摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 3235-3247 I on-I on-S olven t Interactions in Solution Part 4.-Raman Spectra of Aqueous Solutions of Some Nitrates with Monovalent Cations BY RAY L. FROST Chemistry Department, Queensland Institute of Technology, Brisbane, Australia 400 1 AND DAVID W. JAMES* Chemistry Department, University of Queensland, Brisbane, Australia 4067 Received 2 1st December, 198 1 The profile of the band due to the symmetric stretching of the nitrate ion in the Raman spectra of solutions of NH,NO,, Me,N - NO,, LiNO,, KNO, and RbNO, has been component analysed. For the cations NH: and Me,N+ only one component was present. The bandwidth indicates that the ammonium ion causes enhanced dephasing while the tetramethylammonium ion produces essentially no perturbation.Solutions of LiNO, gave evidence of four concentration-dependent species identified as aquated ions, solvent-sepa- rated ion pairs, contact ion pairs and ion aggregates. Solutions of KNO, and RbNO, gave evidence of three species identified as aquated ions, ion pairs and ion aggregates. It was not possible to distinguish between contact ion pairs and solvent-separated ion pairs. Association quotients calculated for LiNO, solutions were Kl( 1 mol drn-,) = I .34 dm3 mol-l and K , (6 mol drn-,) = I .5, for KNO, solutions K( 1 mol dm-,) = 2.0 dm3 rno1-l and for RbNO, solutions K( 1 mol dm-,) = 1 .O dm3 mol-'. Nitrates having monovalent cations produce a variety of effects in aqueous solution. Ammonium nitrate solutions give little indication of disturbance of the anion by the cation1V2 and it has been assumed that the ammonium ion, with a tetrahedral shape, fits well into the hydrogen-bonded network of water. Solutions of the related salt tetramethylammonium nitrate, gave little evidence for a concentration-dependent disturbance of the nitrate ion but the structural disturbance by the large cation was reflected in the infrared spectrum.Spectra of the NO; ion in solutions of lithium nitrate showed a concentration-dependent splitting of the band due to the antisymmetric stretching vibration (v,) in the Raman spectra,, appearance of a second band in the spectral feature due to the in-plane bending vibration (v,) at 738 cm-l in water5 and 734cm-1 in D,0,6 and at high concentrations asymmetry in the band due to the symmetrical stretching vibration ( v , ) .~ The spectra of the nitrate ion in solutions of potassium nitrate show a splitting of v3 which is less than for solutions of LiNO, or NaNO,., In addition, although Y, shows some asymmetry no new features are evident. Studies of the band profile of the v1 band in the Raman spectra*vg have shown that the vibrational relaxation time is almost independent of concentration when the cations are NH,+ or K+ but for Na+ and Li+ the value of z, decreases with concentration, with Li+ > Na+. In this study the Raman band due to the symmetrical stretching vibration was studied for aqueous solutions of the salts LiNO,, KNO,, RbNO,, NH,NO, and Me,N*NO, from dilute solution to saturation. Using the techniques described in the previous paper these spectra have been component band analysed.This analysis will be described for each solute and then the overall trends in results will be examined. 32353236 I 0 N-ION-SO L V E N T I N TER A C T I 0 N S EXPERIMENTAL The salts were all recrystallised twice from water, dried at 110 "C for 48 h and stored over P,O,. All other experimental details and methods of data treatment were as described in the previous paper.") RESULTS A N D DISCUSSION SOLUTIONS OF LiNO, The time correlation functions for solutions of LiNO, are shown in fig. 1 A, which includes the transform for a Lorenzian profile of half-width 3.4 cm-I. The imaginary part of the transform is appreciable and increases rapidly with concentration. This behaviour is compatible with a band composed of more than one component in which the component(s) growing at higher concentrations have a greater half-width than the dilute solution band.The correlation functions at various concentrations also give evidence of weak minima, and in particular the transforms of 8 and 10 mol dm-, solutions have minima at 3.5-4 ps, indicating two components separated by 4-5 cm-1. The time variation of the 8 function derived from the real and imaginary parts of the transform is shown in fig. 1 B. In 1 mol dm-3 solution the 8 function already lies above that of the dilute solution band and the point of intersection is beyond the range of the figure at cu. 7-8 ps, corresponding to a separation between components of 2-2.4 cm-'. The curves for concentrations of 2,4 and 6 mol dm-3 all give evidence of a second minimum at ca.4 ps, corresponding to a peak separation of ca. 4 cm-l, while at the highest concentrations a weak inflection at short times, 1.5-2.5 ps, indicates a separation of ca. 10 cm-l. On this basis the initial band components were positioned at 1047.6, 1050 and 1052 cm-1 with allowance for a fourth component at higher energy. The results of the component analysis of the spectroscopic band at various concentrations are given in table 1 with the variation in band intensity (band area) as a function of concentration being shown in fig. 2. The four components are at 1047.6, 1050.0, 1053.0 and 1065 cm-', and all are dominantly Lorentzian with the half-width increasing from 3.4 cm-l for the lowest-energy band to 6.2 cm-I for the highest-energy band.As for solutions of NaNO, these results are compatible with a three-stage concen tration-dependen t set of equilibria K , K , K3 Li+(aq) + NO,(aq) + Li+ H,O NO;(aq) + Li+ * NO;(aq) + (Li+ NO;),(aq). The values for K, and K2 are included in table 1 . In dilute solution the obtained value of K, of 1.8 dm3 mol-l is smaller than for solutions of NaNO, and similarly the concentration-dependent values for K2 are lower than corresponding values for NaNO, solutions. These comparisons will be discussed in more detail later in the paper. The component-band characteristics illuminate the nature of the association process. Theenergy and band profile of the component assigned to the solvent-separated ion pair is the same as in solutions of NaNO, but the half-width for LiNO, solute is significantly greater (7, = 1.32 ps).This means that the average coulombic field perturbing the NO; and the nature of the relaxation process is the same for both solutes but the symmetry perturbation imposed on the nitrate ion is greater for the lithium cation. This is in accord with the observation that even at low concentrations the splitting of the v3 band is greater in solutions of LiNO, than in solutions of NaNO,. The paradox of the nitrate ion in the solvent-separated ion pair having the sameR. L. FROST AND D. W. JAMES 3237 coulombic perturbation but a greater symmetry perturbation can be resolved by considering the dynamic nature of the solution. Because of its small size the lithium ion is more strongly hydrated that the sodium ion.This means that the residence time for water in the hydration sheath is longer. This in turn means that when the anion I I I 1.54 3.08 4.63 6.17 t i p s 1 1 I 0.94 1.88 2.82 3.75 t/PS FIG. 1.-A, Time correlation functions [concentration/mol dm-3: (a) 0.05; (b) 0.5; ( c ) 4; (6) 6; (e) 8; ( f ) 101 and B, theta functions [concentration/mol drn-,: (u) 0.1; (b) 1; ( c ) 14; (d) 8; (e) 101 from A; spectral bands of aqueous solutions of LiNO,. and cation have a common hydration layer the cation exchanges more slowly and so has a stronger localised perturbation. However, because of the mediation of the water the cation does not have a preferred orientation with regard to the anion and so the coulombic perturbation is still averaged. The stronger aquation of the cation will alsow h, w 00 TABLE 1 .-SECOND MOMENT, MODULATION TIMES AND RELAXATION FOR AQUEOUS (GROUP I) NITRATE SOLUTIONS concentration electrolyte /mol dm-, v,/crn-l vg/cm-l kit,, /(cm-1)2 (Zb/a) x 100 z,/ps z, 4 / 2 n c z,(@/ps z,(e-')/ps LiNO, 1 .o 2.0 4.0 6.0 8.0 10.0 KNO, 0.5 1 .o 2.0 2.75 2.0 3.0 CsNO, 0.5 1 .O RbNO, 1 .o 1048.3 1048.7 1049.0 1050.1 1050.8 1051.9 1048.2 1048.7 1048.9 1049.1 1048.0 1048.7 1048.9 1048.3 1048.7 1048.3 1048.9 1049.8 1050.8 1051.3 1052.5 1048.4 1048.9 1049.1 1049.4 1047.0 1049.1 1049.0 1048.6 1049.0 63.0 68.0 73.2 85.6 95.0 96.0 58.0 58.0 62.0 64.0 56.2 66.0 56.2 59.0 59.0 -0.84 + 1.70 + 3.9 + 4.4 + 8.4 +11.9 + 0.9 1 + 0.63 - 4.02 - 2.50 - 1.15 - 1.22 - 0.89 -0.80 - 0.44 0.76 0.69 0.64 0.62 0.77 0.78 0.87 0.88 0.84 0.83 0.96 0.95 0.96 0.93 0.93 0.80 0.75 0.73 0.77 0.99 1 .oo 0.88 0.90 0.88 0.89 0.96 1.02 1.01 0.95 0.95 1.38 1.29 1.14 1.07 0.92 0.81 1.36 1.36 1.31 1.30 1.34 1.34 1.34 1.36 1.30 1.19 1.21 1.06 0.95 0.77 0.76 1.12 1.10 1.08 1.08 1.05 1.05 1.05 1.06 1.06 0 2: mR.L. FROST AND D. W. J A M E S 3239 favour the formation of isolated ions, a tendency which is evident from the relative concentrations of aquated nitrate ions in the two solutes. The band assigned to the contact ion pair species indicates that both the coulombic and symmetry perturbations are greater for the Li+ counter-ion than for Na+. However, the band remains much more Lorentzian than for NaNO, solute indicating that crystal-type damping is less important for LiNO, solute. The strong aquation of the cation is again a possible cause of these differences.In the contact ion pair the relative orientation of the cation and anion is going to be influenced by the constraint of a common aquation shell. The strong aquation of the lithium ion may well constrain the cation to take up a position 2 4 6 8 10 [ LiNO, I /mol dm-3 FIG. 2.-variation of component-band intensities for the A; spectral bands from fig. 1 : (a) aquated-nitrate band ( 1047.6 cm-'), (b) solvent-separated-ion-pair band (1050 cm-l), (c) contact-ion-pair band (1053 cm-'), (d) ion-aggregate band (1065 cm-I), x = data from ref. (5). which is incompatible with crystal packing. This in turn will make the contact ion pair less favourable than the ion aggregate at high concentrations, as is observed. In fig.2 the previous e s t i m a t e ~ ~ ? ~ of the concentration of 'ion pairs' in this system are included and they are seen to coincide with the intensity of the ion-aggregate band. SOLUTIONS OF KNO, AND RbNO, The time correlation functions obtained on Fourier transformation of the isotropic profile are similar for these two solutes. The functions for potassium nitrate solutions are given in fig. 3A. There is little variation with concentration and no minima are obvious. This means that if there is more than one component the band half-widths are little different and the components will not be widely separated. The theta functions shown in fig. 3 B show deviation from the dilute-solution band at all concentrations with a clearly developed inflection at ca.6.5 ps which corresponds to a band separation of ca. 2.4cm-l. Band component analysis yielded the results collected in table 1, from which the similarity between the two systems is obvious. For both systems there are two main3240 I 0 N-I0 N-SO L V E N T INTER A C T I 0 N S I I I I 1.54 3.08 4.63 6.17 tips tlPS FIG. 3.-A, Time correlation functions [concentration/mol dm-3 (a) 0.01 ; (b) 0.05 (c) 1 ; (6) 21 and B, theta functions [concentration/mol dm-3: (a) 0.01; (b) 0.5; (c) 1; ( d ) 21 from A ; spectral bands of aqueous solutions of KNO,. components at 1047.6 and 1050.0 cm-l and a weak component at much higher energy. The variation of intensity of the three components with concentration for solutions of KNO, is shown in fig. 4. The presence of two rather than three major bands signals behaviour in these solutions which is distinctly different from that in solutions of LiNO, or NaNO,.The two band components are similar in shape with the higher-energy band a little broader. As the size of the cation increases the position of the contact-ion-pair band decreases and so it would be expected that there would be little difference in position of the solvent-separated-ion-pair band and the contact-ion-pair band in solutions of KNO, and RbNO,. Because the band at 1050.0 cm-l has a half-widthR. L. FROST AND D. W. JAMES 324 1 which is of the magnitude expected for a solvent-separated ion pair and because this half-width does not change with concentration, the associated species can be described with some confidence.The change in energy from 1047.6cm-l [NO;(aq)] to 1050.0 cm-’ [NO;(assoc)] can in the first instance be associated with the change in coulombic field with water in the second hydration shell is replaced by K+. The next stage would be the replacement of an H 2 0 molecule in the primary hydration shell with a K+ ion. The coulombic perturbation for these two configurations is clearly similar. However, the symmetry perturbation imposed by a mixed primary shell would certainly cause an increased dephasing and a broadening of the 1050.0 cm-’ band as the concentration increases. That such an increase in half-width is not observed indicates that the relative exchange rates for water and K+ ions with the nitrate ion 1 . 0. 0 . 1 2 3 [KNO,] /mol dm-3 FIG. 4.-Variation of component-band intensities for the A ; spectral bands from fig.3 : (a) aquated-nitrate band (1057.6 cm-’); (b) ion-pair band (1050 cm-l), (c) ion-aggregate band (1067 cm-l). are similar and they are probably also similar to the exchange rates of water molecules in the hydration shell of the K+ ion. This gives rise to a situation where the two species K+ H 2 0 * NO;(aq) and K+ - NO; - H,O(aq) are not distinguishable. Entropic considerations will then lead to the most random arrangement. At higher concentrations the weak driving force towards aquated ions leads to the formation of ion aggregates and low solubility. The solution behaviour can be described in terms of a two-step equilibrium K , K , K+(aq)+NO;(aq)eK+-NO;-H,O(aq)$(K+. NO;),(aq). The constant Kl is not really comparable with Kl or K2 for solutions of NaNO, and LiNO,.The values for KNO, solutions are significantly greater than for RbNO, solutions, which indicates that association is weaker for the larger cation, as might be expected intuitively. SOLUTIONS OF NH,NO, AND Me,N-NO, These two solutes demonstrate distinctly different behaviour to any of the others so far examined. As shown in table 2 neither solute produces aconcentration-dependent shift in the band maximum although the band position for Me,N-NO, is lower at 1046.6 cm-l than that any of the other solutes. The time correlation functions and 105 FAR 78w h, P h, TABLE 2.-cOMPONENT-BAND ANALYSIS FOR THE ISOTROPIC BANDS OF GROUP I NITRATES aquated-nitrate band solven t-separated- ion-pair band shape concentration v M k o .1 4 k 0 . 0 1 ratio area 4 shape electrolyte /mol dm-3 /cm-l /crn-l f 0.00 1 kO.01 vM/cm-l /crn-l ratio area LiNO, 0.5 1 .o 2.0 4.0 6.0 8.0 10.0 KNO, 0.5 1 .o 2.0 2.75 2.0 3.0 RbNO, 1 .o 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 1047.6 3.397 0.852 0.62 1 - 0.561 - 0.575 - 0.37 - 0.327 - 0.1 15 - 0.042 3.395 0.851 0.725 - 0.504 - 0.398 - 0.347 3.395 0.852 0.602 - 0.504 - 0.465 - - - - - - - - - - - 1050.0 1050.0 1050.0 1050.0 1050.0 1050.0 1050.0 0.725 3.95 3.96 4.01 4.03 4.03 4.03 4.03 0.899 0.380 0.844 0.421 0.844 0.388 0.844 0.38 0.844 0.353 0.844 0.297 0.844 0.198 m 0TABLE 2. (con?.) contact-ion-pair band ion-aggregate band concentration 4 shape 4 shape k1/dm3 k,/dm3 electrolyte /mol dm-3 vM,/cm-l /cm-l ratio area vM,/cm-' /cm-l ratio area mol-I rno1-l LiNO, 0.5 1 .o 2.0 4.0 6.0 8.0 10.0 KNO, 0.5 1 .o 2.0 2.75 2.0 3.0 RbNO, 1 .o 1053.0 1053.0 1053.0 1053.0 1053.0 1053.0 1053.0 1050.0 1050.0 1050.0 1050.0 1050.0 1050.0 1050.0 3.60 3.52 3.46 4.90 5.22 5.32 5.16 3.7 3.7 3.7 3.7 3.69 3.69 3.69 0.920 0.889 0.848 0.850 0.862 0.800 0.863 0.852 0.852 0.852 0.852 0.86 1 0.861 0.861 0.001 0.00 1 0.0385 0.24 0.3046 0.529 0.5496 0.275 0.496 0.387 0.623 0.376 0.479 0.531 - 1065.0 1065.0 1065.0 1065.0 - 1067.0 1067.0 1067.0 1057.5 1057.5 1057.5 - 6.0 I 6.42 6.41 6.12 - 7.00 7.05 7.07 6.56 6.60 6.60 - - 0.80 0.01 0.81 0.03 0.83 0.057 0.763 0.207 0.80 0.00 1 0.81 0.01 5 0.80 0.022 0.76 0.0 1 0.76 0.013 0.76 0.016 1.79 - 1.34 0.59 _ _ 0.86 1 .o 0.55 I .5 1.8 - 2.8 1 .O - 1.95 1.9 - 1.9 - 1 .o - 0.95 0.82 - - - - - w Id P w3244 I 0 N-I 0 N-SO L V E N T I N T E R A C T I 0 N S - 0.140 - 0 .2 6 5 a - f -0.445 -0.500 - 0 . 7 4 1 1 . 6 3.2 4.8 6 .I 0.0157 -0.688- -1.35 - -2.84 1 I I I I 1 . 6 3 . 2 4.8 6 . 4 a .a TIPS FIG. 5-A, Time correlation functions [concentration/mol dm-3: (a) 2; (6) 101 and B, theta functions concentration/mol dm-3: (a) 0.01; ( b ) 1 ; ( c ) 2; (d) 41 from A ; spectral bands-of aqueous solutions of NH,NO,. theta functions obtained on Fourier-transforming the isotropic band profile for solutions of NH,NO, are shown in fig. 5. The time correlation functions show little change with concentration and the theta functions are essentially invariant for concentrations of 1-6 mol dm-3 while for 8 and 10 mol dm-, the functions deviate and have an intersection at ca.13 ps, which corresponds to a peak separation of ca. 1.2 cm-l. Because the experimental isotropic spectra showed almost no concentration dependence and the time correlation functions and theta functions were simple, a band component analysis is not reported for these two solutes. For solutions of ammonium nitrate the band half-width of 3.95 cm-l is greater than that for the dilute-solution band, 3.4 cm-l, and this indicates that the ammonium ion enhances the dephasingTABLE 3.-BAND PARAMETERS FOR THE A ; SYMMETRICAL STRETCHING VIBRATION FOR AQUEOUS SOLUTIONS OF AMMONIUM NITRATE AND TETRA-ALKYLAMMONIUM NITRATE F concentration electrolyte /mol dmP3 v,/cm-l vs /cm - NH,NO, 1 .o 2.0 4.0 6.0 8.0 10.0 2.0 3.0 4.0 N(CH,),NO, 1 .o 1047.6 1047.7 1047.7 1047.7 1047.8 1047.8 1046.6 1046.6 1046.8 1046.8 1048.0 1048.1 1048.1 1048.1 1048.1 1048.1 1 046.8 1046.8 1046.9 1046.9 58.0 58.0 58.0 58.0 76.0 90.0 50.0 52.0 56.0 56.0 - 0.47 - 0.48 - 0.40 - 0.48 + 0.82 + 0.90 +4.7 + 4.4 + 1.5 + 1.9 0.85 0.86 0.87 0.89 0.68 0.57 0.79 0.73 0.65 0.64 0.86 0.88 0.88 0.90 0.78 0.72 0.74 0.70 0.65 0.64 1.34 1.36 1.36 1 .36 1.34 1 34 1.53 1.58 1.54 1.54 r .,(e-’)/ps ; z 1.12 2 1.14 4 1.12 + 1.09 p 1.07 1.09 ? 1.45 + 1.48 1.54 m 1.58 u w w R3246 I 0 N-I 0 N-SO L V EN T I N TE R A C T I 0 N S process, probably through an environment perturbation.For tetramethylammonium nitrate solute, on the other hand, the half-width corresponds closely to that of the dilute-solution band. The band position, however, is moved to lower energy by ca.1 cm-l. This movement of the band maximum may reflect the perturbation to the water structure by the large cation. In the group of electrolytes studied three different patterns of interaction can be identified. For the two non-metallic cations the charge is diffused over the ion and the interactions are dominated by size and ability to form hydrogen bonds. The ammonium ion is characterised by a tetrahedral shape which permits its acceptance into the water structure with only a small symmetrq. disturbance to the nitrate ion, which appears as band broadening. Only at very high concentrations does the anion-cation interaction produce any change in the spectrum. The tetramethyl- ammonium cation disturbs the water structure through its hydrophobic nature.It has been previously suggested that this allows preferential solvation of the anion.8 The shift of the nitrate band to lower energy may reflect this change in structure of the hydration shell which is unique to these large hydrophobic cations. The bandwidth in these solutions is the same as that we characterise for the aquated nitrate ion in the absence of cationic perturbation. Note that band component analysis was attempted for both of these solutes but no improvement in fit was obtained. The absence of additional components in these solutions supports the view that the changes we observe are not experimental artifacts but represent real trends. The four remaining solutes give either two (KNO, and RbNO,) or three (LiNO, and NaNO,) main concentration-dependent band components.If the influence of change in water structure is ignored there will be three interactions of importance which are the aquation of the cation, the aquation of the anion and the anion-cation contact interaction. The water-lithium interaction will be the strongest and this is reflected in the dominance of the aquated nitrate ion (which implies a corresponding concentration of isolated aquated cation) at low concentrations. For the sodium cation aquation forces will be less strong and the nitrate takes a position in the second hydration sphere at low concentrations, becoming the dominant species at 1 mol dm-, as against 4 mol dm-, for lithium nitrate solute. The strong aquation of the lithium ion also retards the formation of contact ion pairs which dominate the spectrum of NaNO, solutions from 4 mol dm-, while for LiNO, solute it does not dominate until 8 mol drn-,.For the cations potassium and rubidium the solvent molecules are more weakly held, which causes the exchange rate to be more rapid. The aquation of the nitrate ion, however, is not influenced by the change of cation. The energy of the second band for the potassium and rubidium cations indicates that the perturbation of the nitrate either by a water molecule shared with a cation or by a cation in contact are similar. In terms of exchange rates this is compatible with the water of solvation and the cation having similar exchange rates. This prevents the identification of the solvent-separated ion pair as a distinct species.In dilute solution the associated species is more likely to be a solvent-separated ion pair than a contact ion pair. The appearance of an ion-aggregate species in appreciable concentration, however, makes it likely that anion<ation contact is present in the associated species. The value of the association quotient at 1 mol dm-, concentration reflects the influence of aquation described above. The value for LiNO, solute of 1.35 dm3 mol-1 is lower than for either NaNO, solute or KNO, solute. The strong aquation of Li+, the influence of which extends past the first hydration shell, is undoubtedly responsible for the preferred species not being closely associated with the anion. The remaining three cations show a regular decrease in the association quotient with increasing cation size, which follows the expected trend.R. L. FROST AND D. W. JAMES 3247 The measurement of association quotients for the association equilibria opens the way to study the energy changes in the association equilibria. Some preliminary studies will be examined in the following paper together with the influence of a variety of anions with a range of solution properties. We thank the Australian Research Grants Committee for grants enabling the purchase and maintenance of the Raman spectrometer. Drs R. Appleby and M. T. Carrick are thanked for helpful discussions. P. M. Vollmer, J. Chem. Phys., 1963, 39, 2236. * D. W. James and R. L. Frost, J. Phys. Chem., 1976, 80, 501. D. W. James and R. L. Frost, J. Chem. Soc., Faraday Trans. 1, 1978, 74, 583. D. E. Irish and A. R. Davis, Can. J . Chem., 1968, 46, 943. D. E. Irish, D. L. Nelson and M. H. Brooker, J. Chem. Phys., 1971, 54, 654. J. D. Riddell, D. J. Lockwood and D. E. Irish, Can. J . Chem., 1972, 50, 2951. D. E. Irish and M. H. Brooker, in Advances in Infrared Raman Spectroscopy, ed. R. J. H. Clark and R. E. Hester (Heyden, London, 1976), vol. 2, p. 212. D. W. James and R. L. Frost, Furuday Discuss. Chem. Soc., 1978,64, 48. M. Kosibaa and M. Perrot, C.R. Acad. Sci., Ser. C, 1978, 286, 99. lo R. L. Frost and D. W. James, J. Chem. Soc., Faraday Trans. 1, 1982, 78, 3223. (PAPER 1 /2000)

