摘要:

J. Chem. SOC., Faraduy Trans. I , 1982, 78, 2215-2225 Infrared Spectroscopic Study of CO and CO, Adsorption on Rh-ZrO,, Rh-Al,O, and Rh-MgO BY YUKARI TANAKA, TOKIO IIZUKA* AND Kozo TANABE Department of Chemistry, Faculty of Science, Hokkaido University, Sapporo 060, Japan Received 28th September, 198 1 The adsorption states of CO and CO, on Rh catalysts supported on ZrO,, Al,O, and MgO were studied using i.r. spectroscopy. Upon adsorption of CO on reduced Rh-ZrO,, the carbonate bands due to the reaction 2CO + C+CO,, together with the bands of twin, linear and bridge CO species were observed at moderate temperatures, whereas Rh-MgO gave no appreciable formation of CO, even at higher temperatures (> 100 "C) and Rh-A1,0, showed intermediate behaviour. On the adsorption of CO,, the linear CO band was formed at lower frequency than that on CO adsorption. The linear CO species formed from CO, showed higher reactivity toward hydrogen compared with that from CO adsorption.Infrared spectroscopy is widely used in the study of adsorbed CO on transition metals in relation to their catalytic properties for Fischer-Tropsch synthesis. Carbon monoxide is known to chemisorb in three forms on transition metals. These are the 'bridge' type, the ' linear' type and in some cases the 'twin' type. The stability of the adsorbed species follows the order bridge > linear > twin.' Eischens, reported that metals on which CO chemisorbs in the bridge form (Ni and Pd) show higher catalytic activity in the methanation of CO than metals on which CO chemisorbs in the linear form (Cu and Pt).On the other hand, Heal et aL3 showed that the temperature region for CH, formation on Ni-SiO, is close to that for a quick decrease in the adsorption band of linear CO and concluded that linear CO is involved in CH, formation. Recently, Fujimoto et al., reported that the bridge CO was hydrogenated at lower temperatures than the linear CO to form hydrocarbons consisting mainly of CH, on Rh-Al,O, and Ru-AI,O, and that the twin CO on Rh-Al,O, desorbed without being hydrogenated. On the other hand, the adsorbed states of CO, on the catalyst in the hydrogenation of CO, were not clearly elucidated. Solymosi et aL5 observed linear and bridge CO formed from CO, in the presence of H,, but could not obtain those bands clearly without H,.Dubois and Somorjai6 observed the dissociation of CO, on the surface of Rh metal crystals with LEED and ELS techniques in an ultrahigh vacuum system. In this report we discuss the support effect for CO adsorption on Rh catalysts and the difference in the adsorption states of CO and CO, in relation to their reactivity toward hydrogen. EXPERIMENTAL Supported rhodium catalysts were prepared by impregnating ZrO,, A1,0, or MgO with an aqueous solution of Rh(NO,),. After the evaporation of water, the catalysts were dried in air at 100 "C for 24 h and calcined at 500 O C for 2 h. Zirconium oxide was prepared by the hydrolysis of ZrOC1, with aqueous ammonia, followed by calcining at 500 O C . Aluminium oxide was obtained by calcining the hydroxide at 500 OC, after its preparation by the hydrolysis of Al(NO,), with aqueous ammonia.Magnesium oxide 221 52216 was obtained by heating Mg(OH), at 500 OC. The content of Rh was 2.3 wt% on all the catalysts. For measurement of the i.r. spectrum, a conventional flow-through cell was used. The catalyst was pressed into a thin wafer and placed in the cell. All spectra were recorded in the temperature range 25-250 O C with a Jasco 701G grating spectrometer. The sample in the cell was pretreated by evacuation at 300 OC for 2 h and then reduced for 2 h at the same temperature. Carbon monoxide, carbon dioxide and hydrogen were purified by passing over molecular sieve 4A. I.R. STUDY OF CO AND CO, ADSORPTION ON Rh CATALYSTS RESULTS The i.r. spectra of CO adsorbed on Rh-ZrO, at various temperatures are shown in fig.1 . Upon adsorption at room temperature, the absorption bands at 2095, 2065 and 2030 cm-l in the linear- and twin-type region and at 1870 cm-l in the bridge-type la) 2000 1900 1800 1700 1600 1500 wavenumberlcm-' FIG. 1 .-1.r. spectra of CO adsorbed on Rh-ZrO, previously reduced at 300 "C: (a) background, (6) room temperature, and after heating at (c) 100 "C, ( d ) 150 OC, (e) 200 "C. region were observed. In addition to those bands, the broad signals of carbonate species at ca. 1560 and 1300 cm-l were observed at the beam temperature (ca. 40 O C ) . Upon heating at higher temperatures in the presence of CO, the bands at 1560 and 1300 cm-l and the twin bands at 2095 and 2030 cm-l increased in intensity and the band at 1870 cm-l shifted slightly to lower frequency. After heating at 200 O C in CO, the twin bands almost disappeared, as shown in fig.I (e). When Rh-ZrO, was reduced at 500 OC and evacuated at the same temperature, the development of carbonate and twin bands was also observed, as shown in fig. 2, but the intensity of these bands was slightly less than those in the previous case. With Rh-Ai,O,, the formation ofY. TANAKA, T. I I Z U K A AND K. TANABE 2217 carbonate species was not significant at room temperature, as shown in fig. 3, and only the linear- and bridge-type CO species were observed. Upon heating at higher temperatures ( 2 100 "C), carbonate bands at ca. 1550 and 1300 cm-l appeared, along with twin bands at 2105 and 2035 cm-l, but the intensity of these bands was very weak compared with the case of Rh-ZrO,. On the other hand, the adsorption states of CO on Rh-MgO were very different from those on Rh-ZrO, and Rh-Al,O,, as shown I I I I I 1 2000 1900 1800 1700 1600 1500 wavenumber/cm-' FIG.2.-1.r. spectra of CO adsorbed on Rh-ZrO, previously reduced at 5OOOC: (a) background, (b) room temperature, and after heating at (c) 100 O C , ( d ) 150 O C , (e) 200 O C . in fig. 4. Upon adsorption of CO at room temperature, only the linear-type band at 2020 cm--l and the bridge-type band at 1870 cm-l were observed and the bands of the carbonate species were not observed. After heating the sample in gaseous CO at 100-200 "C, the band at 1870 cm-l increased in intensity and shifted to lower frequency (1825 cm-l after heating at 200 "C), but the carbonate species was not observed after heating at 200 "C.The adsorption of CO, on the Rh catalysts was also examined by i.r. spectroscopy. As shown in fig. 5 (A), carbon dioxide adsorbed on Rh-Al,O, which had been reduced at 300 "C gave the linear CO species at ca. 2020 cm-l. The band intensity increased with the adsorption temperature, but the twin and bridge species were not observed even after heating at higher temperature (> 200 "C). The linear CO band was also observed for CO, adsorption on Rh-ZrO, and Rh-MgO with reduced intensity. Upon adsorption of a mixture of CO, and H,, a strong enhancement of the CO band was observed, as shown in fig. 5(B), but the band frequency was almost the same as that in the case of pure CO, adsorption. 72 F A R 12218 I.R.STUDY OF CO AND CO, ADSORPTION ON Rh CATALYSTS 1 2000 1800 1600 1400 wavenurnberlcm- ' FIG. 3.-1.r. spectra of CO adsorbed on Rh-A1,0, previously reduced at 500 O C , (a) background, (b) room temperature, and after heating at (c) 100 O C , ( d ) 150 "C, ( e ) 200 OC. 1 1 1 1 I I I 2000 1800 1600 1400 wavenumber/cm -' FIG. 4.-1.r. spectra of CO adsorbed on Rh-MgO previously reduced at 300 O C , (a) background, (b) room temperature, and after heating at ( c ) 100 OC, ( d ) 150 OC, (e) 200 "C, (f) CO, adsorption at 200 O C .Y. TANAKA, T. I I Z U K A A N D K. TANABE 2219 When CO was adsorbed on Rh-A1,0, which had been oxidized at 5OO0C, a spectrum similar to that of CO, adsorbed on the reduced surface was observed. The spectra are shown in fig.6. The only difference was the existence of twin species in the case of CO adsorption on the oxidized catalyst. Upon adsorption of CO at room temperature, a sharp band at 2 1 15 cm-l and a small band at 2030 cm-l were observed, 2035 ,2065 , 2105 , I l l 00 2000 1800 1600 : I I 1 I I l 00 2000 1800 1600 wavenum ber/cm -' FIG. 5 . 4 . r . spectra of CO, adsorbed on Rh-Al,O, previously reduced at 300 O C . (A) Adsorption of pure CO, : (a) background, (b) room temperature, and after heating at (c) 100 O C , ( d ) 150 O C , (e) 200 O C , (f) 250 "C, ( g ) CO adsorption at room temperature. (B) Adsorption of CO, and H,: (a) background, (b) room temperature, and after heating at ( c ) 100 OC, ( d ) 150 O C . but the band of the bridge species was not observed. After heating the sample at 100 O C , the linear CO band appeared as a shoulder at 2080 cm-l and the carbonate band due to the reduction of the surface with CO increased in intensity.The linear band shifted to 2055 cm-l upon evacuation at room temperature. The activity of adsorbed CO species on Rh-A1,0, toward H, was examined spectroscopically. When CO species adsorbed on the reduced catalyst were heated in H,, the linear band shifted slightly to lower frequency and the bridge species disappeared at 120- 150 OC, then the linear species disappeared at ca. 175 O C , as shown in fig. 7. The twin-type CO species was stable up to 190 O C . Since these species were observed at 180-200 O C in the absence of H,, the disappearance of bridge and linear CO bands in H, can be ascribed to the reaction with H,.The formation of CH, in the gas phase was detected at ca. 150 OC. In the case of adsorbed CO, in H,, the linear CO band disappeared completely a t 150 OC, as shown in fig. 8. In the presence of CO, 72-22220 I.R. STUDY OF CO AND CO, ADSORPTION ON Rh CATALYSTS and H, in the gas phase, the linear band still existed at 150 OC, as shown in fig. 5 . In this case, the formate species appeared at 1590 and 1390 cm-l at temperatures > 100 OC, but this species was stable after heating at 150-200 "C in gaseous H,. On the oxidized Rh-A1,0,, the linear CO species at 2050-2055cm-l disappeared at 150-175 "C in H,, but the twin species was stable at this temperature. I I 1 I I 2000 1900 1800 1700 1600 wavenum ber/cm-' FIG. 6.-1.r. spectra of CO adsorbed on Rh-A1,0, previously oxidized at 500 O C : (a) background, (b) room temperature, and after heating at (c) 100 O C , ( d ) 150 O C , (e) 200 O C . DISCUSSION ADSORPTION OF co As shown in fig.1, the formation of adsorbed CO, (carbonate) species over reduced Rh-21-0, was observed along with the ordinary CO adsorbed species such as the bridge, linear and twin species upon the introduction of CO. This suggests that CO disproportionates to carbon and CO, on the surface. Upon the adsorption of CO at room temperature, the twin-species bands were weak and the higher-frequency band was observed as a shoulder of the linear band. After heating the sample in CO at 100-150°C, the twin-species band increased in intensity with the increase of the carbonate bands. Since the twin species is known to form on the oxidized Rh site, probably Rh', the development of this species upon heating in CO would indicate oxidation of the Rh site by the dissociation of CO, as reported by Primet.' AfterY.TANAKA, T. I I Z U K A A N D K. TANABE 222 1 the sample had been heated at 200 O C , the twin species almost disappeared and the carbonate band increased in intensity on cooling to room temperature. In this case, the disappearance of the twin-species band was not due to desorption, because gaseous CO was still present in the i.r. cell. The site responsible for the formation of the twin species might be reduced by CO to form CO,. Moreover, since carbonate 1 1 1 I I I 2000 1900 1800 1700 1600 1500 wavenumberlcm -' FIG. 7.-1.r. spectra of CO adsorbed on Rh-Ai,O, in H, at various temperatures: (a) background, (b) room temperature, (c) 100 O C , ( d ) 125 O C , (e) 150 O C , (f) 175 O C .species were also formed on Rh-ZrO, which had been reduced and evacuated at temperatures > 500 O C , we excluded the possibility of CO, formation in the water-gas shift reaction, because the residual water on the surface would be minimized in the evacuation at high temperature. Thus, we concluded that the formation of CO, in CO adsorption was due to the disproportionation of CO to carbon and CO,. In the case of Rh-Al,O,, the formation of CO, was not observed in the adsorption of CO at room temperature, but carbonate bands appeared along with an increase in intensity of the twin-species bands at higher temperatures (2 100 "C).Since the intensities of the carbonate and twin-species bands at 100 O C on Rh-A1,0, were weak compared with those of Rh-ZrO,, the dissociation of CO on Rh-A1,0, was less favourable than on Rh-ZrO,. On the other hand, over the Rh-MgO surface the formation of carbonate and twin species was not observed, even at higher temperatures. This is a marked contrast with the case of Rh-ZrO,. We recently reported that Rh-ZrO, was an excellent catalyst for the hydrogenation reactions of CO and CO, to form hydrocarbons.s The order of activity was Rh-ZrO, > Rh-A1,0, > Rh-2222 I.R. STUDY OF CO AND CO, ADSORPTION ON Rh CATALYSTS c 0 .- in .- E E? c Y b u 1 I 1 I 2000 1900 1800 1700 1600 1500 wavenumber/cm - I FIG. 8.-1.r. spectra of CO, adsorbed on Rh-A1,0, in H, at various temperatures: (a) room temperature, (6) 75 OC, (c) 100 OC, ( d ) 125 O C , (e) 150 OC.SiO, 9 Rh-MgO for the reactions of both CO and CO,. Though we failed to obtain a clear i.r. spectrum of CO adsorption on Rh-SO,, the tendency for CO dissociation corresponded very well with the activity. Recently, experimental evidence has been presented supporting the idea that CO dissociation to form carbon on the surface is an essential initial step in the synthesis of hydrocarbon^.^ Thus, it would be useful to study the surface characteristics of supported Rh catalysts affecting the dissociation of CO. Ichikawa and Kawai recently reported a partial electron transfer from TiO, or ZrO, oxide supports to Rh,(CO),, on the basis of X.p.s. experiments10 They stated that partially reduced states of TiO, or ZrO, stabilize the low oxidation state of Rh and hence the reactivity of CO is enhanced in the synthesis of alcohol from CO and H,.Probably this electron-transfer effect is also operating in our case to stimulate the dissociation of CO over Rh-ZrO,. On the other hand, the reason for the difficulty of CO bond dissociation on Rh-MgO, which probably leads to the inactivity for the hydrogenation of carbon oxides, is not clear yet. One reason for the inactivity of Rh-MgO would be a low dispersion of Rh on MgO. The dispersion values of Rh measured on the basis of H, adsorption were 0.51,0.60 and 0.27 for Rh-ZrO,, Rh-A1,0, and Rh-MgO, respectively.s Since MgO is familiar as a substrate for the epitaxial growth of metal films with low-index orientations on its surface,ll Rh atoms will exist as a metal crystallite of low-index planes mainly on MgO, even in this case.Moreover, the absence of twin CO bands on Rh-MgO in the i.r. spectrum is very similar to that of CO adsorption on Rh metal crystal.12 It is known that the adsorption of CO on Rh metal crystal is classified as non-di~sociative.~~ Thus, the inactivity ofY. TANAKA, T. I I Z U K A AND K. TANABE 2223 Rh-MgO for CO dissociation is ascribed to the segregation of metal which probably has low-index planes on the surface. ADSORPTION OF co, As for the adsorption of CO, over Rh catalysts, previous worker^^^-^^ could not detect the CO band, indicative of the dissociation of CO, on the surface. However, Dubois and Somorjai found the dissociation of CO, on several faces of Rh single crystal by means of LEED or ELS techniques in an ultrahigh vacuum system.6 Owing to the difficulty of detection, Dubois and Somorjai suggested the low sticking probability of CO, and the high rate of the association reaction CO(ads)+O(ads)+ CO,(g) requiring an approximately five- to ten-fold higher exposure for the adsorption of CO, compared with CO to obtain the spectra.6 On the other hand, Primet' and Solymosi et al.5 observed weak CO bands in CO, adsorption on Rh-Al,O,, but this was not reported in their work.In our case on Rh-A1,0,, the dissociation of CO, to CO and 0 atom was observed even at a moderate pressure of CO, (2-15 Torr) and at room temperature. The CO band was also observed on Rh-ZrO, and Rh-MgO, but the intensity of the CO band was weak over those catalysts compared with that on Rh-Al,O,.Probably the dissociation of CO, depends on the nature of the support used, the preparation of the catalyst and the dispersion of the Rh. Note the difference in CO and CO, adsorption on Rh-MgO. The surface of Rh-MgO was inactive for the dissociation of CO to carbon but showed comparable activity to Rh-ZrO, for the dissociation of CO, to CO and 0 atom. This would be the reason why CO was the main product in the reaction of CO, and H, at > 200 O C on Rh-MgO.s Solymosi et al. observed the formation of a CO band from CO, in the presence of H,, but could not obtain the clear band without H, under the same condition^.^ Concerning the main difference between the spectrum of adsorbed CO and that obtained after coadsorption of the H,+CO, mixture, they concluded that (i) the doublet due to twin CO was completely missing and (ii) the linearly bonded CO appeared at lower frequency, 2020-2039 cm-l.They suggested that the adsorbed hydrogen prevents the formation of a twin structure, the hydrogen adsorbed on the metal atom of the carbonyl might donate an electron to CO and consequently the vibrational frequency of CO would shift to lower freq~ency.~! 1 7 7 However, in this work, the twin CO species was also completely missing and the band frequency of the linear species was still lower than that of pure CO adsorption in the absence of H, after a comparable band intensity with that in H, was obtained, as shown in fig. 5. Moreover, since the adsorption of CO was very strong compared with that of hydrogen,8 it would be unlikely that the adsorbed hydrogen prevented the formation of twin species.Thus, it is rather difficult to ascribe the reason for the absence of twin bands and the shift of the linear band to lower frequency to the coadsorption of hydrogen with CO on the same Rh atom. In the case of CO, adsorption on Ni/Aerosil, van Hardeveld et al. reported the shift of CO bands formed from CO, to lower frequency compared with CO adsorption. They explained the shift of the CO bands as follows: the CO, decomposes into CO + 0, which are most probably adsorbed onto adjacent atoms. The adsorption of 0 on Ni results in the oxidation of Ni. The positive charge on the Ni atom polarizes the CO molecule adsorbed on it and/or forces more electrons into the Ni-C bond, thereby weakening the C-0 bond and lowering its stretching frequency.However, in our case with Rh-Al,O,, after the oxidation of the catalyst the linear CO band appeared at higher frequency (ca. 2080 cm-l) than that on the reduced surface, and shifted to 2055 cm-l upon evacuation. Yang and Garland have reported the high-frequency shift of2224 I . R . STUDY OF CO AND CO, ADSORPTION ON Rh CATALYSTS the linear CO band upon the oxidation of Rh.14 Thus, it is also unlikely that the low-frequency shift is due to the oxidation of metal in our case. Since it is well known that the band frequency of CO stretching is a strong function of CO c ~ v e r a g e , l ~ > ~ ~ we are inclined to think that the appearance of CO at a lower frequency in the CO, adsorption compared with pure CO is due to a small covering of CO formed from CO, on the surface.The shift of the CO linear band on the oxidized surface to lower frequency upon evacuation would be due to the removal of weakly adsorbed CO around the linear species having a dipolar interaction with strongly chemisorbed species; however, this linear species still appears at a higher frequency than CO formed from CO,, in spite of the similarity of surroundings (except for the presence of the twin species in the case of CO adsorption on the oxidized surface). Thus, it would be reasonable to ascribe the frequency shift to lower coverage of CO in CO, adsorption. The CO species formed from CO, would be isolated on the surface, and it did not show a shift to higher frequency at higher coverage in the presence of hydrogen.REACTION OF ADSORBED SPECIES WITH HYDROGEN In the reaction of adsorbed CO with H, over Rh-Al,O,, the bridge species reacted first at 120-150 OC, and then the linear species reacted at ca. 175 O C , as reported by Fujimoto et aL4 Although we could not detect a clear bridge band in CO, adsorption, the linear CO species formed from CO, reacted with hydrogen at lower temperatures than pure CO. This corresponds well with the fact that CO, reacted with H, to form methane at a temperature lower than that for the hydrogenation of CO over Rh catalysts.* This phenomenon can be explained as follows: the coverage of CO from CO, is very small compared with that in pure CO adsorption, so the adsorption of H, is more favourable in CO, hydrogenation than the case of CO, which acts as a poison for H, adsorption.For the same reason, the linear CO species in CO adsorption on the oxidized surface showed a slightly higher reactivity with H, compared with CO on the reduced catalyst, corresponding to the fact that the rate of hydrocarbon formation on the oxidized surface was higher than that on the reduced catalyst.* In the reaction CO+H,, the oxidized Rh catalyst showed higher selectivity to higher hydrocarbons (C,-C,) than the reduced catalyst, but only methane was formed in the reaction CO, + H,.* Comparing the adsorbed states of CO on the oxidized surface with those of CO, on the reduced surface, the main difference was the absence of twin species in the case of CO,.Moreover, the twin species was less intense on the reduced surface in CO adsorption. Thus, the CO in a weak adsorption state, such as twin species, would have an important role in the propagation of hydrocarbons in CO hydrogenation on the oxidized surface. H. C. Yao and G . W. Rothschild, J . Chem. Phys., 1978, 68, 4774. R. P. Eischens, in The Surface Chemistry of Metals and Semiconductors, ed. H. C. Gatos (Wiley, New York, 1960). M. J. Heal, E. C. Leisegang and R. G. Torrinton, J. Catal., 1976, 42, 10. K. Fujimoto, M. Kameyama and T. Kunugi, J. Catal., 1980, 61, 7. F. Solymosi, A. Erdohelyi and M. Kocksis, J . Catal., 1980, 65, 428. L. H. Dubois and G. A. Somorjai, Su$. Sci., 1979, 88, L13. ' M. Primet, J . Chem. SOC., Faraday Trans. I, 1978, 74, 2570. T. Iizuka, Y. Tanaka and K. Tanabe, J . Catal., in press. M. Araki and V. Ponec, J. Catal., 1976, 44, 439. lo M. Ichikawa and M. Kawai, Shokubai (CataZysf), 1981, 23, 55. l 1 J. Prichard, T. Catterick and R. K. Gupta, Surf. Sci., 1975, 53, 1 . l2 R. R. Cabanah and J. T. Yates Jr, J . Chem. Phys., 1981, 74,4150. l 3 G. Broden, T. N. Rhodin, C. Brucker, R. Benbow and Z. Hurich, Surf. Sci., 1976, 59, 593.Y. TANAKA, T. I I Z U K A A N D K. TANABE l4 A. C. Yang and C. W. Garland, J. Phys. Chem., 1957, 61, 1504. l5 C. T. Campbell and J. M. White, J. Catal., 1978, 54, 289. l 6 A. C. Collings and B. M. W. Trapnell, Trans. Faraday SOC., 1957, 53, 1436. l7 F. Solymosi, A. Erdohelyi and M. Kocksis, J. Chem. SOC., Faraday Trans. I , 1981, 77, 1003. F. Solymosi, A. Erdohelyi and T. Bansagi, J. Chem. SOC., Faraday Trans. I, 1981, 77, 2645. A. 34. Bradshaw, Surf. Sci., 1979, 80, 215. 2225 (PAPER 1 / 1504)

