摘要:

CF2 Emission During Vacuum Ultraviolet Photodissociation of CF2Br2 BY CHRISTOPHER A. F. JOHNSON" AND (IN PART) MRS. HILARY J. Ross Chemistry Department, Heriot-Watt University, Riccarton, Edinburgh EH14 4AS Received 22nd June, 1978 Photolysis of CF2Br2 at wavelengths shorter than 135 nm leads to the production of electronically excited CF2. At low pressure fluorescence is observed over the range 230-340 nm. Considerable vibrational excitation of the CF2 is evident. The CF2 fluorescence is quenched by H2, D2, CO and C02with rate constants of 5.6 x lo9, 2.7 x lo9, 4 x 1O'O and 3.2 x lo1* dm3 mol-' s-l, respectively. The ultraviolet emission spectrum of the CF2 radical has been examined by a number of workers since its original observation by Venkateswarl~.~ The main emission system lies between about 240 and 340 nm, and has also been examined in absorption under high resolution by MathewsY6 the transition being assigned lB1-'A 1 .Bands at wavelengths longer than 350 nrn have been noted by some groups, and it is possible that some or all of these bands are due to the 3B1-1A1 transition of CF2.3 Gas phase emission studies of CF2 to date have involved electrical or microwave discharge excitation of low pressures of various fluorinated hydrocarbons, while matrix isolation studies employed photolysis of CF2N2 or CF2H2 as the source of CF2, followed by direct U.V. excitation of the matrix isolated CF,. We report the observation of electronically excited CF2 produced by vacuum ultraviolet photo- dissociation of CF2Br2. EXPERIMENTAL The apparatus and technique have been described previously.' Pressures (0-10 Torr, 0-1330 Pa) were measured with a capacitance manometer, and higher pressures (to atmos- pheric) with a resistance strain gauge transducer.Commerical CFzBrz was distilled before use and stored under vacuum. Owing to the weak signal and rapid accumulation of polymeric material on lamp windows, experiments were carried out under low resolution with slit widths of 0.75 mm (band pass w1.2nm) and scan rates of 5 or 10nmmin-'. During quenching experiments the slits were opened to 1.5 111111, and the monochromator set to the broad maximum in the emission (-260 nm). RESULTS AND DISCUSSION During photolysis of CF2Br2 at 121.6, 123.6, 129.5 and 130.6nm emission was observed at x230-350nm. None was observed on photolysis at 147nm (Xe/ sapphire).Although little structure could be resolved, the emission appeared to consist of many closely overlapping bands. The maximum intensity occurred at a shorter wavelength with increasing energy of the incident radiation, from x280 nm at 129.5nm to x255 nm at 121.6nm. There was also a small shift to the blue in the short wavelength limit (< 10 nm) with the higher energy radiation. We assign the emission to CF, for the following reasons. (i) The emission has all the general features of wavelength range and intensity distribution of CF2('B1-'A 1) 2930C. A . F. JOHNSON AND H . J . ROSS I l l 293 1 wavelengthlnm FIG. 1.-Lower trace: emission observed (uncorrected for system response) during the 123 nni photolysis of 0.06 Torr CF2Br2. Upper trace : emission observed during 123 nm photolysis of 0.08 Torr CF2ClN0. Some bands due to NO are also produced.The " baseline " of the upper curve is indicated. i I I I I I I I I 380 360 340 320 300 280 260 210 220 wavelength/nm FIG. 2.-Lower trace : emission observed during the 123 nm photolysis of 0.06 Torr CFzBr2. Upper trace : emission observed during the 123 nm photolysis of 0.06 Torr CF2Br2+720 Tom He. The wavelengths of some known CF2 bands are indicated.6 The intensities of the two spectra are not directly comparable.2932 emission observed by others. (ii) CF,Br, is known to give high yields of ground state CF, when photolysed at longer wavelengths.8 (iii) We obtain a very similar emission on irradiation of CF,ClNO at 123.6 nm, although in this case it is overlapped to some extent by bands due to NO.Only CF, is common to the two molecules. Spectra for comparison are shown in fig. 1. (iv) Addition of He at several hundred Torr causes some rotational and vibrational relaxation. Under these conditions some bands are more clearly resolved (fig. 2) and our estimated wavelengths are similar to literature data for CF2. There are two processes that could give rise to electronically excited CF,, reactions (1) and (2). u LTR A VIO LET P HOTODI sso CI AT ION OF CF,Br, CF2Br2 +hv -+ CF2(lB1) + 2Br[(4~~)~P;] CF2Br2 +hv + CF2(lB1) +Br,(XICT). (1) (2) Using the method and tabulations of Benson to estimate C-Br bond energies, we derive threshold wavelengths of w 135 and x 172 nm for processes (1) and (2) respectively, assuming CF2lB1 is produced in the v = 0 level.Allowing for the vibrational excitation observed by us in photolysis at 130 nm, we estimate from our data a threshold wavelength of 3 138 nm. We therefore conclude that it is reaction (1) we observe in these studies. Vibrational excitation in the bending mode is not surprising in view of the difference between the FCF angle in CF,, lB1 (122.3°)6 and that in the tetrahedral CF2Br, molecule. The short wavelength limit of the fluorescence implies vibrational levels are populated at least to v = 12 at the shorter photolysis wavelengths. 1.0 0.8 0.2 Y I I I I I I I 0.10 0.20 0.30 0.40 pressure/Torr FIG. 3.-Variation of fluorescence intensity with pressure of CF,Br,. The solid curve is a plot of the expression l2 with Pmax set equal to 0.072 Torr.This equation describes the intensity variation with pressure 4lIfmax = ( f ' / P , a x ) ~ X P (1 -f'/f'max) when quenching may be neglected. There are few literature data on the reaction of electronically excited CF2, although it is known that ground state CF2 is unreactive in comparison with, e.g., methylene. Dimerisation [reaction (3)] is the dominant process for removal of CF2 in most gas-phase studies, with k3 21 2.1 x lo7 dm3 mol-l s-l at 298 K.l0 Reaction of ground state CF2 with olefins or 0, is slow l 1 in comparison with reaction (3) CF2 +CF2 + CZF4. (3)C . A . F. JOHNSON AND H . J. ROSS 2933 We have carried out a limited number of quenching experiments using 123 nm incident radiation. For these measurements a wavelength of x260nin was used, with a band pass of w2.5 nm.Over the limited pressure range in which fluorescence was observed (pressures <0.4 Torr), no self-quenching was detected. Fig. 3 demon- strates that the data closely fit the behaviour expected on the basis of no quenching,12 the intensity variation being due entirely to absorption of radiation between the lamp window and the viewing region. Addition of helium at pressures up to 700 Torr resulted in a small reduction in fluorescent intensity at the monitoring wavelength ( ~ 3 0 % at 700 Torr). The complete spectrum obtained under these conditions indicated that the fluorescence maximum had shifted to longer wavelengths with the total intensity unchanged. We assume this to be a consequence of some rotational and vibrational deactivation of the emitting CF,.In all other quenching measurements we assume rotational/ vibrational deactivation of the same magnitude as with helium, and the derived rate constants are corrected accordingly. Quenching half-pressures were determined from Stern-Volmer plots for H2, D2, CO, CO, and N,, a constant pressure (0.05 Torr) of CF,Br, being used throughout. The experimental half-pressures are listed below. Rate constants were derived assuming a fluorescence lifetime for CF, IB1 of 31 ns. Allowance has been made where necessary for absorption of incident radiation by the quenching gas. D2 H2 co coz NZ P, /To r r 190 98 14.4 12.6 1130 k,/109 dm3 mol-1 s-l 2.7k0.15 5.6f0.3 40f2 32f3 <0.1 Quenching due to nitrogen is slow.Indeed, the observed " quenching " could be satisfactorily explained by a rate of vibrational deactivation of CF, 'B1 only about 20 % higher than that observed with helium. Therefore we emphasise that the derived k, for N, is an upper limit. Quenching by H,, D,, CO and C02 is very much faster, the calculated rate constants being two or three orders of magnitude larger than that for dimerisation of ground state CF, [reaction (3)]. The reactivity of electronically excited CF, is comparable with that of singlet methylene itself. For example, Laufer and Bass I3 have reported rate constants for reaction of lCH2 with NO, CO and CH2C0 of 2.4 x lolo, 5.4 x lo9 and 1.9 x 1O1O dm3 mol-l s-l, respec- tively. It would have been of interest to extend these quenching studies to a wider range of molecules.However the present technique limits the quenching molecules to those with small extinction coefficients at 123 nm. Besides the emission attributed to CF, lB1, we also observed weak emission centred at about 285 nm. This emission was very rapidly quenched by H2 or D,, and became sharper in the presence of He or NZ. It is probably the D(lX;)- B(3JI&J transition of molecular bromine, which has been reported in emission by Venkateswarlu and Verma,14 and in absorption by Briggs and Norrish.lS As the D state of Br, lies at 51 800 cm-l above the ground state, we estimate the threshold for process (4) CF,Br, +hv -+ CF,(lAl)+Br,(DIZl) (4) to be at about 137 nm. This value is consistent with our observations of the emission in photolysis at 121, 123 and 130 nm but not at 147 nm.This emission occurs over a relatively narrow range of wavelength, implying very little vibrational excitation of the Br,D state (the B state has a very shallow minimum). It should be noted that the incident radiation is of insufficient energy both to excite Br, to the2934 ULTRAVIOLET PHOTOD I SSOCI A T 1 0 N OF CF2Br, D state and CF2 to the lB1 state at the same time. Therefore process (4) is in competition with process (1) and is apparently of less importance. C . E. Smith, M. E. Jacox and D. E. Milligan, J. Mol. Spectr., 1976, 60, 381. L. Marsigny, J. Ferran, J. Lebreton and R. Lagrange, Compt. rend., 1968, 266, 9, 507. Quach-Tat-Trung, G. Durocher, P. Sauvageau and C. Sandorfy, Chern. Phys. Letfers, 1977, 47,404. 4V. E. Bondybey, J. Mol. Spectr., 1976, 63, 164. P. Venkateswarlu, Phys. Rev., 1950,77, 676. C . W. Mathews, Canad. J. Phys., 1967, 45, 2355. J. P. Simons and A. J. Yarwood, Trans. Faraday SOC., 1961,57,2167. S . W. Benson, Thermochemical Kinetics : Methods for the Estimation of Thermochemical Data and Rate Parameters (Wiley, New York, 1968). F. W. Dalby, J. Chem. Phys., 1964,41,2297. ’ C. A. F. Johnson, V. Freestone and J. Giovanacci, J.C.S. Perkin 11, 1978, 584. lo W. J. R. Tyerman, Trans. Faraday SOC., 1969, 65, 1188. l2 D. W. G. Style and J. C. Ward, Trans. Faraday SOC., 1953,49, 999. l3 A. H. Laufer and A. M. Bass, J. Phys. Chem., 1974,78,1344. l4 P. Venkateswarlu and R. D. Verma, Proc. Indian Acad. Sci. A, 1957, 46, 251. l5 A. G. Briggs and R. G. W. Norrish, Proc. Roy. SOC. A , 1963, 276, 51. (PAPER 8/1153)

ISSN:0300-9599    

DOI:10.1039/F19787402930

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Infrared Spectra of NO, N20, NO2 and 0 2 Adsorbed on SiO Films BY AVIVA LUBEZKY AND MORDECHAI FOLMAN" Department of Chemistry, Technion Israel Institute of Technology, Haifa, Israel Received 28th February, 1978 Infrared spectra of nitrogen and oxygen adsorbed on high surface area SiO films were investigated at different temperatures and surface coverages. On adsorption of NO at 77 K and at low coverage, decomposition of the molecule takes place with simultaneous oxidation of the surface. At higher coverages N20 is formed. On partially or completely oxidized surface absorptions due to NO dimers are obtained. On adsorption of NO2 at 136, 250, 273 and 300 K on freshly prepared SiO films decomposition of the molecule takes place with formation of NO and here again complexes with the surface are formed which are similar to those obtained when NO is adsorbed. SiO is the stable oxide form of Si in the gaseous phase at high temperature. For many years it has been claimed that in the solid form there is no SiO phase, rather Si + SO2.However, in the past few years there has been a general agreement that it is possible to obtain films of SiO that are not a mixture of Si+Si02, if the vaporization is carried out quickly and in a vacuum of 10-5-10-6Torr or in an inert gas atmosphere. These films are amorphous and the Si is not bonded, as in SO2, in the known tetrahedral configuration of quartz.2 SiO has many uses, but till now we know of no work where the adsorptive properties of this compound have been investigated. In the present work adsorption of some nitrogen oxides and oxygen on SiO high surface area films was investigated. EXPERIMENTAL The apparatus and experimental procedure have been described in detaiL3 The SiO films were obtained by fast vaporization of SiO from a tungsten spiral in an inert gas atmosphere (He at 0.150 Torr).The films showed good transparency in the 1200-4000 cm-1 region. In the range 400-1200 cm-1 two very wide and intense bands, known to be those of SiO', were obtained ; one centred at 1000 cm-l, the other at 700 cm-l. O2 and NO were adsorbed at 77 K and spectra were recorded at different temperatures up to room temperature. N20 was adsorbed at 196 K, NO2 at 250,273 and at 300 K. All the spectra were recorded with a model 521 Perkin-Elmer grating spectrophotometer. The calculated spectral slit was 1-1.5 cm-l.RESULTS AND DISCUSSION The frequencies obtained for the absorption bands of the different adsorbates on SiO are given in table 1. The phenomena connected with the adsorption of NO on SiO can be divided into two kinds: (a) the change which takes place in the absorption bands of the film; (b) the appearance of new bands as a result of adsorption. On introduction of NO at 77 K and at equilibrium pressures of 3 x 10-3-0.2 Torr the following changes in the film's absorption bands take place: The band at 1000 1-93 29352936 N AND 8 ADSORPTION ON SiO cm-l shifts to 1100 cm-l and the band at 700 cm-I shifts to 800 cm-l and, at the same time, a new band appears at 450 cm-1 (fig. 1). These shifts are not abrupt, rather a gradual shift of bands can be observed and a gradual growth of the new band at 450 cm-l.The bands at 1100 and 800 cm-l are typical of SO2. Therefore, on the introduction of NO, the first stage is an oxidation of the SiO film to Si02. In TABLE FREQUENCIES OF ADSORPTION BANDS OF NO, N20, NO2 AND 0 2 ADSORBED ON SiO FILMS frequency lcm-1 temperature/K absorption bands of SiO flms absorption bands of SiO films after O2 was adsorbed absorption bands of SiO films after NO was adsorbed additional bands of NO adsorbed on SiO films absorption bands of NO on raising the temperature absorption bands of NzO on SiO films absorption bands of NOz on SiO films 1000; 700 77 1100; 800; 450 77 1100; 800; 450 77 2227; 1862; 1815 1765; 1525; (1270)" 77 1760; 1655; 1295; 660 173 1750; 1655; 1295; 600 300 2222; 1622; 1275 196 2186; 1750; 1655; 1295 ; 660 * This band was seen in the experiment when the oxidation of the SiO film was not yet complete.order to prove this point, oxygen was adsorbed on a fresh film of SiO and it was found that the same spectral changes occurred (fig. 2): a shift of the two bands in the same direction and the appearance of an additional band at 450 cm-l. Oxidation of the film by NO or oxygen caused a decrease in transmittance by 20-30 %. The oxidation of the film by NO also was accompanied by an increase in pressure in the gas phase. 50 t \ '. '.. -.. . 5..4.' 1200 I000 800 600 400 frequency 1cm-l FIG. 1.-Oxidation of SiO by NO. (1) Background, (2) absorption of SiO film, (3) SiO film partially oxidized, (4) SiO fully oxidized.This seems to arise from the nitrogen which was released in the oxidation reaction: SiO +NO -+ Si02 + 3N2. No evidence of formation of surface nitride at 77 K was found (as judged from the spectrum). With adsorption of an additional amount of NO the following bands were obtained : at 1525, 1815 and 2227 cm-l [fig. 3(u)]. At higher pressures two extra bands appeared [fig. 3(b)l: at 1862 and at 1765 cm-I.A. LUBEZKY AND M. POLWAN 2937 These two bands grew with increasing amounts adsorbed and disappeared simul- taneously on evacuation at liquid nitrogen temperature indicating that they belong to the same physically adsorbed species. Owing to the large dead space of the cell and small quantity of the adsorbent, no measurements of the amount adsorbed could be made.The different experiments showed reproducibility as far as the number of bands and their frequencies are concerned. The only difference was the rate with which the film became oxidized. 60 50 - - 20 10 'id00 ' 1200 I000 800 600 400 - - I I I I I I I frequency/cm-' FIG. 2.-Oxidation of SiO by 02. (1) Background, (2) absorption of SiO film, (3) SiO film partially oxidized, (4) SiO fully oxidized. The interpretation of the spectra is hampered by the fact that the adsorbent SiO films undergo oxidation to different degrees and at different rates, and also by the fact that a part of the spectrum is screened by the SiO bands. The differences in the rate of oxidation of the various films can be explained as being due to differences in their specific areas. A great deal of effort was put into the preparation of these films under similar conditions like vaporization rate, temperature, pressure of the inert gas.Despite that the specific surface areas of these films might have differed to a certain extent. The gradual shift of the adsorbent peaks with the introduction of NO (and similarly with the introduction of oxygen) is due to the formation of intermediate oxides of the type Si01.8.4 It can be concluded from the spectra of NO adsorbed on SiO that the first stage is oxidation of the film with the evolution of nitrogen causing an increase in the pressure in the system. The large decrease in transmittance of the oxidized film is due to a change in the index of refraction when going from SiO to SO2. As already mentioned with additional amounts of adsor- bate, bands at 2227,1815 and 1525 cm-l appear.The band at 2227 em-' is ascribed to N20 which is formed as a result of the reaction 2N0 -+ N,O+O. ads ads The justification for this assignment is as follows : When in a series of experiments (at a temperature of 196 K and in a pressure range to 6 Torr) N20 was adsorbed on a fresh SiO film, a band at 2222 em-' was obtained, the frequency of which is very near to that of the band obtained in the case of NO adsorption (fig. 4).2938 N AND 0 ADSORPTION ON sio This band belongs to the assymmetric stretch of N20 which appears in the gaseous phase at 2223 cm-l. The i.r. spectrum of N20 shows two additional bands ; one belonging to the symmetric stretching mode which in the gas phase appears at 1285 cm-l, and the other to the bending mode at 588 cm-l.When N20 was adsorbed, the symmetric mode band at 1275 cm-I could be observed. For comparison purposes the adsorption experiments with N20 should have been performed at 77 K (the temperature at which the adsorption of NO was investigated). However, at that temperature the saturation pressure of N20 is very low and con- densation of the gas, rather than adsorption, could have taken place. Therefore 50 -ly I 7+= 2 90. 80- I 60 - 40 - - I I I I I I 1 I I t 4 0 2 3 ~ 2100 1900 1700 1500 frequency/cm-l FIG. 3.-(a) 1.r. spectrum of NO on SiO film at very low equilibrium pressures (1,2) ; 0,02-0.03 Torr. (b) 1.r. spectrum of NO on SiO film at higher equilibrium pressures (1, 2); 0.05-0.07 Torr. '"g"o--- :i- -__I---A .LUBEZKY A N D M . FOLMAN 2939 the adsorption experiments were performed at 196 K. This temperature difference does not influence our system to a larger extent, apart from a possible small shift in frequency. However, this region is usually screened when NO is adsorbed (fig. 1). In one instance when the oxidation of the film was slow, on adsorption of NO a band at 1270cm-I was observed which appeared simultaneously with the band at 2227 cm-l. Therefore, one can conclude quite certainly that the formation of N20 is a stage in the oxidation process of the surface by NO. In neither case was the bending mode band observed. The adsorption of N20 was reversible as judged from the ease of desorption. On pumping the system the absorption bands at 2222 and 1275 cm-1 disappeared (fig.4, curve 3). In the literat~re,~ when a band frequency of 2200 cm-I appeared on the adsorption of NO on oxides, it was attributed to NO+. This does not seem to be the case here (in addition to the above-mentioned experi- mental finding) since, on slowly raising the film temperature, the band disappears around 175K. This fact is difficult to account for, if the bonding between NO and the surface is in the form of a charge transfer with the formation of NO+. Adsorption of NO on SiOz was investigated by Kortum using reflection spectro~copy.~" A single absorption band at 2253 cm-l was found and has been assigned to NO adsorbed on strained oxygen bridges on the surface. The large shift in frequency (77cm-') as compared with the gas phase, was explained by a partial charge transfer mechanism.The band at 1815cm-I is shifted by 60cm-I to a lower frequency relative to the gas phase. Apart from the bands arising from oxidation, this band appears first and is the most intense. Bands of NO on oxides are known to appear in this frequency range. The different types of possible bonding have been described already in the past by Terenin et aL6 on the basis of what is already known for complexes of NO. The most probable form of bonding is surface)-N=O, that is, a double bond between the oxygen and the nitrogen, the frequency of which falls in this range. The additional band at 1525 cm-1 is in the low frequency range. Various oxygen nitrogen complexes absorb in this range. It is not possible to determine from a single-band which complex is formed The following types of complexes absorb 0 in this region : (1565-1500 cm-l) M-0 I M N = O (1520-1390 crn-l) where each type of complex has additional bands which cannot be seen owing to the absorption of the adsorbent itself.Therefore, in our case, it is not possible to completely characterize the complex. When the film is partially oxidized two additional bands appear on adsorption of larger amounts of NO at 77 K, in addition to the three bands already mentioned. The two bands appear at 1765 and 1862 cm-l, grow and disappear together with the2940 N AND 0 ADSORPTION ON SiO coverage and it can firmly be stated that they belong to the two stretching modes, sym- metric and assymmetric of the cis-dimer (NO),. The frequencies are near to those of the dimers in the liquid and solid phases in matrices and adsorbed phases lo which have been described in previous work.The bands of the dimers also appear when NO is adsorbed onto a film which is first oxidized completely by oxygen. 2200 2000 1800 1600 I400 frequency /cm- Fxo. 5.-1.r. spectrum of NO on SiO film on gradually heating from (1) 77, (2) 145 and (3) 175 K 20 20 frequency /cm-l FIG. 6.-(a) 1.r. spectrum of NO on SiO film at room temperature. (6) I. r. spectrum of NOz on SiO film at room temperature. Generally, aRer a certain period of time during which SiO was in contact with NO, it became completely oxidized and only bands belonging to dimers appeared while other bands disappeared. On the other hand, when the oxidation process was slower, it was possible to preserve a situation where the three bands at 1525, 1815 and 2227cm-1 still appear even though the film was partially oxidized.When at that stage the film is heated starting from 77 K, the spectrum begins to change (fig. 5 )A . LUBEZKY A N D M. FOLMAN 2941 until at about 175 K the following bands appear at 1760, 1655, 1295 and 660 cm-l. These bands are stable up to room temperature with the only change being a shift from 3760 to 1750 cm-l. When NOz was adsorbed at temperatures 250, 273 and 300 K a similar set of bands was obtained (fig. 6) and an additional band at 2186 cm-l. At the same time the film was slowly oxidized. It may be concluded that NO and NO2 adsorbed on a partially oxidized film form very similar complexes with the surface.The two similar situations with respect to the surface arise in the following manner : (1) With regard to NO, the surface is already oxidized and the frequencies of the NO bands shift as already mentioned. These bands are stable up to room temperature. (2) As regards NO2, we begin with a fresh film which is slowly oxidized by NO2 in the reaction NOz + NO+O and here again complexes with the surface are formed, which are similar to those formed when NO is adsorbed. We are unable to characterize the complex formed and to assign these absorptions. On comparing the band frequencies appearing in this work to those described in the literature it can be seen that the following complexes have absorption frequencies close to those in this work : ads 0-NO, - (1630, 1280, 860 and 760 cm-I) (1650-1600, 1225-1170 and 1000 cm-l).The remaining band at 660 cm-l is assigned to the Si-N bond. It is known that this bond gives rise to absorptions in the 570-710 cm-1 region.12 This research was supported by a grant from the United States-Israel Binational Science Foundation (B.S.F.), Jerusalem, Israel. G. Hass and C. D. Calvin, J. Opt. SOC. Amer., 1954,44, 181. J . A. Yasaitis and R. Kaplow, J. Appl. Phys., 1972,43,995. Y . Kozirovski and M. Folman, Trans. Faraduy SOC., 1966,62,808, 1431 ; 1969, 65,244. A. Cachard, J. A. Roger, J. Pivot and C . H. S. Dupoy, Phys. Stat. Sol., 1971,5a, 637. (a) G. Kortiim and H. Quabeck, Ber. Bunsenges. phys. Chem., 1969,73,1020 ; (b) G. Kortiim and H. Knehr, Ber. Bunsenges. phys. Chem., 1973,77, 85. (a) A. Terenin and L. Roev, Spectrochim. Acta, 1959,11,274,946 ; (6) A. V. Alekseev, V. N. Filimonov and A. N. Terenin, Doklady Akud. Nauk S.S.S.R., 1962,147, 1392. D. V. Pozdinakov and V. N. Filimonov, Adv. Mol. Relaxation Proc., 1973, 5, 55. W. A. Guillory and C. E. Hunter, J. Chern. Phys., 1971,50, 3516. lo A. Lubezky and M. Folman, Trans. Faraduy Sue., 1971,67, 3110. l1 (a) G. Blyholder and M. C. Allen, J. Phys. Chem., 1966, 70, 3110; (6) K. Nakamoto, Z.R. l 2 J. Goubeau, Chem.-Ztg., 1973, !?7, 123. * A. L. Smith, W. I. Keller and H. L. Johnston, J. Chem. Phys., 1951,19,189. Spectra of Inorganic and Coordination Compounds (Wiley, N.Y., 1966). (PAPER 8/370)