ISSN:0300-9599    

DOI:10.1039/F19827803235

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Furaday Trans. I , 1982, 78. 3249-3261 Ion-Ion-Solvent Interactions in Solution Part 5.4nfluence of Added Halide, Change in Temperature and Solvent Deuteration on Ion Association in Aqueous Solutions of Nitrate Salts BY RAY L. FROST Chemistry Department, Queensland Institute of Technology, Brisbane, Australia 400 1 AND DAVID W. JAMES* Chemistry Department, University of Queensland, Brisbane, Australia 4067 Received 2 I st December, I98 1 Band-component analysis of the band due to symmetrical stretching of the nitrate ion in aqueous solutions has been used to study association equilibria. Variation of anion composition at constant ionic strength indicates that the association equilibria are weakly dependent on the nature and concentration of counter-ions. Variation of ionic strength at constant anion concentration produced large changes in the associated species.Association was found to be dependent on both the competitive association equilibria with other anions and of the structural disturbance of the solution. The measurement of the concentration of associated species as a function of temperature in solutions of NaNO, showed that the formation of both solvent-separated ion pairs and contact ion pairs is favoured by a reduction in temperature. Values of the enthalpy change associated with the equilibria *HI AH Na+ (aq) + NO; (aq) Na+ . H,O . NO; (aq) e Na+ . NO; (aq) were found to be AHl = -42.2 kJ mol-l and AH, = -64.56 kJ mol-l. In a comparison of association equilibria in H,O and D,O solutions of LiNO,, NaNO, and KNO, the association equilibria were all altered in a way compatible with the solvent D,O having weaker hydrogen bonding and aquating properties than H,O The effect of ionic strength on association equilibria in solution has been assumed to be small in any but dilute ~olutions.l-~ It is assumed that the ionic medium becomes ‘saturated’ and ionic strength is not an important variable at high concentrations.Different electrolytes have been classified as structure making and structure breaking according to their presumed influence on water s t r u c t ~ r e , ~ although this classification may not be appropriate.6 It is to be expected that different electrolytes will produce different effects on the solvent structure and may also have competing equilibria. Thus for a series of solutions having the same ionic strength the association process may be significantly different.In this paper the influence of ionic strength and different counter-anions is reported for association equilibria in aqueous solutions containing LiNO,, NaNO, and KNO,. Studies of the temperature dependence of the association equilibria in solution should yield useful thermodynamic information on solution processes. The limitations of Raman spectroscopy in this context have been recently reviewed.’ A possible major problem which has not been discussed previously is the assignment of spectroscopic components to particular solution species. An incorrect assignment will obviously lead to erroneous results and possibly values of AH and A S having the wrong sign.In spite of these shortcomings the spectroscopic method, with at least a possibility of 32493250 I 0 N-I ON-SO L V E N T I N ’I’ E R A C T I 0 N S identifying particular solution species, has an advantage over bulk measurements where the nature of the associated species is relatively undefined.6 We report in this paper a study of the temperature dependence of the components of the band due to the symmetrical stretching vibration of the nitrate ion in solutions of NaNO,. Water and D,O have been used interchangably as solvents and although attempts have been made to estimate differences in the strength of hydrogen bonding and deuterium bonding the results have not established a clear di~tinction.~ We examine the influence of D,O on the association equilibria for solutions of LiNO,, NaNO, and KNO,.EXPERIMENTAL D,O was used as supplied by the Australian Atomic Energy Commission (99.5%). The alkali halides were recrystallised twice from water, oven dried for 48 h at 110 O C and stored over P,O,. All other experimental and analytical details were the same as previously described.8 RESULTS AND DISCUSSION ASSOCIATION OF MNO, I N THE PRESENCE o’F MX For all solutions studied the correlation functions and theta functions showed strong similarity to those for the corresponding solution in the absence of halide ion. The variations were those expected if variations of component intensities occurred. The spectra were analysed in terms of the same components as were found for the nitrate solutions with the intensity of all components allowed to vary and the band parameters of bands other than the dilute-solution band allowed to vary.The first study was carried out for mixtures of NaNO, and NaCl at constant ionic strength (total concentration “a+] of 5 mol drn-,) and this was followed by a series of studies in which [NO;] was kept constant at 1 mol dm-, and halide was added up to a total concentration ([M+]) of 6 mol drn-,. Ion association at a constant cation concentration was studied for mixtures of sodium nitrate and sodium chloride with “a+] of 5 mol dm-,. The results are collected in table 1 and fig. 1. The relative intensities of the three bands for concentrations of NaNO, from 1 to 5 rnol dm-, are shown in parentheses. It is clear that association between sodium and nitrate does not follow the concentration of nitrate but varies around the values for 5 mol dm-, sodium nitrate.Apparently the association equilibria are strongly controlled by the concentration of water in the solution, which at 5 mol dm-, salt is of the order of 9 molecules of water per molecule of salt. TABLE 1 .-RELATIVE INTEGRATED INTENSITIES OF COMPONENT BANDS FOR NaNOJNaCl AT 5 rnol dmP3 Na+ concentration/ mol dm-3 aquated-nitrate solvent-separated contact NO; c1- ion ion pair ion pair Kl K2 5 0 0.15 0.25 0.52 2.2 2.2 4 1 0.25 (0.25) 0.23 (0.34) 0.52 (0.38) 1.5 2.26 3 2 0.23 (0.32) 0.30 (0.40) 0.46 (0.25) 1.2 1.53 2 3 0.23 (0.38) 0.29 (0.48) 0.47 (0.13) 1.2 1.62 1 4 0.15 (0.42) 0.36 (0.58) 0.47 (0.01) 2.2 1.30R. L. FROST AND D. W. JAMES 0 325 1 F- 1 I 1 I I 1 1 L 1 I I 1 2 3 4 5 concentration/mol drn-3 FIG.1 .-Variation in component-band intensity with anion composition at constant Na+ concentration ( 5 mol dm-3). (a) Aquated-nitrate band; (b) solvent-separated-ion-pair band; (c) contact-ion-pair band. The changes in the concentration of the three species (as indicated by component- band intensities) reflect the influence of the chloride ion competing in the association processes. As the concentration of chloride increases the concentration of solvent- separated ion pairs increases while the concentration of contact ion pairs decreases. This is compatible with contact ion pairing being favoured for the chloride ion over the nitrate ion. In addition, if the nitrate ion is more strongly hydrated than the chloride ion in a mixed solution the solvent-separated ion pair will be favoured for the nitrate ion.TABLE 2.-RELATIVE BAND INTENSITIES FOR BAND COMPONENTS IN SYSTEMS MNO,/MX ~~ solvent-separated- aquated-ni tra te band ion-pair band contact-ion-pair band electrolyte, 4 shape 4 shape mj shape V V V concentration/mol dm-, /cm-' /cm-' ratio area /cm-l /cm-' ratio area /cm-' /cm-' ratio area e.m.s.LiNO,, 1 $0 LiC1,l .O 2.0 3.0 4.0 5.0 6.0 7.0 LiBr, 1 .O 2.0 3.0 4.0 5.0 LiI, I .O 2.0 3.0 4.0 5.0 NaNO,, 1 .O NaC1,O.O 1 .o 2.0 3.0 4.0 LiNO,, 1 .O LiNO,, 1 .O - - 1047.6 3.397 - - - - 1047.6 3.397 - - 0.644 0.636 0.352 0.197 0.150 0.1 15 0.036 0.723 0.475 0.422 0.214 0.086 0.760 0.651 0.566 0.403 0.055 0.389 0.384 0.356 0.278 0.156 4.97 4.94 4.87 5.07 4.88 5.02 5.49 4.86 5.07 5.28 4.77 4.86 5.12 5.07 5.19 4.97 5.05 3.55 3.43 3.28 3.67 3.69 0.756 0.750 0.718 0.712 0.658 0.734 0.621 0.713 0.756 0.813 0.695 0.690 0.73 1 0.726 0.705 0.696 0.675 0.90 0.814 0.755 0.729 0.728 0.345 0.354 0.643 0.663 0.619 0.572 0.574 0.247 0.518 0.553 0.547 0.612 0.220 0.321 0.412 0.460 0.570 0.595 0.439 0.430 0.435 0.496 - - - 8.95 8.90 7.70 7.48 - - 7.95 7.55 7.65 - - 7.85 7.75 7.75 8.1 1 5.00 5.93 5.29 4.87 - - - 0.647 0.647 0.606 0.526 - - 0.735 0.769 0.706 - - 0.696 0.695 0.690 0.857 0.908 0.893 0.903 0.841 0.00 1 0.001 0.001 0.135 0.232 0.293 0.390 0.001 0.00 1 0.024 0.226 0.367 0.00 1 0.001 0.056 0.141 0.375 0.001 0.167 0.200 0.282 0.36 1 1.38 1.30 8.30 0.362 0.69 0.366 2.02 1.429 1.291 6.244 0.016 0.964 1.64 2.75 0.81 1.78 1.63 0.08 0.06 0.04 0.01 0.13NaN0,,4.0 3.0 2.0 1 .o NaNO,, 1 .O KNO,, 1 .O KNO,,l .O KNO,, 1 .O KNO,, 1 .O NaC1,l .O 2.0 3.0 4.0 NaI, 1 .O 2.0 3.0 4.0 KF, 1 .O 2.0 3.0 4.0 KC1,l .O 2.0 3.0 4.0 KBr, 1 .O 2.0 3 .O 4.0 KI, 1 .O 2.0 3.0 4.0 5.0 1047.6 - 3.397 0.852 - - - 1047.6 - - - 3.397 0.852 - - 1047.6 - - 1047.6 - - - 3.397 0.852 - - 3.397 0.852 0.245 0.230 0.23 1 0.145 0.546 0.392 0.259 0.155 0.520 0.391 0.304 0.164 0.509 0.463 0.223 0.157 0.494 0.417 0.384 0.226 0.453 0.486 0.43 1 0.344 0.345 1049.1 3.55 3.56 - 3.56 - 3.56 1049.0 3.56 3.39 3.81 3.60 - - - - 0.90 0.90 0.90 0.90 0.902 0.782 0.719 0.769 0.234 0.307 0.296 0.361 0.420 0.416 0.540 0.48 1 1052.6 4.22 4.54 - 4.43 - 4.72 1052.5 3.26 4.67 5.12 - 4.79 1049.7 3.397 3.397 - 3.397 3.397 3.43 - 3.447 - 3.52 3.60 3.50 - 3.43 3.49 - 3.67 3.79 3.86 3.48 .- 3.61 3.78 - - - - - - - - - - - - __ 0.71 1 0.679 0.672 0.650 0.756 0.906 0.675 0.679 0.826 0.823 0.81 1 0.795 0.799 0.806 0.8 17 0.784 0.837 0.800 0.759 0.820 0.782 0.724 0.824 0.810 0.759 0.519 0.456 0.467 0.491 0.032 0.178 0.189 0.354 0.450 0.596 0.684 0.825 0.477 0.527 0.765 0.823 0.497 0.574 0.619 0.750 0.442 0.507 0.563 0.643 0.650 1.629 1.223 1.145 0.723 0.64 0.08 0.02 0.09 1.26 2.46 1.09 2.52 0.90 1.20 1.19 0.9 1 0.620 1.210 0.327 0.719 1.13 2.05 1.45 1.65 0.633254 I ON-I ON-SO L V EN T INTER ACT I 0 N S For solutions where the concentration of nitrate is kept constant and halide is progressively added, the results may be compared with nitrate solutions containing the same cation concentrations.Results for relative band intensities for baild components are collected in table 2 and these will be discussed according to cation.The results for solutions having potassium cations are summarised in fig. 2. Because of the limited solubility of potassium nitrate comparison of association behaviour in the presence and absence of halide is not possible over all concentrations. The addition of one mole of potassium halide does not change the relative concentration of free and associated nitrate ion. The increase of cation concentration produces considerable reduction in the calculated association quotient. However, there are competing association equilibria in the halide system and also a change in free water and so these association quotients have little absolute significance. The changes in relative intensities of band components do show systematic trends which can be interpreted in terms of solution interactions.In solutions containing fluoride ions the aquated nitrate is reduced relative to those containing iodide, with the tendency being directly related to anion size. A complementary trend is seen for the associated species. Association between the potassium ion and halide ion will be more likely with a larger anion and so the observed trend can be understood in terms of competition for cations between the two anionic species. Solutions of NaNO, (1 mol drn-,) with progressively increasing halide ion concen- tration have been analysed and the results are summarised in fig. 3. At 1 mol dm-, added halide the aquated nitrate is favoured over the associated species and as the concentration of halide increases the solvent-separated ion pairs becomes favoured at the expense of the aquated nitrate ion while the contact ion pair remains lower than for solutions containing only nitrate solute. The variation in the contact ion pairs shows anion dependence, the influence being greater for iodide than for chloride.This is in agreement with the observations in solutions of potassium salts and emphasises the competitive equilibria for contact ion pairs with the largest most polarizable anions preferentially forming the contact ion pairs. The variations seen for the aquated nitrate species and the solvent-separated ion pairs may be due to the disturbance of the water structure by the added halides, both of which can be considered structure breakers due to the large anion size.The disturbances to the water structure facilitate the formation of free aquated ions and as the solute concentration increases the concentration of solvent-separated ion pairs with the nitrate remains high because formation of contact ion pairs takes place preferentially to the halide anion. Similar but more pronounced variations are observed for the systems containing lithium cations, as illustrated in fig. 4. The concentration of the aquated nitrate ion is enhanced by the presence of halide and this enhancement is directly dependent on anion size, being most pronounced for the iodide ion. The concentrations of the solvent- separated ion pairs with nitrate are initially depressed but then are enhanced by the presence of halide ions, with the effect being inversely dependent on the size of the halide ion.Finally, the concentration of contact ion pairs is depressed by the presence of halide with the chloride producing the greatest change. These results may again be understood in terms of competing equilibria and disturbance of the water structure. When the network association between water molecules is disturbed by large anions the formation of aquated cations (and anions) will be enhanced. Thus in the presence of iodide ions there is a pronounced increase in free aquated nitrate ions which occurs at the expence of the formation of solvent-separated ion pairs. At higher concentrations where contact-ion-pair formation is expected such species are formed preferentially with the large polarizable anion and the formation of contact ion pairs with nitrate is depressed. This in turn causes an increase in the proportion of solvent-separated ion pairs, this increase being dependent on the size of the halide ion.R.L. FROST AND D. W. JAMES 3255 0.6 0.4 0.2 0.6 x u w? E e, .- + .s 0.4 aJ .- Y cd - 2 0.2 0.6 0 . 4 0.2 [Na+]/mol dn1-3 FIG. 3.-Variation in component-band with anion composition in the presence of sodium cations. ( a ) Aquated-nitrate band ; (b) solvent-separated-ion-pair band ; (c) contact-ion-pair band. (-) NaNO, ; (- - - -) 1 mol dm-3 NaNO, +x mol dm-3 NaX.3256 I 0 N-I 0 N-S 0 L V E N T INTER ACT I ON S O.( 0.1 0 . i 0 . € h c U .- U .5 0.4 Q) .- Y - m Q) 0.2 0.6 0 .I 0.2 / b l \ 4- -. Cl-,’,/-- ‘ / / /’// / / / A- / ( C l [ Li+l/mol dm-3 FIG.4.-Variation in component-band intensity with anion composition in the presence of lithium cations. (a) Aquated-nitrate band; (b) solvent-separated-ion-pair band; (c) contact-ion-pair band. (-) LiNO, ; (----) 1 mol dm-3 LiN03+x mol dmP3 LiX.R. L. FROST AND D. W. JAMES 3257 The bandwidths and band shapes give evidence of increased perturbation caused by the presence of added halide. The bandwidth is broader and the band shows greater Gaussian character in the presence of halide for both the solvent-separated- ion-pair band and the contact-ion-pair band. The increased width in the solvent- separated-ion-pair band indicates a greater symmetry perturbation leading to more rapid dephasing. This may come from a change in the relative orientation of the anion and cation or possibly from a long-range environmental disturbance due to the additional ionic species.The increased width for the contact-ion-pair band corresponds to a reduction in vibrational relaxation from ca. 1 to 0.65 ps. This increased width is best attributed to a long-range environmental perturbation and the associated increase in Gaussian character is compatible with a change to a more rigid structural character in the solution. This would agree with the greater population of more strongly bound ion pairs due to the presence of halide ion. INFLUENCE OF TEMPERATURE ON THE ASSOCIATION EQUILIBRIA It has been suggestedg that the influence of temperature on the concentration of free and bound species may be quite small.Two studies agree that the formation of inner-sphere complexes ZnNOl is favoured by an increase in temperature,l0? l1 and in solutions of MgSO,, for the formation of contact ion pairs, a AH of ca. 12 kJ mol-l was obtained.12 A study of so1utions13 of Ca(NO,), led to a value of AH = 1.1 kJ mol-1 for a complex which is probably an inner-sphere complex. A solution of sodium nitrate (2 mol dmP3) was studied over the temperature range - 5 to 65 "C and the resulting profiles for the band due to the symmetric stretching vibration were component-band analysed. Three band components were required to describes the profile and these had temperature-independent shape and half-width. The variation in intensity of these three components is shown in fig. 5, from which it is clear that both the solvent-separated species and the contact pair are favoured as the temperature decreases.Although the equilibrium constants calculated from the intensities of band com- ponents are not related to any standard state, estimates of the enthalpy and entropy changes for the equilibria may be obtained from the representation in fig. 6. The 0. .- c 0 . v1 K 2 0 . .- 0 . 0 . 0 . 0 . I 0 10 20 30 40 50 60 temperature/OC' FIG. 5.--Variation in component-band intensity with temperature for 2 mol dm-3 aqueous NaNo,. (a) Aquated-nitrate band ; (h) solvent-separated-ion-pair band; (c) contact-ion-pair band.3258 4 3 2 1 o n 0 1 * - 2 3 4 ION-I ON-SO LVEN T I N TE R A C T I ON S FIG. 6.-Variation of K , with 1/T for the results in fig. 5. obtained values of AH, = -42.2 kJ mol-1 and AH, = - 64.6 kJ mol-l are remarkably larger than expected on the basis of previous studies.l27 l3 The values for the entropy changes of ASl = - 142 J K-l mol-1 and AS2 = - 294 J K-l mol-l indicate that there is a pronounced decrease in entropy on formation of associated species.The results for AH and A S are large in comparison with previous estimates but the change in free energy for the association equilibrium is still quite small because of the balance of energy and entropy terms. Thus at 296 K the value for AG is zero; at higher temperatures AG is positive and the aquated nitrate is favoured while at lower temperatures AG is negative and the solvent-separated ion pair is favoured. The association process involves changes to the total aquation energies which are relatively small (if changes in water-water interactions are excluded).These changes involve the disruption of the solvent sphere about both anion and cation due initially to the disturbance resulting from sharing of a water molecule and secondly to the removal of a solvent to allow direct anion-cation interaction. The major energy changes can be associated with the coulombic interaction between the two ions. This interaction can be visualised as the approach of an anion and a cation to an initial stable separation corresponding to the solvent-separated ion pair and then to a separation corresponding to the contact ion pair. Calculation of such ion-pair binding energies is rendered very difficult because of the planar nature of the nitrate ion and the uncertainty about the stereochemistry of the approach. Assuming an equatorial approach of the cation along a nitrogen-oxygen bond axis the expected Na-N distance would be ca.350 pm.14 If the ions are given central point charges the ion-pair binding energy would be ca. 400 kJ mol-l. The combination of AH, and AH, of - 106 kJ mol-l is much less than this but the relativity of values indicates that the large values of AHl and AH, obtained are not unreasonable on the basis of the expected ion-pair interaction energies.R. L. FROST AND D. W. JAhlES 3259 1 . 0 0.8 0.6 0 . 4 0 . 2 0 . c 0.8 A 4- .& g O * € W c W c) ..-, .- % 0 . 4 - 2 0.2 0 . c 0.8 0.6 0 . 1 0 . 2 0 A '\ \ B / C 1 2 I 6 8 cation concentration/mol dm-3 FIG. 7.-Comparison of component-band intensities in solutions in H,O (---) and D,O (-).A, KNO,; B, NaNO,; C, LiNO,; (a) aquated-nitrate band; (b) solvent-separated-ion-pair band; ( c ) contact-ion-pair band; (d) ion-aggregate band; (e) aquated-ion-pair band. Band parameters for the various bands with the half-width (cm-I) and shape ratio. LiNO,/D,O: band 1 , 3.40, 0.85; band 2, 5.1, 0.84; band 3, 5.6, 0.79; band 4, 6.8, 0.76. NaNO,/D,O: band 1, 3.40, 0.85; band 2, 4.0, 0.90; band 3, 4.5, 0.70; band 4, 9.5, 0.68. KNO,/D,O: band 1, 3.40, 0.84; band 2, 3.5, 0.90; band 3, 6.7, 0.76.3260 I 0 N-ION-SO LVE N T INTER ACT I 0 N S ASSOCIATION I N D,O AS SOLVENT The difference between hydrogen bonding and solvent properties of D20 and H,O is undoubtedly small and in most instances the two solventscan be used interchangeably.Estimates of the difference in hydrogen-bond strength between the two have been inconcl~sive.~ While the time correlation functions for the v1 band of NO, in solutions of LiNO,, NaNO, and KNO, are significantly different in H,O and D20, the theta functions in the two solvents are very similar. The energy band profiles in D20 were analysed assuming that they contained the same number of bands as the aqueous- TABLE 3.-EQUILIBRIUM QUOTIENTS OF GROUP I NITRATES IN H 2 0 AND D20 electrolyte concentration/ mol dm-3 K K2 LiNO, in H,O LiNO, in D20 NaNO, in H 2 0 NaNO, in D,O KNO, in H,O KNO, in D,O RbNO, in H,O 0.5 1 .o 2.0 4.0 6.0 8.0 10.0 0.5 1 .o 2.0 4.0 6.0 8.0 10.0 0.5 1 .o 2.0 4.0 6.0 8.0 0.5 1 .o 2.0 4.0 6.0 0.5 1.0 2.0 3.0 0.5 1 .o 2.0 3.0 1 .o 2.0 3.0 1.79 1.34 0.586 0.862 0.552 2.200 0.926 0.605 2.465 3.81 1 6.6 53.89 115.9 2.74 3.28 I .69 1.34 3.50 8.40 1.5 1.37 0.729 1.38 3.74 11.22 - - - - - - - - - - - 0.0026 0.9 14 0.49 1 5 1 .oo 1.49 1.78 2.776 - - - 0.120 0.500 3.53 2.86 - - 0.269 1.122 3.26 4.35 - - 0.1178 0.843 1.06 1.046 1.95 1.85 1.88 0.440 0.430 0.639 0.582 1.0375 0.9457 0.8205R.L. FROST AND D. W. JAMES 326 I solution bands. In addition, the parameters from the aqueous-solution bands were used as starting parameters in the analysis. The results of the analysis are summarised in fig. 7. The band parameters for the solutions of NaNO, and KNO, are similar in H,O and D,O. However, parameters of band components for solutions of LiNO, in D20 are significantly different from those in H,O.For both the solvent-separated-ion-pair band and the contact-ion-pair band the parameters indicate a more Gaussian shape and a greater half-width in D,O solution. This implies that the perturbation of the nitrate ion is greater in the presence of D,O, which could be due to the influence of the Li+ ion being less ‘moderated’ by the D,O. The component-band intensities (fig. 7) are appreciably different in D,O from those observed in H,O and this difference is emphasised in the equilibrium quotients listed in tableg. For both NaNO, and KNO, solutes the values of the association quotients in D20 are much less than the values in H,O. If the hydrogen bonding in D,O was significantly weaker than in H,O aquation of the cations could be expected to disturb the structure more and this, together with the weaker hydrogen bonding to the nitrate ion, will destabilise the solvent-separated ion pair and will allow the entropic driving force to favour the free aquated ions.The differences in association quotients for LiNO, solute show more complex variation. In dilute solution in D20 the values are much smaller than in H,O, showing that the free aquated ions are again favoured. However, above 2 mol dm-, the association species become more favoured than in water. The formation of associated species in water was influenced by the strong hydration of the lithium ion which moderated the strong polarizing power of the cation. If aquation by D,O is less strong than that by H,O, then this moderating influence would be reduced and the influence of the small cation would be evident.In this regard the k , values for 2 mol dm-, solutions of LiNO,, NaNO, and KNO, in D,O follow the inverse of the order of cation size, i.e. follow the cation polarizing power. The changes in association quotients are all compatible with hydrogen bonding and aquation by D,O being weaker than that by H,O. Note that because of the small values for AG reported in the previous section the changes in bonding and aquation forces could be quite small to produce the changes observed. We thank the Australian Research Grants Committee for grants enabling the purchase and maintenance of the Raman spectrometer. D. E. Irish and M. H. Brooker, in Advances in Infrared and Raman Spectroscopy, ed. R. J. H. Clark and R. E. Hester (Heyden, London, 1976), vol. 2, p.212. J. Nixon and R. A. Plane, J . Am. Chem. SOC., 1962, 84, 4445. J. T. Bulmer, T. G. Chang, P. J. Gleeson and D. E. Irish, J . Solution Chem., 1975, 4, 969. J . Bjerrum, Trans. R. Inst. Technol., Stockholm, 1972, 69, 248. G. E. Walrafen, in Water: A Comprehensive Treatise, ed. F. Franks (Plenum Press, London, 1972), vol. I , p. 151. P. G. Wolynes, Annu. Rev. Phys. Chem., 1980, 31, 345. R. L. Frost and D. W. James, J. Chem. Soc., Faraday Trans. I , 1982, in press. M. H. Brooker, J . Chem., SOC., Faraday Trans. I , 1975, 71, 647. I 1 A. T. G. Lemley and R. A. Plane, J. Chem. Phys., 1972, 57, 1648. l 3 R. E. Hester and R. A. Plane, J . Chem. Phys., 1964, 40, 41 1 . l4 D. W. James, Aust. J . Chem., 1966, 19, 993. ’ B. E. Conway, Ionic Hydration in Chemistry and Biophysics (Elsevier, New York, 1981), p. 552. l o S. A. Al-Baldawi, M. H. Brooker, T. E. Gough and D. E. Irish 1982, 78, 3235. 1202. R. M. Chatterjee, W. A. Adams and A. R. Davis, J. Phys. Chem., 1974, 78, 264. (PAPER 1/2001)