ISSN:0300-9599    

DOI:10.1039/F19827802215

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 2227-2232 Kinetics and Mechanism of Polymerization of Phenylacetylene Initiated by (Phenylmet hox ycarbene)pentacarbonyltungs ten( 0) BY H o H u u THOI, KENNETH J. IVIN AND JOHN J. ROONEY* Department of Chemistry, The Queen’s University of Belfast, Belfast BT9 5AG, Northern Ireland Received 28th September, 198 1 The kinetics of polymerization of phenylacetylene in CDC1, solution, intiated by Ph(MeO)C=W(CO),, has been followed under various conditions by in situ analysis of monomer (A) and initiator (F) using ‘H n.m.r. spectroscopy. In degassed solution at 332 K the polymerization is zero order with respect to A and fractional order (0.7) with respect to F, the variation of rate with temperature (323-343 K) corresponding to an apparent activation energy of 116 kJ mol-’.The reaction is retarded by the presence of CO, 0, or air, but not by N,. Under 1 atm CO or air the reaction becomes first order with respect to F but remains zero order with respect to A. The mechanism of initiation and propagation is discussed in terms of elementary steps involving formation and reactions of metallacarbenes, metallacarbene-monomer complexes and metallacyclobutenes. The polymerization of phenylacetylene (A) initiated by olefin metathesis catalysts such as WCl,, MoCl, and WCl,/Ph,Sn has been extensively studied by Higashimura and Masuda.l The rate of polymerization and molecular weight of the polymer were investigated as a function of solvent, substituent (on the phenyl group of the monomer), cocatalyst and chain transfer agent; it was concluded that propagation proceeded by a coordination mechanism.Recently Katz2 has reported that metallacarbenes such as the so-called Fischer3 compound (F), Ph(MeO)C=W(CO),, and the Casey4 compound, Ph,C=W(CO),, catalyse the polymerization of phenylacetylene and alkylacetylenes to give high- molecular-weight polymers. This provides convincing evidence for the following mechanism where propagation proceeds via formation and rupture of a metallacyclo- butene intermediate. R1R2C CH R1R2C R1R2C-CH R1R2C=CH [Mt] CR3 [MtI+- 111 [Mtl-CR3 [Mt]=CR3 1 -+ etc. CR3 Furthermore, phenylacetylene in conjunction with F also induces the ring-opening polymerization of cy~loalkenes,~ and we have ourselves observed that it accelerates the polymerization of norbornene. There is as yet comparatively little kinetic information with which to test the mechanisms of these reactions.In the present work we have studied the kinetics of polymerization of phenylacetylene using F as initiator, since this compound has the advantages of being relatively stable at room temperature, insensitive to moisture and giving rates of polymerization that are sufficiently slow to be able to follow the disappearance of both monomer and initiator by lH n.m.r. spectroscopy. The 22272228 POLYMERIZATION OF PHENYLACETYLENE problem of irreproducibility and inhomogeneity which are often encountered with the much more active two-component metathesis catalysts were thereby avoided. EXPERIMENTAL Spectrograde deuterated chloroform was used without further purification.Phenylacetylene (B.D.H.) was distilled under reduced pressure. Ph(MeO)C=W(CO), was prepared according to the method of Fi~cher,~ the crude product being purified by column chromatography (silica gel, 1 % dichloromethane in pentane), recrystallized twice from pentane at 195 K, and stored under argon at 253 K. The reaction mixtures were prepared directly in n.m.r. sample tubes which could be connected to the vacuum line for degassing purposes. A small amount of toluene was added to provide an internal n.m.r. standard. After three freeze-thaw cycles the appropriate gas (CO, N, or 0,) was admitted at 1 atm (when required),* the tubes sealed off and then inserted in a thermostat (323-343 K) to start the reaction. At various time intervals the tubes were quenched and the concentration of monomer estimated from the area of the lH n.m.r.peak due to the acetylenic protons (6 = 3.0 ppm), relative to the area of the peak due to methyl protons of the added toluene. In this way the monomer concentration could be determined with a precision of kO.02 mol dm-3. The formation of polymer was indicated by the appearance of a broad peak at 6.3-7.5 ppm. The concentration of F was likewise estimated from its methoxy proton peak (6 = 4.6 ppm); in this case the limit of precision was kO.01 mol dm-3. The formation of carbon monoxide was confirmed by high-resolution mass spectroscopy and that of tungsten hexacarbonyl by 13C n.m.r. RESULTS Typical plots for the conversions of monomer (A) and initiator (F) against time are shown in fig.1 for a degassed solution. Note that the rate of removal of A remains constant up to 75% conversion while the rate of removal of F suffers a sharp fall after 8% conversion. After 120 min the ratio of monomer removed to initiator removed is > 20, indicating a chain reaction. 100 I I I 1 0 0 50 100 150 time/min FIG. l.--Conversion-time plots for (a) monomer (A) and (b) initiator (F) in degassed CDCl, solution at 332 K. [A],, = 0.90, [F], = 0.24 mol dm-3. * 1 atm = 101 325 Pa.H. H. THOI, K. J. IVIN AND J. J. ROONEY 2229 0 100 200 300 4 00 0 50 100 150 200 ti me/ m in FIG. 2.-Concentration of monomer [A] against time for various initial concentrations [A],: (a) 4.9, (b) 1.48, (c) 0.87, ( d ) 0.53 mol dm-, in degassed CDCI, solution at 332 K.[F], = 0.155+0.005 mol dm-,. - 9 -1 0 h - I I/? 0 ,-- .-. 0 0 - E D. -11 . s f: - -12 1 p’ -4 -3 -2 -1 0 FIG. 3.-Logarithmic plot of the rate of polymerization, R,, against the initial concentration of initiator, [F],, in chloroform solution at 332 K: (a) degassed, (b) 1 atm air present, (c) 1 atm CO present, ( d ) 1 atm N, present, (e) 1 atm 0, present. [A], = 0.9 0.1 rnol dm-3. The zero order with respect to monomer is confirmed by the results shown in fig. 2 for experiments at different initial monomer concentrations : the rate of polymeriza- tion R, is constant at (9+2) x lod5 mol dm-3 s-l, over a range of [Ale from 0.5 to 4.9 rnol dmP3, using [F],, = 0.155 & 0.005 mol dm-3. R, as a function of [Fl0 is shown in fig. 3 for the degassed solution [line (a)] and for the reaction in the presence of 1 atm air [line (b)], 1 atm CO [line (c)], 1 atm N,2230 POLYMERIZATION OF PHENYLACETYLENE [point (d)] and 1 atm 0, [point (e)].The reaction is retarded several-fold in the presence of air, CO or O,, but is unaffected by N,; it is evident that the active component of air is 0,. The order with respect to initiator is 0.7 for the degassed solution but unity in the presence of 1 atm air or CO. Although the initial rates of polymerization under 1 atm 0, and 1 atm CO are about the same the reaction slows up under 0, and eventually stops after ca. 3 h. The concentration of initiator decreases much more rapidly under oxygen than under vacuum and falls to nearly zero afer 30 h. Under 1 atm CO or air the order with respect to monomer remains zero.The apparent activation energy for the reaction in the degassed solution is 1 16 kJ mol-l(323-343 K). DISCUSSION The key observations are the zero order with respect to monomer, the fractional order with respect to initiator, and retardation by CO, which is also a product. These observations can be accounted for in terms of the following mechanism. F + I+CO (1) 1+A -+ IA (2) IA-+(IA) -+Pl (3) P,+A e P,A (4), (-4) Pl A -+ ( P A ( 5 ) I denotes Ph(MeO)C=W(CO),, formed from F by loss of one molecule of CO. IA denotes I with one molecule of A coordinated to the vacant site, while (IA) is the intermediate metallacyclobutene. P, is the product Ph(MeO)C=CPh-CH=W(CO), [or Ph(MeO)C=CH-CPh=W(CO)J, formed by intramolecular reaction of IA as indicated in the Introduction.P,A is the species formed by coordination of P, with another molecule of A, while (P,A) is again the corresponding intermediate metallacyclobutene. It is clear that if, over the experimental range of [A], the equilibrium (4), (-4) is well to the right, then for long chains the rate of removal of A will be governed by the rate of step ( 5 ) or (6) and will be independent of [A], as observed. Step (7) represents a termination step of the 16-electron species (P,A) with CO, involving ultimately the formation of a pentacarbonyl species F’ analogous to F. This will also account for retardation by CO; further reactions of F’ can lead to the observed W(CO),. Reaction (1) may be slightly reversible, but not sufficiently so that it competes significantly with step (2), otherwise there would be an appreciable dependence of rate on [A].Competition between CO and A for (P,A) or P, can be omitted for the same reason. The mechanism therefore avoids postulating 20-electron metal complexes as intermediates, but makes a distinction between the various types of 16-electron metal complexes in terms of their relative reactivities towards A and CO or 0,. There is apparently no significant competition between A and CO for reaction with the metallacarbenes, I or P,, since the kinetics are still zero order in A even under 1 atm CO. However, we have to postulate completely opposite behaviour for the metallacyclobutenes, (P, A), which must coordinate CO or 0, readily in the presence of A if the poisoning effects of these diatomics are to be explained.This distinction in the behaviour of the two types of 16-electron metal complexes may lie in the differences between the electron-donor-electron-acceptor properties of A and these gases.H. H. THOI, K . J. I V I N A N D J. J. ROONEY 223 1 Assuming that termination reactions of P, or (P,A) other than those with CO are insignificant, a steady-state treatment, with step (6) rate determining, leads to the following expression which, when CO is present initially in sufficient concentration [CO],, reduces to where R, = k[F] k = k6kl/k7[C0],. (9) The fractional order (0.7) with respect to F in the degassed solution can be explained as follows. In the initial stages a steady-state concentration of CO builds up [reactions ( I ) and (7)]. This concentration of CO is obviously a function of the initial concentration of the Fischer compound, [F],.The denominator in eqn (8) is therefore a function of [Fl0 such that the rate of polymerization has a fractional power dependence. When CO is initially present in much greater amount than that produced in the reaction then a simple first-power dependence on [Fl0 is to be expected. These conditions appear to obtain in the presence of 1 atm CO. The activation energy of 116 kJ mol-1 is unusually high for this type of polymeri- zation but a value of this magnitude is expected if step (1) is of major importance in rate control. The fall-off in the rate of polymerization at very low [A] [fig. 2, line (d)] may be attributed to a significant shift of equilibrium (4), (-4) towards the left.Note that an analogous reaction to step (1) has been postulateds as the first step in the metathesis reactions of alkenes, initiated by Ph,C=W(CO),, and that the metallacarbene analogous to I has been detected by lH n.m.r. The present mechanism is based on the consecutive formation of the 16-electron complexes P,, the 18-electron complex P, A, and the 16-electron metallacyclobutene (P, A). An insertion reaction (lo), of the type shown to occur in the reactions of polar acetylenes with F and assumed to proceed via a zwitterion,, OMe OMe I (CO),Mt CPh (10) 6- s+,OMe I d + + 6- Ph - + I - II ll (CO),Mt=C, (CO),M t-CPh YCECZ Y=C=CZ YC-cz (Y = NEt,, OEt; Z = Me, H; Mt = Cr, W) is ruled out in the present case by the zero order with respect to A. A displacement reaction (1 1) of the type F+A+IA+CO (1 1) can be ruled our for the same reason. In any case reactions (10) and (1 1) do not allow any explanation of the retardation caused by CO. We thank the S.R.C. for financial support.2232 POLYMERIZATION OF PHENYLACETYLENE T. Masuda and T. Higashimura, Macromolecules, 1979, 12, 9 and references cited therein. T. J. Katz, S. J. Lee, M. Nair and E. B. Savage, J . Am. Chem. Soc., 1980, 102, 7940. E. 0. Fischer and A. Maasbol, Chem. Ber., 1967, 100, 2445. C. P. Casey, T. J. Burkhardt, C. A. Bunnell and J. C. Calabrese, J . Am. Chem. SOC., 1977, 99, 2127 T. J. Katz, E. B. Savage, S. J. Lee and M. Nair, J . Am. Chem. SOC., 1980, 102, 7944. C. P. Casey and A. J. Shusterman, J. Mol. Cutal., 1980, 8, 1. ’ H. Fischer and K. H. Doetz, Chem. Ber., 1980, 113, 193. (PAPER 1 / 1 505)

ISSN:0300-9599    

DOI:10.1039/F19827802227

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. Soc., Furuduy Trans. I, 1982, 78, 2233-2238 Self-diffusion in Liquid Acetonitrile under Pressure BY ROBERT L. HURLE AND LAWRENCE A. WOOLF* Research School of Physical Sciences, Australian National University, Canberra, A.C.T. 2600, Australia Received 6th October, 198 1 Self-diffusion measurements are reported for acetonitrile in the temperature range 238-343 K at pressures up to 300 MPa. Lack of reliable high-pressure density data has restricted tests of theory to the temperature range 268-328 K. Over that range the results show a breakdown of rough hard-sphere theory which is attributed to the strongly dipolar nature of the acetonitrile molecule. A correlation approach based on Enskog smooth hard-sphere theory is found to be less sensitive to the dipole-dipole interactions.Recently self-diffusion measurements under pressure have been reported for liquid carbon disulphide,l which has a linear molecular arrangement. Contrary to the predictions of Chandler,2 the results were satisfactorily represented by the rough hard-sphere theory. The limited computer simulation data available3? indicate that a three-centre Lennard-Jones potential seems likely to provide a good prediction of CS, diffusion. This paper reports self-diffusion data in the temperature range 238-243 K for pressures up to 300 MPa for acetonitrile, which resembles CS, in having almost linear molecules5 but differs in having a substantial dipole moment, 11.3 x C m. The data therefore enable different tests of the hard-sphere models of self-diffusion and also provide an opportunity for a computer simulation of the self-diffusion in CH,CN, employing either an extension of the model used by Singer et a1.6 for CO, (using, in this instance, a three-centre Lennard-Jones potential) or a more realistic but computationally more difficult model which explicitly recognises the dipolar character of the molecule.EXPERIMENTAL The majority of the experiments used an n.m.r. method which has been described in detail elsewhere;7.s in the method used here, the value reported for the diffusion coefficient at each pressure is the mean of four n.m.r. experiments. Results were also obtained at 298 K by a high-pressure diaphragm cell techniques* using 14CH,CN as tracer. (There were no differences, beside experimental error, between the results obtained by the two methods.) The acetonitrile was analytical reagent grade dried over a molecular sieve; experiments with this material gave the same results as similar measurements with a specially purified sample of acetonitrile obtained from a different supplier.The labelled material, from the Radiochemical Centre, Amersham, was used without further purification; counting procedures were standards and a high ratio of active/background was obtained by using ca. 0.4 MBq of tracer for each experiment. Density data were obtained from the literaturelo, l1 and at atmospheric pressure are represented within & 0.05% by: p/g cm-3 = 1.0738,--9.1250 x 10-4(T/K)-2.8186 x 10-7(T/K)2 (228 < T/K c 338). (1) At elevated pressures the densities used were those of Srinivasan and Kay” which are probably accurate to +0.2%; they were extrapolated to f 15 K outside the temperature range reported for them by using the virtual linearity of the molar volume of acetonitrile with temperature.22332234 HIGH-PRESSURE DIFFUSION OF A CETON I TR I L E The molar volume data, including the extrapolated points, are represented to k0.57" by (p/MPa, T / K ) : In(V/~m~mol-~) = 5.209,-2.306, x 10-3p+4.7979 x 10-6p2-2.956, x lOP9p3 + ( 103/T) [ -0.6167, + 7.336, x 103/T) +4.OO2, x 10-4p-0.7256g x 10-6p2] (268 < T/K < 328.2). (2) Temperatures were held constant to kO.01 K and pressures were accurate to kO.4 MPa. The overall accuracy ofthe diffusion data is & 2% for the n.m.r. data and f 2.5% for the high-pressure diaphragm cell results.RESULTS AND DISCUSSION The experimental results are given in table 1. They are fitted within the experimental accuracy by (p/MPa, T / K ) : ln(D/10-9m2s-1) = 4.968-3.078 x 10-3p+7.361 x 10+p2 -8.446 x lOP9p3- 1.O49(1O3/T)+4.348 p(103/T). (3) Examination of table 1 shows the excellent agreement between the n.m.r. and diaphragm cell results; this is a further confirmation of the reliability of the two methods as used in these laboratories. There have been no previous measurements reported for the self-diffusion of acetonitrile under pressure but one value has been given by Zeidler12 at atmospheric pressure and 298.2 K. This value of 5.4 x lop9 m2 s-l is very different from the corresponding value in this work, 4.34 x lop9 m2 s-l, and we assume that it is wrong.Czworniak et al.13 have made some imprecise (experimental uncertainty _+ 510%) light scattering measurements of mutual diffusion at atmospheric pressure and 293 K in benzene + acetonitrile and carbon tetracholoride + acetonitrile mixtures; the activity coefficient derivatives13 for these mixtures show they are very non-ideal. Czworniak et al. were unable to use Chandler's rough hard-sphere theory2 to obtain a satisfactory analysis of the data and so could not resolve the question as to whether the dipole-dipole interactions in acetonitrile are sufficiently strong to cause a breakdown of the rough hard-sphere model. According to the model, the self-diffusion coefficient should depend linearly on the molar volume at constant temperature. The n.m.r. data plotted in fig.1 for several of the temperatures of this work show a definite, although not strong, non-linear behaviour. Note that because of the lack of volumetric data only a 60 K range (268-328 K) of high-pressure molar volumes can be used. The rough hard-sphere theory approximates the self-diffusion coefficient D of the real fluid by that of a rough hard sphere and this in turn is equated to that of a smooth hard sphere DsHs modified by a factor A , to allow for interchange of rotational and translational energy : D = A, D,,,. (4) D,,, is the Enskog dense fluid diffusion coefficient corrected for hard-sphere behaviour by incorporation of the molecular dynamics data of Alder et al:l4 DsHs = (2.527 x lo-'/ 6;) (RT/M)g ( Y - 1.384 V,) ( 5 ) where V, = Na3/2/ 2 is the volume of close-packed hard spheres.The rough hard-sphere model enables determination of the factor A, and the equivalent hard-sphere diameter by an iterative procedure which minimises the variation of A , with molar volume V at each temperature. The results for acetonitrile are given in table 2. The 0 values are close to the 0.410 nm estimates by Czworniak et al.13 using the incremental atomicR. L. HURLE AND L. A. WOOLF TABLE 1 .-SELF-DIFFUSION IN ACETONITRILE UNDER PRESSURE 2235 T / K p/MPa D/10-9 m2 s-l T/K p/MPa D/ 1 0-9 m2 s-l 238.2 5.8 10.5 21.9 35.0 55.1 253.2 0.1 5.2 13.3 40.4 75.2 111.3 145.6 268.2 0.1 13.9 43.4 72.8 74.5 110.2 163.8 2 10.9 283.2 0.1 1.2 19.7 54.8 56.0 66.5 104.4 11 1.3 165.1 235.0 1.70 1.67 I .58 1 S O 1.36 2.28 2.23 2.13 1.91 1.66 1.45 1.28 2.87a 2.69 2.40 2.13 2.11 1.87 1.57 1.35 3.54 3.52 3.19 2.79 2.80 2.68 2.34 2.30 1.94 1.60 298.2 0.1 0.1 16.6 37.5 60.9 94.8 120.5 163.4 182.4 186.8 247.0 253.3 302.1 313.2 0.1 13.8 33.8 75.2 117.2 161.4 198.8 251.2 303.6 328.2 4.8 25.5 54.5 94.8 150.0 230.1 301.2 343.2 7.1 27.4 61.8 108.5 175.3 302.5 4.31 4.35b9 4.03c 3.62 3.29 3.0Y 2.73 2.45" 2.24 2.24 1.91 1.93" 1.67 5.01 4.76 4.37 3.75 3.25 2.88 2.59 2.29 2.04 5.70 5.21 4.70 4.1 1 3.45 2.84 2.44 6.63 6.00 5.29 4.5 1 3.80 2.79 a Mean oftwo experiments (2.88,2.87); mean of two experiments (4.34,4.37); high-pressure diaphragm cell experiment.volume method. The temperature dependence of A , is greater than is normally observed for non-polar liquids and indicates a breakdown of the rough hard-sphere theory .In principle, A , should be independent of both temperature and density. However, in associated systems such as methanol and water the increase of AD with temperature is attributed to the corresponding decrease in the influence of hydrogen bonding on the liquid structure. In pyridine, another fluid with a large dipole moment2236 HIGH-PRESSURE D I FFUS I 0 N 0 F ACE TON IT R I L E 5 .o 4 .O " I v) "€ 2 ' 3.0 0. \ 2.0 1 1 1 1 1 1 1 1 1 1 1 I 1 I 1 1 1 1 1 1 I 1 1 46 48 50 52 54 V/cm3 mol-' FIG. 1 .-Self-diffusion coefficient of acetonitrile as a function of molar volume. TABLE 2.-ROUGH HARD-SPHERE AND CORRELATION PARAMETERS FOR ACETONITRILE T/K &/cm3 mol-l a/nm AD &/cm3 mol-l a,/nm 268.2 29.9 0.41, 0.49, 30.2 0.41, 283.2 29.4 0.41, 0.50, 29.7 0.41, 298.2 29.2 0.40, 0.53, 29.2 0.41, 3 13.2 29.0 0.40, 0.53, 28.9 0.40, 328.2 29.0 0.40, 0.55, 28.7 0.40, (7.7 x C m), a strong variation of A , with temperature has been found and attributed to an undefined combination of both (a) changes in quasi-hydrogen-bonding interactions involving the nitrogen atom and ring protons and (6) a strong temperature effect on the re-orientational motions of the molecule.15 Acetonitrile is not associated and the possibility of its non-spherical shape being the cause of the variation of A , seems unlikely, since carbon disulphide has approximately the same shape and the value of A , for that liquid is 0.6 1 k 0.0 1 (268-3 13 K).The principal difference between acetonitrile and CS, is the large dipole moment (1 1.3 x 1 0-30 C m) of the former.BullR . L. HURLE AND L. A. WOOLF 2237 and Jonas15 have interpreted deuteron and nitrogen spin-lattice relaxation times in [2H,]acetonitrile as indicating that the intermolecular potential has little dependence on the orientation of the main symmetry axis. However, an analysis by Bertagnolli and ZeidleP of X-ray and neutron scattering data suggests that the axes of nearest neighbours in acetonitrile tend to be aligned at 90- 125' to that of the central molecule. In CS,, which has no dipole moment, the preferred orientation of nearest neighbours appears to be ~ara1lel.l~ According to Hirschfelder et al. l8 the effective potential energy of dipole-dipole interactions (for large separations) depends inversely on the temperature and the sixth power of the separation.Because A , is virtually independent of density at constant temperature, its variation in table 2 suggests that the large dipole moment of acetonitrile does have a significant influence on its self-diffusion. This inference provides support for Chandler's reasoning2 that strong, rapidly changing inter- actions between particles would adversely affect the validity of the rough hard-sphere theory. This is in contrast to Czworniak et al.13 who interpreted their mutual-diffusion data as indicating that the hard-sphere model was successful for systems containing molecules with dipole moments comparable to that of acetonitrile. An alternative use of hard-sphere theory to correlate self-diffusion data has been provided by Dymond.lg The experimental diffusion coefficient is used to obtain a reduced diffusion coefficient D* = [nD/(nD),,] (V/V,)g (6) where n is the number density for the experiment and the Enskog dilute fluid diffusion coefficient is defined by ( ~ Z D ) ~ = #(RT/zM)+/o2. In terms of experimental quantities A temperature-dependent equivalent hard-sphere correlation diameter oc is determined from the data by establishing the closest fit of D* to a common curve of D* against V/ y0 (= 2/2/no3).The curve obtained for the reduced acetonitrile data is shown in fig. 2 and the values of V/ 6 and oc are given in table 2. Over the five isotherms (268- D* = 1.744 x 106DV-~(M/7y. (7) I I I I / ' I 1 I I I - I I I I 1 1 1.3 1.5 1.7 1.9 FIG. 2.-Reduced self-diffusion coefficient of acetonitrile- A, 268; 0, 283; 0, 298; m, 313; 0, 328 K ; *, reduced diffusion coefficient of equivalent smooth hard sphere.2238 HIGH-PRESSURE DIFFUS I 0 N OF A CE TONI T R I LE 328 K) included in the fit only two points deviate by > 2% (but < 3%) from the curve given by the equation D* = - 0.429, + 0.179,( V/ ?(J + 0.103,( V / b)2.(8) The comparison in table 2 of oC with 0 determined from the rough hard-sphere approach shows that the two sets of diameters are in good agreement. This provides further evidence that the correlation treatment is insensitive to the shape of the molecule; a similar result was found for carbon disulphide. This is the first occasion on which the correlation approach has been used for a molecule with a large dipole moment. There is no indication that the dipole-dipole interactions adversely affect the correlation.This is in contrast to the methanol system20 where there was a less satisfactory correlation of the data and a much greater variation of the correlation diameter. There are few liquids in which the molecule has as large a dipole moment as acetonitrile. The present results show that the rough hard-sphere theory works well despite the obvious effect of dipole-dipole interactions. We thank Mr P. J. Back and Dr K. R. Harris for assistance with the experimental measurements and R. L. H. thanks the Australian Government for a post-graduate research award. L. A. Woolf, J . Chem. Soc., Faraday Trans. I , 1982, 78, in press. D. Chandler, J . Chem. Phys., 1975, 62, 1358. J. H. Dymond, personal communication (December, 1980). D. J. Tildesley, personal communication (March, 1981), see ref. (1). H. Bertagnolli, P. Chieux and M. D. Zeidler, Mol. Phys., 1976, 32, 759. K. Singer, A. Taylor and J. V. L. Singer, Mol. Phys., 1977, 33, 1757. K. R. Harris, R. Mills, P. J. Back and D. S. Webster, J . Magn. Reson., 1978, 29, 473. R. L. Hurle, Ph.D. Thesis (Australian National University, Canberra, 1981). M. A. McCool and L. A. Woolf, High Temp. High Pressures, 1972, 4, 85. and 1965), vol. I and 11. 10 J. Timmermans, Physico-chemical Constants of Pure Organic Compounds (Elsevier, Amsterdam, 1950 l1 K. R. Srinivasan and R. L. Kay, J . Solution Chem., 1977, 6, 357. l 2 M. D. Zeidler, Ber. Bunsenges. Phys. Chem., 1965, 69, 659. l 3 K. J. Czworniak, H. C. Andersen and R. Pecora, Chem. Phys., 1975, 11, 451. B. J. Alder, D. M. Gass and T. E. Wainwright, J. Chem. Phys., 1970, 53, 3813. l 5 T. E. Bull and J. Jonas, J. Chem. Phys., 1970, 53, 3315. H. Bertagnolli and M. D. Zeidler, Mol. Phys., 1978, 35, 177. I i 0. Steinhauser and M. Newmann, Mol. Phys., 1979, 37, 1921. J. 0. Hirschfelder, C. F. Curtis and R. B. Bird, Molecular Theory of Gases and Liquids (Wiley, New York, 1954), p. 28. l9 J. H. Dymond, Physica (Utrecht), 1974, 75, 100. 2o R. L. Hurle and L. A. Woolf, J . Chem. SOC., Faraday Trans. I , 1982, 78, in press. (PAPER 1 / 1547)