ISSN:0300-9599    

DOI:10.1039/F19787402935

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Application of Polymer Theory to Silicate Melts The System MO + MF, + SiO, BY CHARLES R. MASSON" AND WILLIAM F. CALEY Halifax, Nova Scotia, Canada B3H 321 Atlantic Regional Laboratory, National Research Council of Canada, Received 28th February, 1978 Expressions are derived for activities of MO and MF2 as functions of composition in MO + MF2 + SiOz melts. The treatment is based on the assumption that the melts consist of MZC cations and 02-, F- and an array of silicate and fluorosilicate anions of general formula Si,Ojn+ - mFL2nC2-m)- in thermodynamic equilibrium, where 1 Q n < 03 and 0 Q rn < 2n + 2. As in previous treatments, the equilibrium between the oxide and silicate ions is described in terms of an equilibrium constant k, the value of which is independent of the chain length n of the polyions.The equilibrium between the fluoride, silicate, fluorosilicate and oxide ions is described in terms of a second equilibrium constant k' the value of which is assumed to be independent of the degree of substitution of 0- by F on the silicate chains. Theoretical curves are shown for the activity of PbO as a function of com- position in PbO+PbF2+SiOz melts for the simple case in which k' = 0, i.e., for the situation in which PbFz acts solely as a diluent. Comparison with experimental data shows that the value of k' for this system is finite. Accurate determination of k' requires a knowledge of the activities of both MO and MF2. Knowledge of the value of k' allows the ion fraction of any species to be evaluated approximately in the range of composition for which the theory is applicable.In the application of polymer theory to silicate melts it is assumed that the simplest species of discrete silicate ion in the melt is the orthosilicate ion SiO$-, regarded as the monomer, and that these ions can undergo self-condensation with elimination of free oxide ions to yield an array of polysilicate ions in thermodynamic equilibrium. Application of this theory to binary melts MO + SiO, allows the average size and distribution of the silicate ions to be evaluated approximately from thermo- dynamic data in regions of composition where cyclic and network structures may be ignored. Except for the case of densities and molar volume^,^ which depend only on the mean chain length of the silicate ions and not on their size-distribution, extension of this treatment to ternary systems with more than one cation is difficult due to the competitive interactions between the anions and the various cations.This leads to significant deviations from Temkin's Law,4 assumed to hold for the binaries. In principle, however, it is possible to extend the treatment to ternary melts with a common cation. As an example of such a system we consider here melts of the ternary NO + MF2 + SO,. The treatment was developed primarily to gain some insight into the factors which govern the thermodynamic properties and constitution of fluoride-containing melts, although it is applicable, in principle, to other systems in which F- represents any univalent anion. The influence of fluoride additions on the properties of silicate melts has been studied by many investigators '-' Addition of relatively small amounts of CaF, to silicate melts markedly lowers their vis~osity.~-~ Kozakevitch showed that for acidic melts the effect is approximately twice that of CaO and suggested that monovalent F- ions, which have approximately the same and reviewed extensively.' 6-1 2942C.R. MASSON AND W . F . CALEY 2943 radius as bivalent 02- ions, can disrupt the -Si-0-Si- linkages in a silicate net- work, with formation of 02-, which can further disrupt the network : I I I I I I I I I I I I I I I I I I I I (1) (2) -Si-0-Si- + 2F- + -Si-F F-Si- + 02- -Si-O--Si- + 02- + -Si-O- + -0-Si- . Reaction (1) yields fluorosilicate structures wtih Si-F linkages. Bockris and Lowe l 9 proposed that F can replace 0- groups in silicate structures by depolymer- ization reactions of the type (3) with formation of discrete fluorosilicate anions.Direct evidence for Si-F bonding in quenched fluorosilicate melts was obtained by Kumar et aLIO from infrared absorption studies. The loss of fluorine as SiF4 from fluorosilicate melts has also been reported.l0. l1, Si3O$- + 2F- -+ Si03F3- + Si206F5- THEORETICAL It is assumed that the anionic constitution of a ternary MO+MF2+Si02 melt can be described in terms of 02-, F- and an array of silicate and fluorosilicate ions in thermodynamic equilibrium. For a complete description of the system it is necessary to consider all kinds of polyions which can arise by self-condensation of the monomeric units and their interaction with F- ions. This includes cyclic and net- work structures as well as linear and branched chains.Such a complete description is beyond our capability at present due to our inability to formulate exact expressions for the total number of moles or ions in the system when cyclic and network struc- tures are present, a limitation so far inherent in polymer theory generally. In the present treatment we confine attention to the case in which only linear chains are allowed. Although an oversimplification, this facilitates enormously the evaluation of theoretical relationships. As shown previously 20* 21 this approach also provides a reasonable description of the thermodynamic properties of many binary systems at low silica contents (Xsioz < 0.5).For this simple case the silicate ions have the general formula Si,O:2,",+12)- (1 B n < 00) and the equilibrium between the oxide and silicate ions may be written for which the equilibrium constant k , expressed in terms of the ion fractions of the participating species, is given by* and is independent of the chain length n. Activities of MO calculated on this basis for binary MO + SiO, melts are given by the expression SiOt- + S i , O $ ~ ~ ~ ) - + Sin+ 10(2n+4)- 3n+4 + 02- (4) k = Nsi,+,o,,+,No/Nsio,Nsi,o,,+ 1 (5) 1 1 -- - 2+- - 1 Xsioz 1 - UMO 1 + U M O ( l / b - 1) where Xsioz is the mole fraction of SO2. * For simplicity, values of ionic charge are omitted in the expressions for ion fractions.2944 POLYMER THEORY OF SILICATE MELTS For ternary melts MO+MF,+SiO,, the silicate ions again have the general formula SinO'32,",fi2'-.Fluorosilicate ions are considered to arise by substitution of 0- groups by F groups on the silicate chains. The silicate and fluorosilicate ions thus have the general formula Sin03n+l-mF~n+2-m)- where 1 < n < 00 and 0 < m < 2n-1-2. The equilibrium between the silicate and oxide ions is again given by eqn (4). For the equilibrium between the fluoride, silicate, fluorosilicate and oxide ions we may write where the equilibrium constant k' is given by and, for linear chains, is again independent of n. As emphasized previously, the assumption that equilibrium constants for reactions of this nature are independent of the chain lengths of the reacting species must be regarded as an approximation which cannot be expected to hold rigorously for the smallest ions but which should become increasingly reliable as n increases.It is further assumed for simplicity in the present treatment that k' is independent of the value of m in eqn (8). This implies that the reactivity of the 0- groups is independent of their degree of substitu- tion by F groups on the silicate chains. Without these simplifications the treatment appears hopelessly complicated at present. F- +Sin03,+ I-mF~2nf2-m)- 3n-m m + l +02- (7) (8) = Si,O F ( 2 n + l - m ) - k' = NSinO~n-mFm+ iNOINFNSinOsn+i -mFm From eqn ( 5 ) it follows that N s i 2 0 7 = ( k N ~ i 0 s / N o W S i 0 4 ~ S i 3 0 1 0 = ( k N s i o 4 / ~ o ) ~ s i z o , = ( k N S i 0 4 / N 0 ) 2 N S i 0 4 Nsi&13 = ( k N S i 0 4 / ~ o ) ~ s i 3 0 1 0 = ( k N s i o 4 / N o > 3 N s i o 4 or, in general where From eqn (8) we obtain NSin03n + I = An- Nsio4 A = kNsio4/No.or, in general where From eqn (9) and (11) : which yields the ion fraction of any species in terms of Nsio4 if k, k', No and NF are known. For the sum of the ion fractions of anions of chain length n we have, from eqn (13) : Nsin03n + 1 - mFm - - BmNsin03, + 1 (1 1) B = k'NF/No. (12) (1 3) ~ S i , , 0 3 , , + 1 -mFm = An- lBmNSi04 2n+2 2n+2 1 - p + 3 (14) 1-B C NSin03n+l-,nFm = A"-1NSi04 C Bm = An-1NSi04 m=O m u 0C . R. MASSON AND W. F. CALEY 2945 Hence, for the sum of the ion fractions of all silicate and fluorosilicate anions : AS the sum of the ion fractions of all anions is unity so that, from eqn (15) and (16) : 1-No-NF = - which yields NSio4 in terms of k and k' if No and NF for the melt under consideration are known or can be determined experimentally. When k' = 0 (ie., when MF2 acts merely as a diluent) B = 0 in eqn (17) and the expression for Nsio4 becomes Resubstituting for A from eqn (9) yields (NSiO.Jk'=O = (1 -No-NF)(l - A ) which is identical with the expression derived previously [eqn (6), ref.(l)] for the binary MO + SiOz except for replacement of the term (1 -No) by (1 -No - NF). In the general case for which k' is finite, substitution for A from eqn (10) in eqn (17) yields a quadratic expression in Nsio4 whose solution is No(Nal - B5)2 i- K YZ{2N:( 1 + B5) + K YZ])* - (19) NJNO(1 -B5)+KY(1-B)(1 +B2)] 2KB2[N0(B3 - 1) - K Y( 1 - B)] Nsio, = where Y = I-NO-NF (20) and Z = (1 -B)(l- B2).(21) To complete the treatment it is necessary to express the composition of the melt in terms of the ion fractions of the individual constituents. We have : 8 From eqn (14) and (22) : 1 -B5 1-B7 1 - ~ 9 moles Si02 = Nsio4 - +2A&o, - +3A2Nsio4 - 4- ... 1-B 1-B 1-B Similarly it may be shown (see Appendix) that moles MO= N o + L [ Nsio 1 {-+1--}+-{ 1 B B5 -+l+L}] 1 (24) 1-B 1-A 1-A 1-B 1-AB2 1-AB2 1-B2946 POLYMER THEORY OF SILICATE MELTS and B 2 Eqn (23)-(25), in combination with eqn (19), are the expressions required to yield No and NF in terms of k , k' and the composition of the melt. The resulting equations, though cumbersome, may be handled with the aid of a computer. To compare the predictions of theory and experiment it is necessary to know the activities of both MO and MF2 in the system. It has been shown previously that for binary MO+Si02 melts the value of aMO is given approximately by the Temkin equation : For the binary PbO+PbF, it has been shown l5, 22* 2 3 that the activity of PbO, and hence the activity of PbF2, is also given approximately by the Temkin equation and this is substantiated by the results in the accompanying publi~ation.~~ Hence we may write, with some justification : Substitution for No and NF in eqn (19) and (23)-(25) and combination of the resulting expressions with the relationships aMO = No.(26) a,,, = N : . (27) moles MF2 X M F 2 moles SiO, X s i o z - - - and moles MF, X M F 2 moles MO XMo -- - yields aMO and aMF2 in terms of k, k' and the composition of the melt.The complexity of the resulting relationships precludes the derivation of theoretical activity against composition curves for all except the simplest cases. One of these is the situation in which k' = 0 and it is instructive to compare the predictions of theory with experiment for this simple case. MODEL WITH k' = 0 For this case B = 0 in eqn (23)-(25) inclusive and we have Substituting for Nsio4 from eqn (18) and A from eqn (10) in eqn (30) : NFNO ~ ( ~ - N ~ - N F ) ( N ~ + ~ ( ~ - N O - N F ) ) ' Similarly it may be shown thatC. R. MASSON AND W. F. CALBY 2947 For comparison with experiment, use has been made of recent experimental data 23 for activities of PbO in the system PbO+PbF,+SiO,. The value of k for this system has been determined previously as 0.196 from experimental data on the binary PbO+SiO,.The following procedure was used to derive the theoretical activity against composition relationships for this system for k’ = 0. 1.0 0.8 0.6 0.A 0.2 0 0 0.2 0.4 0.6 08 1.0 XPbO FIG. 1 .-Theoretical curves of activity against mole fraction of PbO in PbO+ PbF2 + Si021melts for various ratiosXpbFz/XSio2 = (a) 00, (6) 10, (c) 4, (d)2, (e) 1, (f) 0.5 and (g) 0. The curves correspond to the case k’ = 0, i.e., to the situation in which fluorosilicate ions are absent, and PbF2 acts solely as a diluent. Broken lines indicate compositions for which Xsi02 > 0.3 ; the theory is not expected to hold accurately in this region. k = 0.196. An arbitrary value of aPbO (0 < apbo < 1) was first chosen for a melt of selected ratio XPbF2/XSiO2.This value was substituted for No in eqn (31) and the value of NF calculated, using k = 0.196. The value of NF thus obtained was substituted, along with the value of No, in eqn (33) to yield xpbF2/xpbO. This fixed the composition of the melt. The procedure was repeated for other values of apbo to yield apbo as a function of x p b o for the selected ratio XpbF2/&02. The entire procedure was then repeated for other selected ratios of &bF2/&02 to yield the theoretical curves shown in fig. 1. The curve for &bF2/&02 = 03, which corresponds to the Temkin equation for the binary MO + MF2, is given by ~~Mo~xMF2,xs*o~ = 00 = X M 0 / ( 2 -&lo) (34) and is readily obtained by substituting (1 -No) for NF and (1 - X&) for XMF2 in eqn (33).The curves in fig. 1 predict that, if PbFz behaves solely as a diluent, the effect of replacing SO2 by PbF2 in the binary PbO + Si02 is to cause a slight lowering of the activity of PbO for melts with x p b o > 0.775. For melts with x p b o < 0.775 the2948 POLYMER THEORY OF SILICATE MELTS theory predicts that, at constant XPbO, apbo will be raised by the substitution of PbF2 for Si02 until it eventually attains its value in the binary PbO +PbF2. At Xp,, = 0.775, no change is predicted in the value of amO at all levels of PbF2. The com- position at this point is readily obtained by substituting for apbo from eqn (34) in 0' 0.4 0.6 0.8 I .O XPM) xpbo FIG. 2.-Comparison of theoretical curves for k' = 0 with experimental data for PbO + PbFz + SiOz melts.The points are interpolated from the data of ref. (23). XmFz/XSiOz = (a) 0.5, (b) 1, (c) 4 and (d) 10. As shown previously,1* the theory outlined above (with k = 0.196) describes activities in binary PbO + Si02 melts when Xsioz < 0.3 and it appears reasonable to expect that the theoretical curves in fig. 1 will also be applicable in this region. The range of applicability of the theory expected on this basis is shown by the bold lines in fig. 1. Broken lines indicate compositions for which XsiOz > 0.3.C. R. MASSON AND W. P. CALEY 2949 In form, the curves in fig. 1 closely resemble the experimental data 23 for the system PbO +PbF2 + SO2. In particular, the theory correctly predicts the point of intersection at XpbO = 0.775.A detailed comparison of theory and experiment for melts with various levels of XpbFz/XSiO2 is shown in fig. 2. The theory provides a good representation of the data for melts with XpbO 2 0.8 at all levels of XpbFz/&02. For these melts XsIoz does not exceed 0.14. Similarly, good agreement between theory and experiment is observed for melts with Xpb0 > 0.5 when XpbF2/&02 = 10. For these melts, Xsio2 is always < 0.05. For melts with higher silica contents, however, the theoretical curves for aPbo are lower than the experimental values and the discrepancy becomes greater as Xsio2 increases. These differences are well beyond experimental error or the uncertainties involved in interpolation of the data. Thus the simple theory with k’ = 0 does not adequately describe the effect of PbF2 on the activity of PbO in this system.For a more quantitative representation of the data a finite value of k’ is required. The effect of introducing a finite value of k’ will be to raise the theoretical values of aPbo due to the displacement of 0- groups by F groups on the silicate chains, with liberation of free 02- ions according to reaction (7). The value of k’ required to provide a more detailed test of the theory cannot be estimated from measurements of aPbO alone, due to lack of knowledge of the value of NF required in the computations. For this purpose, values of apbFz must be deter- mined. The experimental determination of apbF2 and a more detailed discussion of the constitution of PbO +PbF2 + SOz melts form the subject of a separate corn- munication.We thank Miss S. Taylor and Mr. J. Uher for assistance with the computations. C. R. Masson, Proc. Roy. SOC. A , 1965,287, 201. C. R. Masson, J. Amer. Ceram. SOC., 1968,51, 134. A. E. Grau and C. R. Masson, Canad. Metal. Quart., 1976, 15, 367. M. Temkin, Zhur. fiz. Khim., 1946, 20, 105. G. H. Herty, Jr., F. A. Hartgen, G. L. Frear and M. B. Royer, US. Bur. Mines Rep. Invest., 1934, 3232. F. Hartmann, Stahl u. Eisen, 1934,54,564 ; 1938, 58,1029 ; Arch. Eisenhiittenw., 1938,10,45. ’ L. Schwerin, Metals and Alloys, 1934,5,118. P. Kozakevitch, Rev. Metal., 1954, 51, 569. P. M. Bills, J. Iron and Steel Inst., 1963, 201, 133. lo D. Kumar, R. G. Ward and D. J. Williams, Disc. Faraday Soc., 1961, 32, 147. l 1 A. Mitchell, Trans.Faraday Soc., 1967, 63, 1408. l2 J. R. Michel and A. Mitchell, Canad. Metal. Quart., 1975, 14, 153. l3 I. D. Sommerville and D. A. R. Kay, Metal. Trans., 1971, 2, 1727. l4 G. J. W. Kor, Metal. Trans., 1977, SB, 107. l 6 M. W. Davies, Chemical Metallurgy of Iron and Steel, ed. B. B. Argent and M. W. Davies l7 B. J. Keene and K. C. Mills, N.P.L. Report Chem., no. 60, 1976. l9 J. O’M. Bockris and D. C. Lowe, Proc. Roy. Soc., 1954,226,423. 2o C. R. Masson, Chemical Metallurgy of Iron and Steel, ed. B. B. Argent and M. W. Davies (Iron and Steel Inst., London, 1973), p. 3 ; J. Iron and Steel Inst., 1972, 210, 89, 369, C. R. Masson, Glass ’77, Proc. XIth Int. Congr. Glass. Prague, ed. J. Gotz (CVTS-Durn Techniky Praha, 1977), p. 1 ; J. Non-Cryst. Solids, 1977, 25, 1.22 M. G. Frohberg, M. L. Kapoor and G. M. Merohtra, Special Electrochemistry (Proc. Inter- national Symposium, Kieve, U.S.S.R., 1972), vol. 2, p. 175. 23 A. E. Grau, W. F. Caley and C. R. Masson, Canad. Metal. Quart., 1976, 15, 267. 24 W. F. Caley and C. R. Masson, J.C.S. Faraday I, 1978,74,2952. P. Perrot and M. F. El Ghandour, Silic. ind., 1976,41, 407. (Iron and Steel Inst., London, 1973), p. 43. K. C. Mills, N.P.L. Report Chem., no. 65, 1977. (PAPER 8/374)2950 POLYMER THEORY OF SILICATE MELTS APPENDIX The following summations were used in the derivations : co 1 -- (0 < x < 1) I-x c xn-I - 11'1 co 1 n m=O For the derivation of eqn (24) and (25) we require to evaluate the summations 2n+ 2 2 n + 2 f c mfvsin03n+l-mFm and n = l m=O n = l m = O 2 c nNsin03,,+1-nlFm* Introducing the value of NSj,103n+l+n,Fm from eqn (13), the first of these is given by From eqn (A4) and (A5) Expanding the summation and using eqn (Al) and (A2) we obtain Similarly, the second summation is given by which, withIthe aid of eqn (A3) yields Again, expanding the summation and using eqn (Al) and (A2) we obtain CA 10) which is the required result. From eqn (15) and (A10) : B5 1 Nsio4 - - -[ { +l}--{-+I}]. 1-AB2 1-AB2 ( A l l ) 1-B 1-A 1-A We now use these relationships to derive the expressions for the number of moles of MO and MF2 in eqn (24) and (25). We first note that all the Si02 in the melt isC. R . MASSON AND W. F . CALEY 295 1 associated with the silicate and fluorosilicate anions. This yields eqn (22) and (23) directly. For polyions with O/Si > 2, the excess oxygen must be ascribed to MO. For polyions with O/Si < 2 the oxygen deficit required for SiOz must be derived from MO and this is reflected as negative terms in the summations. Thus we have A 6 From eqn (A7), (All) and (A13) we obtain moles MO= B which is the result in the text. write As the MF2 is associated solely with the fluoride and fluorosilicate anions we may Substituting for from eqn (A7) in eqn (A15) we obtain -- - B5 { +2+---!-}] (25) B (1-A)(l-B) 1-AB2 1-AB2 1-B which completes the derivations.