ISSN:0300-9599    

DOI:10.1039/F19827803249

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 3263-3279 Ion-Ion-Solvent Interactions in Solution Part 6.-Aqueous Solutions of Metal Nitrates having Cations with Incomplete Valence Shells BY RAY L. FROST Chemistry Department, Queensland Institute of Technology, Brisbane, Australia 400 1 AND DAVID W. JAMES* Chemistry Department, University of Queensland, Brisbane, Australia 4067 Received 2 1st December, 198 1 The Raman spectra of aqueous solutions of AgNO,, TlNO,, Zn(NO,),, Cd(NO,),, La(NO,), and Th(NO,), are analysed as a function of concentration. For all solutes, bands due to aquated-nitrate ions, solvent-separated ion pairs (outer-sphere complexes) and contact ion pairs (inner-sphere complexes) are reported. The nature of the inner-sphere complex leads to a distinctly different band from contact ion pairs for simple monovalent cations.This difference is interpreted in terms of an electron-exchange mechanism. The inner-sphere complexes are discernable at all concentrations studied. The association quotients compare favourably with values from other measurements for solutions of Zn(NO,), and Th(NO,),. Recent studies1 of the vibrational spectra of aqueous zinc nitrate solutions resolved the Raman contour of the A ; symmetrical stretching vibrational band into two bands at 1050.0 and 1035.0 cm-l. The contribution of the 1035.0 cm-l band was found to be 7.3% at 3 mol dm-, Zn(NO,),. In the infrared spectrum, a band at 1047.0 cm-l was observed even at low concentrations (0.89 mol drn-,) as well as the 1050 cm-l band. This 1047.0 cm-l band is infrared forbidden for the unperturbed nitrate ion.Both the Raman and infrared bands showed asymmetry on the low-frequency side over all concentrations, confirming earlier studies with reported low-frequency asymmetry of the 1050.0 cm-l band.2 The cation dependence of the v, region from concentrated aqueous zinc nitrate solutions showed that specific inner-sphere com- plexes or contact ion pairs were not formed’ and this result was substantiated by the lack of appreciable intensity in the 740.0 cm-l region. Hence, it was suggested that discrete solvent-separated ion pairs are formed which exist for long times compared with the vibrational lifetime. If solvent-separated ion pairs are formed then the v, region should be a quartet of bands. Lemley and Plane2 curve resolved their zinc nitrate spectra and found stability constants for the outer-sphere complex Zn(H,O),NO;.An alternative explanation for the splitting of the v, regions is in terms of an asymmetric field splitting model in which the separation of the v, components is merely due to a gradual increase in the field strength of the cation acting on the solvated nitrate group.3 If this second explanation is correct, then outer-sphere complexes might be formed and broken on a time scale too rapid for measurement on the vibrational lifetime scale. Observable splittings of the v1 region however tend to negate such a consideration. No evidence for nitrato complexes was reported as no zinc-oxygen stretching vibration was observed. This suggests that interactions between species of the inner-sphere complex are coulombic rather than covalent.32633264 I 0 N-I 0 N-SO L V E N T INTER ACT I 0 N S The appearance of a 740 cm-l component in the in-plane bending vibrational (v4) band profile of cadmium nitrate solutions indicated that nitrate-cadmium contact species exist . 4 t No outer-sphere or solvent-shared complexes were reported. Based on the curve analysis of the v, region, equilibrium quotients of 0.38 were obtained for the formation of the cadmium nitrato species. Two bands were reported for the v1(A;) spectral region: the 1050 cm-l band assigned to the aquated-nitrate band and the 1043.0 cm-l band which appeared as an asymmetry on the 1050.0 cm-l band. No evidence for outer-sphere complexes of cadmium has been forthcoming.Spectral observations of silver nitrate showed splirting the v, band and reported the v4 region as a single symmetric band.6 A more recent study showed the v,(A;) band profile for silver nitrate, which had long been regarded as a single symmetric band, with peak positions and half-widths which were concentration-dependent.' The peak position and half-width for a 0.5 mol dm-, AgNO, solution were 1047.8 and 8.7 cm-l and for a 9.0 mol dm-, solution the values were 1043.8 and 11.0 cm-l. The v,(A;) spectral band profile of AgNO, showed low-frequency asymmetry and two bands were resolved at 1047.0 and 1041.0 cm-1.6 Infrared spectra showed that the 1041.0 cm-l region was also asymmetric. The conclusion was drawn that only two forms of nitrate in aqueous silver nitrate existed; the free nitrate and the bound nitrate in the form of a contact ion pair.Again no outer-sphere or solvent-separated silver ions were found. The transition to inner-sphere complexes has often been assessed in terms of a second v4 band at 740 cm-l. Whilst new bands have been found in silver nitrate at 750 and 760 cm-l, similar bands did not show appreciable intensity in zinc nitrate and cadmium nitrate solutions. Stability or association constants have been calculated on the assumption that only two bands are present which may be assigned to two species in solution, an inner-sphere or outer-sphere complex and the aquated nitrate ion.6 We report studies of aqueous solutions of the salts TlNO,, AgNO,, Zn(NO,),, Cd(NO,),, La(NO,), and Th(NO,),. On the basis of evidence from the time correlation function and theta function the band due to the symmetric stretching vibration has been component-band analysed.The component bands are interpreted in terms of associated species, and the association equilibria are compared with earlier studies. EXPERIMENTAL The salts were all recrystallised twice from water, oven dried at 110 O C for 48 h and stored over P,O,. Solutions of ThIV were made up at ca. pH 2 to suppress hydrolysis reactions. The remaining experimental and analytical details were the same as previously described. RESULTS The characteristics of the isotropic and anisotropic band profiles together with some calculated parameters are collected in table 1. These results will mainly be discussed in connection with individual solutes.The modulation parameter rC(M2)~/2nc indicates that all of these solutes are characterised by intermediate modulation which is similar to the observation in Group I nitrate solutions. The results of component-band analysis are presented in table 2. The bases for these analyses will be discussed for each solute below. SOLUTIONS OF AgNO, AND TINO, The peak position of the isotropic band in AgNO, solutions shows movement to lower energies with increased concentration (table 1). The second moment of theTABLE 1 .-BAND PARAMETERS, MODULATION TIMES, AND RELAXATION TIMES OF METAL NITRATES WITH UNFILLED VALENCE SHELLS M2,a M2.t 7c ~"(4) Tor(@) ~"(e-') T0Ae-l) concentration va v8 electrolyte /mol dm-3 /cm-' /cm-' /(cm-1)2 /(cm- ) (Cb/a) 100 /ps tCd2/2nc /ps /ps /ps /PS TINO, 0.2 0.4 AgNO, 0.5 1 .o 2.0 4.0 6.0 Zn(No 3) 2 0.5 1 .o 2.0 3.0 Cd"O312 0.5 1 .o 2.0 3.0 4.0 La(NO,)3 0.5 1 .o 2.0 3.0 Ce"O3)3 0.5 1 .o 2.0 3.0 TWO,), 0.25 1 .oo 2.00 2.50 1047.1 1047.2 1047.2 1046.7 1046.4 1045.6 1044.5 1047.9 1048.2 1048.9 1049.7 1048.2 1048.3 1048.8 1048.9 1048.9 1048.3 1048.6 1048.7 1049.0 1048.4 1048.6 1048.8 1049.4 1048.5 1048.8 1049.6 1049.6 1047.4 1047.5 1047.2 1046.9 1046.5 1045.5 1044.0 1048.5 1049.0 1049.0 1050.1 1048.0 1048.0 1048.5 1048.5 1048.3 1048.3 1048.4 1048.6 1048.0 1048.3 1048.5 1049.0 1049.8 1048.8 1049.3 1049.4 1049.8 56.0 56.0 56.0 59.0 64.0 64.0 64.0 72.0 81.0 81.0 81.0 64.0 56.0 64.0 72.0 81.0 64.0 64.0 72.4 72.4 56.0 64.0 64.0 72.0 56.0 64.0 64.0 72.0 128.0 128.0 128.0 168.0 182.0 196.0 256.0 158.0 2 16.0 225.6 306.0 289.0 196.0 256.0 256.0 290.0 256.0 256.0 256.0 256.0 138.0 156.0 240.0 256.0 144.0 156.0 196.0 256.0 - 1.82 -0.59 - 0.52 + 0.48 + 1.48 + 2.20 + 4.58 + 1.76 - 0.44 - 0.09 - 9.69 - 3.7 - 3.48 - 5.48 -6.01 + 1.149 - 2.26 -4.70 - 5.28 - 4.69 - 2.1 - 1.9 - 3.4 - 1.9 - 2.50 - 2.58 - 7.20 - 3.79 1.05 0.99 1.15 1.20 1.18 1.24 1.23 + 0.734 + 0.685 + 0.786 + 0.997 0.75 1 .oo 0.96 0.96 0.943 0.78 0.83 0.85 0.89 0.875 0.820 0.970 0.970 0.90 0.86 1.1 1 1.08 1.02 1.06 1.16 1.23 1.25 1.24 1.23 0.831 0.821 0.940 1.190 0.80 0.99 1.02 1.08 1.13 0.83 0.88 0.96 1 .oo 0.870 0.870 1.03 1.09 0.90 0.91 1.19 1.23 1.36 1.34 1.19 1.1 1 1.08 1.06 1.04 1.33 1.25 1.13 1.03 1.34 1.24 1.07 0.94 0.86 1.34 1.17 0.90 0.74 1.29 1.10 0.87 0.73 1.29 1.09 0.87 0.77 1,83 1.86 2.86 3.12 3.32 3.66 4.08 1.55 1.93 2.53 3.54 1.66 1.88 2.59 2.87 3.62 1.69 1.86 2.59 3.66 1.56 1.80 2.41 3.3 1 1.58 2.02 3.79 7.90 0.95 1 .oo 0.85 0.79 0.75 0.75 0.72 1.063 1.016 0.885 0.698 1.18 1 .oo 0.92 0.8 1 0.74 1.13 1.06 0.92 0.88 1.15 1.08 0.91 0.8 1 1.11 1.03 0.79 0.72 4.87 4.87 1.59 1.40 1.40 1.40 1.31 2.31 2.65 4.41 10.62 1.20 1.54 1.78 2.35 3.45 1.20 1.30 1.59 5.30 1.94 2.72 3.36 4.90 1.73 1.98 4.82 10.65TABLE 2.-RESULTS OF COMPONENT-BAND ANALYSES OF THE V,(A;) BAND PROFILE OF AQUEOUS NITRATE SOLUTIONS OF METALS WITH UNFILLED VALENCE SHELLS solvent-separa ted- aquated-nitrate band ion-pair band concentration V wi shape V 4 shape electrolyte /mol dme3 /cm-' /cm-' ratio area /cm-' /cm-' ratio area TlNO, 0.2 0.4 1 .o 2.0 4.0 6.0 Zn (NO,), 0.5 1 .o 2.0 3 .O 0.5 1 .o 2.0 3 .O 4.0 3 .O 4.0 0.5 1 .o 1.5 2.0 2.5 3.0 TWO313 0.25 0.50 1 .o 1.5 2.0 2.5 0.752 1050.0 0.710 - 0.531 1050.0 0.335 - 0.187 - 0.001 - 0.570 1051.0 0.498 - 0.233 - 0.10 0.607 1050.5 0.445 - 0.297 - 0.208 - 0.150 - 0.155 - 0.1155 - 0.571 1050.5 0.373 - 0.230 - 0.126 - 0.090 - 0.0655 - 0.493 1050.0 0.286 1049.7 0.245 1050.0 0.147 1050.5 0.083 1051.0 0.054 1051.0 - 3.55 3.50 3.47 3.47 3.21 3.47 3.78 3.84 4.50 5.06 3.59 3.59 4.08 4.35 4.66 - - 3.8 3.8 4.37 5.00 5.15 5.25 3.92 4.06 4.31 4.68 5.10 5.20 0.93 0.96 0.905 0.905 0.910 0.904 0.766 0.794 0.744 0.753 0.901 0.901 0.870 0.814 0.808 - - 0.90 0.90 0.865 0.820 0.810 0.802 0.889 0.827 0.799 0.753 0.763 0.789 0.1089 0.085 0.19 0.133 0.0162 0.00 1 0.306 0.343 0.625 0.680 0.305 0.445 0.498 0.469 0.476 0.581 0.500 0283 0.399 0.512 0.556 0.519 0.532 0.392 0.490 0.510 0.513 0.492 0.459 w h, rn rn 7 v, 0 rTABLE 2 ( c m .) second directional-nitrate band directional-nitrate band concentration V shape V 4 shape electrolyte /mg dm-, /cm-* /cm-l ratio area /cm-' /cm-l ratio area e.m.s. TlNO, 0.2 &NO, 1 .o 0.4 2.0 4.0 6.0 Zn(NO,), 0.5 1 .o 2.0 3.0 Cd(N03)2 0.5 1 .o 2.0 3.0 4.0 3.0 4.0 WNO,), 0.5 1 .o 1.5 2.0 2.5 3.0 TWO,), 0.25 0.50 1 .o 1.5 2.0 2.5 3.55 3.55 4.30 4.30 3.76 4.39 3.94 3.46 3.28 6.8 4.74 5.09 5.14 6.06 6.01 - - 5.74 5.74 6.30 6.40 6.50 6.95 4.22 7.88 7.26 7.93 7.60 7.49 0.90 0.90 0.870 0.870 0.901 0.854 0.926 0.845 0.892 0.844 - - - - - - - 0.895 0.895 0.732 0.623 0.623 0.702 0.692 0.803 0.870 0.814 0.756 0.739 0.135 0.196 0.265 0.509 0.694 0.793 0.03 0.079 0.1 11 0.254 0.077 0.176 0.193 0.308 0.353 0.136 0.257 0.136 0.227 0.255 0.302 0.33 0.396 0.056 0.172 0.176 0.242 0.265 0.321 - - 1039.5 - - - 1070.0 - - - - - - - - 1038.5 - - - - - - - - 1032.5 1032.5 1032.5 1032.5 1032.9 - - 3.66 3.66 4.15 3.89 6.06 13.2 10.9 16.6 - - - - - - - - - - - - - - 5.96 5.97 5.99 6.89 6.86 - - 0.757 0.766 0.727 0.743 0.749 0.531 0.783 0.645 - - - - - - - - - - - - - - 0.575 0.576 0.577 0.666 0.663 - - 0.001 0.018 0.086 0.132 0.012 0.027 0.030 0.1 19 - - - - - 0.118 0.1 14 - - - - - - - 0.05 0.068 0.097 0.158 0.166 8.0 16.0 1.10 1.74 13.85 12.31 0.06 0.22 0.64 3.52 - ___ - - - - - 9.0 6.6 0.89 0.57 0.57 13.0 - - - - - -3268 ION-I ON-SO LV EN T INTER ACT I ON S isotropic band varies from 56.0 (cm-1)2 at 0.5 mol dm-3 to 64.0 (cm-1)2 at 6.0 mol drn-,.The second moment of the anisotropic experimental profile shows a very broad band: at 6 mol dm-!j AgNO,, M2(P) is 256 (cm-1)2. The asymmetry index of -0.52 shows that at 0.5 mol dm-, the band is asymmetric on the low-energy side and at 6.0 mol dm-, the value is + 4.58, showing asymmetry on the higher-energy side. Such features are indicative of a system of a number of overlapping bands which alter in intensity at different concentrations. The correlation function of silver nitrate solutions show more curvature than for most nitrate solutions. The anisotropic correlation functions indicate a minimum at 3-3.5 ps, which suggests a peak separation between two major bands of ca. 5.1 cm-l. Theta functions which are complex with a large negative slope show oscillations with nodes occurring every ca.1.6 ps for 4 mol dm-, AgNO,, suggesting a separation between the major components of ca. 8.0 cm-l. Curve resolution of silver nitrate profiles resulted in the characterisation of four bands positioned at 1047.6, 1050.0, 1044.5 and 1039.5cm-'. These bands are attributed to the aquated nitrate ion, the solvent-separated ion pair, the inner-sphere complex of a mononitrato type and a second inner-sphere complex. The position of the inner-sphere complexes is in good agreement with the single band assigned at 1041.0 cm-l by Irish and Chang7 Curve resolution was attempted in terms of a single band at ca. 1044.5 cm-l but this caused a residue at ca. 1038.0 cm-l and a further band was added to take this into account.The dominant feature of the curve analysis of silver nitrate profiles is the intensity of the 1044.5 cm-l band, which increases with increasing concentration and grows at the expense of both the aquated-nitrate-ion band and the solvent-separated-ion-pair band, such that at 6.0 mol dm-3 no intensity is left in these bands, as shown in fig. 1. The intensity of 0. [ AgNO, ] /mol dm-3 FIG. 1 .