ISSN:0300-9599    

DOI:10.1039/F19827802233

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 2239-2249 Acetonitrile on Silica-Magnesia Mixed Oxides Temperature-programmed Desorption and Infrared Study BY GABRIELE RITTER.* HEINRICH NOLLER AND JOHANNES A. LERCHER Institut fur Physikalische Chemie, Technische Universitat Wien, Getreidemarkt 9, A- 1060 Wien, Austria Received 7th October, 193 1 The adsorption of acetonitrile on silica, silica-magnesia mixed oxides and magnesia was studied by means of temperature-programmed desorption (t.p.d.) and infrared spectroscopy. With t.p.d. three types of desorption behaviour were distinguished. On silica, only one comparatively narrow peak was detected (type 1). the mixed oxides containing up to 50mol % magnesia showed one broad peak with pronounced high-temperature tailing (type 2) and the mixed oxides containing 60-90 mol % magnesia showed up to 4 peaks (type 3).The i.r. results gave evidence for at least 3 adsorption structures: acetonitrile bound to surface hydroxyl groups, acetonitrile bound to cations and a surface carboxamide. Assignments of desorption peaks to the different adsorption structures of acetonitrile are suggested. Silica-magnesia mixed oxides have been studied by several authors, with special attention being paid to acid-base properties and related catalytic behaviour. The acidity was attributed to mixed-oxide phases, while the basic properties were related to the magnesia ~ 0 n t e n t . l . ~ Our aim was to investigate the change in the acid-base behaviour over the whole range of composition from silica to magnesia, whereas most of the former studies have dealt with only a few mixed oxides.Acetonitrile was used as a test molecule since it showed a number of well resolved peaks in the thermal desorption spectra (magne~ia/acetonitrile).~ Both e.p.a. (electron- pair acceptor) and e.p.d. (electron-pair donor) sites were involved and the adsorption structures seemed to be very sensitive to pretreatment conditions and hydration states on the surface. Although acetonitrile is a small molecule and should therefore be suitable for studying interactions with the surface, there are not many papers dealing with the adsorption of acetonitrile on oxide surface^.^-^ Most authors found that acetonitrile cannot be used for the characterization of surface properties because of its high reactivity.However, even if acetonitrile is not suitable for the characterization of surface properties, the interactions of nitriles with oxide surfaces of various acidities and basicities should be of potential interest with respect to possible ring-closure reactions and the formation of heterocyclic compounds.lo9 l1 EXPERIMENTAL MATERIALS To obtain the silica-magnesia mixed oxides, Aerosil (Fluka) and magnesium hydroxide (Fluka purum p.a.) were stirred in water at 90 O C for 17 h, centrifuged and dried to constant weight at 120 "C. The samples covered the range 10-90 mol % magnesia. The B.E.T. surface areas are listed in table 1. 22392240 ACETONITRILE O N SILICA-MAGNESIA MIXED OXIDES TABLE 1 .-ACIDITY ( H , d + 3.3) AND B.E.T. SURFACE AREAS OF SILICA-MAGNESIA MIXED OXIDES acidity mol % surface area ( H , d +3.3) sample magnesia /m2 g-I /mmol m-I silica I I1 I11 IV V VI VII VIII IX magnesia 0 10 20 30 40 50 60 70 80 90 100 330 190 240 320 340 190 150 130 143 143 81 0.0085 0.01 15 0.0131 0.01 53 0.0027 0.01 17 0.0 148 0.0075 0.00 16" a Colour change not very clear SiO, 20 40 60 80 MgO mol 70 FIG.1 .-Acidity of silica-magnesia mixed oxides (H, < + 3.3) determined by titration method with n-butylamine. Debye-Scherrer diagrams taken from samples calcined at 650 OC (3 h) showed forsterite besides magnesia in the range 60-90 mol % magnesia, whereas another solid phase, probably talc, was detected in addition to silica in the range 10-50 mol%. The acidity of the samples pretreated at 650 "C for 3 h was measured by titration with n-butylamine according to Benesi using 4-dimethylaminoazobenzene (pK, = + 3.3).The acidity values obtained this way are listed in table 1 and shown in fig. 1. Acetonitrile (Merck Uvasol) was used without any further purification.G. RITTER, H. NOLLER A N D J. A. LERCHER 224 1 APPARATUS TE M PER AT U RE-P ROGR A M ME D DESOR PT ION T.p.d. was carried out in U ~ C U O (ca. 1 Pa) with a heating rate of 10 "C min-' in the range 30-750 "C. The reactor was a quartz glass tube (10 mm in diameter), which was connected to a vacuum pump and to a quadrupole mass spectrometer for detecting the desorbed species (m/e = 1-100). The mass spectrometer and the t.p.d. furnace were coupled to a process com- puter IBM S/7. For further details see ref. (5) and (12). INFRARED MEASUREMENTS The i.r.cell was constructed according to the description of Knozinger et PROCEDURE TEMPER AT U RE-P RO G R A M M E D D E S OR P T I ON The sample (100 mg) was calcined in the t.p.d. reactor at 650 OC, cooled to room temperature and contacted with 10 mm3 acetonitrile for 15 min. The sample was then evacuated at room temperature for 30-60 min, before t.p.d. was started. INFRARED MEASUREMENTS The hydroxides were pressed into thin self-supporting wafers (p = lo8 Pa), which were examined using a transmission technique. Thermal pretreatment of the samples (500 "C for 2 h) was carried out in the reactor before adsorption of acetonitrile. The spectra were taken at room temperature and recorded with a Perkin-Elmer (type 325) grating spectrograph in the range 4000-1000 crn-l.The scanning speed was 0.5 cm-l s-' using slit programme 7, which corres- ponded to a resolution of 3 cm-' at 3600 cm-l. RESULTS AND DISCUSSION TEMPERATURE-PROGRAMMED DESORPTION The temperatures of the desorption maxima of acetonitrile are listed in table 2 and the t.p.d. spectra are shown in fig. 2. The maximum intensities of the t.p.d. curves were adjusted to 1000 mV (by multiplying the intensity values of the curves by a suitable TABLE 2.-DESORPTION OF ACETONITRILE FROM SILICA-MAGNESIA IN THE RANGE 30-750 O C (1 0 OC min-l). TEMPERATURES OF T.P.D. MAXIMA AND MULTIPLICATION FACTOR APPLIED TO INTENSITY VALUES OF T.P.D. CURVES IN FIG. 2. desorption maxima/OC mult. sample peak I peak 2 peak 3 peak 4 factor silica I I1 I11 IV V VI VII VIII IX magnesia 67 76 81 90 135 107 86 78 90 78 70 2.40 1.73 0.8 1 - 1.32 4.35 - 2.49 1.82 1.26 164 260 0.74 I56 247 - 150 244 345 1.38 144 240 326 0.52 - - - - - - - - - - - - ._ - - - - - 150-260 -2242 n 4, ACETONITRILE ON SI LI C A-M A GNE SI A MIXED OX IDES - E 0 - 9 h W P L 0 NG.RITTER, H. NOLLER AND J. A. LERCHER 2243 factor given in table 2) in order to demonstrate how the shape of the curves varied with the composition of the sample. Three types of desorption behaviour may be distinguished : A comparatively narrow peak with its maximum below 100 O C appeared on silica (type 1). The mixed oxides I-V, on the other hand, showed broad peaks with tailing on the high-temperature side (h.t. tailing) (type 2). The temperature of the maximum increased in the order silica, I, 11,111, V, with IV being considered separately because of its exceptional behaviour.Its maximum appeared at the highest temperature and its intensity was the smallest in spite of its high surface area (340 m2 g-l), and the shape of its t.p.d. curve is more symmetrical than those of I, 11, I11 and V. T/"C FIG. 3.-T.p.d. spectrum of magnesia/acetonitrile. (- - -) T.p.d. spectrum without adsorption of water before adsorption of acetonitrile (m/e 41 ; 100% = 2000 mV); (-) t.p.d. spectrum with adsorption of 1 mm3 water before adsorption of acetonitrile (m/e 41 ; 100% = 500 mV). VII-IX and magnesia showed up to 4 peaks (type 3), which are subsequently referred to as peaks 1-4 (table 2). The Debye-Scherrer pattern of VI was similar to those of VII, VIII and IX (magnesia and forsterite), therefore VI was expected to show desorption behaviour of type 3.However, only one well developed peak was observed, although there were slight indications of others. In all cases, only acetonitrile and water were found to be desorbed. With magnesia, t.p.d. was carried out twice with the same sample without further activation between the runs. The t.p.d. spectra were practically identical. So it seems unlikely that any product had been formed during the first run and had remained on the surface. When water (1 mm3) was adsorbed on magnesia before adsorption of acetonitrile, peak 1 was more pronounced than in the case without preadsorption of water (fig. 3). Note that the intensities of peaks 2 and 4 were greatly reduced whereas that of peak 1 was visibly enhanced.The largest relative increase was observed for peak 3.- - I 1 1 I 1 0 , l I l l l l l l l l l r l l l l l r l l . ~ ~ ~ ~ _ _ ~ I N F R A R E D MEASUREMENTS 1.r. spectra are shown in fig. 4-6. S U R F A C E H Y D R O X Y L G R O U P S O F T H E O X I D E S Pure silica and magnesia showed one band (at 3740 cm-l) in the OH stretching region after evacuation (10-l Pa) at 500 O C for 1 h. This band was attributed to the valence vibration of free hydroxyl groups. l4G. RITTER, H. NOLLER AND J. A. LERCHER 2245 100 80 h $ 60 u E * .- E $ LO * 20 0 I 1 , \ 1 1 I \ I \ I 1 I I I I I LOO0 3500 3000 2500 2000 1800 1600 1400 1200 wavenum ber/cm -' FIG. 6.--I.r. spectra of acetonitrile on magnesia. (-) Adsorption of acetonitrile at 4 x lo3 Pa at room temperature; (- - -) after evacuation (lo-' Pa) at room temperature; ( .. . .) after evacuation (10-l Pa) at 200 OC. TABLE 3 .-C-H VIBRATIONS (cm-l) OF ACETONITRILE ADSORBED ON SILICA-MAGNESIA MIXED OXIDES pressure of acetonitrile, temperature sample 4 x lo3 Pa, 25 "C 10-l Pa, 25 "C 10-l Pa, 200 "C silica 3000, 2950 - - I 3000,2950 (2970, 2950, 2920) (2690, 2920) I11 3000,2960,2940 2940 2950,2920 VII 2995,2960,2940 2960,2920 2960,2920 VIII 2990,2930 2960,2920 - magnesia 2980,2940,2920 2980,2960, 2920 2960,2920 The mixed oxides exhibited two hydroxyl bands (3740 and 3670cm-l) after evacuation (lo-l Pa) at 500 O C (fig. 5). The two bands showed almost equal intensity in samples I, I11 and VII, while the band at 3670 cm-l was much weaker in sample VIII.Since the free hydroxyl groups of silica and magnesia have approximately the same i.r. frequency, we concluded that the band at 3670 cm-l was due to hydroxyl groups of talc or forsterite (mixed-oxide phases). This conclusion is in accordance with the results of Koubowetz et aL9 and Kermarec et al.,15 who found the band at 3670 cm-l to increase with magnesia content. The position of the 3760cm-l band was independent of the composition of the sample. We conclude from this that the environment of that (free) hydroxyl group is similar in both solid phases (talc and forsterite).2246 A C ETON I T R I LE ON S I L I C A-M AG NES I A MI XED OX IDES ADSORPTION OF ACETONITRILE The wavenumbers of CH vibrations and CN vibrations are listed in table 3 and 4.Silica. After adsorption of acetonitrile at 40 mbar* the free hydroxyl band of silica (3740cm-l) was weakened and a perturbed, broad hydroxyl band at 3435cm-l appeared (fig. 4), indicating that hydroxyl groups were involved in the adsorption process . TABLE 4.-C-N VIBRATIONS (cm-l) OF ACETONITRILE ADSORBED ON SILICA-MAGNESIA MIXED OXIDES pressure of acetonitrile, temperature sampie 4 x lo3 Pa, 25 O C 10-1 Pa, 25 O C 10-1 Pa, 200 OC silica 2295, 2261 I 2318, 2290, 2260 23 12,2290,2250 I11 1700, 1445, 1375 2310,2282,2245, (2200) VII 1665, 1445, 1375 VIII 2305,2280,2245 2190,2150 - - 2315, 2290, (2260) - 2310,2282 2315,2285, 1700, 1610 1490, 1447, 1375 2195, 1670, 1590, 1450 1700, 1445, 1370 2200, 2190, 1667, 1445 2190,1640, 1570 - (2305), 2295,2275,2250 magnesia 2150, 1610, 1530, 1390 1375, 1368, 1325, 1190 2180, 1610, 1600, 1570 1530, 1457, 1412, 1390 1378, 1192 2180,2160, 1600, 1530 1414, 1390, 1378, 1358 The CH valence bands of the adsorbed acetonitrile were found at 3000 and 2950 cm-l, similar to the bands in the liquid phase (3000 and 2942 cm-l).Two bands (2295 and 2261 cm-l) were observed in the CN stretching region. The band at 2295 cm-l was attributed to a combination band of the symmetrical CH, deformation vibration and the CC stretching vibration, and that at 2261 cm-l was attributed to the CN stretching vibration.16 The corresponding bands in the liquid phase were found at 2293 and 2254cm-l. These bands are known to shift to higher wavenumbers when acetonitrile interacts with an electron-pair acceptor. The stronger the interaction, the larger the shifts.Therefore the interaction of acetonitrile with the surface hydroxyl groups of silica must be assumed to be only slightly stronger than the interaction of acetonitrile molecules with each other (in the liquid phase). After evacuation at room temperture (10-l Pa) the bands due to acetonitrile disappeared and the original intensity of the free hydroxyl band was restored. The low desorption temperature as well as the slight upward shift of the bands in the CN region indicated that acetonitrile interacted rather weakly with silica. Mixed-oxides Iand III. After evacuation at 500 "C for 1 h, acetonitrile was adsorbed at 40 mbar. Only the band at 3740 cm-l was reduced while that at 3670 cm-l was not affected. This might be explained by the asumption that the hydroxyl groups, which gave rise to the band at 3670 cm-l, were not accessible for acetonitrile molecules.The perturbed hydroxyl bands were found at 3450 cm-l with sample I and at 3440 cm-l * 1 bar = lo5 Pa.G. RITTER, H. NOLLER AND J . A. LERCHER 2247 with sample 111. The smaller shift of these hydroxyl bands (295 and 300crn-', respectively) in comparison with that of silica (305 cm-l) indicated that the acid strength of the mixed oxides is lower than that of silica. On sample I the CH valence bands of acetonitrile were at 3000 and 2950 cm-l, which suggested surface properties similar to those of silica. On sample 111, bands occurred at 3000,2960 and 2940 cm-l, indicating that there must be more than one adsorption structure.The larger shift to lower frequencies (2940 cm-l compared with 2950 cm-l) must be caused by a stronger interaction of the methyl group with the surface oxygen.17 Three bands situated at 2318, 2290 and 2260 cm-l with sample I and at 2312,2290 and 2250cm-l with sample I11 were observed in the region of the CN stretching vibrations. According to Krietenbrink16 vibrations of weakly adsorbed acetonitrile and vibrations of acetonitrile coordinated to cations contribute to the band at 2290 cm-I. The bands at 2318 cm-l (and 2290 cm-l) (I) and the bands at 2312 cm-l (and 2290 cm-l) (111) were attributed to acetonitrile molecules coordinated to cations, while those at 2260cm-l (and 2290cm-l) (I) and 2250cm-l (and 2290cm-l) (111) should be due to acetonitrile adsorbed on hydroxyl groups.The band at 2260 cm-I (2250 cm-l) disappeared after evacuation at room temperature, which again indicates the weak interaction of acetonitrile with surface hydroxyl groups. The downward shift in the frequency of the CN stretching vibrations from sample I to sample I11 indicates weaker interaction with cations and hydroxyl groups, which is interpreted in terms of decreasing acid strength. The bands at ca. 1445 and 1370 cm-l were also due to adsorbed acetonitrile molecules. Mixed-oxides VII and VIII and magnesia. For samples VII and VIII the band at 3740 cm-l (due to magnesia or silica hydroxyl groups) was perturbed after adsorption, while the band at 3670 cm-l remained unperturbed. The i.r. spectra of acetonitrile adsorbed on these samples and especially on magnesia showed several bands in the region of perturbed hydroxyl bands, for which Koubowetz et aZ.9 provided a detailed discussion.The bands of the CH stretching vibrations were shifted to lower wavenumbers with increasing magnesia content (table 3). Since that shift, found on all mixed oxides, seems to be too large for a long-range effect, a direct interaction of the methyl group with surface oxygen was assumed. An analogous conclusion was drawn for the adsorption of acetone on magnesialS and was supported by results of Takezawa and K0baja~hi.l~ The CN vibrations at 2310, 2282 and 2245 cm-l on sample VII and at 2305, 2280 and 2245 cm-l on sample VIII were found only during adsorption of acetonitrile at 4 x lo3 Pa. After evacuation at room temperature these bands disappeared.Compared with samples I and 111, acetonitrile bound via cations has less thermal stability. The reason for the disappearance of the CN vibrations could be due either to a decrease in the acid strength (e.p.a. strength) of the cations or to a transformation of the initial surface structures. The decrease in the e.p.a. strength of the cations was shown by the decrease of the frequency for cationically bound acetonitrile (2310 and 2305 cm-l). Further bands situated between 2150 and 2200 cm-l were observed on these samples. The bands were thermally more stable than the CN bands at 2310 and 2305 cm--l. Their intensity increased with increasing magnesia content. This led to the assumption that they were caused by acetonitrile molecules bound to the surface via oxygen and a cation; such adsorption might be a precursor of an acetimidato anion on the surface.In this case the CN stretching frequency was expected to be lowered. The suggestion of Krietenbrink and Knozinger7t l6 that the band at ca. 2200 cm-l is due to condensation products seems unlikely, since acetonitrile was the only desorption product.2248 ACETONITRILE O N SILICA-MAGNESIA MIXED OXIDES After evacuation the band shifted to 2190 cm-I on samples VII and VIII and to 21 80 cm-I on magnesia. Raising the temperature caused further shifts (table 4) of these bands which persisted up to 400 "C. The formation of precursors of acetamide leads us to expect a CN double bond as well as a CO bond. The band situated at 1667 cm-l (sample VII), 1640 cm-l (sample VIII) and 1610 cm-l (magnesia) could be due to a CN double bond.The stronger the bond between the carbon atom and the surface oxygen is, the weaker the CN bond and the lower its vibration frequency. A similar intermediate structure was found in acetonitrile hydrolysis in aqueous medium.19 Some of the acetonitrile molecules seemed to form carboxamide structures on the surface after evacuation at higher temperatures. The antisymmetrical vibrations of such carboxamide structures were found between 1600 and 1580 cm-l, while there was some uncertainty regarding the symmetrical mode, which should be located between 1510 and 1500 cm-l. Nevertheless, note that a water molecule would be needed to form acetamide and it cannot be provided by the hydroxyl group of the surface.Similarly acetone forms carboxylate structures on oxide surfaces, while the desorption product is acetone.'* CORRELATION OF T.P.D. PEAKS WITH ADSORPTION STRUCTURES SILICA The only adsorption structure which was found in the i.r. spectra was acetonitrile bound to surface hydroxyl groups. Therefore the single t.p.d. peak at 67OC is attributed to the desorption of acetonitrile attached to hydroxyl groups. However, there is some discrepancy. The i.r. bands corresponding to adsorption structures via hydroxyl groups disappeared after evacuation at room temperature while the t.p.d. peak attributed to the same surface species appeared at higher temperatures. This behaviour, however, may be due to the difference in the conditions of the t.p.d. and i.r.experiments, in particular to the vacuum for the i.r. experiment being approximately one order of magnitude better than that of the t.p.d. apparatus. It is possible that this difference between t.p.d. and i.r. results is due to most of the weakly bound species being pumped off in the i.r. experiments. MIXED-OXIDES I, 11, 111, I v AND v With samples I, 11, I11 and V a strongly asymmetric t.p.d. peak was found, the maximum temperature of wh-ich went from 76 to 107 OC. The i.r. data (I and 111) give evidence of acetonitrile being bound via both hydroxyl groups and cations. Instead of two peaks (one for the hydroxyl groups, one for the cation), only one peak, with broad h.t. tailing, appeared. Whatever the reason for such tailing may be, e.g. lateral interactions or distribution of the adsorption strength, centres with higher adsorption strength than that of hydroxyl groups must have been formed.Note that the tailing extends up to > 400 O C , whereas on silica desorption is practically complete by ca. 150 O C . As hydroxyl groups can be ruled out as adsorption sites (shown by the h.t. tailing of the desorption peak) it is very likely that cations are involved in this strong adsorption. On silica the number of these cationic sites is practically zero, because of the high shielding of the small Si4+ cation by four 02-, as discussed in former papers. If in the mixtures, especially those with low magnesia content, Mg2+ is substituted for Si4+, the stoichiometry of the compound is changed, in such a way that some oxygen can be omitted, which would mean that the cations are less shielded and hence more accessible to donor molecules.G.RITTER, H. NOLLER A N D J. A. LERCHER 2249 The desorption maximum attributed to desorption from hydroxyl groups increases in the order I < I1 < I11 < V, although the acidic strength of the surface e.p.a. centres decreases in this order. In our opinion a further interaction must be assumed. We believe that interaction of (basic) surface oxygen with CH, groups of acetonitrile contributes to the adsorption of acetonitrile. This type of interaction will become more important as the magnesia content of the sample is raised. Indeed, the i.r. spectrum reveals interaction with the CH, group, shown by the shift of wavenumber, which increased with increasing content of magnesia. Thus we have to take into account two interactions, one decreasing and the other increasing with increasing magnesia content.MIXED-OXIDES VII A N D VIII A N D MAGNESIA Of the four t.p.d. peaks, that with the lowest maximum temperature was attributed to acetonitrile interacting with hydroxyl groups. Beside the low temperature of desorption, the increase of the peak when water was adsorbed before acetonitrile supports our interpretation. The surface species which caused the i.r. absorption between 2200 and 21 50 cm-l was thermally less stable than the surface carboxamide species. Therefore peak 2, which appeared at a lower temperature than peak 3, was attributed to the adsorption form absorbing between 2200 and 2150 cm-l, and peak 3 was assigned to the carboxamide species.For peak 4, we have not yet found a valid interpretation. We thank the ‘ Fonds zur Foerderung der wissenschaftlichen Forschung ’ for providing the i.r. spectrometer. H. Bremer and K-H. Steinberg, 4th Int. Congr. Catalysis, Moscow, 1968 (Akademiai Kiado, Budapest, 1971), p. 1377. H. Niiyama, S. Morii and E. Echigoya, Bull. Chem. Soc. Jpn, 1972, 45, 655. Y. Okamoto, T. Imanaka and S. Teranishi, Bull. Chem. Soc. Jpn, 1973, 46, 4. H. Vinek, H. Noller, J. Latzel and M. Ebel, Z . Phys. Chem. N.F., 1977, 105, 319. G. Hatzl, Diplomarbeit (TU Wien, 1978). H. Knozinger, Forschungsberichte Wehrtechnik, Luft- und Raumfahrt, 1976, 2, 53. H. Krietenbrink and H. Knozinger, 2. Phys. Chem., 1976. 102, 43. G. Della Gatta, J. Calorim. Anal. Therm., 1976, 7, 1/5/1. F. Koubowetz, J. Latzel, H. Noller, J . Colloid Interface Sci., 1980, 74, 322. lo T. Tokuyama, J. Pharm. Sac. Jpn, 1955, 75, 957. l 1 S. Pasynkiewicz, Pure Appl. Chem., 1972, 30, 509. l 2 J . Latzel and G . Kaes, React. Catal. Lett., 1978, 9, 2, 183. l 3 H. Knozinger, H. Stolz, H. Buhl, G. Clement and W. Meye, Chem. Ing. Tech., 1970, 42, 548. l 4 F. H. Van Canwelaert, P. A. Jacobs and J. B. Uytterhoeven, J. Phys. Chem., 1972, 76, 1434. l6 H. Krietenbrink, Diplomarbeit (Ludwig-Maximilian Universitat, Miinchen, 1974). l8 J. A. Lercher, H. Noller and G. Ritter, J. Chem. Soc., Faraday Trans. I , 1981, 77, 621. l9 H. R. Christen, Grundlagen der organischen Chemie (Sauerlander Aarau, Diesterwag Salle Frankfurt- 2o V. Gutmann, The Donor-Acceptor Approach to Molecular Interactions (Plenum Press, New York, M. Kermarec, M. Briend-Faure and D. Delafosse, J. Chem. Soc., Faraday Trans. I , 1974, 70, 2180. N. Takezawa and H. Kobayashi, J. Catal., 1973, 28, 335. am-Main, 1975), p. 619. 1978). R. T. Sanderson, Chemical Bonds and Bond Energy (Academic Press, New York, 1976). (PAPER 1 / 1558) 7 3