ISSN:0300-9599    

DOI:10.1039/F19787402942

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Thermodynamics and Constitution of Silicate Melts The System PbO + PbF, + SiO, BY WILLIAM F. CALEY AND CHARLES R. MASSON" Atlantic Regional Laboratory, National Research Council of Canada, Halifax, Nova Scotia, Canada B3H 3Zl Received 28th February, 1978 Activities of PbFz in PbO + PbF2 and PbO + PbF2 + Si02 melts with up to 40 mol % PbF2 were determined at 1073-1173 K by means of an electrochemical cell with CaF2 as solid electrolyte. The results, together with previous measurements of the activity of PbO, were interpreted in terms of a theory in which the anionic constitution of the melts is represented by 02-, F- and an equilibrium array of silicate and fluorosilicate ions of formula Sin03n+l-mFk2n+ 2 - m)-. The equilibrium constant k' for the formation of fluorosilicate ions by the reaction F-+Si,03n+l-mF~2n+2-m)' = Si,03,-,F(2n+ m + l 1 - d - +02- was estimated to be 0.440.025.This allowed the anionic distribution to be evaluated for basic melts. In addition to 02- and F-, the main anionic constituents in melts with XpbO/XSi02 > 2.5 and XpbFz < 0.25 are SO$-, Si03F3-, SiOzFt-, Si03F-, Si,O$- and Siz06Fs-, with decreasing amounts of more highly fluorinated species. The volatile constituent SiF4 is also present in small amount and becomes increasingly significant as Xsioz/Xpbo is increased. Substitution of PbO by PbF2 causes a net decrease in the mean chain length of the anions. In the preceding communication theoretical expressions were derived for activi- ties of MO and MF, as functions of composition in ternary silicate melts MO+ MF, + Si02 whose anionic constitution was considered to be represented in terms of 02-, F- and an array of silicate and fluorosilicate ions of general formula Sin03n+l-mF(m2n+ 2 - m ) - in thermodynamic equilibrium.The purpose in the present study was to test the applicability of this theory for melts of the system PbO+PbF2 +SO,. Activities of PbO in such melts with up to 20 mol % PbF, have already been reported.2 Here we extend these studies to the measurement of activities of PbF2 over a similar range of compositions. Knowledge of these two activities allows an equilibrium constant k' to be derived for the equili- brium between fluoride, silicate, fluorosilicate, and oxide ions. This, together with an equilibrium constant k previously determined for the silicate-oxide ion equili- brium, provides a quantitative description of the system and allows the distribution of the polyions to be evaluated.The activities were determined by means of a concentration cell without transfer, with a solid CaF, electrolyte. The reliability of CaF, as an ionic conductor in cells of this type, in which the conduction is solely by F- ions, has been established in several investigation^.^-^ EXPERIMENTAL PREPARATION OF MELTS Yellow lead oxide (Fisher, " purified ") of stated purity 99.25 %, lead fluoride (B.D.H., " extra pure ") and silica, obtained by firing Mallinckrodt reagent grade silicic acid over- night at 1273 K, were used. Melts were prepared in Pt-20 % Rh crucibles by firing 50 g 2952W. F. CALEY AND C. R. MASSON 2953 of the mixed reagents in a preheated muffle furnace for 3 h ; firing temperatures were kept as close to the melting point as possible to avoid excessive volatilization. After firing, the melts were quenched between brass plates and crushed to a fine powder in an agate ball mill before use in the cell.CELLS The cells were of the type (A) The design is shown in fig. 1. A calcium fluoride crucible, M 1 cm i.d. x 3.8 cm long, prepared by slip-casting, was used as the solid electrolyte. This crucible contained the melt under investigation. A larger calcium fluoride crucible, M 1.9 cm i.d. x 4.1 cm long long, was used to contain the standard melt. The wall thickness of each crucible was approx. 1.6 111111. These crucibles were suspended concentrically in the hot zone of the furnace by means of an alumina pin H which passed through slots cut in the wall of each crucible and in the lower end of the refractory tube E.The standard melt was a liquid slag of composition 70 mol % PbO, 30 mol %lPbF2 for which the activity of PbF2 had first been determined in Ir,Ar,Pb~), me1 t 1 CaF2(&t andard melt, Pbg,Ar,Ir t o P o t e n t i o m e t e r TC FIG. 1.-Design of cell. A iridium wire, B serum bottle caps, C brass flange, D mullite tube, E ceramic tube to hold crucibles, F alumina tubes, G kanthal shield, H alumina pin, J PbO + PbFz + SiOz slag, K PbO + PbF2 standard slag, L liquid Pb, M CaFz crucibles, N pedestal, P swagelok fitting, Q alumina pushrod, R cooling water, TC thermocouple.29 54 THERMODYNAMICS OF PbO+PbF, +SiO, MELTS the same apparatus by using pure PbFz as the standard melt.NN 6 g slag and 4 g Pb were used in the inner crucible ; the outer crucible contained 12 g Pb and 12 g standard slag. The cell was flushed continuously with argon at atmospheric pressure. Tem. peratures were measured with a Pt/Pt-13 % Rh thermocouple located in the pedestal N- E.rn.f. values were measured by a Leeds and Northrup volt-potentiometer, and could be estimated to kO.01 mV. The cell was held in a platinum-wound resistance furnace controlled to f 1°C. FABRICATION OF CALCIUM FLUORIDE CRUCIBLES The technique used for fabricating the crucibles was similar to that described previo~sly.~ Absolute ethanol was used to suspend the particles rather than an acidic medium as reported elsewhere.O This simplified their manufacture. Fisher " certified " calcium fluoride, used to prepare the slips, was ball milled under absolute ethanol until 75 % of the particles were in the size range 1 to 4pm. The firing temperature was 1233 K, and the shrinkage was of the order of 10.5 %. Final density of the articles was 97 % of theoretical. To avoid formation of a Ca(OH)2 skin, the k e d crucibles were held in a desiccator until ready for use. Thermal stresses were minimal, as only 5 % of the crucibles were lost due to cracking on firing. Thermal expansion in the temperature range of the experiments (1073-1173 K) was of the order of 2 %. A.c. resistivity measurements on the final ceramics were in good agreement with published values [5 x cm-l at 923 K, 1592 Hz (present work) ; 8.5 x R-l cm-I at 923 K, 1592 Hzll]. ELECTRON MICROPROBE ANALYSES At the conclusion of each experiment the cells was sectioned and an electron microprobe was used to examine the contents and the walls of the crucibles.Analyses were performed for calcium in the melt and lead in the crucible. The results yielded supplementary informa- tion concerning the time and compositional range over which the CaF2 electrolyte could operate reliably in contact with corrosive melts of this nature. RESULTS CRUCIBLE-MELT INTERACTIONS The most severe attack of the CaF, electrolyte occurred during experiments with pure liquid PbF2 in the cell and it was for this reason that most of the work was done with a standard PbO+PbF2 melt which contained 30 mol % PbF2 as the refer- ence.Five separate experiments, varying in time from 5 to 10 h, were performed in which the e.m.f. of the standard slag was measured against pure liquid PbF,. Maxi- mum Ca found in the PbF, melt was 3 wt % at the crucible-melt interface; this decreased to a value of 1 % in the slag at a distance of 1 mm from the interface. Pb contamination of the crucible was of the order of 6 wt % at the crucible-melt interface, and 2.8 wt % at a distance of 0.5 mm inside the crucible. Interaction between the crucible and the standard slag resulted in 0 to 5.6 wt % Ca in the slag and 0 to 3 wt % Pb in the ceramic at the slag-crucible interface. In no case was Ca or Pb detected in the slag or electrolyte respectively at a distance of more than 0.5 mm from the interface. For the other melt compositions investigated, slag-crucible interactions occurred only in those samples which contained 20 mol % or more S O 2 or PbF2.Maximum Pb found in the ceramic at the ceramic-slag interface was 2 wt % (60 mol % PbO, 40 mol % PbF2 slag) ; maximum Ca in the slag was 1 % for the same sample. For all other slag compositions, only trace Pb and Ca contamination was found, with no contamination detected more than 0.5 mm inside the slag or ceramic phases. In order to minimize effects due to the crucible-melt reaction, experiments with a particular cell were not allowed to exceed 5 h, during which the cells gave resultsW. F. CALEY AND C. R. MASSON 2955 reproducible to within +1 mV. Temperatures were kept in the range 1097-1137 K. For all experiments reported, the e.m.f.was reversible with respect to temperature and independent of the flow rate of argon. Prolonged use of a cell above 1173 K eventually resulted in a drift of the e.m.f. of the order of 2 mV h-l, which increased with temperature and time. CALIBRATION OF STANDARD MELT The melting point of the standard slag was 913 K. The results of five separate experiments with the cell in the range 1097-1173 K yielded the following expression for the e.m.f., E : These values were reproducible to & 1.5 mV, with a standard deviation of 0.5 mV, when the experiments were performed in accordance with the limitations discussed in the previous section. The e.m.f. values reported below for cells C and D are the measured values, with the appropriate sign, plus the e.m.f.of cell B at the temperature of measurement, as given by eqn (1). All values have thus been corrected to yield the e.m.f. values which correspond to pure PbF2 as the reference melt. (-)Ir,Ar,Pb(l,, standard melt ]CaF,olPbF2(l,,Pb(l,,Ar,Ir'f' (B) E(mV) = -36.53+0.10T (1) THE SYSTEM PbO f PbF2 The results of experiments with the cell Ir,Ar,Pb(l,, (PbO + PbF2),l,ICaF2(s,l(standard melt),Pb(l,,Ar,Ir (C) corrected for the e.m.f. of the standard melt by eqn (1) are shown in fig. 2. The maximum PbFz content for which reliable data could be obtained was 40 mol % ; greater amounts of PbFz resulted in excessive crucible-slag interaction. The equations for each mixture together with their statistical data are given in table 1. TABLE ~.-E.M.F. DATA FOR THE SYSTEM PbO+PbFz (CELLS B AND C) E = A+BTmV standard maximum XPbO XPbFz A B deviationlmv deviationlrnv 0.95 0.05 - 2.7865 0.2172 0.4 1 .o 0.90 0.10 -25.3403 0.1741 0.1 0.2 0.80 0.20 - 32.3270 0.1260 0.1 0.5 0.70 0.30 - 37.0737 0.1020 0.5 1.3 0.60 0.40 - 47.9320 0.0907 0.2 0.4 Activities of PbF2 were calculated from the data by means of the relationship : RT E = - 2F In aPbFz where E is the combined e.m.f. of cells B and C , T is the temperature and 2F is the Faraday constant.These are plotted in fig. 3 as a function of the PbO content at 1173 K. In order to provide an approximate check on the PbF2 activities, a Gibbs-Duhem graphical integration was performed on the PbO activities previously obtained with CaO+ZrO, solid electrolyte cells for the same system., The integra- tion was in terms of activity coefficients,12 and the results are given in fig.3 for com- parison. The excellent agreement between the calculated activities and those deter- mined from the present data provides a check on the reliability of the cells.2956 250 200 > a Q-i E 6 ' 1 150 I I I I I I 1 I - *PbFZ= B+o*O-OO: 0..05 -0- - - - - - 0.10 , ~ o - ~ o ' - ' - : 4-0-0- - -. - - 0.20 o--""--- o ~ o ~ o - o - o 0.4 0.3 0.2 c 0. I 0 0.30 +o-o~~o-O-o-o-o-----o 0.40 5o o - o ~ o - c c c - c - o c c L L o - - ------ - I100 1120 I140 1160 , , , i 0.6 0.7 0.8 0.9 I .o XPbO 1173 FIO. 3.-Activity of PbF2 plotted against mol fraction of PbO for PbO + PbF2 and PbO + PbFz + SiOz melts at 1173 K. PbFz/SiOz = (A)m, (B) 4, (C) 1, @) 0.5. (0) 0.2. x , from apbo by Gibbs- Duhem.T = 1173 K.W . F. CALEY AND C . R . MASSON 2957 2 2 0 200 I80 I60 > E 'u: 140 Ei d - 120 100 ao 60 I080 I100 1120 I140 I I60 1173 temperature /I( to the following compositions : FIG. 4.-E.m.f. plotted against temperature for PbO+PbF2 + Si02 melts. The curves correspond curve XPbO XPbFz xSi02 0.85 0.05 0.10 0.90 0.08 0.02 0.86 0.07 0.07 0.70 0.05 0.25 0.80 0.07 0.13 0.80 0.10 0.10 0.70 0.10 0.20 curve Xpbo (8) 0.80 (9) 0.70 (10) 0 0.66 (11) 0 0.60 (12) 0.70 (13) 0.60 (14) 0.60 XPbF2 xSi02 0.16 0.04 0.15 0.15 0.17 0.17 0.13 0.27 0.24 0.06 0.20 0.20 0.32 0.082958 THERMODYNAMICS OF PbO+PbF2+Si02 MELTS THE SYSTEM PbO + PbF2 + Si02 The temperature dependence of the cell e.m.f. values for fourteen slag compositions (D) with the cell Ir,Ar,Pb(l), (PbO + PbF2 + Si02)(,,lCaF2,,,I(standard melt), Pb(l,,Ar,Ir corrected for the e.m.f.of the standard melt by eqn (1) are shown in fig. 4. The equations corresponding to the lines in fig. 4 are given in table 2, along with the maximum and standard derivations. Activities of PbF2 in these melts, as calculated also by eqn (2), are included in fig. 3 for four PbF2/Si02 ratios, with PbO levels ranging from 60 to 95 rnol %. Error bars for each slag are included in the graph. Because of the high melting points (about 1153 K), and consequent high volatil- ization rates of the slags, the cells containing 94 mol % PbO + 3 rnol % PbF2 + 3 mol % Si02 and 94 mol % PbO +2 mol % PbF2 +4 mol % Si02 were operated only between 1163 and 1178 K. The activities of PbF2 at 1173 K for these compositions are shown also in fig.3. From this figure it may be seen that the PbF2 activity increases with addition of PbF2, at constant PbO levels. Also, for a constant PbF2/Si02 ratio, the activity of PbF2 increases steadily with decreasing PbO content. ISOACTIVITY PROFILES The results in fig. 3 were used to construct isoactivity contours for eight PbF2 activity levels at 1173 K. These are shown in fig. 5, along with the isoactivity con- tours determined previously for PbO. The profiles are slightly concave toward the PbO apex at low PbF2 activities, and away from the PbO apex at higher levels. The contour for apbFz = 0.1 is approximately linear. Table 3 presents the PbF2 activity data taken from fig. 3. XPbO FIG. 5.-Isoactivity contours for -, PbO [ref.(2)] and -a-, PbFz (this work) for the system PbO + PbFz + SiOp at 11 73 K.W. F. CALEY AND C. R . MASSON 2959 TABLE 2,-E.M.F. DATA FOR THE SYSTEM PbO+PbFz+SiOz (CELLS B AND D) XPbO 0.60 0.60 0.60 0.66 0.70 0.70 0.70 0.70 0.80 0.80 0.80 0.85 0.86 0.90 XPbFz 0.32 0.20 0.13 0.17 0.24 0.15 0.10 0.05 0.16 0.10 0.07 0.05 0.07 0.08 x~i02 0.08 0.20 0.27 0.17 0.06 0.15 0.20 0.25 0.04 0.10 0.13 0.10 0.07 0.02 E = A+BT/mV A - 55.6079 -66.1121 - 71.5736 - 70.3822 - 45.0980 - 51.2591 - 41 3220 - 40.7704 - 24.9363 - 4.828 1 - 28.2726 - 7.0429 - 3.6840 24.3271 0.1023 0.1242 . 0.1477 0.1466 0.1150 0,1444 0.1580 0.1945 0.1329 0.1435 0.1833 0.2055 0.1451 0.1716 standard deviahon/mV 0.3 0.6 0.5 0.1 0.1 0.2 0.7 0.7 0.2 0.5 0.2 0.4 0.6 0.5 maximum deviation/mV 0.5 0.9 0.8 0.2 0.2 0.5 1.3 1.4 0.4 0.1 0.4 0.7 1.3 1 .o TABLE 3.-ISOACTIVITY DATA TAKEN FROM FIG.3 FOR CONSTANT apbF2 AT 1173 K aPbF2 XPbO 0.30 0.585 0.61 5 0.20 0.604 0.670 0.700 0.15 0.583 0.646 0.718 0.750 0.10 0.640 0.700 0.775 0.807 0.05 0.71 5 0.772 0.838 0.870 XPbF2 0.332 0.385 0.198 0.264 0.300 0.139 0.177 0.226 0.250 0.120 0.150 0.180 0.193 0.095 0.114 0.130 0.130 0.198 0.066 qbF2 XPbO 0.025 0.700 0.785 0.844 0.888 0.906 XPbF2 0.050 0.072 0.078 0.090 0.094 XSioZ 0.250 0.143 0.078 0.022 0.278 0.177 0.056 0.02 0.805 0.866 0.900 0.91 5 0.065 0.67 0.080 0.085 0.130 0.067 0.020 0.240 0.150 0.045 0.190 0.1 14 0.032 - - 0.01 0.860 0.900 0.940 0.047 0.050 0.060 0.093 0.050 DISCUSSION Following the theoretical treatment outlined in the accompanying communica- tion,l ionic distributions of silicate and fluorosilicate anions in the ternary PbO + PbF2 + Si02 may be calculated from the activities of PbO and PbF2 for each melt composition, if the equilibrium constants k and k’ are known. The equilibrium constant k has the value 0.196,13 and the PbO and PbF2 activities for any melt com- position may be obtained from fig.5. Knowing these three parameters, the follow- ing procedure was used to evaluate k’. Values of apbo (= N02-) and aPbFz (= N$) corresponding to the point of inter- section of a pair of isoactivity lines in fig. 5 were first selected. These values were substituted along with the value of k (= 0.196) and an arbitrary value of k’ in eqn2960 THERMODYNAMICS OF PbO+PbF, +SO, MELTS (19) and (23)-(25) of ref. (1) to yield a calculated value for the composition of the melt.The process was repeated with other values of k‘ until a value was found which yielded a composition close to the experimental value at the point of intersection of the isoactivity lines. The value of k’ thus determined was then used to calculate compositions corresponding to other points of intersection in fig. 5. This value was then adjusted so as to provide the best fit with the bulk of the experimental data over the entire range of compositions. TABLE $.-CALCULATED AND EXPERIMENTAL MELT COMPOSITIONS FOR VARIOUS VALUES OF apbO AND aPbF2 (k = 0.196, k’ = 0.40) UPbO aPbF2 0.4 0.3 0.2 0.15 0.1 0.05 0.025 0.5 0.2 0.15 0.1 0.05 0.025 0.6 0.1 0.05 0.025 0.7 0.05 0.025 0.02 0.8 0.025 0.02 0.01 calculated XPbO 0.564 0.579 0.596 0.620 0.648 0.666 0.665 0.670 0.681 0.697 0.707 0.746 0.749 0.753 0.812 0.807 0.806 0.877 0.873 0.865 XPbFz 0.370 0.268 0.21 5 0.158 0.098 0.063 0.273 0.21 5 0.157 0.097 0.063 0.169 0.102 0.065 0.113 0.071 0.062 0.081 0.070 0.046 Xsioz 0.066 0.153 0.189 0.222 0.254 0,271 0.062 0.115 0.162 0.206 0.230 0.085 0.149 0.182 0.075 0.122 0.132 0.042 0.057 0.089 experimental XPbO 0.578 0.590 0.602 0.615 0.640 0.660 0.670 0.675 0.682 0.693 0.705 0.740 0.744 0.746 0.810 0.802 0.800 0.878 0.873 0.860 XPbFz 0.322 0.191 0.147 0.109 0.060 0.030 0.265 0.192 0.179 0.088 0.050 0.165 0.104 0.061 0.122 0.074 0.060 0.088 0.077 0.046 x5i02 0.100 0.219 0.251 0.276 0.300 0.310 0.065 0.133 0.139 0.219 0.245 0.095 0.152 0.193 0.068 0.124 0.140 0.034 0.050 0.094 In these calculations, which were performed with the aid of a computer, the moles C moles PbO + moles PbFz +moles SiOz following relationship was also required : xc = (3) where C represents PbO, PbF2 or SO2.The value of k’ found by this iterative procedure was 0.40+0.025. The results of these computations are given in table 4, which shows the calculated and experi- mental melt compositions corrresponding to the points of intersection of the iso- activity lines in fig. 5. Fig. 6 illustrates the range of applicability of the theory. The solid lines are the experimentally-determined isoactivity lines for PbO, as shown in fig. 5, and the points, taken from table 4, are calculated compositions for melts of selected aPbO for which the aGtivities of PbO and PbF, were obtained from fig. 5. The agreement between theory and experiment is reasonable for melts with aPbO >, 0.5.In this range ofW. F. CALEY AND C . R. MASSON 2961 XPbO FIG. 6.-Comparison of experimental iso-PbO activity contours (-) with theoretical melt compositions a for melts of selected QpbO and aPbF2, corresponding to the points of intersection of the isoactivity lines in 6g. 5. The compositions were calculated for k = 0.196, k’ = 0.400. 1.0 i XPbF, = O ..- 0. a 0.6 0.2 0.1.0 - I I I I 0.1 0.2 0.3 0.4 XSiOz lXPb0 FIG. 7.-Calcuiated on fractions of 02- and F- in PbO+PbF, +SOz melts at 1173 K.2962 THERMODYNAMICS OF PbO+PbF2+Si02 MELTS compositions the theoretical expressions derived previously are applicable and may be used to evaluate ionic distributions in this system.IONIC DISTRIBUTIONS Calculated ion fractions for 02- and F- are plotted in fig. 7 against XSi02/XpbO for various values of &bF2. These were obtained by interpolation from the experi- mental isoactivity curves in fig. 5, with No2- = apbo and NF = atbF2. 0.25 0.20 0.1 5 I VP 0 .I s 0.10 0.0 5 0 Xsioz lxpw FIG. 8.-Calculated ion fractions of SiOd- in PbO+PbFz+SiOz melts at 1173 K. Values of N02- and NF- from fig. 7 were substituted, along with k = 0.196 and k' = 0.4, in eqn (19) and (13) of ref. (1) to yield the ion fractions of silicate and fluoro- silicate anions. The results for SiO$- and Si20$- are shown in fig. 8 and 9. The addition of relatively small amounts of PbF2 to a melt of fixed Si02/Pb0 ratio causes a marked lowering in the ion fractions of these species.In addition, the maximum in these distributions (which, in the absence of PbF2, occurs at &ioz/&bo = 0.5 for SiOt- and 0.667 for Si20$-) is displaced to lower xsiOz/&bO ratios as XpbF2 is increased. The calculated ion fractions of the monofluorosilicate ions Si03Fi- and Si206F5- are shown in fig. 10 and 11. There is evidence that these ion fractions also exhibit maxima at specific Si02/Pb0 ratios but, except for melts with XPbF2 = 0.15, the results are not sufficiently extensive to illustrate this feature.xSiO2 lXPb0 FIG. %-Calculated ion fractions of Si20:- in PbO + PbF2 + Si02 melts at 1173 K. xSiO2/xPbO FIG. 10.-Calculated ion fractions of Si03F3- in PbO+PbFz+SiOz melts at 1173 K.2964 THERMODYNAMICS OF PbO+PbF2+SiOa MELTS TABLE 5 .-CALCULATED ION FRACTIONS OF OXIDE, FLUORIDE, SILICATE AND FLUOROSILICATE IONS IN MELTS WITH XSiOz = 0.2 ion 02- F- SiO$- Si03F3- Si02F3- SiOF, SiF4 ion fractions for melt compositions (N 3 0.0001) x s i o z = 0.2 xsio2 = 0.2 0.4150 0.693 0.5200 0.2449 0.4472 0.280 0.1780 0.0767 0.0335 0.0330 0.0063 0.0142 0.001 2 0.006 1 0.0002 0.0026 xsio2 = 0.2 XPbo = 0.7 XPbo = 0.6 XPbo = 0.8 XpbF2 = 0.1 XpbFz = 0.2 si2067 - 0.022 0.01 19 0.0028 Siz06F5- 0.0023 0.001 2 Siz05 Fi- 0.0004 0.0005 Si204F:- 0.0001 0.0002 Si203Fa- - o.oO01 Si202F; - - Si20F6 - - Si30:; 0.001 8 0.0008 0.0001 si309~7- 0.0002 - Si308@- I - si40; $j- 0.0001 o.Ooo1 - x 0.005 0.004 0.003 I n % 0, .* z 0.002 0.9969 0.9999 0.9997 X P b F , = 0.10 1 O.OO' 1- xSiOz/xPbO FIG.ll.-Calculated ion fractions of Si20sF5- in PbO+PbF2+Si02 melts at 1173 K.W.F. CALEY AND C. R. MASSON 2965 In the melts under consideration, ion fractions of higher fluorosilicate ions (SiO,Fg-, SOFT, Si205Fi-, Si,O,F3-, etc.) are generally very low, but become increasingly significant as XPbF2 is increased. Table 5 shows the calculated ion fractions of various species when PbF2 is substituted for PbO in a melt of original composition X,,, = 0.8, Xsioz = 0.2. Ion fractions < 0.0001 are not reported. The bulk of the fluorosilicate ions are of low molecular weight, as might be expected for highly basic melts of this nature. Replacing PbO by PbFz leads to an increase in the sum of the free oxide and fluoride ions and a corresponding decrease in the sum of the ion fractions of the " monomeric ", " dimeric ", and higher silicate and fluoro- silicate ions.The net effect of substituting PbFz for PbO at constant silica content is thus to cause a decrease in the mean chain length of the anions, a result which might be anticipated and which is in line with viscosity and other considerations. The calculated ion fractions shown in table 5 together constitute 99.69-99.99 % of the total. 0.08 0.07 0.0 6 0.05 R 8 0.04 & .3 u d .I 0.03 0.02 0.0 I 0 XSiOZ/XPbO FIG. 12.-Calculated ion fractions of Si0.+mFg,4-m) - in PbO + PbFz + Si02 melts with XpbF2 = 0.2 at 1173 K.2966 THERMODYNAMICS O F PbO+PbF,+SiO, MELTS Fig. 12 and 13 show the calculated ion fractions of the monomeric Si04-,,tFc-m-) and dimeric Si,O,-,Fg-")- species in melts for which XPbF2 = 0.2.At the highest silica content (Xsio2/Xpbo = 0.333) these ions constitute 13.74mol % of the total, the remainder being essentially 0,- (41.50 %) and F- (44.72 %). An interesting feature is the emergence of the uncharged molecular species SiF4 as a small but signifi- cant constituent when the silica content of the melt is increased. As this species is volatile, the equilibrium will be displaced. This may partly account for the difficulty in obtaining steady e.m.f. values for melts with higher PbFz contents. The result is in line with the observation of Mitchell l4 that SiF4 is lost from CaO + CaF, + SiO, melts of high CaF, content and of Kumar et a l l 5 that the loss of SiF4 from CaO + CaF, + SO2, Na,O + NaF + SiO, and MgO + CaF, + SiO, melts is strongly dependent on the basicity.0.003C O.OC25 0.00 20 d 2 0.0015 E .- U .- 0.0 0 l o 0.0005 XSiOz/XPbO FIG. 13,-Calculated ion fractions of Si207-mF(,6 in PbO+ PbF2 + SiOz melts with XpbFz = 0.2 at 1173 K. (Ion fractions of Si20F6 were too small for illustration). In view of the simplifying assumptions in the theory, the calculated ionic-distribu- tions reported here must be regarded as approximate, and subject to refinement. An important limitation is the assumption that the reactivity of the 0- groups is not influenced by their degree of substitution by F on the silicate ions. This cannot beW. F. CALEY AND C. R. MASSON 2967 expected to hold rigorously, in view of the difference in electronegativity between the 0 and F atoms (3.5 and 4.0, respectively), so that the value of k’ must vary, to same extent, with the value of m.In addition, restriction of the treatment to linear chains confines the theory to highly basic melts. The results, however, provide a basis for a more complete description of these systems and yield some insight into the factors which govern their thermodynamic properties. The conclusion that discrete fluorosilicate ions are present in PbO + PbF2 + Si02 melts is supported by the results of a recent investigation l6 in which these ions were identified chromatographically as their trimethylsilyl derivatives in extracts of PbO + PbF2 + Si02 glasses by a method described elsewhere?’ We thank Dr. A. E. Grau, who participated in the early stages of this work, and Miss N. Morrison for experimental assistance. We also thank Dr. G. K. Muecke for the microprobe analyses. C. R. Masson and W. F. Caley, J.C.S. Faraalzy I, 1978,74,2942. A. E. Grau, W. F. Caley and C. R. Masson, Canad. Metal. Quart., 1976, 15,267. C. Wagner, J. Electrochem. SOC., 1968,115,933. J. J. Egan, J. Nuclear Materials, 1974,51, 30. T. N. Rezukhina, T. F. Sisoeva, L. I. Holokhonova and E. G. Ippolitov, J. Chem. Thermo- dynamics, 1974,6,883. R. W. Ure, Jr., J. Chem. Phys., 1957, 26, 1363. ’ R. J. Heus and J. J. Egan, 2. phys. Chem. N. F., 1966,49, 38. * J. W. Patterson, J. Electrochem. SOC., 1971, 118, 1033. C. R. Masson, S. G. Whiteway and C. A. Callings, Amer. Ceram. SOC. Bull., 1963,42,745. lo N. A. Haroun and S. M. El-Hout, Amer. Ceram. SOC. Bull., 1976,55, 1063. l 1 J. W. Hinze and J. W. Patterson, J. Electrochem. SOC., 1973, 120, 96. l2 L. S. Darken and-R. W. Gurry, PhysicaZ Chemistry ufMetds (McGraw-Hill, 1953), pp. 260- 264. C. R. Masson, Chemical MetaZZurgy of Iron and Steel, ed. B. B. Argent and M. W. Davies (Iron and Steel Inst., London, 1973), p. 3 ; J. Iron and Steel Inst., 1972, 210, 89, 369. l4 A. Mitchell, Trans. Faraday SOC., 1967, 63, 1408. lS D. Kumar, R. G. Ward and D. J. Williams, Disc. Faraalzy SOC., 1961,32, 147. l6 H. P. Calhoun, W. D. Jamieson and C. R. Masson, J. C. S. Dalton, 1979, in press. (8/715) l 7 J. Gotz and C. R. Masson, J. Chem. SOC. A, 1970, 2683. (PAPER 8/375) 1-94

