogen on Pd/TiO, is not uniform. Weakly chemisorbed hydrogen is pumped off the catalyst by our evacuation pretreatment at 333 K. The uptake found in fig. 2 thus came from the contribution of the hydrogen desorbed during evacuation, including weakly chemisorbed hydrogen on the Pd surface and hydrogen absorbed in the palladium bulk.On the isotherm of Pd/TiO, reduced at 373 K in fig. 2, an illustration is given on the procedure to separate the contribution of weak chemisorption (chemisorbed hydrogen that can be evacuated at 333 K, marked a on the abscissa of the figure) from that of absorption (from A to B on the isotherm). Both chemisorption and absorption were suppressed by reduction at high temperatures. In agreement with literature reports,',* the suppression of absorption is less dramatic than that of chemisorption. Also shown on the isotherm of the 373 K reduced sample is the way to evaluate Ptr for the absorption. The observed Ptp increases with the temperature of reduction. A change in the absorption properties of palladium crystallites with the reduction temperature is therefore indicated.Similar uptake isotherms in fig. 2 have also been measured for these four catalysts at room temperature and at 313 K. Ptr obtained in these isotherms increased with the temperature of hydrogen uptake. Fig. 3 summarizes plots of eqn (2) for catalysts reduced at different temperatures. The slope of these plots decreases with increasing reduction temperature. Table 2 lists the enthalpy of absorption evaluated according to eqn (2). Palladium black has the highest enthalpy of absorption (36.7 kJ). The enthalpy for Pd/TiO, catalysts decreases (from 35.5 to 32.8 kJ) with increasing reduction temperature. Note that these enthalpies are calculated from theC.H . Foo, C. T. Hong and C-t. Yeh 69 5 4 -2 s e 3 5 2 1 I I I I I I 30 3.1 3.2 3.3 3.4 3.5 3.6 3.7 Fig. 3. The isosteric plot of Pt,, the hydrogen pressure at which the phase transition due to absorption is 50 YO accomplished, against the temperature of adsorption measurement for palladium black (+) and 2 % Pd/TiO, reduced at 373 (I), 473 (m), 573 (0) and 673 K (a). I d K/T Table 2. Effect of reduction treatment on the enthalpy of hydrogen absorption into palladium crystallites on Pd/TiO, reduction AH sample temp/K /kJ (mol H2)-' Pd/TiO, 373 35.5 & 0.7 473 34.8 k0.6 573 33.8 & 0.6 673 32.8 & 0.6 Pd black 36.7 f 0.6 dependence of P,, on the uptake temperature. Their values should be free from the interference of weak chemisorption. A change in the particle size of the palladium crystallites during the reduction treatment may be considered as a possible reason for the observed variation in the enthalpy of absorption. A close examination by TEM, however, found that the ayerage particle size of palladium crystallites on all the Pd/TiO, catalysts was > 30A.The enthalpy of hydrogen absorption should not change with the size of Pd in this size range.6 The observed change in AH should not come from the particle size effect. This study therefore reveals a change in the chemical properties of palladium crystallites on TiO, during the reduction treatment. A possible explanation for this change is transferral of electrons from TiO, produced by hydrogen spillover during the reduction treatment to the palladium meta1.l~~ A decrease in the enthalpy of hydrogen absorption in palladium has also been reported for palladium+opper alloy.A flow of electrons from copper to palladium has been proposed as the major cause for this decrease.16 Unfortunately, this study can not verify whether this electronic interaction stems from the formation of Pd-Ti intermetallic compounds. It should be emphasized70 Enthalpy of Hydrogen Absorption here that reduced titania may migrate to the surface of the palladium crystallites. Part of the observed effect could come from the electronic interaction of these migrant TiO, species. The extent of this transference may be small and limited only to a few layers near the surface of the palladium crystallites from the Pd-TiO, interface.17 We acknowledge the financial support of this study by the National Science Council of the Republic of China.The determination of the particle size distribution with TEM by Mr Fu-Hwa Yu of this laboratory is also appreciated. References 1 S. J. Tauster, S. C. Fung and R. L. Garten, J. Am. Chem. SOC., 1978, 100, 170. 2 G. C. Bond and R. Burch, in Catalysis, ed. G. C. Bond and G. Webb, Specialist Periodical Report (The 3 M. G. Sanchez and J. L. Gazguez, J. Catal., 1987, 104, 120. 4 S. J. Tauster, Ace. Chem. Res., 1987, 20, 389. 5 J. M. Herrman, M. Gravelle-Rumeau-Maillot and P. C. Gravelle, J. Catal., 1987, 104, 136. 6 P. Chou and M. A. Vannice, J. Catal., 1987, 104, 1. 7 G. Chen, W. T. Chou and C. T. Yeh, Appl. Catal., 1983, 8, 389. 8 R. Burch and J. D. Bracy, J. Catal., 1984, 86, 384. 9 R. T. K. Baker, E. B. Prestridge and G. B. Mcvicker, J. Catal., 1984, 89, 422. Chemical Society, London, 1983), vol. 6, p. 27. 10 T. C. Chang, J. J. Chen and C. T. Yeh, J. Catal., 1985, 96, 51. 11 T. B. Flanagan and J. F. Lynch, J. Phys. Chem., 1975, 79, 444. 12 E. Wicke and G. H. Nevst, Ber. Bunsenges. Phys. Chem., 1964, 68, 224. 13 D. M. Nace and J. G. Aston, J. Am. Chem SOC., 1957, 79, 3619, 3623, 3627. 14 M. Boudart and H. S . Hwang, J. Catal., 1975, 39, 44. 15 M. Che, C. Naccache and B. Imelik, J. Catal., 1972, 24, 328. 16 R. Burch and R. G. Ross, J. Chem. SOC., Faraday Trans. I, 1975, 71, 922. 17 G. B. Raupp and J. A. Dumesic, J. Catal., 1986, 97, 85. Paper 8/00562A; Received 15th February, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500065

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chern. SOC., Faraday Trans. I , 1989, 85(1), 71-78 Chromia/Silica-Titania Cogel Catalysts for Ethene Polymerisation Polymerisation Kinetics Steven J. Conway, John W. Falconer and Colin H. Rochester* Department of Chemistry, The University, Dundee DDl 4HN, Scotland The kinetics of ethene polymerisation over a series of chromia/silica-titania cogels have been studied. The catalysts were activated by reduction in CO at 623 K. Rates of polymerisation were first order in ethene pressure below 373 K. Above 373 K the rates tended towards second-order kinetics. Arrhenius plots were non-linear and exhibited a maximum, the nature of which was dependent upon the reduction time and the catalyst composition. For polymerisation at 273 K the activity of the catalysts passed through a maximum with increasing reduction time.Catalysts which were inactive at 273 K due to extensive reduction were active at 343 K, and in some cases gave bimodal activity curves. Active catalysts were deactivated by evacuation or reduction in CO at 873 K. Rates of polymerisation at 343 K for catalysts deactivated in uacuo at 873 K showed an initial decrease in activity which was followed by a small increase in activity before further deactivation. The results are consistent with a Langmuir-Hinshelwood polymerisation mechanism. Three types of active site have been identified. Heterogeneous catalysts are often prepared by depositing the active component on a solid support. The support can either act as