ISSN:0300-9599    

DOI:10.1039/F19827802239

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 2251-2257 Surface Tension of Per fluor opropane, Perfluoro-n-but ane, Perfluoro-n-hexane, Perfluoro-oc tane, Perfluorotributylamine and n-Pentane Application of the Principle of Corresponding States to the Surface Tension of Perfluoroalkanes BY I A N A. MCLURE,* V I R G ~ L I O A. M. SOAREST AND BERYL E D M O N D S ~ Department of Chemistry, The University, Sheffield S3 7HF Received 14th October, 198 1 The orthobaric surface tensions of C,F,, n-C,F,,, n-C,F,,, C8FIB, (C,F,),N and n-C,H,, have been measured by the differential capillary-rise technique over various ranges of temperatures. The results for the perfluoroalkanes have been analysed in terms of the van der Waals equation, the Brock and Bird equation incorporating Pitzer’s acentric factor and the phenomenological corresponding-states treatment of Patterson and Rastogi; the data are fitted with varying degrees of success. The phenomenological principle of corresponding states to which the surface tensions for the perfluoroalkanes conform is not identical to that for the n-alkanes and some other homologous series, but it is similar to it, in particular in the need for at least three reduction parameters.The most successful approaches to the interpretation of both the surface thermo- dynamics and the bulk thermodynamics of chain-molecule liquid mixtures rest on the assumption of the conformity of appropriate properties of both the pure components and the mixture to the principle of corresponding states.l12 For mixtures the demonstration of conformity is not straightforward since the assumption of conformity is usually intractably involved with the particular form of mixture theory necessarily introduced at this stage of the analysis.By contrast, the demonstration of the conformity of the pure substances is reasonably simple, requiring essentially a knowledge of suitable physical properties of the pure substances over wide ranges of temperature; for a very demanding test, density data over very wide ranges of temperatures are required. To assist in the analysis of the results of our measurements of the surface tension of some alkane + perfluoroalkane liquid mixtures3$ * we have measured the orthobaric surface tension of perfluoropropane, perfluoro-n-butane, perfluoro-n-hexane, per- fluoro-octane, perfluorotributylamine and n-pentane.We have analysed the results with varying degrees of sophistication in terms of the principle of corresponding states. We have compared the outcome of this analysis with that performed for the n-alkanes and other substances by Patterson and Rastogi.5 t Present address: Instituto Superior Tecnico, Centro de Quimica Estrutural, Complexo I, Av. Rovisco 9 Present address: The Institution of Chemical Engineers, 165-171 Railway Terrace, Rugby CV 21 3HQ. Pais, 1096 Lisbon Codex, Portugal. 13-2 225 12252 SURFACE TENSION OF PERFLUOROALKANES EXPERIMENTAL APPARATUS A N D PROCEDURE The measurements of surface tension were carried out in a closed Pyrex glass cell using the differential capillary-rise technique. The technique was chosen for two reasons.The first is that it yields accurate surface tensions for liquids that wet, as do all those studied in this programme, the walls of the capillaries. The second is that it readily lends itself to closed-cell operation, thus preventing the loss of the materials under study and the ingress of unwanted substances, notably air and (as here for sub-ambient temperature work), water. Avoiding losses is important if, again as here, the materials are volatile, hard-to-obtain in acceptable purity and expensive; furthermore, for mixtures of liquids of differing volatility, for which the same apparatus has been used, losses can lead to undetected changes in composition. If unwanted substances are excluded from the cell the interpretation of the orthobaric data so produced is uncoloured by doubts arising from their presence at the gas-liquid interface. Details of the apparatus and the procedure used are available elsewhere.6, During each measurement the temperature was held constant to within & 0.1 K for temperature below 273 K and f 0.02 K at higher temperatures.MATERIALS Perfluoropropane was supplied by Air Products Ltd with a claimed purity > 95 mol %. Perfluoro-n-butane was supplied by Fluorochem Ltd with a claimed purity > 97.8 rnol %. Perfluoro-n-hexane was supplied by the Imperial Smelting Co with an isomeric impurity level revealed by gas-liquid chromatography to be < 1 mol %. The perfluoro-octane was obtained from a variety of sources: K and K Laboratories, Peninsular Chemicals and Pfalz and Bauer.No sample was of high purity; gas-liquid chromatography suggested an isomeric impurity content of up to 10 rnol %. In view of dismal past experience of separating isomeric perfluoroalkanes no attempt at purification was made and it was hoped that little adverse effect on the measurements would result. The perfluorotributylamine was obtained from Koch-Light Laboratories Ltd with a purity claimed to exceed 99 mol % in terms of perfluorinated (C,F,),N. Proton n.m.r. measurements revealed the absence of partially fluorinated substances. All materials were degassed vigorously in view of the high solubility of air in perfluorochemicals. For measurements below 303 K, Fisons Ltd n-pentane of purity 99.4 rnol % was used after distillation; for measurements above 303 K, Phillips research grade (lot no.1789) of purity > 99.9 mol % was used; both samples were dried with sodium and degassed before measurement. RESULTS The surface tension o was calculated from the equation 0 = ri rj dg[3Ahij + (ri - r j ) ] / 6 (ri - r j ) where Ahij is the difference in the height of the menisci in the capillaries of radius ri and rj, d is the density of the liquid phase and g is the acceleration of free fall in our laboratory (9.813 42 m s - ~ ) , for which an estimate was supplied by the Geological Survey in London. We believe that the error in our results is of the order k0.05 mN m-l for surface tensions reported to within kO.01 mN m-l and of the order f 0.1 mN m-l otherwise. The results obtained at different temperatures for each substance were fitted by a least-squares procedure to two types of equation, namely o = a-bT (2) and ( 3 ) where a, b, oo and p are empirical parameters and T, is the gas-liquid critical temperature of the substance taken from the compilation of Ambrose and Townsend.8I. A.MCLURE, V. A. M. SOARES A N D B. E D M O N D S 2253 TABLE 1 .-ORTHOBARIC SURFACE TENSIONS T/K a/mN m-l a[eqn (2)]/mN m-l o [eqn (3)]/mN m-l a(lit)/mN m-l 233.9 240.9 246.9 253.1 260.1 268.9 233.5 237.7 243.4 247.1 25 1.4 258.9 261.9 265.4 298.0 303.0 3 13.0 3 18.0 323 328 338 303 306 313 318 328 333 337 303 313 318 323 237.9 241.5 249.4 256.7 261.5 266.9 303.0 3 13.0 10.97 9.99 9.33 8.5 1 7.77 6.92 14.00 13.62 12.88 12.41 11.96 11.17 10.94 10.60 11.34 11.0 10.0 9.55 9.1 8.6 7.4 perfluoropropane [density from ref.(9)] 10.89 10.89 10.03 10.01 9.33 9.3 1 8.62 8.60 7.80 7.80 6.80 6.84 13.97 14.00 13.52 13.54 12.40 12.91 12.50 12.52 12.04 12.05 11.23 11.25 10.90 10.92 10.52 10.55 11.46 1 1.46 10.97 10.94 9.99 9.92 9.50 9.42 9.01 8.92 8.52 8.44 7.59 7.49 perfluoro-n-butane [density from ref. (1 O)] perfluoro-n-hexane [density from ref. (1 l)] perfluoro-octane [density from ref. (9)] 13.4 13.41 13.37 13.2 13.18 13.13 12.6 12.64 12.59 12.3 12.26 12.21 11.5 11.49 11.44 11.1 11.11 11.06 10.8 10.80 10.76 16.3 16.30 15.4 15.40 14.95 14.95 14.5 14.5 22.23 22.23 22.22 perfluorotributylamine [density from ref. (1 3)] n-pentane [density from ref. (1 5)] 21.89 21.81 20.81 20.9 1 20.11 20.07 19.56 19.52 18.85 18.89 1 5.0a 14.75 13.9a 13.68 21.80 20.88 20.05 19.50 18.88 14.88 13.80 10.9811 10.0511 9.1 1" 1 3.912 12.712 11.312 16.114 22.1 316 21 .7416 20.8216 20.0616 19.5316 1 8.9416 14.9416 13.8416 a These points were not included in the fitting procedure for eqn (2) or (3).2254 SURFACE TENSION OF PERFLUOROALKANES TABLE 2.-vALUES OF CONSTANTS OF EQN (2) AND (3) a b OSD 6 0 IU OSD substances ~~ C3F8 37.88 0.1156 0.10 43.07 1.22 0.09 n-C,F,o 39.23 0.1084 0.08 42.97 1.21 0.08 51.4 1.37 0.15 n-C,F,, 36.65 0.07668 39.1 1.16 0.1 n-C,H,, 49.56 0.1149 0.06 52.34 1.21 0.06 nGF1, 40.62 0.09786 - (C,F,)3N 43.57 0.0900 - - - - - In table 1 we show the measured values of the surface tensions together with the values calculated from (2) and (3).In table 2 we present the values of a, b, a. and ,u obtained for each compound and the respective standard deviations asu for the calculated surface tensions.No surface tensions have been reported previously for C3F, or n-C,F,,. Skripov and Firsov17 present surface tensions for n-C,F,,, n-C,F,,, n-C,F,,, n-C,F,,, n-C,F,, and n-C10F2,. Their surface tensions are low for n-C,F,, and n-C,F,, and show a temperature dependence different from those of other authors. Since they assumed the constancy of the parachor with temperature it seemed to us that if the densities they used were corrected better agreement might be found. We tried to recalculate their surface tensions for n-C,F,, using correct density values but the improvement was negligible. A close inspection of the boiling points of their material shows them to be lower than those reported by Ambrose and Townsend by as much as 2-12 K, suggesting that the material was highly impure.Therefore we do not include the data of Skripov and Firsov in our discussion. For n-C,F,, the agreement between our results and those of Stiles and Cadyll is very good. The comparison of our results for n-C,F,, with those of Haseldine and Smith12 shows also fairly good agreement, giving some support to our hope that the effect of the presence of the known impurity in our material is small. DISCUSSION The simplest form of expression for the principle of corresponding states for the surface tension of simple fluids is where the function 6(0 of reduced temperature Fis universal for the reduced surface tension 6. The first explicit expression was that of van der WaalP 6 = 6(F) 6 = a/a* = [l -(T/?3]P where o* and p are constants.This equation is a statement of a two-parameter principle of corresponding states. Using combinations of gas-liquid critical constants for a*, simple dimensional analysis yields three choices of o* proportional to V;2/3, P, V:I3 or Tk/3pEi3. Guggenheim found that using the first choice the critical index p for the rare gases is close to 1 1/9.19 For substances similar to the rare gases, for which the critical compressibility 2, is a constant, the choice of form of a* in terms of critical constants is a matter of indifference. For more complex substances 2, is not a constant, and different choices of the form of o* give rise to different functional dependences of a* on F.I. A. M C L U R E , V. A. M. S O A R E S AND B.E D M O N D S 2255 For the perfluoroalkanes we find that the dependence of o on T is equally well described by eqn (2) or (3). This is, however, scarcely an exacting test since at T/T, far from unity the expansion of eqn (3) in powers of T / q produces an expression essentially linear in T. However, within the limits of the test p is not excessively far from the currently accepted best value of 1.26 for T/ q close to unity. More seriously from the point of view of the principle of corresponding states, no combination of critical constants of the kind given above for o* produces a universal function of T/ T, for the perfluoroalkanes studied here. Clearly, therefore, recourse to a more complicated form of corresponding-states principle must be taken.We now turn to a consideration of the most simple of these forms of corresponding- states principle, which involves three parameters characteristic of the substances. Three parameters have been found necessary to bring the bulk properties of n-alkanes, linear dimethylsiloxanes and linear perfluorocarbons in to conformity with the principle of corresponding states. A convenient empirical t hree-parameter form of corresponding-states principle for the surface tension of complex substances has been described by Brock and Bird20 6’ = (a/mN m-1)(pc/atm)-2/3( T,/K)-l13 = (0.491 5 + 0.6529co)( 1 - T)l1/’ where o is Pitzer’s acentric factor defined as CL) = log [p,/lO p(T = 0.7 731. Fig. 1 shows 6’ as a function of co at = T/T, = 0.65 for perfluoropropane to perfluorononane.For comparison, the line of Brock and Bird is also shown. Lack of low-temperature data makes impossible a comparison with the test of Stie121 at = 0.6. For the sake of completeness, we include in fig. 1 even these data which do not enjoy our full confidence. The sources of data not reported in this paper for perfluoropentane are in ref. (22) and for perfluoroheptane and perfluorononane in ref. (12). The points for perfluoropropane fall close to the line of Brock and Bird, but those for perfluoro-octane and perfluorononane fall below the line by 8 and 12%, respect- ively: well above the Stiel criterion of 5% from ‘normal’ liquids. It is unlikely, although not impossible, that these deviations are due to errors in the measurements of o or in the densities needed to obtain o from the observed capillary heights.The question of impurity is relatively simple to settle, since for perfluoropentane and perfluorohexanell at least, branching causes o to increase rather than, as here, to decrease. Although the data for perfluoroalkanes are few we believe that the deviation reveals the onset of the failure of the correlation of Brock and Bird to describe the surface tensions of this homologous series. By contrast the correlation works well for n-alkanes up to n-dodecane. A fuller test of the Brock and Bird expression at more than one temperature is afforded from the plot of the quantity o/pE13 Tk/3 (0.491 5+0.6539 co) against T/Tc. Not surprisingly, a universal curve is found for perfluoropropane to perfluoroheptane, but again the data for perfluoro-octane and perfluorononane fall below the line and increasingly so as T/T, increases.Our (admittedly limited) analysis of the surface tensions of the perfluoro-n-alkanes suggests that at least three reduction parameters are needed for a successful principle of corresponding states. Although the Brock and Bird equation formally meets this criterion it fails to give good predictions for the surface tensions of substances of high chain length or at reduced temperatures. A further objection, although one of principle only, is that co lacks a clear molecular significance similar to that enjoyed by pressure, volume and temperature reduction factors and their combinations.2256 SURFACE TENSION OF PERFLUOROALKANES T 0.23 -2 b. 0.2 1 0.20 I 1 I 1 0.3 0.4 0.5 0.6 w FIG.1.-Reduced surface tension cp;2/3 C1l3 plotted against the Pitzer acentric factor o at (T/T,) = 0.65 following the treatment of Brock and Bird. The points in increasing order of co are C,F, to n-C,F,,. A more searching phenomenological test, in which the reduction parameters do have some physical significance, is that described by Patterson and R a ~ t o g i . ~ In their analysis the reduced surface tension 6” = aPg3 ag3 is a universal function of the measure of reduced temperature apT, where a p is the isobaric thermal expansion coefficient and is the isothermal compressibility. Owing to lack of pVT data for the other perfluorocarbons our analysis using the method of Patterson and Rastogi has been restricted to perfluoro-n-hexane.In the absence of a good reason for discarding either of the discordant pVT data of Stiles and Cady and Dunlap and coworker^^^^^^ 6’’ was calculated using both sets and the results are shown in fig. 2. Clearly, although the results fall close to the curve of Patterson and Rastogi they do not fall on it within the scatter of the points for the substances considered by these authors. They also fall below the theoretical lines obtained by Patterson and Rastogi for various models of the liquid state. On the basis of this test, it looks as if perfluoro-n-hexane surface tensions do not obey exactly the same phenomenological principle of corresponding states as do those of the n-alkanes, linear dimethyl- siloxanes and other homologous series. Without more extensive p VT data little more on this matter can be added.The foregoing analysis suggests that the perfluoroalkanes do require a three- parameter corresponding-states treatment and that it is similar but not identical to that which applies to the n-alkanes and other chain molecule series. Further measurements of surface tension and p VT quantities on material of dependable purity and over wider ranges of temperature and density would be welcome to permit a test of acceptable stringency.I. A. MCLURE, V. A . M. SOARES A N D B. EDMONDS 2257 FIG. 2.-Reduced surface tension aapy;2/3 k-ll3 plotted against a T following the treatment of Patterson and Rastogi for perfluoro-n-hexane; (0) denotes points calculated using the density data of Rohrback and Cady and (0) denotes points calculated using the density data of Dunlap and Scott.The curves (a), (6) and (c) refer to the different models used by Patterson and Rastogi using the following ( m , n) choices for the intermolecular potential function: (a) (6, 12), (b) (3, co) or Flory model, (c) (6, a). 1 2 3 4 5 fi 0 9 10 11 1 2 13 I4 15 16 17 1.3 19 20 21 22 23 1 4 I. Prigogine, The Molecular Theory of Solution (North-Holland, Amsterdam, 1957), chap. 8. R. Defay, I. Prigogine, A. Bellemans and D. H. Everett, Surface Tension and Adsorption (Longmans, London, 1966), chap XII. I. A. McLure and B. Edmonds, J . Chem. Soc., Faraday Trans. I , to be submitted. V. A. M. Soares and I. A. McLure, J . Chem. Soc., Faraday Trans. 1, 1982, 78, in press. D. Patterson and R. A. Rastogi, J . Phys. Chem., 1970, 74, 1067. B. Edmonds, Ph. D. Thesis (University of Sheffield, 1973). J. C. G. Calado, I. A. McLure and V. A, M. Soares, Fluid Phase Equilibria, 1978, 2, 199. D. Ambrose and R. Townsend, Vapour-Liquid Critical Properties (National Physical Laboratory, London, 1975). J. A. Brown, J . Chem. Eng. Data, 1963, 8, 106. J. H. Simons and J. W. Mausteller, J. Chem. Phys., 1955, 20, 1516. V. E. Stiles and G. M. Cady, J . Am. Chem. Soc., 1952, 74, 377. R. N. Haseldine and F. Smith, J. Chem. Soc., 1951, 603. B. Edmonds and I. A. McLure, J . Chem. Eng. Data, 1977, 22, 127. T. M. Reed, J . Phys. Chem., 1959, 63, 1798. International Critical Tables (New York, 1st edn, 4th impression, 1928), vol. 3. J. J. Jasper, J . Phys. Chem. ReJ Data, 1972, 1, 841. V. P. Skripov and V. V. Firsov, Russ. J . Phys. Chem., 1968, 42, 653. J. D. van der Waals, 2. Phys. Chem., 1894, 13, 656. E. A. Guggenheim, J . Chem. Phys., 1945, 13, 253. J. A. Brock and R. B. Bird, AIChE J., 1955, 1, 174. L. I . Stiel, Ind. Eng. Chem., (London), 1968, 60, 50. G. H. Rohrback and G. H. Cady, J . Phys. Chem., 1979, 71, 1938. R. D. Dunlap, R. G. Bedford, J. C. Woodbary and S . D. Furrow, J . Am. Chem. SOC., 1959,81,2927. R. D. Dunlap and R. L. Scott, J. Phys. Chem., 1962,66, 631. (PAPER 1 / 1603)

