摘要:

Adsorption and Polymerization of Acetylene on Oxide Surfaces A Raman Study BY JOHN HEAVISIDE AND PATRICK J. HENDRA* Department of Chemistry, The University, Southampton SO9 5NH PETER TSAI AND RALPH P. COONEY Department of Chemistry, University of Newcastle, N.S.W., Australia 2308 AND Received 20th October, 1977 Laser Raman spectra have been recorded of adsorbed poly(acety1ene) formed by the polymeriza- tion of acetylene on y-alumina and zeolite KX. Spectral evidence indicates that on both y-alumina and zeolite KX monomeric acetylene is physisorbed via its 7~ electron system in a ‘‘ side-on ” orienta- tion. The chain length of the polyene chains formed on y-alumina appear to increase with increasing activation temperature. The line intensities in the Raman spectra of the coloured polyenes are significantly enhanced by resonance effects.The adsorption of acetylene on alumina has been studied previously by infrared rnethod~.l-~ Two distinct modes of interaction have been observed by Yates and Lucce~i,~ and supported by Bhasin et aL5 A stronger “ end-on” interaction involving hydrogen bonding through the acetylenic hydrogen atom and a weaker “ side-on ” interaction involving the n electrons of the acetylene have both been i~lentified.~ As well as adsorbed acetylene, Bhasin et u Z . , ~ reported a coloured product on the surface of alumina. They showed evidence of polymerization by the appear- ance of the C=C stretching band in their infrared spectra. The present study was undertaken to elucidate the nature of coloured surface species resulting from surface induced reactions of acetylene.Previous laser Raman work with the alkali cation exchanged series of zeolites A reveal only a weak “ side-on ” interaction. It has been shown that adsorbed acetylene was more strongly held by zeolite KX than NaX or LiX.’ The tendency of acetylene to polymerize on this surface was therefore also investigated. and X EXPERIMENTAL MATERIALS y-alumina (Grade H supplied by Peter Spence Ltd) was used in the powdered form at a mesh size of 100-200. The B.E.T. N2 surface area is reported by the manufacturers to be z 120 m2 g-l. The zeolite used, was prepared using conventional ion exchange methods from the powdered form of the zeolite NaX, supplied as Linde 13X (by Union Carbide). The amount of potassium exchange was determined by flame photometry and found to be 68 %.Acetylene gas (British Oxygen) was further purified through a purification train consisting of a concentrated sulphuric acid trap and a cold trap (ethanol, 200K). The infrared spectrum of the purified gas showed no evidence of impurity. 2542J . HEAVISIDE, P . J . HENDRA, P . TSAI AND R . P . COONEY 2543 INSTRUMENTATION Laser Raman spectra were recorded on a modified Coderg T800 spectrometer using the green (514.5 nm) or the blue (488.0 nm) emission line of a Spectra Physics model 170 argon ion laser as a source. Laser power levels used ranged from 1 to 130mW at the sample. Plasma radiation was removed from the source using an interference filter (Technical Optics). The laser beam was focused onto the sample surface with a cylindrical lens rather than the more conventional spot focusing lens in order to reduce the light flux at the surface.6 A spectral bandpass of 6 cm-1 was used throughout.Photons were detected using a cooled S20 phototube (EMI-9558) and then processed using a Brookdeal-Ortec 5C1 photon counting system. Supplementary spectra were recorded using a Cary 81 spectrometer in the 90" mode and using an unfocused beam of a Coherent Radiation Laboratories 52G argon ion laser as excitation source. The powdered samples were enclosed in a cell of the type described earlier.g PRETREATMENT OF SAMPLE The powdered y-alumina was activated by heating at temperatures ranging from 300- 1200 K in a vacuum of x mmHg. Samples pretreated in the temperature range 700- 870 K, suffer from a high background fluorescence.To minimize this they were heated in oxygen for 24 h,lO* l1 in addition to the normal pretreatment. All samples were cooled to room temperature under high vacuum prior to adsorption. The zeolite was used in the form of discs pressed at pressures of ~404oOo kg cm-2 in a 1 cm diameter stainless steel split die. The discs were cleaned by slowly heating in vacuo until a pressure of M mm was reached. This precaution was taken to avoid structural breakdown of the zeolite by the sudden removal of water at temperatures above 450 K.12* l3 They were then heated to a final temperature of 820 K in 600 Hg of oxygen for 5-24 h and then cooled to room temperature under a pressure of x mmHg. The X-ray diffraction patterns of the activated zeolite showed no evidence for structural changes as a result of this pretreatment.Using a vacuum system, purified acetylene was introduced at atmospheric pressure to the adsorbents which had been pre-cooled with liquid nitrogen. The precooling of the KX samples (not employed in previous studies)' was important for the production of detectable quantities of polymerized product. Spectra of the adsorption complex were recorded at ambient temperature. RESULTS AND DISCUSSION activated at various temperatures are summarized in table 1. y-alumina and zeolite KX (both activated at 820 K) after adsorption of acetylene. The Raman spectra of acetylene adsorbed on y-alumina and zeolite KX both Fig. 1 shows typical spectra of the adsorbent adsorbate complex formed on PHYSISORPTION OF ACETYLENE )"A L U MI N A Physisorbed acetylene which could easily be removed by evacuation was detected on all samples.The v(C=C) stretching vibration was detected, at about 1962 cm-l in all cases. From table 2 it can be seen that this frequency is considerably shifted downwards from the gas phase value (1974 crn-l)14 and this behaviour is character- istic of acetylene interacting weakly with the substrate, through its n-electron~.~* The alternative " end-on " interaction through the acetylenic proton is also known to occur, 4* but this shifts the v(C=C) stretching frequency to a higher value than in the gas phase. This behaviour has been detected by infrared techniques 4 * although no sign of it was detected in our Raman spectra.P U 0 cn sample (activation temp) TABLE RA RAM AN SPECTRA * OF ACETYLENE ADSORBED ON y-A1203 AND ZEOLITE KX (AP/crn-l) polyene unchanged 645 w 1954 vs 3321 w unchanged 635vw 1961 m unchanged 1962 w cream 1962 w brown 1962 vw purple 1962 vw CH v3 V 2 ddormation mode mode 2’3 v2+”3 2v2 1 0 2 5 1 ~ 1130s 1512s out-of-plane planar planar 2245 w, br 2634 w, br 3027 vw, br 1136m 1520rn 2230 w, br 2608 w, br 3025 vw, br 1013 w 1127 s 1517 s 1122 s 1507 s -2200 w, vbr -2625 w, vbr 3000 vw, vbr 1015 w 1117vs 1501 vs 2225 w, br 2612 w, br 2988 vw, br av, very; w, weak; m, medium; s, strong; br, broad. z 0 1: U cd 0 l? cc2545 The intensity of the band due to physisorbed acetylene appeared to drecrease as the activation temperature of the alumina was raised." As the number of surface hydroxyl groups decreases with increasing activation temperature, we suggest that the physisorbed acetylene interacts through the z-system with them and/or water molecules adhering thereto.This would be an example of sorption to specific sites as distinct from mono- and multi-layer sorption. J. HEAVISIDE, P. J . HENDRA, P. TSAI AND R. P. COONEY I I t 1 3500 3000 2500 2000 1500 1000 500 A;tcm-l FIG. 1 .-Raman spectra of acetylene on (a) y-Al,O, (activated at 820 K) and (b) zeolite KX (activated at 820 K). Exciting line, 514 nm Ar+ (130 mw). Slit width, 6 an-'. TABLE 2.-RAMAN SPECTRA OF THE VARIOUS PHASES OF ACETnENE vibration previous work this work gas liquid 0 crystal KX KX (820 K) y-AlzO3 (300 K) v1 v(G--H) 3374 3341 3324 3318 3321 a v2 v(C=C) 1974 1961 1956 1952 1954 1962 ~ 4 , v i G(HrC-C) 612 630 626 a 645 635 (activation temperature) Q Not observed ; b see ref.(14) ; C see ref. (15) ; d see ref. (16) and (17) ; e see ref. (7). ZEOLITE KX Consistent with previous assignments,6 the bands at Av = 645,1954 and 3321 cm-I are attributed to physisorbed acetylene. Table 2 shows the Raman spectroscopic data for the gas,14 liquid,15 crystalline 16* l7 and adsorbed acetylene on zeolite KX (1954 and 3321 cm-I). The frequencies observed are very similar to those already reported for acetylene on zeolite KA6 and for crystalline acetylene. This would suggest that the adsorbed and crystalline phases are closely related as has been previously suggested. On detailed examination of the spectrum, the band at 1954 cm-l was found to be asymmetric on the lower frequency side suggesting more than one band present.After evacuation, consistent with earlier results for A type zeolites,6 the residual maxima had shifted to 1945 cm-l. Similar effects have been obtained previously 6 e * Fig. l(u) shows the sample activated at 820K in which the band at 1962 cm-' can only just be observed.2546 ADSORPTION AND POLYMERIZATION OF ACETYLENE and have been explained in terms of acetylene bonded to two different types of site in the zeolite framework. H-C= - T C I ! i 0 FIG. 2.-Physisorption of acetylene. POLYENE FORMATION If acetylene is left in contact with activated alumina at room temperature the surface is found to discolour. We therefore studied the Raman spectra of these surfaces with a view to elucidating the nature of the coloured surface species.In the present study, strong Raman lines were found in the 1100 and 1500 cm-l regions in some of the spectra (see table 1). The Raman spectra of polyenes 8 * are known to give rise to two characteristic intense bands in these regions and have been assigned to skeletal vibrations l 8 (see fig. 3). Therefore, we suggest that in this case the two bands in our spectra may be assigned to skeletal modes in a polyene. J J J v2 (1500 cm”) FIG. 3.-Skeletal vibration of polyene chains.J. HEAVISIDE, P. J . HENDRA, P . TSAI AND R. P. COONEY 2547 A weak band can be observed in the spectrum of acetylene on zeolite KX (820 K) at 1025 cm-l and at M 1015 on y-alumina (820 and 1200 K). Shirakawa et aZ.,19 have assigned a band in this region of the trawpoly(acety1ene) spectrum to the trans-CH out-of-plane deformation vibration.The Raman spectra that we observe are hence consistent with that of tran~-poly(acetylene).~~ As well as the funda- mentals, the overtones and combination bands of the symmetric polyene-chain vibrations v3 M 1125 and v2 M 1500 cm-l were observed (see table 1). Ivanova et u Z . , ~ O observed an extended overtone structure of the polyene skeletal mode in the resonance Raman spectra of C2H5COO(CH=CH),COOC2H5, n = 6,7 or 8. They asserted that it is only observed when the exciting frequency falls into the interior of the visible vibrational structure of the absorption band. In this study, a very broad absorption band, ranging from 650 nm into the ultra- violet region, was observed in the reflectance spectrum of the purple polyene on y-alumina (1200K).The two exciting laser lines used (514.5 and 488.0nm) were both in the interior of this absorption band and hence it is likely that the spectrum is resonance enhanced. 1170 1070 \ I I 1 I I 1120 1130 1140 1500 1510 1520 v3 v2 AT1crn-l FIG. 4.-Raman band shifts in the polyene bands with varying activation temperatures of pA1203. The first overtones and combination bands of the skeletal modes were observed in both cases. The intensity of the spectrum was found to be slightly dependent on the exciting line used. In fact, the spectrum was enhanced ,by 60 % using the 514.5 nm line compared with the 488.0 nm line. Gravimetric results show that very little polyene was formed on the surfaces I <0.5 mg g-1 A1203 (1200 K)I, but despite this minute amount, the Raman spectrum of the polyene is very intense.This evidence suggests strongly that the spectrum of the polyene is resonance enhanced. Recently, we have recorded spectra from thermally degraded poly(viny1 chloride).2548 ADSORPTION AND POLYMERIZATION OF ACETYLENE? This material is known to suffer dehydrochlorination to produce unsaturated sectors of the polymer chains containing conjugated double bands. This material produces a Raman spectrum similar to that from sorbed acetylene. In P.V.C. the spectrum is known to be resonance enhanced.21 The v2 and v3 frequencies of the polyene are sensitive to the activation temperature of the y-alumina. These bands shift to lower frequencies with increasing activation temperatures.The plot of frequencies (v, and v,) against the activation temperatures is given in fig. 4. Behringer has attributed the variations in the frequency of the vc=c vibrations in the polyene spectra to differing conjugated chain lengths; an increase in chain length resulting in a decrease of frequency. It appears that the colour changes observed on y-alumina at different activation temperatures (see table 1) are also related to the changes in chain length. Longer chain length increases the conjugation and hence the absorption maxima moves further into the visible region.22 We therefore propose that longer polyene chains are formed on y-alumina activated at high temperatures. It is tempting to speculate on the average length of these polyene sequences produced but we are unable to guarantee the cis: trans isomeric configuration of our product or to be sure that sorption does not perturb the skeletal vibrational frequencies.We feel estimates of chain length could be most misleading. As activation is essential for polymerization, it was thought that the formation of polyenes was linked in some simple way with the acidic surface sites. It has been shown that at higher activation temperatures of y-alumina, an increasing number of active surface sites (of which a large proportion are Lewis acid in nature) are formed.23 This would therefore suggest that the weight of polyene formed should increase markedly with increased activation temperature. As explained above, this cannot be confirmed experimentally, as the amount of polymer formed is minute in all cases.We therefore suspect that the polymerization cannot be taking place at Lewis acid sites activing individually. It is possible that pairs of Lewis sites are involved in the polymerization stage of the reaction, as in olefin isomerisation 24* 2 5 but again we suspect that the amount of polyene produced is too small to satisfy this explanation. It is likely that the polymerization step of the reaction occurs by addition to a terminal C=CH2 group which itself could be absorbed to a Lewis acid pair. Further, the polyene molecules themselves may be adsorbed along their length, thus masking many otherwise active sites, and reducing the amount of polymerization, but we accept that this conclusion is far from satisfactory.Our proposal is supported by the low values observed for the frequencies of the polyene skeletal modes. The authors are grateful to the Office of Naval Research (U.S.N.), ARPA, the S.R.C. and the Australian Research Grants Committee for assistance with funds. J. H. would like to thank the S.R.C. for providing a Postgraduate Award. P. T. thanks the Australian Government for providing a Commonwealth Postgraduate Research Award. N. Sheppard and D. J. C. Yates, Proc. Roy. SOC. A, 1956,238,69. L. H. Little, Infrared Spectra of Adsorbed Species (Academic Press, London, 1966), pp. 143-153. W. A. Pliskin and R. P. Eischens, J. Chem. Phys., 1956,24, 482. D. J. C. Yates and P. J. Lucchesi, J. Chem. Phys., 1961, 35,243. M. M. Bhasin, C.Curran and G. S. John, J. Phys. Chem., 1970,74, 3973. Nguyen The Tam, R. P. Cooney and G. Curthoys, J.C.S. Furaduy I, 1976, 72,2577. Nguyen The Tam, R. P. Cooney and G. Curthoys, J.C.S. Faraday I, 1976,72,2592. * T. A. Egerton and F. S. Stone, J.C.S. Furaduy I, 1973, 69, 22.J . HEAVISIDE, P. J . HENDRA, P. TSAI AND R. P. COONEY lo T. A. Egerton, A. H. Hardin, Y. Kozirovski and N. Sheppard, Chem. Comm., 1971, 841. l1 T. A. Egerton, A. H. Hardin, Y. Kozirovski and N. Sheppard, J. Catalysis, 1974,32,343. l2 R. M. Barrer and W. I. Stuart, Proc. Royal Soc. A, 1959,249,464. l3 J. L. Carter, P. J. Lucchesi and D. J. C. Yates, J. Phys. Chem., 1964, 68, 1385. l4 G. Hemberg, Infrared and Rarnan Spectra of Polyatornic Molecules (Van Nostrand, New York, l5 G. Glockler and M. M. Redrew, J. Chem. Phys., 1938,6,340. l6 G. L. Bottger and D. F. Eggers, J. Chem. Phys., 1964,40,2010 ; 1966,44,4366. l7 M. Ito, T. Yokoyama and M. Suzuki, Spectrochim. Acta, 1970,26A, 695. 2549 P. J. Hendra and E. J. Loader, J.C.S. Trans. Faraday SOC., 1971, 67, 828. 1945), pp. 288-90. J. Behringer, Observed Resonance Raman Spectroscopy, ed. H. Szymanski (Plenum Press, New York, 1967), pp. 168-223. l9 H. Shirakawa, T. Ito and S . Ikeda, Polymer J., 1973,4,460. 2o T. M. Ivanova, L. A. Yanovskaya and P. P. Shorygin, Optika i Spektroskopiya, 1965,18,206. 21 D. L. Garrard and W. Maddams, Macromolecules, 1975, 8, 54. 22 H. H. JafTe and M. Orchin, Theory and Applications of Ultra-violet Spectroscopy (Wiley, 23 P. J. Hendra, I. D. M. Turner, E. J. Loader and M. Stacey, J. Phys. Chem., 1974, 78, 300. 24 W. K. Hall and H. R. Gerberich, J. catalysis, 1966,5, 99. 25 I. D. M. Turner, S. 0. Paul, E. Reid and P. J. Hendra, J.C.S. Faradzy I, 1976,72,2829. New York, 1962), pp. 220-41. (PAPER 711848)