a dispersive medium where both support and active component exist as two separate phases, or the support forms surface complexes with the active component. The Phillips chromium(vr)/silica ethene-polymerisation catalyst is an extreme case of the latter, as unsupported chromia (CrO,) cannot polymerise ethene.' Bulk CrO, decomposes above 473 K, but it is possible to stabilise CrO, up to 1173 K on a silica support by calcination in dry air.This stabilisation is due to the formation of surface chromate and/or dichromate species in which the chromium is linked to the support via Si-0-Cr bonds.2 The incorporation of titania either in or on the silica support results in a chromia catalyst capable of polymerising ethene to yield a polymer with lower average molecular weight than that for polymer generated over unmodified This is thought to be due to a change in the electronic environment of the chromium caused by the formation of Ti-0-Cr as well as Si-0-Cr bonds.The dependence of the polymerisation process on the nature of the chromium and the manner in which the titanium influences this process have here been examined by comparison of the polymerisation kinetics for a series of chromia catalysts supported on silica-titania cogels with titanium contents in the range W.2 wt%. Experimental Catalyst Preparation Silica-titania cogel supports were prepared by a coprecipitation m e t h ~ d . ~ The initial stage involved the coprecipitation of an aqueous sodium silicate solution and a titanium tetrachloride-sulphuric acid mixture. After the gel had been aged for 4 h at 373 K it was 7172 Modijied Phillips Catalyst Kinetics Table 1.Summary of porosity measurements for cogels (after calcination at 873 K) used in the preparation of catalysts pore volume/cm3 g-' titanium surface average pode (wt %) area/m2 g-Ia < 500 A" totalb diameter/A" 0 272 f 5 1.02 2.95 190 1.3 508 f 10 2.34 2.72 212 2.1 273 f 5 1.04 2.34 185 4.2 279 6 1.06 2.18 187 a Determined by N, adsorption. Data refer to a pore diameter of up to 500 A. liquid Determined by washed with a 5% ammonium nitrate solution and several times with distilled water to remove sodium and sulphate ions. Water was extracted from the gel by azeotropic distillation in ethyl acetate and the remaining solvent was removed by drying in air at 373 K for ca. 3 h. A summary of support characteristics are given in table I .The gel was subsequently impregnated with 1 wt% chromium by forming a slurry with a solution of chromium acetylacetonate in ethyl acetate. The resulting slurry was dried on a rotary evaporator. Kinetic Studies Kinetic studies were performed under static conditions using a conventional high- vacuum system linked to a reactor vessel consisting of a cylindrical quartz bulb with diameter 4 cm and volume 148 cm3. Reactant-gas pressures were monitored using a pressure transducer. Samples of catalyst (120 mg) which had been precalcined at 873 K for 16 h in a flow of dry air (50 cm3 min-') were placed in the reactor. The temperature of the catalyst in vacuum was raised to 873 K over a 30 min period and the catalyst was then contacted with oxygen (101 kN m-') for a further 30 min at 873 K.Subsequent evacuation at 873 K for 60 min was followed by a reduction in temperature to 623 K over a 30 min period. The catalyst was then activated by reduction at 623 K in CO (5.33 kN m-2) for various times, with subsequent evacuation for 30min prior to being cooled to the polymerisation temperature. Ethene (ca. 2.2 kN m-') was then admitted to the reactor vessel and the fall in ethene pressure was monitored as a function of time. Reaction temperatures in the range 250-4023 K were maintained with liquid thermostat baths. Two methods of deactivation of reduced catalysts were used; the first involved heating in vacuo at 873 K and the second involved heating in CO at 873 K for various times. Ethene (B.O.C. research grade, > 99.92%) with quoted impurity levels of 1 and 2 ppm of 0, and water, respectively, was subjected to a series of freeze-pumpthaw cycles before use.Carbon monoxide (B.O.C., > 97.5%) was passed through a cold trap (77 K) to remove condensable impurities. Results Calcined supports alone and supported chromia were found to be inactive for low- pressure ( < 3.0 kN m-,) gas-phase ethene polymerisation. However, supported chromia catalysts which had been reduced in CO at 623 K were active. Below 373 K polymerisation rates exhibited a first-order dependence on ethene pressure. Fig. I shows a typical first-order plot. First-order rate constants (k,) wereS. J . Conway, J . W. Falconer and C. H . Rochester 3.2- -- 3.0- ‘E 2 v % 2 2.8- 2 2.6- 0 - 2.4- 73 I I 1 1 - 0 60 120 180 240 300 tls Fig.1. Typical first-order plot for the derivation of rate constants. Example for ethene polymerisation at 343 K over a Cr/silica-titania (2.1 % Ti) cogel reduced in CO (623 K) for 15 min. 3.3 n - 3.2 ‘E 3 v !i 3.1 v) 2 3.0 2.9 4 1.4 - 1.2 b 1.0 2 “E m 2 2 0.8 \ d v 0.6 0 5 I 0 15 20 2 5 30 tlmin Fig. 2. First- and second-order plots for ethene polymerisation at 423 K over a Cr/silica-titania (1.3% Ti) cogel reduced in CO (623 K) for 15 min. deduced per gram of catalyst per second. Above 373 K the pressure dependence of the reaction rate began to deviate from first-order towards second-order behaviour (fig. 2). Polymerisation activities are here expressed as first-order rate constants (kJ. A non- linear dependence of log (polymerisation activity) on reciprocal temperature and the appearance of a maximum polymerisation activity were observed in Arrhenius plots (fig.3). The characteristics of the curves varied with reduction time and support composition. Maximum polymerisation activities were observed at 313 and 273 K for Cr/silica gel and Cr/silica-titania (4.2 O/O Ti) cogel, respectively, reduced in CO for 15 min. Maximum activity was observed at 343 K for both Cr/silica gel reduced in CO for 120 min and Cr/74 Mod$ed Phillips Catalyst Kinetics -1 .o- -Ij - 1 . 8 2 2 I (b) L I I I I 2.5 3.0 3.5 4.0 103 KIT -1 .o- - ly) - 1 . 4 - c I \ M -Y - gJ - 1 . 8 - - - 2 . 2 - Fig. 3. (a) Arrhenius plots for Cr/silica gel reduced in CO (623 K) for I5 min (@) and 120 min (H). (b) Arrhenius plots for Cr/silica-titania (4.2 YO Ti) cogel reduced in CO (623 K) for 15 min (@) and 60 min (m).0 60 120 180 reduction time/min 6 5 " 4 'm - I 53 -Y 2 1 reduction time/min 1 I I I 0 60 120 1 8 0 2 1 0 reduction time/min I 0 s b - reduction time/min Fig. 4. Polymerisation rate at 273 (H), 343 (a) and 373 K (0) over supported chromia catalysts as a function of reduction time in CO at 623 K: (a) 0, (b) 1.3, (c) 2.1, (d) 4.2 wt YO Ti. 0, Reduced 5 h, polymerisation at 273 K. 0, Reduced 12 h, polymerisation at 343 K.75 0 60 120 180 240 0 60 120 180 240 treatment time/min treatment time/min Fig. 5. Polymerisation rate at 273 K (m, 0) and 343 K (0, 0) for two catalysts reduced in CO (15 min at 623 K) followed by heating in U ~ C U O (0, W) or in CO (0, 0) at 873 K. (a) 0, (b) 4.2 wt % Ti.silica-titania (4.2 YO Ti) cogel reduced in CO for 60 min. A marked loss of activity was observed at low polymerisation temperatures (< 283 K). Polymerisation activities