ISSN:0300-9599    

DOI:10.1039/F19827802251

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Furaduy Trans. I , 1982, 78, 2259-2263 Photoinduced Electron Transfer from Triphenylmethyl Anion or Triphenylsilyl Anion to p-Terphenyl BY OSAMU ITO,* TAMOTSU ARUGA AND MINORU MATSUDA Chemical Research Institute of Non-aqueous Solutions, Tohoku University, Katahira, Sendai 980, Japan Received 19th October, 198 1 The steady illumination of triphenylsilyl anion and p-terphenyl (p-TP) in tetrahydrofuran yielded the radical anion of p-terphenyl (p-TP'-), which persisted after cutting off the light. When the triphenylmethyl anion was used, p-TP'- was not observed with the steady illumination; the finding that p-TP'- observed transiently by flash photolysis decayed at a diffusion-controlled rate suggests that the stable triphenylmethyl radical, produced after the donation of an electron to p-TP, accepts an electron from p-TP'-.On the other hand, since the triphenylsilyl radical is able to produce the disilane whose electron-acceptor ability is less than that ofp-TP, back electron transfer fromp-TP'- to the disilane does not occur and thusp-TP' - persists. The photoejection process from organic anions and the electron-trapping mechanism of aromatic compounds have been studied extensively in relation to the intermediates of photostimulated SRN 1 rea~tionsl-~ and the ion-pair phenomena.6* If back electron transfer occurs after the photoinduced process, then we can control the lifetime of the electron or the radical anions of the aromatic compounds by controlling the illumination. We showed in our previous paper that the radical anions of aromatic compounds were formed efficiently by illumination of the phenylthiolate anion and that the back electron-transfer rates correlate with the reduction potentials of the aromatic compounds.* In this paper, we report our findings that the radical anion of p-terphenyl (p-TP ' -) accumulates on steady illumination of the triphenylsilyl anion (Ph,Si-), whereas the radical anion is produced only transiently from the triphenyl- methyl anion (Ph,C-).The cause of these different phenomena is examined. EXPERIMENTAL Commercially available triphenylmethyl chloride, triphenylsilyl chloride and p-terphenyl (p-TP) were purified by recrystallization. Ph,C-,Na+ was produced by contacting Na metal with the chloride in degassed tetrahydrofuran (THF). Ph,Si,K+ was produced by contact with Na-K alloy.The light source was a 500 W high-pressure Hglamp. Absorption spectra of the steady-state solutions were measured by using a Cary 14 spectrophotometer. Absorption spectra of transient species were measured using a flash-photolysis apparatus of standard design ;g the half duration of the xenon flash lamps was ca. 10 ps and the flash energy was ca. 100 J. All measurements were made at 23 1 OC. RESULTS AND DISCUSSION Ph,Si-,K+ was identified from the absorption band at 360 nm [fig. 1 (a)].1° No colour change was observed in the dark by the addition ofp-TP to Ph,Si-,K+. The solution changed to blue with steady illumination using light of wavelength > 3 10 nm [fig. 1 (c) and (41. The absorption bands at ca. 470 and ca.800 nm were attributed 22592260 1.c u S 0.5 2 D 0 ELECTRON TRANSFER T O P-TERPHENYL I I I r I I ( d ) !? ! ! ! ! .f ! ! ! . I \ .,! ! 400 600 800 wavelength/ nni FIG. 1.-Absorption spectra of (a) Ph,Si-,K+ (4 x lop4 mol dm-3) in THF; (b)p-TP (2.5 x lo-, mol dm-,); (c) steady illumination of the mixed solution of (a) and (b) with light of wavelength > 310 nm for ca. 1 h; ( d ) steady illumination with a 500 W high-pressure Hg lamp with light of wavelength > 3 10 nm for 15 min. Optical path of the cell is 2 mm. Insert: decay profile at 468 nm observed by flash photolysis of a mixed solution of (a) and (b) in a 100 mm cylindrical cell. FIG. 2.-Absorption spectra of p-TP'-(,K+); (a) at 23 "C and (6) at - 78 OC in THF. Concentrations of p-TP'-(,K+) is 2.5 x mol dm-3 in a 2 mm cell. to the contact ion-pair of p-TP'-.In fig. 2 the absorption peaks of p-TP'- observed at 23 "C were attributed to the contact ion-pair and those at low temperature to solvent separated ion-pair or free ion.11 The decrease in the absorbance at 360 nm in fig. 1 was ascribed to the consumption of Ph,Si-,K+. After cutting off the light, decay of the absorption bands of p-TP*-,K+ was not observed within ca. 2-3 h. On flash photolysis of the fresh solution, the decay of the absorption band ofp-TP' -, K+ was observed initially, as shown in the insert of fig. 2 . The absorption showing fast decay was also ascribed to p-TP'-,K+; the T-T absorption of p-TP was not observed under our experimental conditions. The initial concentration of p-TP' -,K+0.ITO, T. ARUGA AND M. MATSUDA 226 1 I I I 70 0 800 900 wavelength/nm FIG. 3.-Transient absorption spectrum observed by flash photolysis of Ph&-,Na+ ( lop4 mol dmp3) in the presence of p-TP ( 5 x lo-' mol dm-+) in THF; light of wavelength > 3 10 nm was used. The absorbances 30 p s after flash were depicted. Insert: the second-order plot of the decay curve ofp-TP'-(,Na+) at 845 nm (optical path is 100 mm). produced by one flash was estimated to be ca. 4.3 x mol drn-,; ca. 40% ofp-TP'-, K+ decayed within 300ps and the rest persisted for a long time. With steady illumination of the solution containing Ph,C-,Na+ and p-TP, no colour change was observed. Flash photolysis of the solution yielded a transient absorption spectrum as shown in fig. 3. The absorption peak at 830 nm was ascribed to the contact ion-pair (p-TP'-,Na+) and the absorption peak at 915 nm to the solvent separated ion-pair (p-TP'-//Na+) or the free ion.Both species co-exist because the concen- trations of the ions are low (ca. lo-' mol dm-,). Decay of the transient p-TP'-(,Na+) was very rapid; the second-order plot of the decay curve was linear (insert of fig. 3) and the slope yielded k/E = 2.5 x lo5 cm s-l at 845 nm, where k and E refer to the rate of constant of the decay process and the molar extinction coefficient of p-TP'- (,Na+), respectively. The linear second-order plot suggests reaction between p-TP ' - (,Na+) and species of the same concentration as p-TP'-(,Na+). For both Ph,Si-,K+ and Ph,C-,Na+, dark electron transfer from these anions to p-TP did not occur; this suggests that the reduction potential of p-TP (- 2.36 V us.SCE)12 is more negative than that of Ph,Si' or Ph,C'. Photoinduced electron transfer may occur via direct photoejection from the anions [mechanism (l)] or via electron abstraction from the anions by the excited p-TP [mechanism (2)]. Sincep-TP'- was produced by exciting only the anion (Ph,M-), mechanism (1) is more probable than mechanism (2); however, on the flash photolysis of the anions in the absence ofp-TP, no absorption band for a solvated electron (ego,,) was observed in the visible region. (1) I Ph,M- A Ph,M. + e,,, esolv +p-TP -p-TP ' - hv p-TP -p-TP * p-TP* +Ph,M--p-TP'-+Ph,M'.2262 ELECTRON TRANSFER TO P-TERPHENYL The decay process of p-TP'- produced transiently by flash photolysis may include back electron transfer from p-TP'- to Ph,M' [reaction (3)], kb p-TP'- + Ph,M * ---L.I)-TP + Ph,M- (3) since the reaction system of Ph,C-, Na+ andp-TP was reproducible by successive flash photolysis.Decay of p-TP'- can be expressed by eqn (4) -d[p-TP'-]/dt = k,[p-TP'-] [Ph,M']. (4) If Ph,M' is not consumed by other reactions such as recombination to produce (Ph,M),, the concentration of Ph,M' is equal to that ofp-TP'- at any time; thus eqn (4) can be converted into eqn (5) The k value observed in fig. 3 was ascribed to the rate constant of back electron transfer (k,) from p-TP'-(,Na+) to Ph,C'; the value of k, was estimated to be 5 x lo9 mol-1 s-' using the value of E reported in the 1iterature.ll The value of k, is close to the diffusion-controlled limit.The decay curve of the initial part of fig. 1 may also be attributed to back electron transfer from p-TP'-(,K+) to Ph,Si; although the value of k, could not be estimated precisely, the rate constant may be similar to the above value. In the case of the photoreaction system of Ph,Si-,K+ and p-TP, Ph3Si' is known to be able to produce the disilane and disiloxane [reactions (6) and (7)]13-15 kr 2Ph,Si' - (Ph,Si), 2Ph,Si ' + THF- (Ph,Si),O etc. (7) The recombination rate constant (k,) close to the diffusion-controlled limit was estimated by analysing the decay curve of the 330 nm absorption band of Ph,Si' observed by flash photolysis of hexaphenyl disilane at 23 O C ; although the E value was not estimated precisely, the k, value was calculated to be 3.0 x lo9 dm3 mol-l s-l from 2k,/~ = 1.5 x lo5 cm s-l (at 330 nm) using the assumption E = lo4 dm3 mol-l cm-l.The ratio of the disiloxane to disilane was reported to be ca. 0.1 from the product analysis. 13-15 When (Ph,Si), was mixed with p-TP'-(,K+) produced by contact with K metal, electron transfer from p-TP'-(,K+) to (Ph,Si), was not observed in the dark. This suggests that the reduction potential of (Ph,Si), is more negative than that of p-TP. Our finding that p-TP' -(,K+) accumulates on steady illumination suggests that Ph,Si' formed by the donation of an electron is converted into the disilane (or the disiloxane) before the return of the electron from p-TP'-(,K+). Electron transfer in the dark was observed by mixing of Ph,Si-,K+ with perylene (El,, = - 1.66 V), anthracene (El,, = - 1.94 V) or pyrene (El,, = -2.06 V);13 thus aromatic compounds with more negative reduction potentials than that of Ph,Si' can be used as electron acceptors for the photoinduced electron-transfer system.The p-TP'- produced by the photoinduced process from anions such as the phenylthiolate and triphenylstannyl anions decayed at a diffusion-controlled rate; this is due to the fact that the reduction protentials of diphenyl disulphide (El,, z - 1.80 V)8 and hexaphenyl distannane (El,, = -2.30 V)16 are less negative than that of p-TP. In conclusion, we may summarize the conditions necessary for getting a high steady concentration of the radical anions with the photolysis of the anion as follows: (a) spontaneous electron transfer from the anion to aromatic compound does not occur in the dark and (b) the radical formed from the anion is rapidly converted into a substance having a more negative reduction potential than the aromatic compound.0.ITO, T. ARUGA AND M. MATSUDA 2263 P. B. Ayscough and F. P. Sargent, J . Chem. SOC., 1963, 5418. H. J. S. Winkler and H. J. Winkler, J . Org. Chem. 1967, 32, 1965. A, B. Pierini and R. A. Rossi, J. Organomet. Chem., 1978, 144, C12. H. A. Fox, Chem. Rec., 1979, 79, 253. ,3 E. E. van Tamelen, J. I. Brauman and L. E. Ellis, J . Am. Chem. SOC., 1971, 93, 6141. fi H. C. Wang, E. D. Lillie, S. Slomkowski, G. Levin and M. Szwarc, J . Am. Chem. Suc., 1977,99,4612. ' H. C. Wang, G. Levin and M. Szwarc, J . Am. Chem. SOC., 1978, 100, 6137. H. Tagaya, T. Aruga, 0. Ito and M. Matsuda, J. Am. Chem. SOC., 1981, 103, 5484. G. Porter and M. A. West, Techniques of Chemistry, ed. A. Weissberger (Wiley, New York, 1974), p. 367. lo A. G. Evans, G. Salamah and N. H. Rees, J . Chem. SOC., Perkin Trans. 2, 1974, 1163. l 1 K . H. Buschow, J. Dieleman and G. T. Hoijtink, J . Chem. Phys., 1965, 42, 1993. I 2 C. K . Mann and K. K . Barnes. Electrochemical Reactions in Non-aqueous Systems (Marcel Dekker, l 3 A. G. Evans, M. L. Jones and N. H. Rees, J. Chem. Soc. B, 1969, 894. l4 W. E. Fearon and J. C. Yong, J . Chem. SOC. B, 1971, 272. l 6 R. E. Dessy and R. L. Pohl, J . Am. Chem. SOC., 1968, 90, 2005. New York, 1970), chap. 5. H. Sakuri, in Free Radicals, ed. J. K. Kochi (Wiley, New York, 1973), vol. 11, chap. 25. (PAPER 1 / 1627)