-Variation in intensity of the component bands of the A ; symmetrical stretching vibrational band for aqueous solutions of AgNO, : (a) aquated-nitrate band; (b) solvent-separated-ion-pair band; (c) directional-nitrate band. the solvent-separated-ion-pair band becomes negligible after 2 mol drn-,. The vibrational relaxation times of the aquated-nitrate-ion band the solvent-separated- ion-pair band are 1.56 and 1.51 ps, in contrast to the value for the directional-nitrate band at 1044.5 cm-l, which is 1.23 ps.The difference is due to the more efficient relaxation mechanism brought about by the symmetry perturbation of the nitrate. At high concentrations (> 6 mol dm-3), there is appreciable intensity in the region of 1039.5 cm-l. The high molar intensity observed for silver nitrate solution^,^ in contrast to nitrate solutions of Groups I and 11, is indicative of increased polarization of theR. L, FROST AND D. W. JAMES 3269 nitrate ion. Studies of silver nitrate melts have suggested that electron donation occurs between the anion and cation.s* The low energies of peak positions of the inner-sphere complexes may occur because of this electron-exchange mechanism and this will be further examined shortly.The equilibrium quotients calculated in this study assume that the molar intensity for the aquated nitrate ion and the associated nitrate species are all the same. Previous studies of ionic melts and aqueous solutions have indicated that the intensity of the v,(A;) mode of the nitrate ion is sensitive to polarization by the ~ a t i o n . ~ ~ lo It is then likely that the intensities will vary in the order: aquated nitrate ion < solvent-separated ion pair < contact ion pair. The differences are likely to be greatest between the aquated nitrate ion and the solvent-separated ion pair. There is at present no basis for suggesting what the real molar intensities are and we are presently addressing this problem.The equilibrium quotients calculated on the assumption that the molar intensities of all nitrate species are the same are likely to be too large, with Kl being the most influenced. In the sequential equilibria which are utilised in this study the solvent-separated ion pair is considered separately from contact associated species. It is likely that measurements made in dilute solution using transport measurements also examine this equilibrium. Other measurements such as ion-exchange measurements cannot distinguish between a contact ion pair and a solvent-separated ion pair and so the equilibrium constants may involve a combination of the two equilibria. The equilibrium between solvent-separated ion pairs and contact ion pairs has not been studied previously.In most of the equilibria the water is involved in the equilibrium and because of the concentrations involved the concentration of water should be included. Because this has not been done the equilibrium quotients are of qualitative significance only. l2 of 0.9 dm3 mol-l is probably most closely related to our Kl value of 0.67 dm3 mol-l. A previous value of 0.08 estimated from a Raman spectroscopic study appears to underestimate the intensity of the band at 1042 cm-l. A value of 0.22 dm3 mol-l has been reported’, also for an equilibrium corresponding to Kl. For thallium(1) nitrate the Kl value is less than and the K4 value is greater than those for silver nitrate. A value for thallium nitrate of 0.3 dm3 mol-l has been reported13 for thallium nitrate solutions and although this is not very different from Kl at 0.4 mol dm-, it probably relates to a combination of Kl and K4.Whilst thallium nitrate has been studied in melts and aqueous melts, no detailed study of aqueous thallium nitrate has been published, probably because of its low solubility (0.4 rnol drn-,).l4 The peak position for 0.2 mol dm-, TlNO, at 1047.1 cm-l is slightly less than the nitrate dilute solution band at 1047.6 cm-l and the asymmetry index of - 1.82 compared with -0.47 for NH4N0, shows that the band profile is asymmetric on the lower-energy side. The second moment value of 56 (cm-1)2 is typical of dilute nitrate solutions and the vibrational relaxation measured from the long-time slope of the correlation function is 0.95 ps.The theta function displays slight curvature over the 6 ps time-scale and is typical of low-concentration theta functions. Curve resolution of thallium nitrate profiles yielded the results listed in table 2. The dominant band at this concentration as expected is the aquated-nitrate-ion band with a relative intensity of 0.75. However, significant contributions of bands attributed to the solvent-separated ion pair and directional nitrate still persist at 0.2 mol drn-,. The directional-nitrate band and solvent-separated-ion-pair band have relaxation times of 1.49 ps, which is slightly shorter than the value of 1.56 ps for the band attributed to the aquated nitrate ion. The value reflects a more efficient vibrational relaxation for the cation perturbed nitrate even at these low concentrations.A value of an equilibrium constant calculated from conductance3270 I 0 N-I 0 N-S 0 L VE N T INTER ACT I 0 N S SOLUTIONS OF ZINC NITRATE The correlation functions for aqueous solutions of Zn(NO,), show increasing slope with concentration. In addition the functions for both the isotropic and anisotropic bands show an interference in the time region 4-4.7 ps, suggesting a peak separation between components of 3.6-4.2 cm-l. Even though correlation functions of the anisotropic spectra are affected by low signal-to-noise ratios, a second interference appears at ca. 3 ps, which infers a peak separation of ca. 5 cm-l. The theta functions shown in fig. 2 yield more information. The slope of the theta function for 0.4 mol 0.0 3.2 tips 6 .FIG. 2.-Theta functions for zinc nitrate solutions. Concentration/mol dm-3: (a) 0.5; (6) 1 ; (c) 2; (d) 2.5; (el 3. dmP3 Zn(NO,), solution is - 0.245 rad ps-l and relative to the origin of transformation (1050.44 cm-l) the value indicates a phase shift of I .30 cm-l and a peak maximum of 1047.9 cm-l. The 2 mol dmP3 zinc nitrate solution theta function exhibits a node at ca. 5.1 ps and this suggests a spatial separation of ca. 3 cm-l between the major components. The 3 mol dm-, solution shows two nodes at 3 and 5 ps and such features typify a theta function composed of more than two bands. Component-band analysis of 0.5 mol dm-3 zinc nitrate solution was carried out in terms of two bands centred on the peak positions of 1047.6 and 1051.0 cm-l, which are assigned to the aquated nitrate ion and the solvent-separated ion pair.This analysis showed no appreciable intensity in the 1052-1056cm-1 region but residuals occurred at ca. 1045.0 and 1070.0 cm-l. The observed band profiles of zinc nitrate are not simple and the curve resolution was carried out on a basis of four bands centred on 1047.7, 1051 .O, 1045.0 and 1070.0 cm-l, with the results listed in table 2. The four bands are attributed to the aquated nitrate ion, the solvent-separated ion pair or outer-sphere complex, the ' directional ' nitrate or inner-sphere complex and an ion aggregate. In the Raman spectrum of hexa-aqua zinc nitrate crystal, the peak frequency of v,(A;) was 1057.0 cm-1 and on dehydration of the crystal a new peak at 1047.0 cm-l was observed in addition to the 1057.0 cm-l band, and on further dehydration increased intensity at 1045.0 and 1047.0 cm-l was also found.15 It wasR.L. FROST AND D. W. JAMES 327 1 proposed that the new line at 1047.0cm-l was due to the nitrate entering the first coordination sphere of zinc. Hester and Scaife in a study of variably hydrated zinc nitrate16 also found the symmetric stretching frequency at 1048 and 1045 cm-l. The variation in intensity of the bands was accounted for in terms of an equilibrium between species in the molten hydrate. The variation of intensity with concentration for the four components is shown in fig. 3. [Zn(NO,),] /mol dm-3 FIG. 3.-Variation in intensity of the component bands of the A ; symmetrical stretching vibrational band for aqueous solutions of Zn(NO,),.(a) Aquated-nitrate band; (b) solvent-separated-ion-pair band; (c) directional-contact-ion-pair band ; ( d ) ion-aggregate band. The predominant band throughout is the solvent-separated-ion-pair band, which has appreciable intensity over a wide concentration range. Both this band and the 1045.0 cm-l band grow at the expense of the 1047.6 cm-l band. The 1045.0 cm-l band is attributed to an inner-sphere complex in which the cation is bound directionally to the nitrate ion. This may be contrasted with the contact ion pair formed between lithium and nitrate, where the interaction is coulombic and non-directional. Appreciable intensity (0.1 19) of the 1070.0 cm-l band is shown at a concentration of 3 mol dm-3. This band, which is broad, and the band assigned at 748.8 cm-l (the v4 region), which is also broad and weak, may be attributed in total or in part to the ion-pair aggregate.2 Of all the band profiles of nitrate solutions studied, zinc nitrate proved the most difficult to curve analyse.The band assigned to the directional nitrate at 1045.0 cm-l may bedecomposed into two bands at 1045.0and 1042.5 cm-l, with slight improvement to the e.m.s. value. In a previous study this region was analysed into two bands at 1050.0 and 1035.0cm-l with half-widths of 15.0 and 8.0 cm-l. In contrast, the half-widths determined in this work of the inner-sphere-complex band and the solvent-separated-ion-pair band were found to be 6.8 and 5.0 cm-l. The intensity we observe at 1035.0 cm-l is small and less than one percent of the total.Such a band is better assigned to isotope bands of the nitrate ion.3272 ION-I ON-SO L VEN T INTER ACT I 0 N S Equilibria in aqueous zinc nitrate solutions may be expressed as: K , K4 Zn(H20),2+ + NO;(aq) f Zn(H,O),NO;(aq) Zn NO,(H,O),+(aq) aquated nitrate outer-sphere inner-sphere complex or solven t-separated ion pair c o m p 1 ex K5 K3 Zn(NO3)2(H,O)g(aq) [Zrl(NO3)Jx (H2o)y. dinitrato inner- ion aggregate sphere complex The equilibrium constants Kl and K4 for the association of zinc and the nitrate as the solvent-separated ion pair and the inner-sphere complex are 1.89 dm3 mol-l and 0.14. The previously reported value', of 0.15 is close to the K4 value which we report. Both Kl and K4 are concentration-dependent, as seen from table 3. TABLE 3.-EQUILIBRIUM QUOTIENTS FOR THE ION ASSOCIATION OF THE NITRATES OF METALS WITH UNFILLED VALENCE SHELLS concentration electrolyte /mol dmP3 KJdm3 mol-l K4 TlNO, 0.2 0.4 &NO, 0.5 1 .o 2.0 4.0 0.5 1 .o 1.5 2.0 2.5 3.0 0.5 1 .o 2.0 3 .O 4.0 0.5 1 .o 1.5 2.0 2.5 3.0 0.25 0.50 1 .o 1.5 2.0 2.5 0.96 0.42 1.31 0.674 0.592 0.115 1.89 1.88 2.7 4.44 - - 1.65 2.24 2.82 3.6 5.2 1.74 2.87 4.83 17.3 - - 6.43 10.66 12.15 16.03 - - 1.19 2.3 0.76 1.39 3.82 - - 0.13 0.14 0.26 0.37 0.252 0.396 0.388 0.657 0.742 0.48 0.568 0.498 0.543 0.635 0.744 0.144 0.35 0.345 0.472 0.54 0.70 -R.L. FROST A N D D . W. JAMES 3273 The shape of the v4(E’) profile is worthy of comment. The predominant band is centred at 719 cm-l with a small intensity contribution at 745 cm-l. The band at 719 cm-l is asymmetric without being split and may be decomposed into two bands at 715 cm-l and 725 cm-l which may be attributed to the v4(E’) nodes of the free aquated nitrate ion and the solvent-separated ion pair.The intensity of the 745 cm-1 band is weak and may be best attributed to a low-intensity band such as the ion-pair aggregate as with LiNO, solutions. If this is the case the v4 band for the so-called inner- sphere complex must be coincident with the position of the solvent-separated-ion-pair band. Certainly the v4 region of zinc nitrate is very different from that of cadmium nitrate, which is discussed later. The vibrational relaxation time of the nitrate ion in the solvent-separated ion pair is concentration dependent. This may indicate that in this species the cation approaches the anion more closely as the concentration increases but that the interaction is still coulombic rather than an electron-exchange process.The vibrational relaxation of the inner-sphere complex is 1.38 ps, which is comparable to the relaxation of the free nitrate of 1.56 ps. SOLUTIONS OF CADMIUM NITRATE The correlation functions of the v,(A;) band profile for cadmium nitrate solutions show similar characteristics to the correlation functions of zinc nitrate. The isotropic correlation functions show increased negative slope with concentration and long-time linearity with some interference at 4.3 ps, which is also reflected in the anisotropic spectra. The theta function plots for cadmium nitrate are complex at the higher concentrations. The slope of the 0.5 mol dm-3 theta function suggests a peak position of 1048.3 cm-l.The 0.5 and 1.0 cm-l theta functions have a node at ca. 3.8 ps, which implies a phase interference of 4.4 cm-l, i.e. a separation between the major peaks of ca. 4.4 cm-l. A simplistic approach to the higher-concentration theta functions suggests a node at ca. 2.7 ps and the implication is that a peak separation of ca. 6.2 cm-l between the major peaks is evident. The cadmium nitrate solutions were curve resolved with peaks centred at 1047.6 and 1051.5 cm-l for the lower concentrations. Large error mean squares were obtained on a simple two-band fit with residuals at 1042.5 and 1033.0 cm-l. Curve resolution was then carried out in terms of three major bands at 1047.6, 1050.5 and 1042.5 cm-l.No bands of any intensity were found in the 1052.0-1056.0 cm-l region, although some residual occurs at 1066 cm-l which may be attributed to low concen t ra t i ons of ion aggregates. The 1 042.5 cm-1 band corresponds well with the band of the bound nitrate in the C,, nitrate complex reported by Davis and Plane at 1043.0cm-1.5 Table 2 reports the results of the curve analyses. An alternative analysis was carried out with an additional band at 1038.5 cm-l but little improvement in e.m.s. was obtained. The difference in the inner-sphere-complex peak position between cadmium and zinc is a measure of the increased covalency between the metal and the nitrate.16 The increased intensity of the 1042.5 cm-1 peak compared with the intensity of 1045.0 cm-l peak of zinc nitrate adds weight to the concept of increased directional bonding.The vibrational relaxation time for the aquated nitrate ion is 1.56 ps and, as for Zn(NO,), solutions, the vibrational relaxation time for the solvent-separated ion pair decreases with concentration from 1.47 at 0.5 mol dm-, to 1.13 ps at 3.0 mol dm-,. The variation in concentration of the various solution species is shown in fig. 4. The relative concentration (%) of solvent-separated ion pair remains reasonably con- stant over a wide concentration range at a value of ca. 0.44, which shows that nearly half the nitrate molecules are involved in solvent-separated ion pairing. The inner-sphere complex has a relaxation time of 0.88 ps, which is shorter than the relaxation time 106 F A R 783274 I 0 N-I 0 N-SO L V E N T I N T E R A C T I 0 N S x-x-x- I 0 .ol I I I I 1 2 3 L [Cd(NO,), 1 /mol dm-3 FIG.4.