for four catalysts at various polymerisation temperatures are shown in fig. 4 as a function of reduction time. At 273 K a maximum activity was attained after a short reduction time (usually 15 min) at 623 K. Further reduction resulted in a loss of polymerisation activity at 273 K, with the loss being promoted by increased titanium content. Upon increasing the polymerisation temperature to 343 K a second maximum at ca. 60-120 min reduction time was observed, after which the activity began to decrease slowly. The appearance of bimodal activity curves with an increase in temperature are shown most clearly for 0 and 2.1 YO titanium.An overall reduction in activity was observed on increasing the polymerisation temperature to 373 K for 0 and 4.2% titanium. High-temperature evacuation or reduction in CO at 873 K of a standard reduced catalyst (standard reduced being defined as reduction in CO for 15 min at 623 K) was found to reduce the polymerisation activity at 273 and 343 K (fig. 5), the activity at 343 K passing through a maximum after an initial loss of activity when treated in uacuo. High-temperature treatment in CO caused a greater decrease in polymerisation activity than heat treatment in uacuo. Also a greater loss of activity at 273 K was observed for a catalyst with a higher titanium content. Discussion The first-order behaviour of the polymerisation rate on ethene pressure (at temperatures < 373 K) has been explained by a Rideal-Eley mechanism whereby polymerisation occurs by reaction of gas-phase ethene with the adsorbed polymer chain (or monomer) or by the Langmuir-Hinshelwood mechanism in which polymerisation proceeds by the reaction of an adsorbed ethene molecule with an adjacently adsorbed polymer chain or monomer. A comparison of theoretical predictions for both mechanisms with experimental observations indicates the Langrnuir-Hinshelwood mechanism to be the more probable mechanism6 for polymer growth.For such a mechanism the propagation reaction is normally envisaged as occurring in two steps,’ uiz. (1) reversible ethene76 ModiJied Phillips Catalyst : Kinetics adsorption at a vacant coordination site and (2) subsequent ethene insertion into the active Cr-C bond: transition state Above 373 K the reaction order showed a tendency towards second-order kinetics (fig.2). At increasing temperatures desorption of coordinated ethene molecules from the active site and a destabilisation of the transition-state complex is expected. In the case of destabilisation of the transition-state complex, insertion of the weakly complexed ethene into the growing polymer chain may be facilitated by the coordination of a second ethene molecule : Such behaviour would result in a second-order dependence of polymerisation rate on ethene pressure.8 The existence of a maximum in the dependence of polymerisation rate on temperature (fig. 3) has been reported before.Groeneveld et al.’ believe this to be a consequence of depolymerisation, although Clark’’ points out this is unlikely at such low temperatures (353 K). Zakharov et a1.l’ proposed that the maximum is a result of the presence of feedstock poisons which are adsorbed on the support at low temperatures, become mobile at higher temperatures and irreversibly deactivate the catalyst. It has been demonstrated that this poisoning mechanism is improbable.’ The dependence of catalytic activity on polymerisation temperature is complex in nature. Some of the features in the Arrhenius plots can be rationalised in terms of a Langmuir-Hinshelwood mechanism. At low temperatures a high surface coverage will be obtained and the rate-determining step will be the insertion of the coordinated ethene.With an increase in temperature the rate constant of the insertion reaction increases. However, when the temperature is high enough, desorption of coordinated ethene predominates and ethene coordination becomes rate-determining. This results in falling polymerisation rates, and thus a maximum is observed. In the temperature regions (up to 373 K) where these features are noted, the reaction is strictly first-order and no deviation towards second-order kinetics occurs. Thus the maximum polymerisation rates and positive activation energies are apparently due to changes in the kinetic parameters and not to a change in the polymerisation mechanism. At low temperatures a marked fall-off in polymerisation activity is observed (fig. 3). This is thought to indicate a possible change in the initiation step of the polymerisation mechanism.12 Low- temperature stabilisation of the Cr-C bonds in the possible initial coordination complexes2 may reduce the ability for rearrangement into a propagation centre (P), this rearrangement being a prerequisite to propagation.S.J . Conway, J . W. Falconer and C . H. Rochester 77 For polymerisation at 273 K the activity of the catalysts passed through a maximum with increasing reduction time (fig. 4). Upon increasing the polymerisation temperature at 343 K, catalysts which were previously inactive began to display increased activity. At 343 and 373 K the activity curves were bimodal in nature. This is consistent with the presence of two active sites for which there are two possible alternative explanations.Case 1 Two oxidation states of active Cr are present. The passing of polymerisation activity at 298 K through a maximum with increasing reduction time in CO at 623 K has been reported. Beck and Lun~ford'~ correlated the polymerisation activity at 298 K with CrI'I concentration. Later Myers and L~nsford'~ found the catalytic activity at 298 K to be inversely proportional to Cr" concentration as indicated by the chemiluminescent intensity associated with the reaction of oxygen with supported Cr". Upon re-evaluation of their catalyst, Myers and Lunsford12 found that the Cr" catalyst which was inactive at 298 K began to exhibit polymerisation activity as the temperature was increased (> 303 K). They concluded that extensive reduction of the Phillips catalyst results in reduction of Cr"' to Cr", and hence causes a change in the nature of the active polymerisation site.Wittgen et al.15 pointed out that CO reduction of chromate surface species can result in the formation of chromium with even valencies, IV and 11. In contrast, dichromate reduction can be expected to result in the formation of valencies from v to 11. For a mechanism involving the reduction of a dichromate species the concentration of CrI'I is expected to go through a maximum when Crrv species are totally reduced, while the concentration of Cr" should steadily increase. Although CO is normally assumed to be a two-electron reducing agent, without a more detailed knowledge of the surface electrochemistry and surface species, it is not possible to assign either chromate or dichromate species as the precursor of the active site.Case 2 Two types of active Cr" are present. Several types of Cr" on CO-reduced catalysts have been reported by different groups, although not all are believed to be active. Three types of Cr", each varying in their coordination number, have been identified using spectrosopic and calorimetric techniques. 