ISSN:0300-9599    

DOI:10.1039/F19787402542

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Complexes of Ammonia and Ethylenediamine with CuI' on Zeolite A BY ROBERT A. SCHOONHEYDT," PAUL PEIGNEUR AND JAN B. UYTTERHOEVEN Centrum voor Oppervlaktescheikunde en Colloidale Scheikunde, Katholieke Universiteit Leuven, De Croylaan 42, B-3030 Heverlee, Belgium Received 15th December, 1977 Adsorption of NH3 on a dehydrated CuA zeolite gives Cu(NH3)$+ and a small amount of tetrahedrally coordinated CuII ions. These complexes are unstable and decompose by vacuum heating in the range 295-373 K to yield lattice-bonded CuII ions characterized by an e.p.r. signal with gll = 2.3087, All = 0.015 09 cm-l, g1 = 2.067 and A L = 0.000 96 cm-'. Above 373 K this signal transforms to gll = 2.3715, All = 0.013 63 cm-l, g1 = 2.062 and A 1 = 0.001 92 cm-'. Adsorption of ethylenediamine on a dehydrated CuA zeolite gives only Cu(en)", stable up to 473 K.However, in the presence of water, Cu(en):+ and Cu(en)$+ can also be synthesized on the surface. The latter two complexes are unstable and decompose by vacuum heating to the mono-complex. Cu(en)'+ can also be loaded on NaA by ion-exchange from aqueous solution, but Cu(en)a+ cannot, due to space requirements and the instability of the bis-complex in the supercages of zeolite A. First it is necessary to exchange the mono-complex and to synthesize the bis-complex in situ by adding large excesses of ethylenediamine to the aqueous suspension. The e.p.r. and reflectance spectra of the various complexes depend slightly but consistently on the environment in the supercages of zeolite A. Their analysis, however, shows that this does not induce major changes in the CuII-N bonding characteristics. The complex formation of Cu" with ethylenediamine (= en) on montmorillonite is strongly favoured with respect to aqueous s01ution.l'~ The absorption band maximum of Cu(en)i+ on these minerals is situated around 20 000 cm-l, 2000 cm-l above its band maximum in aqueous solution.The corresponding gain in crystal field stabilization energy (C.F.S.E.) is 17-25 kJ mol-l .4 On faujasite-type zeolites Cu(en)l+ is destabilized. The degree of destabilization depends on the cation exchange capacity (C.E.C.) of the mineral and the loading with Cu(en)g+.l* Thus, Y zeolites can be loaded with Cu(en)g+ up to -70 % of the C.E.C. without decom- position of the complex. On X-type zeolites the exchange with Cu(en)f+ gives a mixture of Cu(en)i+, Cu(en)2+ and aqueous Cu" ions on the solid.Clearly the negative charge density of the surface plays a major role in these phenomena. To substantiate this idea we extended these studies to the complexation of Cu" with ethylenediamine and with ammonia on A-type zeolites. In this case the Si : A1 ratio equals 1, whereas it is 1.25 and 2.50 for X- and Y-type zeolites respectively. There- fore, A-type zeolites have the highest charge density of this series of high surface area solids. EXPERIMENTAL PREPARATION OF CUII-EXCHANGED A-TYPE ZEOLITES kg NaA, as obtained from Union Carbide's Linde division, were exchanged for 24 h in 1 mol dm-3 NaCl solution at room temperature, washed 4 times and subsequently exchanged overnight in 0.9 dm3 0.003 mol dm-3 CuCI2 at room temperature.The pH of 2550 5 xR . A . SCHOONHEYDT, P . PEIGNEUR AND J . B . UYTTERHOEVEN 2551 the exchange solution before mixing with NaA was -6. The pH of the supernatant after exchange with CuIr was 7.90. The sample was washed Cl--free (AgN0,-test), dried in air at room temperature and stored in a polyethylene bottle. Its water content was 23.4 :$ and its exchangeable cation content was 6.32 meq g-l Naf and 0.99 meq g-l Cu". This sample is indicated by the symbol CuA-10. A second sample, CuA-1, was prepared in identical conditions but with 10 times less CuII in the exchange solution in order to obtain -1 % exchange, suitable for e.p.r. work. These samples were used for adsorption of ethylenediamine and ammonia from the gas phase. Exchange levels above 10 % give partial lattice destruction.Such samples were not investigated in this study. EXCHANGE OF Cu(en)$+ (x = 1,2) ON A-TYPE ZEOLITES NaA, as obtained from Union Carbide's Linde Division, was exchanged twice in 1 mol dm-3 NaCl, washed C1--free and air-dried at room temperature. This material was mixed with solutions with Cu : en ratios equal to 1 or 2 containing enough CuII to achieve 1 and 10 % exchange levels. These materials are denoted as Cu(en)A-1, Cu(en)A-10, C~(en)~A-l and Cu(en)2A-10. Bis-complexes on A were also prepared at 1 and 10 % exchange levels from a suspension with a Cu: en ratio = 1 by adding ethylenediamine progressively. E.p.r. and reflectance spectra of these samples were recorded as such, and after degassing overnight up to 523 K with 50 K intervals.GAS PHASE ADSORPTION OF ETHYLENEDIAMINE AND NH3 ON CUII-EXCHANGED CuA-1 and CuA-10 were degassed in vaciio at 673 or 573 K for 8 . 6 4 ~ lo4 s. O2 was added at these temperatures, the samples cooled and evacuated at room temperature. NH,, purified by freezing and evacuation was adsorbed at room temperature from a calibrated volume in doses of ratio NH3/Cu11 = 1. After addition of each dose the equilibration time was 1.728 x lo5 s prior to recording the spectra. Direct saturation of the samples with 4.666 x lo4 N m-2 NH3 was also performed. Ethylenediamine, previously dried over metallic Na, was adsorbed at room temperature until saturation of the samples, previously degassed as described above, was reached.E.p.r. and reflectance spectra were recorded of these en-saturated samples and after each 5OK interval in the subsequent degassing process. The maximum temperature was 523 K. In another experiment en was allowed to adsorb overnight on the samples degassed as described above, followed by adsorption of water vapour and again ethylenediamine. After each adsorption step e.p.r. and reflectance spectra were recorded. ZEOLITES A TECHNIQUES E.P.R. E.p.r. spectra were recorded at 77 K in the X-band on a Varian E-3 spectrometer. The spectra were simulated with the SIM-13 computer program written by Lozos, Hoffman and Franz of Northwestern University, under the assumption of axial symmetry. g-Values were determined relative to a DPPH standard (g = 2.0036). REFLECTANCE SPECTROSCOPY The spectra were recorded between 2000 and 360 nm on a Cary 17 spectrometer equipped with a type I reflectance unit.The integrating sphere was coated with MgO. The reference was MgO. Both sample and reference were in matching vacuum cells.6 The spectra were digitalized and stored on paper tape with a Hewlett Packard 3480C digital voltmeter and a Hewlett Packard 3489A data punch respectively. They were processed in the IBM 370/158 computer to obtain plots of the Kubelka-Munck function against wavenumber with a Calcomp 502 off-line plotter. The baseline was processed in the same way and could therefore be substracted from the experimental spectra. 1-8 12552 Cu COMPLEXES ON ZEOLITE-A RESULTS COMPLEXES WITH AMMONIA The stepwise adsorption of NH3 on the dehydrated CuA-10 was followed by reflectance spectroscopy.The spectra are shown in fig. 1. After adsorption of 1 NH3/Cu" the 3 band spectrum of dehydrated CuA-10 reduces to a single broad band in the range 10 000-17 000 crn-l. With increasing loadings however, a shoulder n 8 0.128 0.096 0.064- 0.032- 0.000, 5000 15000 25000 wavenumber /ern-' FIG. 1 .-Reflectance spectra of the stepwise adsorption of NH3 on CuA-10 : 1, dehydrated at 550 K ; 2, 0.99 N H ~ / C U I I ; 3, 2.07 N H ~ / C U I I ; 4, 4.00 NH,/CuII ; 5, 12 NH,/CuII. at 15 500 cm-l appears, which grows continuously at the expense of the spectrum of the dehydrated sample. At saturation the latter is only present as a weak shoulder. The 6580cm-l is an overtone of the NH3 vibrations. Fig. 2 shows some typical e.p.r.spectra of CuA-1 saturated with NH3, and after various NH3 desorption steps. Simulated spectra are shown as dotted lines. Table 1 summarizes the g-values and TABLE VALUES AND HYPERFINE SPLITTING CONSTANTS OF THE COMPLEXES OF CUT' WITH NH3 ON ZEOLITE A treatment II A p m - 1 91 AYlcrn-1 A rlcrn-1 structure evacuation at 373 K 2.3087 0.015 09 2.067 0.000 96 - ( 0 1 ) 3 ~ " evacuation at 473 K 2.3715 0.013 63 2.068 0.001 92 - (Od3 cU1I saturation withNH3 2.2480 0.017 61 2.0575 0.001 05 0.001 14 Cu(NH&+ hyperfine splitting constants derived from the simulated spectra. After room temperature saturation with NH, two species are present : most of the CuI1 was in the form of a tetragonal complex with 911 = 2.271 and All = 0.017 89 cm-l, a small amount was tetrahedrally coordinated [shown by the symbol T in fig.2(1)] with approximate parameters 911 = 2.0079, A , , = 0.0132 cm-l and g1 = 2.258. The overlapping spectra prevented a more accurate determination of the parameters. Prolonged saturation with NH3 at room temperature or at 373 K did not alter this spectrum, except for a small decrease in 911 for the tetragonal complex : gI1 = 2.263 and 2.258, respectively after prolonged saturation at room temperature and at 373 K. Evacuation at room temperature and at 323 K produced a new tetragonal species,R . A. SCHOONHEYDT, P . PEIGNEUR AND J . B . UYTTERHOEVEN 2553 I I 2 1 1 DPPH FIG. 2.-Experimental and simulated e.p.r. spectra of CuA-1 in the presence of NH3 : 1, saturation at room temperature for 4 days ; 2, 1 +evacuation at 373 K ; 2', simulated spectrum of 2 ; 3, 2+ evacuation at 473 K ; 4, 3 +evacuation at 773 K in O2 +saturation with NH3 after removal of 0 2 at room temperature ; 4, simulated form of 4.T is the tetrahedral spectrum. The right hand sides of spectra 3 and 4 are enlarged perpendicular regions of spectra 3 and 4 respectively.2554 Cu COMPLEXES ON ZEOLITE-A which became the only one present after evacuation at 373 K (fig. 2, spectrum 2). Its g-values, derived from the simulated spectrum, are shown in table 1 together with the hyperfine splitting constants. A third tetragonal signal is produced by evacuation at 423 K and this in turn became the only species present after evacuation at 473 K. The spectrum is shown in fig. 2 with the perpendicular region on an enlarged scale.This signal remained unchanged upon further evacuation except for a decrease in intensity above 673 K. The sample turned white after evacuation at 773 K. The signal intensity was reduced to 3 its original value but could be restored with 0,. Saturation with NH3 reproduced the original spectrum but much less tetrahedral complex was present. If the spectrum was generated in 250 Torr NH3 all the CU" was converted to the tetragonal complex and N-superhyperfine structure was visible in the perpendicular region. This is shown in fig. 2 (spectrum 4) on an enlarged scale. The g-values and hyperfine splitting constants are shown in table 1. This spectrum did not change upon addition of water. COMPLEXES WITH ETHYLENEDIAMINE The complexation of Cu" with gaseous en was followed by e.p.r.on CuA-1 and by reflectance spectroscopy on CuA-10. Fig. 3 shows the e.p.r. spectra. After 4 days saturation with gaseous en at 298 K the sample turned blue. The e.p.r. spectrum of Cu(en)2+ was then generated, with parameters similar to those of the monocomplex TABLE 2.*-VALUES AND HYPERFINE SPLITTING CONSTANTS OF THE COMPLEXES BETWEEN CU" AND ETHYLENEDIAMINE ON ZEOLITE A sample treatment complex gll ATicm-1 gI AYIcm-1 CUA-1 saturated with en Cu(en)2+ 2.239 0.0182 2.047 0.002 02 CUA-1 en+ H20 Cu(en)$+ 2.205 0.0191 2.050 0.002 85 CuA-1 en+ HzO+ en Cu(en)z+ 2.214 0.0172 2.048 0.00076 Cu(en)zA-l Cu(en)$+ exchange+ en Cu(en)$+ 2.1888 0.0201 2.045 0.001 90 Cu(en),A-1 Cu(en) exchange+ en+ Cu(en)$+ 2.1924 0.0197 2.048 0.002 85 Cu(en)2A-10 Cu(en) exchange+ en+ Cu(en)i+ 2.2075 0.01 94 2.047 0.002 66 Cu(en)A-10 Cu(en) exchange evacuated Cu(en)2+ 2.2667 0.0179 2.053 0.001 43 Cu(en)A-10 Cu(en) exchange, evacuated Cu(en)*+ 2.2485 0.0172 2.053 0.002 95 Cu(en)A-10 Cu(en) exchange, evacuated Cu(en)2+ 2.2595 0.01 59 2.055 0.002 57 NaA Cu(ea)2+ exchange at Cu(H,O);+ 2.3545 0.0151 2.072 O.OO0 96 pH = 5.8 washings NaOH addition to exchange solution NaOH at r.t.during 15 min at 381 K at 439 K in dehydrated X-type zeolite^.^ The e.p.r. parameters are shown in table 2. This complex was stable up to 473 K in vacuu. The spectral intensity decreased upon evacuation at 523 K and the sample turned white. Cu(en)2+ was also generated upon saturation with en at 363 K. When, however, the sample was subsequently equilibrated overnight at room temperature with 20 Torr water vapour, it turned violet and the e.p.r.spectrum was that characteristic of Cu(en)z+ (fig. 3 and table 2). The monocomplex could be regenerated by evacuation at 323 K for 1 h. Saturation with water vapour again yielded Cu(en)$+, showing the reversibility of the complexationR. A. SCHOONHEYDT, P. PEIGNEUR AND J . B. UYTTERHOEVEN 2555 reaction. When the violet sample was saturated with en, which had not been dried over Na prior to adsorption, the sample turned blue and the e.p.r. spectrum was characteristic for the tris-complex Cu(en)$+ (fig. 3 and table 2). This complex was unstable and evacuation at 323 K resulted in the formation of the mono-complex. FIG. 3.-E.p.r. spectra of Cu(en)2+ (l), Cu(en)i+ (2) and Cu(en)<+ (3) formed by adsorption of en and water from the gas phase on dehydrated CuA-1. The dotted lines represent the simulated spectra, the full lines the experimental spectra.2556 Cu COMPLEXES ON ZEOLITE-A Qualitatively the same phenomena were observed on CuA-I0 both by e.p.r.and reflectance spectroscopy, but the conversions dehydrated -+ mono -+ bis -+ tris were only partial. Fig. 4 gives some reflectance spectra for illustration. Thus, a 4 day adsorption period of en on CuA-10 even at 338 K converts Cu" only partially to Cu(en)'+. This complex absorbs at 14 540 cm-', while the band around 11 300 cm-I is reminiscent of lattice-bonded CU".~~ * Subsequent adsorption of water produces Cu(en)$+ with a band maximum at 18 400 cm-' but the spectrum shows shoulders at 11 500 and 14 500 cm-', characteristic of lattice-bonded Cu" and Cu(en)2+.Adsorption of en on this violet sample turned it blue but both e.p.r. and reflectance spectra were ill-defined. The band maximum was at 16 700 cm-I, but the band was asymmetric towards its low frequency side. In any case, evacuation above room temperature produced the mono-complexes again, besides the originally present lattice-bonded Cur'. The former was gradually converted to the latter upon prolonged evacuation at 473 K. 0.064r 14OOO 18500 23000 wavenumber /cm-' FIG. 4.-Reflectance spectra of the complexes of CuII with ethylenediamine on CuA-10 : 1, saturated with en vapour at 338 K ; 2, 1 +water vapour ; 3,2+en. EXCHANGE OF THE CU"-ETHYLENEDIAMINE COMPLEXES If an aqueous solution with Cu" : en = 1 at pH = 5.8 was mixed with an aqueous suspension of NaA, previously brought to the same pH with HN03, and if the pH was kept in the range 5.8-6.3 during exchange, the resulting zeolite contained aqueous Cu" ions as shown by the e.p.r. parameters in table 2.Addition of en to this suspen- sion increased the pH, the zeolite became blue and both e.p.r. and reflectance spectro- scopy indicated the presence of aqueous Cu" and Cu(en)2+. At pH = 7.5 Cu(en);+ was formed on the zeolite too but it was only at pH > 10 that nearly all the Cu" was converted to the bis-complex. If this suspension was left standing in the laboratory, the supernatant became violet and the zeolite white, indicating that Cu" migrates out of the zeolite to form Cu(en),"+ in solution.If the exchange was performed directly in a Cu(en),"+ solution, no Cu(en)f+ was found on the zeolite even at pH = 11, and only after repeated washings with en solutions was it possible to record minor quantities (-0.5 % of the C.E.C.) of Cu(en);+ on the solid. The e.p.r. para- meters of these Cu(en)2A systems, determined from the simulated spectra, are summarized in table 2. There was a small but definite dependence of these quantitiesR . A . SCHOONHEYDT, P. PEIGNEUR AND J . B . UYTTERHOEVEN 2557 on the mode of preparation. The corresponding reflectance spectra are shown in fig. 5. The suspension of Cu(en),A-10, prepared by exchange of the mono-complex and addition of en until pH = 11.2, is characterized by a broad band with maximum at 17 600 cm-I.Upon evacuation, the band becomes narrower and shifts to n 8 wavenumber Icrn-' FIG. 5.-Reflectance spectra of Cu(en)zA-fO : 1, suspension prepared by addition of en until pH = 11.2 to Cu(en)A-10 (the real spectral intensity is 10 x larger than shown) ; 2, evacuated at room temperature during 300 s ; 3, evacuated at 323 K during 8.64 x lo4 s. 18 000 cm-l. Upon heating, the band maximum moves again to smaller wave- numbers and the band becomes strongly asymmetric at its low frequency side. This indicates that a decomposition of Cu(en)i+ occurred with formation of the mono- complex and above 373 K also lattice-bonded Cu". This was confirmed by the wavenumber /cm- l FIG. 6.-Reflectance spectra of Cu(en)A-10 : 1, Cu(en)A-10, evacuated at room temperature for 900 s ; 2, evacuated at 329 K for 8.64 x lo4 s ; 3, evacuated at 381 K for 8.64 x lo4 s ; 4, evacuated at 439 K for 8.64~ lo4 s.2558 Cu COMPLEXES ON ZEOLITE-A analogous e.p.r.experiments on Cu(en),A-1 . A comparison of the decomposition of Cu(en),A-1 and Cu(en),A-10 revealed that up to 473 K the ratio mono : bis was higher for the 1 % exchange level (55 : 45) than for the 10 % exchange level (33 : 60). Above 473 K Cu" was reduced and the zeolite turned white. The decomposition of the mono-complex on Cu(en)A-10 is illustrated in fig. 6. Besides the combination band of N20 at 5270cm-l and the overtones of the NH and OH stretching vibrations in the region 6800-7000 cm-l, the spectrum contains an asymmetric band due to Cu" at 14 500 cm-l with a shoulder around 17 150 cm-l.Upon evacuation the spectrum intensifies and the band maximum shifts to 13 730 cm-l. After evacuation at 381 K it is again situated at 14 500 cm-l but is asymmetric towards the low frequency side. At 439 K a strong reduction of intensity occurred with no significant change in the band position. The e.p.r. parameters of these monocomplexes are given in table 2. In every case only one species was detected, except after room temperature evacuation. Besides a signal due to Cu(en),+ we saw a signal due to CU" with oxygen ligands. The g-values and hyperfine splitting constants of Cu(en)*+ on these ion exchanged samples are different from those of Cu(en),+ on the dry zeolite. The differences are more pronounced than for Cu(en),"+. No spectra were recorded above 473 K because Cu" was reduced.DISCUSSION COMPLEXES OF CU" WITH NH3 CU" readily coordinates dry, gaseous NH3 as shown by the disappearance of the characteristic spectrum of dehydrated CuA8 even at very small loadings. For NH3/Cu11 < 1 the broad band in the region 10 000-17 000 cm-l reflects the presence of several species Cu(NH,);+ (x = 1-4) and lattice-bonded Cu". Above 1 NH,/CU~~ the situation is clearer. There is a systeniatic growth of a band at 15 500 cm-l at the expense of the 10 800 em-' band, characteristic of lattice-bonded Cu". The position of the band maximum of the ammonia complex is identical to that of the tetraammine in X- and Y-zeolites with small Cu" contents. We therefore assign the 15 500 cm-l band to Cu(NH,)i+. This is, however, not the only species present, Some Cu" ions are tetrahedrally coordinated, probably in the form [(OJ,-CU-NH,]~+ where O1 is a lattice oxygen of the six-membered rings.Evidence for this comes both from the e.p.r. spectra which showed a species with gL N 2.257 and gll z 2.0079 and from the reflectance spectra in the form of a low frequency shoulder on the 15 500 cm-l band. The e.p.r. parameters of Cu(NH,);+ in A-type zeolites depend slightly but in a systematic way on the environment. Thus, during the process of exposure to NH3 glr diminishes from 2.271 to 2.258. Mter destruction of the complex and regeneration of the Cu"A sieve readsorption of NH3 gives Cu(NH,)i+ with gll = 2.248 and a clear N-superhyperfine structure in the perpendicular region. This environmental dependence of the characteristics of the Cu-ammine complexes is a general observation in zeolite^.^^ However, the tetra-ammine is less stable in A-type sieves than in X or Y.In the latter cases, degassing produced an intermediate [(O,),-CU-NH,]~+, stable in the range 356-383 lo Room temperature evacua- tion of NH,-saturated A produces an e.p.r. signal with 911 = 2.3087, reminiscent of CuI1 surrounded by oxygen-containing ligands. This is therefore lattice-bonded CuI1 and the only species present after degassing at 373 K. At high temperature a second lattice-bonded Cu" species is produced with 911 = 2.3715, which became the only species present at 473 K. The same two Cu" signals were also detected after degassing a hydrated Cu"A sieve at 673 K.8 In the configuration with the largest gll, Cu" is quantitatively converted toR.A . SCHOONHEYDT, P . PEIGNEUR AND J . B . UYTTERHOEVEN 2559 Cu(NH,)i+ upon adsorption of NH3. This was not the case after the initial NH3 adsorption, i.e. when the two e.p.r. signals of lattice-bonded Cu" were present. We conclude that the configuration giving rise to 911 = 2.3715 is fully accessible to NH3, whereas the configuration with 911 = 2.3087 corresponds to Cu" ions on such sites that a small amount of Cu" remained tetrahedrally coordinated upon NH3 adsorption. We suggest that the signal with 911 = 2.3715 is due to Cu" on the hexagonal rings in the supercages and the signal with gll = 2.3087 to Cu" on the hexagonal rings inside the sodalite cages. Two similar Cu" signals are detected in dehydrated CuY zeolites, one with 911 = 2.38 and one with 911 = 2.32.The former has been attributed to Cu" on sites 11, the latter to Cu" on sites I'.I1 This assignment is analogous to ours. COMPLEXES OF CU" WITH ETHYLENEDIAMINE The adsorption of dry, gaseous ethylenediamine on a dehydrated Cu"A sieve is peculiar with respect both to the NH3 adsorption, discussed in previous paragraphs and to the adsorption of ethylenediamine on X and Y zeolite^.^ Thus, in the latter cases tris-complexes were synthesized even at small en-loadings. In the present case only the mono-complex is formed on the surface. This preference of A-type zeolites for the mono-complex is further substantiated by our experimental observations that bis- and tris-complexes can only be made in the presence of water. Moreover they are very unstable, as simple degassing at 323 K restored the mono-complex, which itself is stable up to 473 K.At higher temperature CuI1 is reduced. The e.p.r. parameters and the reflectance spectra of these complexes are very similar to those of the same complexes on the surface of X- and Y-zeolites. Therefore, this peculiar behaviour of the Cu" ethylenediamine complexes in A-type sieves with respect to X- and Y-zeolites reflects, in the first place, differences in environment. The critical diameter of the supercage of A-type zeolites (1.14 nm) is smaller than that of X- and Y ( N 1.3 nm), but more important, the surface negative charge density is higher and so is the number of exchangeable Na+ ions per supercage.Thus, without water the mono-complex is formed with Cu" still coordinated to 3 oxygens of the six-membered rings. A second ethylenediamine molecule is unable to approach the Cu" of the mono-complex within coordinating distance, but the small water molecule can. As soon as the mono-complex is freed from the surface by water, a second ethylenedi- amine molecule coordinates to Cu" to form the bis-complex, because ethylenediamine is a stronger ligand than water. In this picture we have neglected the role of the residual Na+ ions, because their coordinating power is several orders of magnitude less than that of CuII, at least in aqueous solution. This may be different on the surface of a dehydrated sieve, with its much smaller electrical permittivity, but no experimental data are available to substantiate this idea. In any case, the bis- and tris-complexes in the supercages of CuA are unstable and water is needed to stabilize them, i.e.to shield them from the attracting power of the negatively charged surface. Therefore, removal of water destroys them and restores the mono-complex again. In an aqueous suspension not only the water molecules, but also the dimensions of the complexes in solution, with respect to the free diameter of the 8-membered rings, and the pH, determine the complex loading of zeolite A. Thus, the mono- complex can be exchanged directly onto the A-type sieve above yH = 6. This is impossible or almost impossible for the bis-complex because the dimensions of the complex (0.47 x 0.77 nm) exceed the free diameter of the 8-membered windows (-0.44 nm).Another factor is the instability of the bis-complex in the supercage of zeolite A : the bis-complex can by synthesized only in aqueous suspensions by first exchanging Cu" or the mono-complex and then adding large excesses of2560 Cu COMPLEXES ON ZEOLITE-A ethylenediamine to the suspensions (pH The instability of the bis-complex is further substantiated by our observation that Cu'I moves out of the zeolite into the en solution upon standing overnight. The e.p.r. parameters and band maxima in the reflectance spectra of these hydrated Cu(en)A and Cu(en),A samples are significantly different from those of the dry sieves, which again indicates the effect of the environment on the complexes in the supercages. This idea is further evidenced from the behaviour of Cu(en)A-10 upon evacuation (fig.6). Before evacuation some Cu(en)$+ is present on the surface, as shown by the high frequency shoulder on the 14 500 cm-l band of Cu(en)2+ in the reflectance spectrum. After evacuation at room temperature no trace of Cu(en)z+ is left but some Cu" ions are in an oxygen environment. This is shown by the shift of the 14 500 cm-l band to lower wavenumbers and the appearance of the characteristic e.p.r. signal. Any heating produces only Cu(en)2+, with e.p.r. parameters strongly dependent on the treatment (table 2). This observation could not be matched with Cu(en),A samples as the bis-complex partially decomposed to Cu(en)2+. However, this reaction was not quantitative, in contrast to the behaviour of these complexes on the dry sieves.Higher complex loadings and water seem to favour the bis- complex. 11). CHARACTERIZATION OF THE CU" AMINE COMPLEXES On the dry sieves the band maxima of Cu(NH,)i+, Cu(en),+ and Cu(en)d+ are located at 15 500,14 540 and 17 400 cm-l respectively, significantly below those of the same complexes in aqueous solution (16 600 cm-l, 15 500 and 18 200 respectively). TABLE 3.--M.O. COEFFICIENTS OF THE 3d ORBITALS OF Cur' IN THE M.O. EXPRESSIONS OF THE CUII-AMMINE COMPLEXES ON A-TYPE ZEOLITES 8 2 ( h Z , dY4 U out-of-plane n complex in-plane n c?2(dXZ-Y2) 8 ? ( h Y ) Cu(NH)&+ on A 0.89 0.92 0.88 on X, Y : low loading 0.88 0.90 0.66 high loading 0.87 0.93 0.69 &(en):+ gas phase adsorption of wet Cu(en)2A-l 0.88 0.85 0.83 wet Cu(en)2A-10 0.85 0.93 0.87 dry Y 0.85 0.89 0.73 wet Y 0.84 0.92 0.76 wet X 0.81 0.90 0.77 solution 0.85 0.89 0.82 en+H20 on A 0.84 0.92 no convergence In fully hydrated sieves the band maximum of Cu(en),+ remained around 14 500 cm-1 but that of Cu(en)j+ shifted towards 18 000 cm-l.These observations, in combina- tion with the variation of the e.p.r. parameters of these complexes with their environ- ment in the supercages, prompted us to examine these complexes in more detail. This is given in the Appendix. The general conclusion emanating from table 3 is that there is no significant effect of the various preparations of Cu(en)l+ in A-type zeolites on the Cu-N bonding characteristics. There is, however, a tendency for smaller p2 values in X- and Y-zeolites than in A.This is the coefficient of the out- of-plane n-orbital and therefore susceptible to direct interaction with surface orbitals.R . A . SCHOONHEYDT, P. PEIGNEUR AND J . B . UYTTERHOEVEN 2561 This is in agreement with our previous conclusions.12 The effect is small. The zeolite matrix should be considered as a solvent with weak coordination power, at least with respect to the coordination of the ammines considered. As a solvent it drastically influences the complex formation when compared with water. R. A. S . is indebted to the N.F.W.O. (Belgium) for a grant as '' Bevoegdverklaard Navorser ". P. P. acknowledges financial support from the Belgian Government. The technical assistance of F. Pelgrims is greatly appreciated. The authors thank Prof. Smets for the use of his e.p.r.machine. This work was made possible through the finances of the Belgian Government (Staatssekretariaat voor Wetenschapsbeleid). APPENDIX It is a reasonable assumption that the effective symmetry of Cu(NH3)$+ and Cu(en);+ is D4h. Cu(en)2+ and Cu(en);f have lower symmetries, but reflectance and e.p.r. techniques on these powdered materials do not resolve them, As a consequence we have omitted them in this calculation. The experimental band maxima of Cu(NH3)$+ and Cu(en)$+ are assumed to be the average of the 3 possible spin-allowed, parity-forbidden transitions in O4h (2B2g c 2Blg, 'Alg +- 2Blg, 2Eg c 2Blg). The m.0. coefficients were calculated with the theory of Moreno and Barrius0,~~9 l4 which was successfully applied to Cu(en)$+ on X and Y zeolites and on Camp Berteau montmorillonite. l2 Our previous data for Cu(NH3)$+ on synthetic faujasites were recalculated and incorporated in table 3 together with the present results. In view of the approximations involved in the band assignments and our estimate of the overlap integrals l2 the m.0. coefficients of table 3 can only be used for comparison purposes. P. Peigneur, Adsorption A$nity and Stability of Transition Metal Ammine Complexes on Alumina-silicates, Ph. D. Thesis (K. U. Leuven, 1976). A. Maes, P. Marijnen and A. Cremers, J.C.S. Faraday I , 1977, 73, 1297. A. Maes, P. Peigneur and A. Cremers, J.C.S. Faraday I, accepted for publication. F. Velghe, R. A. Schoonheydt, J. B. Uytterhoeven, P. Peigneur and J. H. Lunsford, J. Phys. Chern., 1977,81,1187. P. Peigneur, J. H. Lunsford, W. De Wilde and R. A. Schoonheydt, J. Phys. Chem., 1977, 81, 1179. F. Velghe, R. A. Schoonheydt and J. B. Uytterhoeven, Clays and Clay Min., 1977, 25, 375. ' P. A. Jacobs, W. De Wilde, R. A. Schoonheydt, J. B. Uytterhoeven and H. Beyer, J.C.S. Furuday I, 1976,72, 1221. R. A. Schoonheydt, P. Peigneur and J. B. Uytterhoeven, to be published. W. De Wilde, R. A. Schoonheydt and J. B. Uytterhoeven, A.C.S. Symp. Ser., 1977, 40, 132. lo E. F. Vansant and J. H. Lunsford, J. Phys. Chem., 1972, 76,2860. l 1 W. Morke, F. Vogt and H. Bremer, 2. anoug. Chem., 1976,422,273. l 2 R. A. Schoonheydt, J. Phys. Chem., 1978, 82, 497. l 3 M. Moreno and M. T. Barriuso, Anales Fisica, 1975, 71, 205. l4 M. Moreno and M. T. Barriuso, Solid State Comm., 1975, 17, 1035. (PAPER 7/2194)