-Variation in intensity of the component bands for aqueous solutions of Cd(NO,),. (a) Aquated-nitrate band ; (h) solvent-separated-ion-pair band; (c) directional-nitrate band. of the other nitrate-ion species. The inner-sphere complex with zinc and cadmium cations, and indeed all cations reported here, is characterised by a peak energy lower than the dilute solution band. It has been suggested elsewhere that linkages of this sort involve electron transfer from the metal to the nitrate Recent studies of the n.m.r. of solutions of Cd(NO,), show concentration-dependent change in the chemical shift of Cd2+.17 This change is not observed for cations with a filled valence shell.18 The observations of the 740 cm-* band over all concentrations5 supports the concept of inner-sphere complexes existing at all concentrations. The v, band profile is split for solutions > 1 mol drn-,.There is an increase in intensity at ca. 750 cm-l at higher concentrations (3 and 4mol drn-,) which may be attributed to the appearance of ion aggregates (band component at 1063 cm-l) or an additional band component at 1038.5 cm-l. The sharpness of the component at ca. 740 cm-l confirms that the Cd-ONO, interaction is specific and directional. Equilibria for cadmium nitrate solutions may be described by the following series of equilibria K , K4 Cd2+ + NO;(aq) Cd(H,O),NO,+(aq) e Cd(NO,)(H,O)$(aq) free aquated outer-sphere ions complex inner-sphere complex The equilibrium quotients K, and K, are listed in table 3.The values for Kl are in the main less than the values for solutions of Zn(NO,),. Previously reported values13 suggest that this should be so though the actual values reported are closer to K4 than Kl. The values for K4 are significantly larger for solutions of Cd(NO,), than for Zn(NO,),. This reflects an increasing stability for the inner-sphere complex which may be reflected in the lower energy for the component band in Cd(NO,), solutions.R. L. FROST A N D D. W. JAMES 3275 SOLUTIONS OF LANTHANUM NITRATE The peak positions for dilute solutions of lanthanum nitrate (0.5 mol drn-,) are at 1048.3 cm-l, slightly above the value for the infinite-dilution band. In the concentration range 0.5-2 mol dm-, the isotropic peak frequency is greater than the anisotropic peak frequency.The isotropic bandwidths as measured by the second moment show the bands are not as broad as the isotropic bands of zinc and cadmium nitrates. This feature is also reflected in the anisotropic second moments. The M2@) value for Zn(NO,), at 3 mol dm-, is 306 (cm-1)2 compared with the value of 3 mol dm-3 La(NO,), which is 256 (cm-1)2. The peak position and half-width measurements indicate that contact ion pairing is minimal and that solvent-shared ion pairs may be preferred. The changes in asymmetry on the lower-energy side give a good indication that the experimental profile is composed of overlapping bands whose intensity alters with concentration. The correlation functions for lanthanum nitrate show linearity at long time whereas the correlation functions for 1 , 2 and 3 mol dmP3 exhibit interference in the power spectrum: 1 mol dm-, at 3.3 and 4.7 ps, 2 mol dmP3 at 2.35 and 4.3 ps and 3 mol dm-3 at 2.35 and 4.7 ps.Although it is not possible to predict component-band spacing when more than two overlapping bands make up the experimental profile, the interference times indicate band separations of 7.09 and 3.54 cm-'. Correlation functions do show that the overall profile is becoming broader and is altering in shape with concentration. Such features are further illustrated by the theta functions for lanthanum nitrate, which are shown in fig. 5. Fig. 5 also illustrates the complexity of a theta function formed by a number of overlapping bands.Band-component analysis was carried out in terms of three components centred at tlPS FIG. 5.-Theta functions for lanthanum nitrate solutions. Concentration mol ~ i m - ~ : (a) 0.5; (b) 1 ; (c) 2 ; ( d ) 2.5; (e) 3.0.3276 I 0 N-I 0 N-SO L V E N T INTER ACT I 0 N S 0 . o l 1 I I 1 2 3 [ La(NO,), I /mol dm-3 FIG. 6.-variation in intensity of the component bands of the A ; symmetrical stretching vibrational band for aqueous solutions of La(NO,),. (a) Aquated-nitrate band ; (h) solvent-separated-ion-pair band ; (c) directional-nitrate band. 1047.6, 1050.5 and 1044.5 cm-l, which are attributed to the free aquated nitrate, the solvent-separated nitrate and the directional nitrate. No bands were found in the region 1051-1053.0 cm-l. Fig. 6 shows the variation in intensity of these components as a function of concentration.The striking feature of the figure is the relatively high intensity of the solvent-separated-ion-pair band over all concentrations. A second feature is the existence of the directional-pair band at low concentrations. The relative integrated intensity of the ‘directional’ nitrate band at 0.5 mol dm-3 is 0.136 and grows to a value of 0.4 at 3.0 mol dm-3 at the expense of the aquated- nitrate-ion band, which decreases from 0.571 at 0.5 mol dm-3 to 0.06 at 3 mol dm-3. The variation in intensity of the component bands is quite similar to that shown by solutions of Cd(NO,),, except that the solvent-separated species is preferred to the aquated species. The vibrational relaxation time of the band attributed the solvent- separated ion pair decreases from 1.2 ps at 0.5 mol dm-3 to 1.01 ps at 3 mol dm-3, which is parallel to the behaviour seen for solutions of Zn(NO,), and Cd(NO,),.The nitrate is more strongly perturbed by the cation which is directionally interacting with the anion along a fixed direction. This increased perturbation of the inner-sphere complex is reflected in the vibrational relaxation time of 0.9 ps for the band attributed to the complex. This interaction is closer to that seen for Zn(NO,), (v = 1045 cm-l) than Cd(NO,), (v = 1042.5 cm-l). The equilibrium quotient K, of 1.74 dm3 mol-1 at 0.5 mol dm-3 is compatible with an earlier value13 of 1.26 dm3 mol-l. The value of K4 shows little concentration dependence and at the concentrations studied may approach a true equilibrium constant.SOLUTIONS OF THORIUM NITRATE The peak position of the experimental profile varies from 1048.5 cm-l at 0.25 mol dm-3 and increases to 1049.6 cm-l at 2.5 mol dm-3, i.e. over a 2.25 mol dm-3R. L. FROST AND D. W. JAMES 3277 concentration the profile position increases by only 1 cm-l yet the band at high concentrations shows striking asymmetry, as indicated by the asymmetry index which is always large and negative. At 2.0 mol dm-3 the index has a value of - 7.20, which is one of the largest negative indices we have measured. The second moments of the isotropic band vary between 56 (cm-1)2 at 0.25 mol dm-3 and 72 (cm-1)2 at 3 mol dm-3, and the second moments of the anisotropic bands vary from 144 (cm-1)2 to 256 (cm-1)2 over this concentration range.The second-moment values provided no indication of the pronounced asymmetry. The correlation functions are non-linear at all concentrations and show interference at ca. 1.0 ps, suggesting peak separation of ca. 16.0 cm-l between two peaks. The correlation functions also show some interference at 4.3 ps. The theta function is complex and not open to simple analysis. Although attempts were made at curve resolution in terms of three components, this proved unsuccessful. The simplest analysis requires decomposition of the experimental profile into four bands. The 2.5 mol dm-3 spectra show two features on the lower-energy side which together with the bands assigned to the aquated nitrate and the solvent-separated nitrate make a four-band basis of analysis.The bands were centred at 1047.6, 1050.5, 1041.9 and 1032.5 cm-l. The latter two bands are attributed to directional-nitrate bands in the form of a monodentate and a bridging nitrato ligand. The nitrato-thorium interaction 0 . 0 . [Th(NO,),] /mol drn-3 FIG. 7.-Variation in intensity of the component bands for aqueous solutions of Th(NO,),. (a) Aquated- nitrate band; (b) solvent-separated-ion-pair band; ( c ) first directional-contact-ion-pair band, ( d ) second directional-ion-pair band. is becoming increasingly covalent, as is reflected in the band positions. Fig. 7 shows the variation in relative integrated intensity with concentration for the four bands. As with the other nitrates studied, the analysis shows a significant intensity of the band attributed to the solvent-separated nitrate over a wide concentration range and the existence of the directional-nitrate species at all concentrations. Thus the intensity of the 1051.0 cm-l band increases from 0.392 at 0.25 mol dm-3 to a maximum of 0.513 at 1.5 mol dmW3 and decreased to 0.46 at 3 mol dmP3. The 1041.9 band has an intensity of 0.056 for 0.25 mol dmP3 increasing to 0.321 at 2.5 mol dm-3.The 1032.5 cm-l band3278 I 0 N--I 0 N-S 0 L V E N T , I N TE R A C T I 0 N S shows intensity growth over the concentration range from 1.0 to 2.5 mol drnP3. The growth of the bands attributed to directional nitrate is at the expense of the aquated-nitrate-ion band. A previous study on the structure of aqueous thorium nitrate solutions analysed the v1 region into two bands at 1052.0 and 1038.0 cm-l.19 The bands are in reasonable agreement with this study which in effect has taken these two bands and split them into two other bands.The vibrational relaxation of the solvent-separated-ion-pair band is 1.05 ps and for the directional-nitrate bands is 0.746 and 0.647 ps. The influence of the cation on the nitrate is exemplified by these very short relaxation times as a very strong perturbation of the nitrate. The intense polarizing power of the thorium cation resulting from the high charge/radius value causes a large symmetry perturbation of the nitrate ion which reduces the symmetry of C,, or C, for the two types of directional nitrate. This symmetry perturbation is reflected in the short relaxation times of the directional nitrate bands.The equilibria for aqueous thorium nitrate solutions may be represented: aquated solvent-separated inner-sphere ion pair complex The curve analysis of the thorium nitrate experimental profile has been made in terms of four bands which may be attributed to the aquated nitrate ion, the solvent-separated ion pair, the directional nitrate or inner-sphere complex and a second directional nitrate, perhaps a bridging nitrato complex. There are four equilibria shown with five species. It is likely that other equilibria are also present in which two or more of the nitrate ions are complexing the thorium cation. There have been a number of measurements of equilibrium constants for solutions of thorium nitrate. These are in general for an equilibrium written as Th4+(aq)+ NO;(aq) --+ ThNO!+(aq).Previous values from Raman spectralg assign bands to a contact ion pair and so these may be assumed to encompass both Kl and K4. Based on our analysis a value of 0.78 dm3 mo1-1 would be calculated at 0.5 mol dm-3, which compares with a value of 1.0 dm3 mol-l previously reported.lg Other determinati~nsl~. 2o using, for example, ion exchange, report concentration-dependent values from 4.7 dm3 mol-l at 0.5 mol dm+ to 16 dm3 mol-l at 2 mol dm-3. These values are obviously related to the values of Kl which we report. This is perhaps expected and illustrates that the first association determined using equilibrium and transport measurements is for the solvent-separated ion pair. It has been shownlg that statistically the ratio K4/K5 for solutions of thorium nitrate should be between 2.0 and 2.4. The ratio of the band areas for the two directional nitrates is 2.0 at 2.5 mol dm-3, which agrees well with the predicted value of K 4 / K 5 . We thank the Australian Research Grants Committee for grants enabling the purchase and maintenance of the Raman spectrometer. Y. Sze and D. E. Irish, J . Solution Chem., 1978, 7, 395. A. T. Lemley and R. A. Plane, J . Chem. Phys., 1972, 57, 1648. D. E. Irish and M. H. Brooker, in Advances in Infrared and Raman Spectroscopy, ed. R. J. H. Clark and R. E. Hester (Heyden, London, 1976), vol. 2, chap. 6, p. 212.R. L. FROST AND D. W. JAMES 3279 T. Moeller and P. W. Rhymer, J . Phys. Chem., 1942, 46, 477. A. R. Davis and R. A. Plane, Inorg. Chern., 1968, 7, 2565. R. E. Hester and R. A. Plane, Inorg. Chem., 1964, 3, 769. ’ T. Chang and D. E. Irish, J . Solution Chem., 1974, 3, 175. J. H. B. George, J. A. Rolfe and L. A. Woodward, Trans. Faraday. SOC., 1953, 49, 375. D. W. James, R. D. Carlisle and W. Leong, Aust. J . Chem., 1970, 23, 779. lo P. M. Volmer, J. Chem. Phys., 1963, 39, 2236. l 1 C . W. Davies, Trans. Faraday SOC., 1927, 23, 351. l 2 W. H. Banks, E. D. Righellato and C. W. Davies, Trans. Faraday SOC., 193 1, 27, 621. l 3 R. M. Smith and A. E. Martell, Stability Constants (Plenum Press, New York, 1976). l 4 D. J. Gardiner, R. W. Jackson and B. P. Straughan, J. Mol. Struct., 1979, 54, 31. l 5 M. H. Brooker and D. E. Irish, Can. J. Chem., 1971, 49, 1510. l6 R. E. Hester and C. W. J. Scaife, J . Chem. Phys., 1967, 47, 5253. l 7 R. L. Frost and D. W. James, unpublished results. l9 B. G. Oliver and A. R. Davis, J. Inorg. Nucl. Chem., 1972, 34, 2851. 2o V. V. Fromin and E. P, Maiorova, Zh. Neorg. Khim., 1956,1,1703,2749; P. H. Tedesca, V. B. DeRumi and J. A. Gonaleg Quintana, J. Inorg. Nucl. Chem., 1968, 30, 987. D. W. James and R. E. Mayes, Aust. J . Chem., in press. (PAPER 1 /2002)