ISSN:0300-9599    

DOI:10.1039/F19787402952

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Surface Properties and Catalytic Activity of a Mo-fixed Catalyst' Structure of the Active Site and Mechanism for Selective Oxidation of Ethyl Alcohol BY YASUHIRO IWASAWA," YASUO NAKANO AND SADAO OGASAWARA Department of Applied Chemistry, Faculty of Engineering, Yokohama National University, Ooka-cho, Minami-ku, Yokohama, Japan Received 13th March, 1978 The physico-chemical properties of the surface of a Mo-fixed catalyst obtained by a ready reaction between Mo(n-C3H5)4 and surface OH groups of SiOa were investigated by X-ray photoelectron spectroscopy, photoluminescence technique, U.V. diffuse reflectance spectroscopy, i.r. spectrometry and Hz/02 uptake. The difference between the surface properties of the fixed catalyst and the con- ventional impregnation catalyst is discussed.The Mo-ked catalyst showed higher activity and selectivity than the conventional impregnation catalyst in the ethyl alcohol oxidation. Selective oxidation to form acetaldehyde proceeded by a two-stage redox (Mo6+ + Mo4+) mechanism. Coordinatively unsaturated dioxomolybdenum of uniform nature in a tetrahedral coordination is the active species in the selective oxidation of ethyl alcohol. The surface properties of heterogeneous catalysts have been studied extensively.2 However, the surface state of an impregnation catalyst which has important effects on the catalytic activity, selectivity and reaction mechanism is in general complex and still indefinite: the nature of the catalyst surface strongly depends upon the kind of support, the method of preparation of catalysts, the preparation condition, etc.Consequently, the mechanism of surface reaction on a conventional impregna- tion catalyst becomes ambiguous and conflicting results may sometimes be drawn. summarized the method of preparation of supported cataIysts using metal n-ally1 complexes and their catalytic activities. Metal n-ally1 complexes selectively react with the OH groups of solid supports like SiO, to form stable surface complexes under oxygen-free atmospheres. reported high activity for the metal oxide or metal catalysts supported on metal n-ally1 or methallyl complex bases in ethylene hydrogena- tion, butadiene polymerization, metathesis and hydrogen oxidation. A synthetic catalyst with a well-defined surface is very suitable for obtaining clear information on the active site and reaction mechanism.The present study contributes to our understanding of the surface properties of molybdenum-fixed silica catalyst, its catalytic activity and the action of the active site in the selective oxidation of ethyl alcohol, chosen as a typical reactant, in com- parison with a conventional impregnation catalyst. Ballard Yermakov and coworkers EXPERIMENTAL PREPARATION OF Mo-FIXED CATALYST The Mo-fixed catalysts were obtained from the ready reaction between a tetrakis-n- allylmolybdenum complex, Mo(n-CSH&, and the OH groups on the surface of silica (silica gel, 30-60 mesh, Nishio Industry, Japan, surface area of 567 m2 g-'). All procedures 2968Y . IWASAWA, Y . NAKANO AND S. OGASAWARA 2969 for synthesis were carried out in a specially devised Pyrex-glass apparatus under vacuum (base pressure of 5 x lW5 Torr) or in a flow of high purity (99.999 %) argon at atmospheric pressure.M o ( ~ ~ - C ~ H ~ ) ~ was synthesized according to the following equation described by Wilke et al. ; 5 -lS°C, 12 h diethyl ether C3H5Cl + Mg (large excess) C3H5Mg C1 -35"C, 1,5h -78'C, 20 h -3OOC 4CSHSMgClf MoClj -+ -+ + Mo(z-C~H~)~. (2) diethyl ether pentane The Mo-fixed catalyst was prepared by a similar procedure to that described by Yermakov et aL6 as folIows ; +Si-O\ 0 2 , 20°C Mo e psi-o>Mo=O 011) (bluish black - black) ( 5 ) +Si-O/ H2, 600°C b S i - 0 s!-o>Mo<z (IV) (white). %Si-O\ 0 2 , 3OOOC +Si-o/ H2, 450°C bs1-0 Mo-0 + The procedure was carried out in a vacuum of base pressure 1 x Torr.Silica gel was heated at 550°C for 1.5 h in oxygen (70 Torr) with a liquid N2 trap and evacuated for 1 h at the same temperature prior to Mo fixation. The surface complex (I) formed was washed three times with pentane in vacuum, in order to remove the Mo-complex, with no chemical interaction with the silica surface. Some Mo remained in the pentane solution after the Mo-fixing and the washing solution was analysed by both gravimetry as Moo3 (requiring careful treatment because of easy sublimation) and colorimetry at 460nm using a thiocyanate- SnClz method. Thus, by subtracting the final M o ( ~ - C ~ H ~ ) ~ concentration left in the pentane from its initial concentration, the amount of molybdenum fixed on the silica gel surface was determined. PREPARATION OF CONVENTIONAL IMPREGNATION CATALYST The impregnation catalyst was obtained by an ordinal impregnation method using aqueous ammonium paramolybdate solution.REACTION AND PRODUCT ANALYSIS The decomposition of ethyl alcohol to acetaldehyde, diethyl ether, ethylene and water in the presence or absence of oxygen was investigated in a conventional closed-circulating system (dead volume : 170 cm3) using 0.39 g of catalyst. The reaction products were analysed by gas chromatography using a 2 m column of dioctyl sebacate at 70°C. SPECTROSCOPIC STUDY The i.r. absorption spectra were measured using a 36.5 mg disk of 20 mm Mo catalyst fixed on to aerosil, under various conditions, in a glass-cell directly connected to a closed circulating system using a JASCO IRA-2 i.r.spectrometer. The U.V. diffuse reflectance spectra were taken in the range 210-700 nm after the catalysts had been oxidized at 550°C for 2 h in a separate system. The X.P.S. emissions were monitored by means of McPherson ESCA 36 spectroscopy, where the binding energies were referred to 83.7 eV of the Au 4f712 level taking account of the charge effect. Photoluminescence was investigated in a quartz cell2970 PROPERTIES OF Mo-FIXED CATALYST directly connected to a vacuum system using a JASCO FP-4 spectrophotometer. The reproducibility of the peak position at maximum intensity was within 3 nm. The e.s.r. spectra were measured on a JEOL e.s.r. spectrometer after quenching the reaction by cooIing the catalyst to room temperature during the ethyl alcohol oxidation.RESULTS Mo-FIXING PROCESS Ballard and Yermakov et aL6 demonstrated that the metal n-ally1 complexes easily react with the surface OH groups of Si02 at room temperature. The decrease in the intensity of the OH stretching vibration in fig. 1 indicates that the Mo ions combined with the silica surface through the oxygen atom of the OH groups. wavenumber 1crn-l (b) SiOz, (c) the difference spectrum of (a)-@). FIG. 1.-Intensity change in the OH vibration band by Mo-fixing; (a) the fixed catalyst (IV), Yermakov et a1.’ suggested from an i.r. spectrum that the ligands of the surface complex (I) are of the 0-ally1 type. The reaction of Mo(n-C3H,), and the surface OH groups was more rapid when the silica containing the larger number of OH groups was employed. Thus the rate of Mo fixation depended upon the surface state of silica.At the silica surface with a low content of OH groups, in addition to eqn (3), the following step may also participate in the Mo-fixing process ;6 + Si-OH 20ec +Si-0 Si-0>Mo(c3H5)2 +C3H6(g) (7) * Si\, + Mo(Tc-C~H~)~ -+ + Si/ 3 Si-C3H5 REDOX BEHAVIOUR OF Mo-FIXED CATALYST The reversible cycle of reduction-oxidation between the fixed catalysts (11), (111) and (IV) was observed as shown in eqn (5) and (6). The oxidation of (11) to (111) by oxygen was very rapid at room temperature. The amount of oxygen consumed was 1.1 0-atoms per Mo atom. More than stoichiometric uptake of oxygen may be caused by adsorption of oxygen and/or partial over-oxidation of Mo2+ to Mo6+,Y . IWASAWA, Y . NAKANO AND S .OGASAWARA 297 1 although this is small. Further oxidation from (111) to (IV) proceeded slowly at 100°C and was very ready at 280-300°C. The catalyst (IV) was reduced to (111) with hydrogen at 450"C, at which stage the uptake of hydrogen molecules per Mo atom was unity with an error of 6%. (111) was converted to the reduced form (11) in 3 h at 600°C under 80 Torr of hydrogen, the ratio of the amount of dihydrogen uptake to the number of Mo atoms being 2k0.1 as shown in fig. 2. Thus the fixed catalysts, (II), (111) and (IV), can be produced individually by controlled reduction- oxidation steps. The irreversible surface reconstruction of the Mo-fixed catalyst employed was not observed during repeated redox treatments, judging from the redox behaviour of the catalyst and the reproducible catalytic activity.reduction time/h FIG. 2.-Amount of Hz consumed per Mo atom during reduction with H2 ; Mo/Si02 = 1.7 wt %. PHYSICOCHEMICAL PROPERTIES OF Mo-FIXED CATALYST (IV) The X.P.S. of the fixed catalyst (IV) showed characteristic doublet peaks of a hexavalent molybdenum ion. The values were found to be 235.7 &- 0.3 eV for the Mo 3 4 orbital and 232.7 k0.3 eV for the Mo 3 4 level and the full width at half maximum height of the Mo 3 4 emission peak was x 2.4 eV. MOO, as a standard sample gave the well-known doublet peaks at 235.5 and 232.5 eV for 3d3 and 3d3, respectively. E m . spectroscopy showed no signal at any temperatures, Mo5+ species not being contained in catalyst (IV). The U.V. diffuse reflection spectrum showed absorption peaks at 290 and 225 nm, neither peak nor shoulder at longer wavelengths being found.An illumination in the charge transfer band of the Mo=O double bond was followed by a strong emission in the energy range (20-25) x lo3 cm-l. The luminescence from the Mo-impregnated catalyst depended upon the energy of the exciting light as shown in fig. 3. The emissions are divided into two. Photoexcitationat (34.4-35.7) x lo3 cm-l generated an emission peak at 21.8 x lo3 cm-l, while illumination at (30.8-32.3) x lo3 cm-l was accompanied by luminescence at 22.9 x lo3 cm-'. On the other hand the emission energy of the Mo-fixed catalyst (IV) was independent of the exciting energy, indicating the presence of only one kind of molybdenum emitting species. The typical transmittance i.r.spectra of the fixed catalyst (IV) are given in fig. 4(a) and (b), where absorption peak at 916 cm-l was observed, SiO, having no peak near this position. The absorption is due to the stretching vibration of the Mo=O bond. The reactivity of this molybdenum-oxygen bond was checked by i.r. spectrometry. Fig. 4(c) is a variation of the band intensity of the Mo-0 bond vibration with ex- posure to ethyl alcohol, methyl alcohol or oxygen at given temperatures. The Mo=O double bond readily reacted with ethyl alcohol at 150°C ; the intensity markedly decreased from 1 to 2, in fig. 4(c), by reduction with ethyl alcohol in the absence of2972 PROPERTIES OF Mo-FIXED CATALYST oxygen. The absorption band at 916 cm-l was restored to its original level by oxidation with oxygen. The decrease in the band with ethyl alcohol was much faster than its increase with oxygen [l+2 and 2 4 3 in fig.4(c)]. Methyl alcohol behaved similarly to ethyl alcohol except at the higher temperature studied. The 2 20 emission energy/103 cm-l FIG. 3.-Photoluminescence of (a) fixed catalyst (IV) and (b) impregnation catalyst ; exciting light, 1, 285 ; 2,295 ; 3, 320 nm. 4 0 90 ,;- wavenumber /cm-' FIG. 4.-1.r. absorption spectra of the Mo-fixed catalyst ; (a) Difference spectrum between the Mo- fixed catalyst (IV) and Si02, where a reference Si02 disk of adequate weight was chosen by trial and error. (6) (1) fixed catalyst (IV), (2) Si02, (3) calculated difference spectrum between (1) and (2). (c) (1) fked catalyst (IV), (2) spectrum after reaction with ethyl alcohol for 20 min at 150°C in the absence of 02, (3) after oxidation with O2 for 1.2 h at 150°C after (2), (4) in a steady state under ethyl alcohol : Oz = 1 : 11 at 155"C, (5) after reaction with methyl alcohol at 180°C in the absence of 02.Y .IWASAWA, Y . NAKANO AND S . OGASAWARA 2973 steady-state level in the i.r. absorption band intensity depended upon the ratio of ethyl alcohol to oxygen. The behaviour of the 916cm-l absorption band was compatible with that expected from the kinetic data. DEHYDROGENATION OF ETHYL ALCOHOL IN THE ABSENCE OF OXYGEN The reactivity of the active sites, the Mo=O bonds, in the catalyst was studied in a closed circulating system in the temperature range 100-200°C in the absence of oxygen. The catalysts were treated at 550°C for 1.5 h under 70 Tom of oxygen with a liquid N2 trap followed by evacuation for 1 h prior to dehydrogenation; when necessary, the fixed catalyst (IV) was converted to catalysts (11) or (111) as shown in eqn (5) and (6).The products of the reaction were acetaldehyde, diethyl ether, ethylene and water. Evolution of hydrogen was not observed in the course of dehydrogenation. The formation of acetaldehyde on both the fixed (IV) and the impregnated catalysts are shown in fig. 5, where reactions on the fixed catalysts having tetra- and di-valent molybdenum structures are also shown. The dehydrogena- tion activity of the fixed catalyst (IV) was found to be much higher than those of the impregnated and other fixed catalysts. The fixed catalyst with Mo2+ was essentially inactive.Ethyl alcohol dehydrogenation on the conventional impregnation catalyst below 150°C was not significant, while on the fixed catalyst (IV) it took place even below 100°C. The activity of dehydrogenation decreased in the following order : 10 30 5 reaction timelmin FIG. 5.-Dehydrogenation activities of the fixed catalysts with three different structures and of the impregnation catalyst in the absence of oxygen ; reaction temp. = 150"C, catalyst = 0.39 g, Mo/SiOz = 1.7 wt %. (1) Fixed catalyst (IV), (2) fixed catalyst (III), (3) fixed catalyst a), (4) im- pregnation catalyst. Acetaldehyde was produced more selectively on the fixed catalyst than over the impregnated one as shown in table 1. The selectivity towards acetaldehyde formation on the fixed catalyst was 100 % at 100°C.A difference in the distribution of the inter- and intramolecular dehydration products on both catalysts was also observed : diethyl ether was predominantly produced on the fixed catalyst, while on the irn- pregnated one ethylene was mainly formed at higher temperatures. The activation energies of dehydrogenation were determined from the initial rates of reaction to be 51.0k5.4 kJ mol-' on the fixed catalyst (IV) and 46.0k5.4 kJ mol-1 on the im- pregnated one.2974 PROPERTIES OF Mo-FIXED CATALSYT The dependence of the initial rates of acetaldehyde formation upon the number of Mo6+ ions remaining in the fixed catalyst which could be determined from the amount of Mo6+ converted to Mo4+ by the controlled reduction mentioned above, is shown in fig.6 in which the initial rates change linearly as the amount of Mo6+ decreased. This result indicates that the active species is a hexavalent molybdenum of uniform nature. TABLE SE SELECTIVITY AND ACTIVATION ENERGY OF THE DECOMPOSITION OF ETHYL ALCOHOL ON THE MO-FIXED CATALYST (Iv) (A)* AND THE IMPREGNATION CATALYST (B)* IN THE reaction temperature 1°C 100 125 150 175 200 activation energy /kJ mol-l ABSENCE OF OXYGEN ylectivity/ % A B A B A B acetaldehyde diethyl ether ethylene - 0 0 100 - 0 - 96.2 I 3.8 - 86.7 82.6 8.5 9.6 4.8 7.8 82.7 73.5 11.3 11.8 6.0 14.7 67.4 - 14.2 18.3 - 51.0k5.4 46.0k5.4 75.2k7.5 62.7k7.5 - 81.1k8.4 * Catalyst = 0.39 g, Mo/SiOz = 1.7 wt%. X 4 4 3 - 2 0,- 0.5 e a x v Mo6+/total Mo FIG. 6.-Dependence of the initial rates of CHJCHO formation upon the amount of Mo6+ in the fixed catalyst (IV) at 150°C.OXIDATION OF ETHYL ALCOHOL WITH OXYGEN The rate of ethyl alcohol dehydrogenation in the absence of oxygen decreased with reaction time. When oxygen was admitted on to the catalyst during the reaction as shown in fig. 7, the oxidation activity was recovered. The oxidation of ethyl alcohol to form acetaldehyde in the presence of oxygen attained a steady state. The selectivity in the presence of oxygen was higher than in its absence. Table 2 shows a higher selectivity on the fixed catalyst (IV) than on the impregnation catalyst. The rate equation under experimental conditions was approximately expressed by the power rate law as follows; r + kP&2P&5, where PE and Po, represent the partial pressures of ethyl alcohol and oxygen, respectively. When acetaldehyde formation from ethyl alcohol in the presence of oxygenY .IWASAWA, Y . NAKANO AND S . OGASAWARA 2975 proceeds by the two-stage redox mechanism, the following scheme and rate equation for the steady state procedure may be expressed ; :$I:> Mo <: +C2H50H + k1 4si-o "-"> Mo=O + CH3CH0 + H,O (8) Mo=O+&O,+ :I:> Mo //o No (9) k2 + Si-O\ 3 Si-O/ PElr = l/ki +P&k2 J P . 2 (10) PE/r is plotted against P E I J c 2 in fig. 8, where both PE and Po, are varied and the rates measured. The correlation between both variables, PE/r and P,/,/X, was found to have good linearity. I .- reaction timelmin FIG. 7.-02 effect on CH&HO formation on 0.39 g catalyst (Mo/Si02 = 1.7 wt%) at 150°C; -: in the presence of 02, - - - : in the absence of 02.TABLE 2.-sELECTIVITY AND ACTIVATION ENERGY OF THE OXIDATION OF ETHYL ALCOHOL WITH OXYGEN ON THE FIXED CATALYST (Iv) (A)* AND THE IMPREGNATION CATALYST (B)* reaction selectivity/ % temperature acetaldehyde diethyl ether ethylene /"C A B A B A B 100 100 - 0 - 125 97.1 - 2.9 - 150 94.5 89.2 4.5 7.0 1 .o 2.8 175 93.5 87.2 5.3 7.6 1.2 5.2 200 - 85.1 - 8.1 - 6.8 - 0 0 - activation energy /kJ mo1-I 62.3+ 5.4 58.9+ 5.4 - 56.4+ 8.4 - - * Catalyst = 0.39 g, Mo/SiOz = 1.7 wt%. The rate constants, kl and k2, could be calculated from fig. 8. The values at 150°C are given in table 3. The initial rates of reduction of the Mo6+ species with ethyl alcohol [eqn (S)] and of oxidation of the Mo4+ species with oxygen [eqn (9)] were obtained independently by gas chromatography and by volumetry of O2 uptake, respectively.The rate constants determined from the rates thus obtained are also given in table 3.2976 PROPERTIES OF Mo-FIXED CATALYST PE/d% FIG. 8.-correlation of pE/r with pE/dK2 at 150°C. TABLE 3.-RATE CONSTANTS, kl AND kZ, IN STEPS (8) AND (g), RESPECTIVELY * 10kdmin-1 103k2/[m3 (s.t.p.)]+ min-1 1 1.7+ 0.3 7.6+ 0.8 2 1 . 5 k 0 . 3 7.1k0.9 * 1 : Obtained from eqn (10) ; 2 : independently obtained from the rates of steps (8) and (9). E.S.R. MEASUREMENT When the catalyst surface was exposed to ethyl alcohol at 165°C in the absence of oxygen, an asymmetric signal assigned to Mo5+ species (1.93, for g1 and 1.8g5 for gji) was observed. However, the maximum amount of the Mo5+ ions was found to be 0.5+0.12 % of the total molybdenum contained in the catalyst.The amount of Mo5+ ion in the steady state oxidation was only 0.3 %. On the other hand about 20 % of the Mo6+ ions under similar conditions were converted to Mo4+ judging from the amount of acetaldehyde formed. It seems that the Mo5+ species does not play a role in the oxidation of ethyl alcohol. DISCUSSION The conventional impregnation method followed by calcination for catalyst activation has been employed extensively in the preparation of supported catalysts. For a molybdenum-supported catalyst, inorganic acids or their salts, like ammonium paramolybdate in an aqueous solution, are generally used. Oxymolybdenic species in aqueous solution are in equilibrium between and (MoxOy)n-, depending upon the pH, temperature and concentration of the solution.Consequently, a non-uniform distribution of active sites and heterogeneous properties of the catalyst surface may be produced during the immersion and calcination. On the other hand the fixed catalyst was synthesized molecularly taking advantage of the ready reaction between Mo(n-C,H,), and the surface OH groups as expressed in eqn (3)-(6). The decrease in the peak intensity in the OH vibration region and the fact that Mo(n-C,H,), did not form a significant surface complex on the support with no OH groups available under the present experimental conditions, j- indicates that the surface OH groups participate in the Mo-fixing process. The Mo-fixation should be governed by the population of OH groups at the sup- port surface.Therefore the distribution of molybdenum could be controlled by the number and topography of the surface OH groups. The silica gel used was found to have nearly one OH group per 100 A2 on average, according to the method reported by Sat0 et aL9 using (C2H5),A1. Peri demonstrated, on the -f Reaction of MO(~T-C~H~)~ with the silyl ether bond of the SOz surface may occur, but its rate seems to be small.Y. IWASAWA, Y. NAKANO AND S . OGASAWARA 2971 basis of a stoichiometric reaction of AlCl, or SiC14 with surface OH groups, that silica gel largely holds paired OH groups even after the heat-treatment at 600°C. This indicates that Mo-Wng mainly proceeds by eqn (3) and that step (7) may play a lesser role in the fixing reaction under the present conditions.Consequently, the molybdenum ions may be atomically dispersed at the silica surface with chemical bondings through oxygen atoms of the silanol groups : the mean distance between the nearest-neighbour molybdenum ions in the fixed catalyst is estimated to be w 20& assuming ideal distribution of the paired OH groups. Indeed, all the molybdenum ions participated in the uptake of H2 or O2 in the stepwise redox process on the fixed catalysts, (IV) e (111) e (11), as shown in fig. 2 : the behaviour observed in the redox steps was very similar to that found by Yermakov et aL6 The molybdenum ion thus fixed [catalyst(IV)] was confirmed to be in a hexavalent level by X.P.S. emission data. The interaction of molybdic acid with the surface OH groups on the supports has been postulated ideally l1 on the basis of an acid-base interaction accompanied by the elimination of water as follows ; H H I I 0 0 + 0 0 - \& 0 ' ' 0 I H I H However, since an anion of paramolybdenic acid is usually used in the impregnation, the scheme according to Yamagata et aE.,13 must be modified as follows: H H I I + n OH'.anion of para- "-A paramolybdenic + (mdybdcnic acid) This form then follows activation by calcination. Yamagata l2 also demonstrated that molybdenum ions are supported on the sites at which anionic OH groups are located. The presence of acidic OH groups resulted in the precipitation and deposition of some molybdenum species. Castellan et aZ.I3 showed the presence of silicomolybdic acid on the Mo-impregnated silica catalyst where a tetrahedral molybdenum species was also found.With higher molybdenum content a poly- molybdate species was also observed. Thus these investigations show the hetero- geneous surface state of the impregnation catalysts. The situation of the Mo-fixed catalyst is different because Mo(n-C,H,), may react with the hydrogen (which is protonic) of the OH groups rather than the basic OH groups, followed by the removal of propene molecules. The Mo-fixed catalyst may have active sites in a different environment from the conventional impregnation catalyst. The dependence of the luminescence from the impregnation catalyst upon the excitation energy of the charge transfer band of the Mo=O bond, as shown in fig. 4, indicates that there must be at least two different emission centres.14 On the other hand the emission energy of the Mo-fixed catalyst (IV) was independent of the excital tion energy change, reflecting the homogeneity of the molybdenum emitting species at the surface.It is concluded from observations of the photoluminescence that2978 PROPERTIES OF Mo-FIXED CATALYST the active species in the Mo-fixed catalyst has a uniform environment, where the surface of the conventional impregnation catalyst was heterogeneous. The quench- ing behaviour of oxygen and propene molecules clearly support this.I5 The U.V. reflection spectrum of the fixed catalyst (IV) with absorption peaks at 290 and 225 nm shows that the Mo ion is situated in a tetrahedral position :I6 the 290 and 225 nm peaks are assigned to the transition of 3t, c tl and 3t2 - 2t,, respectively, in a tetrahedral c~mplex.~’ This is in agreement with the structure expected from the synthesis scheme.The absorption band at 916 cm-1 could be assigned to an asymmetric stretching vibration of the molybdenum-oxygen bond of either the tetrahedral or octahedral oxomolybdenum(6 + ) species. Mitchell et aZ.I demonstrated that the frequency of a Mo=O bond vibration depended upon the number of the coordination and the terminal oxygen atoms on the molybdenum : the oxomolybdenum species with two terminal oxygen atoms in tetrahedral coordination had a wavenumber near 920 cm-l. The frequency also depends upon the degree of distortion of the tetrahedral structure. Thus the absorption peak at 91 6 cm-’ indicates a tetrahedral dioxomolybdenum structure.Consequently, the collective results of U.V. reflectance spectroscopy, i.r. spectro- metry, the photoluminescence study, the X.P.S. emission peak and the volumetric H2/0, uptake confirm, as expected from the step-wise synthesis procedure for the Mo-fixed catalyst preparation, that a molybdenum ion in the fixed catalyst (IV) which is highly distributed in a hexavalent level has the dioxo-structure with homo- geneous character, and is in tetrahedral coordination at the silica surface, while a molybdenum ion in the conventional impregnation catalyst is heterogeneous in the surface chemical state. These differences in the surface properties of the fixed and impregnated catalysts should correlate with the differences in their catalytic activities and their selectivities in ethyl alcohol oxidation. Investigation of a fixed catalyst with a well-defined active site give clearer information on the reaction mechanism.Fig. 5 and table 1 show that the Mo-fixed catalyst (IV) is very active and selective in ethyl alcohol dehydrogenation compared with the conventionally impregnated catalyst. Lower-valent molybdenum ions showed much less activity. Fig. 6 shows that hexavalent Mo ions are active species for dehydrogenation. Again the linearity in the relation of the amount of Mo6+ with activity indicates that Mo ions in the fixed catalyst have uniform activity. The rate of oxidation to form acetaldehyde was reduced as the number of Mo6+ ions decreased in the absence of oxygen, while the activity was restored by the admission of oxygen into the system in the course of the ethyl alcohol dehydrogenation as shown in fig.7. Steady-state oxidation pro- ceeded in the presence of oxygen. Thus ethyl alcohol oxidation seemed to proceed by the two-stage redox mechanism expressed in the eqn (8) and (9). In the steady- state procedure the reaction rate is given by eqn (lo), which is confirmed by fig. 8. The fixed catalyst (111) with a tetravalent Mo ion could be produced as needed as above mentioned, while in the case of a conventional impregnation catalyst, controlled formation of uniform Mo4+ species may be difficult. The rate of oxidation of the Mo4+-fixed catalyst with oxygen was determined volumetrically. The rate of reduction of the Mo6+ species to Mo4+ was obtained under oxygen-free conditions.Accordingly, the rate constants, kl and k2, in eqn (8) and (9) could be determined independently. Table 3 shows good agreement in the values obtained from the steady-state eqn (10) and the individual measurements of the two steps. These results prove that ethyl alcohol oxidation took place in the redox reaction between the Mo6+ and Mo4+ species.Y. IWASAWA, Y. NAKANO AND S. OGASAWARA The reaction scheme may be shown in more detail as follows : 2979 0 0 ' \o I Si 1 Si /I\ /I\ 0 II 4 " O ' . I Si Si /I\ /I\ + C~HSOH H H I 0 (p: Si /I\ Si " f2H5 I 0 0 \I' /O Mo 0 / \o I Si I Si /I\ /I\ CH3CHO ( 3 ) Si Si /I\ /I\ Dissociation of the 0-H bond in alcohols on metal oxide catalysts is generally easy. Indeed, the i.r. spectra showed a rapid decrease in the intensity of the Mo=O vibration band with exposure to ethyl alcohol even below 100°C, at which tempera- tures formation of acetaldehyde was very slow.This implies that step (l), which is the dissociation of the OH bond of ethyl alcohol accompanied by the change of the Mo=O double bond to an Mo-0 single bond, is not rate-controlling. The absorp- tion bands in the steady-state oxidation of ethyl alcohol followed by evacuation at 150°C appear at 1446, 1385 and 1373 cm-1 in the region of the CH3 and CH2 de- formation modes besides the 2982,2935 and 2900 cm-l bands for the C-H stretching vibrations. These absorption peaks arise from the CH,CH,O-group.l The silica support showed no peak with significant intensity in this region under similar experimental conditions.Accordingly ethyl alcohol molecules dissociatively adsorb on dioxomolybdenum(6 + ) species forming ethoxide group. The ethoxide group in the i.r. spectrum decreased with reaction time at 150°C. The rate-determining step in oxidation of ethyl alcohol to form acetaldehyde should be the abstraction of a hydrogen atom (hydride) from the C-H bond, as shown in the step (3) of the scheme. The formation of water was also observed in the absence of oxygen, indicating that this process must occur prior to oxidation of the Mo4+ species by oxygen. No evolution of hydrogen was observed during the oxidation in spite of the presence or absence of oxygen. Consequently, ethyl alcohol oxidation takes place in conjunction with the redox cycle of Mo6+ + Mo4+, where the dioxomolybdenum(6+) structure is the active species.The oxidations of methyl alcohol and ethyl alcohol have been explained on the basis of a two-stage redox mechanism on MoO~,~O V205-M003,21 Fe-molybdate,22 Th-molybdate, 23Sn02-Mo03,24 e t ~ . ~ Our conclusion gives clear evidence for the mechanism of alcohol oxidation. Trifiro and Pasquon 26 demonstrated that the selective oxidation of methyl2980 PROPERTIES OF MO-FIXED CATALYST alcohol could be connected with the presence of a Mo=O terminal bond which has the character of a labilized double bond. The present results on a well-defined k e d catalyst prepared by stepwise synthesis clearly support this idea. Furthermore, the coordinatively unsaturated dioxomolybdenum(6 + ) which has uniform dis- tribution in a tetrahedral coordination is demonstrated to be the active species in the selective oxidation of ethyl alcohol.Ethyl alcohol molecules readily attack the unsaturated coordination sphere of the tetrahedral dioxomolybdenum ion and are activated for dehydrogenation. The structure and environment of the Mo6+ ion may affect the electronic property (a degree of c~valency)~’ of the Mo-0 bond as the active site : hence the abstraction of hydride from the methylene group of ethyl alcohol becomes easier on an Mo-fixed catalyst than on a conventional impregnation catalyst. The high activity of the coordinatively unsaturated dioxomolybdenum, which has uniform character in the tetrahedral position seems to be general in de- hydrogenation processes with various reactants.Investigation of the allyl-type oxidation of propene will confirm this.l5 The present study on an Mo-fixed catalyst with a well-defined surface state has many advantages for investigation of the essential factors of catalysis, the structure and environment of the active site, and the reaction mechanism. The authors thank Prof. Y. Kondo of Rikkyo University for measurements of the U.V. diffuse reflectance spectra and Dr. M. Soma of Tokyo University for measurements of the X.P.S. emission spectra. Part of this work was financially supported by the Asahi Glass Foundation. Y. I. thanks the Sakko-kai Foundation for a subsidy. The preliminary results were reported by Y. Iwasawa et al., Shokubai, 1977, 19,4. The term “ fixed catalyst ” is proposed to indicate a well-defbed supported catalyst with chemical bonding between the transition element and the support in order to distinguish this type of catalyst from conventional impregnation catalysts.For examples of Mo-containing catalysts, W. K. Hall and M. LoJacono, Proc. 6th Int. Congr. Catalysis (London, 1976), 1977,1,246 ; N. Giordano, J. C. J. Bart, A, Vaghi, A. Castellan and G. Martinott, J. Catalysis, 1975,36,81 ; T. Fransen, 0. Van Der Meer and P. Mars, J. Catalysis, 1976,42,79 ; G. T. Pott and W. H. J. Stork, in Preparation of Catalysts, ed. B. Delmon, P. A. Jacobs and G. Poncelet (Elsevier, Amsterdam, 1976), p. 548 ; V. M. Vorotyntsev, V. A. Shvets and V. B. Kazanskii, Kinetics and Catalysis, 1972, 12, 1108. G. H. Ballard, Advances in Catalysis (Academic Press, N.Y., 1973), vol.23, p. 263. Yu. I. Yermakov, Catalysis Rev.-Sci., 1976, 13, 77 ; Yu. I. Yermakov, B. N. Kuznetsov and Yu. A. Ryndin, J. Catalysis, 1976, 42,73, and references therein. G. Wilke, B. BogdanoviC, P. Hardt, P. Heimbach, W. Keim, M. Kroner, W. Oberkirch, K. Tanaka, E. Steinriicke, D. Water and H. Zimmermann, Angew. Chem. Ini. Edn, 1966,5, 151. Yu. I. Yermakov and B. N. Kuznetsov, Prep. 2nd Japan-Soviet Catalysis. Seminar, (Tokyo), 1973, p. 65. ’ Yu. I. Yermakov, B. N. Kuznetsov, L. G. Karaktchiev and S. N. Derbeneva, Kinetica i Kataliz, 1973, 14, 709. G. A. Tsigdinos and C. J. Hallada, Molybdenum Chemicals-Chemical Data Series Bulletin, 1969,14 Feb. ; A. K. Babko and G. I. Gridchina, Zhur. neorg. Khim., 1968,13,61(Eng). M. Sato, T. Kanbayashi, N. Kobayashi and Y.Shima, J. Catalysis, 1967, 7, 342. lo J. B. Peri and A. L. Hensley Jr., J. Phys. Chem., 1968, 72,2926. l1 M. Dufaux, M. Che and C. Naccache, J. Chim. phys., 1970, 67, 527 ; G. S. John, M. J. Den Herder, R. J. Mikorsky and R. T. Wasters, Advances in Catalysis (Academic Press, N.Y., 1957), vol. 9, p. 225. l2 N. Yamagata, Y. Owada, S . Okazaki and K. Tanabe, J. Catalysis, 1977, 47, 358. l3 A. Castellan, J. C. J. Bart, A. Vaghi and N. Giordano, J. Catalysis, 1976, 42, 162. l4 G. T. Pott and W. H. J. Stork, CataIysis Reviews (Marcel1 Dekker, 1976), p. 163 ; S. Coluccia, M. Deane and A. J. Tench, Proc. 6th Int. Cong. Catalysis (London, 1976), 1977, 1, 171 ; V. B. Kazansky, A. M. Gritscov and V. A. Shvets, Prep. 2nd Japan-Soviet Catalysis Seminar (Tokyo), 1973.Y . IWASAWA, Y . NAKANO AND S . OGASAWARA 298 1 l5 Y. Iwasawa and S. Ogasawara, to be published. l6 M. Che, F. Figueras, M. Forissier, J. McAteer, M. Perrin, J. L. Portefaix and H. Praliand, Proc. 6th Int. Cong. CatalysisLondon (1976), 1977, 1,261 ; G. N. Asmolov and 0. V. Krylov, Kinetics Catalysis, 1971, 11, 847. l7 M. Wolfsberg and L. Helmholz, J. Chem. Phys., 1952, 20, 837. l 8 P. C. H. Mitchell and F. Trifiro, J. Chem. SOC. (A), 1970, 3183. l9 Y. Soma, T. Onishi and K. Tamaru, Trans. Faraday SOC., 1969, 65, 2215 ; R. G. Greenler, 2o G. K. Boreskov, B. I. Popov, V. N. Bibin and E. S. Rozischnikova, Kinetika i Kataliz, 1968,9, 2 1 R. S. Mann and M. K. Dosi, J. Catalysis, 1973, 28, 282. 22 J. Edwards, J. Hicolaidis, M. B. Cutlip and C. 0. Bennett, J. Catalysis, 1977, 50, 24; M. Carbucicchio and F. Trifiro, J. Catalysis, 1976, 45, 77 ; P. Jiru, B. Wichterlova and J. Tichy, Proc. 3rd Znt. Cong. Catalysis (Amsterdam), 1965, 1, 199. 23 V. Srihari and D. S. Vismanath, J. Catalysis, 1976, 43,43. 24 M . Niwa, M. Mizutani and Y. Murakami, Chem. Letters, 1975, 1295. 25 J. Novakova, P. Jiru and V. Zavadil, J. Catalysis, 1971, 21, 143 ; P. L. Villa, A. Szabo, F. 26 F. Trifiro and I. Pasquon, J. Catalysis, 1968, 12,412. 27 G. K. Boreskov, Proc. 5th Int. Congr. Catalysis, 1972, 2, 981. J. Chem. Phys., 1962,37,2094. 796 ; J . Novakova, P. Jiru and V. Zavadil, J. Catalysis, 1970,17, 93. Trifiro and M. Carbucicchio, J. Catalysis, 1977, 47, 122. (PAPER 8 1462)