16. l7 Other analytical techniques have identified at least two main types of Cr11.'E-20 McDaniel and Welch2' proposed two types of Cr", one producing low-molecular-weight polymer, the other producing polymer of high molecular weight. Merryfield et aZ.22 reported that Cr" becomes deactivated by reduction at high temperatures (> 673 K). This deactivation is not due to a change in oxidation state since the average state remains constant at 2.0.CO chemisorption experiments showed that the amount of chemisorbed CO decreases with increasing reduction temperature (> 623 K). This loss of coordinative unsaturation is thought to be due to a structural rearrangement of the chromium. Such a phenomenon may occur on prolonged reduction at 623 K (fig. 4). It is feasible that a mechanism based on two types of active Cr", each varying in their rate of formation and deactivation, may operate. The incorporation of titania appears to promote either the deactivation or reduction78 Modijied Phillips Catalyst Kinetics of the low-temperature active site (fig. 4). At high polymerisation temperatures high titania (4.2 % Ti) concentrations reduce the activity of the low-temperature active site.Plausible causes of these effects include a change in surface geometry (table 1) or the generation of Lewis-acid sites by titania, which may affect the physical and/or chemical properties of the active chromium. The possibility of Ti-0-Cr bonds being responsible for these effects cannot be discounted. Heating active catalysts at high temperature (873 K) in uacuo or in CO deactivates the catalyst towards ethene polymerisation (fig. 5). Such conditions have previously been observed to result in catalyst deactivation, which was attributed to loss of coordinative unsaturation by aggregation. Deactivation in CO is particularly severe as compared to deactivation in uacuo. This may be due to the formation of thermally stable Cr-carbonyl complexes,23 which allow the Cr greater surface mobility, enabling the Cr to rearrange itself on the surface in an inactive form.Titania also appears to influence the stability of the low-temperature active site. Catalysts treated in vacuo pass through a maximum at 343 K. This is likely to be due to the formation of a new surface complex, possibly involving a change in surface coordination. l2 The existence of two types of active site, each possessing different kinetic parameters, together with the complexity of mechanism will contribute to the non-linearity of the Arrhenius plots. This would also influence the position of the maximum, depending on the relative concentrations of each active site. We thank the S.E.R.C. for a CASE Studentship, BP Chemicals for financial support, and Dr G.W. Downs for helpful discussions. References 1 D. R. Witt, in Reactivity, Mechanism and Structure in Polymer Chemistry, ed. A. D. Jenkins and 2 M. P. McDaniel, Adv. Catal., 1985, 33, 47. 3 T. J. Pullukat, R. E. Hoff and M. Shida, J. Polym. Sci., Polym. Chem. Ed., 1980, 18, 2857. 4 M. P. McDaniel, M. B. Welch and M. J. Dreiling, J . Catal., 1983, 82, 118. 5 R. E. Dietz, U.S. Patent 3887494, 6/1975. 6 A. Clark, Ind. Ing. Chem., 1967, 59, 29. 7 J. P. Hogan, J. Polym. Sci., 1970, 8, 2637. 8 W. Cooper, in Comprehensive Chemical Kinetics, ed. C. H. Bamford and C. F. H. Tipper (Elsevier, 9 C. Groeneveld, P. P. M. M. Wittgen, J. P. M. Lavrijsen and G. C. A. Schuit, J . Catal., 1983, 82, 77. 10 A. Clark, in Polymerisation and Polycondensation Processes Symp., Ind. Eng. Div., 155th National 11 V. A. Zakhorov, Y. I. Ermakov, L. P. Ivanov and V. B. Skomorokhov, Kinet. Catal., 1968, 9, 499. 12 D. L. Myers and J. H. Lunsford, J . Catal., 1986, 99, 140. 13 D. D. Beck and J. H. Lunsford, J. Catal., 1981, 68, 121. 14 D. L. Myers and J. H. Lunsford, J . Catal., 1985, 92, 260. 15 P. P. M. M. Wittgen, 16 B. Fubini, G. Ghiotti, L. Stradella, E. Garrone and C. Morterra, J . Catal., 1980, 66, 200. 17 G. Ghotti, E. Garrone, G . Della Gatta, B. Fubini and E. Giamello, J . Catal., 1983, 80, 249. 18 H. L. Krauss, B. Rebenstorf, U. Westphal and D. Schneeweiss, in Preparation of Catalysts, ed. B. Delman, P. A. Jacobs and G. Poncelet (Elsevier, Amsterdam, 1976), p. 489. 19 P. Morys, U. Gorges and H. L. Krauss, Z . Naturforsch., Teil B, 1984, 39, 458. 20 H. L. Krauss and R. Hopfl, Proc. Eur. Symp. Thermal Anal., 198 1, 2, 175. 21 M. P. McDaniel and M. B. Welch, J . Catal., 1983, 82, 98. 22 R. Merryfield, M. P. McDaniel and G. Parks, J . Catal., 1982, 77, 348. 23 J. N. Finch, J . Catal., 1976, 43, 111. 24 M. P. McDaniel and T. D. Hottovy, J . Colloid Interface Sci., 1980, 78, 31. A. Ledwich (Wiley, Chichester, 1974), p. 43 1. Amsterdam, 1976), p. 155. Meeting A.C.S. (1968). C. Groeneveld, P. J. C. J. M. Zwaans, H. J. B. Morgenstern, A. H. von Heughten, C. J. M. von Heumen and G. C. A. Schmit, J. Catal., 1982, 77, 360. Paper 8/00722E; Received 23rd February, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500071

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J. Chew. SOC., Faraday Trans. I, 1989, 85(1), 79-90 Chromia/Silica-Titania Cogel Catalysts for Ethene Polymerisation Infrared Study of Nitric Oxide Adsorption Steven J. Conway, John W. Falconer and Colin H. Rochester* Department of Chemistry, The University, Dundee DDI 4HN, Scotland Infrared spectroscopy has been used to study the adsorption of NO on a chromia/silica-titania cogel (4.2 O h Ti) reduced in carbon monoxide. A mononitrosyl surface complex gave a band at 1807 cm-'. Four types of dinitrosyl complexes have been identified by pairs of infrared bands at I856/ 1742, 1868/ 1736, 1875/ 1755 and 1886/ 1766 cm-' and are ascribed to Cr'' sites of differing structural type. A Crl*I site also gave infrared bands at 1875/1755 cm-' due to a dinitrosyl complex. The effects of reduction conditions, chromia content, water, oxygen and ethene on NO adsorption have been investigated.Recent kinetic and polymer studies have indicated that the surfaces of chromia/silica gels and chromia/silica-titania cogels are heterogeneous. '* Heterogeneity could be attributed to several factors, which include the amorphous nature of the silica creating a broad distribution of adsorption site rea~tivities,~ the possible existence of chromate and dichromate structures of supported chromia4 and the formation of reduced chromia in different oxidation and coordination ~ t a t e s . ~ - ~ Infrared-spectroscopic examination of NO adsorption on reduced catalyst has proved useful for the characterisation of chromia catalysts, and several types of site which form dinitrosyl complexes have been identified.7 The present study was carried out in an attempt to gain further evidence about the oxidation and coordination states of chromium ions present on the surface, and to establish the nature of the catalytically active sites for ethene polymerisation.Experimental The catalyst, a chromia (1 YO Cr)/silica-titania (4.2 Yo Ti) cogel was identical to that used previously.' Infrared spectra of adsorbed NO (equilibrium pressure 133 N m-2) were recorded using a Perkin-Elmer model 580B spectrometer. Catalysts in the form of self-supporting discs were