ISSN:0300-9599    

DOI:10.1039/F19787402550

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Radiotracer Studies of Self-diffusion in the Plastic Solids Norbornylene and Norbornane BY ALAN V. CHADWICK" AND JACQUES W. FORREST 1- University Chemical Laboratory, University of Kent, Canterbury, Kent CT2 7NH Received 28th December, 1977 Radiotracer self-diffusion measurements have been made for norbornane (bicyclo[2,2, llheptane) and norbornylene (bicycIo[2,2,l]hept-2-ene) in their h.c.p. plastic crystalline phases. In the case of norbornane the measurements were limited by the poor quality of the specimens; however, the results are consistent with existing n.m.r. data and support the view that self-diffusion occurs via a mono- vacancy mechanism. More extensive measurements were possibJe for norbornylcne, and in the temperature range 251 to 308 K the atmospheric pressure results fit the equation Measurements on oriented norbornylene specimens showed that any anisotropy of the self-diffusion was within the experimental error.Pressure effect studies for norbornylene at 281 and 301 K yielded an activation volume equal to 0.840.1 times the molar volume. These results are in excellent agreement with existing n.m.r. data and confirm a monovacancy mechanism of self-diffusion. The present work suggests that in cases where there is a disagreement between radiotracer and n.m.r. measurements this cannot be explained simply in terms of the entropy of fusion of the material. Translational self-diffusion in plastic crystals, molecular solids in which the molecules are globular and undergo rapid endospherical reorient ation, has been monitored by radiotracer, nuclear magnetic resonance (n.m.r.) and plastic deformation (creep) measurements.' 9 Trends in the diffusion parameters have been associated with variations in the entropy of fusion, ASf, and the entlialpy of self-diffhion, AH, is usually compared with the latent heat of sublimation, E,.The activation enthalpies from n.m.r. measurements, AHn, exhibit an apparent dependence on ASf ;1-4 AH, increases from 1 L, to 2 L, as ASf increases from 1 PP to 2.5 R, where R is the gas constant. The activation enthalpies from tracer measurements, AH,, are consistent with the values from creep experiments and are usually equal to -2Ls irrespective of AS,.'. Thus for materials with AS, < 1.7 R values of AH, are usually greater than those of AHn. This could be taken as an indication that the two techniques are not monitoring the same process,2 although it has been noted that the differences between the n.m.r.and radiotracer self-diffusion coefficients, D, and D,, are not large. More measurements for low ASf materials are needed to understand the diffusion behaviour in these systems. Since the plastic crystals have close-packed structures, the predominant point defects are expected to be lattice vacancies and therefore self-diffusion is usually assumed to occur by a monovacancy mechanism.l* Detailed calculations of AH are not available for plastic crystals, but to a first approximation it would be expected to be similar to that found for rare gas crystals. In these latter materials theoretical calculations of AH for a monovacancy mechanism yield AH !Y 2 L, and this is in f Present address : Materials Section, Building Research Station, Garston, Herts.2562A . V . CHADWiCK AND J . W. FORREST 2563 agreement with experiment.6 In addition to a simple comparison of AH and L,, three other methods have been employed to obtain more direct information on the diffusion mechanisms in plastic crystals ; (i) measurement of the isotope mass effect on self-diffu~ion.~* * and (iii) measurement of the effect of hydrostatic pressure, P, on diffusion.1°-12 The first two methods make use of the fact that in the solid state the translational diffusion of a particular molecule is not necessarily a random process and successive jumps of the molecule may be ~0rrelated.l~ The correlation factor, f, is a measure of the degree of correlation and can be calculated for a given diffusion mechanism, lattice structure and experi- mental diffusion technique.In method (i) the correlation factor for tracer diffusion, ft, is determined and compared with the theoretical values. The ratio D,/D, is equal toft&, wheref, is the correlation factor for n.m.r. diffusion, and in method (ii) this ratio is compared with the theoretical values. This latter method is less direct and at presentf, is not well-known for all structures. It also requires careful calibration of the temperature scales in the two experiments. Studies of D as a function of P allow evaluation of the activation volume for self-diffusion, A V,. Information on the diffusion mechanism can be obtained by comparing the ratio of AV, and the molar volume, kT,, with theoretical estimates for the various diffusion mechanisms and with values of this paranieter in other systems in which the mechanism has been identified from other types of experiment.In this paper we report the results of radiotracer studies of self-diffusion in norbornam (bicyclo[2,2,l]lieptane) as a functioa of temperature, T, and in norborny- lene (bicyclo[2,2,l]hept-2-eiie) as a function of T and P. Norbornyiene has one plastic phase l4 (129 K-melting point) in which the structure is close to being ideally hexagonal close-packed (h.c.p.) with c/n = 1.61 l 5 and ASI. = 1.22 R.'" Norbornane has two plastic phases and AS, = 1.53 R.14 In the lower phase (131-306 K) the structure is h.c.p. with cia = 1.50 l 5 and in the higher phase (386 K-melting point) the structure is face-centred cubic (f.c.c.).For both these materials n.m.r. spin lattice relaxation times in the rotating frame, Tlp, have been measured as a function of T 4 and P.lo* Since these are low ASf materials the comparison of tracer and n.m.r. diffusion parameters will be important. (ii) Comparison of D, and D, 6 * EXPERIMENTAL MATERIALS The details of purification of the niaterials and preparaiion of the tritium-labelled tracers have been described elsewhere.16 Vapour phase chromatography indicated that norbornylene and norbornane contained < 1000 p.p.m. and < 10 p.p.m. impurity, respec- tively. The major impurity in norbornylene was its isomer nortricyclene (tricyclo[2,2,1 ,03s5]- heptane) which we were unable to remove completely; however, since the norbornane was prepared by hydrogenation of the norbornylene it seems reasonable to assume that the levels of other impurities were -c 10 p.p.m.Good crystals of norbornylene (10 mm diameter x 50 mm long) were gown from the melt by the Bridgman-Stockbarger technique and by sublimation growth. The orientations of the melt-grown crystals were determined by X-ray methods. Samples were cut parallel and perpendicular to the c axis, the maximum angle of misalignment being about 10'. A crystal deliberately doped with 7000 p.p.m. nortricyclene was melt-grown. Norbornane crystals grown from the melt shattered on cooling through the phase transition at 306 K. Good crystals of the h.c.p. phase were obtained by sublimation growth; however, they were sensitive to thermal shock.Cooling these crystals to 243 K or cutting them with a razor blade introduced cracks and although these annealed out after several hours, the samples were considered to be of a poor quality for diffusion experiments.2564 SELF-DIFFUSION IN PLASTIC SOLIDS DIFFUSION PROCEDURE The basic sample preparation, annealing and sectioning procedures were similar to those used in previous work.5 The tracer deposit was applied to the sample in the form of a saturated solution in n-pentane and the sectioning was performed in a deep-freeze maintained at 243 K. During an anneal the diffusion couple was encapsulated in mercury to eliminate evaporation, Diffusion anneals under pressure were performed in an apparatus similar to that described by McKay and Sherwood.12 The essential features of the system were a commercial pressure vessel (Pressure Product Industries), a liquid pressure transmitting system, a manual hydraulic pump (Enerpac P228-7) and a calibrated Budenberg gauge.The temperature and pressure variations of the sample during anneals were < +O.l K and < +2 MN m-2, respectively. Standard procedures were used to correct the anneal times for the heating and cooling, and pressurising and de-pressurising periods. RESULTS NORBORNANE The solution of the diffusion equations for this type of tracer sectioning experiment is C(x) = ____ ‘ exp ( -x2]4D,t). (zDtf)* Here C(x) is the specific activity of tracer that has penetrated to a depth x into the crystal, Q is the total amount of tracer initially deposited on the sample and t is the time of the diffusion anneal.100 I x2/10-’ m2 FIG. 1.-Typical diffusion profile for norbornane : 9, raw data points for a run at T = 303.5 K with t = 233 200 s ; x , the same points after subtraction of the tail (see text). The straight line fit yields Dt = 9.8 x 10-ls m2 s-l. A typical profile for norbornane is shown in fig. 1 and clearly this cannot be fitted to eqn (1) with a unique value of D,. The long tail is typical of pipe diffusion,’. l7 i.e. diffusion along bulk defects such as grain boundaries and dislocations. This wasA . V . CHADWICK AND J . W. FORREST 2565 to be expected from the poor quality of the samples. By subtracting a linear extra- polation of the tail off of the raw data points it was possible to estimate the contribution from lattice diffusion to the profile.The fitting error in D, from this procedure was +30 % ; however, the absolute error was difficult to assess. At worst the values of Dt obtained represent an upper limit to the true values. The poor quality of the sample and the low diffusion coefficients meant that measurements were restricted to a very short temperature range, 293-304 K. In fig. 2 the present results are compared with the n.m.r. data. lo3 KIT FIG. 2.-Self-diffusion coefficients for norbornane plotted as a function of reciprocal temperature : a, radiotracer results ; x , n.m.r. data.4 T.P. is the h.c.p.-f.c.c. transition temperature. NORBORNYLENE Typical diffusion profiles for norbornylene are shown in fig.3. In contrast to norbornane there was only a slight tailing at deep penetration. Diffusion coefficients were evaluated by neglecting these tails and fitting the remainder of the profile to eqn (1). The fitting errors gave an uncertainty in D, of about 10 %. The profiles x2/10-' m2 FIG. 3.-Typical diffusion profiles for norbornylene : the lines are best fits to the points and the arrows indicate 9Dtt. m2 s-' ; 0, T = 251.3 K, t = 240 4320 s, P = atmospheric, Dt = 1.45 x 10-15 m2 s-I ; +, T = 301.2 K, t = 226 440 s, P = 34.5 MN m-2, Dt = 2.67 x x , T = 281.0 K, t = 272 160 s, P = 0 , T = 299.3 K, t = 150 600 s, P = atmospheric, Dt = 6.49 x m2 s-' ; atmospheric, Dt = 1.75 x m2 s-l.2566 SELF-DIFFUSION IN PLASTIC SOLIDS were linear within the region x < 3Jzt which is an indication that the values of D, were representative of lattice diffusi0n.l When it was possible this point was further checked by showing that D, was independent of the anneal time.Unfortunately this was not possible when D, was -10-15 m2 s-I since the arsneal times necessary to obtain the minimum reasonable penetration were around the maximum limit for convenient and accurate measurements, i.e. one month. The anneal times used were similar to those employed by Hampton and Sherwood in their careful tracer study of cyclohexane. Thus precautions were taken to minimize the effect of pipe diffusion. If these procedures were not completely successful then D, would be an over-estimate of the lattice diffusion and the error would increase with decreasing values of D,.Consequently AH, and A V, would be underestimated. 103 KIT FIG. 4.-Self-diffusion coefficients for norbornylene plotted as a function of reciprocal temperature : 0, 8 and + are the radiotracer results for crystals with faces (1 c-axis, I c-axis and with random orientations, respectively ; x , n.m.r. data.4 M.P. is the melting point. The atmospheric pressure results and the corresponding n.m.r. data are shown in fig. 4. The results for oriented crystals show that any anisotropy of self-diffusion is less than the experimental error. Since the structure of norbornylene is close to ideal h.c.p. this result was not unexpected. At 273 K, D, for the nortricyclene doped crystal was the same as that for the nominally pure crystals. Again this is not an unexpected result due to the similarity of the structures of dopant and host and the relatively low doping level employed.A least-squares fit of the tracer results to an Arrhenius equation in the form D, = D,, exp (-&H,/RT) yielded + 3.3 D,, = k . 6 - lo-’ m2 s-l and AH, = 49.2f2 kJ mol-’. Within the limits of the present experiments AH, is independent of 7‘. The results of the pressure experiments are shown in fig. 5. coefficient can be expressed as l 8 The self-diffusionA . V . CHADWICK AND J . W . FORREST 2567 where c1 is the lattice parameter, v is a vibrational frequency (usmlly taken as the Debye frequency), /? is a geometric factor, and AGd is the Gibbs free energy for diffusion. The activation volume, AVd, is defined as = [a(AGd)/ap]T (4) and hence from eqn (3) The second term on the r.h.s.of eqn (5) is negligible and AVd is obtained from the slope of a plot of log D against P. The present results yielded AVd = (82+6) x m3 mol-1 at 301 K and AVd = (75k5) x m3 mol-1 at 281 K. Within the limits of this study AYd is independent of P and T. P/MN m-” FIG. 5.--Self-diffusion coeffcients for norbornylene plotted as a function of pressure : 0 and 9 are the radiotracer results in randomly oriented crystals at 301.2 and 281.0 K, respectively; x , n.m.r. data at 296 K. The lines are best fits to the points. DISCUSSION The available diffusion data for norbornane and norbornylene are collected in table 1. The present tracer results for norbornane are not sufficient in themselves to yield information on the mechanism of self-diffusion.The n.m.r. data are consistent with a monovacancy mechanism of self-diffusion in norbornane. The value of AHn/Ls is I .76 5 0.06 which is only slightly smaller than the theoretically predicted value for this mechanism. However, the prediction is an approximation based on an analogy with the rare gas crystals and the difference may not be significant. The strongest evidence for this mechanism being operative comes from the n.m.r. pressure data. Hard-sphere calculations for close-packed solids (f.c.c. and h.c.p.) predict an activation volume for vacancy formation equal to V, and the corresponding migration volume equal to -0.8 V,. Experimental determinations of AVd/Vm in systems where the vacancy mechanism has been confirmed by other techniques are considerably less than 1.8, e.g.0.7 to 0.95 in metals,19* 2o 0.72 to 1.3 in plastic crystals l2 and 1.2 in naphthalene.21 It is usually accepted that AVd/Vm < 1 is due to some inward relaxation of the molecules surrounding the vacancy and this would explain the result for norbornane. It would be unwise to evaluate AH, from the few data points in this study; however, they roughly parallel the temperature2568 SELF-DIFFUSION I N PLASTIC SOLIDS dependence of the n.m.r. data. The value off,/& for vacancy diffusion in a h.c.p. structure is not known but it would not be expected to be very different from that for a f.c.c. structure, i.e.ft/fn N 1.4. Thus the present estimate of 1.5k0.2 is extra support for monovacancy self-diffusion occurring in norbornane.TABLE 1 .-SELF-DIFFUSION PARAMETERS FOR h.c.p. NORBORNANE AND NORBORNYLENE The present norbornane reference norborny lene 1.53 14 1.22 3 1 f l 4 33+ 1 - - 1.49+ 0.1 1 1.76+_ 0.06 4 1.475 0.06 1.5k0.2 this work, 4 1.5k0.2 100.2 15 95.8 0.99+0.01 (299 K) 11 0.88s40.02 (296 K) 0.86+0.06 (301 K) 0.7840.05 (281 K)} reference 14 4 this work 4 this work, 4 15 10 this work tracer measurements for norbornylene yield AH&, as 1.49 & 0.1 1 which is reasonable evidence for a monovacancy mechanism even though it is smaller than the theoretically predicted value. Again this difference may be due to a weakness in the theoretical model. There is also the possibility that AHt has been under- estimated due to an undetected contribution to D, from pipe diffusion.It is difficult to eliminate this possibility although a comparison with other tracer studies of plastic crystals would suggest that it is unlikely. Unless norbornylene has a different bulk defect nature to other materials of this type then the present experiments should have minimized the effect of pipe-diffusion and AHt should not be significantly smaller than the true activation enthalpy for lattice diffusion. Similarly, the value of A Vd/ V,, -0.8, should be representative of lattice diffusion and, by analogy wiih other materials, support the view that self-diffusion in norbornylene involves a monovacancy which is slightly relaxed. The only directly comparable tracer study of a close-packed plastic crystal has been made for f.c.c. pivalic acid l2 where A Vd/ V, was found to be 1.3.For four f.c.c. plastic crystals, including pivalic acid, A Vd/ V, has been determined by the creep technique l2 and the values are in the range 1 .O to 1 .3. The fact that norbornylene lies outside this range may be a consequence of the difference between f.c.c. and h.c.p. structures although considerably more data for both lattice types would be needed to verify this point. In addition, it should be remembered that the measured AV, is not the sum of the local defect formation and migration volumes. It is these latter volumes that ought to be used in comparing various materials ; however, their evaluation from A Vd requires a knowledge of Poisson's ratio.Ig Since this is not available for the plastic crystals that have been studied it is not meaningful to try to interpret small variations of AV,/V,.For norbornylene there is a good agreement between AHt and AHn, the latter being obtained from data points over a much wider temperature range.4 The n.m.r. data show a region of apparently increased activation energy at the highest tempera- tures, as can be seen in fig. 4. This has been observed in other low ASf plastic crystals. It has been attributed to the use of the Torrey model to evaluate D, from the Tlp data and is not due to any change in the diffusion process.22 As with norbornane, the ratio of DJD, is 1.5+_0.2 and provides additional support for monovacancy self- diffusion. The agreement between the tracer and low pressure n.m.r. values of AV, for norbornylene is excellent. The n.m.r. study lo revealed a region of low AV, at pressures which would be beyond the scope of a tracer study.A .V . CHADWICK AND J . W. FORREST 2569 In summary, norbornane and norbornylene are two low ASf plastic crystals for which tracer and n.m.r. studies of diffusion yield concurrent results; in the case of norbornylene these include D, A H and AVd, which can be interpreted in terms of monovacancy self-diffusion. It may be significant that both these materials are relatively brittle.23 Differences between AH,, and AH, in some soft, low ASf, materials are now well-substantiated, notably in cyclohexane, * although differences in D and AVd from the two techniques are not large. There is still no complete explanation of this effect. Recent work 22 would suggest that the n.m.r.data does not monitor pipe-diffusion and the isotope-mass effect suggests that the lattice diffusion processes monitored by the tracer methods are the same for all plastic crystals irrespective of ASf. One major obstacle in resolving the discrepancy is the fact that the tracer measurements are limited to D, > lO-I5 m2 s-l and therefore cannot be used to investigate the possibly interesting regions at high P and low T. The conclusion of the present work is that the source of the discrepancy is not simply the numerical value of ASf and that the explanation needs to be sought in terms of the nature of the plastic crystalline state. One of us (J. W. F.) thanks the University of Kent for the award of a graduate assistantship. The authors acknowledge valuable discussions held with Drs.N. Boden, R. Folland and J. H. Strange. A. V. Chadwick and J. N. Sherwood, Point Defects in Solids, ed. J. H. Crawford, Jr. and L. M. Slifkin (Plenum Press, New York, 1975), vol. 2, p. 441. J. N. Sherwood, Surface and Defect Properties of Solids (Spec. Period. Rep., Chemical Society, London, 1973), vol. 2, p. 250. P. Bladon, N. C. Lockhart and J. N. Sherwood, Mol. Cry~t. Liq. Cryst., 1973, 19, 315. R. Folland, R. L. Jackson, J. H. Strange and A. V. Chadwick, J. Phys. Chem. Solids, 1973, 34, 1713. A. V. Chadwick, J. M. Chezeau, R. Folland, J. W. Forrest and J. H. Strange, J.C.S. Faraday I, 1975,71,1610. 6A. V. Chadwick and H. R. Glyde, Rare Gas Solids, ed. M. L. Klein and J. A. Venables (Academic Press, London, 1977), vol. 2, p. 1151. ’ N. C. Lockhart and J. N. Sherwood, Farday Symp. Chem. SOC., 1972,6,57. * E. M. Hampton and J. N. Sherwood, J.C.S. Faraday I, 1976,72,2398. P. Bladon, N. C. Lockhart and J. N. Sherwood, MoZ. Phys., 1971, 22, 365. lo R. Folland and J. H. Strange, J. Phys. C, 1972,5, L 50. l1 R. Folland, S. M. Ross and J. H. Strange, Mol. Phys., 1973, 26,27. l2 P. McKay and J. N. Sherwood, J.C.S. Faraduy I, 1975,71,2331. A. D. LeClaire, PhysicaZ Chemistry-un Advanced Treatise, ed. H. Eyring, D. Henderson and W. Jost (Academic Press, New York, 1970), vol. 10, p. 261. 14E. F. Westrum, Molecular Dynamics and Structure, ed. R. S . Carter and J. J. Rush (U.S. Department of Commerce, National Bureau of Standards, 1969), N.B.S. Special Publication 301, p. 459. l 5 R. L. Jackson and J. H. Strange, Acfa Cryst. B, 1972,28, 1645. l6 M. Y. Al-Shaker, A. V. Chadwick and J. W. Forrest, J. Labelled Compounds, 1975,11, 242. l7 J. N. Sherwood and D. J. White, Phil. Mag., 1967, 16,975. l 8 P. G. Shewmon, Diffusion in Solids (McGraw-Hill, New York, 1963). D. Lazarus and N. H. Nachtrieb, Solids under Pressure, ed. W. Paul and D. M. Warschauer (McGraw-Hill, New York, 1963), p. 43. 2o N. L. Peterson, Solid State Physics, ed. F. Seitz, D. Turnbull and H. Ehrenreich (Academic Press, New York, 1968), vol. 22, p. 409. 21 E. Hampton and J. N. Sherwood, J.C.S. Faraday I, 1975, 71, 1392. 22 N. Boden, J. Cohen and R. T. Squires, MoZ. Phys., 1976, 31, 1813. 23 R. M. Hooper and J. N. Sherwood, Surface and Defect Properties of Solids (Spec. Period. Rep., Chemical Society, London, 1977), vol. 6, p. 308. (PAPER 7/2275)