ISSN:0300-9599    

DOI:10.1039/F19787402968

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Preparation of High Surface Area Reduced Molybdenum Oxide Catalysts BY ROBERT BURCH* Department of Chemistry, The University, Whiteknights, Reading RG6 2AD Received 20th March, 1978 High surface area molybdenum oxides have been prepared by the thermal decomposition and reduction of molybdenum (VI) oxalate. It has been found that the initial trioxide formed is highly oxygen deficient and that the composition depends on the method of preparation, varying from for an oxide prepared in vacuum. The rate of reduction was found to be dependent on: the method used to decompose the oxalate ; the temperature of the decomposition ; the partial pressure of hydrogen ; and the partial pressure of water vapour. Vacuum-prepared trioxides are reduced in a single stage to MOO^-^. Nitrogen-prepared trioxides reduce first to Mo401 or MozOs depending respectively on whether " wet " or '' dry " conditions are used.In most cases the reduction-time curves are essentially sigmoidal in character. Possible reduction mechanisms and rate-determining steps are discussed. for an oxide prepared by decomposing the oxalate in nitrogen, to ~ ~~ ~~ The platinum metals have been used for many years as catalysts for reforming and isomerization reactions of hydrocarbons. Recent attempts to improve the existing platinum catalysts have centred on multimetallic catalysts of which the platinum- rhenium reforming catalysts are perhaps the best known and most important. Both platinum and rhenium, however, are rare and expensive metals, so there are good reasons for trying to find alternative materials which could be used to catalyse reforming and isomerization reactions.Recently published work on clean metal films indicates that the group 6 transition metals have high activity for catalysing some hydrocarbon reactions. Supported oxides of these metals are difficult if not impossible to reduce to the metallic state,2 and unsupported oxides, although they can be reduced, usually have low surface areas. Consequently, little is known about the catalytic activity of these metals in a highly dispersed form. The reduction of low surface area molybdenum trioxide has been investigated previously by a number of workers 3-9 but no results have been published on the reduction of high surface area oxides. Low surface area oxides reduce very slowly, and molybdenum metal can only be prepared in reasonable times by using high temperatures (typically > 900 K) which, of course, produce a low surface area material. The thermal decomposition and reduction of molybdenum (VI) oxalate produces high surface area oxides.lo* These high surface area unsupported molybdenum metalfmolybdenum oxide materials have been shown previously l2 to have high activity and selectivity for the isomerization of straight chain hydrocarbons to the corresponding branched chain isomers.The origin of the catalytic activity in these materials and the mechanism of the isomerization reaction are currently being investigated. In this paper we describe the preparation of the high surface area reduced oxides and show how the kinetics of reduction are affected by the experimental conditions.2982R . BURCH 2983 Subsequent papers will describe the physical structure of the solids,13 and how this is created during the decomposition and reduction reactions, and the variation in cat- alytic activity with the physical and chemical properties of the materials. EXPERIMENTAL MATERIALS Molybdenum (IV) oxalate (Climax Molybdenum), nominal composition H2(M003(C204)(H20)) H20, was used as a starting material. This was crushed and sieved, and the powder having particle sizes between 150 and 250 pm was used in all the subsequent experiments. Hydrogen and nitrogen (B.O.C. 99.999 % purity) were passed through a liquid nitrogen trap before use. No further purification appeared to be necessary judging by the fact that the reduced samples were invariably extremely pyrophoric.EQUIPMENT Most of the thermal decomposition experiments were performed in situ in a Cahn vacuum Samples varying in weight from 20 to 100mg were loaded into a platinum crucible Occasional batches of trioxide were prepared by decomposing about 10 g of the oxalate micro balance. (diameter 8 mm, depth 5 mm) suspended by a fine nichrome wire from the balance arm. in a tube furnace either under vacuum or under static nitrogen. RESULTS Several different experimental parameters have an effect on the rate of reduction of molybdenum trioxide. These are : (i) the method used to decompose the molyb- denum oxalate; (ii) the temperature of reduction; (iii) the partial pressure of hydrogen ; (iv) the partial pressure of water.The effect of these various parameters is summarised in fig. 1. METHOD USED TO DECOMPOSE THE OXALATE Decomposition at 573K in nitrogen produces a small degree of oxygen de- ficiency MOO^.^^), whereas decomposition at 573 K under vacuum gives a compound of nominal composition MOO^.^^. At 673 K, decomposition in nitrogen produces a compound of composition MOO^.^^. These different " trioxides " also have important structural differences from each other. Decomposition at 573 K in nitrogen (or in hydrogen) gives a mainly meso- porous material with a B.E.T. surface area of about 100 m2 g-l. The oxide pre- pared under vacuum contains mainly micropores, and has a B.E.T. surface area of 40m2g-l. It should be noted that, because of the microporosity of both the nitrogen (and vacuum) prepared oxides, the B.E.T.surface areas must be regarded as upper estimates of the true surface area. This will be particularly true for the vacuum-prepared oxide, for which the true surface area may be perhaps 20 % smaller than the B.E.T. surface area. In contrast to the porous trioxide prepared at 573 K, decomposition at 673 K gives an essentially nonporous oxide with a small surface area (10 m2 g-l). The difference in surface area between the nitrogen (and vacuum) prepared tri- oxides might be expected to result in a similar difference in reactivity during reduction with hydrogen. In fact, as can be seen from fig. 2(a), exactly the reverse trend is observed. The vacuum-prepared oxide is significantly more reactive than the nitrogen- prepared material even allowing for the fact that the compositions of the starting2984 MOLYBDENUM OXIDE CATALYSTS 10 m2g-1 FIG.1.-Products formed from the decomposition and reduction of molybdenum oxalate under different experimental conditions. OF---- time/h FIG. 2 . 4 2 ) Reduction of vacuum-prepared (curve A) and nitrogen-prepared (curve B) molybdenum trioxide at 573 K and lo5 Pa hydrogen. (b) Reduction of nitrogen-prepared trioxides at 673 K (curve C) and 573 K (curve D) and lo5 Pa hydrogen. (c) Reduction of molybdenum trioxide at 573 K at various hydrogen pressures : E, 91.3 ; F, 70.7 ; G, 49.4 ; H, 23.2 kPa. (d) Reduction of molybdenum trioxide at 673 K at various hydrogen pressures : J, 100.2 ; K, 71.4 ; L, 48.4 ; M, 17.7 kPa. (e) Reduction of molybdenum trioxide in dry conditions at 573 K and lo5 Pa hydrogen.ct is the extent of reduction to molybdenum dioxide.R . BURCH 2985 materials are different. Thus, the time required for reduction from 20 to 100 % to occur is 3.2 h for the nitrogen-prepared oxide and 2.2 h for the vacuum-prepared oxide. Clearly the reactivity is not related only to the surface areas of the samples. The vacuum-prepared oxide must contain a greater concentration of active sites at which nucleation of the product phase can occur. This may well be associated with the microporosity of these materials and with the high degree of oxygen deficiency in the initial trioxide (MOO,. 2). The large oxygen deficiency in the vacuum-prepared trioxide does not appear to arise from any kinetic effects during the thermal de- compositions of the oxalate because, in separate experiments in which a nitrogen- decomposed oxide is subsequently evacuated at 573 K for several hours, the maximum oxygen deficiency corresponds only to MOO^.^^.A very high degree of oxygen deficiency appears only to be possible if it is incorporated into the oxide during the decomposition of the oxalate. EFFECT OF CHANGING THE REDUCTION TEMPERATURE Fundamental changes in the physical structures of the samples caused by heating at higher temperatures prevented the evaluation of the activation energy for the reduction reaction. For example, it was observed that on heating a nitrogen- prepared oxide sample from 573 to 673 K the surface area decreased from 100 to about 10 m2 g-l, and the porosity of the oxide was completely eliminated.The consequences of these structural transformations in terms of the reactivity during reduction are illustrated in fig. 2(b), which compares the rate of reduction for two samples taken from the same batch of nitrogen-decomposed oxalate. Although the oxide at 673 K has a much higher oxygen deficiency the rate of reduction at 673 K is only marginally faster than at 573 K. In fact, if we compare the reduction curves for the two oxides which have similar oxygen deficiencies, that is curve A in fig. 2(a) and curve C in fig. 2(b), we observe that the reduction is faster at 573 than at 673 K. Although these two samples have approximately the same initial chemical com- position they have very different physical structures. The 573 K oxide is micro- porous whereas the 673 K oxide is nonporous. We presume that this indicates that the oxygen deficiency is accommodated in different ways in the two oxides.In the 673 K oxide the deficiency may be contained almost exclusively by shear plane formation MOO^.^^ corresponds to the shear phase Mo,02,), whereas in the 573 K oxide there may be a high concentration of oxygen vacancies at or close to the surface of the micropores. This would suggest that the rate of reduction is controlled by a surface process and is therefore dependent on the number of nucleation sites. If oxygen diffusion was the rate controlling step we would expect to find a much higher rate of reduction at the higher temperature. Other evidence for a surface-controlled reaction will emerge later.EFFECT OF HYDROGEN PRESSURE A series of experiments was performed on fresh samples taken from a batch of nitrogen-decomposed oxalate, in order to determine the effect on the rate of reduction of varying the hydrogen partial pressure. The rate of reduction both at 573 K [fig. 2(c)] and at 673 K [fig. 2(4] increases linearly with the hydrogen pressure. We shall discuss later the implications of these results in terms of reaction mechanisms.2986 MOLYBDENUM OXIDE CATALYSTS EFFECT OF WATER VAPOUR An experiment was performed in which the sample pan in the microbalance was surrounded by about 20 g of previously outgassed high surface area alumina. The purpose of this experiment was to see whether the rate of reduction would be affected by removing, as efficiently as possible, all the water molecules in the vicinity of the sample.The reduction curve obtained in this experiment is shown in fig. 2(e). Instead of the normal sigmoid curve obtained with earlier samples we observe a double sigmoid curve. Furthermore, at point X, when the reduction had proceeded to MOO^.^, the microbalance was evacuated with the sample at 573 K. Hydrogen was reintroduced, and the rate of reduction was initially much faster than previously. The evacuation treatment may have created new active sites on the surface of the reduced oxide, or it may have removed adsorbed water which was inhibiting the reduction reaction. The physical structure of the reduced oxide prepared under dry conditions is also quite different from the structure of any of the other reduced materials in that it contains a large concentration of micropores.The presence of water during reduc- tion seems to have three effects. It affects the shape of the reduction curve, and presumably the mechanism of the reduction reaction, it affects the pore size distribu- tion of the reduced product, and it appears to inhibit the reduction reaction. DISCUSSION OXYGEN DEFICIENCY I N MO03-x The first unusual aspect about this work which requires some explanation is the high degree of oxygen deficiency which is observed in all the trioxide starting materials, but particularly in the case of the vacuum-decomposed oxalate samples. For low surface area molybdenum trioxide samples which contain isolated oxygen vacancies the limiting composition is about Higher levels of oxygen deficiency are accommodated by the elimination of vacancies with the creation of shear ~1anes.l~ Various shear compounds of molybdenum oxides are known (Mo401 M O ~ O ~ ~ , Mo9OZ6 etc.) but none of these would be expected to form merely by heating moly- bdenum trioxide under vacuum at 573 K.A few disordered shear planes might be created under these conditions but the overall composition would not fall below, say, M002.g9. On thermodynamic grounds, no lower oxide should be formed in the absence of a reducing agent. One possible reason for the high degree of oxygen deficiency of our samples could be the reducing action of hydrocarbon molecules originating from the grease used on the ground joints of the microbalance.We do not think that this is correct for a number of reasons. First, the maximum oxygen deficiency is observed on samples which have been heated for only a few minutes at 573 K. Other samples have been heated for longer periods and yet have not been reduced to the same extent. Second, the degree of oxygen deficiency appears to be determined mainly during the decomposition reaction and not as a result of outgassing at 573 K. Third, in all cases the weight of the oxygen-deficient samples becomes constant after only a few minutes at 573 K, and then remains constant for several hours at least. If reduction by hydrocarbon vapours was important we would not expect a reaction which was initially very rapid to suddenly stop completely. An alternative, and we think a more acceptable, explanation of the high degree of oxygen deficiency is that this arises because of the high surface area and porosity of our materials.We consider that the oxygen deficiency i s accommodated partlyK. BURCH 2987 by the formation of shear planes and partly by the creation of oxygen vacancies in or near the surface of the oxide particles. The evidence in support of this proposal is twofold. First, we observe the highest oxygen deficiency for the vacuum-prepared trioxide which in addition to being very microporous is also the most reactive during reduction. Second, even if we assume that for the compound (surface area 40m2 g-') all the oxygen vacancies occur in the surface this would still be reasonable because of the large surface area of the oxide.It would be equivalent in terms of the overall composition of having a composition of M002.995 for an oxide with a surface area of 1 m2 gl. Further support for the concept of having a large number of oxygen vacancies in the surface of the trioxide is found in the published l i t e r a t ~ r e . l ~ - ~ ~ Both Haber et aZ.15 and Cimino et a1.16 have used X-ray photoelectron spectroscopy to study the surface of molybdenum trioxide after very mild reduction treatments. The con- clusions of both sets of workers are essentially the same. They have shown that in the very early stage of the reduction lattice oxygen ions close to the surface are removed to create isolated oxygen vacancies, with the electrons being localised on an adjacent Mo6+ ion forming Mo4+. When the concentration of oxygen vacancies exceeds an unspecified critical value, strings of edge-sharing MOO octahedra are created from the original corner-sharing octahedra present in stoichiometric MOO,.(This is the precursor state for the development of the shear compounds which are characteristic of bulk MOO,.) The essential feature of this published work is that it shows that the surface of molybdenum oxide need not necessarily have the same chemical or physical structure as the bulk oxide. Even a very mild treatment can create in a few minutes a surface reduction which would require many hours to produce throughout the bulk oxide. This is further substantiated by the observation l5 that a thin layer of molybdenum metal is formed on low surface area Moo3 under conditions of temperature (823 K), hydrogen pressure (5 x Pa) and time (1 h) where the rate of reduction is so slow that no bulk metal would be formed.We conclude from this that although only a very small concentration of oxygen vacancies can be accommodated in the bulk of MOO,, a much higher concentration of oxygen vacancies can be present on the surface of an oxide particle. If, as in our case, the oxides have large surface areas and provided that the oxygen deficiency is built in during the genesis of the trioxide, then large oxygen deficiencies are possible. We may contrast the compositions of the nitrogen-prepared trioxide MOO^.^,) and the vacuum-prepared trioxide MOO^.,^) and speculate that MOO^.^^ cor- responds to the limiting composition due to shear plane formation in a high surface area oxide, and that therefore the lower oxygen content of reflects a higher concentration of oxygen vacancies at the surface.This, as we shall see presently, should result in the two types of oxide having different reactivities. SHAPE OF REDUCTION AGAINST TIME CURVES There are several features of the reduction against time curves shown in fig. 2 which require discussion. Curve A has a smooth sigmoidal shape which is charac- teristic of a reaction in which the rate is controlled either by the nucleation and growth of the product phase, or by an autocatalytic reaction. Curve B has a discontinuity at a composition in the vicinity of MOO^.,^, which would correspond to the shear phase Mo4011, and it may be that on reduction MOO^.^, goes first to Mo4OI1 and then to Moo2-,.We note also that the reduction does not stop at MOO, but con- tinues until a composition of about MOO,., is reached, at which point a much slower2988 MOLYBDENUM OXIDE CATALYSTS linear reduction is obtained. Reduction towards bulk molybdenum metal does continue, nevertheless, even at these low temperatures. Curve A does not show a discontinuity at MOO^.,^. This may simply be because it is difficult to observe a discontinuity so close to the start of the reduction of On the other hand it is also possible that the reaction does not pass through Mo,Oll because of the different structure of the starting material. An unusual reduction against time curve is obtained [fig. 2(e)] when the reaction is performed at 573 K in dry conditions.The reduction seems to proceed through a highly oxygen deficient Mo205 phase as an intermediate compound rather than Mo,O, , as before. We appear to have observed three types of reduction process; Moo2.,, going directly to MOO,-, in a single step; MOO,.,, going first to Mo40,, and then to Moo2-, under wet conditions; MOO^.^^ going to Mo205 and then to MOO,-, in dry conditions. The differences between the reduction of Mo02.B2 and M002.g8 axe perhaps not surprising bearing in mind the different structures, surface areas, and chemical compositions of the starting materials. However, the differences between the wet and dry reductions of MOO^.^^ are certainly unexpected. The product of the dry reduction is much more microporous than the product of the wet reduction. This indicates that water vapour is rapidly removed from the oxide surface. Possibly the efficient removal of water vapour from the surface stimulates the oxide to lose further oxygen and thus by-pass Mo4OI1.Reduction at 673 K (curves C and J) appears to give a parabolic curve, but we suspect that this is only because the initial acceleratory step is very short and not easily detected. Thus, when reduction at 673 K is performed at lower partial pres- sures of hydrogen (curves L and M) we again observe a smooth sigmoidal reduction curve. In all cases the reduction proceeds in a single step from MOO^.,^ to MOO,-,. A further interesting observation in the reduction experiments at 673 K is that at low hydrogen partial pressures where the rate of reaction is slow the final product (taken to be the composition at which the reduction curve almost levels off) is much more highly reduced than are any of the other products obtained at higher hydrogen pressures.This is a reproducible effect. We have here the unusual situation that an initially slower reduction reaction can eventually result in a more highly reduced material. We do not have any certain explanation of this effect, although it is possible that, because reduction is slow, water vapour produced as a product is more readily desorbed from the surface, and therefore causes less damage to the active surface of the reduced product so that it in turn becomes reducible. MECHANISM OF REDUCTION It has been pointed out earlier that there are significant differences in the physical structures of the three types of trioxide which we have been investigating.We must be cautious, therefore, in analysing the kinetic data in terms of a single reaction mechanism since the rate-determining step may change from one structure to another. Bearing in mind this reservation it is at least possible to produce a plausible interpreta- tion of the kinetic data based on a common rate-determining step. However, before discussing our results further it will be useful to summarise what is known * at the atomic level about the mechanism of the reduction of molybdenum trioxide. In the first stages of reduction a high concentration of oxygen ion vacancies are created in the surface of the oxide. These vacancies diffuse relatively slowly into the bulk at the temperatures used in the reduction experiments.X-Ray photoelectron spectroscopy shows that Mo4+ (but not Mo5+) ions are created at the same time asR. BURCH 2989 oxygen vacancies are produced. Clusters of edge-sharing octahedra containing Mo4+ ions are formed, and some metal-metal bonding between Mo4+ ions is indicated by the appearance of Mo2+ signals in the electronic spectrum, and by the nature of the MOO, structure which places alternate pairs of metal atoms close together. Further surface reduction to metallic molybdenum occurs relatively easily. The large concentration of shear planes in Moo3-,, is taken by Haber to indicate that the creation of the MOO, lattice, which also contains edge-sharing octahedra, should be facile. He therefore proposes that nucleation of MOO, will be rapid, and the rate-determining step is considered to be the dissociative chemisorption of hydrogen.The results of our work described earlier which are pertinent to a discussion of the reduction reaction are; the observation of a sigmoidal reduction against time curve ; the dependence of the rate of reduction on the hydrogen partial pressure both at 573 and at 673 K ; the high activity of a highly oxygen deficient vacuum-prepared oxide ; the inhibition by water vapour ; and the effect of ageing on the reactivity of the trioxides. All these factors point to a surface process being rate-determining, rather than say oxygen diffusion. Nucleation and growth of a product phase would be the most usual interpretation, but the dissociative adsorption of hydrogen coupled with an autocatalytic reaction would also account for the observed results.On the assumption that nucleation is rate-determining we can explain the vari- ations between the various samples as follows. We suggest that for the vacuum- prepared trioxide there are a large number of oxygen vacancies at the surface, which can migrate and combine to germinate a nucleus of the reduced product. Rather fewer oxygen vacancies exist in nitrogen-prepared samples, and even fewer survive after raising the temperature from 573 to 673 K. The overall result is that reduction of a vacuum-prepared trioxide at 573 K is fast because it contains a large number of nuclei, reduction of a nitrogen-prepared trioxide at 573 K is slower because it contains fewer nuclei ; and reduction of the same trioxide at 673 K is slow because it contains very few nuclei, and the higher inherent activity of each nucleus only just compensates for the much smaller number of nuclei.On the other hand, if we assume that the chemisorption of hydrogen is rate- controlling we would expect a linear reduction against time curve. However, the observed sigmoidal curve could be produced by the autocatalytic effect of small amounts of low valent, or even metallic molybdenum, formed on the surface during the earliest stages of reduction. It is certainly possible that low valent molybdenum will have the ability to activate hydrogen. We do not think that it is possible, from the information available, to distinguish between nucleation and hydrogen adsorption as possible rate-determining steps.We are grateful to the S.R.C. for an equipment grant. We are grateful to the Climax Molybdenum Company, who first suggested this research, for their continuing support. We thank Mr. N. B. Mason for his assistance with some of the early experiments. J. F. Taylor and J. K. A. Clarke, 2. phys. Chem., 1976,103,216. J . R. Anderson, Structure of MetaZlic Catalysts (Academic Press, London, New York, 1975). J. von Destion-Forstmann, Canad. Metallurg. Quart., 1965, 4, 1. Ph. A. Batist, C. J. Kapteijns, B. C. Lippens and G. C. A. Schuit, J. CcztczZysis, 1967, 7, 33. Kh. Vasilev and T. Pencheva, Khim. Ind. Sofia, 1970,18,202. M. J. Kennedy and S. C. Bevan, Proc. 1st Int. Conf. Molybdenum, Reading, 1973, ed. P. C. H. Mitchell, p. 11.2990 MOLYBDENUM OXIDE CATALYSTS 0. Bertrand and L. C. Dufour, Compt. rend. C, 1974,278,315. Chemical Systems, ed. P. Barnet (Elsevier, Amsterdam, 1975), p. 696. T. Fransen, P. C. van Berge and P. Mars, React. Kinetics Catalysis Letters, 1976, 5, 445. lo G. A. Tsigdinos, C. J. Hallada and R. W. McConnell, U.S. Patent 3,912,660 (to Amax:Inc.), October 14, 1975 ; German Patent 2,451,778 (to Amax Inc.), May 28, 1976. l1 G. A. Tsigdinos and W. W. Swanson, Ind. andEng. Chem., to be published. I2 R. Burch and P. C. H. Mitchell, J. Less-Common Metals, 1977, 54, 363. l 3 R. Burch, J.C.S. Faraday I, 1978,74,2991. l4 L. A. Bursill, Proc. Roy. SOC. A, 1969, 311, 267. l5 J. Haber and J. Janus, in Reaction Kinetics in Heterogeneous Chemical Systems, ed. P. Buret (Elsevier, Amsterdam, 1975), p. 737 ; J. Haber, W. Marczewski, J. Stoch and L. Ungier, Ber. Bunsenges. Phys. Chem., 1975, 79, 970 ; J. Haber, Proc. 2nd Int. Con$ Molybdenum, Oxford 1976, ed. P. C. H. Mitchell and A. Seaman, p. 119 ; J. Haber, A. Koslowska and J. Sloczynski, in Reactivity of Solids, ed. J . Wood, 0. Lindquist, C. Haelgesson and N.-G. Vannerberg (Plenum Press, London and New York, 1977), p. 331. * A. Castellan, J. C. J. Bart, A. Bossi and N. Giordano, Reaction Kinetics in Heterogeneous l6 A. Cimino and B. A. de Angelis, J. Catalysis, 1975, 36, 11. l7 J. Grimblot and J. P. Bonnelle, Compt. rend. C, 1976, 282, 399. (PAPER 8/520)