prepared by compacting ca. 70 mg of catalysts in a 25 mm die at ca. 10 MN m-2. The cell and apparatus used were similar to those described elsewhere.8 Following calcination in dry air (50 cm3 min-') at 873 K for 16 h and evacuation at 623 K for 30 min, the catalyst was then activated by reduction at 623 K in CO (5.33 kN mP2) for 15 min, with subsequent evacuation at 623 K for 5 min prior to cooling to beam temperature.This treatment is referred to as standard reduction. NO was purified by a distillation m e t h ~ d . ~ Results No infrared bands due to adsorbed NO were observed when either calcined catalyst with no reduction, or calcined and reduced support alone was exposed to NO. Fig. 1 ( a ) shows the infrared spectra of NO (equilibrium pressure 133 N m-2) adsorbed 7980 Modijied Phillips Catalyst : I. R. Study 10% I 7 1807 I 1 1856 ' /- I :.i I I, I I , , 1742 :a) I I I 2000 1900 1800 1700 60 3) 1742 I I 1 2000 1900 1800 1700 w avenumber/an- Fig.1. (a) Infrared spectra of NO adsorbed on a standard reduced catalyst: (-) background spectrum, (-O-) equilibrium pressure 133 N m-2, and after evacuation for 5 min (-. . - ) and 20 min (----) at beam temperature. (b) Infrared spectra showing the effect of chromium content on NO adsorption on standard reduced catalysts : (-O-) 0.4 % Cr, (--. - .) 1 .O % Cr and (----) 5.0% Cr. on a standard reduced catalyst. The predominant feature of NO adsorbed on a standard reduced catalyst is a pair of broad unsymmetrical bands with maxima at 1856 and 1742 cm-'. Bands of this type are normally assigned to the symmetric and asymmetric stretching frequencies of dinitrosyl complexes, Cr(N0),.7 Present as minor features are a weak band at 1807 cm-', which disappeared when the sample was evacuated, and a shoulder at 1868 cm-', which is responsible for the asymmetry of the 1856 cm-' band.The 1807 cm-' is assigned to a mononitrosyl complex, Cr(NO).' Fig. l(b) shows infrared spectra of NO in contact with three standard reduced catalysts containing 0.4, 1 .O and 5.0 YO chromium. Increasing the chromium con- centration resulted in an increase in intensity of the symmetric component of the dinitrosyl bands and its broadening to higher wavenumbers, a broadening of the asymmetric component and an increase in intensity of the 1807 cm-' band. Thermal desorption of the adsorbed NO species responsible for the bands at 1856 and 1742 cm-' bands was incomplete even at 473 K [fig. 2(a)]. The asymmetry of the 1856 cm-' band towards high wavenumbers was probably due to a band at 1870 cm-' [fig.1 (a)] which is assigned to an overtone of the Si-0 stretching mode for the catalystS. J. Conway, J. W. Falconer and C. H. Rochester 81 lO0/0 I 1742 la) 50 'i 1 I I , 4 1856 1742 - 1736 (b) 2000 1900 1800 1700 2000 1900 1800 1700 1600 wavenumber/cm-' Fig. 2. (a) Infrared spectra showing the effect of thermal desorption of NO adsorbed on a standard reduced catalyst: (-O-) after evacuation at beam temperature, and at (-- - - ) 433 K and (----) 473 K for 15 min. (b) Infrared spectra showing the effect of water on NO adsorbed on a standard reduced catalyst: (-O-) after evacuation at beam temperature for 15 min; (-a - .) to (-0-) after exposure to increasing water vapour pressures, and subsequent evacuation.support." High-temperature evacuation promoted loss of the shoulder at 1868 cm-' and a narrowing of the 1742 cm-l band, particularly on the low-wavenumber side. Adsorption of water on an NO-treated standard reduced catalyst removed the bands at 1856 and 1742 cm-l, revealing a second pair of bands at 1868 and 1736 cm-' [fig. 2 (b)]. A weak band indicative of adsorbed water was present at 1630 cm-l. Spectra recorded during a series of exposures to oxygen of an NO-treated standard reduced catalyst, and after evacuation for 5 min are shown in fig. 3(a) and (b). This resulted in the disappearance of the bands at 1868 and 1856 cm-', the appearance of a band at 1875 cm-' and a progressive shift of the maximum at 1742 cm-l to 1755 cm-' [fig. 3(a)]. This was accompanied by the appearance of a band at 1625 cm-l, the intensity of which increased with increasing exposure to oxygen [fig.3(b)]. When an NO-treated standard reduced catalyst was allowed to stand in NO for 12 h a shift in maxima from 1856 and 1742 cm-l to 1875 and 1755 cm-l [fig. 3(c)] and the appearance of an intense band at 1625 cm-' (not shown) was observed. Infrared spectra of NO adsorbed on catalyst reduced for various times [fig. 4(a)]1 0 O/e I 1875 Modified Phillips Catalyst : I.R. Study I 1742 a) 0 1900 1800 1700 I I I 50 50 50 50 50 50 50 1 i 0 1600 60 2c 1 I I . , 1856 1755 3) 0 1900 1800 1700 I I I wavenurnberIcn-' Fig. 3. (a) and (b) Infrared spectra showing the effect of oxygen on NO adsorbed on a standard reduced catalyst: (-O--) after evacuation at beam temperature for 15 min; (a) (-- - a ) to (-@-) and (6) (-a - - ) to (--..) after increasing time exposure to oxygen (1 atm) and subsequent evacuation. (c) Changes in the infrared spectra with time of NO adsorbed on a standard reduced catalyst : (-O-) immediately after contact with NO and (-- . .) after 720 min. exhibited only intensity differences of the 1856 and 1742 cm-' bands for reduction times up to 60 min. Increasing reduction periods (> 60 min) resulted in a shift of the maximum dt 1856 cm-' to 1870 cm-'. Exposure of standard reduced catalysts to ethene at 293, 343 and 413 K for 3 min before adsorption of NO [fig. 4(b)] resulted, except for catalyst treated at 41 3 K, in lower intensities of bands due to adsorbed NO than the corresponding intensities for catalyst which had not been exposed to ethene.The fall in band intensity was greatest for the catalyst treated at 293 K. The band at 1807 cm-' gave the biggest fractional decrease in its intensity. Spectra measured during a series of oxygen treatments of an ethene-treated standard reduced catalyst which had been exposed to NO are shown in fig. 4(c). Initial oxygen treatments resulted in loss of the 1856 cm-' band to give a pair of bands at 1868 and 1739 2 cm-'. Further exposure to oxygen reduced the intensity of both bands. Standard reduced catalysts further reduced in CO at high temperature (873 K) for various times and exposed to NO exhibited changes in band positions. Reduction for 2 h shifted the maxima from 1856 and 1742 cm-' to 1873 and 1746 cm-'.After furtherS. J . Conway, J . W. Falconer and C. H. Rochester 83 6( h $5 3 B 3 b 10% I 7; f 1870 - Ip 1742 1856 4 2000 1900 1800 1700 60 1 856 1742 I I I 2000 1900 1800 1700 wavenumber/an- ' 60 1868 1856 741 .. . 2000 1900 1800 1700 Fig. 4. (a) Infrared spectra of NO adsorbed on a catalyst reduced (CO, 623 K) for (-O-) 3, (- . . .) 15, (----) 60, (-@-) 180 and ( * - .) 720 min. (b) Infrared spectra showing the effect of ethene adsorption/polymerisation on NO adsorption on a standard reduced catalyst : (-O-) NO on a standard reduced catalyst, and after exposure to ethene at ( - - .) 293, (----) 343 and (--. . .) 