ISSN:0300-9599    

DOI:10.1039/F19827803263

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 3281-3286 Production of CO from CO, by Reduced Indium Oxide BY KIYOSHI OTSUKA,* TAKAO YASUI AND AKIRA MORIKAWA Department of Chemical Engineering, Tokyo Institute of Technology, Ookayama, Meguro-ku, Tokyo 152, Japan Received 4th January, 1982 The forward and backward reactions of the process In,O,(s) + xCO(g) $ In20a-z(s) + xCO,(g) proceed rapidly and repeatedly at 673 K when the degree of reduction of In,O, is low. However, the yield of the backward reaction decreases as the degree of reduction rises above ca. 40%. This is ascribed to a decrease in the number of surface-active sites for the decomposition of CO,. The free-energy change of the reaction evaluated from the ratio of partial pressures P(CO,)/P(CO) at equilibrium agreed well with results obtained by other workers, assuming indium metal as the reduced species.Since carbon dioxide is the most stable and most highly oxidized state of carbon, temperatures > 2000 K are needed to decompose it thermally into carbon monoxide and oxygen. Recently we have reported that the decomposition of water proceeds smoothly at or above 573 K on reduced indium oxide.' The preliminary reduction of indium oxide has been carried out by various reductants such as carbon, biomass,l hydrogen or carbon monoxide.2T We may expect that the decomposition of carbon dioxide should also occur on reduced indium oxide according to the following (1) reaction : In,O,-, + xC0, -+ In,03 + xC0 where In203-, represents crystalline In,O or a mixture of indium metal and In,03. In this report we examine the conditions for the decomposition of CO, on a reduced sample of In,O,.The existence of crystalline In20 has been suggested by several w o r k e r ~ . ~ - ~ However, there is still controversy concerning its exi~tence.''~ Hence the reduced state of In,03 will be described on the basis of the free-energy change for eqn EXPERIMENTAL (1). The In,03 used was a reagent-grade powder (purity > 99.9%) obtained from the Wako Pure Chemical Co. The surface area of the fresh In,O, measured by nitrogen adsorption (using the B.E.T. method) was 13.9 m2 g-l. Carbon dioxide was purified three times by trap-to-trap distillation. Carbon monoxide was purified by passing it through a silica-gel adsorbent cooled at 195 K. The experiments were carried out using a conventional mercury-free gas-circulation apparatus of cu.320 cm3 volume. After the partial reduction of In,O, by CO or H,, the gases in the system were pumped out and the decomposition of CO, was started by adding CO, (21.3 x lo3 Pa). The progress of the reaction was followed by measuring the amounts of CO and CO, present by gas-chromatographic analysis. RESULTS AND DISCUSSION Fig. 1 shows repeated forward and backward reactions of step (1) at 673 K. Before the experiments the In,O, sample was reduced by hydrogen at 673 K to a degree of 328 13282 PRODUCTION OF CO USING In,O, reduction of 27.8%, as determined on the basis of the oxygen atoms removed. The right-hand vertical axis indicates the percentage recovery of the reduced oxide to In,O, evaluated from the quantity of CO produced.Decomposition of CO, proceeded rapidly at 673 K and the reaction was almost complete within 50 min. The condensation of CO, at 77 K (at A) caused the backward reaction; i.e. the reduction of the oxide was initiated by CO which had been formed before point A. Evaporation of the condensed CO, at point B reinitiated the decomposition of CO,. The results in fig. 1 show that both the forward and backward reactions of equilibrium (1) pro- ceed smoothly and repeatedly at 673 K. 6.0 - 0 E -3 4.0 2 z P I 0 - 3 E .- 2 2.0 cc 0 4- t: m F 0 t B I 1 I 50 100 150 ti me / m in 100 m 0 N e 0 x 0 I u 50 & 0) w * E & 0) a 1 FIG. 1 .-Repeated forward and backward reactions of equilibrium (1) at 673 K. The weight of In,O, used was 0.20 g.Fig. 2 shows the percentage recovery of the reduced oxide to In,O, by CO, at 673 K as a function of time for various samples with different initial degrees of reduction. The positions of the downward and upward arrows indicate the length of time for which the temperature of the reactor increases and the time at which the temperature reaches 773 K, respectively. Although the rate of CO production for the sample having an 85% degree of reduction was very slow at 673 K, the reaction was completed rapidly when the temperature was raised to 773 K. Fig. 3 shows the effect of CO, pressure on the rate of formation of CO at 673 K for samples of 1 1.4 and 8.1 % degree of reduction. The rate depends on the pressure of CO,. This suggests that the rate-determining step is not the solid-state diffusion process of a reduced species of the oxide but the reaction of CO, with this species on the surface.The large difference observed between the apparent activation energies of the decomposition of water (50.1 kJ rn~l-')~ and of CO, (ca. 110 kJ mol-I, see fig. 5 ) also supports the above considerations. The quantity of CO produced and the percentage recovery of the reduced sample obtained from fig. 2 are plotted in fig. 4 as functions of the degree of reduction of In,O,. The maximum in the amount of CO produced in 20 min lies at 2540% reduction. For the sample having a low degree of reduction, the rate of CO productionK . OTSUKA, T. YASUI AND A. MORIKAWA 3283 FIG. 2.-Percentage recovery of the reduced oxide to In,O, as a function of time for samples with different initial degrees of reduction.Initial degree of reduction and weight of the sample as follows: V, 4.9% (0.50 g); 0 , 9 . 6 % (0.50 g); D, 17.7% (0.20 g); V, 19.1% (0.20 g); A, 27.8% (0.20 g); 0,42.6% (0.20 g); 0, 55.3% (0.50 g); 0, 65.1 % (0.20 g); A, 79.6% (0.20 8); a, 84.8% (0.10 g). 0 5 10 15 20 p(CO,)/ lo3 Pa FIG. 3.-Effect of CO, pressure on the rate of CO production for the following degrees of reduction of In,O,: 0, 11.4%; A, 8.1 %. is fast enough to complete the reaction in 50 min, as can be seen in fig. 2. Hence, the quantity of CO produced for samples of low degrees of reduction is restricted by the quantity of reduced species able to react with CO,. On the other hand, the quantity of CO produced is governed by the rate of reaction of CO, with the reduced species for the sample having a high degree of reduction.Both the percentage recovery and the quantity of CO produced in 20 min decrease sharply when the degree of reduction is increased above ca. 40%. The rates of decomposition of CO, have been measured3284 PRODUCTION OF CO USING In,O, at two different temperatures (598 and 648 K) for samples with the same degree of reduction. Sets of rates were obtained for samples of different degrees of reduction (1 0.9-80.2%). The apparent activation energies and pre-exponential factors, obtained from an Arrhenius plot of the rates at the two different temperatures, are plotted as functions of the degree of reduction in fig. 5 . The results in the figure indicate that the activation energies do not change with the degree of reduction but that the pre-exponential factors drop sharply above ca.30% reduction. This suggests that the nature of the reduced In,O, species active in CO, decomposition does not change with the degree of reduction up to ca. 80%. The drop in the values of the pre-exponential factor strongly suggests that the sharp decrease in the yield of CO above ca. 40% reduction level (fig. 4) can be ascribed to the decrease in the number of surface-active sites to react with CO,, probably owing to the decrease in the surface area of the reduced particles caused by their sintering. 8 8 6, 0 2 x Y 0 50 100 degree of reduction (%) FIG. 4.-Effect of the degree of reduction on the quantity of CO produced and the percentage recovery to In,O,.Quantity of CO produced: A, in 5 min; A, in 20 min. Percentage recovery to In,O,: 0, in 20 min; 0, final value at 773 K. The final percentage recovery to In,O, at 773 K (closed circles in fig. 4) becomes lower as the degree of reduction decreases, which can be ascribed to the presence of highly reducible surface oxygen atoms on fresh In,O,: such oxygen atoms cannot be reformed by the oxidation with CO,., The dashed curve in fig. 4 is the least-squares plot for the closed circles assuming the following equation: Y = 100(1-;) (2) where a is a constant, Y is the percentage recovery to In,O, and X is the degree of reduction of the sample. The number of active surface oxygen atoms evaluated from the value of a is 2.0 x mol (g In203)-l. The crosses in fig.4 are the percentages of final recovery to In20, observed in the case of water decomp~sition.~ The number of active surface oxygen atoms obtained from water-decomposition data is 1.8 x mol g-l. Both values are in fair agreement with each other.K. OTSUKA, T. YASUI A N D A. MORIKAWA 3285 150 - I - 0 E 21 --. x $100 S a, s 0 m .A Y .- + U 50 2 a a 0 4,000 100 50 10 50 70 0 degree of reduction (%) FIG. 5.-Apparent activation energy of CO production and pre-exponential factor as functions of the degree of reduction of In,O,. TABLE 1 .-EXPERIMENTAL VALUES OF P(CO,)/P(CO) AND THE FREE-ENERGY CHANGE OF EQN (3) AGO of eqn (3)/kJ temp./K P(CO,)/P(CO) this work ref. (10) ref. (1 1 ) ref. (12) 550 0.053 _+ 0.008 - 40.2 -41.2 - 38.9 - 44.4 600 0.088 f 0.003 - 36.3 - 38.7 - 35.6 -41.0 646 0.1 13 f 0.003 - 35.2 - 36.3 - 32.6 - 38.0 673 0.134 0.008 - 33.7 - 34.9 - 30.8 - 36.2 700 0.158 f 0.006 - 32.9 - 33.5 - 28.9 - 34.4 773 0.245 &0.001 - 27.2 - 29.7 - 23.8 - 29.5 Broch and Christensen8 could not prove the existence of crystalline In,O after detailed X-ray analysis of reduced In,O,.Our X-ray spectroscopic studies on the partially reduced In,O, could only indicate the existence of In,O, and indium metal. Hence, the reduced state of the oxide has been studied by measuring the equilibrium pressures of CO, and CO over the partially reduced indium oxide (< 51 % degree of reduction) at different temperatures. The ratios of the observed pressures of CO, and CO are shown in the second column of table 1 and yield the free-energy change of reaction (3) as shown in the third column of table 1.Stubbs et aLIO and Hochgeschwender and Ingrahaml' determined the free-energy change for reaction (4) 2 In(1) + 3 CO,(g) In,O,(s) + 3 CO(g) (3) using gas-equilibration techniques. Combining their data with literature data for the free energies of formation of water vapour and of CO, gas,12 the free-energy changes3286 PRODUCTION OF co USING h203 of eqn (3) were calculated at different temperatures. They are listed in the fourth and fifth columns of table 1. The values listed in the last column were obtained from the free energies of formation of CO, and those of In203 determined by Newns and Pelmore13 using e.m.f. measurements with the galvanic cell In(l), In,03(s)~0.85Zr02 + O . 1 5CaO(Ni(s), NiO(s). The values obtained in this work agree well with those of other workers and support the previous suggestion3 that the In203-,(s) in equilibrium (1) is a mixture of indium metal and In203. K. Otsuka, Y. Takizawa, S. Shibuya and A. Morikawa, Chem. Lett., 1981, 347. K. Otsuka, T. Yasui and A. Morikawa, J. Catal., 1981, 72, 389. K. Otsuka, T. Yasui and A. Morikawa, Bull. Chem. Soc. Jpn, 1982, 55, 1768. W. Klemm and H. U. V. Vogel, Z. Allg. Chem., 1934, 219, 45. K. A. Klinedinst and D. A. Stevenson, J. Chem. Thermodyn., 1973, 5, 21. T. J. Anderson and L. F. Donaghey, J. Chem. Thermodyn., 1977, 9, 617. L. Brewer, Chem. Rev., 1953, 52, 38. N. C. Broch and A. N. Christensen, Acta Chem. Scand., 1966, 20, 1996. A. J. Van Dillen, J. W. Geus and J. H. W. de Wit, J . Chem. Thermodyn., 1978, 10, 895. K. Hochgeschwender and T. R. Ingraham, Can. Metall. Q., 1967, 6, 293. I. Barin and 0. Knacke, Thermochemical Properties of Inorganic Substances (Springer-Verlag, Berlin, 1973 and 1977). lo M. F. Stubbs, J. A. Schufle and A. J. Thompson, J. Am. Chem. Soc., 1952, 74, 6201. l3 G. R. Newns and J. M. Pelmore, J. Chem. Soc. A, 1968, 360. (PAPER 2/007)