ISSN:0300-9599    

DOI:10.1039/F19787402982

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Structural Characterisation of High Surface Area Reduced Molybdenum Oxide Catalysts BY ROBERT BURCH* Department of Chemistry, The University, Whiteknights, Reading RG6 2AD Received 20th March, 1978 High surface area reduced molybdenum oxides have been prepared by the thermal decomposition and reduction of molybdenum (w) oxalate. Nitrogen physisorption isotherms, analysed by the a,-method, have been used to determine the physical structure of the oxides. The effect of tempera- ture, water vapour pressure and the degree of reduction, on the physical structure of the oxides has been studied. It is found that reduction creates initially slit-shaped micropores, which on further reduction, or in the presence of water vapour, broaden into mesopores. The changes in structure which accompany the passivation or the annealing of reduced oxides are described. High catalytic activity for the hydrogenation, dehydrocyclisation, isomerization, etc., of organic compounds is restricted mainly to the platinum group metals. Platinum itself is a very effective catalyst for a wide range of hydrocarbon reactions, but it is expensive so there is considerable incentive to search for alternative catalysts.There are two approaches to this problem. One is to prepare more efficient, more stable platinum alloy catalysts. Another is to consider whether other transition metals might have useful catalytic activity. Binary compounds of chromium, molybdenum and tungsten (oxides and sulphides) have moderate activity for catalysing hydrocarbon reactions, but very little is known about the catalytic properties of the corresponding highly dispersed metals.This is mainly because of the difficulty in preparing dispersed high surface area materials. In the case of supported catalysts there is a strong interaction between the metal ions and the support (silica or alumina) which makes it extremely difficult to reduce the supported metal ions to the metallic state.l It may be possible to overcome these problems by using different supports, for example graphite, or by deactivating the oxide supports before impregnation with solutions of the metal ions. Before embarking on such research it would seem appropriate to determine whether or not these transition metals have any useful catalytic properties which could be exploited by using new or modified supports.Published work on clean metal films indicates that a number of the early transition metals have high activity for catalysing hydro- carbon reactions. In order to determine whether this high activity could be generated under less stringent experimental conditions it was decided to investigate the prepara- tion of high surface area, unsupported molybdenum metal/molybdenum oxide catalysts, prepared by the thermal decomposition and reduction of molybdenum(v1) ~ x a l a t e . ~ ~ These materials have been shown previously to have high activity and selectivity for the isomerization of n-hexane to branched chain isomers. The origin of the catalytic activity and the mechanism of the isomerization reaction are currently being investigated. The kinetics and mechanism of the thermal decomposition and reduction of molybdenum(vI) oxalate have been described elsewhere.6 We describe in this paper the way in which the pore structure and the surface area of the reduced molbydenum oxides evolves during the progressive reduction of 29912992 MOLYBDENUM OXIDE CATALYSTS molybdenum trioxide at low temperatures.We also describe how the physical structure of the materials is affected by the experimental conditions used during the preparation. Changes in structure have been followed by measuring nitrogen adsorption isotherms. EXPERIMENTAL MATERIALS The molybdenum metal/molybdenum oxide samples were prepared by the thermal decomposition and reduction of molybdenum(v1) oxalate, Ht [MoO3(C2O4)H2O] . HzO (Climax Molybdenum).Details of the kinetics of the reduction reaction have been described elsewhere.6 Hydrogen for reduction, and nitrogen for the adsorption isotherms, were B.O.C. 99.999 % purity, and were used without further purification, except to bleed the gases slowly into the apparatus through a liquid nitrogen trap. The reduced oxides were invariably extPemely pyrophoric, which indicated that the purity of the gases was completely satisfactory. EQUIPMENT All the decomposition, reduction and nitrogen adsorption experiments were performed in situ in a Cahn R.G. electrobalance. Samples of molybdenum oxalate were ground and sieved (< 150 pm mesh size) and portions of about 100 mg placed in the bottom of a platinum crucible (8 111111 diameter by 5 mm deep) suspended by a nichrome wire from the arm of the microbalance.The microbalance was fitted with a vacuum jacket and could be evacuated to Pa using an Edwards E02 oil diffusion pump. For the nitrogen adsorption isotherms pressures were measured with a Bell and Howell pressure transducer. The temperatwe of the liquid nitrogen bath was measured periodically with a nitrogen vapour pressure thermometer. To ensure that the temperature of the sample corresponded as cIosely as possible with the temperature of the cryostat, the sample was suspended just above the flat bottom of a narrow glass hangdown tube (outside diameter 14 mm). The hangdown tu& in turn was immersed to a constant depth of 25 cm in the liquid nitrogen. RESULTS AND DISCUSSION DECOMPOSITION PRODUCTS OF MOLYBDENUM OXALATE The decomposition of molybdenum oxalate was performed by warming the samples slowly to 573 K in nitrogen (sample MX-l3N), hydrogen (MX-l2H), or vacuum (MX-29V).In the latter case the rate of heating was controlled so as to maintain the pressure below 5 x Pa. The trioxide formed in each case is slightly oxygen deficient, and the compositions determined from the weight increase observed after heating the samples to constant weight in air at 800 K were MOO^.^^ (nitrogen), Moo2+ (hydrogen), MOO^.^^ (vacuum). When decomposing the oxalate in hydrogen the reaction was stopped at M002.g6 by quenching the sample. The nitrogen adsorption isotherms obtained for these oxides are shown in fig. l(a). The isotherms have little in common except for the presence of a hysteresis between the adsorption and desorption branches of the isotherm, indicating some mesoporosity in all the samples. However, both the size and particularly the shape of the hysteresis loops are quite different.The hysteresis loop for MX-13N approxi- mates to type A in the de Boer classification which would indicate that the pores were cylindrical, whereas the hysteresis loops for MX-12H and MX-29V are more like type B corresponding to slit shaped pores. The difference between MX-13N and MX-12H may be due to the higher oxygen deficiency in MX-12H. Although the bulk compositions only differ by 0.02 in the oxygen to molybdenum ratio the surface may have undergone a proportionately much larger change.R. BURCH 2993 The overall shape of the isotherms again show differences.The isotherm for MX-13N is almost type I1 in the Brunauer Deming, Deming and Teller classification,* MX-12H is type IVY and MX-29V is almost type I. These differences in shape reflect differences in the pore size distribution. The isotherm for MX-13N rises steeply throughout the range of relative pressures 0.3 c P/P, c 0.95, indicating the presence of pores right across the size range from small micropores (diameter (2.0 nm) to large mesopores (diameters up to 50 nm). The isotherm for MX-12H has a similar upward curvature at low relative pressures, but levels off ai high vdues of PIP,. The pore distribution in this case is shifted towards smaller pores. For MX-29V the whole balance of the pore distribution has tipped towards microporosity and there are now very few mesopores.The subtle differences in the pore size distributions of these trioides can be described far more precisely by using the a=-method of isotherm analysis. I I I I 0 4 08 0.k 0.8 PiPo FIG. 1 .-Nitrogen physisorption isotherms for molybdmum oxides : (a) molybdenum trioxides prepared by decomposing molybdenum oxalate in hydrogen (A), in nitrogen (B), and in w c u m (C) ; 0 adsorption, x desorption ; (b) samples MX-3(a) and MX-3(b) ; (c) sample MX-7 ; (4 sample MX-10 ; (e) standard isotherms for iron molybdate (1) and non-porous molybdenum trioxide (2) ; (f) samples AM-1N and AM-2N ; (g) samples MX-1 l(a) (curve E), and MX-1 l(c) (curve F) ; (h) passf- vated and active reduced oxides : G, active sample ; €3, passivated sample ; J, annealed sample.2994 MOLY BDE N UM OXIDB CA'IALY S TS (&-METHOD The as-method 9-13 provides a simple graphical method of comparing the various features of the isotherm of an uncharacterised material with the isotherm obtained under the same experimental conditions for a non-porous standard reference material.A suitable reference material is one whose surface chemistry corresponds as closely as possible to the surface of the unknown sample. This ensures that the interaction between adsorbate molecules and the external surface of the unknown sample will be similar to the interaction between adsorbate molecules and the surface of the non-porous reference. Differences between the two isotherms can therefore be attributed to differences between the physical structure of the unknown sample and the non-porous reference material.From the standard isotherm the amount adsorbed (WJ at each value of relative pressure (PIP,) is converted into a " reduced " adsorption by dividing W, by Wo.4, where Wo.4 represents the amount adsorbed at a reference pressure P/Po = 0.4. The values of Ws/Wo., thus obtained are called as-values. The isotherm for the unknown sample (2) is analysed by plotting Wz against Q. Deviations of this a,-plot from linearity are interpreted in terms of micropore filling or capillary condensation depending on the nature of the deviations. The choice of a good standard isotherm is crucial to the as-method of isotherm analysis. As a check we have used two isotherms, one for an iron molybdate sample and one for a molybdenum trioxide sample which was prepared by decomposing ammonium dimolybdate. Both isotherms appear to have the correct characteristics [fig.l(e)]. The shapes of the isotherms and the absence of hysteresis are typical of non-porous materials. The samples have low surface areas (14.5 and 11.0 m2 g-l, respectively) which is also consistent with a non-porous structure. Furthermore, the surface of both of the reference materials should be fairly similar chemically to the surfaces of our molybdenum samples. Qs-CURVES FOR MOLYBDENUM TRIOXIDES The a,-curves for samples MX-l3N, MX-12H and MX-29V are shown in fig. 2(a) together with the as-curve for sample AM-1N. This latter sample, which has been referred to earlier, was prepared by the thermal decomposition of ammonium dimolybdate in nitrogen at 623 K.The linear as-plot for AM-1N confirms that the isotherm for iron molybdate is a satisfactory standard. Curve C shows a small upward deviation from linearity at PIP, = 0.15 which indicates micropore filing in pores whose diameters range from very small through into the mesopore region. This is followed by some capillary condensation in mesopores but the upper limit where the mesopores are all filled only extends to a diameter of 7 nm. The final almost flat portion of curve C corresponds to multilayer adsorption on a small external surface. Curve A is similar to curve C in exhibiting on upward deviation at PIPo = 0.15, but the upper limit on the mesopore diameter is slightly higher at about 10 nm. Multilayer condensation on a small external surface is again observed.Curve B exhibits an upward deviation at much higher values of relative pressure (PIP, = 0.35), so the proportion of micropores is very much smaller in this sample. The upper limit on the pore diameter in the mesopore range is now >20 nm. There- fore, in sample MX-13N there is a distribution of pore sizes ranging from micropores right through to very large mesopores. Some of the properties of these trioxides are summarised in table 1. A comparisonR. BURCH 2995 between SBET and Sp, and between Vo.9J and Vp, confirms that the degree of micro- porosity increases in the order MX-13N < MX-12H < MX-29V. The external surface area decreases in the reverse order, with the decrease from MX-13N to MX-12H being particularly noticeable.The external surface area of MX-13N (table 1) is 26 m2 g-l. From X-ray line broadening we have estimated an average particle size for this material of 40-45 nm which corresponds to a surface area of 30m2g-l for cubic particles. The agreement between these two surface areas is very good. From the nitrogen adsorption isotherms on the trioxide samples we can obtain valuable information on the way in which the preparative conditions affects the structure of the trioxide. In wet nitrogen an essentially mesoporous trioxide is US FIG. 2.--as-plots for molybdenum oxides : (a) A, sample MX-12H ; B, sample MX-13N ; C, sample MX-29V; D, sample AM-1N; (b) reduced oxides x, MX-10; 0, MX-7; 0, MX-3(a); 0, MX-3(b) ; (c) samples MX-ll(a), (curve E), and MX-1 l(c), (curve F) ; (d) active (curve G), passivated (curve H), and annealed (curve J) samples; (e) samples AM-1N (lower curve) and AM-2N (upper curve).TABLE PROPERTIES OF THE TRIOXIDES MX-l3N, MX-l2H, MX-29V, AM-1N surface areafmz g-1 pore volumefcm3 8-1 sample sBBTa SETb Sme VP* YO.95. MX-13N 99 26 100 0.12 0.12 MX-12H 119 9.5 90 0.09 0.11 MX-29V 41 3.6 20 0.02 0.03 AM-1N 11 11.6 - - a SBET is the surface area, calculated from the nitrogen isotherm using the B.E.T. method. S ~ T is the external surface area, calculated from the as-plots. C Sp is the surface area of all the pores except micropores, calculated from the nitrogen desorption isotherm. d Vp is the volume of the mesopores, calculated from the nitrogen desorption isotherm. e Vo.9s is the volume of all the pores, calculated from the nitrogen isotherm.2996 MOLYBDENUM OXIDE CATALYSTS obtained.In wet hydrogen a slightly more oxygen deficient trioxide is obtained which contains some mesopores but also a significant number of micropores. Under vacuum conditions a very microporous material with a small external surface area is obtained. We conclude from this that hydrogen reduction generates slit-shaped micropores in mesoporous molybdenum trioxide, and that water vapour causes a significant proportion of the micropores to be widened out to mesopore dimensions. DEVELOPMENT OF STRUCTURE DURING HYDROGEN REDUCTION REDUCTION I N THE PRESENCE OF WATER VAPOUR 1. OXALATE-DERIVED POROUS TRIOXIDES. Fig. l(b), (c) and (d) show the nitrogen adsorption isotherms, and fig. 2(b) shows the corresponding a,-plots, for reduced samples prepared from " nitrogen-decomposed " oxalate according to the conditions of temperature and time summarised in table 2.These samples were prepared by reduction in static hydrogen so that water vapour produced in the reaction would remain in the vicinity of the samples. TABLE 2.-sAMPLES OF REDUCED MATERIALS PREPARED IN THE PRESENCE OF WATER VAPOUR heat treatment time at temperature time at higher mol % sample 573 K/h raised to/K temperaturelh Mo metal MX-7 23 - 18.4 MX-3(a) 20 773 1 52.0 MX-10 8 673 1.5 52.6 - a MX-3(b) 798 5 90.0 aMX-3(b) was prepared by further reduction of MX-3(a). bA~~urning only Mo and Mooz are present in reduced oxides. Both the temperature and time of reduction affect the structure of the reduced oxides, although the temperature is more important.Thus MX-7 and MX-3(a) have almost identical isotherms and a,-plots, whereas MX-10 is less microporous and more mesoporous. We note also that the degree of reduction does not appear to be a critical factor in determining the physical structure of the reduced oxides, at least up to a nominal composition of 50 % Mo metal. The main differences between MX-7, MX-3(a) and MX-10 appear to originate in the different thermal treatments. The development of the pore structure during reduction appears to occur in the following way. The initial reduction creates slit-shaped micropores in the trioxide. Providing the temperature is kept low (573 K) these micropores are widened into a range of pore sizes by a surface reconstruction which is catalysed by water vapour.If the temperature is raised to 673 K the average pore size is increased and the effective external surface area is also increased. If the temperature is raised to 773 K some sintering occurs with a concomitant loss of external surface area and a decrease in the number of mesopores. When reduction is continued at temperatures above 773 K there is a considerable loss of mesoporosity but fresh reduction retains the micropore content of the sample. The overall effect, shown by sample MX-3(b), [fig. l(b) and 2(b)], is to reduce the surface area of the sample and to increase the relative proportion of micropores. The aS-plot for MX-3(b) shows very little additional pore filling above PIPo = 0.5, which indicates an upper limit on the pore diameter of about 5 nm.The reduced product has a structure, therefore, which depends on temperature and, at higher reduction levels, on the degree of reduction. Low temperature reduction produces a porous material containing mainly micropores but someR . BURCH 2997 mesopores. Reduction at intermediate temperatures causes the average pore diameter to increase. Prolonged reduction at high temperatures results in sintering with a loss of external area and some blocking-off of the mesopores, but, because new micropores are continuously being created by further reduction, the micropore content remains high. Table 3 summarises some of the relevant properties of these reduced oxides. Comparisons between SBET and Sp and between V,,.95 and Vp confirm the general pattern of the structural changes described above. TABLE 3.-PROPERTIES OF REDUCED OXIDES sample SBET a SEXT a S p a VP" v0.95 " MX-7 158 11.9 110 0.095 0.144 MX-3(a) 144 11.9 110 0.096 0.140 MX-3 (b) 78 7.1 31 0.026 0.058 MX-10 162 26 106 0.107 0.159 MX-ll(a) 174 9.5 68 0.054 0.115 MX-ll(C) 111 8.3 62 0.040 0.070 a See caption to table 1.2. NOWPOROUS TRIOXIDES. The trioxide samples described above had porous structures and although it has been implied that hydrogen reduction initially creates microporosity this does not show up very clearly because of the complexity of the structure of the starting material. The effect of hydrogen can be demonstrated more clearly by considering the reduction of a non-porous trioxide (AM-1N). On reduction at 733 K a sample (AM-2N) containing nominally 6 % molybdenum metal was obtained. The isotherms in fig.l(f) and the a,-plots in fig. 2(e) show very clearly the creation of micropores but the complete absence of mesopores. It does not appear from sample AM-2N that micropores broaden into mesopores during the normal reduction process. Mesopores may be created by a secondary structural transformation involving the coalescence of adjacent micropores. The presence of mesopores in the high surface area reduced oxides, but not in the low surface area oxides, presumably arises because in small particles the average spacing between the micropores is small and so there is a good chance of finding micropores close enough together to coalesce. A simple 3-dimensional model made to represent as closely as possible the physical properties of sample MX-10 shows that between 30 and 40 % of the micropores would be close enough to another micropore to coalesce.REDUCTION IN ABSENCE OF WATER VAPOUR A set of experiments was performed in which the sample pan was surrounded by previously outgassed alumina pellets to maintain a dry environment around the sample. The nitrogen isotherms obtained for the samples reduced under dry conditions, at the temperatures shown in table 4, are plotted in fig. l(g), with the corresponding a,-plots shown in fig. 2(c). Some of the physical characteristics of these samples are summarised in table 3. There are important differences between, for example, MX-ll(a) and MX-10 after approximately the same heat treatment.MX-ll(a) is very much more micro- porous, it has fewer mesopores, and a much smaller external surface area. Further reduction of MX-ll(a) to MX-ll(c) causes a decrease in the B.E.T. surface area, a decrease in microporosity, but very little change in the mesoporosity.2998 MOLYBDENUM OXIDE CATALYSTS The differences between samples MX-ll(a) and MX-10 suggest that in the presence of water vapour some micropores are broadened into mesopores, possibly by coalescence as described earlier. A similar effect of water vapour was noted earlier when we described the differences between the trioxides decomposed under vacuum and in hydrogen. The dry oxide which was prepared under vacuum was much more microporous. TABLE 4.-sAMPLES OF REDUCED MATERIALS PREPARED UNDER DRY CONDITIONS heat treatment sample 573 K/h raised to /I( temperaturelh Mo metal time at temperature time at higher mol % MX-ll(a) 4 623 1.1 24.4 MX-ll(c) a 673 22.0 80.0 a MX-1 l(c) was prepared by further reduction of MX-1 l(a).PASSIVATION OF REDUCED OXIDES The reduced oxides described above are extremely pyrophoric. They can be passivated by treatment with very dilute oxygen+argon mixtures. It is of interest to examine the effect of passivation on the physical structure of the samples. Fig. l(h) shows the nitrogen isotherms (curve G is a desorption isotherm) before (curve G) and after (curve H) a passivation treatment for a sample having a nominal composition of 59 % Mol41 % Moo2.* The corresponding a,-plots are shown in fig. 2(d). The similarity between curve G [fig.l(h)] and the earlier isotherms for samples MX-11 and AM-2N shows that the pyrophoric material has a structure which is dominated by microporosity. Curve G [fig. l(h)] is almost exactly 50 % larger than curve H across the whole range of relative pressures. This shows that the passivation treatment has affected all of the surface features of the sample equally. Approximately one third of (i) the external surface area, (ii) the number of mesopores, and (iii) the number of micropores, has each been eliminated by the passivation reaction. The B.E.T. surface area is also reduced by a similar amount, from 124 to 84 m2 8-l. EFFECT OF ANNEALING After heating the passivated sample in nitrogen at 823 K for 19 h the adsorption isotherm shown as curve J in fig. l(h) was obtained.The corresponding a,-plot, in fig. 2(d) shows that there have been significant changes in the structure as a result of this annealing treatment. The a,-plot has a slightly shorter initial linear portion and at a, = 0.8 the curve deviates upwards from linearity. This indicates that the balance between microporosity and mesoporosity has shifted in favour of meso- porosity. The annealing treatment has caused the elimination of many of the micropores, and possibly a small proportion of the micropores have been widened to mesopore dimensions. However, the external surface area has not been much affected by the annealing treatment, as evidenced by the close similarity in the a,-plots for samples H and J over the range 1.6 < a, < 2.5. SUMMARY The evolution of the structure of molybdenum oxides obtained from the thermal decomposition and reduction of molybdenum oxalate can be summarised as follows.* I am grateful to the Climax Molybdenum Co. Ltd for supplying a sample of the passivated material, and for the desorption isotherm on the active material.R. BURCH 2999 In the absence of water vapour decomposition of molybdenum oxalate at 573 K produces an essentially microporous material, but in the presence of water vapour both micropores and mesopores are obtained. Reduction of the trioxide at 573 K creates micropores initially, but further reduction, or again possibly the presence of water vapour, causes some pores to be widened into mesopores. In contrast, the reduction of a low surface area trioxide produces only micropores.Passivation of the active surface of a reduced product results in almost identical decreases in the number of micropores, the number of mesopores, and the B.E.T. surface area. Finally, annealing reduced samples at 823 K results mainly in a loss of microporosity and some increase in mesoporosity. I thank the S.R.C. for an equipment grant. I am grateful to the Climax Molybdenum Co. Ltd who first suggested this work, for their continuing support. We are grateful to Prof. K. S. W. Sing for many stimulating discussions concerning this work. J. R. Anderson, Structure of Metallic CataZysts (Academic Press, London and New York, 1975). J. F. Taylor and J. K. A. Clarke, 2. phys. Chem., 1976, 103,216. ’ G. A. Tsigdinos, C. J. Hallada and R. W. McConnell, U.S. Patent 3,912,660 to Amax Inc. October 14, 1975 ; German Patent 2,451,778 (to Arnax Inc.), May 28, 1976. G. A. Tsigdinos and W. W. Swanson, Ind. and Eng. Chem., to be published. ’ R. Burch and P. C. H. Mitchell, J. Less-Common Metals, 1977, 54, 363. R. Burch, J.C.S. Faraday I, 1978,742982. ’ J. H. de Boer, The Structure and Properties of Porous Materials (Butterworth, London, 1958), p. 68. S. Brunauer, L. S. Deming, W. S. Deming and E. Teller, J. Amer. Chem. SOC., 1940, 62,1723. K. S. W. Sing, in Proc. Int. Symp. Surface Area Determination (Butterworth, London, 1970), p. 25. lo F. S. Baker, J. D. Carruihers, R. E. Day, K. S. W. Sing and L. J. Stryker, Disc. Faruday SOC., 1971, 52, 173. M. A. Alaria Franco and K. S. W. Sing, AnaZes de Quim., 1975,71,296. K. S . W. Sing, Ber. Bunsenges.phys. Chem., 1975, 79, 724. (Academic Press, London and New York, 1976). l 3 K. S. W. Sing, in Characterisation ofPowder Surfaces, ed. G. D. Parfitt and K. S. W. Sing (PAPER 8/521) 1-95