413 K for 3 min. (c) Infrared spectra showing the effect of oxygen on NO adsorbed on a standard reduced catalyst pretreated with ethene (343 K) : (-O-) after evacuation at beam temperature for 15 min; (--.. - ) to (-@-) after increasing time exposure to oxygen, and subsequent evacuation. reduction at 873 K the bands moved to 1868 and 1739 cm-' and also suffered losses in intensity [fig. S(a)]. Fig. 5(b) shows spectra of NO adsorbed on a standard reduced catalyst which had been evacuated at high temperature (873 K) before admission of NO. Band intensities decreased with increasing treatment times at 873 K. The maxima also shifted to higher wavenumbers. For short treatment times the intensity of the band at 1807 cm-l was greater than the corresponding intensity for standard reduced catalysts. The changes with time in the infrared spectrum of NO adsorbed on a high temperature (873 K) evacuated catalyst are shown in fig.6(a) and (b). The spectra show the intensity of the 1807 cm-' band to decrease with time and the two major bands to shift to high wavenumbers. After 12 h two new bands with maxima at 1886 and 1766 cm-' were observed.84 Modijied Phillips Catalyst : I. R . Study I 1746 a) 10 1900 1800 1700 60 2c '- /--\ 1865 I 1745 3) 0 1900 1800 1700 I wavenumber/cm- Fig. 5. (a) Infrared spectra showing the effect of high-temperature reduction (873 K) in CO of a standard reduced catalyst on NO adsorption: (-O-) 120, (--. . a ) 360 and (----) 720 min reduction. (b) Infrared spectra showing the effect of high-temperature evacuation (873 K) of a standard reduced catalyst on NO adsorption: (-O--) 15, (--. a ) 180 and (----) 360 min evacuation. Discussion Infrared spectra of NO adsorbed on standard reduced catalysts [fig.l(a) and (b)] are characteristic of NO adsorbed on reduced silica-supported ~ h r o m i a . ~ The weak central band at 1807 cm-', easily removed by evacuation at beam temperature, is commonly assigned to a mononitrosyl specie^.^ Infrared and calorimetric investigations have shown that the sites characterised by the two more intense bands are the result of two- coordinate complex f~rrnation.~ Analogues of these complexes are also formed on reduced molybdena-alumina catalysts. l1 The infrared spectra of these surface complexes have been assigned to both dinitrosyl', '' and dimeric12* l3 complexes, with the balance of the present evidence currently favouring dinitrosyl complexes. l1 For a dinitrosyl complex the two bands are interpreted as being the symmetric and asymmetric stretching modes.7 The broad and unsymmetrical nature of the present infrared bands suggest NO was adsorbed on several types of different site.This conclusion is substantiated by the band resolution and enhancement and shifts induced by various catalyst treatments which were carried out.S. J. Conway, J. W. Falconer and C. H. Rochester 85 I 1744 4 60 1766 b) 2000 1900 1800 1700 2000 1900 1800 1700 Fig. 6. (a) Changes with time in the infrared spectra of NO adsorbed on a standard reduced catalyst evacuated for 15 min at 873 K: (-O--) immediately after contact with NO, and after (--. . .) 30, (----) 60, (-@-) 120 and (. * .) 240 min. (6) Changes with time in the infrared spectra of NO adsorbed on a standard reduced catalyst evacuated for 180 min at 873 K: (-O-) immediately after contact with NO and (--.- .) after 720 min. Thermal desorption of NO adsorbed on a standard reduced catalyst showed the complex characterised by the 1856/1742 cm-' band pair was thermally stable in comparison to other NO complexes [fig. 2(a)]. Krauss and H0pfl14 have shown, using temperature-programmed desorption, that NO is desorbed partially as NO, partially as N, and partially as N,O with accompanying oxidation of Cr". A desorption mechanism involving the decomposition of the Cr(NO), species and partial oxidation of chromium has also been rep~rted.'~ The interaction of water with NO adsorbed on a standard reduced catalyst revealed a pair of bands at 1868 and 1736 cm-l [fig.2(b)]. The interaction of propionitrile, pyridine, H,O, CO and NH, with NO adsorbed on reduced chromia-silica catalysts has been investigated previously. 15, l6 The original bands assigned to Cr(NO), were replaced by new pairs of bands which were attributed to the coordination of an extra ligand forming three-coordinate complexes. L86 Modified Phillips Catalyst : I.R. Study Such a mechanism appears unlikely for water, as evidence of the 1868 and 1736 cm-' pair was noted in the absence of water [fig. l(b), 4(c) and 5(a)]. An alternative explanation is a ligand displacement reaction. However, the formation of three coordinate complexes with molecules other than water cannot be discounted. Treatment of NO adsorbed on a standard reduced catalyst with oxygen resulted in the formation of a pair of bands at 1875 and 1755 cm-', with accompanying loss of all other bands [fig.3(a)]. In a similar experiment carried out by Krauss and Weis~er,'~ oxidation of adsorbed NO to adsorbed NO, occurred at low temperature, while at ambient temperature reoxidation of the catalyst resulted. As an absorption band at 1625 cm-' was observed [fig. 3(b)] similar to the band resulting from NO, adsorption on a reduced catalyst, it can be assumed that some NO oxidation had occurred at beam temperature. These absorption bands were found for a standard reduced catalyst after exposure to NO for 12 h (band at 1625 cm-l not shown). The ingress of oxygen into the system, which would be capable of oxidising chromium as well as adsorbed NO, may account for the shift in maxima.8 In contrast to oxygen-treated standard reduced catalyst exposed to NO no 1875 and 1755 cm-' bands were observed for such catalysts pretreated with ethene [fig.4(c)]. Absorption bands corresponding to the 1868/1736 cm-l were observed, suggesting the site attributed to these bands was oxidised to the site characterised by the 1875/ 1755 cm-' pair. An adsorption site giving rise to a mononitrosyl species has been shown to coordinate an additional NO, forming a dinitrosyl which gives infrared bands at 1887 and 1765 cm-l.' Similar bands (1 886/ 1766 cm-') attributable to this fourth site were observed when a catalyst heated in vacuo at 873 K was exposed to NO for 12 h [fig. 6(b)]. Three types of adsorption sites which form dinitrosyl species exhibiting average stretching frequencies [$( vsym + vasym)] at 1806, 18 17 and 1826 cm-' have previously been identified.7 These values are close to these observed in this study (i.e.1799, 1802, 1815 and 1826 cm-l). This is normally taken to indicate that the formal charge is increasing along the series, which is in disagreement with the evidence that CO-reduced catalysts contain primarily Cr" with small amounts of unreactive Cr''' as a-chromia.6 This was explained by assuming that the 'actual' charge depends on the number of surface oxygens coordinated to the divalent chromium, the oxygen ligands acting as strong electron acceptors . The relationship between average stretching frequency, chromium electron density and dinitrosyl behaviour can be discussed in terms of a metal-nitrosyl bond consisting of a 0 bond between the nitrogen lone pair and an empty orbital of the chromium, together with a n back donation from the occupied metal d orbitals to the unoccupied n* orbitals of N0.16 A high electron density on the chromium