ISSN:0300-9599    

DOI:10.1039/F19827803281

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 3287-3296) Velocity Correlations in Aqueous Electrolyte Solutions from Diffusion, Conductance and Transference Data Application to Concentrated Solutions of Cadmium Chloride BY REGINALD MILLS Diffusion Research Unit, Research School of Physical Sciences, Australian National University, Canberra, A.C.T. 2600, Australia A N D HERMANN GERHARD HERTZ* Institut fur Physikalische Chemie und Elektrochemie der Uni\ ersitit Karlsruhe, 7500 Karlsruhe, West Germany Received 15th January, 1982 Tracer-diffusion and intradiffusion coefficients have been measured for the Cd and C1 ion constituents and for water molecules in aqueous solutions of cadmium chloride in the concentration range 0.1-5.8 mol dm+ at 298.15 K . These measured coefficients have been combined with transport data from other sources to calculate velocity cross-correlation coefficients (VCCs).In contrast to a fully dissociated 1 : 2 salt such as magnesium chloride, the VCCs for CdC1, show much greater positivity, and in particular the anion-anion VCC is positive over the whole concentration range. An explanation for the negative transference numbers in concentrated CdCl, solutions is offered using the linear-response-theory formalism. In previous studies of this kind' velocity cross-correlation coefficients (VCCs) for aqueous electrolyte solutions have been calculated from conductance, transference- number, mutual- and tracer-diffusion and activity-coefficient data. The definition of VCCs and their mode of calculation from the primary transport data are detailed in the above papers.The electrolyte systems studied so far comprise a number of both I : 1 and 1 : 2 salts. Within the latter group have been BaCl,, CaCl,, MgCl, and NiCl,, and all of these show little or no complex formation. It is therefore of considerable interest to examine an electrolyte such as CdCl, for which there is strong evidence for complexation in aqueous solution. The evidence for complex formation in CdCl, solutions stems from a variety of observations. For example, the fact that in transference-number experiments5* the value for the cation constituent becomes negative at high concentrations has been explained in terms of its incorporation into complex negative ions. Reilly and Stokes7 have used activity-coefficient data to examine complexation in cadmium chloride solutions and have calculated stability constants for complex formation.There is also a considerable amount of spectroscopic evidence for the existence of complexes in these solutions.*. The Experimental section of this paper concerns the measurement of the tracer- diffusion coefficients of the two ion constituents and the solvent in CdCl, solutions, and is the first comprehensive diffusion study of such coefficients for a complexed system. Note that Paterson et aI.lo1 l 1 have measured tracer-diffusion coefficients for the two ion constituents (not the solvent) in aqueous CdI, solutions. However, the concentration range studied was much lower than that reported here for CdCl,, and their data were analysed in the irreversible-thermodynamics formalism.32873288 TRACER DIFFUSION I N CdC1, SOLUTIONS EXPERIMENTAL Analytical-grade CdCl, (Univar) was used without further purification. The radioactive tracers were Il5Cdrn and ,,Cl, both being obtained from the Radiochemical Centre, Amersham, U.K. The diaphragm-cell technique was used to measure tracer-diffusion coefficients for the Cd and C1 constituents; the procedures are described fully by Mills and Woolf.12 The self-diffusion coefficients of water in the solutions were determined by the n.m.r. spin-echo method and details of this technique have been given by Harris et a/.', TABLE 1 .-TRACER-DIFFUSION COEFFICIENTS OF ION CONSTITUENT SPECIES (Cd AND C1) AND H,O IN AQUEOUS CdC1, SOLUTIONS AT 298.15 K c/mol dmP3 Cd c1 H,O 0 0.057, 0.1 13, 0.204, 0.508, 0.565, 1.134 2.045 3.043 4.076 5.787 0.712 0.734 0.733 0.723 0.666 0.566, 0.404, - 0.155, 0.047, 2.030 1.420 1.299 - - 0.945 0.713 0.466, 0.170, 0.050, 2.299a 2.275 2.241 2.191 2.065 1.828 1.496 1.139 0.826 0.433, - a Value of Mills.', RESULTS The tracer-diffusion data for all constituents in the CdCl, + H,O system at 298.15 K are listed in table 1.The ion-constituent coefficients have an average precision of +0.4% and the solvent coefficients In fig. 1 we have plotted the data in table 1 and for comparison the corresponding tracer-diffusion data for the MgC1, + H,O system.15 Magnesium chloride is a strong 1 : 2 electrolyte and shows no evidence of complex formation. The difference between the two sets of results is striking. We take the C1- results first; in the CdC1, case, its tracer-diffusion coefficient decreases very steeply from the infinite-dilution value as compared with MgC1,.This rapid decrease may be attributed to the fact that even at very low concentrations, the normally mobile C1- is bound to the slower moving Cd ion constituent in complexes such as CdC1+ and CdCI,. Cd2+ has practically the same limiting diffusion coefficient as ME,+, but as the concentration increases its value rises above the latter and shows a small maximum at ca. 0.05 mol dm-3. The reason for this slight maximum in the Cd diffusion pattern probably is due to the fact that in complexes such as CdC1+ and CdCl,, the faster moving C1 constituent increases their mobility when compared to the bare Cd2+. Thereafter with increasing concentration the Cd constituent coefficients approach slowly to those of Mg2+ and coincide with them at ca.3 mol dmP3. It is also 1 %.R. MILLS AND H. G. HERTZ 3289 worth noting that the C1 ion constituent values also coincide with those of Cd and Mg2+ at ca. 5 mol dm-3. Another important difference between diffusion in the CdCl, and MgC1, systems is that in the former the mobility of water molecules is considerably greater than in the latter. A simple explanation is that Mg2+ is regarded as a strong structure-making ion with both primary and secondary water coordination shells. Cd2+ is larger than Mg2+ and even in its uncomplexed state would be less of a structure-maker than Mg2+. ... 0 1 2 3 4 5 6 c/mol dm-3 FIG. 1.-Tracer-diffusion coefficients of ion constituents and of water in CdCI, and MgCI, solutions from 0.06 to 5.8 mol dm-3 ( T = 298.15 K).However, in addition if it forms complexes with C1- such as CdCl+ and CdC1, then its structure-making capacity would be diminished even further. The diminution of structure in water, of course, leads to higher mobility of the water molecules. VELOCITY C RO S S-CO R R E LAT I ON C 0 E F F I C I EN T S GENERAL In this section the data reported above are combined with other transport and thermodynamic data to calculate generalized transport coefficients which have their basis and definition in linear response theory. One could have chosen to use the formalism of irreversible thermodynamics and calculated ionic lij coefficients, as was done in the CdI, case by Paterson et a1.l07 l1 However, it is felt that the VCC approach3290 TRACER DIFFUSION I N CdCl, SOLUTIONS is a more powerful one for complex-forming electrolytes in that the behaviour of complexes is reflected better by a dynamic variable such as the velocity.The calculation of VCCs for CdCl, solutions follows the procedures laid down in the series of papers by Hertz et previously cited. Sources of data, other than for tracer diffusion, used in these calculations are as follows : conductances, McQuillan ;6 transferences numbers, McQuillan ;6 mutual-diffusion coefficients, Rard and Miller ;16 activity coefficients, Robinson and Stokes.17 0.6 0.5 0 1 2 3 4 - cs/mol dm-3 FIG. 2.-Ionic velocity correlation coefficients faa, fa, and f,, in an aqueous solution of CdC1,.For comparison the ionic VCCs in a MgCl, uij) are given as dotted curves. The dot-dashed curves represent the predicted or standard VCCS,~’, of the constituents of CdC1, if the latter were present as a completely undissociated molecule. For further details see text ( T = 298.15 K). IONIC VELOCITY CORRELATION COEFFICIENTS The anion-anion, cation-cation and anion-cation VCCs,f,,,f,, andf,,, which have been computed from the experimental data according to the methods outlined in two previous papers3> are plotted against the concentration (on the molarity scale) in fig. 2. In the same figure we have included the corresponding coefficients of the classical fully dissociated 1 : 2 electrolyte MgCl, as dotted curves. All theLj curves for CdC1, are shifted to more positive values relative to those of MgCl,.Thus there are more positive velocity correlations between the various constituents of the solute CdCl, ; i.e. if a Cd2+ ion has a given positive (or negative) velocity at t = 0, then shortly after this instant, the C1- ions on the average still have a positive (or negative) velocity, and likewise, if a C1- at t = 0 has a given positive (or negative) velocity, shortly afterR. MILLS AND H. G . HERTZ 329 1 this instant the mean representative of all the other C1- ions has a positive (or negative) velocity. This seems to be a simple consequence of the complex formation; the constituents are (at least partly) tightly bound to each other and thus they move together, i.e. they have positive velocity correlations. In order to get a deeper and more quantitative insight it is instructive to use velocity-correlation terminology and examine the consequences of treating CdCl, as a purely molecular compound dissolved in water.1 We first divide thehi into an intramolecular and an intermolecular part, i.e. f i j = (fijlinter + (hjlintra i,.i = a, c (1) where the subscripts a, c refer to anion and cation. Then for velocity correlations within the molecule itself it can be shown29l8 for CdC1, that (fijlintra = Dc = Da (2) where the Di are self-diffusion coefficients of constituent i. for the intermolecular VCCs. as under We now formulate standard values (those predicted under certain ideal conditions) where components 1 and 2 refer to water and CdCl, and Mi and xi are the molecular masses and mole fractions of component i, respectively.Introducing eqn (2)-(4) into eqn ( I ) we obtain In fig. 2 these standard values, calculated on the assumption that CdCl, is a molecular compound, are compared with the experimentally determined values. Thus eqn (5)-(7) are shown as dot-dashed curves. First it is seen that, at high concentrations, fac andfcc are comparatively close to our predicted standard curves for the molecular species, both being more positive than the corresponding standard curves. In contrast to this faa is markedly more positive than the standard quantity fza. Moreover, for the MgC1, solution there is always a shift of thefij which is much stronger and in the negative direction relative to thef?' curves. (If computed with masses and diffusion coefficients of the MgC1, solution, the f: values for the hypothetical MgCl, molecule would come out much the same.) However, remarkably enough, the order remains the same, i.e.fcc < faa <far. To account for the deviations in the experimentalf,' from the standard molecular values, we turn now to examine the conductance, which is one of the factors in their3 2 9 2 TRACER DIFFUSION IN CdC1, SOLUTIONS computation. The ‘ reduced ’ equivalent conductance has previously been defined as where A is the equivalent conductance, z, is the charge of the cation Faraday constant. This ‘reduced’ conductance has the same units as the and velocity-correlation coefficients. In fig. 3 we show the value of A* of the concentration. This figure also gives A* for the MgC1, solution, and F is the self-diffusion as a function and the shift 0 1 2 3 ‘Es/mol dm-3 FIG.3.-Reduced equivalent conductance A* = ART/z, F 2 of CdCI, and MgCl, in water as a function of the concentration (molarity scale). For comparison the cationic self-diffusion coefficient, D,, of Cd*+ is also shown ( T = 298.15 K). of the coefficients faa,fac and fCc to more negative values as compared with those for CdCl, is accompanied by a much stronger increase in the reduced equivalent conductance A*. The reduced equivalent conductance can be separated in terms of transference numbers so that A* = t, A* -k t, A* withR. MILLS A N D H. G . HERTZ For Cs 2 3.5 mol dmP3, t, < 0, and considering eqn (9) we must have fac > D c +fee* If we assume, considering eqn (I), (5) and (7), that fcc = (fachnter then it follows from inequality (1 1) that we must have (fachntra > D c (1 1') and with eqn (9) we get (faclintra = D c - tc A*+ Application of experimental data tells us that at high concentrations (i.e.Cs = 4- 4.5 mol d m 3 ) (faJintra is only ca. 1 % larger than D,. So a very small positive deviation of (fac)intra from the cationic self-diffusion coefficient is sufficient to explain why the transference number is negative. In the case of MgCl, the cationic transference number is positive; thus we have fa, < D c +Sac (faclintra < D c - which means that We have seen that the deviation from neutral-molecule behaviour, which determines the contribution tc A* to the reduced equivalent conductance, is very small, and now turn to the anionic contribution.We use eqn (l), and determine ( faa)inter: (faahnter = faa - (faalintra where (fa.Jintra is the same figure as derived from our t,A* values. The results for the four concentrations 3, 3.5, 4.0 and 4.5 mol dmP3 are -0.21 x lop9, -0.18 x lop9, - 0.14 x and - 1.2 x lop9 m2 s-l, respectively. However, according to eqn (3), for the molecular species we should have (faa)inter = 2fcc, the corresponding four fCc values are -0.13 x -0.12 x -0.092 x and -0.072 x m2 s-l (see fig. 2); thus the (faa)inter values are too positive. This may clearly be seen from fig. 2.T Thus we obtain the result that the electric conductivity of a concentrated CdCl, solution can be related to the fairly strong positive excess velocity correlations of the anions among themselves. VELOCITY CORRELATION COEFFICIENTS INVOLVING THE WATER MOLECULE In fig.4 we show the solute-molecule-solvent-molecule velocity correlation coefficient fsw as a function of the (analytical) CdC1, concentration. As in our previous work1* 2 * it has been calculated from the relation d In y* Dp = - 3 ~ ( l + c W , ) 3 ( ~ + ~ : czlc? M, dc,* )fsw where D,, is the mutual diffusion coefficient and y' the activity coefficient of the solute CdCl, (component 2). In this formula the molarities (c,) and molalities (c;) refer to 1 cm3 and 1 g, respectively. Fig. 4 also gives the water-water velocity correlation coefficient, fw,, which again has been calculated in the same way as described previously1* 2 v t D, = 0.160 x m2 s-l at E, = 4 mol dm-3.3294 TRACER DIFFUSION I N CdC1, SOLUTIONS The dashed curves in fig.4 represent the corresponding standard VCCs which one would predict if the solution were an ideal mixture18 (apart from an activity factor 1/3 occurring in the expression for f",) with and c , f i w I 1 I 1 I , 0 1 2 3 G 5 zs /mol dmA3 FIG. 4.-Solute-water (f,,) and water-water uww) VCCs in an aqueous solution of CdC1, as a function of the CdC1, concentration. The standard VCCs cr",) are also shown ( T = 298.15 K). It will be seen that forf,, the deviation from the standard behaviour is appreciable, which is a consequence of the marked peculiarity of the activity factor (1 + c; d In y* /dc;). DISCUSSION The most interesting point which should be discussed is the comparison of the conventional interpretation of the negative cationic transference number with the one given within the framework of linear-response theory.Let Ai be the equivalent conductance of the ith ionic species; then according to Spirolg the transference number of the ion constituent Cd is given by the expressionR. MILLS AND H. G . HERTZ 3295 where we have neglected the contributions of the neutral CdCl, molecules and of the very small amount of CdClj- ions.' With the stability constants as reported by Reilly and Stokes7 one calculates at c, = c, = 4 mol dmP3: cCp+ z 0.02 mol dm-3, cCdC,+ z 1.2 mol dm-3, cCdCl2 z 2.0 mol dmP3, ccdc,<; z 0.6 mol dm--3 and cC1- z 0.6 mol dm-3, neglecting the influence of the activity coefficients (which are also reported in the paper of Reilly and Stokes).The general physical rituation is little changed if these activity coefficients are incorporated.* If the stability constants of Reilly and Stokes, which were based o n data measured to ca. 1.2 rnol dm-3 in the CdCI, concentration, are valid up to 4 rnol dm-3 then an interesting anomaly arises. In order to obtain negative transference numbers one must assume that the conductance of a CdC1; ion is about twice as large as that of CdCI+. This seems rather unlikely. However, Lutfullah and Paterson,20 using the procedures of Reilly apd Stokes have calculated the concentrations of the various species present in CdI, solutions. At ca. 0.28 rnol dmP3, where the transference numbers become negative, the concentrations of negative and positive Cd complexes are almost equal and it becomes unnecessary to postulate a large mobility difference.Further e.m.f. measurements of the type performed by Reilly and Stokes to 4 mol dm-3 in CdC1, solutions are obviously necessary to clarify this matter. In conclusion, we wish to describe the particular behaviour of CdCl, (and similar solutions) in terms of single-particle (or distinct) velocity correlation integrals (u~')(O)c~')(t))dt k , l = c,a, i,j= 1,2. c: We recall that om VCCs are defined f k l = N , lx (~(1') ( 0 ) ~ f ) ( t ) ) dr then for the case of a molecular solute according to eqn (4), (6) and (7) we derive 0 1 " 1 M , P( 1 -x, 41) D2 0 M c;c = { ( u p (0) UP) ( t ) ) o dt = -- V (14) 1 " 1 M , V ' ( 1 - ~ , 8 2 ~ ) D, D2+- 0 M N2 czc, = (da)(0) d c ) ( t ) ) , dt = -- V where V' = V/NO is the mean particle molecular volume and M = x, M , +x, M,; N , = N,,n,, n, is the number of moles of solute added to the system.Next we write the corresponding integrals for the case of a fully dissociated electrolyte such as MgCl,. (15) The result is18 * Since this paper was written Prof. Stokes was kind enough to compute the molarities of the various complex species present for given stoichiometric CdC1, concentrations. The computed data indicate that this statement is correct.3296 TRACER DIFFUSION I N CdC1, SOLUTIONS The constants Bi and Ci (i = 1-3) are specific quantities determined by the system; the constants Ai (i = 1-3) are probably of a more general nature but at this stage not exactly known. We may summarize the statements given in eqn (12)-(17) in the following way.Consider a set of N, particles of kind C and 2N, particles of kind A (e.g. elements) which are dissolved in a certain solvent; the volume is fixed and it has the value V. Then if the N, dependence of the single particles' (distinct particles) velocity correlation integrals is described by the limiting behaviour for N2 -, 0 as given by eqn (12)-(14), the solute compound is a molecular one; if, on the other hand, the limiting behaviour as N, -+ 0 is described by eqn (15)-(17), the solute compound is present in the ionized state. Now it will be seen that CdCl, (and similar substances) behave in a particular manner which lies in a characteristic way between the two standard types just mentioned. At low concentrations (small values of N,) c,, lies between the patterns given by eqn (1 3) and (16), that means that c,, is approximately proportional to N;:, whereas according to eqn (13) and (16) it should be proportional to N i and N;Z, respectively. At high solute concentrations c,, follows fairly well the prediction of eqn (1 3), as shown in fig.2. In the same way at small concentrations la, lies between the behaviour of expressions (14) and (1 7). Whereas according to the latter equation c,, again should have a proportionality to N;$ and according to eqn (14) should be K N;l, experimentally one finds something like c,, oc N;;; again the behaviour of la, at high concentrations is fairly well described by eqn (14). Finally, the most marked deviation from the typical standard behaviour occurs for caa.According to the ionic-type behaviour at small concentrations caa should be negative [eqn ( 1 5 ) ] . However, within the range we were able to study, c,, was always found to be positive; this is the sign predicted by eqn (1 2). However, at higher concentrations Caa remains more positive than predicted by eqn (12). So the strongly positive (distinct) particle velocity correlations between the chloride ions are the typical and very specific physical effect which can be related to the interesting anomalies in this type of electrolyte solution. It is the velocity analogy of what from another point of view is described as auto-complex formation, i.e. as a strong crowding of anions around most of the cations. H. G. Hertz, Ber. Bunsenges. Phys. Chem., 1977, 81, 660. H. G. Hertz and R. Mills, J . Phys. Chem., 1978, 82, 952. A. Geiger and H. G. Hertz, J . Chem. SOC., Faraday Trans. I , 1980, 76, 135. A. Geiger, H. G. Hertz and R. Mills, J . Solution Chem., 1981, 10, 83. W. E. Lucasse, J . Am. Chem. SOC., 1929, 51, 2605. A. J. McQuillan, J. Chem. SOC., Faraday Trans. I , 1974, 70, 1558. ' P. J. Reilly and R. H. Stokes, Aust. J . Chem., 1970, 23, 1397 * J. E. D. Davies and D. A. Long, J . Chem. SOC. A , 1968, 2054. H. A. Brune, Thesis (Karlsruhe, 1960). lo R. Paterson and Lutfullah, J . Chem. SOC., Faraday Trans. I , 1978, 74, 93, 103. R. Paterson and C. Devine, J . Chem. SOC., Faraday Trans. I , 1980, 76, 1053. l2 R. Mills and L. A. Woolf, The Diaphragm CeZZ (ANU Press, Canberra, 1968). l 3 K. R. Harris, R. Mills, P. Back and D. S. Webster, J . Magn. Reson., 1978, 29, 473. l4 R. Mills, J. Phys. Chem., 1973, 77, 685. l5 K. R. Harris, H. G. Hertz and R. Mills, J . Chim. Phys., 1978, 75, 391. l6 J. A. Rard and D. G. Miller, unpublished work. l7 R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 2nd edn, 1959). la H. G. Hertz, in Diflusion and Conductance in Ionic Liquids, 2. Phys., Chem. N.F. Suppl., in press. l9 M. Spiro, J . Chem. Ed., 1956, 33, 464. *O Lutfullah and R. Paterson, J . Chem. SOC., Faraday Trans. I , 1978, 74, 484. (PAPER 2/076)