ISSN:0300-9599    

DOI:10.1039/F19827802259

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chern. SOC., Faraday Trans. 1, 1982, 78, 2265-2274 Temperature and Density Dependence of the Self-diffusion Coefficient of n-Hexane from 223 to 333 K and up to 400 MPa BY KENNETH R. HARRIS? Diffusion Research Unit, Research School of Physical Sciences, Australian National University, Canberra, ACT 260 1, Australia Received 20th October, 198 I Self-diffusion coefficients for liquid n-hexane measured by the n.m.r. spin-echo technique are reported at 223, 248, 273, 298 and 333 K at pressures up to 400 MPa. Owing to a marked non-linear volume dependence, the results could not be directly fitted to the rough hard-sphere model of diffusion, but reduced diffusion coefficients were fitted as a function of a reduced volume, expressed as the ratio of the molar volume to a ' hard-core volume '.The temperature dependence of an equivalent hard-sphere diameter calculated from this hard-core volume was found to be adequately represented by an equation origmally proposed by Protopapas, Andersen and Parlee for liquid metals: CT = a,[l -B(T/T,):]. The constant B, 0.072, is of similar magnitude to that found to fit CH, and C,H, self-diffusion data, 0.069, and is consistent with diameters derived from the viscosity measurements of Brazier and Freeman. Between 273 and 333 K where the diffusion and viscosity results overlap, the group Dvt is constant, at a given temperature, with t = 0.97: consequently the Stokes-Einstein equation is not obeyed. As part of a continuing study of the effect of temperature and pressure on diffusion in dense liquids and gases, measurements of the self-diffusion coefficient of n-hexane, made between 223 and 333 K at pressures up to 400 MPa, are reported.and other laboratorie~~*~ has been concerned with 'simple' fluids of small rigid-core molecules, or water.5 The flexible nature of the n-alkane chain gives rise to an equilibrium mixture of conformers of different shapes which is liquid over a wide range of pressure. This equilibrium between the conformers is known to be pressure dependent,6 the percentage of gauche configurations increasing with pressure, whereas in the solid only trans configurations are found. It is therefore of interest to see whether these factors have any influence on the transport properties of chain molecules, and what deviations from the predictions of hard-sphere theories generally applied to the rigid-core molecules can be observed in this case.Most recent experimental work from this1* EXPERIMENTAL The n-hexane sample used was Merck (Darmstadt, Germany) Uvasol spectroscopic grade, distilled and dried over molecular sieves. A gas chromatographic analysis using a Supelco (Bellefonte, Penn., USA) column (5% SP-1200+ 1.75% Bentone 3) showed no evidence of unsaturated or non-homologous impurities. The self-diffusion coefficients were measured by the n.m.r. spinxcho method : the apparatus used and technique followed have been described in earlier papers.', The sample was confined in a glass and 316 stainless-steel bellows cell. In some of the experiments reported here, a phase-lock loop detector, linear to better than 0.2%, was used in place of the diode detector t Present address : Department of Physical and Inorganic Chemistry, University of Adelaide, Adelaide, South Australia 5001. 22652266 SELF-DIFFUSION OF LIQUID n-HEXANE of earlier work.? Calibration of the quadrupole coil used for generating the magnetic field gradient was carried out by measuring the positions of the first minima of the n.m.r.echo. Excellent agreement (1 %) was found between these results and results obtained earlier using tracer diffusion coefficients for water and benzene as standards.’ Temperatures were measured to f0.05 K with calibrated Pt resistors and pressures to f0.02 MPa using calibrated Heise Bourdon gauges (Dresser Industries, Newtown, Conn., USA.).$ RESULTS Density data from the literature*-l1 were fitted to an equation of state of the form 4 K = z aipi i=o ai = aoi + a I i / T where K is the linear secant modulus K = v,(P-Pa)/(v, - V ) (2) and Va is the molar volume at atmospheric pressure, pa.The temperature depend- ence of ai is based on that of the constant B in the Tait equation of state.12 This polynomial expression was found to give a better fit to the data than that equation which has only two parameters. The values of Vo were fitted to the equation 3 i - 0 Vo = C pi( T - 273.1 5)i (3) To reduce systematic errors in the amalgamation of several sets of data, values of K were calculated using the reported values of Yo for the different sets, but these were normalised relative to a datum at 298.15 K taken as a reference point (131 .612 cm3 m01-l)~~ before being fitted to eqn (3).The molar mass of n-hexane was taken as 86.178 g mol-l. The accuracy of the densities is ca. 0.2%. The coefficients aij and pi are given in table 1. This equation of state was used to obtain extrapolated densities above 200 MPa at 248 K and 150 MPa at 223 K. The experimental results are given in table 2 and, reduced by the factor 1 /d T, are shown plotted in fig. 1. The precision is estimated to be k 1.5% and, taking into account the calibration, the accuracy _+ 2.5%. The spin-echo measurements of Douglass et a/. obtained at atmospheric pressure between 182 and 333 K1* and at room temperature to 50 MPa15 are consistent with the data reported here. The average value of D at 0.1 MPa and 298.15 K, 4.18 x m2 s-l, lies between the two tracer values in the literature,l67 l 7 4.26, and 4.09, x lop9 m2 s-l.Three measurements were also made on n-heptane at atmospheric pressure using the sample described in ref.(16), theresultsbeing0.397,1.46and3.04 x lop9 m2 s-lat 195.5,250.0and299.7 K, respectively. These values are consistent with those of Douglass and McCalll* and the tracer value at 298.15 K, 3.016 x loy9 m2 sP1.l6 The parallel nature of the isotherms in fig. 1 suggests that the data could be fitted to an equation of the form D / d T = E l + % K/(l+E3/K) (4) It is a pleasure to thank Professor N. J. Trappeniers, Dr K. 0. Prins and Mr T. Jongeneelen of the Van der Waals Laboratorium, University of Amsterdam, for making available the design of this detector, and Mr P. Smith of the School Electronics Unit for its construction. $ The calibrations were made by the Division of Applied Physics, CSIRO, Sydney, N.S.W.It was found that the calibrations of Heise pressure gauges may differ by as much as 1 % from the factory calibration on delivery, and consequently these were checked periodically.K . R. HARRIS 2267 where with 298.15 K as the reference temperature, T,. Eqn (4) was more sxcessful than a four-term polynomial. The standard deviation of the fit, obtained by a non-linear least-squares method using weighting factors corresponding to the experimental errors for D, T and V, was 1.2%. That the fit was satisfactory is shown by the deviation plot, fig. 2. The coefficients E~ and 8i are given in table 3. V , = V - el( T - T,) - 0,( T - T,)2 ( 5 ) TABLE COEFFICIENTS FOR EQN (1) AND (2) 1 a0i Q l i Pi 0 -0.100 475 x 104 0.482 533 x lo6 0.127 292 x lo3 1 0.498 727 x 10' 0.605 314x 10' 0.162 834 3 0.403 108 x lop4 -0.866 104 x lop2 0.157 563 x 4 -0.241 280 x lo-' 0.493 288 x - 2 -0.208 482 x lo-' 0.428 990 x 10' 0.335 495 x 10-3 standard deviation 1 1 MPa 0.04 cm3 mol-' DISCUSSION COMPARISON WITH MOLE CU L A R-D Y N A M I CS CALCULATIONS FOR THE HA R D-S PHE R E F L U ID Some of the results reported here have been used in a discussion by Dymond and WoolP* of the diffusion of various solutes at infinite dilution in n-hexane solutions.The dependence of the ratio of the solute intra-diffusion coefficient to the solvent self-diffusion coefficient on the density and on the solute-solvent mass and size ratios was interpreted in terms of a rough hard-sphere model based on a comparison of experimental data with molecular-dynamics computer simulation results for tracer diffusion in the hard-sphere fluid It proved possible to interpret much of the self-diffusion data for rigid-core molecules using similar rough hard-sphere models.20 The density dependence of the self-diffusion coefficient of the hard-sphere fluid is reproduced by using a temperature dependent core diameter, 0, as an adjustable parameter.The diffusion coefficient, D,,,, calculated at a given state point using this hard-sphere diameter, is found to differ from the experimental value by a factor which is often independent of both temperature and density. This second disposable parameter is taken to represent the slowing of translational diffusion due to the inelastic collisions of real polyatomic molecules and it is known as the translational-rotational coupling constant, A .That it should be independent of density and very nearly independent of temperature for essentially spherical molecules has been argued by Chandler. 2o Thus the self-diffusion coefficient is given by D = ADsHs A d I where n is the number density and C is a 'correction' factor relating the self-diffusion coefficient, DSHS, obtained directly from molecular-dynamics calculations for the2268 SELF-DIFFUSION OF LIQUID II-HEXANE TABLE 2.-sELF-DIFFUSION COEFFICIENTS OF Il-HEXANE 248.15 273.15 298.15 223.15 0.1 48.7 51.4 100.7 150.4 200.5 252.9 292.3 0.1 0.1 25.8 55.7 100.8 154.4 197.5 250.3 292.9 305.0 308.0 350.6 391.2 0.1 25.0 25.2 99.8 152.3 198.7 248.6 249.5 302.8 344.0 393.8 0.1 0.1 0.1 0.1 27.6 50.0 101.4 105.3 153.0 200.1 249.5 290.8 295.8 299.1 353.5 8.347 8.649 8.664 8.89 1 9.075 9.229b 9.367b 9.460b 8.103 8.103 8.303 8.494 8.725 8.943 9.089 9.244b 9.356b 9.386b 9.395b 9.493b 9.584b 7.856 8.084 8.085 8.557 8.792 8.964 9.125 9.128 9.278 9.384 9.503 7.598 7.598 7.598 7.598 7.889 8.073 8.397 8.418 8.645 8.834 9.005 9.133 9.148 9.157 9.306 1.32" 0.88," 0.86," 0.6 1 0.433a 0.31," 0.1 7," 2.08 2.08 1.68 1.37 1.02 0.76, 0.59," 0.443" 0.358" 0. 342" 0.339" 0.263" 0.2 1 6" 2.96 2.47" 2.49 1.51" 1. 14" 0.9 1," 0.7 1 , 0.70ga 0.56," 0.47," 0.38," 4.15 4.21 4.21 4.16 3.32 2.85 2.08 2.05 1.66 1.32 1.07 0.89," 0.89," 0.86," 0.70," 0.22,"K. R.HARRIS 2269 TABLE 2.-continued 333.15 0.1 21.9 42.0 69.8 99.1 99.3 148.3 203.9 249.0 296.2 347.6 392.4 7.215 7.512 7.719 7.944 8.136 8.137 8.399 8.643 8.814 8.972 9.124 9.242 5.97 4.85 4.19 3.56 3.09 3.02 2.43 1.91 1.62 1.35 1.14" 0.962a a Experiments in which diode detection was used. The phase-lock loop detector was used Extrapolated densities lying outside region of data fitted by eqn (1) and for the remainder. 333K 3 110 120 130 V/cm3 mol-' 1c0 FIG. 1 .-Self-diffusion isotherms (D/Th) as a function of molar volume for n-hexane. hard-sphere fluid to that given by simple Enskog theory.21 This latter quantity is a function of the reduced density, n* = no3. The series of experimental isotherms is therefore fitted to what is really a family of hard-sphere potentials given by a(T).The density dependence of D,,, at reduced densities of the order found in liquids is such that DsHs is essentially linear in the molar volume, V , and DymondZ2 has given2270 6 4 2 h -k L o Q -2 -4 -6 SELF-DIFFUSION OF LIQUID n-HEXANE 100 110 120 130 1 40 150 I.',/cm3 mol-' FIG. 2.-Deviation plot for self-diffusion data fitted to eqn (4) and (5). 0, 223; @, 248; a, 273; ., 298; A, 333 K. TABLE 3.-cOEFFICIENTS FOR EQN (4) AND (5), EQN (9) AND (lo) AND EQN (9) AND (13) lo9 D:eqn (4) & (5) D*:eqn (9) & (10) (q*)-l:eqn (9) & (13) (q*)-':eqn (9) & (13) E, -0.287 010 x 10' ll -0.269 799 x 10' -0.390 101 -0.355 281 E , -0.477 509 x 10, c3 -0.469 259 x 10, -0.482 531 x 10, -0.478 343 x 10, c2 0.150 482 x 10-1 [, 0.143 694 x 10-l 0.202 714 x lo-, 0.186 267 x 8, -0.495 595 x 10-l 5, -0.447 498 x -0.511 774 x lop3 -0.546 286 x lop3 8, 0.224 741 x C2 0.200 519 x lop5 0.273 280 x 0.142 285 x standard deviation (%) 1.2 1.6 1.1 2.8 T range/K 223-333 223-333 298-37332 27 3-3 3 3" a simple equation expressing this.The diffusion isotherms of simple liquids, such as ethylene,23 often show the same behaviour. However, the non-linear dependence of D for n-hexane (see fig. 1) contrasts strongly with that predicted by the model. That a similar effect is observed for substances such as benzene, tetramethylsilaneZ4 and c h l o r ~ f o r m ~ ~ suggests that it is quite general at high densities. Because of this non-linearity, it is not possible to obtain a good estimate of the coupling constant, A , in this particular case, though this can be done in favourable circumstances if the isotherms are sufficiently linear at low density.24 It is convenient to represent the diffusion results in the form of the reduced diffusion coefficient introduced by Dymond26 where the superscript co denotes the dilute gas value of the product (nD) for the hard-sphere fluid, i.e.3/8 ( R T / ~ M ) ~ o - ~ , and y0 is the volume of random close packing, La3/2/2. After substitution of these values in eqn (7), D* can be expressed in terms of experimental quantities as D* = aD DV-t(M/RT)i (8)K. R. HARRIS 227 1 where an = 5.029 x los molf. If the hard-sphere model applied exactly, the product AD* would have the same functional dependence on n* as the molecular-dynamics data, wirh the appropriate choice of 0.Though this is not the case for n-hexane, the D* isotherms do seem to form a family of curves. Over the range of state points covered by these experiments, D* can be fitted as a function of molar volume by condensing the data onto a single 28 through the transformation Q cg -2 -1 - 6 - where T, is again a chosen reference isotherm (298.15 K). It was found that D* could then be satisfactorily fitted to an equation similar to eqn (4): O O e 5 0 O A - A 0.. 0 A a 0 A - I I I I I The coefficients Ci and ti are given in table 3 and the residuals are plotted in fig. 3. 4 The fact that the hexane results conform to eqn (lo), and that the D* isotherms depend only on the reduced density, i.e. on the ratio of the molar volume to a hard-core volume, albeit temperature dependent, assigned to the molecules themselves, could be interpreted as support for a free-volume model of diffusion.This is what one expects and finds for substances which can be modelled as hard spheres, rough or smooth. However, the quite empirical eqn (4) fits equally well and data of quite high precision would be needed to decide between eqn (4) and (10). The form of eqn (9) for the close-packed volume is quite arbitrary. However, it is possible to make an estimate of the dependence of 0, and hence V,, on temperature. In one of the earliest comparisons of the diffusion coefficients of the hard-sphere fluid with that of real liquids, Protopapas et al.29 attempted to relate the hard-sphere diameter 0 to the intermolecular potential function.With the assumption that the bottom of the well was parabolic, they derived a simple formula for 02272 SELF-DIFFUSION OF LIQUID n-HEXANE where a, is the separation at the well minimum and T, the melting temperature. B is a constant which can be related to the ratio of the vibrational amplitude of a molecule to the intermolecular separation at the melting temperature. For the large number of metals examined by Andersen and his coworkers, B was equal to 0.1 12. In earlier work,2s using comparisons between hard-sphere and experimental values for the quantity D*, estimates of a have been obtained for CH, and C2H,, for both of which A is ca. 1. These are shown plotted against Ti in fig. 4. A good correlation is obtained with eqn (1 1) in both cases, with B = 0.069.The line for CH, is slightly concave: this reflects the wide range of temperature for which data are available (up to 1.7 TJ, i.e. well beyond the range for which the approximation of a parabolic well shape might be expected to hold. Values for n-hexane are also shown (table 4). These were calculated from eqn (10) assuming 0 at 298.15 K to be 0.566 nm, a value obtained by making an approximate fit of the low-density data to the hard-sphere equations.'? 26 0.55 I X 15 20 10 15 FIG. 4.-Temperature dependence of the (equivalent) hard-sphere diameters of CH, (go = 0.430, nm), (0); C,H, (go = 0.447, nm), (A); n-C,H,, (oo = 0.624, nm) (El, D ; x , ~ 7 ; ~ ~ W, v ' O ) .K. R. HARRIS 2273 TABLE 4.-EQUIVALENT HARD-SPHERE DIAMETERS FOR n-HEXANE [a (298.15 K) = 0.566 nm] 223.1 5 0.5747 (298.15) 248.15 0.5713 323.15 273.15 0.5683 348.38 (298.15) (0.5660) 373.36 333.15 0.5635 273.15 (298.15) 303.15 333.15 (0.5660) 0.5639" 0.5625a 0.56 17" 0.56Bb (0.5660) 0.5655b 0.5628b Calculated from falling-slug-viscometer re~ults.~2 Calculated from rolling-ball-visco- meter results.lo Again a straight line is obtained, with the parameter B taking the value of 0.072.It appears therefore that eqn (1 1) is suitable for correlating hard-sphere diameters for liquids conforming to the model and equivalent hard-sphere diameters for those for which D* is a function of reduced density alone, and that this can be done over a fairly wide range of temperature. COMPARISON WITH THE VISCOSITY It is of interest to compare the density dependence of the diffusion coefficient with that of the viscosity.At a given temperature the product (Dq) for the intra-diffusion coefficient of a particular solute in a range of solvent of differing viscosity almost always increases as the viscosity, q, increases3O9 31 and the following relationship, recurrent in the literature, has been proposed Dqt =f(q t < 1. (12) This equation, which is more general than the Stokes-Einstein equation, may also be applied to self-diffusion. Measurements of the viscosity of n-hexane have been reported for temperatures between 298 and 373 K by Dymond et aZ.32 and between 273 and 333 K by Brazier and Freernan.'O The former were obtained with a falling-slug viscometer, the latter with a rolling-ball viscometer. The two sets of data are, however, not concordant, differing by as much as 10% at 333 K and 400 MPa, whereas the experimental error is of the order of 2%.Both sets are in reasonable agreement with the capillary- viscometer results of Kuss and P01lmann~~ at 313 K and 0.1 to 150 MPa, and as it is not clear which set is to be preferred both are used in this analysis. The viscosities have been fitted to an equation analogous to eqn (10) where V,' is again given by eqn (9). The reduced viscosity q* is given byz6 where q7"0 is the dilute gas value for the hard-sphere fluid and eqn (14) may be reduced, in a similar way to that for D*, to q* = a,, ~ v % ( R M T ) - ~ (1 5 )2274 SELF-DIFFUSION OF LIQUID n-HEXANE with a, = 6.0348 mol-4. The coefficients of eqn (13) are given for the two different sets of viscosities in table 3.The fit of the falling-slug results is better than that of the rolling-ball data. Approximate diameters calculated as described above are also plotted in fig. 4. Those obtained from the data of Brazier and Freemanlo are in very good agreement with the diameters calculated from the diffusion coefficients (cf. CH,28), but those calculated from the viscosities of Dymond et af.32 are not, giving the much lower value of 0.046 for B. The closeness of the value of B obtained from the combined self-diffusion and former set of viscosity data for n-hexane to that found for methane and ethylene suggests that any change in conformation with pressure has little effect on the apparent hard-sphere diameter. Both sets of viscosities were also fitted to eqn (1 2), the falling-slug and rolling-ball results yielding values of 0.93 0.01 and 0.97 f 0.01, respectively, for the exponent t.Again the former gave the better fit. With either set, it is apparent that the Stokes-Einstein relationship is not obeyed. K. R. Harris, Physica, 1978, 94A, 448. L. A. Woolf, J. Chem. SOC., Faraday Trans. I , 1982, 78, 583. J. Jonas, D. Hasha and S. G. Huang, J . Phys. Chem., 1980, 84, 109. D. L. Hogenboom, K. Krynicki and D. W. Sawyer, Mol. Phys., 1980, 40, 823. K. R. Harris and L. A. Woolf, J. Chem. SOC., Faraday Trans. I , 1980, 76, 377. P. E. Schoen, R. G. Priest, J. P. Sheridan and J. M. Schnur, Nature (London), 1977, 270, 412; J. Chem. Phys., 1979, 71, 317. H. E. Eduljee, D. M. Newitt and K. E. Wade, J . Chem. SOC., 1951, 3086. F.1. Mopsik, J. Res. Natl Bur. Stand., Sect. A , 1967, 71, 287. lo D. W. Brazier and G. R. Freeman, Can. J . Chem., 1969, 47, 893. l 1 J. H. Dymond, K. J. Young and J. D. Isdale, J. Chem. Thermodyn., 1979, 11, 887. l 2 A. Kumagai and H. Iwasaki, J. Chem. Eng. Data, 1979, 24, 261. l 3 K. R. Harris and P. J. Dunlop, J. Chem. Thermodyn., 1970, 2, 813. I4 D. C. Douglass and D. W. McCall, J . Phys. Chem., 1958, 62, 1102. l5 D. C. Douglass, D. W. McCall and E. W. Anderson, Phys. Fluids, 1959, 2, 87. l6 K. R. Harris, C. K. N. Pua and P. J. Dunlop, J. Phys. Chem., 1970, 74, 3518. ’ K. R. Harris, R. Mills, P. J. Back and D. S. Webster, J . Magn. Reson., 1978, 29, 473. K. Aoyagi and J. G. Albright, J. Phys. Chem., 1972, 76, 2572. J. H. Dymond and L. A. Woolf, J. Chem. SOC., Faraday Trans. I , 1982, 78, 991. l9 B. J. Alder, W. E. Alley and J. H. Dymond, J. Chem. Phys., 1974, 61, 1415. 2o D. Chandler, J. Chem. Phys., 1975, 62, 1358. 21 B. J. Alder, D. M. Gass and T. E. Wainwright, J . Chem. Phys., 1970, 53, 3813, 22 J. H. Dymond, J . Chem. Phys., 1974, 60, 969. 23 B. Arends, K. 0. Prins and N. J. Trappeniers, Physica, 1981, 107A, 307. 24 H. J. Parkhurst Jr and J. Jonas, J . Chem. Phys., 1975, 63, 2698. 25 M-K. Ahn, K. R. Harris and L. A. Woolf, unpublished results. 26 J. H. Dymond, Physica, 1974, 75, 100. 2’ J. H. Dymond and T. A. Brawn, Proc. 7th Symposium Thermophys. Props (Am. SOC. Mech. Eng., New 28 K. R. Harris and N. J. Trappeniers, Physica, 1980, 104A, 262. 29 P. Protopapas, H. C. Andersen and N. A. D. Parlee, J . Chem. Phys., 1973, 59, 15. 30 D. F. Evans, T. Tominaga and C. Chan, J . Solution Chem., 1979, 8,461 ; D. F. Evans, T. Tominaga 31 H. J. V. Tyrrell, Sci. Prog., 1981, 67, 271. 32 J. H. Dymond, K. J. Young and J. D. Isdale, Znt. J . Thermophys., 1981, in press; J. D. Isdale, 33 E. Kuss and P. Pollmann, 2. Phys. Chem. N . F., 1969, 68, 205. York, 1977), p. 660. and H. T. Davis, J. Chem. Phys., 1981, 74, 1298. J. H. Dymond and T. A. Brawn, High Temp. High Press., 1979, 11, 571. (PAPER 1 / 163 1)

ISSN:0300-9599    

DOI:10.1039/F19827802265

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chern. SOC., Faraday Trans. I , 1982, 78, 2275-2278 Osmotic Coefficients of Urea + Guanidinium Chloride Mixtures in Water at 298.15 K BY TERENCE H. LILLEY* AND DENNIS R. TESTER Chemistry Department, The University, Sheffield S3 7HF Received 2nd November, 198 1 The osmotic coefficients of urea + guanidinium chloride mixtures in aqueous solutions have been determined at 298.15 K using the isopiestic vapour-pressure technique. A comparison is made between the results obtained and the corresponding results for other urea+ salt aqueous solution mixtures. There is a correspondence between the urea + salt interactions and the tendency for the salts to destabilise protein structures. It has been known for some time that aqueous solutions containing either urea or guanidinium chloride (GuCl) can induce the denaturation of pr0teins.l More recent work2 has shown that solutions containing either or both of these solutes have a marked effect on the rates of denaturation and renaturation.The present work was stimulated by the observation3 that mixed solutions of urea + GuCl are considerably less efficacious in their effect on the rate of denaturation of lysozyme than solutions containing only one of these solutes. The thermodynamic properties of aqueous urea solutions have been investigated in some considerable detail and there are possibly more precise thermodynamic data available for the urea+water system than for any other system containing a non-electrolyte in water.4 It is only relatively recently that attention has been turned towards the properties of aqueous GuCl ~olutions.~-~ In the present paper we present an investigation of the osmotic coefficient behaviour of binary-solute aqueous mixtures containing urea and GuCl.The results obtained indicate rather pronounced interactions between the two solutes. EXPERIMENTAL AND RESULTS The apparatus and experimental procedure have been described previously.8 The urea was of ultrapure quality and the GuCl was of analytical reagent grade and was twice recrystallised from water. Both substances were intensively dried before use. The experimental results obtained, presented as the molalities of the solutes in isopiestic equilibrium, are given in table 1. In view of the precision to which the osmotic coefficients of urea solutions are known, solutions containing this solute were used as the reference solute in the isopiestic experiments. DISCUSSION It is convenient8 when treating systems such as that considered here to define a term (1) Nm4) = mref 4ref-2mg 4E-mn 4; where, in the present instance, mref is the molality of the urea reference solution in a given experimental run and 4ref is the osmotic coefficient of this reference solution.22752276 OSMOTIC COEFFICIENTS OF UREA+GUC~ TABLE 1 .-EXPERIMENTAL RESULTS : MOLALITIES OF ISOPIESTIC SOLUTIONS rn/mol kg-I m/mol kg-l urea GuCl - A(m4) urea GuCl - A(m4) 13.0462 11.9080 10.7220 8.1091 6.7483 5.3910 2.7056 1.3244 0 8.6956 7.891 5 7.1074 5.3719 4.4763 3.5663 1.8121 0.8913 0 0 0.8420 1.6599 3.3576 4.1947 4.9945 6.5039 7.2487 7.9260 0 0.5580 1.1003 2.2242 2.7824 3.3225 4.3562 4.8749 5.3500 - 0.5389 0.8 172 1.0523 1.0550 0.9809 0.6505 0.3816 - - 0.3314 0.5408 0.7121 0.7146 0.67 13 0.4448 0.248 1 - 6.957 1 6.3602 5.6785 4.3099 3.6022 2.8929 1.4489 0.7528 0 3.3874 3.4353 3.0534 2.2983 1.9178 1.5385 0.7719 0.4019 0 ~~ ~~ 0 0.4122 0.871 1 1.7550 2.1947 2.627 1 3.48 11 3.8735 4.2934 0 0.2227 0.4684 0.9362 1.1684 1.3972 1.8547 2.0679 2.2975 - 0.2477 0.421 1 0.5672 0.5686 0.5338 0.3552 0.2040 - 0.1078 0.1907 0.2583 0.2606 0.2438 0.1597 0.0889 TABLE 2.-cOEFFICIENTS AND THEIR 95 % CONFIDENCE LIMITSa FOR THE REPRFSENTATION OF A(m4) BY EQN (2) A,l/mol-l kg A21/mo1-2 kg2 A,2/mo1-2 kg2 A,,/m01-~ kg3 A,,/molP kg3 A,,/m01-~ kg3 A 14/m01-4 kg4 ~ , , / m ~ l - ~ kg4 - 0.1907 (0.0 127) 2.4341 (0.5623) x lop2 3.7332 (0.3737) x -2.3293 (0.3145) x -2.821 1 (0.7091) x - 1.9235 (0.9351) x lop3 6.2514 (4.3590) x 1.1544 (0.7815) x lop4 a The 95% confidence limits are shown in parentheses.The molalities of GuCl and urea in the binary solute mixtures are denoted by mg and m,, respectively, and 4: and 4; are the osmotic coefficients of GuCl and urea in single-solute-containing solutions at the molalities mg and m,. The osmotic coefficients of urea solutions were obtained from stoke^'^ critical compilation. The osmotic coefficients of single-solute solutions containing GuCl were found to agree very well with those obtained by Schrier and S ~ h r i e r , ~ and consequently their parametric expression was used to obtain values for the osmotic coefficients of GuCl solutions. The agreement between the present results for GuCl single-solute solutions and those calculable from the results obtained in two recent investigations by Bonner' and by Bates and coworkers6 are acceptable by most standards. We have included theT.H. LILLEY A N D D. R. TESTER 2277 calculated values of A(m4) for the various mixtures in table 1. As is now customaryg the values of A(m4) were fitted to an equation of the form Table 2 gives the values of the coefficients ( A t j ) , along with their 95% confidence limits, which were required to fit the results adequately. In a notional sense, Aij represents a measure of the interaction of i species of GuCl with j species of urea; e.g. A,, corresponds to interactions between two GwCl species with two urea species. There are several manipulations which are possible given these coefficients. For example, by appropriate integration and differentiation of eqn (1) and (2) we obtain the following expression for the logarithm of the ratio of the activity coefficient of GuCl in urea+GuCl mixture ( y g ) to that in a solution containing only GuCl ($) at the same molality as the GuCl in the mixture: The corresponding expression for the logarithm of the activity coefficient ratio for the urea is In (y,/y:) = x D/(i+j- 1) A i j mk mip1.j-1 i-1 (4) Eqn (3) and (4) in association with the coefficients given in table 2 and the activity coefficients of the solutes in solutions containing only one solute component4. may be used to calculate the activity coefficients of urea and GuCl in mixtures. As was mentioned in the Introduction, urea + GuCl solutions denature proteins rather effectively.However, it has also been suggested1" that urea mimics, to some degree, an average peptide group in proteins. Given this assertion it is convenient to calculate the free energy of transfer of urea from water to GuCl solutions under the condition that the urea molality approaches zero. This free energy of transfer (A&?) f I I I I 1 0 2 4 6 8 salt molality/mol kg-' FIG. 1.-Variation of the free energy of transfer of urea (A&) with salt molality.2278 OSMOTIC COEFFICIENTS OF UREA+GuCI will then give an indication of how a peptide group interacts with GuCl in aqueous systems. The expression for the free energy of transfer under the above condition is A&? = RT C (Ail/i>mt.i-1 In fig. 1 we illustrate the dependence of the free energy of transfer of urea on the molality of added salt for GuCl. Included in this figure are the corresponding results for urea in solutions of NaC1,ll LiCPO and CaC1,.lo The results for all four salts indicate an attractive interaction with urea (i.e. the free energies of transfer are negative) and the ranking order in terms of such interactions is NaCl < LiCl < GuCl < CaCl,. This is the same as the order observed' for the destabilising effect of salt solutions on protein structures. In other words urea + salt interactions in aqueous solution appear to follow the Hofmeisterl series. The implication of this is that, given the common features of structure between urea and the peptide bond, the denaturation of proteins by salts is related to how salts interact with peptide bonds. We acknowledge help in various ways from K. G. Davis. See P. H. Von Hippel and T. Scheich, Acc. Chem. Res., 1969,2,257 and F. Franks and D. Eagland, Crit. Rev. Biochem., 1975, 3, 165, for references to early work. C. Tanford, Ado. Protein Chem., 1968, 23, 121. B. Robson and T. H. Lilley, to be submitted for publication. The data have been compiled by R. H. Stokes, Aust. J . Chem., 1967, 20, 2087. M. Y. Schrier and E. E. Schrier, J . Chem. Eng. Data, 1977, 22, 73. J. B. Macaskill, R. A. Robinson and R. G. Bates, J. Chem. Eng. Data, 1977, 22, 411. T. H. Lilley and R. P. Scott, J. Chem. SOC., Faraday Trans. I , 1976, 72, 184. T. H. Lilley and R. P. Scott, J. Chem. SOC., Furaday Trans. 1, 1976, 72, 197. V. E. Bower and R. A. Robinson, J. Phys. Chem., 1963, 67, 1524. ' 0. D. Bonner, J. Chem. Thermodyn., 1976, 8, 1167. lo M. Y. Schrier, A. H. C. Ying, M. E. Ross and E. E. Schrier, J. Phys. Chem., 1977, 81, 674. (PAPER 1 / 1707)