ISSN:0300-9599    

DOI:10.1039/F19787402562

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Infrared Study of CO Chemisorption on Zeolite and Alumina Supported Rhodium BY MICHEL PRIMET* Institut de Recherches sur la Catalyse, 79 Boulevard du 11 November 1918, 69626 Villeurbanne Cedex, France Received 5th January, 1978 The chemisorption properties of supported rhodium with respect to carbon monoxide are investi- gated by infrared speclrometry. Zeolite and alumina supported rhodium is obtained following hydrogen reduction. Highly dispersed rhodiuin is prepared on the zeolite support whatever the temperatme of reduction over the range 280-600°C. For rhodium supported on alumina, the dispersion is less good and decreases when the conditions of reduction become more drastic. For particles of larger diameter than 20& chemisorption of CO leads mainly to gem dicarbonyl species, the site of adsorption being one rhodium atom in the oxidized state.These gem dicarbonyl complexes are not formed at low temperature, but develop on warning up to room temperature. Dissociation of carbon monoxide is then assumed, CQ in excess adsorbs on the surface Rh-0 species to give gem dicarbonyl complexes, probably of Rhr, which interact with thc lattice in the case of zeolite support. Infrared spectroscopy of adsorbed species was extensively used to characterize various catalysts. In tbie case of srtpported metals, the specific adsorption, on the metal, of compounds like CO or NO has given useful information concerning factors frequently involved in catalytic process : effect of the support,l particle size effect,2 modifications of surface properties by alloying, etc.Unsupported metals, in the form of ribbons, foils, f i l m or monocrystals, were used in order to examine a wider spectral range or to remove the expected support effect. In the case of metals like Ni, Pt, Pd or Fe, the results obtained for supported or Unsupported systems are in good agreement both for bands positions and the number of bands. This analogy in behaviour is not observed for CO adsorption on rhodium :4-7 bands attributed to linear and bridged species are detected on unsup- ported rhodium,4* whereas gem dicarbonyl species are mainly observed for alumina or silica supported rhodium.6* This discrepancy between supported and unsupported rhodium catalysts prompted us to investigate CO adsorption on supported rhodium upon varying the nature of the support, the metal particle size and the temperature of the adsorption experiment.The use of zeolite as a support enables us to examine the spectral range below 1200 cm-l from which information concerning the zeolite structure and the effect of the reduction on the lattice are expected. EXPERIMENTAL The rhodium form of Y zeolite was prepared from a sample of NaY zeolite with a Si/AI ratio of 2.4 (Linde, S.K. 40). Cation exchange was carried out by stirring the NaY zeolite with an aqueous solution of [Rh(NH3)5C1]2+ complex. The sample was carefully washed with distilled water and dried at 323 K. Rhodium chemical analysis indicated that the sample contained 3.1 % wt of rhodium. Ammonia and chlorine coordinates were removed 2570M. PRIMET 2571 by a thermal treatment at 623 K in a flow of oxygen in accordance with previous E.S.C.A measurements. * Reduction was performed under hydrogen over the temperature range 473-573 K.Bare surfaces of rhodium were obtained by removal of adsorbed hydrogen under vacuum at 623 K. Electron microscope determinations, carried out on the samples, showed the presence of highly dispersed metal with a particle size between 4 and 8 w . Rh/A1203 catalysts were prepared by stirring the alumina (Aluminum Oxid P from Degussa) with an aqueous solution of rhodium trichloride. Water was eliminated by evaporation under vacuum, then the solid was dried at 393 K. Chemical analysis gives a Rh content of 3.68 % wt. Reduction was performed under a flow of hydrogen either at 773 K, or at 1273 I<.In both cases adsorbed hydrogen was removed at 773 K under vacuum. For 500°C hydrogen reduction (Rh/A1203 773, electron microscope measurements showed the presence of rhodium particles ; 85 % of them had a diameter between 5 and 20 A with some crystallites of 50A. For 1273 K hydrogen reduction (Rh/A1203 1273), 70 % of the metal was found in particles of 20-40A diameter, some crystallites in the range 100-140A were also detected, but no particles of diameter lower than lOA were observed. Samples were pressed to form thin pellets 18 m in diameter and weighing 30-40 mg for Rh/A12Q3 catalysts, or about 15 mg for RhNaY samples for experiments in the range 4000-1300cin-'. Very thin pellets of about 6-8 mg were necessary to investigate the 1000-500 tin-l range of the zeolite framework vibrations. The pellet was then introduced in a cell where the reduction treatments were performed in situ.Depending on the type of cell used, infrared spectra were recorded either at 300 K or at a lower temperature using a Fourier Transform spectrometer (Digilab F.T.S. 14) with a resolution of 4 crn-l. Every spectrum was the sum of 200 scans and was stored in the disk memory. The initial spectrum of the sample was used as a reference for the following adsorptions in order to remove the bands due to the background of the support (especially in the case of zeolite). RESULTS AND DISCUSSION ADSORPTIONS OF co ON BARE SURFACES Carbon monoxide was admitted under a pressure of 0.3 kN m-2 onto RhNaY samples. Whatever the temperature of hydrogen reduction in the range 473-873 K, the infrared spectra showed the same features after CO adsorption.A typical spectrum is presented in fig. 1. Two main bands around 2100 and 2040 cm-I developed during the contact time under CO. The intensities of both bands increased with time, but the intensity ratios remained constant. These two absorptions were split into two components 21 15-2100 and 2050-2025 cm-l. In addition, a broad band was observed between 1920 and 1750 cm-l; it is possible to resolve it into bands located at 1870-1830 and 1760 cm-'. In the range 1000-500 cm-l, a band at 910 cm-I was detected and its intensity grew with the time of contact under CO. After oxygen treatment at room temperature, the bands between 1920 and 1750 cm-1 completely disappeared whereas the main bands near 2100 and 2040 cm-' and that at 910 cm-I were unaffected.Nevertheless, the 2100 and 2040 cm-l bands were now better resolved, which means that the species vibrating between 2100 and 2040 cm-1 was removed by room temperature oxygen treatment. Indeed, the comparison of spectra recorded before and after reaction with oxygen revealed the disappearance of a band close to 2070 crn-I. Carbon monoxide adsorbed onto Rh/A1,03 1273 gave a single strong band at 2065 cm-I with a small shoulder at 2100 cm-I and an absorption at 1935 cm-I (fig. 1). By oxygen treatment at 300 K, the two low frequency bands disappeared and the spectrum showed new bands at 2128-2100-2038 cm-I in addition to absorption due to carbonated species bound to the support.Adsorption of CO onto Rh/Al,O3 773 gave an intermediate state between that on the two preceding samples, i.e., appearance of the 2110-2062-2038 cm-' bands2572 co ON SUPPORTED RHODIUM with comparable intensities and of a broad band at 1875 cm-l (fig. 1). As with the previous solids, the 21 10 and 2038 cm-l bands did not disappear by oxygen treatment at 300 K in contrast to the 2062 and 1875 cm-l bands. 2110 - 2200 m 1800 (a wavenumber/cm- 2200 2000 (a FIG. 1.--Infrared spectra of CO adsorbed at 300 K on supported rhodium (A) RhNaY sample treated under oxygen, then under hydrogen and finally under vacuum at 623 K. (B) Rh/A1203 sample treated under flowing hydrogen at 773 K and desorbed at the same temperature. (C) Rh/A1203 sample treated under flowing hydrogen at 1273 K and desorbed at 773 K.Spectra (a) correspond to CO irreversibly adsorbed at 300 K. Spectra (b) were obtained after treating the previous samples with 15 kN m-2 oxygen at 300 K. In comparison with the results obtained for CO adsorption on rhodium we attribute the 2060-2070 cm-l band to a linear Rh-CO species. The bands situated in the range 1920-1760 cm-1 are assigned to multicentred species. The band at 2128 cm-l which developed after oxygen treatment at 300 K is due to Rh / O ‘co surface complexes. According to the pioneer work of Garland et aZ.,6 the two bands near 2100 and 2040 cm-l are attributed respectively to the asymmetric and symmetric modes of a gem dicarbonyl complex of metallic rhodium Rh co <,, From these assignments, it appears that particles of diameter smaller than lOA are highly suitable for the formation of Rh(CO)z species, since this surface complex is observed mainly on RhNaY samples.When the mean particle size is >20 A (Rh/Alz03 1273), this complex is no longer observed : only linear and multicentred forms are detected.M. PRIMET 2573 Without further experiments, it is difficult to find an explanation for the splitting of the bands due to the gem dicarbonyl species in RhNaY samples. It is probable that the specific properties of rhodium in zeolites must be considered since this phenomenon is not encountered for Rh/Al,03 catalysts. I I I I I 2200 21 00 2000 1900 1800 wavenumber /cm-' FIG. 2.-Influence of preadsorbed oxygen or hydrogen on the spectrum of CO irreversibly adsorbed at 300K on Rh/A1203 773 sample.(a) Bare sample, (b) sample precovered by oxygen at room temperature, (c) sample precovered by oxygen at 473 K, (d) sample precovered by hydrogen at 300 K.2574 co ON SUPPORTED RHODIUM According to the previous results, the i.r. spectra of CO adsorbed on rhodium should give a useful clue concerning the dispersion of the metal : the appearance of the doublet 2100-2040 cm-l for highly dispersed rhodium and formation of the 2070-2060 cm-1 band for poor dispersions, i.e., high temperatures of reduction or rhodium films. The assignment of the 2100-2040 cm-I doublet to two CO molecules adsorbed on the same site is in agreement with the experimental data, except in the zerovalent state where the site is identified with one metal atom. We bear in mind the following facts: (i) the doublet is not affected by treating the sample under oxygen at 300 K, which is unusual for carbon monoxide adsorbed on a metal, (ii) hydrogen treatment at room temperature led to a strong decrease in the high frequency band (2100 cm-I), the low frequency one (2040 cm-l) being overlapped by the shift in the 2070-2060 cm-1 band towards lower wavenumbers ; (iii) the frequen- cies observed for this gem dicarbonyl species are very close to the frequencies of [Rhr (@O),Cl], complex in the solid state 2089-2033 cm-1 lo or adsorbed on a NaY zeolite 21 12-2095 and 2050-2030 cm-l." Considering these observations, it seems unlikely that a metallic rhodium atom could be the adsorption site involved in the formation of the gem dicarbonyl species.ADSORPTION OF co O N HYDROGEN OR OXYGEN PRECOVERED SURFACES Rh/A1,03 773 sample was chosen as reference in the investigation of the effect of preadsorbed oxygen or hydrogen on the metal. If the surface was precovered by oxygen at room temperature, CO adsorption gave intense bands due to the gem dicarbonyl species, whereas the bands due to linear and multicentred forms decreased (fig. 2). If the rhodium surface was oxidized by heating under O2 at 473 K, the subsequent adsorption of CO at 300K produced only the bands due to the gem dicarbonyl species, i.e., the surface of the catalyst was fully oxidized after oxygen treatment at 473 M. On the other hand, if CO was adsorbed on Rh/A1203 which had been first exposed to hydrogen, then evacuated at 300 M, CO adsorption gave a spectrum which showed weaker bands due to gem dicarbonyl species (fig. 2).In addition, a cycle of 0,-CO treatments was performed at room temperature on the Rh/Al,03 1273 sample. Every adsorption of oxygen led to a spectrum of adsorbed CO with an increase of the intensity of the doublet close to 2100-2040 cm-l. ft appears that the formation of the doublet at around 2100 and 2040 cm-l by CO adsorption must be correlated with the presence of oxygenated species bound to the metal. Indeed, the higher the amount of chemisorbed oxygen, the more dominant the doublet in the infrared spectrum. In contrast, the presence of preadsorbed hydrogen reduces the extent of rhodium oxidation. ADSORPTION OF CARBON DIOXIDE The chemisorption of carbon monoxide on supported rhodium was then accom- panied by an oxidation process for highly dispersed metal.Two types of reaction can be taken into account to explain this result : (i) Disproportionation of CO followed by dissociation of CO, : 2 c o 3 co, + C(S) co, -+ C(s) +20(s) co + C(s)+O(s). (iif dissociation of CO 111 bcpth cases, carbon monoxide in excess was adsorbed on the oxygenated surfaceM. PRIMET 2575 species to give the gem dicarbonyl complex. In the hypothesis of C 0 2 formation as an intermediate reactant, adsorption of CO, must be favoured in the case where the widest doublet was observed, i.e., that with the zeolite support. At 300 K, C 0 2 adsorption on RhNaY sample did not give v CO bands in the i.r. spectrum. Under the same conditions with Rh/A1203 773, bands were observed at 2025 and 1860 cm-l ; they disappeared on oxygen treatment at 300K and C02 then dissociated on the Rh/A1,03 catalyst.The frequencies observed were lower than those found for GO adsorption on the same solid; weak CO coverage can be explained by the position of these bands. If the Rh/AI2O3 sample was precovered with O2 at 300K, the dissociation of C02 did not occur. Thus, the dismutation of CO must be ruled out, since C 0 2 is not dissociated on ihe catalysts which produced the highest content of oxidized form after CO adsorption. LOW TEMPERATURE EXPERIMENTS Fischer-Tropsch reactions involved the dissociation of carbon monoxide on metals and a carbide intermediate was postulated. This dissociation of CO occurred during the methanation of CO on nickel catalysts ;12* l3 it was found that alloying nickel with copper strongly reduced the CO dissociation and in the same way the rate of methanation.14 On nickel, we have found that the CO dissociation occurred via a Ni4C0 intermediate which led to Ni3C and NiO surface species.12 For rhodium films, CO adsorption is considered as non-dissociative. 1.r.measurements by transmission 4 9 or by reflection absorption confirm this statement. When the conditions of absorption became more severe, CO dissociation was observed. Marbrow and Lambert l6 have shown that CO was adsorbed in a single state at 300 K on Rh (1 lo), but under electron impact carbon monoxide was dissociated, i.e., surface carbon was shown to inhibit CO chemisorption, whereas surface oxygen fed to the formation of a new more tightly bound form of CO.Hydrogenation of CO and C 0 2 between 523 and 623 K over polycrystalline foils of rhodium confirmed the decomposition process, since under these reaction conditions, the surface is covered with a catalytically active carbonaceous deposit while some oxygen is located below the surface. Dissociation of CO on highly dispersed Rh seems to be the only way to account for our experimental data. This reaction must be temperature dependent and the nature of adsorbed species is expected to change drastically with the adsorption temperature. In order to investigate modification of the surface species with the temperature, CO adsorption was performed below 300K on Rh/Al,03 773 and RhNaY samples. Adsorption of CO at 77 K under a pressure of 0.27 kN rn-, on Rh/A1203 gave only the bands at 2075 and 1920-1870 cm-l (fig.3) assigned to linear and multicentred forms, respectively. In addition, small bands at 2190 and 2160 cm-l were detected and attributed to CO interacting with the support. When the temperature of adsorption was increased, bands at 2095 and 2028 cm-l developed and reached their maximum intensities after one hour contact at 300 K. The resulting spectrum was identical to that obtained by direct adsorption of CO at room temperature. If the temperature of adsorption was raised to 350K, the 2095 and 2028cm-l bands increased. The doublet attributed to the gem dicarbonyl species is not observed at tempera- tures below 193 K. Its progressive formation on increasing the temperature of adsorption must be correlated with dissociation of CO on rhodium followctl by adsorption of CO on the oxidized surface groups.2576 co ON SUPPORTED RHODIUM -1 I 2190 2160 --la--- > 10% LILlq4q 2095 I I I I I 2200 ZOO0 1800 .wavenurnber/cm-l sample. Carbon FIG. 3.-1.r. spectra of CO adsorbed at low temperature on Rh/A1203 773 monoxide was introduced under a pressure of 0.27 kN m-2 at 77 K and then the temperature was raised to (a) = 103 K, (b) = 203 K, (c) = 300 K, (d) = 353 K. A similar experiment was carried out on very thin pellets of RhNaY catalyst. CO adsorption under a pressure of 0.27 kN m-2 gave, between 77 and 143 K, a spectrum which had two main bands at 2040 and 1910 cm-1 ; CO adsorption on Naf ions l 8 was also detected by a peak at 2170 cm-' (fig.4). Thus, the spectrum observed under CO at 143 K is quite different from the spectrum obtained at 300 KM . PRIMET 2577 and particularly the doublet which is split for the RhNaY sample is not detected. Nevertheless this spectrum is difficult to explain. Indeed, the frequencies are slightly different from those mentioned for U/Al2O3 at 143 K (2075-1920-1870 cm-l). The differences noticed can tentatively be explained by the presence of peculiar sites in the small particles of rhodium in zeolites or by the formation of rhodium carbonyl complexes, as already evidenced by Ozin et aZ.19 By increasing the temperature above 193 K, the 21 15-2098 and 2048-2020 cm-1 bands developed similarly. At the same time, a band appeared at 910 cm-l and grew as the temperature increased.1898 1830 I 2041 I I I I 2200 2100 2ooo 1900 1800 I 910 wavenumber /cm-' FIG. 4.-1.r. spectra of CO adsorbed at low temperature on RhNaY sample reduced at 773 K. Carbon monoxide was introduced under a pressure of 0.27 k N m-2 at 77 K, and then the temperature was raised to (a) = 143, (b) = 153, (c) = 228, (d) = 300 K. As for the Rh/A1,03 sample, the appearance of the gem dicarbonyl complex must be associated with the dissociation of carbon monoxide. Hindered at low temperature, this process is achieved to a large extent at 300 K. In a previous study * of the properties of RhrJr ions in NaY zeolite, it was shown that the reduction of Rh"' ions is performed at 300 K in presence of carbon monoxide and water. E.S.C.A. measurements indicate that rhodium is in the state Rh' after reaction.The doublet due to Rhr species was found at the same position as it was for CO adsorption on supported rhodium zeolite. This analogy strongly co <co2578 co ON SUPPORTED RHODIUM suggests that dissociation of CO on RhNaY leads to rhodium in the plus one oxidation state which is able to chemisorb two CO molecules to form a gem dicarbonyl complex. Indeed, recent work performed here by U.V. reflectance spectrometry has shown that chemisorption of CO at room temperature on reduced RhNaY samples leads to the formation of a large number of planar Rh' complexes.20 The dissociation of CO on supported rhodium is enhanced for the highly dispersed metallic phase, whereas this process is strongly limited for poor dispersions. Metal surface areas were often measured by chemisorption of H2, O2 or CO. In the latter case, the stoichiometry of the adsorption must depend strongly on the particle size of rhodium.Actually, Wanke et found an increase of the CO/Rh ratio with decrease in the particle size. INTERACTION WITH ZEOLITE LATTICE In the case of the RhNaY sample, a band at 910 cm-I appears by adsorption of CO. The evolution of its intensity follows that of the doublet near 2100 and 2040 cm-l. Such a band has been already observed at 900 cm-l after removal of the NH3 ligands present in a NaY zeolite exchanged with palladium or rhodium ammine complexes.22 This band disappeared during hydrogen reduction at 150°C.22 For Cu" exchanged zeolite, a band at 900 cm-I was also created by removal of water.22 TABLE 1 .-ZEOLITE FRAMEWORK INFRARED ASSIGNMENTS 2J* 26 internal tetrahedra (structure insensitive) external linkages (structure sensitive) asym.stretch 1250-950 cm-l asym. stretch 1050-1150 cm-l sym. stretch 720-650 cm-l sym. stretch 750-820 cm-1 T-0 bend 420-500 cm-' double ring 650-500 cm-' pore opening 300-420 cm-l The band at around 900cm-' is not very sensitive to the nature of the cation Mn+ (Rh, Pd, Cu) bound to the zeolite lattice. Thus, it is difficult to attribute it to a Mn+-oxygen vibration. It appears more realistic to consider a local deformation of the T (T = Si or Al) oxygen bonds of the zeolite due to interaction of the oxygen atom with the cation. Indeed, X-ray measurements have shown that interaction of cations such as Pd" 23 or CoI1 24 with the lattice of a NaY zeolite lead to an important lengthening of the T-0 bond which distorts the tetrahedron.This increase causes a weakening of the force constant of the T-0 bond which displaces the T-0 vibration towards lower frequencies. The bands in the vibrational spectra of the zeolite framework have been assigned by Flanigen et aZ.25* 26 Some of these assign- ments are summarized in the table 1. In Y-zeolites, internal tetrahedron TO4 vibrations include the two most intense bands : the strongest at around 1000 cm-1 (T-0-T asymmetric stretching mode) and the other of medium intensity at around 469 cm-' (0-T-0 bending mode). From infrared data and unit cell parameters of synthetic alumino silicate sodalites, Taylor et al.27 observed a linear relationship between the T-0 distance and the wavenumber of the v,, mode : d(T-O)(A) = 2.195- (5.263 x 10-4)(~,, T-0-T).By applying this relationship to a NaY zeolite for v,, (T-0-T) = 1035 cm-', we find a T-0 distance of 1.65 A which is close to the mean value observed by X-ray diffra~tion.~~ If we consider the frequency detected for Pd"Y zeolite, we find A(T-0) = 1.72A. PdI1 ions in Y-zeolite are mainly located on the SIP site, which means that they are interacting with Oc3) oxide ions ; X-ray diffraction measurementsM. PRIMET 2579 performed on such systems have given the value 1.720 A for the T-O,,, distance, which is identical to that deduced from infrared data. Thus, the cation lattice interaction with lengthening of the T-0 bond from 1.65 to 1.72 A can explain the shift of the 1035cm-l band towards 900cm-l.The appearance of a band near 900cm-’ must be connected with the presence of charged species interacting with the zeolite framework. Since the increase in the (21 12-2098)-(2048-2020) cm-1 bands is associated with the increase in the 910 cm-l band, we can conclude that CO adsorption on RhNaY sample leads to the formation of oxidized species of rhodium interacting with the lattice. CONCLUSION It appears that adsorption of CO on rhodium is more complex than on supported platinum or nickel. For Rh, the nature of the adsorbed CO is strongly dependent on the metal particle size. For large particles, CO is mainly adsorbed in a linear Rh-CO complex and the CO/Rh ratio is close to 1. When the particle size decreases, gem dicarbonyl complexes are formed and the smaller the diameter, the greater the amount of gem dicarbonyl species ; so the CO/Rh ratio is > 1.Not only the number of CO molecules linked to one rhodium atom, but also the state of rhodium after CO adsorption is particle size dependent. On large particles (films or high reduction temperatures), adsorption of CO is not dissociative. When the particle size decreases below lOA, carbon monoxide is dissociated. This process is strongly hindered below 173 K and is mainly observed at room temperature for zeolite supported rhodium. Subsequent adsorption of CO leads to the formation of gem dicarbonyl species of oxidized rhodium probably in the Rh’ state. In such a case, the RhNaY samples would be very efficient catalysts for the hydrogenation of CO and CO, since on these solids the dissociation of CO is already achieved to a great extent at room temperature.Thanks are due to Mrs. I. Mutin and C. Leclerq for taking electron micrographs and to Mr. Marc Dufaux for his help with the infrared experiments. The author is grateful to Dr. G. Coudurier, Dr. G. Naccache, Dr. M. V. Mathieu and Dr. B. Imelik for helpful discussions. F. Figueras, R. Gomez and M. Primet, Adv. Chem. Ser. Molecular Sieves, 1973, 121, 480. ’ M. Primet, J. M. Basset, E. Garbowski and M. V. Mathieu, J. Amer. Chem. Soc., 1975,97,3655. J. A. Dalmon, M. Primet, G. A. Martin and B. Imelik, Surface Sci., 1975,50,95 ; M. Primet, M. V. Mathieu and W. M. H. Sachtler, J. Catalysis, 1976, 44, 324. C. W. Garland, R. C. Lord and P. F. Troiano, J. Phys. Chem., 1965,69, 1188. R. Queau and R. Poilblanc, J. Catalysis, 1972, 27, 200. A. C. Yang and C. W. Garland, J. Phys. Chem., 1957,61, 1504. ’ H. Arai and H. Tominaga, J. Catalysis, 1976,43, 131. * M. Primet, J. C. Vedrine and C. Naccache, J. Mol. Catalysis, in press. M. Primet, to be published. C. W. Garland and R. J. Wilt, J. Chem. Phys., 1962,36, 1094. l 1 M. Primet, unpublished results. l 2 G. A. Martin, M. Primet and J. A. Dalmon, J. Catalysis, in press. l 3 M. Araki and V. Bonec, J. Catalysis, 1976,44,439. l4 W. L. Van Dijk, J. A. Groenewegen and V. Ponec, J. Catalysis, 1976, 45,277. M. G. Wells, N. W. Cant and R. G . Greenler, Surface Sci., 1977, 67, 541. l6 R. A. Marbrow and R. M. Lambert, Surface Sci., 1977, 67,489. l7 B. A. Sexton and G. A. Somorjai, J. Catalysis, 1977, 46, 167. ’* C. L. Angel1 and P. C. Schaffer, J. Phys. Chem., 1966,70, 1413. l9 L. A. Hanlan and G. A. Ozin, J. Amer. Chenz. SOC., 1974,96,6324.2580 co ON SUPPORTED RHODIUM 2o H. Praliaud, personal communication. 21 S. E. Wanke and N. A. Dougharty, J. Catalysis, 1972, 24, 367. 22 M. Primet, unpublished results. 23 P. Gallezot and B. Imelik, Adv. Chem. Ser. Molecular Sieves, 1973, 121, 66. 24 P. Gallezot and B. Imelik, J. Chim. phys., 1974, 71, 155. 2 5 E. M. Flanigen, H. Khatami and H. A. Szymanski, A h . Chem. Ser., 1971, 101, 201. 26 E. M. Flanigen, in Zeolite Chemistry and Catalysis, ed. J . A. Rabo, A.C.S. Monograph, 1976, 27 C. M. B. Henderson and D. Taylor, Spectrochim. Actu A, 1977, 33, 283. 171, 80. (PAPER 8/018)

ISSN:0300-9599    

DOI:10.1039/F19787402570

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Reviews of Books Hydrogen: Its Technology and Implications. Vol. &Hydrogen Properties. By R. D. MCCARTY. Ed, K. E. Cox and K. D. WILLIAMSON. (C.R.C. Press, Cleveland, Ohio, 1975). Pp. 321. Price $46.00. In recent years we have seen the publication of extensive, accurate thermodynamic tables of many of the simpler substances, especially the permanent gases. These tables of the proprties of hydrogen are based on ones prepared by R. D. McCarty under the auspices of the National Aeronautics and Space Administration and they fit into this category. In view of the author’s experience and his use of the most up-to-date procedures in making his correlations, the result must be considered as representing the present state of the art as far as the generation of tables of properties of this sub- stance is concerned.Properties are tabulated for both normal and parahydrogen over the whole fluid range up to 1 kbar and 3000K in SI units. They include volume, enthalpy, entropy, C,, Cv, speed of sound, thermal conductivity, viscosity and dielectric constant. Various derivatives are listed also although the headings to some of these are incredibly cryptic. One labelled “ ISOTHERM DERIVATIVE CU M-MPA/KG ” in most of the tables but labelled “ Isotherm derivative, m3/MPA/kg ” in the saturation table was particularly confusing. Nowhere in the text is the term explained and from its proximity to the tabulation of volume one might expect it to be -(aV/W)~with units m3/(MPa kg), often erroneously written m3/MPa/kg. In fact, the numbers tabulated make sense only if one interprets them as ( W l a p ) ~ , the units in the saturation table being the ones that are wrong.Unfortunately, this ambiguity is not an isolated case. There are graphs with insufficient informa- tion on them (and no explanation in the text), properties tabulated twice but with two different values. Indeed, the whole presentation has the appearance one would expect if one person were independently to write the text, a second collect the graphs and a third draw up the tables, and at the price asked, one is justified in expecting better than this. As a final blow, this reviewer’s copy was defective: one whole page was missing and in its place was printed a duplicate copy of another. Clearly this expensive book is going to have a very limited sale and one cannot consider it good value for money, but if one’s work involves frequent access to tables of the properties of hydrogen, and investigators into the so-called “ hydrogen economy ” might fit into this category, then these tables will be as invaluable as they are unique.G. SAVILLE Received 5th October, 1977 Chemistry and Physics of Solid Surfaces. Ed. RALPH VANSELOW and S. Y. YONG. (CRC Press, Cleveland, Ohio, U.S.A., 1977). Pp. xi+ 386. Price $20.80. This book contains a set of review articles of lectures delivered at the second (biennial) Inter- national Summer Institute in Surface Science, 1975, University of Winsconsin, Milwaukee. The aim is to provide a forum to discuss and evaluate the latest developments by leading research workers. An outstanding contribution by (the late) Erwin W.Miiller on all aspects of ion microscopy is accompanied by four articles on surface analysis by electron and ion beams and also photoemission and magnetic resonance techniques; there are theoretical papers on the quantum physics and chemistry of surfaces, on field-emission energy distribution of electrons from clean metals, and a survey of modern chemisorption theory; various aspects of physisorption, surface (and bulk) diffusion of metal atoms, adsorption and crystal growth, and molecular beam epitaxy, are critically examined in other reviews. Finally, technological applications concerned with surface effects in the electronic industry and in thermonuclear fission, together with the interaction of gases with solid surfaces in catalysis, are the subject of three other articles. Most of the writers are foremost scientists attached to American research establishments or universities, but four reviews by well-known English, French and German specialists are also included. The reviewer was present at both the 1973 and 1977 ISISS meetings and can commend the excellence of the contributions and the vitality of the Discussion (unfortunately not included in the printed volume), and would recommend the present volume to graduates and mature postgraduates pursuing advanced courses and research in surface chemistry concerned with the gas-solid interface. UnhappiIy the reviewer’s copy had the unique feature that the binding was upside down and back to front ! F. C . TOMPKINS Received 30th November, 1977 2581