would increase the n back- donation to the n* orbital of NO, weakening the N-0 bond and consequently shifting the frequency of the infrared bands to lower wavenumbers, while increasing the strength of the M-N bond.This could account for NO adsorbed on the 18561 1742 cm-' site being resistant to desorption as a consequence of a strong M-N bond and apparently easily oxidised due to a weakening of the N-0 bond, while NO adsorbed on the 1875/1755 cm-' site is more resistant to oxidation. The small variation in average stretching frequencies between the 1856/ 1742 cm-' and 1868/ 1736 cm-' species suggest a similarity in electron density on the chromium. Two possible alternatives are Cr(NO), complexes based on chromate (1) and dichromate-like (2) structures in oxidation state two.ON NO ON NO ON NO \ /S. J. Conway, J. W. Falconer and C. H. Rochester 87 The proposition that the 1868/ 1736 cm-' site can be oxidised to the 1875/ 1755 cm-l site could be envisaged as dichromate oxidation from oxidation state two to three. ON NO 9 ON NO ON NO \ / Cr-0- '0' ON-!h<I>Ir-NO - 0 V - 0- I -0 I I /////// /A///// This assignment of oxidation states is plausible since the average stretching frequency of the proposed Cr"' species is similar to an e.s.r.-active CrIII dinitrosyl (1 8 16.5 cm-I) identified by Beck and Lunsford,'* and the formation of Cr' and CrO is most improbable.l9 Cr" coordinated to three surface oxygens and characterised by N-0 stretching frequencies similar to the CrIII dinitrosyl was proposed to be present on the surface of CO reduced catalyst^.^ This site may account for the shift of the symmetric component to higher wavenumbers upon prolonged reduction [fig. 4(a)] and high- temperature (873 K) reduction [fig. 5(a)], as CrIII is unlikely to survive such conditions. Although the assignment of the 1886/1766 cm-' site to a high oxidation state of chromium would be appropriate, no evidence of oxidation occurring at high temperatures has been presented.6 Alternatively Cr" coordinated to four surface oxygens characterised by a high average stretching frequency, indicating a low electron density, has been p r ~ p o s e d .~ The conditions and the nature of the reduced chromium would favour the formation of this site. A highly coordinatively unsaturated chromium ion bound to the surface would become mobile at high temperatures, enabling the chromium to rearrange itself in a different coordination state, possibly in a vacancy site on the surface. The assignments of absorption maxima to differing Cr sites on the catalyst (table 1) permit an interpretation of fig. 1 (b). The main features of this figure are: (a) a broadening of the symmetric stretching component to higher wavenumbers, (b) a broadening of the asymmetric stretching component and (c) an increase in mononitrosyl absorption with increasing Cr concentration.Hogan20 showed that as the Cr concentration of a catalyst was reduced the efficiency per Cr for ethene polymerisation increased. This was thought to arise from there being a limited number of locations on the silica surface which generate high-activity sites.2' It is suggested that the order in which the sites are formed is first the tetrahedral Cr" and then later the dichromate species, and possibly Cr" coordinated to three surface oxygens. The formation of dichromates would be expected to be favoured over chromate at higher Cr loadings where the Cr to surface hydroxyl ratio is increased and the Cr atoms are closer to one another. The increase in mononitrosyl at higher Cr loadings may be due to the formation of a site (or sites) not energetically or sterically favoured at lower loadings.Recent evidence suggests at least two types of site are active for low- and high-pressure ethene polymerisation, the relative concentrations of which are dependent upon CO reduction time.''2 Both Cr" and C P ' are thought to be catalytically active, extensive reduction being responsible for a change in the dominant active site from CrII' to Cr'1.2,5 This reduction results in a change in the nature of the active site, whereby the active site changes from one active at low temperature to one which is active only at higher temperatures. Both active sites are deactivated by heating in vacuo or in CO at high temperature (873 K). The activity at 343 K of a catalyst treated in vacuo decreases initially then increases slightly with treatment time before falling off further.This maximum was thought to indicate the formation of a new type of site active only at high polymerisation temperatures.88 Modijied Phillips Catalyst : I.R. Study Table 1. Summary of infrared band assignments for NO adsorbed on chromia/silica-titania catalyst (4.2 wt % Ti) structure of oxidation structure of coordination band position/cm-' complex state site number 1807 mononitrosyl 2 chromate a 1856/ 1742 dinitrosyl 2 chromate 2 1868/ 1736 dinitrosyl 2 dichromate 2 1875/ 1755 dinitrosyl 2 chromate 3 1875/ 1755' dinitrosyl 3 dichroma te 3 1886/ 1766 dinitrosyl 2 chromate 4 ~~ ~~~~ a Coordination number thought to be variable.' ' Oxidising conditions favour Crrrr formation, which in the presence of NO is believed to give bands similar to those for one of the Cr" complexes.Supporting evidence for this suggestion has been reported by Ghiotti et af. Extensive reduction [fig. 4(a)] resulted in a shift of the symmetric component to higher wavenumbers and a loss in intensity of both components consistent with the loss of the site giving bands at 1856/1742 cm-' in the presence of NO. The loss of these bands assigned to NO liganded to tetrahedral Cr" is not necessarily identified with the loss of the low-temperature active site for several reasons. A comparison of experimental observations with model predictions for polymer characteristics2 and a direct correlation of [Cr(N0)J3+ concentration detected by e.s.r. spectroscopy with low-temperature activity'* suggest that Cr'II is the low-temperature active site.If Cr"' and Cr" coordinated to three surface oxygens, which are thought to exhibit very similar absorption bands,7' l8 are formed, then the inability to discriminate between the two sites using infrared spectroscopy would compromise their correlation with activity. Infrared spectra of NO adsorbed on a standard reduced catalyst following exposure to ethene at 293 and 343 K indicated all sites were affected by ethene pretreatment, while exposure at 413 K showed little change within experimental variation [fig. 4(b)]. Strong ethene chemisorption on inactive and active sites occurs on the reduced catalyst,22 therefore the loss of bond intensity may be partly due to ethene chemisorption preventing coordination of NO to the chromium.The changes in bond intensity are consistent with the tendency of chemisorbed molecules to be desorbed with increasing temperature. Comparison of NO adsorbed on a standard reduced catalyst treated with oxygen [fig. 3(a)] and one pretreated with ethene prior to exposure to NO and oxygen [fig. 4(c)] showed that no 1875 or 1755 cm-' bands were present in the latter case, while bands corresponding to the 1868/ 1736 cm-' site were observed. Although this