ISSN:0300-9599    

DOI:10.1039/F19827802275

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. 1, 1982,18, 2219-2283 Stabilization of Cuprous Oxide Photocathode in Aqueous Thiocyanate Solution by Aliphatic Alcohols K. TENNAKONE,* W. G. D. DHARMARATNA AND S. C. JAYEWARDENA Department of Physics, Ruhuna University College, Matara, Sri Lanka Received 16th November, 198 1 Long-chain normal aliphatic alcohols are found to suppress photocorrosion and increase the efficiency of a cuprous oxide photoelectrochemical cell with aqueous thiocyanate solution as the electrolyte. It is suggested that alcohols inhibit corrosion by surface adsorption. Photoelectrochemical cells with semiconductor electrodes seem to be one of the most promising methods for solar-energy conversion. However, they are plagued with instability resulting from photocorrosion, and methods of suppressing corrosion have attracted much attenti0n.l l1 In this work we describe our observations of the stabilizing influence of adsorbed normal aliphatic alcohols on a photocathode of cuprous oxide in aqueous thiocyanate solution.The mechanism of stabilization is discussed. EXPERIMENTAL Cu,O plates are made by heating 7 x 11 cm copper sheets in an oven to 850 "C. The outer film of CuO is removed by etching in dilute acetic acid. As the cell performs better in the back-wall mode, the anode used is a copper window (same dimensions as the cathode) coated with cupric sulphide. To obtain a reproducible dark open-circuit voltage, the anode is prepared by the following procedure. The cleaned copper electrode is immersed in a 5% solution of Na2S and then left overnight in a saturated solution of H2S.Two electrodes, separated 1 cm from each other, are supported in a glass vessel (12 x 8 x 3 cm). The electrolyte is an alkaline solution of KCNS (0.5 mol dmP3 KCNS, 0.2 mol dma3 NaOH). RESULTS AND DISCUSSION Cu,O is a p-type semiconductor of band-gap Eg = 2.3 eV. When a semiconductor surface in contact with an electrolyte is illuminated, the electron-hole pairs created by photons (hv > E,) are separated by the surface-barrier electric field. In a p-type material, electrons tunnel across the barrier into the electrolyte and holes carry the positive current in the interior of the semiconductor.l2? l3 Electrons tunnelling into the solution cause chemical reductions near the semiconductor surface. Unless a redox system with good electron-accepting properties is present in the solution, the semiconductor is reduced.The photoreduction of a Cu,O cathode is rapid in most electrolytes. An acidic medium dissolves Cu,O or reduces it instantly to metallic copper. In an alkaline medium, the familiar redox systems either produce a weak photo- response or attack the electrode in the dark. After testing several electrolytes we found 22792280 c U 2 0 PHOTOCATHODE STAB I LI Z ATION 2 2 4 E + l .. 1 .5 .o .5 .o 0 10 20 30 40 50 60 tlmin FIG. 1.-Decay of the short-circuit photocurrent: (a) without alcohol, (b) 10-2 mol dmP3 ethanol, (d) lop3 mol dmP3 propan-1-01, (e) loP3 mol dm-3 pentan-1-01, ( f ) hexan-1-01, (s) mol dm-3 methanol, (c) mol dmP3 mol dmP3 octan-1-01. (Intensity of illumination z 10 W m-2, plate area = 60 an2.) > E L- 0 2 4 6 8 10 12 14 16 tlmin FIG.2.-Time development of the open-circuit voltage: (a) without alcohol, (b) lo-* mol dm-3 methanol, ( c ) lo-* mol dmP3 ethanol, ( d ) mol dm-3 propan-1-01, (e) mol dm-3 pentan-1-01, ( f ) rnol dm-3 octan-1-01. (Intensity of illumination x 10 W m-*, plate area = 60 cm2,) mol dmP3 hexan-1-01, ( g ) that a very high photoresponse can be obtained with a solution of KCNS. Neutral or alkaline solutions of KCNS do not react with Cu,O. However, if the short-circuited cell is exposed to light, the Cu20 plate deteriorates rapidly, forming a grey deposit of CuO and CuCNS. The time dependence of the short-circuit current is indicated in fig. 1 (a). The cell decays to complete extinction in ca.25 min (intensity of illumination x 10 W m-2 from a tungsten-filament lamp). Normal aliphatic alcohols (NAAs) have a remarkable effect on the stability of the cell. When minute quantities of NAAs (< mol dm-3 for methanol or ethanol, < lop3 rnol dm-3 for higher NAAs) are added to the electrolyte, the lifetime of the cell is increased from 25 min to several hours (fig. 1). The effectiveness of NAAs as stabilizing agents increases progressively with the chain length. The slow decay of the photocurrent (fig. 1) is almost entirely due to deterioration of the copper sulphideK. TENNAKONE, W. G. D. DHARMARATNA A N D S. C. JAYEWARDENA 2281 anode. If the anode is replaced, the cell recovers to practically the same efficiency. This process can be repeated for ca.10 h, when it shows a 25% drop in the starting photocurrent. In the absence of NAAs the open-circuit voltage of the cell also decays with time (fig. 2). Long-chain NAAs fully stabilize the saturation voltage. It is seen that the higher the chain length, the shorter is the time taken to reach saturation. Again NAAs are found to increase the efficiency of the cell. Fig. 3 gives ZV against Z 91 I N 3 2 1 I I I 1 I 0 10 20 30 40 50 60 //PA cm-2 FIG. 3.-ZVagainst Zcurves: (a) without alcohol, (6) mol dm-3 ethanol, (d) rnol dm-3 propan-1-01, (e) lop3 mol dmp3 pentan-1-01, (f) lop3 mol dmp3 hexan-1-01, (g) mol dm-3 octan-1-01. Efficiencies at the maximum power point are (a) 0.28, (6) 0.45, (c) 0.48, ( d ) 0.52, (e) 0.56, (f) 0.61 and (s) 0.65. (Intensity of illumination sz 10 W m-3r plate area = 60 cm2.) mol dm-3 methanol, (c) curves.It is apparent that long-chain NAAs are more effective. An explanation for the above facts will not be given. KCNS solution acts as a redox couple owing to the existence of CNS- and (CNS); ions. Possible reactions occurring at the electrodes are as follows. At the photocathode (CNS); ions accept electrons to yield CNS- ions, i.e. (CNS); + e -+ 2CNS-. Near the anode CNS- ions discharge electrons producing CNS free radicals, which combine with CNS- ions in the solution to regenerate (CNS); ions, i.e. CNS--e -+ CNS CNS + CNS- -+ (CNS);. Photocorrosion is probably due to the presence of anodic regions in the Cu,O surface. Uneven illumination, difference in surface light absorption coefficients and other non-uniformities create regions of different potential in the photocathode; 74 FAR 12282 c U 2 0 PHOTOCATHODE ST A B I LI ZATION short-circuiting across these regions deteriorates the anodic regions by the following mechanism.In the anodic region, CNS- ions discharge electrons yielding CNS free radicals, which sometimes, instead of combining with (CNS)- to regenerate (CNS);, react with Cu,O to form cuprous thiocyanate and cupric oxide, i.e. C U , ~ + CNS + CuO + CuCNS. The fact that the deposit formed on the photocathode is a mixture of CuCNS and CuO supports the above hypothesis. Two other observations favour this argument. When a metallic copper electrode is used as the anode, it quickly corrodes with deposition of CuCNS. The other observation is that if a circular patch of light is focused into the cathode, the dark region just outside the periphery of the illuminated patch corrodes immediately leaving the bright region untarnished.The illuminated region is cathodic: the anodic region in the immediate vicinity is the dark area around the periphery. Note that photocorrosion in an anodic region destroys the photosensitivity, making it even more anodic; this enhances corrosion which then spreads throughout the surface. Our observations strongly suggest that NAAs prevent photocorrosion by adsorption. Traces of NAAs arrest the deposition of CuCNS and CuO on the photocathode. A concentration of pentan-1 -01, hexan- 1-01 or octan- 1-01 of < mol dmP3 is sufficient for this purpose. There is no evidence that the adsorbed NAAs are chemically changed.If the purity of the electrolyte is maintained the NAA layer (probably a monolayer) remains active almost indefinitely. Addition of soap or other detergent at once restores photocorrosion. Obviously this is caused by removal of the adsorbed layer by the detergent. The adsorbed NAA probably inhibits photocorrosion because it prevents the direct contact of CNS free radicals with the anodic regions. Electrons, however, could easily tunnel across the NAA layer. The long-chain NAAs are better adsorbed14 and produce more effective barriers against ions and free radicals. Although NAAs completely suppress photocorrosion leading to the formation of CuO and CuCNS, at higher intensities of illumination a certain degree of photoreduction of Cu,O to metallic copper takes place.The adsorbed NAA possibly screens only larger ions and free radicals. Hydrogen ions may perhaps penetrate the barrier and reduce Cu,O to metallic copper. Again, as expected the reduction of Cu,O to metallic copper occurs more slowly with long-chain NAAs. We mentioned earlier that the open-circuit voltage of the cell decays in the absence of NAAs. This observation can be explained as being due to short-circuiting across anodic and cathodic regions, which quickly causes the anodic regions to corrode and expand. In the presence of NAAs the enhancement of anodic regions by corrosion is stopped, and the cell attains saturation voltage. The same effect explains why the efficiency is increased by NAAs. When the short-circuiting in the cathode is arrested, the photocurrent is driven almost entirely through the external circuit.The slow decay of the photocurrent results from a deterioration of the copper sulphide anode. This is an oxidation process that deposits Cu,O. Since Cu,O is a very bad electron acceptor, once this oxide layer is formed CNS- ions fail to deliver electrons efficiently to the anode. We have not been able to find a better material for the anode. Familiar stable electrodes have electrode potentials which are either too low or too high compared with that of the photocathode. Another observation we cannot explain is that only NAAs inhibit photocorrosion in Cu,O. Branched aliphatic alcohols have the opposite effect; traces of theseK. TENNAKONE, w. G. D. DHARMARATNA AND s.c. JAYEWARDENA 2283 compounds accelerate corrosion. Aromatic alcohols are ineffective, but they do not accelerate corrosion. The above cell is not suitable as a practical device for the conversion of solar energy. In the present form the efficiency does not exceed 1 %. Also it may not be completely regenerative, as the redox action of KCNS involves several steps. Nevertheless, our investigations suggest that agents which are adsorbed at the photoelectrodes might inhibit corrosion in more efficient practical photoelectrochemical cells. A. Heller, K. C. Change and B. Miller, J. Electrochem. Soc., 1977, 124, 697. A. J. Bard and M. S. Wrighton, J. Electrochem. SOC., 1977, 124, 1706. H. Gerischer and J. Gobrecht, Ber. Bunsenges. Phys. Chem., 1978, 82, 520. H. Hodes, Nature (London), 1980, 285, 29. Y. Nakato, S. Tonomura and H. Tsubomura, Ber. Bunsenges. Phys. Chem., 1976, 80, 1289. Y . Nakato, T. Ohnishi and H. Tsubomura, Chem. Lett., 1975, 883. J. G. Mavroides, J. C . Fan and H. G. Zeiger, in Photoefects at Semiconductor-Electrolyte Interfaces, ed. A. J. Nozik, A.C.S. Symp. Ser. No. 146 (American Chemical Society, Washington, D.C., 1981), p. 217. M. S. Wrighton, R. G. Austin, A. B. Bocarsly, J. M. Bolts, 0. Hass, K. D. Legg, L. Nadjo and M. C. Palazzto, J. Am. Chem. Soc., 1978, 100, 1063. F. Decker and T. Freund, J. Chem. Phys., 1967, 47, 1543. lo V. A. Tyagai, Electrokhimya, 1965, 1, 381. l 1 D. S. Ginley and M. A. Butler, J. Electrochem. Soc., 1978, 125, 1968. l2 R. Williams, J. Chem. Phys., 1960, 32, 1505. l 3 S. R. Morrison and T. Freund, J. Chem. Phys., 1967, 47, 1543. l 4 A. W. Adamson, Physical Chemistry of Surfaces (Interscience, London, 2nd edn, 1967). (PAPER 1 / 1774) 74-2