ISSN:0300-9599    

DOI:10.1039/F19787402581

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Hydration, Dehydrative Counter-ion Binding and Helix Formation of Charged Poly(a-amino acid)s in Aqueous Alcohol as Revealed by a Preferential Binding Study BY TOSHIO MORI, JIRO KOMIYAMA" AND TOSHIRO IIJIMA Department of Polymer Science, Tokyo Institute of Technology, Ookayama, Meguro-ku, Tokyo 152, Japan Received 8th September, 1977 The preferential binding behaviour of one of the components of aqueous alcohol solutions to sodium poly(or-L-glutamate) (PLGNa) and poly(a-L-lysine HBr) (PLLHBr) at 25°C has been studied by differential refractometry. Theaqueous mixtures have 1-propanol(1 -PrOH), 2-propanol(2-PrOH), t-butanol (t-BuOH) and 1,2ethanediol (EG) as their second component. The electrical conductance and the or-helix content of these polymers were also measured to interpret the preferential binding behaviour.Aqueous monohydric alcohols show a characteristic dependence on the alcohol concentration : the initial trend of the increasingly negative binding of the alcohols is reversed at certain alcohol concentrations, giving clear breaks. The results are interpreted invoking an operational model of these polymers, in which polar and less polar portions undergo solvation separately. PLGNa residue is hydrated by 20 water molecules, of which 12 are taken up by the polar portion and the rest by the less polar portion. PLLHBr residue is hydrated by 27 water molecules, of which 15 are taken up by the polar portion. With increase in the alcohol concentration, the hydration of the polar portions of these polymers is lost upon counter-ion binding, which is followed by the helix formation of the polymers.5 water molecules are released from around the less polar portions upon formation of the helix. In a previous paper, we reported the preferential binding behaviour of one of the components of aqueous organic solvents or organic solvent binary mixtures to PLLHBr.l The variation in the preferential binding parameters for the nine binary solvent mixtures was examined over different composition ranges and was interpreted by invoking an operational model of PLLHBr, in which polar and less polar portions of the polymer undergo solvation separately. Definition of these portions is rather diffuse, but we assume that the charge conveying group of the polymer and the counter-ion are included in the polar portion, and the main chain groups and the methylene groups of the side chain are included in the less polar portion.Of particular interest among the results is the preferential binding of 2-PrOH from the aqueous mixtures : the binding is increasingly negative with concentration until this trend is reversed from a break at around a mole fraction of 0.30 of the alcohol in the solvent mixture. We have shown that the negative preferential binding before the break is explained if the polymer residue is exclusively hydrated by 27 water molecules, against the composition change in the solvent mixture. We have speculated that the reverse trend may be associated with dehydration due to the counter-ion binding as well as to the a-helix formation of the polymer. Such extensive hydration of PLLHBr should be established firmly.Extended research on the preferential binding behaviour to another charged poly(a-amino acid) from aqueous alcohol solutions may give some information about the hydration as influenced by the changes in the states of the relevant polymers. In the present study, the preferential binding 1-82 25832584 PREFERENTIAL BINDING TO CHARGED POLY (AMINO ACID)S measurements for PLLHBr were supplemented for the aqueous t-BuOH system and the binding to PLGNa from aqueous I-PrOH, 2-PrOH and t-BuOH was investigated. For comparison, measurements for PLGNa in aqueous EG and dimethyl sulphoxide (DMSO) were included in the study. EXPERIMENTAL MATERIALS PLGNa was obtained by saponification of poly(y-methyl-a-L-glutamate) which was prepared according to Oya et aL3 The polymer, after being washed with acetone and dried, was dissolved in water and recovered by precipitation in 2-PrOH at 5°C.The impurities and low molecular weight fraction of the polymer were removed by electrodialysis of the aqueous solution for 50 h, followed by ultrafiltration through a Diafilter G 10 T (Bio- engineering) which is impermeable to polymers of molecular weight >lo4. The polymer was recovered by freeze drying and stored in a desiccator until use. The molecular weight as determined by the viscosity measurements was 1.1 x lo5. The preparation and purifica- tion of PLLHBr was described in the previous paper. The organic solvents (analytical grade) obtained from Wako Pure Chemicals were dried by anhydrous Na2S04 and fractionally distilled using an efficient column under the appropriate pressures.The solvents were stored in desiccators ; their densities agreed within & 5 x g ~ m - ~ with the literature values in table 1, in which other physical constants of the solvents are listed. Laboratory supply water, which was distilled after deionization, was used ; its conductance was (1-2) x Q-' cm-l. TABLE 1 .-PROPERTIES OF SOLVENTS HzO 1-PrOH 2-PrOH t-BuOH EG dielectric constant, (25°C) 78.4 20.3 19.9 12.5 37.7 dipole moment (25°C) 1.85b 3.09 1.66 1.66 2.28 (20°C) (30°C) (3OoC) (2OoC) refractive index, (25°C) 1.3325 1.3837 1.3752 1.3851 1.4306 density/g ~ m - - ~ (25°C) 0.9970 0.7998 0.7813 0.7812 1.1100 a Taken from ref. (5) unless specified; b ref. (6); C ref.(7). DMSO 46.7 3.89 1.4773 1.0958 APPARATUS AND METHODS The apparatus and procedure of equilibrium dialysis have been described e1sewhere.l The PLGNa concentration in the dialysis equilibrium was determined by adding an aliquot of 1 /200 residue mol dm-3 aqueous glycol chitosan solution, followed by colloidal titration with 1 I400 residue mol dm-3 aqueous potassium poly(vinylsulphonate), using toluidine blue as indicator.8 Permeation of PLGNa through the dialysis bag (Visking) was found to be <1 % of the total polymer. Circular dichroism measurements were performed using a Jasco J 20 automatic spectro- polarimeter at 250°C. The instruments were calibrated with (+ )-10-camphorsulphonic acid. The helix content, xH, of PLGNa and PLLHBr in the aqueous alcohol solutions was determined by measuring the variation of the residue ellipticity of the 1 .2 0 ~ residue mol dm-3 solutions at 222 nm, [6]222/deg cm2 dmol-l and using the r e l a t i ~ n , ~ Conductance measurements were made with a Yokogawa Universal Bridge BV-Z-103B at 1 kHz at 25.00"C. The cell constant was determined before each set of experiments using standard aqueous potassium chloride solution. The specific conductances were determined for solutions of both polymers whose concentration was fixed at 1.65 x residue mol dm-3, and therefore the equivalent conductances, A, were calculated on the residue mole bases.T. MORI, J . KOMIYAMA AND T. IIJIMA 2585 Corrections due to the solvent conductances were applied to all the conductance measure- ments and these amounted to <10 % of the conductances of the solutions in most cases.The viscosities of aqueous EG,l0s l1 l-PrOH,lO* l1 2-PrOH l2 and t-BuOH l3 were taken from 1iterat~re.l~ PREFERENTIAL BINDING PARAMETER The preferential binding parameter (in mole alcohol per residue mol polymer) of an alcohol from the aqueous mixture to polymer at 25°C was determined according to eqn (2) l5 Here subscripts 1, 2 and 3 refer to water, polymer and alcohol, respectively. n is the refractive index, rn is the molarity of the respective component in water and C the concentra- tion of solution in g ~ r n - ~ . Vis the partial specific volume, ,u is the chemical potential, A4 is the molecular weight (per residue for the polymers) and the superscript indicates infinite dilution of the polymer.Partial specific volumes were calculated graphically from specific volume against concentration plots. The specific volume was measured by an Anton Parr precision density meter DMA 02C, at 25.00"C. Though the preferential binding parameters have been expressed by the above quantity in this report in conformity with the previous one, the values were found to be negative irrespective of the alcohol species over the concentration ranges investigated. The following relation may be used to obtain (aml/dmz)~,,p3, the preferential hydration parameter, from the values reported here l5 The refractive index difference was measured at 25.00"C by a Carl Zeiss Jena laboratory interferometer. (dn/aC2)p1,f13 and (an/aC,),, were measured for the polymer solutions in the concentration range (1-5) x g cm-3 when the alcohol volume of the solvent mixture was <60 %.Greater than this alcohol content, the polymers would not dissolve in the solvent mixture, giving the solution a slightly opaque appearance. The solution was prepared therefore initially by dissolving the polymer in water and then by adding alcohol gravimetrically to the solution with care to attain sufficient accuracy. Measurements of the refractive index increments were made for the duplicate polymer solutions in the concentration range (0.5-2) x g ~m--~. RESULTS AND DISCUSSION the systems PLGNa in aqueous 1-PrOH, 2-PrOH, t-BuOH, EG and DMSO are shown in table 2. The values for the system PLLHBr in aqueous t-BuOH are listed in table 3. The preferential binding parameters for PLGNa in aqueous organic solvents are plotted against x3, the mole fraction of the organic solvent component in the solvent mixture, in fig. 1.The plots of the parameters for PLLHBr in aqueous t-BuOH are shown in fig. 2, together with those in aqueous 2-PrOH reported previously. Fig. 1 shows that the preferential binding parameter of EG to PLGNa increases with EG concentration, and after passing a small maximum at x3 N 0.1, it is inverted by a change to preferential hydration at x3 N 0.25. The preferential binding of DMSO to PLGNa is found to be increasingly negative with the content, though the extent is moderate. In this case, the polymer precipitated if x3 exceeded 0.28. Fig. 1 shows the remarkably strong negative dependence of the preferential binding parameters of the three monohydric alcohols to PLGNa.Before the breaks which appear at x3 N 0.27, the points for each alcohol are found close together giving a decreasing trend, while after the breaks, they give different increasing trends. The values of C3, J!3, (an/ac2)fl,,p3, ( a n / a c 2 ) m 3 , (anlacs)m, and (am,lam2),",,,3 for2586 PREFERENTIAL BINDING TO CHARGED POLY (AMINO AC1D)S TABLE 2.PREFERENTIA.L BINDING OF ORGANIC SOLVENTS TO PLGNa FROM AQUEOUS MIXTURES solvent (3) vol. % x3 10 20 30 40 60 70 75 85 10 20 30 40 50 60 70 75 10 20 30 50 60 70 75 80 10 20 30 40 50 60 70 80 90 10 20 30 40 50 60 0.026 0.057 0.093 0.138 0.265 0.359 0.41 9 0.575 0.025 0.055 0.091 0.135 0.190 0.260 0.354 0.41 6 0.020 0.043 0.066 0.158 0.219 0.304 0.359 0.429 0.034 0.074 0.121 0.176 0.241 0.326 0.423 0.561 0.739 0.028 0.060 0.097 0.143 0.198 0.272 C3/g cm-3 0.08 16 0.1626 0.2446 0.3266 0.4894 0.5696 0.6095 0.6878 0.0789 0.1585 0.2407 0.3224 0.4030 0.4764 0.561 1 0.6015 0.0772 0.1535 0.2152 0.3977 0.4775 0.5559 0.5949 0.6343 0.1095 0.2223 0.3335 0.4461 0.5587 0.6733 0.7789 0.8934 1.0014 0.1 105 0.2221 0.33 1 1 0.4439 0.5535 0.6703 (A) H,O(l)+ 1-PrOH(3) 1.1211 0.174 1.1477 0.143 1.1740 0.133 1.2078 0.111 1.2172 0.089 1.2288 0.103 1.2321 0.107 1.2370 0.1 19 (B) H20(l)+ 2-PrOH(3) 1.1140 1.1340 1.1597 1.1920 1.2098 1.2370 1.2522 1.2687 0.167 0.140 0.130 0.109 0.103 0.097 0.117 0.119 (C) HzO(l)+ t-BuOH(3) 1.1305 1.1511 1.1539 1.2285 1.2444 1.2596 1.2596 1.2668 0.171 0.142 0.127 0.084 0.082 0.090 0.100 0.103 0.8640 0.8682 0.8731 0.8787 0.8871 0.8906 0.8441 0.8982 0.8997 0.182 0.176 0.166 0.150 0.138 0.126 0.120 0.112 0.loo 0.191 0.181 0.171 0.157 0.141 0.136 0.131 0.128 0.185 0.174 0.166 0.151 0.143 0.135 0.129 0.127 0.191 0.176 0.168 0.138 0.132 0.128 0.124 0.123 0.172 0.161 0.152 0.142 0.137 0.132 0.130 0.126 0.121 (E) H20( 1)+ DMSO(3) 0.8592 0.157 0.180 0.8591 0.131 0.170 0.8628 0.120 0.161 0.8634 0.092 0.151 0.8749 0.073 0.141 0.8859 0.045 0.130 /cm3 g-1 0.091 0.084 0.064 0.056 0.048 0.043 0.040 0.036 0.097 0.081 0.073 0.062 0.046 0.036 0.032 0.032 0.111 0.098 0.074 0.057 0.047 0.040 0.038 0.035 0.090 0.090 0.089 0.088 0.087 0.084 0.083 0.082 0.081 0.118 0.133 0.133 0.148 0.152 0.157 (am ,/am 2); 1 ,ps /mol mol-1 - 0.52 - 1.40 - 2.09 - 3.41 - 6.74 - 6.43 - 6.06 - 4.21 - 0.51 - 1.29 - 1.72 - 2.77 - 4.27 - 5.78 - 3.43 - 2.65 - 0.40 - 0.91 - 1.50 - 3.78 - 5.35 - 6.42 - 5.14 - 5.93 0.30 0.50 0.54 0.36 0.06 - 0.43 - 0.97 - 2.25 - 6.37 - 0.42 - 0.70 - 0.84 - 1.25 - 1.68 - 2.58T .MORI, J . ROMIYAMA AND T . IIJIMA 2587 TABLE 3.-PREFERENTIAL BINDING OF t-BuOH TO PLLHBr FROM AQUEOUS MIXTURES ( ~ a ~ 2 ) ~ ~ , ~ ~ (an/acz),, ( a n / a ~ ) , , , ~ (8m318m2)E~,~~ /cmJ g-1 /cm3 g-1 /mol mol-1 - solvent (3) vol. % x3 C3/g cm-3 Y3/cm3 g-1 /cm3 g-1 (F) H20(1)+ t-BuOH(3) 10 0.020 0.0772 1.1305 0.160 0.178 0.111 - 0.50 20 0.043 0.1535 1.1511 0.123 0.160 0.098 - 1.29 40 0.110 0.3166 1,1990 0.086 0.140 0.064 - 3.83 50 0.158 0.3977 1.2285 0.077 0.134 0.057 - 5.52 60 0.219 0.4775 1.2444 0.080 0.130 0.047 - 7.40 70 0.304 0.5559 1.2561 0.085 0.126 0.040 - 9.58 80 0.429 0.6343 1.2668 0.111 0.124 0.033 - 5.33 Such characteristic preferential binding behaviour has been reported for PLLHBr in aqueous 2-Pr0H.I Fig.2 shows that the preferential binding of t-BuOH to PLLHBr from the aqueous solution also conforms to this behaviour pattern. Summarizing these results, the charged poly(a-amino acid)s in aqueous monohydric alcohol solutions 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 x3 FIG. 1 .-Dependence on solvent composition of preferential binding of organic solvent to PLGNa ; in aqueous 1-PrOH (O), ZPrOH (A), t-BuOH (0), EG (0) and DMSO (V). The lines for aqueous EG and DMSO were drawn through the experimental points. The line for aqueous monohydric alcohol solutions was drawn assuming a hydration number of 20 for PLGNa. See text for details.Tanford l6 and Inoue and Timasheff have related the preferential binding parameter to the actual solvation number as (d1nJ/am2);1,pj = n’3-(x3ixl> (4) j j where n{ and n’J are the solvation numbers ofjth portion of the polymer by component 1 and 3, respectively. In terms of this equation, the linear relation shown in fig. 3 means that En4 = 0, i.e., the polymer is not solvated by the alcohols and the slope gives the hydration number which is constant with the change in solvent composition. Linear regression analysis gave the slopes as 19, 20 and 21 for aqueous 2-PrOH, t-BuOH and 1-PrOH, respectively. However, the shortage of points for each series and the errors inherent in these measurements exclude attachment of any significance to these differences ; hence we conclude that PLGNa is hydrated by20 water molecules2588 PREFERENTIAL BINDING TO CHARGED POLY (AMINO ACID)S in aqueous monohydric alcohol solutions.The hydration number, 27, of PLLHBr in aqueous 2-PrOH,' was confirmed for aqueous t-BuOH as shown by the plot in fig. 4. The persistence of these hydrations indicates that the hydrated layers of polymer including the counter-ions are not perturbed unless the alcohol content of 0 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 x3 FIG. 2.-Dependence on solvent composition of preferential binding of organic solvent to PLLHBr ; in aqueous 2-PrOH (A), and t-BuOH (0). The line was drawn assuming a hydration number of 27 for PLLHBr. See text for details. the solvent mixtures exceeds certain thresholds. It is also recalled that analysis using a definite solvation number was successfully applied to the preferential binding behaviour of PLLHBr in aqueous organic and organic solvent binary mixtures in the previous rep0rt.l The bulk dielectric constants of the solvent mixtures investigated in this and the previous study never fell below 40, except for the aqueous monohydric 0 r( 4 I 8 -2.0 8 CI -- -4.0 3 0; n (0 $' -6.0 c$ -8.0 \ m 0 A 0 0 A 00 0 I I I I I I 1 ' <L 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 1.3 1.4 x3 /XI xni = 20.FIG. 3.-Plot of preferential binding parameter against x3/x1 for PLGNa in aqueous 1-PrOH (o), 2-PrOH (A) and t-BuOH (0). The line was drawn according to eqn (2) assuming 2n3 = 0 andT. MORI, J . KOMIYAMA AND T. IIJIMA 2589 alcohol solutions at their higher alcohol contents.From the electrostatic point of view, the lowering of the dielectric constant is primarily concerned with the binding of the counter-ions by the polymer charges. If the binding is such that dehydration of the pairing ions occurs, the preferential binding behaviour must reflect this, and we discuss below the preferential binding behaviour after the breaks in this context. Therefore no dehydrative counter-ion binding takes place for these charged poly(a- amino acid)s when the bulk dielectric constant of the solvent mixture is above a threshold, ~ 4 0 . This argument has been added here because dehydrative counter- ion binding has so often been referred to in discussions of the electrolytic behaviour of poly(carboxy1ic acid)s, 8-21 including PLGNa, in aqueous solution.I 0 - 2.0 I 2 -4.0 - M \ 8 g -10.0 m -6.0 5 2 -8.0 0s (D \ -12 .o 0 A A I I I I I I I ' c 4 c L 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.80.9 1.3 x3 1x1 FIG. 4.-Plot of preferential binding parameter against x3/x1 for PLLHBr in aqueous 2-PrOH (A) and t-BuOH (0). The line was drawn according to eqn (2) assuming En3 = 0 and Cnl = 27. The exceptionally large hydration numbers assigned to PLGNa and PLLHBr in the aqueous monohydric alcohol solutions need some comment. Kauzmann and Kuntz 22 pointed out that the hydration numbers of polypeptides depend critically on the method of investigation and the state of the polymer being measured. Kuntz 23 gave 4.5 and 7.5, respectively for PLGK and PLLHCl residue hydrations, by n.m.r. measurements. Breuer and Kennerly 24 and recently Rochester and Westerman 25 calculated these numbers as 3 and 4 for PLGNa and PLLHBr from the water sorption isotherms of these polymers.From the buoyant density measurements of PLGCs and PLLHBr in aqueous CsCI, Ifft and coworkers 26 have reported values of 6.5 and 16.5 for the residue hydrations. These instances show that at least several water molecules are " strongly " bound to the charged groups and the counterions of these poly(a-amino acid)s. In addition, the hydration of the peptide bond has often been referred to in the literature. Though the large hydration numbers found for PLGNa and PLLHBr in the present study in principle include the excess water molecules in unspecified regions around the polymers, they are probably located in the immediate neighbourhood of the polymer entities, by thermodynamic preference, not necessarily by strong interaction with the charges or the polar groups of the polymers.Their differences from the hydration numbers cited above may be explained by this fact. It is questionable whether the hydrations found in this study persist in aqueous solution.2590 PREFERENTIAL BINDING TO CHARGED POLY (AMINO AC1D)S TABLE 4.-ELECTRICAL CONDUCTANCE OF PLGNa AND PLLHBr IN AQUEOUS ALCOHOL SYSTEMS alcohol A* VOI. % x3 /cmz Ja-1 equiv.-1 q/mW (A) PLGNa (1) H20+ 1-PrOH 0 10 20 30 40 50 60 70 80 90 10 20 30 40 50 60 70 75 80 85 90 10 20 30 40 50 60 70 75 80 85 10 20 30 40 50 60 70 80 90 O.OO0 0.026 0.057 0.093 0.138 0.192 0.265 0.359 0.491 0.681 43.8 31.9 23.2 17.2 13.2 9.24 6.61 2.72 1.12 0.26 (2) H20+ 2-PrOH 0.025 0 055 0.091 0.135 0,190 0.260 0.354 0.416 0.484 0.571 0.679 34.3 22.7 15.8 10.8 7.40 5.15 2.30 1.27 0.77 0.45 0.19 (3) H20+ t-BuOH 0.020 0.043 0.066 0.110 0.158 0.219 0.304 0.359 0.429 0.519 0.034 0.074 0.121 0.176 0.241 0.326 0.423 0.561 0.739 31.0 19.5 13.4 8.96 5.40 2.66 1.16 0.82 0.30 0.10 (4) HzO+EG 34.4 25.5 19.5 15.6 11.3 7.81 5.75 4.21 2.80 8.91 12.8 16.3 21.0 24.4 26.2 27.1 26.9 25.0 22.9 12.9 18.1 23.2 27.8 30.1 31.0 30.1 28.6 27.1 25.1 23.2 13.0 20.0 25.8 34.4 40.4 45.1 47.8 48.2 47.8 46.6 11.5 15.0 20.5 26.6 34.1 48.0 68.0 104.0 160.0T.MORI, J . KOMIYAMA AND T. IIJIMA 259 1 TABLE 4.-contd. alcohol A*/ vol. % x3 /an2 a-1 equiv.-1 q/mPt (B) PLLHBr (1) H20+ 2-PrOH 0 O.Oo0 54.9 8.91 20 0.055 28.1 18.1 40 0.135 14.3 27.8 60 0.260 8.31 31.0 70 0.354 5.78 30.1 80 0.484 3.92 27.1 90 0.679 1.95 23.2 (2) HzO+ t-BuOH 10 20 30 40 50 60 70 75 80 90 0,020 0.043 0.066 0.110 0.158 0.219 0.304 0.359 0.429 0.632 37.2 25.8 18.8 13.1 7.84 4.95 3.20 1.94 1.10 0.37 13.2 20.0 25.8 34.4 40.4 45.1 47.8 48.2 47.8 45.3 * Measured for the solutions of PLGNa and PLLHBr at 1.65 x residue mol dm-3.Visco- sities were taken from ref. (14). Recently, Inoue and Izumi 27 reported negative preferential binding of lY4-dioxan to poly[N5-(a-hydr~xypr~py1)-~-glutamine], a neutral poly(a-amino acid), from aqueous mixtures. Replotting their results in terms of eqn (4) indicated that 12 or more water molecules are in excess around the polymer residue in 0 to 70 weight % dioxan solutions. It seems that the large excess of water molecules are not specific to the charged poly(a-amino acid)s in aqueous monohydric alcohol solutions.The reversal of the decreasing trends in preferential binding in aqueous mono- hydric alcohol solutions will now be discussed. Fig. 1 and 2 show that breaks appear at x3 - 0.27 for PLGNa and at x3 - 0.29 for PLLHBr, irrespective of the alcohol species. The dielectric constants of the solvent mixtures at these points are 40, 40 and 29 for aqueous 1-PrOH, 2-PrOH and t-BuOH at x3 = 0.27, and are 38 and 27 for aqueous 2-PrOH and t-BuOH at x3 = 0.29, re~pective1y.l~ After these points, the preferential binding of the alcohol begins to increase, or the preferential hydration to decrease, with the alcohol content. The decreasing trend differs according to the alcohol species ; it is much sharper for PLGNa in aqueous 2-PrOH than in the other aqueous alcohol solutions and is also sharper for PLLHBr in aqueous t-BuOH than in aqueous 2-PrOH.If we compare this behaviour with the helix contents of the polymers shown in fig. 7 and 8, in some cases, the helix formation starts at a signifi- cantly higher value of x3 than that where the breaks are located and this parallels the less drastic decrease in preferential hydration. These results suggest two possible reasons for the reversed trends : (i) by dehydration accompanied by counter-ion binding and (ii) by the helix formation of the po1ymers.l. 2 8 In terms of the solvation model proposed in the previous paper, water molecules bound to the polar portion of the relevant polymer are released upon counter-ion binding and a fraction of those bound to the less polar portion are released upon helix formation.According to2592 PREFERENTIAL BINDING TO CHARGED POLY (AMINO ACID)S eqn (4), the remaining hydration numbers at each stage are estimated from the slopes of the straight lines connecting the relevant points to the origin in fig. 3 and 4 and are shown in fig. 7 and 8. The difference from the full hydration number gives the hydration loss, which can be resolved into two contributions if the degree of the F - 8 3.0 b 0 2.0 1.0 I I I I I I I 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 x3 FIG. 5.-Dependence on solvent composition of Walden products of PLGNa in aqueous 1-PrOH (O), 2-PrOH (A), t-BuOH (0) and EG (0). * Represents the value in water.The broken line indicates 1/10 of the constant value at lower alcohol concentrations. counter-ion binding is known. The electrical conductance of the polymers in aqueous alcohol solutions was measured for this purpose. The results are given in table 4 and are shown in fig. 5 and 6 as the Walden products. In aqueous EG, whose dielectric constant is relatively high at over 40, the products for PLGNa are essentially constant. In aqueous monohydric alcohol solutions, the products for both polymers x3 FIG. 6.-Dependence on solvent composition of Walden products of PLLHBr in aqueous 2-PrOH (A) and t-BuOH (0). * Represents the value in water. The broken line indicates 1 /lo of the constant value at lower alcohol concentrations.T . ICIORI, J . KOMIYAMA A N D T.IIJIMA 2593 deviate downward from x3 - 0.07, become about a half of the constant value as x3 is increased to 0.18-0.24, and finally fall to < 1/10 at the highest alcohol contents. In fig. 7 and 8, the numbers in the symbols indicate the points which were measured when the corresponding Walden products became <1/10 of the constant value at lower alcohol contents. The small values suggest that the dehydrative counter-ion binding is close to completion. If this is assumed, the points denotecl by *1 and *2 in fig. 7 should give the hydration number of PLGNa, which has essentially full counter-ion binding and also maximum helical conformation. Subtraction of this 2- 9'' f 1 I I I 4' I I I I ; ; I 1 ; 0 0.1 0.2 0.3 0.4 0.5 0.6 x3 FIG. 7.-Dependence on solvent composition of hydration number and helix content of PLGNa in aqueous 1-PrOH (o), 2-PrOH (A) and t-BuOH (0).number, 3, from 20 gives the cumulative dehydration number by these processes as 17. Points indicated by *3 and "4 in fig. 7 give 8 as the hydration number of PLGNa with the same counter-ion binding but with only a small fraction of helical content ; hence the dehydration number due to the counter-ion binding is 12. The difference between these numbers gives 5 as the dehydration number upon the helix formation of PLGNa. In terms of the solvation model, the total hydration number, 20, of PLGNa, is divided into 12 for the polar portion and 8 for the less polar portion, of which 5 water molecules are released upon helix formation. The same calculation can be applied to PLLHBr data.The points denoted by *6 in fig. 8 should give the2594 PREFERENTIAL BINDING TO CHARGED POLY (AMINO ACID)S hydration number of helical PLLHBr with full counter-ion binding. This number, 7, gives the dehydration number due to the two processes as 20, which may be divided into 15, previously given as the hydration number for the polar portion and 5, as the dehydration number due to helix formation. The last number is comparable with 6 , the desolvation number for PLLHBr upon helix formation in dimethyl- formamide + DMSO mixtures.l The comparison of hydration, conductance and helix content, sheds further light on the influence of the polymer charge on helix formation. Fig. 7 shows that, in aqueous 1-PrOH and t-BuOH, the helix formation of PLGNa takes place after the C I I I Q r I 0 0.1 0.2 0.3 0.4 0.5 0.6 i o o 80 6 0 4 0 2 0 x o ' = 3 8 2 - .Y .-( x3 FIG.8.-Dependence on solvent composition of hydration number and helix content of PLLHBr in aqueous 2-PrOH (A) and t-BuOH (0). dehydrative counter-ion binding is completed ; while in 2-PrOH, the formation precedes completion and fig. 8 shows that this is the case for PLLHBr in both aqueous alcohol solutions. For the last three cases, the hydration numbers at the 5 % helix level were read from the figures as 10 for PLGNa in aqueous 2-PrOH, 18 for PLLHBr in aqueous t-BuOH and 19 for PLLHBr in aqueous 2-PrOH. The differences between these numbers and the total solvation numbers of the respective polymers give the dehydration numbers, from which the degree of counter-ion binding is obtained as 80 % for PLGNa in aqueous 2-PrOH, 60 and 50 % for PLLHBr in aqueous t-BuOH and 2-PrOH.In all cases, the dehydrative counter-ion binding proceeds in preference to the helix formation, implying that fractions of the chargesT. MORI, J . KOMIYAMA AND T. IIJIMA 2595 on the polymer side chains are rendered dormant by the binding, whereby the electro- static repulsion amongst the charges is reduced. In a given aqueous alcohol, PLLHBr begins to form a helix under a higher charge density than PLGNa: a reasonable result considering that PLLHBr has the longer side chain by two methylenes. In the past, many investigations have been made on the helix formation of charged poly(a-amino acid)s in aqueous monohydric alcohol solutions.Barteri and Pispisa studied the helix formation of PLLHBr in aqueous 2-PrOH. Though they interpret this as " fully charged helix formation " and this is the current view of the helix in aqueous alcohol the term seems to be misleading. Their speculation that extensive dehydration from around the charged sites on the polymer induces the helix formation is valid in the sense that such a dehydration is accompanied by counter-ion binding and so is indirectly concerned with the helix formation. Among the three aqueous monohydric alcohol solutions, aqueous t-BuOH has the lowest dielectric constant at a given composition. However, as found in fig. 7, the dehydrative counter-ion binding for PLGNa in aqueous 2-PrOH takes place at lower alcohol contents than in aqueous t-BuOH, indicating the ease of dehydration in the former mixture.The dehydration from around the less polar portion of the polymer should also be connected with the energetics of the helix formation. Fig. 7 shows that the helix formation of PLGNa in aqueous 2-PrOH takes place at lower alcohol contents than in aqueous 1-PrOH. Since the dielectric constants of these aqueous alcohol solutions are the same at any composition investigated, it is suggested that this dehydration in aqueous 2-PrOH is favoured at lower alcohol contents than in aqueous 1 -PrOH. This research was supported in part by the Scientific Research Fund of the Japanese Ministry of Education. The authors are indebted to Dr. Hisashi Uedaira of the Research Institute for Polymers and Textiles for the use of a Carl Zeiss inter- ferometer.J. Komiyama, T. Mori, K. Yamamoto and T. Iijima, J.C.S. Faraday I, 1977, 73,203. Y. Koiwa and Y. Fujimoto, 22nd Disc. Meeting SOC. ofPolymer Sci (Tokyo, 1973). Y. Iwakura, K. Uno and M. Oya, J. Polymer Sci., 1967,5,2867. R. J. Hawkins and A. Holtzer, Macromolecules, 1972,5,294. J. A. Riddick and W. Bunger, Techniques of Chemistry, Vol. 11, Organic Solvents, ed. A. Weissberger (Wiley-Interscience, New York, 1970). Handbook of Chemistry and Physics, ed. R. C. Weast (Chemical Rubber Co., Cleveland, Ohio, 48th edn, 1968). C. W. N. Cumper, J. F. Read and A. I. Vogel, J. Chem. SOC., 1965, 5323. G. Holzwarth and P. Doty, J. Amer. Chem. Soc., 1965,87,218. * H. Terayama, J. Polymer Sci., 1952, 8,243. lo A. E. Dunstan, 2. phys. Chem. A, 1905,51,732. l1 A. E. Dunstan, J. Chem. Soc., 1905,87,11. l2 A. L. Olsen and E. R. Washburn, J. Phys. Chem., 1938,42,275. l3 E. P. Irany, J. Amer. Chem. Soc., 1943,65,1396. l The Physico-chemical Constants of Binary Systems in Concentrated Solutions, ed. J. Timmermans (Interscience, New York, 1960), vol. N. l5 H. Inoue and S. N. Timasheff, J. Amer. Chem. SOC., 1968,90,1890. l6 C. Tanford, J. Mol. Biol., 1969,39, 539. l7 H. Inoue and S. N. Timasheff, Biopolymers, 1972, 11,737. l 8 G. S. Manning and A. Holtzer, J. Phys. Chem., 1973,77,2206. l 9 A. Ikegami, Biopolymers, 1968,6,431. 'O C . Tondre and R. Zana, J. Phys. Chem., 1972,76,3451. 21 J. Komiyama, M. Ando, Y . Takeda and T. Iijima, Europ. Polymer J., 1976, 12,201. '' I. D. Kuntz and W. Kauunann, Adv. Protein Chem., 1974,28,239. 23 I. D. Kuntz, J. Amer. Chem. Soc., 1971,93,514. 24 M. M. Breuer and M. G. Kennerley, J. Colloid Interface Sci., 1971,37, 124.2596 PREFERENTIAL BINDING TO CHARGED POLY (AMINO ACID)S 2 5 C. H. Rochester and A. V. Westerman, J.C.S. Faraday I, 1976, 72, 2753. 26 R. Almassy, J. S. V. Zill, L. G. Lum and J. B. Ifft, Biopolymers, 1973, 12, 2713. 27 H. Inoue and T. Izumi, BiopoZymers, 1976, 15,797. 28 G. Nemethy and H. A. Sheraga, J. Chem. Phys., 1962,36, 3401. 29 M. Barteri and B. Pispisa, BiopoZymers, 1973, 12,2309. 30 R. H. Liem, D. Poland and H. Scheraga, J. Amer. Chem. Soc., 1970,92,5717. (PAPER 7/1599)