indicates a change in the nature of the 1868/ 1736 cm-' site it does not mean necessarily that it was the active site. The reason for the change is also unknown. High-temperature evacuation of CO reduced catalysts has been studied before.6, ' 9 23 A decrease in the concentration of the site assigned to tetrahedral Cr" accompanied by a decrease in sorptive capacity for CO and reduced reactivity of Cr" towards reoxidation was interpreted as a loss of polymerisation activity and tetrahedral Cr" assigned as an active site.However, Myers and L u n ~ f o r d ~ ~ found that their catalyst was more active, and upon NO adsorption observed that the concentration of the dinitrosyl previously assigned to tetrahedral Cr" was reduced and the mononitrosyl concentration was increased. They concluded that the active site on Cr" catalysts is a Cr" moiety characterising mononitrosyl formation. We previously reported that after an initial fall in activity an increase in activity occurs before the activity falls off with increasing treatment time.' The maximum was attributed to the formation of a new type of site.Spectra of adsorbed NO [fig. 5 (b)] indicate an increase in mononitrosyl concentrationS . J. Conway, J. W. Falconer and C. H. Rochester 89 and fall in dinitrosyl concentration, particularly the 1856/ 1742 cm-l dinitrosyl, suggesting that the site characterised by mononitrosyl formation is responsible for the increase in activity. Here a correlation between the fall in activity and the loss of the 1856/ 1742 cm-' site is observed; nevertheless this cannot be an unambiguous assign- ment. For ethene polymerisation proceeding by the Langmuir-Hinshelwood mechanism two vacant coordination sites are required. Ghiotti et aL7 identified a mononitrosyl which was capable of coordinating a second NO by an activated process thought to involve a ligand-displacement reaction in which at least one surface oxygen is detached, forming a dinitrosyl characterised by absorption bands at 1887 and 1765 cm-'.The changes with time of NO adsorbed on a catalyst evacuated at high temperature are consistent with the formation of a similar dinitrosyl [fig. 6(a) and (b)]. A comparison of high-temperature reduction and evacuation showed the former treatment to be more effective for catalyst deactivation. This was thought to be due to the formation of thermally stable CO-Cr complexes, allowing the Cr greater surface mobility and subsequent rearrangement into an inactive f0rm.l Consequently, the comparative stability of the site (1868/ 1736 cm-l) assigned to a dichromate species may be due to CO coordinated in such a way as to block vacant coordination positions and prevent rearrangement.Furthermore, dichromates have been suggested to have a good geometrical fit with the silica surface, which may promote thermal stability.21 Since only ca. 10 % of the total Cr is active on a catalyst containing 1 wt YO Cr3, probing of inactive sites is inevitable. Inactive catalysts deactivated by high-temperature reduction still exhibit intense bands [fig. 5(a)]. However, the presence of bands characteristic of the 1868/ 1736 and 1875/ 1755 cm-l sites on an inactive catalyst cannot be interpreted as meaning these sites are always inactive, as Cr species must be considered as families containing sub-species in which the behaviour of a species has to be considered as an average.6, 7 9 25 Welch and McDanie126 favour a deactivation mechanism involving the aggregation of Cr into aggregates of CrO or Cr20,, whereas othersGy7 favour one involving the coordination of surface oxides to the Cr.In both cases a loss of coordinative unsaturation occurs, but a comparison of fig. 5(a) and (6) suggest that differing mechanisms occur with high-temperature evacuation and high-temperature CO treatment. In fig. 5 (b) the significant absorption due to mononitrosyl species suggests that some coordination of surface oxides occurs during high-temperature evacuation. NO adsorbed on different coordinatively unsaturated Cr is characterised by specific absorption bands, showing NO to be sensitive to the nature of the Cr site. Dinitrosyl species typical of these observed on a reduced Cr/silica-titania (4.2 YO Ti) cogel were also observed on a reduced Cr/silica gel.This indicates that the same adsorption sites are present on both catalysts and no observable Ti-0-Cr bonding occurs on the cogel, as adsorption sites bonded to different support atoms would be expected to generate dinitrosyl species with different absorption ~pectra.~' We thank the S.E.R.C. for a Studentship, BP Chemicals for support through the CASE scheme and Dr G. W. Downs for helpful discussions. References 1 S. J. Conway, J. W. Falconer and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1989, 85, 71. 2 S. J. Conway, J. W. Falconer and C. H. Rochester, submitted. 3 M. P. McDaniel, Adv. Catal., 1985, 33, 47. 4 M. P. McDaniel and M. B. Welch, J .Catal., 1983, 82, 98. 5 D. L. Myers and J. H. Lunsford, J. Catal., 1985, 92, 260. 6 B. Fubini, G. Ghiotti, L. Stradella, E. Garrone and C. Morterra, J . Catal., 1980, 66, 200. 7 G. Ghiotti, E. Garone, G. Della Gatta, B. Fubini and E. Gianello, J. Catal., 1983, 80, 249. 8 M. A. Sutton, Ph.D. Thesis (Nottingham University, 1981).90 ModiJied Phillips Catalyst : I. R . Study 9 R. E. Nightingale, A. R. E. Downie, D. L. Rotenberg, B. Crawford Jr and R. A. Ogg Jr, J . Phys. Chem., 1954, 58, 1047. 10 D. D. Eley, C. H. Rochester and M. S . Scurrell, Proc. R. SOC. London, Ser. A, 1972, 329, 375. 11 R. P. Rosen, K. Segawa, W. S. Millman and W. K. Hall, J . Catal., 1984, 90, 368. 12 E. L. Kugler, R. J. Kokes and J. W. Gryder, J . Catal., 1975, 36, 142. 13 E. L. Kugler and J. W. Gryder, J. Catal., 1975, 36, 152. 14 H. L. Krauss and R. Hopfl, Proc. 2nd Eur. Symp. Therm. Anal., 1981, p. 175. I5 A. Zecchina, E. Garrone, G. Ghiotti and E. Borello, in Catalysis, Heterogeneous and Homogeneous, ed. B. Delmon and G. Jannes (Elsevier, Amsterdam, 1975), p. 243. 16 E. Garrone, G. Ghiotti, S . Coluccia and A. Zecchina, J. Phys. Chem., 1975, 79, 984. 17 H. L. Krauss and B. Weisser, 2. Anorg. Allg. Chern., 1975, 412, 82. 18 D. D. Beck and J. H. Lunsford, J. Catal., 1981, 68, 121. 19 C. Groeneveld, P. P. M. M. Wittgen, A. M. van Kersbergen, P. L. M. Mestrom, C. E. Nuijten and G. C. A. Schmit, J . Catal., 1979, 59, 153. 20 J. P. Hogan, J . Polym. Sci., 1970, 8, 2637. 21 D. R. Witt, in Reactivity and Mechanism and Structure in Polymer Chemistry, ed. A. D. Jenkins and A. Ledwich (Wiley, New York, 1974), p. 431. 22 R. Merryfield, M. P. McDaniel and G. Parks, J . Catal., 1982, 77, 348. 23 A. Zecchina, E. Garrone, G. Ghiotti, C. Morterra and E. Borrello, J. Phys. Chem., 1975, 79, 966. 24 D. L. Myers and J. H. Lunsford, J . Catal., 1986, 99, 140. 25 H. L. Krauss, B. Rebenstorf, U. Westphal and D. Schneeweis, in Preparation of Catalysts, ed. B. 26 M. B. Welch and M. P. McDaniel, J . Catal., 1983, 82, 110. 27 S . J. Conway, unpublished results. Delom, P. A. Jacobs and G. Poncelet (Elsevier, Amsterdam, 1976), p. 489. Paper 8/01532E; Received 19th April, 1988

ISSN:0300-9599    

DOI:10.1039/F19898500079

出版商:RSC    

年代:1989

数据来源: RSC