ISSN:0300-9599    

DOI:10.1039/F19827802279

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. I , 1982, 78, 2285-2296 Infrared Study of the Adsorption of Fluoroalcohols on Silica BY ANTONIO R. ACOSTA SARACUAL, SUNITA K. PULTON AND GEORGE VICARY Department of Chemistry, The University, Nottingham NG7 2RD AND COLIN H. ROCHESTER* Department of Chemistry, The University, Dundee DD1 4HN Received 17th November, 198 1 Infrared spectra have been recorded of ethanol, 2,2,2-trifluoroethanol and 1 1 1,3,3,3-hexafluoropropan- 2-01 adsorbed on silica at the solid/vapour interface and on silica immersed in carbon tetrachloride. The predominant surface-adsorbate interactions at ca. 306 K involved the formation of hydrogen bonds between surface silanol groups and hydroxy groups in adsorbed molecules. The strongest adsorption of this type occurred for single alcohol molecules adsorbed on to pairs of adjacent silanol groups. Isolated silanol groups acted as donors of hydrogen bonds, each to a single alcohol acceptor molecule.For ethanol, at high surface coverages lateral interactions between adjacent adsorbed ethanol molecules occurred and led to the build-up of aggregates of ethanol on the silica surface. Some chemisorption occurred for all three alcohols primarily through addition reactions at exposed strained siloxane groups. Infrared studies of the adsorption of methanol on silica at room temperature have established that the predominant interaction involves the formation of hydrogen bonds between surface silanol groups and hydroxy groups in adsorbed molecule^.^-^ At high surface coverages lateral interactions between adjacent methanol molecules also became important.2 A small degree of chemisorption results in the formation of surface methoxy groups possibly via esterification of silanol groups4! but more probably by the addition of methanol molecules to strained siloxane groups in the silica surface.'? 2* The formation of hydrogen bonds between silanol groups and adsorbed molecules gives shifts AFoH in the position of the infrared band assigned to the OH-stretching vibrations of silanol groups.For the adsorption of pyridines electron-donating substituents enhanced the basicity at the hydrogen-bond acceptor site (N atom). This led to a strengthening of the hydrogen-bonding interactions which was reflected in bigger AVoH value^.^ Hydroxy groups in organic molecules can act as either acceptors or donors of hydrogen bonds.Complexes (I) and (11) therefore represent two possible modes of adsorption of hydroxy compounds (ROH) on silica. Increasing the basicity of the hydroxy (ROH) group in complex (I) would be expected to strengthen the hydrogen- bonding interaction and, by analogy to the results for pyridines,' should lead to a bigger shift for the silanol group. Conversely, for complex (11) AVOH would be reduced because an increase in basicity would be accompanied by a decrease in the hydrogen-bond donor strength of the organic hydroxy group and therefore a weakening of the hydrogen-bonding interaction. The spectroscopic shifts which accompany the adsorption of ethanol (pK, 15.9 in water at 298 K),8 2,2,2-trifluoroethanol (pK, 1 2.4)97 lo and 1, 1,1,3,3,3-hexafluoro- 22852286 FLUOROALCOHOLS-kSILICA R I 9, H- H 0’ I / I \ Si R I H.H O\ ..O’ propan-2-01 (pK, 9. 3)11 at the silica/vapour and silica/carbon tetrachloride interfaces are reported here. A related study of the adsorption of substituted phenols at the silica/n-heptane interface12 gave results which were complicated by steric factors and modes of adsorption involving interactions between silanol groups and aromatic ring systems. l3 EXPERIMENTAL Ethanol was purified by distillation. The fluoroalcohols were commercial samples which gas-liquid chromatographic analysis showed to be > 99% pure. Carbon tetrachloride was dried over calcium sulphate, decanted and distilled from calcium sulphate. Alcohol vapours were adsorbed on to Degussa Aerosil 200 silica discs (surface area 176 m2 g-l) mounted in a Pyrex infrared cell with fluorite optical windows14 and glass-blown to a vacuum apparatus capable of maintaining a dynamic vacuum of N m-2.Discs were heated at 673 K in oxygen (27 kN m-2, 21 h) and vacuum (2 h) and cooled to ambient temperature (ca. 306 K) in the spectrometer (Perkin-Elmer 157G) beam before the admission of alcohol vapour. Adsorption at the solid/liquid interface was studied by adding carbon tetrachloride or solutions of an alcohol in carbon tetrachloride to a silica tube containing silica which had been preheated at 1073 K (1 h) in a vacuum. Shaking produced homogeneous dispersions which were transferred to a conventional infrared liquid cell for spectroscopic examination.Spectra were also recorded of solutions in carbon tetrachloride containing a fixed concentration of alcohol and a series of concentrations of diethyl ether. The aim was to measure the effect of alcohol pK, on and the equilibrium constant for the formation of hydrogen bonds between alcohol donor molecules and a given acceptor molecule. Solutions were at 298 K while spectra were being recorded. Diethyl ether was dried with sodium and distilled. RESULTS The adsorption of 2,2,2-trifluoroethanol on silica resulted in decreases in the intensity of the maximum at 3745 cm-l due to isolated surface silanol groups and the concomitant appearance of a broad maximum due to hydroxy groups perturbed by hydrogen-bonding interactions (fig. 1). The broad band shifted from 3500 at low surface coverages through 3430 cm-l at half coverage to 3400 cm-1 at high coverages.A maximum at 3660 cm-l could be entirely ascribed to the OH-stretching vibrations of 2,2,2-trifluoroethanol molecules in the vapour phase [fig. 1 (g)]. Bands associated with CH-stretching or CH-bending vibrations were at spectral positions which closely corresponded to maxima at 3000, 2953, 2900, 1460, 1416 and 1365 cm-l [fig. 1 (g)] for 2,2,2-trifluoroethanol vapour, and contained contributions from the spectra of both vapour and 2,2,2-trifluoroethanol molecules adsorbed via hydrogen-bonding interactions with the surface. Desorption of 2,2,2-trifluoroethanol by evacuation at ambient temperature revealed that some chemisorption had occurred. Residual maxima at 2975, 2910, 1467, 1427ACOSTA SARACUAL, PULTON, VICARY AND ROCHESTER I I I 38 oi0 4 g, I 1 3000 3600 2287 wavenumber/cm-' FIG. 1 .-Spectra of silica (a) in vacuum, (b)-cf) in contact with increasing pressures of 2,2,2-trifluoroethanol vapour.(s) Spectrum of 2,2,2-trifluoroethanol vapour (staggered for clarity in the 2800-3700 cm-' regon) which in contact with silica gave spectrum 0. I I I 1 I I 3600 3000 r----- 1500 1300 wavenumber/cm-' FIG. 2.-Spectra of silica (a) in vacuum, (b) after exposure to 2,2,2-trifluoroethanol vapour [fig. 1 v)] and evacuation (ca. 306 K, 70 min). ( c ) Spectrum of liquid 2,2,2-trifluoroethanol.2288 FLUOROALCOHOLS+SILICA and 1380 cm-l [fig. 2(b)] were due to CH-containing chemisorption products. The intensity of the maximum at 3745 cm-l [fig.2(b)] did not return to its initial value [fig. 2 (a)] before the admission of 2,2,2-trifluoroethanol vapour. However, small increases in adsorption intensity had occurred in the spectral range 3300-3700 cm-l. Subsequent heat treatment at 673 K (24 h) resulted in the disappearance of bands, apart from weak shoulders at 1405 and 1370 cm-l, due to chemisorbed hydrocarbon species and the loss of enhanced intensity in the 3300-3700 cm-l range. The intensity of the band at 3745 cm-l was not affected by the thermal treatment, but a new weak maximum appeared at 3705 cm-l. The spectrum of liquid 2,2,2-trifluoroethanol was recorded for comparison with the spectrum of 2,2,2-trifluoroethanol in the adsorbed state. Maxima at 1458, 1428 and 1376 cm-l exhibited different relative intensities [fig.2 (c)] than the corresponding maxima in the vapour spectrum [fig. 1 (g)]. In particular the enhanced intensity of the band at 1458 cm-l was compatible with the maximum at 1467 cm-l in spectra of chemisorbed species [fig. 2 (b)]. Bands at 2970 and 2895 cm-l exhibited similar relative intensities to the values for the bands at 2975 and 2910 cm-l [fig. 2(b)] for adsorbed species. Spectra of silica in the presence of 2,2,2-trifluoroethanoI vapour contained no evidence for a band corresponding to that at 3642 cm-l [fig. 2(c)] due to non-hydrogen-bonded hydroxy groups in the liquid alcohol. The broad maximum at 3365 cm-l due to hydrogen-bonded hydroxy groups in the liquid phase occurred at a lower wavenumber than the broad maximum at 3400 cm-l [fig.1 cf)] in spectra of silica exposed to high vapour pressures of 2,2,2-trifluoroethanoI. In spectra of 1,l , 1,3,3,3-hexafluoropropan-2-o1 on silica the maximum at 3745 cm-l was progressively replaced by a broad band shifting from 3490 cm-l at half surface coverage to 3450 cm-l at high coverage (fig. 3). Bands at 3666 and 3626 cm-l were I I I ' 3 600 3000 1500 1300 wavenumber/ cm-' FIG. 3.-Spectra of silica (a) in vacuum, (b)-(f) in contact with increasing pressures of 1,1,1,3,3,3- hexafluoropropan-2-01 vapour. (g) and (h) Spectra of I,], 1,3,3,3-hexafluoropropan-2-01 vapour which in contact with silica gave spectra (d) and v), respectively.ACOS T A SARA C U AL, PU LTON, V I C ARY AND ROCHESTER 2289 due to OH-stretching vibrations of alcohol molecules in the vapour phase [fig.3(g)]. Maxima at 2990 and 2950 cm-l, due to CH-stretching vibrations, and at 1372 cm-l, due to CH-bending vibrations, contained contributions from both adsorbed species and vapour phase. However, a shoulder at 1422 cm-l [fig. 3 (b)-(f>] was more intense than expected by comparison with the intensity of a weak band at 1427 cm-l in the vapour-phase spectrum. The shoulder corresponds to a band at 1427 cm-l in spectra of the liquid alcohol [fig. 4(c)] which has been assigned to an OH-deformation vibration of associated 1, 1 , 1,3,3,3-hexafluoropropan-2-o1 molecule^.^^ 3600 3000 1500 1300 wavenumber/cm-' FIG. 4.-Spectra of silica (a) in vacuum, (b) after exposure to 1, 1,1,3,3,3-hexafluoropropan-2-o1 vapour [fig. 2 cf)] and evacuation (ca.306 K, 1 h). (c) Spectrum of liquid l , l , 1,3,3,3-hexafluoropropan-2-o1. Desorption of 1, 1 ,1,3,3,3-hexafluoropropan-2-o1 by evacuation removed molecules held to the surface by hydrogen-bonding interations, but residual chemisorbed species were evidenced by bands at 2957 and 1378 cm-l [fig. 4(b)] due to CH-stretching and CH-deformation vibrations, respectively. The maximum at 3745 cm-l nearly returned to its original intensity before the admission of alcohol vapour [fig. 4(a)]. However, the chemisorptive reactions led to a general increase in adsorption intensity in the range 3300-3700 cm-l [fig. 4(b)]. Evacuation of the disc at 423 K completely returned the band at 3745 cm-l to its initial intensity but had a negligible effect on the remainder of the spectrum. After subsequent evacuation at 548 K the spectrum in the range 3300-3800 cm-l was identical to the spectrum in fig.3(a). The bands at 2957 and 1378 cm-l remained and were only slightly reduced in intensity by further heat treatment of the sample at 673 K for 30 min. Hydrocarbon species were desorbed by subsequent heatingin oxygen (673 K, 27 kN m-2, 72 h), and the two bandsdisappeared. A very weak maximum at 3705 cm-l was present in spectra after the thermal treatment in oxygen. The broad maximum due to OH-stretching vibrations of associated l,l, 1,3,3,3- hexafluoropropan-2-01 molecules in the liquid phase was at 3425 cm-l [fig. 4(c)].2290 1OC h 5 E 60 f e .- 111 .3 E U 20 FLUOROALCOHOLS+SILICA 3600 3000 1500 1300 wavenumber/cm-' FIG. 5.-Spectra of silica (a) in vacuum, (b)-(g) in contact with increasing pressures of ethanol vapour.(h), (i) and (J] Spectra of ethanol vapour which in contact with silica gave spectra (4, v) and ( g ) , respectively. Hydrogen-bonding interactions between surface silanol groups and adsorbed ethanol molecules led to similar spectroscopic changes (fig. 5) to those reported for the adsorption of methanol on silica.2y3*6 Two broad bands due to vibrations of perturbed hydroxy groups appeared at 3460 and 3350 cm-l together with a shoulder at 3610 cm-l probably due to stretching vibrations of unperturbed2 or weakly perturbed OH-groups in adsorbed ethanol molecules. Stretching vibrations of alkyl groups in adsorbed species gave maxima at 2982, 2940 and 2904 cm-l [fig. 5(d)]. At moderate surface coverages there was negligible contribution to the bands from the spectrum of the gas phase [fig.5(h)]. At low surface coverages the most prominent maxima due to CH-bending vibrations were at 1485 and 1380 cm-l [fig. 5(b)], the former band being absent from spectra of ethanol vapour [fig. 5(j)] or liquid [fig. 6(c)]. At high coverages a triplet of bands at 1445, 1400(sh) and 1380 cm-l [fig. 5(g)] resembled the triplet at 1448, 1425 and 1380 cm-l for liquid ethanol [fig. 6(c)]. The latter also gave a broad maximum at 3330 cm-l. Desorption of weakly bound ethanol molecules revealed infrared bands at 2988, 2942,2905(sh), 1493, 1427 and 1380 cm-l [fig. 6(a)] due to vibrations of chemisorbed species which could not be completely desorbed by thermal activation up to 673 K in vacuum or oxygen [fig.6(b)]. The enhancement of absorption intensity in the range 3300-3700 cm-l [fig. 5(a) and 6(a)] arising from ethanol chemisorption was progressively lost by heat treatment at increasing temperatures up to 673 K. The band at 3745 cm-l was not completely returned to its initial intensity before the adsorption of ethanol. Heat treatment at 673 K in oxygen generated surface silanol groups for which the OH-stretching vibrations gave a band at 3705 cm-l [fig. 6(6)] which was more intense than similar bands observed after the adsorption of either of the two fluoroalcohols.ACOSTA SARACUAL, PULTON, VICARY AND ROCHESTER 229 1 looIj 7 1 I 1 I 3800 3500 3600 3000 wavenumbedcm-' 1500 1300 FIG. 6.-Spectra of silica after (a) exposure to ethanol vapour [fig.5(g)] and evacuation (ca. 306 K, 1 h), (b) further evacuation (ca. 306 K, 24 h; 423 K, 1 h; 548 K, 1 h; 673 K, 72 h), heat treatment in oxygen (27 kN m-*, 673 K, 20 min) and evacuation (ca. 306 K, 5 min). Spectra of ethanol (c) liquid, (d) vapour. 7 60 20 3 600 3200 3600 3200 wavenumber/ cm -' I I I I I 1 3 600 3200 FIG. 7.-Spectra of silica immersed in (a) carbon tetrachloride, and solutions in carbon tetrachloride of (b) and (c) ethanol (0.01 7 and 0.054 mol dm-3, respectively), (d) 2,2,2-trifluoroethanoI (0.013 mol drnp3) and (e) 1, 1,1,3,3,3-hexafluoropropan-2-o1 (0.014 mol drn-9.2292 FLUOROALCOHOLS+SILICA Spectra of silica immersed in solutions of the alcohols in carbon tetrachloride are shown in fig. 7. Isolated silanol groups give a band at ca.3690 cm-l when at the silicalcarbon tetrachloride interface. l6 Alcohol solutions exhibited maxima due to OH-stretchingvibrations at 3635 cm-1 for ethanol, 3620 cm-l for 2,2,2-trifluoroethanol and 36 15 and 3580 cm-l for 1, 1 , 1,3,3,3-hexafluoropropan-2-01. Bands due to vibrations of perturbed hydroxy groups when the alcohols were adsorbed onto silica appeared at 3490(sh) and 33 10 cm-l for ethanol, 3430 cm-l for 2,2,2-trifluoroethanol and 3390 cm-l for 1 ,I, 1,3,3,3-hexafluoropropan-2-01. The spectroscopic shifts AVOH which occurred when the three alcohols formed hydrogen bonds with diethyl ether in carbon tetrachloride were 145 cm-l for ethanol, 240 cm-l for 2,2,2-trifluoroethanol, and 370 cm-l for 1,1,1,3,3,3-hexafluoropropan- 2-01.The equilibrium constants for the association of alcohol+ether to give a hydrogen-bonded complex at 298 K were 1.3, 5.8 and 65.6 mol-1 dm3, respectively. DISCUSSION The AVoH values for the formation of hydrogen bonds between alcohols and diethyl ether compare favourably with previous values of 144 and 232 cm-l for ethano117 and 2,2,2-trifluoroethano1,l8 respectively. An equilibrium constant of 5.2 mol-1 dm3 for the association of 2,2,2-trifluoroethanol and diethyl ether in carbon tetrachloride at 297 1 K has been reported.18 Equilibrium constants of 1.1 (at 294.9 K),17 6.4 (298 K, AVO, = 251 cm-l)19 and 36.4 (298 K, 382 cm-l)19 have been quoted for the association of ethanol, 2,2,2-trifluoroethanol and 1 ,1,1,3,3,3-hexafluoropropan-2-o1, respectively, with di-isopropyl ether in carbon tetrachloride.The corresponding enthalpy changes, AH, accompanying hydrogen-bond formation were - 17.5, - 21.3 and - 29.6 kJ m01-i.i73 l9 Fig. 8(a) illustrates the magnitude of the changes in ACoH which resulted I I I I 400 - ' 300 5 2 \ X 0 200 100 10 12 14 16 PK, FIG. 8.-Correlation between AS,, and pK, data for the formation of hydrogen bonds between alcohols and (a) diethyl ether, (b) surface silanol groups.ACOSTA SARACUAL, PULTON, VICARY AND ROCHESTER 2293 from variations in the acid strength of the hydrogen-bond donor molecule. Increased acidity enhanced the strength of the hydrogen-bonding interaction and more than doubled the resulting spectroscopic shift, AiiOH. The interpretation of AVoH values for hydrogen-bonding interactions between surface silanol groups and adsorbed alcohol molecules is complicated by the involvement of hydroxy groups in both the donor and acceptor species.An 100 mg silica disc with 2.54 cm diameter gives an absorbance value of ca. 1.05 at 3745 cm-l. The surface population of isolated silanol groups was ca. 2 nm-2.20 Hence the extinction coefficient (referred to conventional solution chemistry units) at the 3745 cm-’ absorption maximum is 84 cm-l mol-l dm3, which is of the same order of magnitude as the extinction coefficients at the absorption maxima due to OH-stretching vibrations of the monomer alcohols in carbon 22 The broad maxima between 3300 and 3600 cm-l in fig. 1 , 3 and 5 contain contributions due to not only perturbed silanol groups but also perturbed alcohol hydroxy groups.Despite this, values of the differences AVoH between the positions (in cm-l) of the broad bands and 3745 cm-l should still suggest whether complex (I) or complex (11) represents the predominant mode of interaction between adsorbed alcohols and the silica surface. Increasing alcohol acidity would strengthen the hydrogen bond in complex (11), leading to greater perturbations of both the alcoholic hydroxy group and the silanol group. The band positions for OH-stretching vibrations of alcohol monomers are similar [fig. 1 ( g ) , 3 ( g ) and 5 ( j ) ] , and therefore the separate spectroscopic shifts for alcohol hydroxy groups and silanol groups would both contribute in the same sense to a shift to lower wavenumber (AVO, bigger) of the overall band envelope.For complex (I) increasing alcohol acidity should weaken the interaction, perturb both hydroxy groups to a lesser extent, and lead to smaller spectroscopic shifts for the separate hydroxy groups and hence for the overall band envelope. The shifts AVO, measured from 3745 cm-l at half coverage of isolated surface silanol groups were 255, 315 and 395 cm-l for the adsorption of 1,1,1,3,3,3- hexafluoropropan-2-01, 2,2,2-trifluoroethanol and ethanol, respectively. Increasing acidity of the alcohols led to a smaller spectroscopic shift [fig. 8(b)]. Structure (I), in which the alcohol molecule acts as an hydrogen-bond acceptor, must therefore represent the predominant surface-adsorbate complex. The preference for structure (I) over structure (11) is consistent with estimates which suggest that pK, is ca.7 for isolated silanol groups in water at 298 K.23- 34 Silanol groups are more acidic than the three alcohols and are therefore more likely to act as the donor species in hydrogen- bonding interactions. Fig. 9 shows the relationships between losses in absorbance at 3745 cm-l due to the perturbation of isolated silanol groups and the appearance of broad infrared bands due to vibrations of perturbed hydroxy groups. Three aspects of these plots give information about the absorption of alcohol vapours on silica. The curvature at low coverages is characteristic77 25-28 of adsorption of single alcohol molecules on to two, or possibly three,2 adjacent surface silanol groups. The latter were responsible for a shoulder at 3680 cm-l in spectra of silica [fig.1 (a)] before adsorption took place and have been shown to be stronger sites than isolated silanol groups for the adsorption of ketones,’. 2 5 7 28 esters2’ and carboxylic acids26 on silica. The linear sections of the plots in fig. 9 all have slopes of ca. - 0.77 and characterize a one-to-one interaction between isolated surface silanol groups and adsorbed alcohol molecules in accordance with structure (I). The possibility that a single adsorbate molecule might interact simultaneously with two isolated silanol groups29 is unlikely to have been significant since hydrogen-bonding interactions between phenol hydroxy groups and isolated silanol groups have been proved primarily to involve one2294 FLUOROALCOHOLS+SILICA I 0 0.4 0-8 absorbance FIG.9.-Correlations between absorbance losses at 3745 cm-1 and absorbance gains at 3430, 3490 and 3350 cm-I for the adsorption of (a) 2,2,2-trifluoroethanoI, (b) 1,1,1,3,3,3-hexafluoropropan-2-ol and (c) ethanol, respectively. adsorbed molecule per silanol group.2o The assumption of a one-to-one interaction for phenol adsorption led to consistency between the number of molecules of phenol adsorbed and the population of isolated silanol groups per unit area of surface. The lines in fig. 9(a) and (b) remained straight up to the highest surface coverages for which, spectra were recorded. At high coverages the adsorption of each 2,2,2- trifluoroethanol or 1 , 1 , 1,3,3,3-hexafluoropropan-2-01 molecule continued to cause the perturbation of one further surface silanol group. Although the broad band maxima shifted slightly to lower wavenumbers with increasing coverage there were no marked changes in band shape and the band positions were not as low (in cm-l) as those for the liquid fluoroalcohols.Lateral interactions between adjacent adsorbed fluoro- alcohol molecules may have been responsible for the band shifts but the extents of the interactions were not as great as in the case of methanol adsorption on silica, for which the formation of adsorbed polymeric chains has been proposed.2 The non-linearity at high coverages of the relationship plotted in fig. 9(c) suggests that, like methanol,2 ethanol formed polymeric aggregates on the silica surface. Comparison of the spectra in fig. 5(g) and 6(c) shows that for ethanol on silica at high coverages there was enhanced absorption intensity within the 3000-3300 cm-l spectral range, which may be assigned to OH-stretching vibrations of ethanol molecules, not only linked to surface silanol groups through hydrogen bonds but also involved in lateral hydrogen-bonding interactions with neighbouring ethanol mole- cules.The downward curvature of the line in fig. 9(c) occurred at ca. 60% coverage of silanol-group adsorption sites. Ethanol molecules perturbed by interaction with silanol groups and other ethanol molecules apparently contributed less to the overall absorption intensity at 3350 cm-1 than molecules which only interacted with silanolACOSTA SARACUAL, PULTON, VICARY A N D ROCHESTER 229 5 groups [structure (I)].However, at the highest coverages of silanol-group sites the growth of absorbance at 3350 cm-l again became significant compared with the losses at 3745 cm-l [fig. 9(c)]. This effect is ascribed to the onset of multilayer adsorption to give ethanol aggregates akin to liquid ethanol, for which the absorption maximum due to OH-stretching vibrations was at 3330 cm-l [fig. 6(c)]. The maximum at 3460 cm-l in spectra of ethanol on silica [fig. 5(e) and (f)] was compatible with a corresponding band at 3500 cm-l for methanol on silica.2*3,5,6 A second band at 3500 cm-l appears when methanol chemisorbs on to strained siloxane surface groups to form adjacent hydroxy and methoxy groups which interact laterally through a hydrogen-bond linkage.'* The absorbance increases in the range 3300-3700 cm-l, which resulted from adsorption of ethanol [fig.6 (a)], 2,2,2-trifluoroethanol [fig. 2(6)] or l , l , 1,3,3,3-hexafluoropropan-2-o1 [fig. 4(b)] followed by subsequent evacu- ation, may be ascribed to similar chemisorption reactions. An alternative explanation5 would be that the absorbance increases resulted from the chemisorption of water formed by esterification of isolated surface silanol groups to give surface alkoxy or fluoroalkoxy groups. This explanation is less acceptable, particularly for 1 , l ,1,3,3,3- hexafluoropropan-2-01, the adsorption of which, followed by evacuation at ca. 306 K, resulted in only a 2% decrease in the intensity of the maximum at 3745 cm-l (fig. 4). Further evacuation at 423 K completely restored the maximum to its original intensity, confirming that negligible esterification of silanol groups had occurred.More significant decreases (ca. 16%) in band intensity at 3745 cm-l resulted from adsorption of ethanol or 2,2,2-trifluoroethanol followed by evacuation at ca. 306 K. For 2,2,2-trifluoroethanol further evacuation at 673 K completely restored the band intensity. However, for ethanol evacuation at 423 and 673 K led to 11 % and 5% deficits, respectively, in the absorbance value. Some esterification of silanol groups had probably occurred when ethanol was adsorbed on silica. The bands at 2988,2942 and 2905 cm-l in spectra of ethanol chemisorbed on silica correspond to similar bands previously observed at 2986,2938 and 2907 cm-l and reported to be compatible with the infrared spectrum of ethyl ortho~ilicate.~~ The weak band at 3705 cm-l which appeared after high-temperature treatment of chemisorbed alcohols on silica was most intense after the adsorption of ethanol [fig.6 (b)]. A similar band at 3700 cm-l, which resulted from high-temperature treatment of the products of chemisorption of methyl fluorosulphate on silica, was ascribed to silanol groups which had +SiF groups adjacent to them on the silica surface.31 The same explanation would be tenable for the fluoroalcohols, providing that the chemisorbed fluoroalkoxy groups formed at ambient temperatures were decomposed at 673 K rather than being desorbed intact. However, for ethanol the infrared band may be compared with the maximum at 3705 cm-l in spectra of silica immersed in liquid heptane13 or 2,2,4-trimethylpentane,28 and by analogy is assigned to isolated surface silanol groups perturbed by interaction with ethyl groups chemisorbed on adjacent sites either as --+ SiOEt (with an OH to OEt separation which was too great for OH * - 0 E t hydrogen bonding) or possibly as + SiEt.The 5% loss of intensity of the maximum at 3745 cm-l which resulted from the sequence of treatments [fig. 5 (a)-6(b)] is therefore ascribed to these perturbation effects rather than to esterification of the silanol groups themselves. The adsorption of ethanol at ca. 306 K may have caused esterification of a few silanol groups, but the latter were regenerated by high-temperature treatment. We thank the University of Oriente, Venezuela, for financial support.2296 FLUOROALCOHOLS+SILICA E.Borello, A. Zecchina and C. Morterra, J. Phys. Chem., 1967, 71, 2938. M. J. D. Low and Y. Horano, J. Res. Znst. Catal. Hokkaido Univ., 1968, 16, 271. R. S. McDonald, J. Phys. Chem., 1958, 62, 1175. M. Baverez and J. Bastick, J. Chim. Phys., 1969, 66, 935. A. V. Kiselev, V. I. Lygin and K. L. Shchepalin, Kinet. Catal., 1971, 12, 154. P. Ballinger and F. A. Long, J. Am. Chem. SOC., 1960, 82, 795. P. Ballinger and F. A. Long, J. Am. Chem. SOC., 1959, 81, 1050. * E. Borello, A. Zecchina, C. Morterra and G. Ghiotti, J. Phys. Chem., 1967, 71, 2945. ' D. M. Griffiths, K. Marshall and C. H. Rochester, J. Chem. SOC., Faraday Trans. 1, 1974, 70, 400. lo C. W. Roberts, E. T. McBee and C. E. Hathaway, J. Org. Chem., 1956, 21, 1369. l 1 W. J. Middleton and R. V. Lindsey, J. Am. Chem. SOC., 1964, 86, 4948. l2 C. H. Rochester and D-A. Trebilco, J. Chem. SOC., Faraday Trans. 1, 1978, 74, 1137. l3 C. H. Rochester and D-A. Trebilco, J. Chem. SOC., Faraday Trans. 1, 1978, 74, 1125. l4 P. Jackson and G. D. Parfitt, Trans. Faraday Soc., 1971, 67, 2469. l5 J. Murto, A. Kivinen, R. Viitala and J. Hyomaki, Spectrochim. Acta, Part A, 1973, 29, 1121. l6 C. H. Rochester, Adv. Colloid Interface Sci., 1980, 12, 43. l7 I. Motoyama and C. H. Jarboe, J. Phys. Chem., 1967, 71, 2723. l8 A. D. Sherry and K. F. Purcell, J. Phys. Chem., 1970, 74, 3535. l9 A. Kivinen, J. Murto and L. Kilpi, Suom. Kemistil. B, 1967, 40, 301. 2o K. Marshall and C. H. Rochester, J. Chem. Soc., Faraday Trans. 1, 1975 71, 2478. *l A. Kivinen and J. Murto, Suom. Kemistil. B, 1967, 40, 6. 22 J. Murto and A. Kivinen, Suom. Kemistil. B, 1967, 40, 14. 23 M. L. Hair and W. Hertl, J. Phys. Chem., 1970, 74, 91. 24 K. Marshall, G. L. Ridgewell, C. H. Rochester and J. Simpson, Chem. Znd. (London), 1974, 775. 25 K. Marshall and C. H. Rochester, Faraday Discuss. Chem. SOC., 1975, 59, 117. 26 K. Marshall and C. H. Rochester, J. Chem. SOC., Faraday Trans. I, 1975, 71, 1754. 27 S. N. W. Cross and C. H. Rochester, J. Chem. SOC., Faraday Trans. I, 1979, 75, 2865. 28 C. H. Rochester and D-A. Trebilco, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 2211. *' A. D. Buckland, C. H. Rochester, D-A. Trebilco and K. Wigfield, J. Chem. SOC., Faraday Trans. 1, 30 L. Abrams and A. 0. Allen, J. Phys. Chem., 1969,73, 2741. 31 D. D. Eley, G. M. Kiwanuka and C. H. Rochester, J. Chem. SOC., Faraday Trans. 1, 1975,71,2340. 74, 2393. (PAPER 1 / 1786)

ISSN:0300-9599    

DOI:10.1039/F19827802285

出版商:RSC    

年代:1982

数据来源: RSC