ISSN:0300-9599    

DOI:10.1039/F19787402583

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Tin Oxide Surfaces Part 4.-Infrared Study of the Adsorption of Oxygen and Carbon Monoxide + Oxygen Mixtures on Tin(1v) Oxide, and the Adsorption of Carbon Dioxide on Ammonia- pretreated Tin(1v) Oxide BY PHILIP G. HARRISON* AND EDWARD w. THORNTON Department of Chemistry, University of Nottingham, University Park, Nottingham NG7 2RD Received 25th October, 1977 The transmission infrared technique has been used to study the adsorption of oxygen and carbon monoxide+oxygen mixtures on tin(rv) oxide. Adsorption of oxygen results in the appearance of two weak bands at 1155 and 1020 cm-' which did not shift using 180-enriched oxygen. No assign- ment of the species responsible for these bands was made, but the accompanying broad intense band below 900 cm-' is associated with surface Sn-O-Sn bridges.Surface carbonate and bicarbonate species are formed slowly when the oxide is exposed to carbon monoxide+ oxygen mixtures, essentially independent of the gas mixture composition in the range 10-70 % CO. No bulk reduction of the oxide was observed, in contrast to the behaviour previously observed with carbon monoxide alone. Hydration of the oxide surface prior to treatment with adsorbent severely inhibits the adsorption of carbon dioxide and carbon monoxide + oxygen mixtures. Preadsorbed ammonia has a similar inhibiting effect with carbon dioxide, but leads to the formation of a low stability carbamate salt of the unsubstituted carbamic acid, H2NC02H. Tin(1v) oxide gel is an effective low-temperature catalyst for the oxidation o carbon monoxide by oxygen. The mechanism of the catalysis, proposed from kinetic data, involves successive reduction and reoxidation of the oxide catalyst,l whilst our previous infrared investigations have demonstrated the formation of a surface unidentate carbonate species prior to the irreversible reduction of the bulk oxide, which causes darkening and loss of transmission.The same species, along with a surface bicarbonate, is formed when carbon dioxide is chemisorbed on tin(1v) oxide. In this paper, we report infrared studies of the adsorption of oxygen and carbon dioxide mixtures on tin(1v) oxide. EXPERIMENTAL The apparatus, the tin(1v) oxide sample, the procedure for the preparation of discs suitable for examination, the infrared cell, the spectrometer, oxygen, carbon monoxide, carbon dioxide, and ammonia have been described previously.2* All the discs used had a " thickness " of M 11.2 mg cm-2.The ordinate axes of fig. 2 and 4-6 have been displaced so as to avoid overlapping of spectra. The numbers of the ordinate axes are the transmittance at the highest wavenumber shown, RESULTS AND DISCUSSION OXYGEN ADSORPTION The infrared spectrum of oxygen adsorbed on tin@) oxide has been observed by Gundrizer and Da~ydov,~ who observed a number of absorption bands in the range 1200-900 cm-I, and attributed them to adsorbed ion radical forms of oxygen of the type 0;. However, none of their assignments could be confirmed using *O-enriched oxygen. 25972598 TIN OXIDE SURFACES Fig. 1 shows spectra recorded during the adsorption of oxygen on a tin(1v) oxide disc which had been evacuated at 723 K.On a sample which had been evacuated overnight at 723 K, oxygen gave an increase in transmission above 900cm-l and an intense absorption shoulder appeared below 9OOcm-l which extended to the region where oxide absorption becomes intense ( R! 500 cm-l). In addition, two weak absorption bands appeared at 1155 and 1020 cm-l. These features were not immediately removed by re-evacuation at 723 K, and fig. l(3) shows a shoulder still remained at 800-700 cm-l after 2 h. The admission of 80-enriched oxygen (l80 = 63.8, 1602 = 26.9, 160180 = 18.5, 1 8 0 2 = 54.6 %) produced a very similar spectrum to that produced by normal oxygen. The two bands at 1155 and 1020 cm-l were also produced during 180-enriched oxygen adsorption on a tin@) oxide disc which had been allowed to stand in an 180-oxygen atmosphere at 723 K in order to enrich the surface with l80.w avenumber lcm-' FIG. 1.-Adsorption of oxygen on tin@) oxide. (1) Sample evacuated at 723 K. (2) During exposure to oxygen (800 Nm-2, 320 K); (3) after subsequent evacuation (723 K, 2 h) and (4) during exposure to l*O-enriched oxygen (63.8 % leg) at 1.33 kN m-2 at 320 K. The species responsible for the 1 155 and 1020 cm-l bands (which varied in intensity from sample to sample but always in concert) may be due to adsorbed oxygen species as Gundrizer and Davydov suggest. However, the inability to observe an isotopic shift with l80 suggests that the bands have a different origin, but it is difficult to establish its nature.The broad absorption band below 900 cm-l [which was present in all the spectra in ref. (4)] is associated with surface oxide ions and probably with the Sn-0-Sn bridge structure since the Sn-0-Sn bond vibration in (polymeric) tin(1v) species occurs in the range 735-755 ~ m - l . ~ Bands associated with the siloxane Si-0-Si bridge in silica have also recently been observed.6 ADSORPTION OF CARBON MONOXIDE-kOXYGEN MIXTURES During the adsorption of CO + O2 mixtures, the spectra recorded closely resembled the spectra produced by carbon dioxide adsorption on discs activated at the sameP . G . HARRISON AND E . W . THORNTON 2599 65- 75- 65- 65- - 10% - 1 wavenumber /cm-' FIG. 2.-Adsorption of 10 % CO in O2 on tin(Iv) oxide at 320 K and 3.0 kN m-2. Evacuation temperatures and periods of exposure to the gas were (1) 395 K, 130 min ; (2) 503 K, 120 min ; (3) 623 K, 10 min ; (4) 623 K, 120 min.Samples (3) and (4) were oxygen treated after evacuation. Spectra recorded during the chemisorption of C02 on tin(Iv) oxide (at 320 K, 2.0 kN m-2) after evacuation at ( 5 ) 508 and (6) 618 K. t 20,_J 0 1600 1400 1 200 lo00 lo00 1600 1400 1 200 wavenum ber/cm-' FIG. 3.-Adsorption of 20 % CO in O2 on tin(1v) oxide at 320 K and 1.33 kN m-2. (1) Sample evacuated at 723 K. During subsequent exposure to (2) O2 at 320 K and (3) reaction mixture at 320 K for 30 min. (4) After evacuation and treatment with I80-enriched O2 at 723 K. ( 5 ) During subsequent adsorption of a '*O-enriched CO+02 mixture at 320 K ; O2 63 % "0, CO 60 % '*O- enrichment.2600 TIN OXIDE SURFACES TABLE 1 .-CHARACTERISTIC BAND FREQUENCIES OBSERVED FOR SURFACE CARBONATO SPECIES adsorbate/( T / K ) characteristic band frequencies/cm- 1 surface species ref.1380 1380 1300 1225 1370 1295 1225 1275 1380 1300 1225 unidentate carbonate bicarbonate * unidentate (carbonate 2 bicarbonate bidentate (carbonate unidentate carbonate bicarbonate { this work 1290 bidentate lo20 {carbonate thiswork a Shifts to 1560 cm-1 on a deuterated surface ; remains unchanged on a deuterated surface. temperature. However, the spectra were less intense and the peaks due to the unidentate carbonate were relatively more intense. Also, the spectra increased in intensity over a period of ~2 h, whereas the bands produced by carbon dioxide adsorption reached equilibrium after only a few minutes.In contrast to the experi- ments for pure CO adsorption,2 the sample disc was not discoloured and did not suffer an overall Ioss in transmission during adsorption of a CO+O, mixture with 69 - 78- 43- 55- 8 *I f: E * 43- 10% 45- __ - 1 I I I 1 1 )o 1600 1400 1 wavenumber/cm-' FIG. 4.-Adsorption of 50 % CO in O2 on tin(1v) oxide at 4.1 kN m-2. (1) Sample evacuated (518 K, 19 h). Spectra recorded after reaction with the gas mixture (2) at 320 K for 0.5 h ; (3) at 398 J for 20 h, and (4) after subsequent evacuation at 320 K for 2 h. (5) Sample evacuated, exposed to ammonia (1.33 k N m-2, 320 K, 10 min), and re-evacuated (320 K, 2 h). (6) After subsequent exposure to the gas mixture.P . G . HARRISON AND E . W. THORNTON 260 1 between 10 and 70 % CO.Some typical spectra recorded during exposure of tin(1v) oxide discs to CO+O, mixtures at 320 K are shown in fig. 2 and 3, while typical characteristic frequencies for surface carbonato species are listed in table 3. Fig. 3 shows spectra recorded during the adsorption of CO + 0, and '0-enriched CO + 0, on a tin(1v) oxide disc which had been evacuated at 723 K prior to the adsorption. A new product is indicated by three absorption bands at 1580, 1290 and 1020 cm-' shifting to 1460, 1265 and 1020 cm-1 on *O substitution. These can be identified as being due to a bidentate carbonate complex, similar to that formed by C 0 2 adsorption on tin(1v) oxide discs which had been thermally activated at 718 K and exhibit bands at 1590 and 1275 ~ m - l . ~ wavenumber /cm-' FIG.5.-Effect of re-hydration on the adsorption of carbon dioxide. (1) Background spectrum after evacuation at 516 K for 3 h. (2) During subsequent exposure to carbon dioxide (2.0 kN m-2, 320 K). (3) Background after activation as for (1) followed by exposure to water vapour (1.33 N m-2, 0.5 min, 320 K) and evacuation (24 h, 320 K) ; (4) during subsequent exposure to carbon dioxide (2.0 kN m-2, 320 K). Some qualitative observations derived from the above and other experiments can be summarised for the adsorption of CO + 0, mixtures during the first two hours of exposure in a closed cell compartment: (a) bicarbonate maxima formed only in conditions where CO, was produced and the surface had been treated at a sufficiently high temperature (> 500 K).(b) The carbonate bands formed in relatively larger proportions at equilibrium (compared to C 0 2 adsorption), but the bicarbonate maxima more rapidly established their full intensity. (c) For adsorption at 320 K, all of the bands formed much more rapidly on a surface which had been oxygen treated at the pretreatment evacuation temperature. Reaction with a highly outgassed surface was very slow (several hours). ( d ) For samples pretreated above about 500 K the carbonate maxima (1440 and 1380 cm-l) and the shoulder at 1580 cm-l increased considerably in intensity on heating in a CO+O, mixture above 320 K (fig. 4 for example). (e) The spectra produced were essentially independent of the gas mixture composition in the range 10-70 % COY provided that the experiment was not carried on so long as to seriously deplete the oxygen present, when some loss in transmission occurred.Thus, at 320 K a CO + O2 mixture behaves more like COz than CO in its chemi- sorption reactions, but on heating in the gas mixture a large amount of surface2602 TIN OXIDE SURFACES carbonate is produced showing absorption maxima at 1440 and 1380 cm-l. These bands are much reduced in intensity by evacuation at 320 K (fig. 4), and the unidentified shoulder at 1580 cm-' is removed. The remainder of the carbonate can only be removed by evacuation at 400-500 K. The carbon monoxide is probably initially adsorbed at an oxide site rather than a Lewis acidic metal ion [observation (c) above]. The carbonate species may be the intermediate in the Sn0,-catalysed oxidation of carbon monoxide.I I I I I 1 I I 1 800 1600 1400 1200 wavenumber /cm-I FIG. 6.-Chemisorption of carbon dioxide on tin(rv) oxide having preadsorbed ammonia. (1) Sample evacuated (508 K, 2 h) and then exposed to COz (2.0 kN m-2, 320 K). Sample then activated as for (l), exposed to ammonia at room temperature and then evacuated under the following conditions before cooling to room temperature and recording the spectrum in a COz atmosphere (2.0 kN m-2, 320 K) : (2) 320 K, 2 h ; (3) 508 K, 0.5 h ; (4) 508 K, 17 h ; (5) 563 K, 3 h ;L(6) 613 K, 3 h ; (7) 611 K, 15 h plus oxygen treatment at 61 1 K. The adsorption of CO + O2 on a surface which had been oxidised at 723 K resulted in the formation of a species similar to that found by Ismailov et aZ.* for CO adsorp- tion on 773 K evacuated SnO, (this work : 1580+ 1290 cm-1 ; Russian authors quote 1630+1260cm-l).However the admission of CO to a tin(1v) oxide disc which had been outgassed at x723 K resulted in such a massive loss in transmission that no spectra could be recorded, hence the situation in the two experiments are evidently not directly comparable. EFFECT OF PREADSORBED WATER AND AMMONIA ON CHEMISORPTION OF COz, CO AND CO+O, MIXTURES We have previously shown that treatment of the tin@) oxide surface with trimethylchlorosilane severely inhibits adsorption of carbon monoxide and carbon dioxide, presumably by steric hindrance of surface bare metal and oxide sites.gP. G . HARRISON AND E. W. THORNTON 2603 When a tin(1v) oxide disc was rehydrated after thermal activation, the chemisorp- tion of carbon dioxide was severely inhibited and the reaction products formed were those expected for a low temperature surface (fig.5). Very similar results were observed for the adsorption of CO + O2 mixtures (not illustrated). The effect of ammonia on the adsorption of a CO+02 mixture had a similar effect in that adsorption was inhibited, the intense carbonate bands were not produced, and no catalytic oxidation to C02 occurred (detected in the gas phase infrared spectrum during adsorption on the pure oxide) (fig. 4). However, preadsorbed ammonia reacted with carbon dioxide to produce a low stability carbamate salt of the unsubstituted carbamic acid, H2NC02H. This is illustrated in fig. 6, where it can be seen that the absorption spectrum of C02 adsorbed on an ammonia-covered surface was due to the carbamate ion.After the removal of some ammonia by evacuation at temperatures up to x 573 IS, a composite spectrum resulted due to carbamate, carbonate and bicarbonate species. Complete removal of ammonia resulted in the restoration of the expected spectrum due to C02 adsorption on a clean surface. The four bands characteristic of the adsorbed carbamate ion occurred at 1609 cm-l [G(NH,)], 1510 cm-' [v,,(COO)], 1450 cm-l [v(C-N)] and 1245 cm-l [6,,(NH2)]. These positions agree with the positions reported for ammonium carbamate lo at 1624,1525,1404 and 1261 cm-l, respectively. The formation of a low stability N-substituted salt of carbamic acid during the adsorption of carbon dioxide on a microporous ion exchanger containing 2-amino- ethylamine groups has also been reported recently.l CONCLUSIONS The data show that a surface bidentate carbonate and a bicarbonate species are chemisorbed on to tin@) oxide from CO+02 mixtures, and are presumably the intermediates formed in the tin(1v) oxide-catalysed oxidation of carbon monoxide. The site of the chemisorption (of CO from the gas mixture) is most probably an oxide ion since (a) chemisorption takes place more readily on a preoxidised surface, and (b) preadsorption of water or ammonia, which coordinate to bare metal sites on the surface, severely inhibits chemisorption from CO + 0, mixtures. We thank the S.R.C. and the International Tin Research Institute for support in the form of a CASE Award (to E. W. T.). M. J. Fuller and M. E. Warwick, J. Catalysis, 1973, 29,441. E. W. Thornton and P. G. Harrison, J.C.S. Faraday I, 1975, 71,461. P. G. Harrison and E. W. Thornton, J.C.S. Faraday I, 1975,71, 1013. T. A. Gundrizer and A. A. Davydov, Reaction Kinetics and Catalysis Letters, 1975, 3, 63. B. Jezowska-Trzebiatowska, J. Hanuza and W. Wojciechowski, Spectrochim. Acta A, 1967, 23, 2631. A. Zecchina, personal communication. ' P. G. Harrison and E. W. Thornton, unpublished observations. * E. G. Isvailov, V. F. Anufrienko, N. C. Maksimov, A. A. Davydov and V. D. Sokolovskii, Doklady Akad. Nauk S.S.S.R., 1974, 216, 847. P. G. Harrison and E. W. Thornton, J.C.S. Farad'ay I, 1976, 72, 1310. V. M. Kirytenko, A. V. Kiselev and B. I. Lygin, Kolloid. Zhur., 1975, 37, 382. lo D. L. Frasco, J. Chem. Phys., 1964,41,2134. (PAPER 7/1873)

ISSN:0300-9599    

DOI:10.1039/F19787402597

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Tin Oxide Surfaces Part 8.t-Infrared Study of the Mechanism of Formation of a Surface Isocyanate Species on SnO, .055 CuO during Catalysis of the Oxidation of Carbon Monoxide by Nitric Oxide BY PHILIP G. HARRISON* AND EDWARD W. THORNTON Department of Chemistry, University of Nottingham, University Park, Nottingham NG7 2RD Received 2nd November, 1977 An experimental study has been carried out on the mechanism of formation of a surface isocyanate species on the mixed oxide catalyst SnOz. 0.55 CuO during the initial stages of catalysis of the CO+NO reaction at %470 K. Using infrared spectroscopy, the isotopic shifts of the 2189 cm-1 pseudo-antisymmetric stretching vibration have been measured for 13C, 15N and lSO substitution. The oxygen atom of the surface isocyanate has been shown to originate from NO rather than CO as was previously assumed.This observation has been interpreted in terms of a mechanism involving initial dissociative chemisorption of CO followed by the formation of a fulminate via reaction of NO with the surface carbon atom and subsequent rapid isomerisation to the isocyanate : 23 +co +]-c + 3-0 1-C+NO-+ 1 -CNO 1-CNO + 1-NCO. The presence of surface isocyanate species on catalysts during the catalysis of the CO +NO reaction has been well established by means of infrared spectros~opy.l-~ For noble metal catalysts Unland 1-3 has suggested the surface isocyanate as an intermediate in ammonia formation in a damp CO+NO reaction mixture. For copper metal catalysts, Rewick and Wise interpreted their kinetic data in terms of a strongly bound and thermally stable -NCO adspecies on the copper surface exhibiting primarily the function of a catalyst inhibitor.Similarly, London and Bell in their study of a silica-supported CuO catalyst were able to incorporate the formation of a surface isocyanate species into their reaction mechanism and interpret this as strongly supporting the postulate that nitric oxide can dissociate upon adsorption. These studies lm5 were all carried out at elevated temperatures ( 3 -420 K), and the authors did not investigate the mechanism(s) whereby the surface isocyanates were formed. In contrast, the mechanistic study of Brown and Gonzalez ti on silica-supported ruthenium was carried out at x 300 K. In this case, the proposed mechanism involves the interaction of an adsorbed NO molecule with two adjacent CO molecules to form an isocyanate and a molecule of carbon dioxide, without the dissociative chemi- sorption of either reactant.The study of Arai and Tominaga on alumina-supported rhodium at 473-573 K concluded that a reactive Rh-N species was formed from the CQ surface complex by elimination of C02 and then could subsequently react / \ Rh NO with CO to form Rh-NCO or NO to form Rh-N20. t Part 7. P. G. Harrison and E. W. Thornton, J.C.S. Faraday I, 1976, 72, 2484. 2604P . G . HARRISON AND E . W. THORNTON 2605 involved isotopically substituted reactants, only 13C-enriched CO and 15N-enriched NO have been used to demonstrate that both the C and N atoms in the product isocyanate are derived from the reactants, and that the isotope shifts are consistent with an isocyanate surface species.Con- sequently, in the present study we have extended the scope of the isotope substitution experiments to include l 8 0 and hence elucidate the origin of the oxygen atom in the surface isocyanate species, whether from CO, NO, or the oxide catalyst surface. This study has been carried out at ~ 4 7 0 K for reaction mixtures withp(CO)/p(NO) > 1.5 on a mixed oxide catalyst formed by complete ion exchange (during coprecipita- tion) of CU',~, on to hydrous tin@) oxide gel. The catalyst has been described by Fuller and Warwick for their kinetic studies during catalysis of the CO+O, and CO +NO reactions. Although some of the investigations EXPERIMENTAL APPARATUS The infrared cell and vacuum line have been described previously.O Infrared spectra were recorded with a Perkin-Elmer model 577 spectrometer, using abscissa expansion as required over the range 2400-2000 cm-l. The ambient beam temperature was 320 K. MATERIALS The oxide catalyst sample was kindly supplied by Dr. M. J. Fuller of the International Tin Research Institute, and came from a batch which was used by Fuller and Warwick for their kinetic studies of the CO+02 reaction.8 The Cu/Sn atomic ratio was 0.55 in the calcined material, this composition corresponding to the full utilisation of the parent ion exchange capacity of the hydrous Sn02 towards Cuxl. Oxygen, nitrous oxide and carbon monoxide (B.O.C., grade X) were used as received, carbon dioxide (Distiller's) was purified by passing through a stainless steel coil 1 m in length and 2 mm diameter held at 211 K by freezing 2-ethoxyethyl acetate, nitric oxide (B.O.C.) was purified by passing through the same trap at 197K (acetone+solid C02).1802 ( ~ 6 3 atom % 13C0 (-91 atom % 13C), ClSO (-60 atom % l80), 15N0 (m95 atom % 15N) and N180 (-41 atom % l80) were obtained from B.O.C. Prochem and used without further purification. PROCEDURE For infrared experiments the oxide catalyst was ground with an agate pestle and mortar and pressed into self-supporting discs using a stainless steel die; the discs were 2.5 cm in diameter with a " thickness " of 12.2 mg cm-2. Before performing adsorption experiments, the discs were evacuated (w N m-2) in the infrared cell for 2-24 h at temperatures in the range room temperature to 630 K, usually followed by oxygen treatment at the evacuation temperature (1.3 kN m-2, 1 h) and re-evacuation for 0.5 h.The exact procedure varied for each experiment, and will be described in the Results section. For accurate measurement of the positions of the bands shown in fig. 6, the spectra were recorded with 5x or lox abscissa expansion at scanning speeds of 100 or 50 cm-l min-l. The wavenumber markers were calibrated with gaseous carbon monoxide. The positions are estimated to be accurate to _+2 cm-l. RESULTS DEHYDRATION Following evacuation at room temperature, the initially pale blue disc turned pale green, and the infrared spectrum showed a broad, intense hydroxyl stretching band (fig. 1) and a weak shoulder at = 1600 cm-l due to molecular water.The2606 TIN OXIDE SURFACES v(0H) band decreased on evacuation up to 420K and the disc became darker in colour. At higher evacuation temperatures (> 550-570 K), the v(0H) band was removed almost entirely, a weak band remaining at 3600cm-l. This band could be only partly exchanged with D,O vapour, and the surface hydroxyl groups could not be restored by exposure to water vapour (at room temperature or 570 K). At this stage, the disc was black in colour. The background spectrum in the range 100, I 1 I I 4000 3500 3000 2: wavenumber /cm-' io FIG. 1.-Infrared spectra of SnOz .0.55CuO disc after evacuation at (1) room temperature for 0.75 h ; (2) 373 K for 0.5 h ; (3) 508 K for 1 h ; (4) 513 K for 14 h ; and (5) 623 K for 2 h followed by oxygen treatment at 623 K (130 N m-2, 2 h).3500-900 cm-l was a sloping straight line showing no spurious absorption bands due to impurities. Treatment with oxygen had no effect on the colour, but at higher evacuation temperatures (> 570 K) the transmission of the disc was improved by oxygen treatment. Evacuation of the disc above ~ 7 2 0 K resulted in a severe loss in infrared transmission, which could not be restored by oxygen treatment, thus making infrared experiments impossible. ADSORPTION OF CARBON DIOXIDE The adsorption of carbon dioxide was studied at 320 K as a function of the disc pretreatment evacuation temperature in the range 320-650 K. After evacuation at 320 K, the catalyst disc showed three broad, strong infrared bands at 1530-40, 1380 and 1060cm-1 during exposure to CO, (fig.2). These decreased in intensity on evacuation at 320K and further decreased in intensity on evacuation at higher temperatures, being removed altogether at about 570 I<. The positions of the absorption bands are quite different from the carbonate bands on tin(1v) oxide,l O neither do they closely resemble the complex set of bands observed by Amerikov and Kasatkina during C 0 2 adsorption on silica-supported Cu0.l The three bands may be assigned to the stretching bands of a coordinated carbonate where the 1060 cm-l band is the v2(A,) [v,(XY) for planar XYJ vibration, and the 1530-40 and 1380 cm-l bands are due to the v,(A,) and v,(B,) [the split va(XY) for planar XY3] vibrations.12P . G . HARRISON AND E . W. THORNTON 2607 The splitting of the degenerate vibrations (Av = 150cm-') is more typical of a unidentate than a bidentate carbonate ; this range of splittings has been reported as approximately AY = 230-100 cm-' for a unidentate carbonate and Av = 370-320 cm-' for a bidentate ~arb0nate.l~ Further support for the assignment to a unidentate wavenumber /crn-' FIG.2.-Infrared spectra of a Sn02. 0.55CuO disc exposed to carbon dioxide. (1) Background after evacuation and treatment with DzO (s.v.P., 0.5 h) at 393 K. (2) During subsequent exposure to C02 (400 N m-2) at 320 K, and (3) after evacuation at 320 K for 1 h. carbonate comes from the position of v2 band at 1060 cm-l ; this lies in the middle of the reported range of frequencies for unidentate carbonates (1080-1040 cm-l), whereas the corresponding bands for bidentate carbonates are found at 1030- 1020 crn-l.l3 2.0 n 8 1.0 k 0 a\- evacuation temperature/K FIG.3.-PIot of the intensity of carbonate bands produced by adsorption of carbon dioxide on SnOz .0.55CuO as a function of pretreatment evacuation temperature. F(C0,) is the sum of the optical densities of the 1535 and 1380 cm-I bands. Points refer to the intensity of carbonate bands after evacuation ( x ) and during subsequent exposure to C02 at 320 K (0).2608 TIN OXIDE SURFACES As the pretreatment evacuation temperature was increased, the intensity of the carbonate bands both in equilibrium with C 0 2 at 320 K and after evacuation at 320 K decreased. This is shown graphically in fig. 3, where the intensity of the carbonate bands is measured by F(CQ,) which is the sum of the optical densities of the 1530-40 and 1380 cm-l bands.Both curves intersect the abscissa at ~ 7 2 0 K, which implies that no surface carbonate should form during exposure to C 0 2 of samples evacuated above this temperature. None of the catalyst sample discs showed a band at ~ 2 3 5 0 cm-I due to linear, adsorbed CQ2 during exposure to CQ2. This contrasts with silica-supported CuO where a band has been observed in this region in two previous studies.5’ l1 ADSORPTION OF NITROUS OXIDE Following evacuation in the range 320-650 K (and oxygen treatment above ~ 5 7 0 K), exposure to N20 at 320 K failed to produce any absorption bands due to adsorbed species in the range 4000-900 cm-l. 1 21 20 rl E 2. 3 2110 - I ; 21 00 / / / / I I I I 400 500 600 700 L 300 calcination tempera t ure/K FIG.4.-Plot of the stretching frequency of carbon monoxide adsorbed on SnO, .0.55CuO at 320 K against the pretreatment temperature for evacuation (0) and evacuation and oxygen treatment ( x ). ADSORPTION OF CARBON MONOXIDE Carbon monoxide was adsorbed to produce an intense, sharp band in the range 2100-2120cm-l together with weak carbonate bands at 1530 and 138Ocm-l. The latter two bands showed a progressive decrease in intensity as the pretreatment calcination temperature was increased, and extrapolation showed that no carbonate formation should occur for samples pretreated by evacuation on oxygen treatment above ~ 7 0 0 K, similar to the observation for CO, adsorption above. The band at ~ 2 1 0 0 cm-I can be readily assigned as due to the stretching vibration of adsorbed carbon monoxide.The exact position of the band was sensitive to pretreatment conditions : as the evacuation temperature was increased from 320 to 673 K the position of v(C0) changed continuously from 2100 to 21 10 cm-l, but when the sample was oxygen treated after evacuation above ~ 5 0 0 K, the frequency increased still further at 2121 cm-I (fig. 4). The absorption band became noticeably weaker and sharper for higher pretreatment temperatures, and could be removed completely from all samples by evacuation at 320 K for a few minutes.P . G . HARRISON AND E. W. THORNTON 2609 The position of the CO stretching band is more typical of CO adsorbed on copper metal, where the band has been variously reported at abound 2110 cm-1,14-20 than of CO adsorbed on copper(I1) oxide, where the vibration has been observed at 2130-2140 21* 22 On oxidised copper (where the adsorption site was presumed to be Cu+), the band has been reported at 2113 cm-l,19 whilst the CO complex with copper(1) chloride in water has v(C0) = 2112 Thus it would seem more probable that the adsorption site in the present study is a coordinatively unsaturated copper atom or ion in the 0 or +1 oxidation state, but which can be oxidised to Cu" at higher temperatures by molecular oxygen.On heating in CO above ~ 5 0 0 K the transmission of the sample disc decreased to such an extent that infrared experiments were no longer possible. Similarly, following evacuation or evacuation followed by oxygen treatment above M 730 K, the transmission of pressed discs was too low for infrared experiments to be carried out.wavenumber /cm-' FIG. 5.-Infrared spectra of nitric oxide adsorbed on to a SnOz .0.55CuO disc. (1) Background after evacuation at 508 K for 3 h ; during subsequent exposure to NO at 320 K at (2) 660 N m-2 and (3) 3.3 kN m-2. (4) Adsorption of NO (320 K, 4.0 kN m-2) on a disc after evacuation and oxygen treatment at 602 K. ADSORPTION OF NITRIC OXIDE No new infrared absorption bands were detected during exposure of catalyst discs preevacuated at -= M 380 K to nitric oxide at 320 K. For sample discs evacuated above x380 K NO was adsorbed to produce a weak band at 1873 cm-l (403 K evacuation), shifting as the pretreatment temperature was increased to 1880 cm-' (643 K evacuation and oxygen treatment) (fig.5). The absorption band was removed by pumping off the gas phase at 320 K, and was also displaced readily and completely by addition of CO to the gas phase at 320 K. The position of the band is close to that reported for NO adsorbed on silica-supported CuO (1890 cm-1).22 The position of the v(N0) absorption band may be taken as indicative of the type of bonding of the adsorbate to the surface, and from the classifications of Terenin et uZ.24-27 (reviewed more recently by Shelef and Kummer),28 it can be surmised that the bonding in the present case is either double bond ionic : -l=&=o or coordinative : J 1- :NO2610 TIN OXIDE SURFACES where ] represents the surface site, probably a coordinatively unsaturated copper ion or atom since separate experiments showed that no new bands were formed in this region during exposure of pure tin@) oxide discs (pretreated by evacuation and oxygen exposure at ~6623 K) to NO at 320 K.No bands due to oxidised species (such as NO2, NO; or NOT) were observed. ADSORPTION OF CO+NO MIXTURES After exposure of a 623 K evacuated and 0,-treated disc to a CO +NO mixture at 320 K for 1 h, new absorption maxima appeared at 1500 and 1350 cm-l due to a carbonate species and at 2120 cm-l due to adsorbed CO [fig. 6(a)]. Following 2500 r,,,, 21 0 2 wavenumber /cm-' (6) 0 K) FIG. 6.-Spectra recorded during the adsorption of CO+NO mixtures on SnOz .0.55CuO after activation in oxygen at 623 K. (a) (1) Background. After exposure to CO (6.7 kN m-2, 320 K) (2) and after exposure to 23 % NO in CO (8.65 kN m-2) at (3) 320 K for 1 h ; (4) 388 K for 0.25 h ; and (5) 473 K for 0.25 h.(b) Continued from (a) (5). Spectrum recorded after evacuation at 320 K for 5,10,15,20,25,30 and 35 mh. (c) (1) Background. After exposure to CO (6.7 kN m-2, 320 K) (2) and after exposure to 16.6 % NO in CO (8.0 kN m-2) at (3) 320 K for 0.25 h ; (4) 473 K for 0.5 h ; and 473 K for 1.5 h. subsequent heating in the mixture at 388 K for 0.3 h, the carbonate bands weakened slightly and a further band appeared at 2189 cm-l. On further heating at 473 K for 0.25 h, the carbonate bands shifted to 1480 and 1360 cm-' and weakened still further, whilst the band at 2189 cm-l grew stronger and the v(C0) band shifted to 2112 cm-l. On evacuation at 320 K the 2112 cm-1 band was removed, but the 2189 cm-l band was unchanged [fig.6(b)]. During the reaction, absorption bands due to N20 and C02 appeared in the spectrum of the gas phase. On heating in a CO+NO mixture for longer periods such that the gas phase became depleted in NO, the overall transmission of the disc decreased and the spectra at 2000-2200cm-l decreased in intensity considerably [fig. 6(c)].P. G . HARRISON AND E. W. THORNTON 261 1 The 21 89 cm-l band was formed readily on heating the catalyst discs in CO +NO mixtures [p(CO) > p(NO)] at ~ 4 7 0 K. In order to investigate the nature of the surface species responsible for this band, further experiments were carried out with isotopically substituted reactants. The spectra obtained are illustrated in fig.7(a)-( f ). The full lines are spectra recorded after heating in the appropriate reaction mixture at 463-483 K for 0.3 h and cooled without evacuation. During the reaction the CO invariably underwent a decrease in stretching frequency. Typical observed fre- quencies and shifts were : v ( C 0 a ) before Av during spectrum gas phase composition reactionlcm-1 reaction Icrn-1 7(a) normal mixture 2120 -8 7(b) (91 % 13C)CO+N0 207 1 -8 7(4 (95 % ~~N)NO+CO 2120 -6 7(d) (91 % 13C)C0+(95 % ISN)NO 2071 -7 7(e) (63 % 180)CO+N0 2070 -7 21 14) -41 The dashed lines show spectra recorded after removal of the gas phase at 320 K. The magnitudes of the AI3C, AISN and A13C+ 15N shifts for the 2189 cm-I band (table 1) are consistent with a surface isocyanate species by comparison with the shifts observed in other studies.When l*O-enriched CO was used, no isotopic shift was observed [fig. 7(e)], the band remaining sharp and symmetrical at 2189 cm-l. However, when 40 % l 8 0 - enriched NO was employed [fig. 7 ( f ) ] the band shifted to 2184 cm-l and developed TABLE 1 .-INFRARED ABSORPTION BANDS AND ISOTOPIC SHIFTS (cm-') FOR ISOCYANATE SPECIES ADSORBED ON CATALYST SURFACES AND FOR MODEL COhlPOUNDS catalyst or model compound Ru/AI,O~ Rh/A1903 PtIA1203 PtJA1203 CuO/Si02 d cu Ru/Si02 f Rh/A1203 g Si+NCO h Sn02 .0.55 CuO Si+CN Si+NC NCO- in KCI Hg(CN0)2 j vas(NC0) 2238 2264 2200 2380 21 80 2235 2315 21 89 221 8 2100 2195.8 2200 2220) 60 17 60 10 13 62 11 17 73 51 30 39 36 58.7 17.4 7.9 76.3 30 4 Shift to lower wavenumber compared with figure in column 2.b Ref. (2) and (3). Ref. (1). d Ref. (5). e Ref. (4). f Ref. (6). g Ref. (7). h B. A. Morrow and I. A. Cody, J.C.S. Farday I, 1975, 71, 1021. i V. Shettino and I. C. Hisotsuma, J. Phys. Chem., 1970, 52, 9. j Ref. (29).2612 TIN OXIDE SURFACES wavenumber /cm-I (4 (4 FIG. 7.Tdrared spectra recorded after exposure of SnOz .0.55CuO discs to isotopically-substituted CO+NO mixtures at 463-483 K for 20-30 min. Full = before evacuation, dotted = after evacua- tion at 320 K. The gas mixture compositionsand pressures were 16.6 % NO, 8.0 kNm-2 for (a)-@) and (f), 21.7 % NO, 6.12 kN m-2 for (e). (a) Normal abundances. (b) CO = 91 atom % 13C, normal NO. (c) NO = 95 atom % 15N, normal CO. (d) NO = 95 atom % 15N, CO = 91 atom % 13C. (e) CO = 60 atom % "0, normal NO.(f) NO = 41 atom % l*O, normal CO. a shoulder on the low wavenumber side. In order to investigate this band further, the spectral trace and background were digitised at 2cm-l intervals from 2100 to 2300 cm-l using a 10 x expanded abscissa scale. The resultant data were converted to an absorbance spectrum and analysed by adapting the data to a form suitable for input to a least-squares curve-fitting program. The envelope was thus resolved into %Po L I I 21 00 2200 2300 wavenumber /cm-l FIG. 8.-Pseudo-antisymmetric stretching bands of isotopically normal and1 80-substituted isocyanate species formed on the surface of SnOz .0.55CuO in the conditions as for fig. 7 ( f ) ; experimental points, fitted band profile and calculated components.P .G . HARRISON A N D E. W. THORNTON 2613 two overlapping Lorentzian line shapes centred at 2189 and 2172cm-l, with the low wavenumber band accounting for 43.5 % of the area under the curve (fig. 8). The final sum of squares was rather high, but attempts to fit the data to one band only resulted in a much worse fit, and inclusion of more than two bands resulted in unreal computed parameters (e.g. negative width or depth) for one of the bands. In different calculations the position of the higher wavenumber component was nearly constant (+2 cm-I), but both the position and intensity of the isotopically shifted band were sensitive to the estimated starting parameters used. Consequently the uncertainty in the band position is about &5 cm-l. The calculated half-height widths of the two bands [r+(2189) = 51 cm-l and r,(2172) = 42crn-l-J are similar to the half-height widths of the other isocyanate bands in fig.6 which all have in the range 35-55 cm-l. The isocyanate band was stable to evacuation up to ~ 4 4 0 K, but was removed entirely by evacuation at 480 K. It was also stable to oxygen and nitric oxide (1.3 kN m-2) at 320-440 K, but could be removed by either reagent at higher tempera- tures, and was removed entirely by exposure to water vapour (130 N m-,, 320 K, 0.5 h). Exposure of a disc exhibiting a preformed isotopically normal band at 2189 cm-l to 180-enriched 0, (63 % or 80-enriched NO (41 % at 383 K for 1 h did not shift the band or produce any detectable asymmetry. In order to investigate possible incorporation of oxygen from the oxide surface into the isocyanate, a sample disc was evacuated and then treated with 180-enriched O2 (63 % l80), all at 623 K.After subsequent reaction using isotopically normal NO and 80-enriched CO (60 % at 473 K, a single symmetrical band remained at 2189 cm-l on cooling and evacuation at 320 K. DISCUSSION Fuller and Warwick have studied the kinetics of the oxidation of carbon monoxide catalysed by coprecipitated SnO, + CuO gels both by oxygen Briefly, for near optimum composition catalysts (SnO, xCu0, 0.5 < x < 0.55) increasing calcination pretreatment temperature increased the catalytic activity for both reactions up to ~ 7 5 0 K, followed by a decrease in activity and surface area for further increase in calcination temperature.This decrease was attributed to the onset of CuO crystallisation as a separate phase above ~ 7 2 0 K, as confirmed by powder X-ray crystallography. The present infrared results show that a number of changes occur on calcination in vacuo. After evacuation at 320 K the (green) disc adsorbs CO with a stretching frequency more characteristic of CO adsorbed on a partially reduced CU' site than a Cu" species. Sites suf3ciently active to adsorb nitric oxide are only produced by evacuation above x400 K, and the oxide surface is essentially completely irreversibly dehydroxylated by evacuation above z 600 K, although some hydroxyl groups remain in the bulk or trapped in closed pores. The change in properties on heating above 720-50 K (either in vacuo or in oxygen) was also manifest in the present study, both as a loss in infrared transmission and, as indicated by lower temperature experiments, with the extent of C 0 2 adsorption (fig.3). The nature of the surface and bulk rearrangement accompanying CuO crystallisation is undoubtedly complex and out of the scope of the present study. Whether the large loss in infrared transmission is associated with the CuO crystallisation or with the conduction electrons of the partially reduced Sn02 support [tin-1 19m Mossbauer spectroscopy of a 740 K-evacuated disc showed no evidence for a tin(@ species, in contrast with the previously reported results for partially reduced (by carbon mon- oxide) SnO, where a Sn" species was detected]l* is uncertain. The surface changes and by nitric oxide.2614 TIN OXIDE SURFACES continuously with increasing calcination temperature in the range 300-700 K.In particular, the sites for carbonate formation during carbon dioxide adsorption (presumably oxide ions in the ion-exchanged Cu-containing surface layer) decrease in concentration and eventually disappear at about the CuO-cry stallisation temperature. This decrease and eventual loss of surface oxide ions may be a factor responsible for CuO crystallisation, the oxide ions helping to stabilise the surface structure at lower temperatures. The environment of the copper ions responsible for CO adsorption also changes [as reflected by the change in v(CO,,,)]. After evacuation, the copper ions are in a partially reduced state, but can be returned to the +2 oxidation state by oxygen treatment above x550 K.During the initial stages of catalysis of the CO +NO reaction by a 623 K evacuated disc, a carbonate species is produced on the surface as well as adsorbed carbon monoxide exhibiting a stretching frequency (2 120 cm-I) characteristic of adsorption on an oxidised copper (Cu2+) site. As the reaction proceeds, however, the carbonate is removed and the v(CO,,,) band shifts to lower frequency, characteristic of a partial reduction of the surface and a lowering in the oxidation state of the copper ions. This corresponds to the partial reduction reported during the kinetic studies of the CO + O2 and CO +NO reactions catalysed by this oxide. The present infrared spectroscopic results also show the existence of an isocyanate surface species, formed rapidly during the initial stages of the reaction at 670 K.Previous studies reporting this surface species have assumed that it was formed by initial dissociative chemi- sorption of NO : 21 +NO -N 1-N+]-0 where ] represents some general surface site, followed by chemisorption of CO : 1-N+CO 1-NCO. (2) The 1-N surface species also provide a route to N20 formation as a product : 1-N+NO 4]+N20. (3) Alternatively, mechanistic studies at 300 K on supported ruthenium and at 473- 573 K on supported rhodium ' both invoked more complex reaction schemes. The latter study still postulated the formation of the 1-N intermediate and subsequent reaction via routes (2) and (3) above. If reaction (2) were the route to isocyanate formation in the present study, then the oxygen atom of the -NCO species would be expected to be derived from the carbon monoxide reactant.This, however, cannot be the case since reaction with 80-enriched CO resulted in an isotopically normal surface species. In contrast, when the reaction was carried out with 180-enriched NO, the absorption band was shifted slightly and developed a shoulder on the low-wavenumber side. Total isotopic shift of the band is not realised due to the paucity of the l80-enrichment in the NO (41 atom % Resolution of this envelope into two components gave an isotopic shift A l s o = 17 cm-l, in reasonable ageement with the result reported for a -Si-NCO species on silica (table 1). Separate experiments failed to show any isotopic exchange of the isocyanate species either with 80-enriched O2 or 80- enriched NO.\ /P . G . HARRISON AND E . W. THORNTON 2615 These results imply that the mechanism described in eqn (1) and (2) above cannot be operating in this case. The most probable alternative route for the formation of surface isocyanate is one involving initial dissociative chemisorption of CO : 2]+co +]-c+]-0 (4) followed by reaction of the adsorbed carbon atom with nitric oxide : 1-C + NO -+ 1-CNO (5) and subsequent isomerisation of the resultant fulminate to the stable isocyanate structure, probably via a cyclic oxaziranyl-type intermediate : 1-CNO + +]+ This type of rearrangement is well known to occur for alkali metal fulminates as well as nitrile oxides and covalent organometallic fulminates on cautious heating2 The identity of the surface site involved in these reactions is not known, but it is reasonable to assume that it is a coordinatively unsaturated copper ion in the 0 or + 1 oxidation state.Dissociative chemisorption of CO on such a site at 470 K is not unreasonable especially as the lowering of the vads(CO) frequency below that of the gas phase implies some weakening of the C=O bond. The mechanisms proposed in eqn (4)-(6) do not mean that eqn (1) and (3) are not operating since these seem by far the most likely reactions leading to N20 formation? and N20 and C02 were detectable in the gas phase by infrared spectroscopy during the reaction. Kinetic results on this reaction have shown that this catalyst is very selective towards NO reduction to N2, but any comparison between the present infrared study and kinetic results should be made with caution, since the kinetic results refer mainly to a steady state partially reduced catalyst while the infrared results deal with the freshly activated or only slightly reduced catalyst.Since the steady state catalyst, like the thermally sintered oxide, exhibits negligible infrared transmission, the surface species present on the steady state catalyst are inaccessible to infrared spectroscopy. Thus, the mechanism found for isocyanate formation during the initial stages of this reaction is one which previous authors have apparently not considered and have not tested for. We thank the International Tin Research Institute and the S.R.C. for support in the form of a CASE Award (to E. W. T.). M.L. Unland, J. Phys. Chenz., 1973, 77, 1952. ’ M. L. Unland, J. Catalysis, 1973, 31, 459. M. L. Unland, Science, 1973, 179, 567. R. T. Rewick and H. Wise, J. Catalysis, 1975, 40, 301. J. W. London and A. T. Bell, J. Catalysis, 1973, 31, 96. M. F. Brown and R. D. Gonzalez, J. Catalysis, 1976, 44, 477. ’ H. Arai and H. Tominaga, J. Catalysis, 1976, 43, 131. * M. J. Fuller and M. E. Warwick, J. Catalysis, 1974, 34, 445. M. J. Fuller and M. E. Warwick, J. Catalysis, 1976, 42, 418. lo E. W. Thornton and P. G. Harrison, J.C.S. Faraday I, 1975, 71,461. 1-8 3261 6 TIN OXIDE SURFACES l 1 V. G. Amerikov and L. A. Kasalkina, Kinetics and Catalysis, 1971, 12, 137. l2 K. Nakamoto, Infrared Spectra of Inorganic and Coordination Compounds (Wiley, London, l3 L. H. Little, Infrared Spectra of Adsorbed Molecules (Butterworth, London, 1966). l4 R. P. Eischens, W. A. Pliskin and S. A. Francis, J. Chem. Phys., 1954, 22, 1786. l5 A. W. Smith and J. M. Quets, J. Catalysis, 1965, 4, 163. l6 A. W. Smith, J. Catalysis, 1965,4, 172. l7 J. Pritchard and M. L. Sims, Trans. Faraday SOC., 1970,66,427. l8 A. M. Bradshaw and J. Pritchard, Proc. Roy. SOC. A, 1970,316, 160. l9 H. G. Tomkins and R. G. Greenler, Surface Sci., 1971, 28, 194. 2o M. A. Chesters, J. Pritchard and M. L. Sims, Chem. Comm., 1970, 1454; and references 21 D. A. Seanor and C. H. Amberg, J. Chem. Phys., 1965,42,2967. 22 J. W. London and A. T. Bell, J. Catalysis, 1973, 31, 32. 23 J. 0. Alben, L. Yen and N. J. Farrier, J. Amer. Chem. SOC., 1970,92,4475. 24 A. N. Terenin and L. M. Roev, Acres du Deuxizme Congds International de CataZyse (Technip, 25 L. M. Roev and A. N. Terenin, Optika i Spektroskopia, 1959, 7,759. 26 L. M. Roev and A. V. Alekseev, Elementary Photoprucesses in Molecules, ed. B. S. Neoprent 27 A. N. Terenin and A. V. Alekseev, J. Catalysis, 1965,4,440. 28 M. Shelef and J. T. Kurmner, Chem. Eng. Progr. Symp., 1971, 67, 74. 29 W. Beck, Organometal. Chem. Rev. A, 1971, 7,159. 1970). therein. Paris, 1961), vol. 2 ; and references therein. (Consultants Bureau, New York, 1968). (PAPER 7/1924)