ISSN:0300-9599    

DOI:10.1039/F19787402991

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Spectroscopic Investigation of the Structure of a Novel Zerovalent Cobalt Nitrosyl in Zeolite Matrices BY H ~ B N E PRALIAUD, GIS~LE F. COUDURIER AND YOUNES BEN TAARIT* C.N.R.S., Institut de Recherches sur la Catalyse, 79, Boulevard du 11 Novembre 1918, 69626 Villeurbanne Cedex, France Received 12th April, 1978 Cobalt (II) hexahydrate introduced in synthetic zeolites by conventional ion-exchange was shown to give up readily five of its water molecules upon evacuation at room temperature. Subsequently the cobalt ion was coordinated by three lattice oxide ions from the zeolite lattice and by a single water molecule so as to acquire a slightly distorted tetrahedral symmetry. This tetrahedral CoII species readily reacted at room temperature with nitric oxide to yield a zerovalent paramagnetic cobalt dinitrosyl within the zeolite framework.The complex was identified on the basis of its i.r., U.V. and e.p.r. spectra as a C2h distorted octahedral complex with a dX2-,,2 ground state with a co- ordination sphere made up of three lattice oxide ions [possibly O(4) ions], a water molecule and two NO ligands which transferred their odd electrons to the cobalt d orbitals. In recent years, transition metal exchanged zeolites attracted much attention and were studied using a myriad of techniques. The prominent feature of these solids resides in the ease with which they form within their framework a variety of complexes upon interaction of transition metal zeolites with various organic and/or inorganic reagents. In fact, even when initially located in remote sites, transition metal ions were conclusively shown to migrate to accessible sites in order to co- ordinate adsorbed molecuIes.1-7 These phenomena have been observed by both direct and indirect methods. E.s.r. spectroscopy appeared to be one of the most efficient and informative means to investigate such cation behaviour within the zeolite lattice, though powder patterns could provide valuable information on the symmetry and electronic structure of the complexed ion and generally on the location of the complex. As to their structure, most transition metal complexes were shown to behave in a similar fashion to their solution analogues, although specific interaction with the zeolite framework could also modify this behaviour. In particular, the zeolite lattice appeared to be a stabilising matrix for a variety of cations in unusual coordination or oxidation states.2* 4* 8*12 We report in this study the formation of a novel zerovalent cobalt nitrosyl complex generated within the zeolite framework. EXPERIMENTAL AND MATERIALS The starting zeolite material was a synthetic faujasite type zeolite (both X and Y) with respective unit cell formulae (NaA102)86(Si02)1 6xH20 and (NaA102)5 6(Si02)1 6yH20.The cobalt zeolites were obtained by conventional ion-exchange of the parent materials with dilute cobalt (11) sulphate solutions. Several exchange levels (2.9, 6, 8.5, cobalt %) were reached upon repetitive procedures. The samples were then thoroughly washed with deionised water. E.s.r. spectra were recorded on a Varian E9 spectrometer equipped with a dual cavity operating in the X band mode and DPPH was used as a standard for g value determination.Q band spectra were obtained using a variable temperature accessory, Mn2+ impurities in the zeolite were used as a convenient internal reference. 3000H. PRALIAUD, G. F. COUDURIER AND Y. BEN TAARIT 3001 Diffuse reflectance spectra were recorded, at room or liquid nitrogen temperature in the range 220 to 2500 nm, on an Optica Milano CF4N1 spectrometer using a differential method, the reference being MgO or the starting zeolite. Infrared measurements were performed on a Perkin Elmer spectrometer model 125. The zeolite sample was compressed into a thin wafer at a pressure of 1 ton cm-2 and inserted into a sample holder. A conventional i.r.cell, as described elsewhere,13 was used. Nitric oxide was purified by the freeze-pump-thaw technique and finally distilled from a liquid nitrogen+pentane slurry (138 K) to remove NOz impurities. RESULTS Fully hydrated cobalt zeolites were light pink in colour and exhibited only a broad e.s.r. signal at g x 3.3 at whatever temperature, The U.V. spectrum showed two absorption maxima at 525 and 1200 nm [fig. l(a)J. Upon evacuation at room temperature (1-2 h) to a final pressure of Torr, the cobalt zeolite turned purple and the e.s.r. spectrum was not significantly affected. By contrast the U.V. spectrum was modified drastically : the absorption maximum at 525 nm decreased in intensity and an absorption occurred in two ranges : 1200-1600 nm on the one hand and 520-640 nm with maxima at 540, 575 and 620 nm on the other hand, as shown in fig.l(b). The infrared spectrum still showed the 6H,O strong band at 1640 cm-l characteristic of molecular water and at 1605 cm-l due to H20 specifically co- ordinated to Co2+ ions.14 I I I 1 I 1 1 . 200 400 600 800 A/nm FIG. l.-(a) U.V. diffuse reflectance spectrum of COII(H~O)~ in faujasite type zeolites. (b) U.v spectrum of CoII zeolite degassed at room temperature for 2 h. (c) U.V. spectrum of the nitric oxide adduct. Adsorption of nitric oxide resulted in the growth of two strong vNO bands at 1890 and 1806 cm-l (fig. 2). Both bands grew at the same rate upon further contact with NO. Their growth did not affect the band at 1605 cm-l. The sample now turned grey at room temperature and greenish-grey at 77 K.Under these con- ditions a new e.s.r. spectrum could be recorded at 77 K (fig. 3). On progressively warming the sample, the signal intensity decreased and the features broadened. At around 153 K, the signal was broadened beyond detection. The X band pattern shown in fig. 3 is characteristic of an intense hyperfine interaction of the electron with a spin 712 cobalt nucleus. The g and hyperfine principal components which3002 I . R . , U.V., E.P.R. OF COO NITROSYL IN ZEOLITES could be extracted from the X band spectrum are listed in table 1. As the con- siderable line overlap set a lower limit for the accuracy of the g and A values, Q band spectra were sought and they confirmed the departure of the g tensor from axial symmetry (fig.4). However, in spite of the fairly low temperature achieved in the variable temperature device, the Q band spectrum was obscured by appreciable line broadening thus precluding any improvement in the accuracy of the gx and gr components. Again the g and A values are also listed in table 1. They are in fair agreement with those deduced from the X band spectrum. In the U.V. spectrum, absorptions at 240, 360, 460-70, 750 and 1200 nm could be seen [fig. l(c)] at all temperatures. - * 1806 1 I 9 Y FIG. 3.-X band spectrum of the nitric oxide adduct recorded at 77 K.€I. PRALIAUD, G. F. COUDURIER AND Y. BEN TAARIT 3003 t * I I l I I J gu FIG. 4.-Q band spectrum of cobalt nitrosyl recorded at 108 K. TABLE 1 . C . s . ~ . PARAMETERS OF COBALT (0) ,NITROSYL X band 2.13 (0.05) 2.11 (0.05) 2.36 (0.05) +118 (5) +118 (5) -75 (5) mode gx 0, Bz AxIG-1 AxlG-1 A.IG-1 Q band 2.14 (0.05) 2.10 (0.05) 2.29 (0.05) +120 (5) +12O (5) -75 (10) DISCUSSION The magnitude of theihyperfine coupling to the cobalt nucleus together with the strong g shift from the free electron g value, clearly indicate that such species as NO, NO2, NO;-, etc., should be ruled out and that the odd electron is most certainly a cobalt d electron. However, the low spin d7C0" complexes are usually octahedral or square- pyramidal with a d,Z ground state.6* I 5 - l 7 The gz value for such complexes was typically very close to the free spin g value while the other two g components were usually larger than both gz and ge.This is obviously not what is observed for this particular nitrosyl complex.Furthermore, the formation of a d7 nitrosyl compound would require the presence of two NO molecules and a formal dismutation of the NO ligand into NO- and NO+ species to ensure the d7 configuration : this is definitely inconsistent with the experimental infrared spectrum. A single NO ligand could also secure a d7 cobalt configuration, but the odd electron of the NO molecule would still be unpaired and we would have to deal with a triplet state situation. COORDINATION OF COBALT SPECIES The hydrated cobalt zeolite showed a U.V. spectrum characteristic of octahedral COII(H~O)~ in the large cavities of the zeolite, the 525 and 1200 nm maxima being respectively due to the 4T,,(F) + 4T1,(P) and 4T2,(F) transitions. Following partial dehydration and prior to NO adsorption, the U.V.spectrum showed the existence of cobalt (11) species in tetrahedral symmetry, as could be deduced from the appearance and development of the two ranges of absorption between 520-640 and 1200-1600 nm at the expense of the absorption due to [Co(H20),l2+. The existence of the two absorption ranges is due to transitions from the non-degenerate ground state 4A,(P) to the triply degenerate excited 4T1(P) and 4T1(F) states. The splitting of the bands corresponding to the excitation to the 4T1(P) level with maxima at 540, 575 and 620 nm could be due to dynamic Jahn-Teller effect or to a low symmetry3004 I.R., U . V . , E . P . R . OF COO NITROSYL ILN ZEOLITES perturbation which lifts the excited state degeneracy. The tetrahedral Co" cation should therefore be located in an SII position at the hexagonal windows within the supercage.Upon adsorption of nitric oxide the decrease in the 520-640 nm absorption and the simultaneous occurrence of bands at 360, 460-70 and 750-60 nm are indicative of changes in the coordination and electronic configuration of the cobalt complex. The infrared spectrum of the nitric oxide adduct is consistent with a dinitrosyl species. In fact the two vNO bands are similar to those observed in the case of well defined nitrosyl compounds where the bands at 1890 and 1806 cm-l are due to the symmetric and antisymmetric NO vibration modes. The presence of a residual water molecule in the coordination sphere of the cobalt (11) ions as witnessed by the presence of the 1605 cm-1 band in addition to three oxide ions prior to NO adsorption is consistent with the tetrahedral symmetry inferred from the U.V.spectrum. Addition of two NO molecules as deduced from the i.r. spectrum leads to an octahedral type symmetry. It is highly probable that both the zeolite lattice and the specific ON-Co NO bond angle contribute to impose drastic alteration of the octahedron. The occurrence of the vNO vibrations at 1890 and 1806 cm-l indicative of a dinitrosyl species and the following analysis of the U.V. spectrum strongly favour a dgCoO dinitrosyl structure arising from the transfer of both odd electrons of the two NO ligands to the initial Co" tetrahedrally coordinated ion thus forming a Coo octahedral dinitrosyl. Indeed the 460-70 and 750-60 nm transitions characteristic of the NO adduct are in agreement with a distorted algCo0 complex : the 460-70 nm absorption is assigned to an internal Coo transition from the ground 2D(3d9) state to the excited 2F(3d84s1) configuration.l* On the other hand the 750-60 nm transition is more difficult to assign but also more informative.Copper" d9 distorted Oh complexes are known to exhibit d-d transitions in the 620-930 nm range.19 In particular [CU(H~O)~]~+ in zeolite absorbs at 800 nm.20 Substitution of water by N-ligands such as NH, shifts the absorption maximum to around 620 to 570 nm.21 Similarly dgN? complexes have been identified in zeolite matrices following either Hz reduction of Ni" ions 22 or addition of nitric oxide to Ni" exchanged In the former instance Nil species gave rise to an absorp- tion maximum around 740 nm ascribed to an internal transition of the Nif ion.In the latter case octahedrally distorted Ni' nitrosyl identified on the g basis of its e.p.r. spectrum (911 = 2.365 and gl = 2.193) and its i.r. spectrum (vNO at 1892 cm-l) gave rise to transitions at 360, 650,740 and 1650 nm.23 As previously 22 the 740 nm band is characteristic of the internal (d9-d84s1) transition of Nil while the 650 nm absorption is due to d-d transition of 3d9Ni1 distorted octahedral complex. As the charge on the central atom of the octahedrally distorted complex having N ligands de- creased from 2 to 1 Cu" and Ni', the energy of the corresponding d-d transition was also expected to decrease, as was actually observed, from 570 to 650 nm.Therefore, the transition energy of the isoelectronic Coo octahedral nitrosyl complex, con- sidering the zero charge on the central atom, should be even lower ; hence, the 750 nm maximum seems to be a likely candidate for such a d-d transition. In fact Coo complexes have already attracted considerable interest in recent years. In particular, paramagnetic cobalt tetracarbonyl has been reported to form upon sublimation of CO,(CO)~ onto a cold finger held at 77 K in the microwave cavity of an e.p.r. spectr~meter.~~ This radical species was also studied by Symons et aZ.25 and its reactivity towards O2 demonstrated. More recently Ozin and coworkers26 reported an i.r., e.p.r. and U.V. study of Co(CO), and CoCO),. In a more recentH . PRALIAUD, G. F .COUDURIER AND Y. BEN TAARIT 3005 study, Symons and coworkers 2 7 succeeded in isolating a substituted cobalt (0) carbonyl [Co(CO), pb PhJ- which is thought to retain the structure of the parent carbonyl, i.e., trigonal bipyramidal. Coo paramagnetic complexes were also pre- pared by chemical or electrochemical reduction of Co' organometalh compounds and again a distorted tetrahedral structure was favoured.28 Lastly a mixed nitrosyl dicobalt carbonyl [Cp,(Co),NOCO] and the Cp2(Co),(CO), anion were shown to be paramagnetic and relevant g parameters to be close to those reported for Co(CO), etc. Simultaneous X-ray structural determinations showed both com- pounds to have a distroted tetrahedral stru~ture.~ The overriding conclusion emanating from all these studies was the general agreement on the magnitude and ordering of the g values as can be seen from table 2: g1 > gz N ge. In all cases this was interpreted as due to a dgCoo complex in a C,, distorted tetrahedral sym- metry consistent with a d,t ground state.TABLE 2 . C . S . R . PARAMETERS OF DISTORTED TJ3TRAHEDRAL coo COMPLEXES compound reference g, gY gz AJG-1 Ay/G-l Az/G-I Co(CO), in C O ~ ( C O ) ~ matrix 24 2.134 2.134 2.02 56 56 67 Co(CO), in CO matrix 26 2.128 2.128 2.007 55 55 58 Co(CO), Pb Ph; in Co(CO), Pb Ph3 27 2.020 2.017 1.996 -48 -48 f 3 7 As one can see from table 1 our g values differ from those reported in table 2 in the ordering, with an additional splitting of the perpendicular component into two distinct values both of them larger than ge, and gz far larger than gx and gr.On the basis of the relative magnitude of the g tensor components st distorted tetrahedral symmetry with a dz2 ground state should be definitely ruled out, which is reasonable in view of the ligand count. By contrast, this particular ordering of the g values is similar to that observed for d9Cu2+ complexes in distorted octahedral symmetry and to that reported for Ni' nitrosyl complexes formed within the zeolite cavities 4* and also to NO2 adduct also forming Ni' specie^.^ All these complexes were shown to have distorted octahedral symmetry with a dxz-y2 ground state. Additional evidence for a CjX2-,,2 ground state could be provided upon resolving the hyperfine tensor. Within experimental error, the hyperfine tensor appeared as nearly axial, and could be resolved into its isotropic and dipolar parts in four possible ways depending on the choice of sign for Al and All, i.e., All = +75 and AL = 11 8 G.Taking the signs as all positive or all negative results in too small an aniso- tropy to account for the 3d character of the electron. Taking A , , = +75 and Al = - 118G yields Aiso = - 53.7 and 2B = + 128.7. This is consistent with a dz2 ground state, which had to be ruled out on the basis of the ordering and magnitude of theg values with respect tog,. The alternative choice, which we favour, gives Aiso= 59.7 and 2B = - 128.7. This result is consistent with a high population in thedx2-,2 orbital and is also consistent with the observed g tensor and the U.V. spectrum. The residence time in this orbital, considering that for unit occupancy 2B0 = - 152.3 G,27 is quite reasonable = 84.5 %.On the other hand the usual value for Aiso for cobalt com- plexes is around - 100 G. To achieve a positive value would imply an appreciable admixture of the dX2-,,2 orbital with the outer 4s shell. This is quite reasonable in the case of D2,, or CZh distortion of the octahedral symmetry. Such a distortion3006 I . R . , U . V . , E . P . R . OF COO NITROSYL I N ZEOLITES within the zeolite framework could scarcely be avoided. Thus the approximate 4s population that could be deduced from this positive value is within 11.5 %. Hence 96 % of the odd electron is exclusively delocalized on the metal atom. This is consistent with the absence of any detectable nitrogen superhyperfine coupling.The departure of the g tensor from axiality is also consistent with a C,, or DZh deviation from octahedral symmetry. The rigid zeolite matrix can of course undergo an appreciable Si A1 alteration in order to accomodate various ions and complexes ; however, it is doubtful whether these alterations could be important enough to secure perfect symmetries in the case where oxide ions of the lattice are part of the coordination shell. Indeed X-ray studies 30 showed that 0-Co-0 angles in Co-Y zeolites are far larger than 90" : they are 97" to 117" depending on the nature of other attached ligands. Furthermore the vNO vibration bands relative intensities give a useful estimate of N-Co-N bond angle using the formula :31 4 = 2 arcotan JR(sym)/R(asym) where R is the intensity of the particular vibration mode, the N-Co-N bond angle deduced from the experimental R(sym)/R(asym) ratio of 0.37 is about 117" and clearly indicates a large departure from D2,, symmetry, which is expected considering the structural nature of the host.In conclusion the isolated complex is best described as a d9 Coo dinitrosyl complex with a ~,z-,J ground state in a C2h distorted octahedral symmetry. This is the first example of an octahedral zerovalent paramagnetic nitrosyl complex. The occurrence of such a complex is obviously favoured by the ability of the zeolite matrix to accommodate and stabilize such unusual valence and coordination states. Note that the complex was simply produced by NO addition and did not require any extreme treatment to be formed, thus emphasizing the suitability of the zeolite matrix for this particular type of chemical study.0 \ / \ / / \ / \ The authors gratefully acknowledge discussions with Prof. Che of the Pierre and Marie Curie University and with Dr. B. Imelik and Dr. C. Naccache. Thanks are due to G. Wicker and J. F. Dutel for technical assistance. C. Naccache and Y. Ben Taarit, Chem. Phys. Letters, 1971, 11, 11. C. Naccache and Y . Ben Taarit, J. Catalysis, 1971, 22, 171. P. Gallezot, Y . Ben Taarit and B. Imelik, J. Catalysis, 1972, 26, 295. C. Naccache and Y. Ben Taarit, J.C.S. Faraday I, 1973,69, 1475. P. Gallezot, Y . Ben Taarit and B. Imelik, J. Phys. Chem., 1973,77, 2364. E. F. Vansant and J. H. Lunsford, J.C.S. Chem. Comm., 1972, 830 and Trans. Faraday Soc., 1973,69,1028. ' K.A. Windhorst and J. H. Lunsford, J. Amer. Chem. SOC., 1975,97, 1407. * C. Naccache, Y. Ben Taarit and M. Boudart, A.C.S. Symposium series, Molecular Sieves 11, 1977, 40,156. P. H. Kasai and R. J. Bishop Jr., J. Amer. Chem. SOC., 1972,94,5960. lo C. Naccache, J. F. Dutel and M. Che, J. Catalysis, 1973, 29, 179. l1 M. Che, J. F. Dutel, P. Gallezot and M. Primet, J. Phys. Chem., 1976, 80, 2371. l2 E. Garbowski and J. C. Vedrine, Chem. Phys. Letters, 1977,48,550. l3 M. V. Mathieu and P. Pichat, in La Catalyse auLaboratoire et dansl'lndustrie (Masson, Paris, l4 G. F. Coudurier, to be published. l5 (a) J. P. Maher, J. Chem. SOC. A, 1968, 2918 ; (b) M. E. Kimball, D. W. Pratt and W. C. l6 M. C. R. Symons and J. G. Wilkinson, J. Chem. Soc. A, 1971,2069. 1967), p. 319. Kaska, Inorg. Chem., 1968, 7, 2006.H. PRALIAUD, G. F. COUDURIER AND Y. BEN TAARIT 3007 l7 F. D. Tsay, H. B. Gray and J. Danon, J. Chem. Phys., 1971,53,3760. l9 B. J. Hathaway and D. E. Billig, Coord. Chem. Reu., 1970,5,143. 2o J. D. Mikheikin, V. A. Shvets and V. B. Kazanskii, Kinetika i Kataliz, 1970, 11, 747. 'l J. H. Anderson, J. Catalysis, 1973, 28,76. 22 E. Garbowski, M. V. Mathieu and M. Primet, Chem. Phys. Letters, 1977, 49,247. 23 H. Praliaud, to be published. 24 H. J. Keller and H. Wawersik, 2. Naturforsch, 1965, 20b, 938. 2 5 S. A. Fieldhouse, B. W. Fullam, G. W. Neilson and M. C. R. Symons, J.C.S. Dalton, 1974,567. 26 L. A. Hanlan, H. Huber, E. P. Kiindig, B. R. McGarvey and G. A. Ozin, J. Amer. Chem. SOC., 27 0. P. Anderson, S. A. Fieldhouse, Ch. E. Forbes and M. C. R. Symons, J. Orgammetal. 28 H. Van Willigen, W. E. Geiger Jr. and M. D. Rausch, Inorg. Chem., 1977, 16, 581. '' I. Bernal, J. D. Korp, G. H. Reisner and W. A. Herrmann, J. Organometal. Chem., 1977,139, 30 P. Gallezot and B. Imelik, J. Chim. phys., 1974, 155. 31 W. Beck, A. Melnikoff and R. Stahl, Chem. Ber., 1966,99, 3721. C. E. Moore, Atomic Energy Leuels mat. Bur. Stand. Circ. 467, Washington D.C., 1952), vol. 2. 1975,97, 7054. Chem., 1976,110,247. 321. (PAPER 8/695)