ISSN:0300-9599    

DOI:10.1039/F19787402604

出版商:RSC    

年代:1978

数据来源: RSC

摘要:

Electrical Force Between Two Permeable Planar Charged Surfaces in an Electrolyte Solution BY ANTONY J. DUNNING-/-, JAMES MINGINS,* BRIAN A. PETHICA~ AND PETER RICHMOND Unilever Research, Port Sunlight, Wirral, Merseyside L62 4XN Received 8th November, 1977 Numerical calculations of the pressure and electrical contribution to the Helmholtz interaction free energy for two uniformly charged sheets interacting across 1 : 1 electrolyte solution are presented. The model differs from that considered by most other workers in that the electrolyte is allowed to penetrate the charged surfaces. Further, in addition to the cases of constant charge and constant potential, the case where the surfaces contain reversibly ionised groups is considered. In order that the results may have some relevance to cell-membrane interactions, all the numerical calculations are done using parameters which represent approximately the conditions in physiological saline solution.The results for constant potential are identical to those obtained for impenetrable interacting charged sheets. The results for constant charge differ in an essential way. Specifically, the pressure in the limit of zero separation tends to a constant value whereas for impermeable surfaces it is, for small separations, inversely proportional to the square of the separation and thus diverges as the separation approaches zero. Consequently, the difference between constant potential and constant charge for this model is not as marked as differences obtained by other workers who studied impermeable surfaces.As one might anticipate, the results for reversibly ionised surfaces lie between the two extremes of constant charge and constant potential. Present calculations differ from those of Gingell in that the complete Poisson-Boltzmann equation is solved numerically. The electrical double layer force between impenetrable charged surfaces, of interest to colloid scientists, has been thoroughly studied in recent years. Constant wall charge and constant potential as well as the case where the surfaces reversibly ionise as the surfaces approach have all been considered using smeared potential and charge m~dels.l-~ Extension of the constant charge case to systems with discrete surface charges has been made by Richmond and a general thermodynamic treatment of interacting semi-infinite surfaces in equilibrium with any number of adsorbing species has been given by Hall.' The case where electrolyte may penetrate the surfaces, which is of considerable relevance for membrane biologists, has not been studied in such detail.To the best of our knowledge, the " two-sided " double layer was first treated by Verwey and Niessen in the context ofthe polar oillwater interface and later applied to semi-conductor systems such as the silver iodidelwater interface. ' 9 lo Using the linearised Debye-Hiickel approximation the electrical potential for partially penetrable surfaces or " fuzz " layers at an oil/electrolyte solution interface has been calculated by Gingell l1 who, in collaboration with Parsegian,12 later derived an analytic expression for the pressure between two completely penetrable charged sheets, again using the linearised Debye-Hiickel approximation.Other work on penetrable wall charges includes that of Bell, Levine and Pethica l3 who applied the Haydon and Taylor l4 model in the calculation of the electrical potential across a penetrable -t Present address : Directorate General Scientific and Technical Information and Information Management, Biltiment " Jean Monnet ", Plateau de Kirchberg, Luxembourg (Grand Duchy). $ Present address : Clarkson College, Potsdam, New York, 13676, USA. 261 7261 8 INTERACTION OF PERMEABLE CHARGED SURFACES surfactant layer at an oillwater interface. From the above account it can be seen that analysis of one important case is missing, namely the calculation of the pressure petween two penetrable charged plates containing reversibly ionised groups using the full Poisson-Boltzmann equation.In this paper we treat this case and compare the results with those from constant charge and constant potential assumptions. THEORETICAL We consider the model system illustrated in fig. 1 which consists of two similar plane parallel thin uniformly charged sheets (surface charge density a) surrounded by an electrolyte, taken here as 1 : 1 valent but which could be any symmetrical electrolyte. The electrostatic potential, $, is related to the charge density p(x) via Poisson’s equation d2* 4n dx2 & - - P W -= where E is the dielectric constant of the electrolyte solution. We make the well tried assumption that self consistency can be made using the Boltzmann distribution P ( 4 = qn b P (- Pq$) - exp (Pq$)l (2) where n is the charge number density in bulk solution (where $ = 0), q is the electronic charge and P = (kT)-l.Tis the absolute temperature and k is Boltzmann’s constant. Thus we obtain the Poisson-Boltzmann equation It is straightforward to show that for our system, the solution is given in terms of the transcendental equations du = lcz s” uo J ~ ( c o s ~ u -cash u,) (4) and A = $[(cash U I - cosh u0)* + (cosh U Z - l)+] = 2[(sinh2 uJ2 - sinh2 u0/2)3 + sinh u,/2] (5) It can be shown by a number of different methods ‘ 9 exerted on the charged plates is that the electrical pressure, p , p(1) = 2n kT[cosh Pq$(O) - 13. (7) The electric free energy of interaction A , per unit area may then be obtained by integrat- ing eqn (7) in the usual manner.Thus A,(Z) = 2 Jl p(x) dx = 1’ p ( x ) dx. - IA . J . DUNNING, J . MINGINS, B . A. PETHICA AND P . RICHMOND 2619 The expression (7) for the pressure between the charged sheets is a consequence of our use of the Poisson-Boltzmann equation. However, depending upon whether we assume constant potential, constant charge or some criterion which takes into account ionisable groups on the surfaces, different values of the mid-plane potential at a given separation, 21, arise for the same number density of ionisable groups on the surface and the same electrolyte concentration. These lead to different pressure against distance curves in the three situations. FIG. 1 .-Diagram representing two negatively charged plane parallel membranes, and the electrical potential surrounding them.#(I) is the membrane potential, u is the surface charge density on the membrane and #(O) is the mid-plane potential. NUMERICAL PROCEDURE CONSTANT POTENTIAL [$(2)] If the surface potential is kept fixed, eqn (4) can be used to calculate the plate separation, 21, for a chosen set of values of the mid plane potential. The charge on the surfaces is then given by eqn (5). CONSTANT CHARGE (a) If the surface charge is fixed then the surface potential will vary as the plate separation varies. Now for a chosen value of cr and a chosen set of values of $(O), the value of $(I) is given by eqn (5). Using eqn (4) with the appropriate values of $(O) and $(I) will yield 1 and enable the [$(O), I ] curve to be plotted.IONISABLE SURFACES [BOTH $ ( I ) AND VARYING] In the case where the surfaces or model membranes are composed of ionisable groups, an extra condition must be introduced to account for the equilibrium between an ionisable molecule and its ionised state. We follow previous workers and assume for simplicity the reaction at the surface takes the form HA + H++A- (9)2620 INTER4CTION OF PERMEABLE CHARGED SURFACES where HA and A- represent neutral and negative surface sites respectively. Their relative concentrations are determined by the concentration of the H+ ion at the surface [H:], i.e., where K is the effective acid ionisation constant for equilibrium. If a is the fraction of dissociated molecules with charge q, and S is the total number of molecules per unit area then 101 = aqs.25 20 5 0 0.1 0.6 1.1 f/nm FIG. 2.-Mid-plane potential #(O) plotted against I for constant charge (C) and constant potential (P). Curves A, B and D are for a at equilibrium varying with separation of the membranes for am = 0.1 (A), am = 0.5 (B) and am = 0.9 (D). Introducing the familiar notation p . . . = - log,, . . . and using Pe . . . for the loge . . . equivalent we may write eqn (10) in the form Since the ions in solution obey the Boltzmann law, clearly we also have PeHs = PeHbulk-PqlCI(O*A . J . DUNNING, J . MINGINS, B . A . PETHICA AND P . RICHMOND 2621 From eqn (12) and (1 3) we have a = [exp - ~ A - - P ~ ~ ( O l ~ + 13-' (14) A = Peffbulk- PeK* (15) where When the sheets are infinitely far apart @(O) = 0 and from eqn ( 5 ) we have A, = 4 sinh (u,/2).I 1 1 0 K-1 1.5 3 Ilnm FIG. 3.-Comparison of normalised degrees of ionisation (cC/a,) for constant charge (C) and constant potential (P) for a varying at equilibrium with a , = 0.1 (A), a , = 0.5(B) and a, = 0.9 (D). Now if a, is the corresponding degree of dissociation we have from eqn (14) and (16) that A = l n (1 - :;J+um and 2n12, 2nA Ica, ica s = la,l/a,q = - = -. From eqn (9, (14), (16) and (18) it follows that We may now proceed as follows : for a set of values a, and u, compute S, A and A, using eqn (16)-(18). Set ul and obtain a from eqn (14). The values of a and S may be used to obtain a from eqa (18). Using the same value of ul, uo may be obtained from eqn (19). Finally uo and uz may be used to obtain I from eqn (4).2622 INTERACTION OF PERMEABLE CHARGED SURFACES For each of the three situations we can, therefore, obtain curves of p(Z) and A,(Z) as a function of A.A comparison of the results obtained is given in the next section. RESULTS AND DISCUSSION In comparing the three cases outlined above we have chosen parameters which represent approximately the conditions in physiological saline solutions. So, in an idealised sense, we are answering questions about the electrical force and energy 1.25 1.0 7 0.75 E 2 c, 1 X \ v) !i 0.5 0.25 0 0.5 K-1 1.0 1.5 Zlnm FIG. 4.-Comparison of the electrical pressure, p , between the charged surfaces as a function of the separation distance, I, for constant charge (C) and constant potential Cp) for a varying at equilibrium with am = 0.5 (B).The dotted lines show the variation of the distances at which selected values of the mid-plane potential ($,,) are obtained for the three cases. relationships in biological cells. The study of a more detailed model in which the cell walls have thickness, structure and discrete ionisable molecules on each side of the membrane is a more complex subject worthy of future investigation but for the present purpose of examining the differences between the three types of assumption we are probably justified in our choice of a simple model. Thus the 1 : 1 valent electrolyte concentration is 0.145 mol dm-3, T = 293.16 K, E = 78.54 and the surface potential [$(a)] when I is infinite is taken as - 18 mV.A . J . DUNNING, J . MINGINS, B .A . PETHICA AND P . RICHMOND 2623 The results of calculating the mid-plane distances, the wall potentials and the degree of ionisation as a function of mid-plane potential for the cases of (i) constant charge, (ii) constant potential and (iii) ionisable membranes with am = 0.1, 0.5 and 0.9 are given in fig. 2 and fig. 3, respectively. When the sheets are a distance k--l apart there is a 12 % increase of the mid-plane potential on going from the constant potential to the constant charge case. The case of ionisable membranes lies between these two situations. The differences between the various cases increases quite markedly as the separation, I, decreases. We note here that for our system k--l = 0.8 nm. The corresponding pressures and free energies of interaction are shown in fig. 4 and 5.5 0 0.5 K-1 0.9 FIG. 5.-Electrical contribution to the Helmholtz free energy of interaction A&) as a function of the separation distance, I, for constant charge (C), for a varying at equilibrium with am = 0.1 (A), am = 0.5 (B). The dotted lines show the variation of the distance at which selected values of the mid-plane potential ( #o) are obtained for the three cases. It is of interest to compare the results for our model with those obtained by other workers who considered the interaction between impermeable dielectric half spaces. For the case of constant potential, both models yield identical results. If we fix the membrane potential then from eqn (4) we see that uo(Z) and hence the pressure will take the same values as for the usual model.The surface charge obtained from eqn (5) will, however, differ from the usual situation where it is determined by the relation II = 2[sinh2 ul/2-sinh2 u0/2]*. (20)2624 INTERACTION OF PERMEABLE CHARGED SURFACES The case of constant charge, however, differs in an essential way. As the plate separation, 2, becomes small, clearly uo --* ul and in the usual case it follows by expanding eqn (4) and (20) that u1 -uo N (g sinh uo 2 and 3, 2c 2[-?) sinh 14.1'. From eqn (21) and (22) we obtain sinh u0 2: A/d, i.e. cosli uo N [I + ( A / K ~ ) ~ ] * and, therefore, p 2: 2nkT{[l+(Ay]i-l} + 2nkT -. a 1+0 K l Thus the pressure in the limit of small separations is divergent. For our system we obtain from eqn (9, again by expanding in powers of (ul - u,) and retaining leading terms, that Therefore Thus the pressure in the limit of small separations tends to a limiting value.This is physically very reasonable, and indeed one might imagine for some fluid-like mem- branes that when the separation is small the ions may penetrate the membrane and the electrical pressure will saturate as indicated above. 3, N, 2 sinh u0/2. (24) p 21 nkTA2. (25) B. V. Derjaguin and L. D. Landau, Actu Physicochim., 1941, 14, 633. E. J. W. Verwey and J. Th. G. Overbeek, Theory of the Stability ofLyophobic Colloids (Elsevier, Amsterdam, 1348). B. Ninham and V. A. Parsegian, J. Theor. Biol., 1971, 31, 405. G. M. Bell and G. C. Peterson, J. Colloid Interface Sci., 1972, 41, 542. D. Chan, J. W. Perram, L. White and T. W. Healy, J.C.S. Furuday 11, 1975, 71, 1046. P. Richmond, J.C.S. Furuduy 11, 1974, 70, 1061 ; 1975, 71, 1154. ' D. G. Hall, J.C.S. Furuduy II, 1977, 73, 101. * E. J. W. Verwey and K. F. Niessen, Phil. Mag., 1939, 28,435. T. B. Grimley and N. F. Mott, Disc. Furaduy SOC., 1947,1,3 ; T. B. Grimley, Proc. Roy. SOC. A , 1950, 201,40. lo E. P. Honig, Truns. Furuduy Soc., 1969, 65,2248. D. Gingell, J. Theor. Biol., 1967, 17, 451 ; 1967, 19, 340. l2 V. A. Parsegian and D. Gingell, Biophys. J., 1972,12,1192. l3 S. Levine, G. M. Bell and B. A. Pethica, J. Chem. Phys., 1964, 40, 2304. l4 D. A. Haydon and F. H. Taylor, Phil. Truns. A, 1960, 252, 225. (PAPER 7/1967)