ISSN:0300-9599    

DOI:10.1039/F19787403000

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Rayleigh Scattering Depolarization Ratio and Molecular Polarizability Anisotropy for Gases BY MARTIN P. BOGAARD,t A. DAVID BUCKiNGHAM,* RAYMOND K. PIERENS$ AND ALLAN H. WHITE$ University Chemical Laboratory, Lensfield Road, Cambridge CB2 1 EW Received 14th June, 1978 The depolarization ratio for Rayleigh scattering by 39 molecular species are reported for A = 488.0, 514.5 and 632.8 nm. The literature has been critically surveyed for refractivities which are used in conjunction with the observed depolarization ratios to yield the anisotropy of the molecular polariz- ability of the molecules. Results are given for C02, OCS, CS2, C2H6, C3Hs, isobutane, cyclo- propane, C2H4, allene, C2H2, methylacetylene, dimethylacetylene, spiropentane, cyclohexane, CHC13, CF2CI2, CH3Br, CH31, t-butylchloride, t-butylbromide, ethylene oxide, (CH3)2O, (CH3)2S, H2S, S02, Nz, N20 and CO.C6Ht5, C~HSF, 1,3,5-CsH3F3, C6F6, CH,CN, (CH,)ZCO, CH3F, CHZF2, CHF3, CH3C1, CHZC12, The anisotropy of the electric dipole polarizability is a molecular property which, in addition to its intrinsic interest,l is needed in the evaluation of molecular quadru- pole moments,2 magnetizability anisotropies 3* and molecular hyperpolariz- abilities 1* from measurements of the birefringence induced in gases by electric or magnetic fields. It is also relevant to the evaluation of long-range intermolecular forces. The frequency dependence of the polarizability anisotropy can be used6* ' to deduce orientationally dependent intermolecular forces. The molecular polarizability anisotropy can be determined from measurements of the depolarization ratio of Rayleigh scattered light.Although the advent of laser light sources has made such measurements at well defined wavelengths a relatively easy procedure, the literature contains only a few results. *-l This paper reports measurements, generally for three wavelengths, of the de- polarization ratio of Rayleigh light scattered by a number of gases. The measure- ments have been made over a number of years on molecules which have been of interest here. The results are combined with refractive-index data to yield the anisotropy in the molecular polarizability. EXPERIMENTAL APPARATUS AND PROCEDURE The optical system is very similar to that described by Bridge and Buckingham except that the scattering cell is outside the laser cavity.For measurements at 632.8 nm a 50 mW He+Ne laser was used ; all other measurements were made with an argon ion laser. Both lasers have a beam divergence of < 1 mrad and this introduces negligible errors in the determination of the depolarization ratio. A Glan-Thompson prism sets the plane of j- Present address : $ Present address : S Present address : Australia. 6009, Australia. School of Chemistry, University of New South Wales, Sydney, N.S.W. 2033, School of Chemistry, University of Sydney, N.S.W. 2006, Australia. Department of Chemistry, University of Western Australia, Nedlands, W.A. 3008BOGAARD, BUCKINGHAM, PIERENS AND WHITE 3009 polarization of the approximately vertically polarized laser beam at 90+0.02" to the scattering plane, The silica scattering cell has the entrance and exit windows for the laser beam set at the Brewster angle.It has a blackened horn opposite the viewing window. The cell is connected to a vacuum line through a Millipore filter of 0.10p.m pore size; provision is made for measuring the gas pressure in the cell. After traversing the scattering cell the laser beam falls onto either a monitoring photocell or a light trap. The solid angle over which scattered light is collected is defined by two circular stops; one immediately outside the viewing window of the cell and one in front of the photo- multiplier. The aperture of these stops is chosen to minimise " geometrical errors " ;' the maximum divergence of light accepted by the detector is < 3". The degree of polarization of the scattered light is determined with a second Glan-Thompson prism mounted in a graduated circle fitted with a micrometer-controlled tangent drive ; this enables the orienta- tion of the prism to be set with an accuracy better than 10'.The entrance face of the prism is set accurately perpendicular to the scattered light direction. Both this prism and the prism in the laser beam have an extinction ratio better than An EM1 9558B photo- multiplier is used to measure the scattered light intensity ; the dependence of its sensitivity on polarization direction was checked and found to be negligible. Two methods of measurement have been used: an analogue method and a photon- counting method. No systematic variation in the depolarization ratio was found for those substances investigated with both methods.For the analogue method the laser beam incident on the cell is intensity-modulated with a chopper. A lock-in amplifier with a differential input is used as a null detector; it detects the balance between the signal derived from the photomultiplier and that derived from a photocell which monitors the laser beam after it has passed through the cell. The latter signal is accurately attenuated by a factor F with a voltage divider until balance is achieved ; F-l is proportional to the scattering cross-section of the gas in the scattering cell. Measurements are made for the scattered light polarized parallel and perpendicular to the scattering plane. These measurements are then corrected by subtracting the corresponding values obtained with an evacuated cell and their ratio then yields the depolarization ratio.The voltage divider is constructed from two decade boxes and has sufficient resolution and accuracy to yield depolarization ratios accurate to one part in lo5. The photon-counting method uses an SSR 1120 amplifier/discriminator and SSR 11 10 counter. The laser intensity is suffciently stable over a period of minutes to make continuous monitoring of the laser intensity unnecessary. After traversing the cell the laser beam is caught in a light trap. Depohrization ratios are obtained as the ratio of the counts measured, in a chosen time interval, for the horizontal and vertical components of the light scattered by the gas-filled cell less the relevant counts measured for the evacuated cell.Care is taken to keep the count rate below 5 x lo5 s-l to avoid " pile-up " errors due to the dead-time (10 ns) of the discriminator. The compounds used were commercial samples. Liquids were degassed by several freeze-pump-thaw cycles and by trap-tetrap distillation. The cell was filled by passing the vapour above the degassed liquid through the 0.1 pm filter ; care was taken that the vapour pressure in the cell did not exceed one-half the room temperature saturation vapour pressure. Gases were taken directly from commercial cylinders and passed through the filter. Clean- ing procedures for the cell have been described previously.' The effect of impurities on the observed depolarization ratio, P&s, can be appreciable. If a and b refer to the species and impurity respectively where R = dbat/d,a: and p, a and d are respectively the depolarization ratio, mean polar- izability and number density.The species for which Ap is found to be significant will now be listed; the important impurities, with the upper limit of concentration (mole fraction) as stated by the supplier, are shown in parenthesis : ethane (ethylene 3 %), methylacetylene (dimethyl ether 1.2 %, allene 0.5 %), fluoromethane (dimethyl ether < 1 % by g.l.c.), rifluoromethane (nitrogen 1 %), dimethyl ether (carbon dioxide 0.25 %, methanol 0.5 %),molecule carbon dioxide carbonyl sulphide carbon disulphide ethane propane isobutane cyclopropane ethylene allene acetylene methylacetylene dimethylacetylene spiropentane cyclohexane benzene fluorobenzene 1,3,5-trifluorobenzene hexafluorobenzene acetonitrile acetone TABLE 1 .-DEPOLARIZATION RATIOS, POLARIZABILJTXES AND POLARIZ ABILITY ANISOTROPIES 488.0 4.12 k0.02 4.00 f0.02 7.69 f0.03 0.190 kO.003 0.214 k0.003 0.153 50.003 0.1675 2 0.002 4.225 kO.03 1.266 20.005 1.897 50.005 2.29 kO.01 2.82 kO.01 0.322 kO.OO1 0.165 k0.003 1.98 kO.01 2.11 k0.02 2.27 kO.05 2.46 kO.02 loop0 at Alnm : lO40a/C2 m2 J-* at I/nm : a 514.5 4.085 kO.02 3.95 k0.02 7.56 k0.03 0.188 k0.004 0.208 kO.006 0.1495 5 0.002 0.162 kO.002 4.17 k0.03 1.889 k0.005 2.27 kO.01 2.79 k0.01 0.315 rt0.002 0.163 k0.003 1.96 kO.01 2.09 kO.02 2.24 k0.05 2.44 kO.02 I .2475 2 0.005 632.8 488.0 4.05 k0.02 2.965 3.88 20.05 5.86 7.18 k0.07 10.06 0.166 kO.OO1 5.07 0.195 k0.007 7.165 0.137 kO.010 9.219 0.1425 f 0.002 6.38 1.207 k0.002 4.78 4.00 k0.10 7.08 1.851 kO.004d 3.96 2.25 kO.02 6.49 - 8.32 - 8.96 0.136 k0.003 12.4 1.89 kO.01 11.87 2.006 20.006 11.65 - 11.6 - 11.87 1.725 3.0.05 1.715 f0.05 1.64 k0.05 5.04 0.603 20.006 0.5965f0.004 0.530 k0.005 7.24 514.5 2.957 5.85 9.96 5.06 7.147 9.186 6.35 4.76 7.03 3.94 6.45 8.29 8.93 12.3 11.79 11.58 11.5 11.82 5.02 7.22 632.8 2.933 5.79 9.65 5.01 7.075 9.087 6.28 4.70 6.91 3.88 6.35 8.19 - 12.2 11.56 11.37 11.3 11.65 4.96 7.14 lO4O(ct11 --crl)/C2lm* J-* at Ilnm : 488.0 514.5 632.8 2.398 2.381 2.350 4.67 4.63 4.53 11.40 11.18 10.53 0.858 0.851 0.791 (1.29) (1.26) (1.21) -1.40 -1.38 -1.30 - 1.01 1 - 0.991 - 0.91 8 (2.103) (2.077) (2.014) 5.80 5.72 5.50 2.140 2.129 2.072 3.86 3.82 3.75 5.52 5.47 - - 1.973 1.946 -1.95 -1.93 -1.74 -6.56 -6.48 -6.23 (6.65) (6.58) (6.32) - -6.87 -6.78 -7.33 -7.27 - 2.59 2.57 2.49 (2.19) (2.17) (2.02)TABLE 1 (cont ) .-DEPOLARIZ ATION RATIOS, POLARIZABILITIES AND POLARIZABILITY ANISOTROPIES molecule loop0 at A / n m : 488.0 514.5 lO4Oa/C2 m2 J-1 at I/nm : a lO40(all -al)/C2m* J - 1 at I/nm: a s * 632.8 488.0 514.5 632.8 b 488.0 514.5 632.8 fluoromet hane - - 0.094 k0.002 2.937 2.929 2.904 [3], 2 - - 0.345 difluoromethane 0.118 k O .0 0 4 0.117 fO.004 0.095 k0.003 3.07 3.06 3.04 c (0.409) (0.406) (0.363) trifluoromethane 0.071 kO.01 0.072 fO.01 0.050 50.004 3.145 3.139 3.119 [3] -0.32 -0.32 -0.27 chloromethane 0.787 k0.006 0.779 k0.008 0.755 fO.009 5.12 5.10 5.04 [3], 2 1.768 1,751 1.705 dichloromethane 1.164 k0.005 1.151 k0.005 1.081 k0.003 7.40 7.37 7.26 121 (3.117) (3.084) (2.947) t richlorome t hane 0.691 k0.003 0.689 kO.004 0.652 k0.005d 9.61 9.58 9.47 [3, 281 -3.109 -3.094 -2.974 dichlorodifluoromethane - - 0.687 f0.005 7.54 7.52 7.49 C - - (2.41 5) 2.327 2.308 2.253 bromomethane 0,890 k0.006 0.885 k0.005 0.865 kO.01 6.33 6.30 6.22 [31 3.091 3.062 2.946 iodomethane 0.849 k0.003 0.842 f0.003 0.809 k0.004 8.61 8.57 8.41 ~31 t-but ylchloride 0.156 k0.003 0.152 k0.004 0.129 k0.004 11.5 11.5 11.3 T 1.75 1.73 1.58 t-butylbromide 0.345 kO.004 0.342 20.005 0.329 k0.007 12.9 12.9 12.7 T 2.94 2.93 2.82 ethylene oxide 0.340 k0.004 0.336 k0.004 0.295 kO.01 4.97 4.96 4.91 [31 (1.126) (1.117) (1.034) dimethyl ether 0.375 20.002 0.371 k0.002 0.35 kO.01 5.87 5.85 5.81 [3, 461 (1.397) (1.384) (1.33) dimethyl sulphide 0.509 f0.004 0.506 k0.004 - 8.56 8.51 8.40 T (2.37) (2.35) - hydrogen sulphide 0.0625 k 0.004 0.061 k 0.005 0.044 & 0.003 4.290 4.267 4.199 [27, 40, 561 (0.42) (0.41) (0.34) sulphur dioxide 1.86 kO.01 1.85 fO.01 1.79 k0.01d 4.411 4.389 4.326 [40] (2.359) (2.341) (2.269) - 0.783 nitrogen - - 1.042 k0.006 1.984 1.979 1.967 [6, 491 - nitrous oxide 6.115 k0.02 6.08 k0.02 5.95 f0.04 3.365 3.354 3.318 [5, 14, 601 3.362 3.341 3.267 carbon monoxide 0.521 k0.007 0.519 k0.007 0.480 k0.005d 2.231 2.223 2.200 [37] 0.626 0.622 0.592 a For comparison with literature data in c.g.s.units : 1 e.s.u. (or cm3) = 1.1 12 64 x 10-l6 C2 m2 J-I. References [ ] are selected from Landolt- Bornstein, T refers to liquid data from Timmermans, CSee refractivities in the experimental section. dTaken from Bridge and Buckingham.e Values in brackets are 1040 X 3 4 K] /C2 m2 J-'.3012 MOLECULAR POLARIZABILITY ANISOTROPY FOR GASES hydrogen sulphide (propylene 0.3 %, propane 0.1 %), carbon monoxide (nitrogen 1 %, carbon dioxide 0.9 %). To avoid obscuring the variation of the depolarization ratio with wavelength, the uncertainty shown in table 1 does not include the effect of impurities ; how- ever, Ap fbllows readily from relation (1). REFRACTIVITIES The vaIues of a, the mean molecular polarizability, reported in table 1 have been ob- tained by a critical evaluation of literature data. The mean molecular polarizability a is related to the refractive index n by where Vm is the molar volume, N is Avogadro’s number and co is the permittivity of a vacuum (4x80 = 1.112 65x 10-l’ C V-l m-l = 1 e.s.u.).For gases at low densities n % 1 and (2) reduces to a = (2c0 Vm/N)(n- 1). (3) Most of the mean polarizabilities in table 1 have been obtained from gas-phase refractivities quoted in Landolt-Bornstein l2 (the original literature references, quoted here in square brackets, are given in Landolt-Bornstein ; because of errors in this reference, the original literature was consulted in most cases). Usually the refractivity is reported in a manner which allows for deviations from the ideal gas laws ; where such corrections were not made Vm was calculated using virial coefficients tabulated by Dymond and Smith.13 The quanti- ties n2,60, ni5760 and nd are defined by Landolt-Bornstein ;12 modern values l3 of the virial coefficients of hydrogen yield (nf,760- 1) = 1.000 61 (nd- 1).The following relations obtain : a/C2m2J-l = 6.5910~ (nE760- 1) = 7.1942~ (r~:,57~~-1). Theresultsin table 1 have been selected on the basis of consistency between authors. For some of the compounds studied gas-phase refractivities are not available and in these cases a was calculated from liquid data with the aid of eqn (2). Where a can be calcu- lated from both gas and liquid data agreement is usually within 1 %. Except where noted below values of the refractive index and density of liquids were taken from the com- pilation of Timmermans ;14 these a values are referenced in table 1 as T. The polarizability entry for some compounds, referenced (c) in table 1, needs individual mention.They are, in order : carbon disulphide: Lowrey’s corrections for non-ideal gas behaviour are too large (this was noted by Alms et al.),1° modern values of the virial co- efficients l3 yield the result in table 1 ; dimethylacetylene: based on liquid data,” dispersion assumed to be the same as for methylacetylene ; spiropentane: based on liquid data,16 dispersion assumed to be the same as for cyclopropane ; lY3,5-trifluorobenzene: based on liquid data,l dispersion estimated from results for benzene and hexailuorobenzene ; hexa- Jluorobenzene: based on liquid data ;l difluoromethane: estimated as the average for methane [5, 16, 431 and tetrafluoromethane [5] ; dichlorodifluoromethane: estimated as the average of tetrachioromethane [3, 451 and tetrafluoromethane [5].RESULTS AND DISCUSSION The results of our determinations of the depolarization ratio po, for incident light linearly polarised perpendicular to the scattering plane, are summarised in table 1 together with the mean molecular polarizabilities extracted from the literature. The error estimates attached to the po values are a combination of the reproducibility of results obtained within a series of measurements performed on a given day and those series done at widely separated intervals during which the apparatus was dismantled and reassembled. The error estimate includes, therefore, the uncertainty due to unavoidable differences in alignment. For those species where the presence of impurities leads to a significant uncertainty in po, the error limit may need to be increased accordingly (see experimental section).Table 2 compares the present results with other recent determinations usingBOGAARD, BUCKINGHAM, PIERENS A N D WHITE 301 3 laser light sources ; it shows that results from the principal investigators rarely differ by > 4 %. This figure is an order of magnitude larger than the precision usually attained in the measurement; it is, however, comparable to the 3 % variation in po found by some workers on realignment of the apparatus. The treatment of Bridge and Buckingham shows that the effects of divergence of the incident laser light beam and the angular aperture of the scattered light detector should result in errors < 0.1 % even for CHF3. Slight misalignment of the polarizing prism in the incident beam gives negligible errors.Strain birefringence in the optical windows used is typically rad which results in a negligible error (< 0.1 %) from the window which passes the incident beam. A similar strain birefringence in the window which passes the scattered light yields an absolute error of about 4 x in p o ; this is important only for those species with small depolarization ratios. Particularly significant are the results for COz, OCS, CS2 and N20. These species are readily obtained in a high state of purity and are highly anisotropic, strong scatterers for which strain birefringence and stray light effects are small. Yet the present results differ from those reported by Burnham and co-workers lo* l9 by up to 5 %. Pressure- dependent background scattering may account for these differences but the effects of such scattering are difficult to establish and eliminate with certainty.A slight change (< 1 %) of po with decreasing gas pressure was detected for allene, bromo- methane, chloromethane and isobutane and is mentioned without comment. For these compounds the entries in table 1 are averages for a pressure range 0.1-1 atm. In conclusion, it seems best to assume that the accuracy of depolarization ratios is no better than NN + 3 %. Nonetheless, for a set of po values taken from a given investigator the precision is considerably better than this and the variation of po TABLE 2.-LITERATURE VALUES OF DEPOLARIZATION RATIOS (1oopo) FOR 488.0 AND 632.8 nm. The numbers in parentheses are experimental uncertainties in the last quoted figure.N2 N2O so2 co co2 cs2 ocs CH3CH3 CH, CH2CH3 H2C=CH2 HC=CH CHSCdH C6H6 CH3F CHzF2 CHF3 CH3CI CHC13 CH3Br CH31 CH3CN ref. (10) ref. (19) 1.07 (2) - 6.28 (3) - 1.99 (2) - 4.18 (4) 4.00 ( 5 ) 7.98 (2) - 4.20 (2) - - 0.20 (1) - - - - - - 1.87 (1) - 2.12 (1) 2.10 (5) 1.99 (2) - - 0.10 (1) - - - - 0.78 (1) 0.76 (1) - 0.85 (1) 0.80 (1) 1.68 (2) 1.65 (2) 0.685 (10) - - 488.0 nm others a 1.08 5.5 - 3.8 I - 0.52 0.82 - 1.90 (3)b - - I - - - - this work 6.12 (2) 1.86 (1) 0.52 (3) 4.12 (2) 7.69 (3) 4.00 (2) 0.190 (3) 0.214 (3) 1.266 (5) 1.897 (5) 2.29 (1) 1.98 (1) 0.118 (4) 0.07 (1) 0.787 (6) 0.691 (3) 0.890 (6) 0.849 (3) - 1.535 (3)b 1.73 (5) 632.8 nm ref. (8) this work 1.018 (5) 1.042 (6) 5.96 (2) 5.95 (4) 1.79 (1) 0.480 (5) - 4.03 (1) 4.05 (2) - 7.18 (7) - 3.88 (5) 0.198 (1) 0.166 (1) - 0.195 (7) 1.210 (5) 1.207 (2) 1.851 (4) - 2.25 (2) 1.90 (1) 1.89 (1) - 0.094 (2) 0.179 (5) 0.095 (3) 0.050 (5) 0.050 (4) 0.766 (4) 0.755 (9) 0.652 (5) - - 0.865 (1) - 1.64 (5) - 0.809 (4) (1 Unless otherwise indicated these results are taken from ref.(9) or G. M. Aval, R. L. Rowel1 and J. J. Barrett, J. Chern. Phys., 1972, 57, 3104. From ref. (11).3014 MOLECULAR POLARIZABILITY ANISOTROPY FOR GASES with wavelength is significant. For the dispersion of po the present results are in good agreement with those reported earlier.1° For a gas consisting of randomly oriented molecules in the absence of electronic angular momentum and magnetic fields, the depolarization of Rayleigh scattered light is determined by K the anisotropy of the polarizability tensor aSr : For incident light vertically polarized to the scattering plane, the depolarization ratio p o for Rayleigh scattering at 90" to the incident beam is given by the classical relation The limits of applicability of this expression have been discussed in some detail by Bridge and Buckingham who conclude that for other than the lightest molecules eqn (5) remains valid provided the whole of the rotational Raman spectrum is in- cluded in the measurement of p o ; however, vibrational Raman scattering must be excluded.Observations on the isotropic molecules SF6 and CC14 suggest that the contribution to po from vibrational Raman scattering is of order 2 x This contribution may be significant therefore for all molecules with p o < A recent determination 2o of the depolarization ratio at 514.5 nm for CHF3 explicitly excludes the vibrational Raman contribution and yields po = (2.67 & 0.32) x The result reported here is (7.2 & 1 .O) x K 2 = (3as,as, - %!?a,,)/2~ssa,,.(4) po = 3rc2/(5+4rc2). (5) For molecules with a three-fold or higher axis of symmetry eqn (4) reduces to where all and al are respectively the polarizability components parallel and per- pendicular to the axis of symmetry and a = $apS is the mean polarizability. Applica- tion of eqn (5) and (6) together with the mean polarizability yields the values of (all -al) listed in table 1. The sign of (all -al) has been assigned using experience gained in the examination of the polarizabilities obtained from the Kerr effect in gases and dilute solutions.21 Molecules which belong to the point groups C2, and DZh have three unique polarizabilities axx, a,, and azz, where x, y and z are the principal axes of inertia, and a and K~ are insufficient to determine all three.For these molecules table 1 lists 13arc1. It has been shown 22* 2o that for these asym- metric tops relative intensities of features in the rotational Raman spectrum can yield the magnitude and sign of (axx-a,,,,)/(a-3azZ). Together with a and K : ~ this can yield all three polarizability components. These are available for ethylene 22 and water.2o In favourable circumstances the second-order Stark effect can also yield infor ma- tion on the polarizability anisotropy and, if the mean polarizability is known, all three polarizability components of C2, asymmetric tops can be determined.This has been done for ozone,23* 24 Polarizabilities obtained from the Stark effect are static polarizabilities ; these can be related to the electronic polarizability if the intensity and polarization of the infrared absorption bands are known. There is reasonable agreement O between electronic polarizabilities determined from the static values corrected by subtracting atomic polarizabilities and from the frequency dependence of the light scattering. K = (all -a1)/3a, (6) Two of us (R. K. P. and A. H. W.) acknowledge with gratitude the receipt of a Royal Society and Nuffield Foundation Commonwealth Bursary. M. P. Bogaard and B. J. Orr, Electric Dbole Polarizabilities of Atoms amd Molecules, in International Review of Science, Physical Chemistry, Series 2, ed.A. D. Buckingham (Butter- worth, London, 1975), vol. 2, p. 149.BOGAARD, BUCKINGHAM, PIERENS AND WHITE 301 5 A. D. Buckingham and R. L. Disch, Proc. Roy. SOC. A , 1963, 273, 275. A. D. Buckingham, W. H. Prichard and D. H. Whiffen, Trans. Faruday Soc., 1967,63,1057. M. P. Bogaard, A. D. Buckingham, M. G. Corfield, D. A. Dunmur and A. H. White, Chem. Phys. Letters, 1971, 12, 558. A. D. Buckingham and B. J. Orr, Trans. Faraday SOC., 1969, 65,673. G. A. Victor and A. Dalgarno, J. Chem. Phys., 1969,50,2535. ’ P. R. Certain and L. W. Bruch, Intermolecular Forces, in M.T.P. International Review of Science, Physical Chemistry, Series 1, ed. W. Byers Brown (Butterworth, London, 1972), vol. 1, p. 113. R. L. Rowell, G. M. Aval and J. J. Barrett, J. Chem. Phys., 1971, 54, 1960. * N. J. Bridge and A. D. Buckingham, Proc. Roy. SOC. A, 1966,295, 334. lo G. R. Alms, A. K. Burnham and W. H. Flygare, J. Chem. Phys., 1975,63,3321. l2 Landolt-Biirnstein, Zahlenwerte und Funktionen (Springer, Berlin, 1962), Band 11, Teil8. l 3 J. H. Dymond and E. B. Smith, The Virial Coe$cients of Gases (Clarendon, Oxford, 1969). l J. Timmermans, Physico-Chemical Constants of Pure Organic Compounds (Elsevier, Amsterdam, 1950, 1965), vol. I and 11. lS P. Pomerantz, A. Fookson, T. W. Mears, S. Rothberg and F. L. Howard, J. Res. Nat. Bur. Stand., 1954, 52, 51. l6 V. A. Slabey, J. Amer. Chem. Sac., 1954, 76, 3603. l7 G. C. Finger, F. H. Reed and J. L. Finnerty, J. Amer. Chem. SOC., 1951, 73, 153. l8 M. E. Baur, D. A. Horsma, C. M. Knobler and P. Perez, J. Phys. Chem., 1969, 73, 641. l9 A. K. Burnham, L. W. Buxton and W. H. Flygare, J. Chem. Phys., 1977, 67,4990. 2o W. F. Murphy, J. Chem. Phys., 1977,67, 5877. 21 See for example : C. G. Le Fkvre and R. J. W. Le Fkvre, in PhysicaZ Methods of Chemistry, ed. A. Weissberger and B. Rossiter (John Wiley, New York, 1972), chap. VI, vol. 1, part 3C. 22 G. W. Hills and W. J. Jones, J.C.S. Faraday 11, 1975, 71, 812. 23 W. L. Meerts, S. Stolte and A. Dymanus, Chem. Phys., 1977, 19,467. 24 K. M. Mack and J. S. Muenter, J. Chem. Phys., 1977,645, 5278. l F. I. Panachev, E. Yu. Korableva and M. I. Shakhparonov, Rum. J. Phys. Chem., 1976,50,1130. (PAPER 8/1116)

ISSN:0300-9599    

DOI:10.1039/F19787403008

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Mutual Reaction of Isopropyl Radicals BY P. ARROWSMITH Department of Physical Chemistry, University of Cambridge, Lensfield Road, Cambridge L. J. KIRSCH" AND Shell Research Ltd, Thornton Research Centre, P.O. Box 1, Chester CHI 3SH Received 17th February, 1978 The second order mutual reaction between isopropyl radicals has been studied over the temperature range 301-424 K. Kinetic measurements of the overall termination rate, using a molecular modula- tion spectrometer, have been combined with product analysis studies to determine Arrhenius para- meters for the recombination (kr) and disproportionat ion reactions. The values obtained : )] cm3 molecule-1 s-1 - 1340 J mol-l ( RT kr = 1 . 4 ~ lo-" give good agreement with product analysis studies over a wide range of temperatures. These rate coefficients are correct to +20 % over the stated temperature range, and the error in the activation energies is f2600 J mol-I (95 % confidence limits).The parameters for kf extrapolate to a value that is about twice the value measured by the very low pressure pyrolysis method at 750 K. Until recently the rates of alkyl radical recombination were the subject of some controversy in the 1iterature.l This situation has been remedied by the develop- ment of techniques capable of more direct, absolute measurements of the rates of these processes. Thus data produced by the very low pressure pyrolysis (VLPP) method of Golden and co-workers and by the molecular modulation (MMS) method of Parkes and Quinn have generally been in acceptable agreement by comparison with the order of magnitude discrepancies that have existed hitherto.For the recombination of isopropyl radicals : kr 2C3H7. + C6HI4. Golden et aL2 report log (k,/cm3 molecule-l s-l) = - 1 1.3 & 0.2 over the temperature range 683-808 K ; the MMS technique yields log (kr/cm3 molecule-l s-l) = - 11.1 +0.1 at 298 K. In the present note we report measurements of isopropyl radical recombination over the intermediate temperature range 301-424 K based on kinetic studies using the MMS method. The MMS technique has been described in detail by Parkes et aL3 and was used in the present experiments without essential modification. Briefly, isopropyl radicals were generated by square wave modulated photolysis of azo-isopropyl (AIP) and were detected by absorption spectroscopy using the transition at 235 nm.The modulated absorption signal lags behind the photolysis lamp modulation by an amount depending upon the radical removal rate and the modulation frequency. 3016TABLE ~ . - E X P E ~ N T A L RESULTS FOR THE MUTUAL COMBINATION OF ISOPROPYL RADICALS ref. temperature/K kt/cm3 molecule-1 s-1 o/cm2 molecule-1 kci/kr kr/cmJ molecule-1 s-1 ka/cm3 molecule-1 s-1 present work 301 (9.1k6.6)~ (3.6IfI:l.O)~ 0.60f0.01 (0.62) (8.63.2.0)x 10-12 b 5 . 2 ~ 10-12 Parkes and Quinn 298 1 . 4 ~ 10-l' 3.9 x 10-1 0.65k0.05 (0.62)) present work 323 0.57 f 0.02 (0.60) present work 382 (1.3k0.3)~ 10-l1 (4.25k0.6)~ 0.545+0.01 (0.57)[373 K] (8.2k1.9)~ 4 . 6 ~ present work 424 (1.5+0.3)x 10-l' (5.0k0.6)~ 0.52 C (0.545) (9.91 1 . 9 ) ~ 10-l2 5.1 x Golden et aL2 683 1.0 (0.48) 4.1 x 10-l2 4.1 x 10-l2 Golden et aL2 723 0.75 (0.48) 5.1 x 3.8 x 6 .4 ~ Golden et aL2 768 1.4 (0.47) 4 . 6 ~ 10-l2 Golden et aL2 808 1.5 (0.47) 5.9x 8 . 9 ~ ment of kd/k,. C By extrapolation. a Values in parentheses computed from the expression given by Klein et d4 b Derived from the value of kt from ref. (1) and the present measure-301 8 MUTUAL REACTION OF ISOPROPYL RADICALS This signal was digitized and fed into two up-down counters. The first of these (in phase) counts up when the lamps are on and down when they are off. The second count (in quadrature) lags behind the in-phase count by 90". Such data are obtained over a range of modulation frequencies and substrate pressures. The relationship between the in-phase and in-quadrature counts and their magnitudes permit both the ouerd termination rate coefficient, k,, and the radical adsorption coefficient, 0, to be determined, provided that the photolysis rate is known.The latter was calculated by measuring the decline in the absorption signal (due to removal of the substrate) over an extended photolysis period. The temperature of the reaction vessel was controlled by circulating hot air through a coaxial annular vessel : a temperature of -470 K could be achieved in this manner, although in the present experiment there was a practical limit of -423 K because of the onset of thermal reactions. reduced photolysis period, 7 2/B/cm-2 molecule3 s+ FIG. 1 .-Molecular modulation data for the second order termination reaction between isopropyl Values of k, and cr are shown in table 1.The results at 301 K were obtained from a brief series of experiments to ensure compatibility with the earlier data of Parkes and Quinn.2 At 378 and 424 K much more extensive studies over a wide range of modulation frequencies and substrate pressures of 0.66 - 1.33 kN m-2 were carried out to determine k, and 0 with more precision. In all experiments nitrogen diluent was added to a total pressure of 0.1 N m-'. The results at 424 K are plotted in fig. 1. Here, reduced absorption counts [counts (@)-l], both in-phase and in- quadrature, are plotted against reduced photolysis period (@). The theoretical relationships connecting these reduced variables for second order recombination kinetics have been derived previously.They contain, as unknown parameters, the values of k, and cr alone and these may, therefore, be obtained by a non-linear least squares minimization procedure. The curves shown in fig. 1 were computed using the values of k, and 0 shown in table 1. For the present results, the error limits shown for k, and o are an overestimate of the 95 % confidence limits. radicals at 424 K. (a) u = 5.5 x 10-l8 ; (b) u = 5 . 0 ~ ; (c) u = 4.5 xP. ARROWSMITH AND L. J . KIRSCH 3019 The removal rate constant, measured by the MMS method, is given by k, = kd + k,, where kd describes the rate of the disproportionation reaction : kd 2C3H7.3 CsHG+C3Hs. The disproportionation/recombination ratio must therefore be known to derive ICr from our present kinetic measurements. Klein et aL4 have measured this ratio over the temperature range 77-380 K in both solid and gaseous phases and describe their results by the expression : Because our present data require some extrapolation from the temperature range of Klein et aZ.’s measurements in the solid phase, and because their gas phase measure- ments showed some deviation from their overall expression, we have measured kd/k, over the temperature range 298-273 K in the gas phase.The system used, to be described in greater detail el~ewhere,~ was a coaxial photolysis system designed for geometric similarity with that used in the MMS kinetic experiments. Samples were withdrawn after known periods of extended photolysis and the principal products propane, propene and 2,3-dimethylbutane, were estimated by g.1.c.analysis using a Pye Unicam GCD gas chromatograph equipped with a Porapak Q column, flame ionization detector and a DP 88 computing integrator. Experiments were carried out at 298, 323 and 373 K. At 298 and 323 K the yields of propane and propene were equal within experimental error ; at 373 K there was a slight excess of propane ( N 3 %) which probably arises because of the abstraction reaction :G CSH, +AIP + C,H,+AIP 0 . The ratio kd/k, was therefore measured from the ratio of the propene to 2,3-dimethyl- butane yields. The results obtained are shown in table 1. They confirm the decrease in the value of kd/k, with increasing temperature that was found by Klein et aL4 and, more recently, at higher temperatures, by McKay et aLG Various measurements of kd/k, in the literature have been reviewed by the latter workers.The present results are in good agreement with previous studies at similar temperatures (298-400 K), particularly with the formula given by Klein et al. Extrapolation of our own results to the temperatures used by MacKay et al. (518-573 K) gives a slightly lower range of values (0.490-0.478) than they report (0.52-0.49). Fig. 2 shows an Arrhenius plot of the values of k, obtained in the present work (by combination of the experimental determination of k, and k,/k,, the latter extrapolated at 424 K) and the results reported by Golden et aL2 The expanded ordinate in this plot accentuates the errors of both experiments. However, there is no evidence from the present measurements of a negative activation energy for k, that would directly reconcile the MMS and VLPP work.Rather, the trend is in the opposite sense, an error-weighted statistical analysis of all the MMS/product analysis results yielding a small positive activation energy of 1340 J mol-1 with the 95 % confidence limits at - +2640 J mol-’. The corresponding extrapolation of k, implies a discrepancy of about two at the temperature of Golden et aZ.’s measurements. We believe the reason for this difference may be that the VLPP method determines k, significantly below the high pressure limit, as was indeed suggested by Golden et aL2 in the discussion of their results. There are two pieces of evidence supporting this view. First, the values O f kd/kr obtained directly from the VLPP experiments (table 1) are substantially higher than the value of -0.45 at 750 K predicted by extrapolation of the various3020 MUTUAL REACTION OF ISOPROPYL RADICALS lower temperature product analysis studies.Secondly, the established temperature dependence of kd/k,, taken in conjunction with a negative activation energy for kr, would imply an even more negative activation energy for kd, which is unacceptable for a reaction of this type. The body of evidence, both kinetic and from product analysis, is better reconciled by attributing the temperature dependence of kd/k, primarily to the small positive activation energy for the recombination reaction that our present experiments suggest. We then conclude that the rate of the dispro- portionation reaction is almost independent of temperature. Significantly, the values of kd that may be derived from the present work at 301-424 K and from the VLPP method at 683-808 K are in acceptable numerical agreement (table 1).[The VLPP method yields k, and an approximate value of kd/k, directly. The apparent trend with temperature of kd (table 1) is insignificant in view of the error (kO.5) in kd/kr.] 12 1.1 1.0 0.9 0.8 0.7 0.6 0.5 0.4 - - - - - 0 - 0 0 - - I 1 I I a 1.0 1.5 ' 2.0 25 3.0 - FIG. 2.-Arrhenius plot of data for the recombination rate constant (kr) for isopropyl radicals. In conclusion, we may derive from our results the following Arrhenius parameters to describe the combination of isopropyl radicals over the temperature range 0, Parkes and Quinn ;I 0, present work; 0, Golden et aL2 301-424 K : k, = 1 .4 ~ [ exp ( - 340 mol- l)] cm3 molecule-1 s-1 RT 210 J mol-' kd = 5.0 x 10-12 [exp ( RT )] cm3 molecule-1 s - 1 . These parameters yield rate constants correct to +20 % over the stated temperature range and the approximate error in the activation energies is +2600 J mol-f (95 % confidence limits). They give good agreement with various product analysis measurements of kd/kr over the temperature range 77-573 K. They are also in good agreement with VLPP measurements at -750 K, subject to the recognition that the VLPP determinations of k, measure this rate constant a factor of approximately two below its high pressure limiting value.P. ARROWSMITH AND L. J . KIRSCH 3021 The authors are grateful for useful discussions with Dr. D. A. Parkes (Shell Research B.V., Amsterdam) and to Mr. K. P. Donnelly (D.P.M.M.S., University of Cambridge) for his assistance with the statistical analysis of the results. D. A. Parkes and C. P. Quinn, J.C.S. Faraday I, 1976,72, 1952. D. M. Golden, L. W. Piskiewicz, M. J. Perona and P. C. Beadle, J. Arner. Chem. SOC., 1974, 96, 1645. D. A. Parkes, D. M. Paul and C. P. Quinn, J.C.S. Faraday I, 1976,72, 1935. R. Klein, M. D. Scheer and R. Kelley, J. Phys. Cltem., 1964,68,598. L. J. Kirsch and D. A. Parkes, to be published. G. McKay, J. M. C. Turner and F. Zark, J.C.S. Faraday I, 1977,73, 803. (PAPER 8/281)

ISSN:0300-9599    

DOI:10.1039/F19787403016

出版商:RSC    

年代:1978

数据来源: RSC