ISSN:0300-9599    

DOI:10.1039/F19787402617

出版商:RSC    

年代:1978

数据来源: RSC

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Kinetics and Mechanism of Fast Metal-Ligand Substitution Processes in Aqueous and Micellas Solutions Studied by means of a Dye Laser Photochemical Relaxation Technique BY BRIAN H. ROBINSON" AND NEAL C . WHITE Chemical Laboratory, University of Kent at Canterbury, Canterbury, Kent CT2 7NH Received 4th January, 1978 Rapid metal-ligand substitution reactions have been studied in aqueous and sodium lauryl sulphate (SLS) micellar solutions by means of a dye-laser induced photochemical relaxation method. Perturbation of the system involving Ni2+(aq) and the bidentate nitrogen ligand PADA produces two relaxation effects in both aqueous and micellar solutions. The faster relaxation process is concentration independent and is identified with a ring-closure reaction to reform the initial Ni(PADA)2+ complex.Rate constants and activation energies for this process are found to be similar in both media. The slower relaxation time is associated with the recombination reaction between Ni2+(aq) and PADA, for which kf298*2 in water is found to be 1.4~ lo3 dm3 mol-' s-', which is in good agreement with previous measurements. In the surfactant solutions, the rate of complexation is greatly enhanced close to the c.m.c. followed by a decrease as the concentration of SLS is further increased. This is due to the increased local concentration of Ni2+ at the micelle surface as the c.m.c. is approached. It is shown that it may not be valid to make the assumption kRC > k-2 in the kinetic scheme of the reaction. However, the effect on kf is small, such that kf is still dominated by the rate of loss of water from the metal ion.The faster reaction between Zn2+(aq) and PADA was also studied, but only one concentration dependent relaxation time was observed. This was consistent with the complex re-formation reaction and kZgsm2 was found to be 6.4 x lo6 dm3 mol-I s-l. The most frequently used relaxation technique for the study of fast reactions involving ground-state species in solution is temperature-jump relaxation spectro- scopy. The temperature-jump is usually generated by means of a capacitor discharge, a microwave pulse or a laser pu1se.l In non-aqueous solvents, an adiabatic expansion (in a pressure-jump apparatus) can produce a rapid temperature drop in 50-100 ,us.2 The temperature- and pressure-jump methods perturb the equilibrium position of a chemical reaction indirectly through a perturbation of the reaction medium, so that relatively high input energies are required. Furthermore, the resulting relaxation amplitude depends on the position of the equilibrium.However, the extent of a chemical reaction initiated by direct light absorption into a reacting species is dependent on the quantum yield for photodissociation and should therefore in favourable cases require less input energy. Also the equilibrium condition and related thermodynamic constraints (e.g. non-zero AHc for a temperature-jump perturbation) are eliminated. Since the light is absorbed by a specific reagent in the chemical system, the technique can be used to generate high concentrations of selected intermediates.This cannot be achieved by other perturbation techniques, which perturb the solvent, so the technique may be helpful for the analysis of complex reaction schemes, as is indicated in this paper. However, it is not easy to predict in advance whether a 26252626 FA S T M ETA L -L I G A ND S U B S T I T U T I 0 N given reaction can be studied by the photochemical method, since the relative probabilities of the various de-excitation pathways cannot easily be quantified. Early examples of a photochemical perturbation of a chemical equilibrium used a flash-lamp for e~citation,~ but recently several papers have referred to the use of a ruby or neodymium laser as the excitation source. An early exaaple using a Q- switched neodymium laser for excitation of vibrational modes was reported by Goodall and Greenh~w.~ More recent examples 5 - 9 involve studies of the structural inter- conversions of Ni" complexes in water and acetonitrile, processes which occur in the microsecond time range.have discussed the factors which influence the amplitude of the photochemical perturbation, and have compared photochemical and temperature-jump methods for studies on the same system. HasinofFIO employed a ruby laser for photochemical excitation in a study of the reaction between carbon monoxide and haemoglobin as a function of pressure, and a variation of the method, using indirect photochemical excitation, hax been reported by Hubbard et aZ.,I1 for proton transfer reactions. Meyer et al.24 have recently demonstrated the applicability of the flash perturbation relaxation technique to electron-transfer reactions which have essentially gone to completion.Several other instruments and applications have been discussed in a recent publication.' In this paper, we report results obtained (by means of a 3 p s pulse from a Rhodamine-6G dye laser) on the photochemical dissociation and subsequent recom- bination of metal complexes. The systems studied demonstrate the applicability of the method and give some indication of the origin of the relaxation perturbation. The ligand used was pyridine-2-azo-p-dimethylaniline (PADA) (fig. l), and the kinetics of complexation have been studied in bulk water with both Ni2+(aq) and Zn2+(aq). In addition, the influence of micelles formed by sodium lauryl sulphate (SLS) on the rate of reaction between Ni2+(aq) and PADA has been investigated.The kinetics of the association reaction between Ni2+(aq) and PADA have been studied previously by both stopped-flow and joule-heating temperature-jump methods. Results obtained by the laser method can therefore be compared with those obtained by other well-established techniques. The reaction between Zn2+(aq) and PADA is much faster, and has been studied only by the temperature-jump method. l4 The complexation reaction between a metal ion and a bidentate ligand involves several discrete steps : as shown in Scheme 1, step 1 involves the rapid (diffusion- controlled) formation of an outer-sphere complex ; step 2 represents monodentate complex formation, for which the forward rate constant can be identified with that for water exchange (kX) as measured by n.m.r.methods. Since PADA is a bidentate ligand, ring closure (RC) is required to form the final product, and this is usually assumed to be rapid compared with dissociation of the monodentate complex (i.e. kRc % k-2). This mechanism is generally referred to as the Eigen-Wilkins mechanism. Sutin and Creutz SCHEME 1 N KO, N ke, f i (H20)5Ni2+(H,0) + ) f (H20)5Ni2+(Hz0) ) + (H20),Ni2+N N+ H20 N N k - 2 (2) monodentate complex (3) N N k - 3 11 kRc (H20)4Ni2+ ) + H 2 0 bidentate complexB . H . ROBINSON AND N . C . WHITE 2627 Provided that [Ni2+] $ [PADA], and KO, < [Ni2+]-l, we have (for metal complex formation induced by mixing the separate reagents in a stopped-flow experiment) the complete kinetic expression : When kRc & k--3, the ratio of monodentate to bidentate complex at equilibrium is insignificant, and it is therefore difficult to determine kRc by conventional relaxation methods. However, as will be shown later, this is possible using the laser technique. FIG.1 .-Pyridine-2-azo-p-dimethylaniline (PAD A). The reaction between Ni2+ and PADA in the presence of SLS micelles has been studied previously by the stopped-flow method.16 Results reported in this paper using the laser technique extend the range of measurments to the region around the critical micelle concentration (c.rn.c.) of the detergent, where the reaction is too fast for the st opped-flow method. EXPERIMENTAL MATERIALS AND METHODS Purification of SLS and PADA and buffering of the solutions were carried out as described previously.16 Ni(N03)2 6H20 and ZII(NO~)~ 6H20 were both B.D.H.AnalaR grade, and were used without further purification. The stopped-flow apparatus used in the pH- jump studies has been described previously. l7 An Electrophotonics SUA-9 dye laser coupled to a spectrophotometric detection arrangement was used for perturbation of the chemical systems under investigation. A block diagram of the apparatus is shown in fig. 2. Untuned emission was obtained at 580nm (f 5 nm) or 470 nm (f 5 nm) using the dyes Rhodamine-GG or Esculin respectively in " Ultra " methanol (Hopkin and Williams). The laser beam diameter was 5 mm and the pulse-width at half-height was found to be ~3 ys. The energy output was measured by an ITL power meter (Model LOMB 102/K101), and laser energies in the region of 100 mJ were generally employed, The sample cell was a fluorescence-type rectangular cell (Hellma 6040F) located in a thermostatted holder such that the detection path length is 10 mm and the laser path length is 4 mm.Temperature control was to k0.05 K and the temperature in the cell was measured directly by a Comark electronic thermometer (1604 Cr/Al). The monitoring light source was a 50 W pre-focus tungsten lamp (Wotan A1/17). The output was adjusted to focus at the input end of a 3 mm diameter quartz light-guide (Schott- Jena). A shutter and heat filter are located inside the lamp housing. There is also provision for inserting an optical filter in front of the light-guide to prevent exposure of the sampIe to the full spectral output of the lamp The other end of the light guide is attached to the sample cell-holder, and a second light-guide transmits the emergent beam to a Bausch and Lomb high intensity u.v./visible grating monochromator.The photomultiplier (IP21) housing and control circuitry (for amplification and time constant control) were fitted directly to the exit side of the monochromator. The exit slit-width of the monochromator was set to 1.3 mm, which gives a spectral bandwidth of %5 nm. The narrow bandwidth was required to eliminate scattered laser light when it was necessary to work at wavelengths2628 F A S T MET A L-LI G A ND SUB ST IT U T I 0 N close to that of the laser output. The fastest relaxation which can be studied is limited by the laser pulse-width and is ~5 p.The transients were recorded on a Tektronix 549 storage oscilloscope for photography and trace-matching or on a Datalab 905 transient recorder linked to a Telequipment D65 monitoring oscilloscope. Stored digitised data from the transient recorder were transferred to paper tape for subsequent computer analysis. Errors in relaxation times and amplitudes are standard deviations of approximately six transients. ORIGIN OF THE PHOTOCHEMICAL PERTURBATION The dye laser pulse induces photochemical dissociation of the metal complex, a process which is well known in inorganic photochernistry.l8 The reaction scheme may be described generally as follows [where S is a solvent (water) molecule and L-L is a bidentate ligand] : 2s + SS-SSML L monoden tate k - 2 k i ' SsMSL L n L L+SGM, B - 1 outer-sphere complex The excited-state complex S4& ') can in general be deactivated by radiationless decay L (thermal deactivation), by fluorescence or by " chemical reaction ", i.e. photosubstitution processes (A) or (B), where (A) respresents concerted loss of both ligand groups from the metal ion.When kRC >> k-z, the outer-sphere complex can only be produced by (A). When kRC - k--2 the outer-sphere complex can also be produced via (B) and dissociation of the monodentate complex. After perturbation the ground state species SsML L or S5MSL L must be obtained rapidly to ensure that the subsequent recombination reactions are ground state processes. Any S5MSL L formed will rapidly produce SsM and L L by diffusion apart of the reagents, since k-l/k,, - lo5 for reactions involving Ni2f.Therefore as a result of the photochemical perturbation, we can predict that two relaxation times might be observed in m A nB . H . ROBINSON A N D N. C. WHITE 2629 a favourable case when the various complexes have different spectra. A faster relaxation would be associated with ring-closure and a slower one (identical to that measured by stopped-flow and temperature-jump techniques) would be identified with recombination of free metal ion with ligand. Therefore, as a result of the photochemical perturbation, we would hope to be able to measwe ~ R C , ( k ~ c k ~ ~ x kl/k-l)/(k-z+kRC) and k-2k-3@-2+k~c). The relaxation amplitudes of the two processes are determined by: (i) the optical properties of the absorber in the system, (ii) the quantum efficiencies 41 + 42 of the photo- substitution processes (A) and (B), (iii) the ratio of the rate constants k ~ ~ / k - ~ .monochromator FIG. 2.-Block diagram of the laser apparatus with spectrophotometric detection. From a semi-quantitative consideration of our results, it is clear that the quantum efficiency of process (B) is greater than for process (A). Following Sutin et aZ.,’ we can calculate the overall quantum efficiency, +l+q52, to be of the order of 1 % for the incident photon intensity (as used) of 7x 10l6 photons cm-2 and &580 nm of (H20)4Ni2+N) of 2x lo4 dm3 mol-’ cmn-’. A precise treatment is difficult because of the uncertainties regarding the extinction coefficient of the monodentate complex.N RESULTS AND DISCUSSION Under most conditions two relaxation times are observed, as predicted in the preceding section. All individual transients were found (within experimental error) to be exponential decay curves when the systems were studied under pseudo-first-order conditions (i.e. [Niz+] $ [PADA]). Following perturbation, the absorbance of the solution always returned to its initial value; this indicates that there is no overall photodecomposition of the complex as a result of perturbation. When PADA alone was irradiated with laser light of 580 nm, no relaxation was observed. However, when laser light of 470nm was employed (the free ligand absorbs strongly at this wavelength), a relaxation was observed, decaying in a time ( N 100 ,us) which was independent of ligand concentration and temperature.2630 FAST METAL-LIGAND SUBSTITUTION THE SYSTEM Ni2++PADA I N WATER The Ni2+ + PADA system, perturbed at 580 nm, gave rise to two relaxation times, well separated on the time axis, which are independent of the detection wavelength over the spectral absorption range.The faster relaxation time (-40 ps) is indepen- dent of Ni2+ and PADA concentrations, but is dependent on temperature. The slower relaxation time (10-1000ms) is dependent (when [Nil, 9 [PADA].) on the concentration of Ni2+ (but not PADA) and temperature [subscript T = total (initial) concentration]. C wavelength Inm FIG. 3.-Spectra of (a) PADA, (b) NiPADA2+ and (c) PADAH+. The visible absorption spectra of PADA and Ni(PADA)2+ are shown in fig.3. The dependence of the amplitudes of the two separate transients on detection wave- length (fig. 4) shows that the species which are involved in the two relaxation processes are different, isosbestic points being located at 490 nm (slow z) and 535 nm (fast z). Zero amplitude is observed for the slow relaxation time at a wavelength which corresponds to the isosbestic point between PADA and Ni(PADA)2+ (fig. 3). The magnitude of the relaxation amplitude, however, does not depend on the position of equilibrium, but is directly proportional to the concentration of Ni(PADA)2f complex and the incident laser energy (fig. 5).B . H. ROBINSON AND N. C . WHITE 263 1 420 440 460 480 500 520 5 4 0 560 wavelength/nm FIG. 4.-Amplitude wavelength dependence for the fast ( x - X ) and slow (0-0) relaxation processes. pH = 6.5, T = 298.2 K.[Ni2+]~ = 1 X mol dm-3; [PADAIT = 2 X mol dmd3. In bulk water, at 298.2 K, the fast relaxation time (T~) has a value of 35 +4 ps corresponding to a first order rate constant of 2.9k0.3 x lo4 s-l. This fast process will be identified with ring-closure [step (3) of Scheme 11, following photo- chemical generation of the monodentate complex. The activation energy (E,)f is 8 - 6- 4 - 2 - /-/-- / X Y X o e I I I I I I 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0 absorbance (A) FIG. 5.-Relaxation amplitude as a function of NiPADA2+ absorbance. pH = 6.5, T = 298.2 K, 1 = 550 nm. ( x - X ) fast relaxation process, (0-0) slow relaxation process.2632 FAST METAL-LIGAND SUBSTITUTION 40 Ifr 4 kJ mol-l, so that the process is clearly different from that measured in solutions containing only PADA, for which E, 21 0 for laser excitation at 470 nm.The slow relaxation time (7,) shows a linear dependence on Ni2+ concentration as shown in fig. 6 and the data fit the simple rate expression : where and 110 100” 90 80 70 - - - - * ?, 60- 0” A? 50- 4 0 30 20 - - - (7, ’) = k(obs)s = kf[Ni2+] + kb ” 1 2 3 4 5 6 7 8 9 [Ni2+]/10-2 mol dm-3 FIG. 6.-Plot of k(&s)s against [Ni2+] for the system Ni2++PADA-!-Hz0. (-0-O), T = 308.2 K ; (- x-x), T = 298.2 K ; (-n-n), T = 288.2 K ; pH = 6.5. [PADAIT = 2 x r n ~ l d m - ~ . Excellent agreement for k,2 8.2 (= 1.4 x lo3 dm3 mol-1 s-l) and kE9 8 * 2 is obtained on comparison with other fast reaction techniques (table 1). The value of (&JS is 57 & 5 kJ mol-l, again in excellent agreement with previous determinations.13 This agreement, with the additional evidence from the isosbestic point for the slow process, allows us to conclude that, following photochemical excitation, the complex must dissociate to give separated Ni2+(aq) and ligand.B .13. ROBINSON AND N. C . WHITE 2633 THE SYSTEM Zn2++PADA I N WATER When Zn2+ is substituted for Ni2+, only one relaxation proportional to the Zn2+ concentration is observed. The complexation rate constant, kf, is found to be 6.4 x lo6 dm3 mol-1 s-l at 298.2 K (fig. 7) which is again in excellent agreement with the value obtained by the joule-heating temperature jump method, l4 confirming that the rate of reaction is not significantly ionic strength dependent.'i- [Zn2+]/10-3 rnol dm-3 FIG. 7.-Plot of kobs against [Zn2+] for the system Zn2++PADA+Hz0. (T = 298.2 K), pH = 6.5. [PADAIT = 2 X mol dm-3. INFLUENCE OF MICELLAR SLS ON THE REACTION BETWEEN Ni2+ AND PADA From previous stopped-flow studies,16 it has been found that the presence of micellar sodium lauryl sulphate (SLS) has a pronounced effect on the association rate constant for the reaction between Ni2+ and PADA. It was shown that the reaction occurs at the negatively-charged surface of the anionic micelle, since PADA is a relatively hydrophobic ligand and hence is located in the surface region of the micelle. Ni2+ is also readily adsorbed onto the surface.16 Using the laser perturba- tion method, two relaxations are again observed which have qualitatively the same features as those for the bulk reaction.The fast relaxation time is independent of Ni2+ and PADA concentrations and the isosbestic point is at 535 nm. The fast relaxation time (7;) is loo+ 10 ps at 298.2 K corresponding to a first order rate process with k; - 104s-', which may be again identified with kRc for the complexation reaction on the micelle surface. This is slightly slower than the value (3 x lo4 s-l) measured in the absence of SLS. The slow relaxation time, corresponding to (kLbs)C1 is independent of the monitoring wavelength, and the wavelength corresponding to zero amplitude is at 485 nm. 7; depends linearly on the nickel ion concentration, and shows a dramatic dependence on the SLS concentration, especially in the region of the critical micelle concentration (c.m.c.).I N WATER2634 FAST METAL-LIGAND SUBSTITUTION A plot of log (kLb& against log [SLS] is shown in fig. 8, and a plot of (k&& against ([Ni2+]/[SLS], - c.m.c.) for concentrations of SLS above twice the c.m.c. is shown in fig. 9. ([SLS], is the weighed-in concentration of SLS). The results obtained previously for (k;& by the stopped-flow method l6 under the same final conditions are also plotted in fig. 8 for comparison. Excellent agreement is again obtained by the two methods. [SLS]/mol dm-3 FIG. 8.-Plot of log,, (k&& against log,, [SLS]. (-0-0), T = 308.2 K ; (- x - x), T = 298.2 K ; (-A-A), T = 288.2 K ; (-@-a), results obtained by stopped-flow method at 298.2 K ; pH = 7.7. [Ni’f]~ = 1 X mol drnd3.mol dm-3; [PADAIT = 2 X { [NiZf]/([SLS], - ~.m.c.)}/lO-~ FIG. 9.-Plot of (k&bs)s against { [Ni2+]/([SLSl0 -~.m.c.)>/lO-~ for sodium lauryl sulphate (pH = 7.7).B. H. ROBINSON AND N. C. WHITE 2635 The rate increases sharply from the bulk value when the concentration of SLS exceeds the c.m.c. (-7.5 x mol d n r 3 at 298.2 K). Close to but above the c.m.c., (k&& is approximately constant but at higher concentrations of micelle it falls gradually to values close to that measured in bulk water. Above the c.m.c., the reaction occurs on the micelle surface, and a simple argu- ment suggests that the association rate constant should be inversely proportional to the concentration of micellar binding sites at the surface, and first-order in the excess reagent.This behaviour is confirmed by fig. 9. Detailed considerations of the mechanism of the reaction are given in a previous paper.16 TABLE 1 .--SUMMARY OF THE KINETIC DATA kr/103 dm3 mol-1 s - ' kb/S-l kill04 s-1 (E&/kJ mol-1 (E&/kJ mol-1 Ni2+ 1.3 a 0.1 a 55+4 a S.F. Ni2+/SLS 1.1" 52.3+4 S.F. Zn2+ 6.6 x lo3 2.2 x lo4 35.8 T. J. 1.4 0.1 2.9 57+5 40k4 L. 1.1" 1.0 55+4 40+4 L. 6 . 4 ~ 103 c 2.1 x 104 c L. a Ref. (16), b ref. (16), C this work, dref. (14), * calculated from eqn (15) of ref. (16). S.F. = stopped-flow. T.J. = temperature-jump. L = laser photochemical perturbation. Activation energies are (E& - 40 +4 kJ mol-l and (E& - 55 +4 kJ mol-I. These measurements are made at a constant SLS concentration of lo-' mol dm-3 assuming a temperature independent c.m.c.(Fig. 8 shows the actual effect of temperature on the kinetics in the c.m.c. region). The value of (I?;), is very similar to that measured in bulk solution (table 1) suggesting that the energetics of the reaction on the micelle surface are not significantly changed compared with the bulk aqueous phase. In particular the enthalpies of activation for water-exchange and ring closure appear to be little affected by the change in environment. The results further confirm that the dependence of zs on the concentration of micelles is due to an entropic effect involving the concentrative effect of the micelle surface on the reagents. DISSOCIATION KINETICS INVOLVING THE NiPADA2+ COMPLEX A T 298.2 K Using the stopped-flow technique in a pH-jump mode, it is possible to determine k-3 independently, using the procedure originally developed by Basolo et al.' Measurements were made over a range of final pH values in the region 3.5-5.5.The value of k-3 obtained (0.3 s-l) was found to be independent of pH. Values of kb (= k-2k-3/k-2+kRC) can either be measured as the intercept on the plot of (ko& against [Ni2+] (fig. 6) or by reacting the NiPADA2+ complex with Hg2+. In both cases a value of k, of 0.1 s-I was obtained. By means of the value of kRc obtained from the laser perturbation, it is possible using eqn (2) and (3) to calculate all the postulated rate and equilibrium constants for complexation. We find, if outer-sphere complex formation is assumed to be diffusion-controlled : K,,,k,, = 1.5 x lo3 dm3 rnd-1 s-1 kRC = 2.9 x 104 s-' k-3 = 0.3 S-' k-, = 1.4 x lo4 s-' K2 = 2.1 K3 = 0.97 x 10'.agreement with that obtained by spectrophotometry l 3 of 1.2 x lo4 dm3 mol-'. The kinetic value of K( = Kl x K2 x K3 = 1.02 x lo4 dm3 mol-I) is in excellent2636 FAST METAL-LIGAND SUBSTITUTION It seems from our analysis that the assumption k-2 < kRc should not be made. Ring closure can in some instances be rate limiting, as was observed by Hoffinann 21 in studies of nickel malonate complex formation. However, the effect on kf of the term kRC/(k-2+kRC) is small, such that kf is still dominated by the rate of water loss from the metal ion. k,, obtained from our kinetic measurements will be of the same order as that obtained from n.m.1. measurements (3 x lo4 s - ~ ) . ~ * However, it should be noted that published data 22 on the stability constant (K') for formation of the monopyridine nickel(I1) complex (which would be expected to be a good model for monodentate complex formation with PADA)23 suggests that k e x / L 2 - 800 since the value of K' measured experimentally was 80 dm3 mol-l.We thank the S.R.C. for the provision of apparatus associated with this work, and for a postdoctoral research assistantship (N. C. W.). We especially thank Prof. E. F. Caldin for guidance and helpful discussions. I E. F. Caldin, Chem. in Britain, 1975, 11, 4. H.-J. Buschmann, W. Knoche, R. A. Day and B. H. Robinson, J.C.S. Faraday I, 1977,73,675. e.g. R. L. Strong and J. E. Willard, J. Amer. Chem. SOC., 1957, 79, 2098 ; C. Breitschwerdt and A. Weller, 2. Electrochem., 1960, 64, 395 ; Q. H. Gibson, Biochem. J., 1959, 71, 243. D. M. Goodall and R. C . Greenhow, Chem. Phys. Letters, 1971,9, 583. K. J. Ivin, R. Jameson and J. J. McGarvey, J. Amer. Chem. SOC., 1972,94,1763. H. Hirohara, K. J. Ivin, J. J. McGarvey and J. Wilson, J. Amer. Chem. SOC., 1974, 96,4435. ' J. J. McGarvey and J. Wilson, J. Amer. Chem. SOC., 1975, 97, 2531. * L. Campbell and J. J. McGarvey, J.C.S. Chem. Comm., 1976, 746. C. Creutz and N. Sutin, J. Amer. Chem. Soc., 1973, 95, 7177. B. B. Hasinoff, Biochem., 1974, 13, 3111. C. D. Hubbard, C. J. Wilson and E. F. Caldin, J. Amer. Chem. SOC., 1976,98, 1870. Joussot-Dubien (Elsevier, Amsterdam, 1975). "1. Gianini, in Lasers Phys. Chem. Biophys. Proc. 27th Int. Meet. SOC. Chim. Phys., ed. J. l3 M. A. Cobb and D. N. Hague, J.C.S. Faruday I, 1972, 68,932. I4 G. R. Cayley and D. N. Hague, J.C.S. Faraday I, 1971, 67, 786. M. Eigen and R. G. Wilkins, in Mechanism of Inorganic Reactions, ed. R. F. Could, Adv. Chem. Series No. 49. B. H. Robinson, N. C . White and C. Mateo, Adu. Mol. Relax. Proc., 1975, 7, 321. Concepts of Inorganic Photochemistry, ed. A. W. Adamson and P. D. Fleischauer (Wiley Interscience, New York, 1975). (Amer. Chem. SOC., Washington, D.C., 1965), p. 55. I6 A. D. James and B. H. Robinson, J.C.S. Faruduy I, 1978,74, 10. l9 F. Basolo, J. C . Hayes and H. M. Neumann, J. Amer. Chem. Soc., 1954, 76, 3807. 2o T. J. Swift and R. E. Connick, J. Chem. Phys., 1962, 37, 307. H. Hoffmann, Ber. Bunsenges. phys. Chem., 1969, 73, 432. 22 L. 6. SillCn and A. E. Martell, StabiZity Corzstants of Metal-Ion CompZexes (Chem. SOC. Spec. Publ., London, 1964). 23 M. W. Grant and C . J. Wilson, J.C.S. Farauby I , 1976, 72, 1362. 24 R. C. Young, F. R. Keene and T. J. Meyer, J. Amer. Chem. Soc., 1977, 99, 2468. (PAPER 8/01 1)

ISSN:0300-9599    

DOI:10.1039/F19787402625

出版商:RSC    

年代:1978

数据来源: RSC