ISSN:0300-9599    

DOI:10.1039/F19827803287

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 3297-3306 Reactions Involving Electron Transfer at Semiconductor Surfaces Part 12.-Nature and Origins of Photoactivity on Oxides of 3d Transition Metals for Elimination Reactions of Secondary Alcohols BY JOSEPH CUNNINGHAM,* BENJAMIN K. HODNETT, MOHAMMAD ILYAS, EDWARD M. LEAHY AND JOHN P. TOBIN Chemistry Department, University College, Cork, Ireland Received 25th January, 1982 Activity for conversicns of the vapours of propan-2-01 and butan-2-01 into products corresponding formally to elimination of H,, H,O or a (C,-C,) bond from the parent alcohol were compared for oxides of the 3d transition metals under thermal- and photo-activation. Use of a gas-chromatographic, continuous reactant-flow technique with f.i.d. detection favoured the detection of products corresponding to a large turnover per surface site (t.a.p.s.).Under these conditions photoenhancement of the ( -H2) and (C,--C,) products at significant levels was detected only over ZnO and TiO, and in the presence of gaseous oxygen. Such photocatalytic activity, and the contrasting absence of photoactivity continuing to high turnover over oxides featuring cations with partially filled 3d levels, is attributed to predominance and persistence of an 0--type character and reactivity only for holes photogenerated in the diamagnetic ZnO and TiO, samples. Photoassisted reaction of alcohol with several oxides at t.a.p.s. N 1 , i.e. with the first monolayer of those oxides, could be observed using a more sensitive mass-spectrometric technique.This also revealed incorporation of oxygen- 18 into acetone produced from propan-2-01 under these conditions, and the origins of this exchange at low t.a.p.s. are considered. Poisoning experiments employing a gas-chromatographic pulsed reactant technique at low t.a.p.s. provide evidence that photoassisted activity on ZnO invo ves photoinitiated one-electron transfer processes. The differing types and degrees of activation attainable, respectively, by photoacti- vation and by thermal activation upon surfaces of powdered samples of high-purity ( 3 99.99%) oxides of the iron-group transition metals are compared in this study using as probe molecules the secondary alcohols propan-2-01 and butan-2-01 admixed with 0,. The study is part of a programme of work on metal-oxide-catalysed elimination reactions of alc~holsl-~ which includes amongst its objectives the critical evaluation of possibilities that photoactivation may enable elimination reactions to be effected at lower temperatures and with greater selectivities than under conventional thermal activation.Realization of selective photocatalytic pathways based on such possibilities would have two requirements: (i) photoactivation of sites on the metal oxide should selectively activate them for one particular elimination reaction, as might for example occur if photoactivation were to drive dehydrogenation (- H,) along a pathway involving radical intermediates formed by photoinitiated one-electron transfer4+ at the metal-oxide surface, whereas dehydration ( - H,O) was thermally activated and proceeded solely by classical Lewis acid-base electron-pair interactions;69 (ii) in addition, conversion of alcohol to photoproducts on any photoactivated sites should be rapidly followed by removal of products from the sites, as might for example result at low temperatures via photodesorption of photoassisted products,s thereby making sites available for further conversions.This coukd allow high values of turnover accomplished per photoactivated site (taps*.) even at low temperatures. With 32973298 ELECTRON TRANSFER AT SEMICONDUCTOR INTERFACES respect to the latter point, our recent results on photoassisted oxygen-isotope exchange on zinc oxides at room temperature led to recognition of the possibility that one single-photon activation of a site involving photoinitiated localization of an electron thereon, could in favourable circumstances allow it to turn over many reaction events (ie.t.a.p.s*. % 1) before further photoactivati~n.~ Possibilities for satisfying the favourable conditions necessary for attaining such multiplier effects upon t.a.p.s*. appear lower for alcohol-metal-oxide interfaces than for ( lSO, + lS0,) ZnO interfaces. Activation of adsorbed alcohol at interfaces between u.v.-illuminated ZnO or TiO, and electrolyte solutions has been reported to involve an initial localization of a photogenerated hole to yield a radical, followed by a secondary one-electron transfer process.1o The overall effect of such processes, which lead to a doubling of the photocurrent upon addition of alcohol, has been dubbed ‘current doubling’.Detection of the latter at illuminated alcohol-metal-oxide interfaces has been taken to indicate occurrence of photoinitiated one-hole-transfer events. Alcohol conversions emanating from such localization of photogenerated holes would be limited to t.a.p.s*. < 2 per single-photon activation of a surface site. Unless products were desorbed from the active site and allowed it to be reactivated repeatedly, such processes would not satisfy a criterion suggested by Childs et d . l l for true photocatalytic activity, uiz. continuation unchanged up to t.a.p.s*. > 10. EXPERIMENTAL It has been widely reported1Y3.l1 (a) that activity of TiO, and ZnO for selective oxidation of alcohol vapours can readily be measured using gas-chromatographic procedures (with flame ionization detection) on the basis of the steady-state level of conversion to ( - H,O) or (- H,) or (C,-C,) product in the exit gases from a continuous-flow microcatalytic reactor; and (b) that the presence of molecular oxygen is essential in the gas phase if a continuing photocatalytic conversion to (- H,) etc. is to be achieved.We therefore adopted such procedures for initial comparisons of the extent to which the metal oxides listed in table 1 can activate alcohols to undergo elimination reactions under thermal- or photo-stimulation. The high-purity metal oxides (2 99.99% purity w.r.t. metal) and operation of the gas-chromatography system in both the continuous flow and pulsed reactant modes were described previously in full.3 In the present study, gaseous oxygen at p ( 0 , ) z 380 Torrt was premixed with the alcohol-vapour reactant, and surfaces of the powdered metal oxides were preoxidised by exposure to p ( 0 , ) z 380 Torr at the reaction temperature for several hours prior to introduction of oxygenated alcohol vapour.Photoactivation was achieved by exposure of the metal-oxide powder (spread as a thin layer on a fritted disc) to the output of a 125 W medium-pressure Hg arc lamp filtered through a Pyrex water-cooled jacket. RESULTS AND DISCUSSION PHOTOEFFECTS AT HIGH T.A.P.s*. Results obtained with 20 Torr of propan-2-01 as the alcohol reactant, flowing together with argon and with 380 Torr of 0, over the indicated metal oxides, are summarised in table 1, the second column of which lists any thermally assisted conversions detected at 350 K.At this relatively low temperature, thermal activation of the (alcohol+O,)-metal-oxide interfaces in the dark was apparently too weak to bring about a significant level of elimination reactions during the short contact-time allowed by the flow procedure (except for MnO, surfaces, which yielded small but continuing activity for dehydrogenation at a rate of 1.2 pmol min-l per 200 mg of MnO,). 1- 1 Torr = 101 325/760 Pa.TABLE 1 .--STEADY-STATE CATALYTIC ACTIVITIES OF METAL-OXIDE SURFACES FOR ELIMINATION REACTIONS FROM PROPAN-2-OL IN THE DARK AND UNDER U.V. ILLUMINATION at 350 K, at 475 IS, U.V. illumination surface U.V. illumination dark mrtal area oxide /m2 g-' A*(C,-C,)" A*( - H,)" ( - H d b (- H,O)b A*(C,-Cp)" A*( - H,)" A*( - HSOIa ( 1 ) (2) (3) (4) ( 5 ) (6) (7) (8) (9) TiO," 50 1.212 0.98 V,O/*e 2.6 MnOZd 27 - (0.35) Fe,03d 4.6 - - NiOd 1.9 CuOd 2.2 - - Znof 4.1 0.14 0.04 - - Cr,03d 55 - (0.03) CO,O," 3.4 - - - - 5.2 0.35 3.2 2.7 0.04 1.3 0.9 0.9 0.3 0.9962 5.5 1.4943 1.9924 0.88 0.19 2.6 3.1 - (0.046) - 0.2 - - -0.3387 0.68 - - (0.454) - (0.265) 0.007 0.30 1.225 - 0.147 - - - - - - - - - - - - Dark rates expressed as ymol min-' per 200 mg sample; photoenhancement expressed as ymol min-' per 200 mg sample; Degussa (P-25); Spex Note: for the whole experiment: p ( 0 , ) = 380 Torr; p(R0H) = 20 Torr; feed-gas flow rate = 40 cm3 min-', photon flux = 5 x 10l6 s-l at wavelengths (09.99('<); rates for V,O, are at 415 K; f NJZ-SPSOO.300-800 nm.c,3300 ELECTRON TRANSFER A T SEMICONDUCTOR INTERFACES At 350 K and in the absence of illumination none of the oxides caused detectable steady-state rates of conversion to acetaldehyde product, whose formation would formally require (C,-C,) bond cleavage. The third and fourth columns of table 1 summarise, under the headings A*(C,-C,) and A*( - H,). respectively, any new or enhanced activity for these processes detected whenever the interfaces were exposed to a flux of ca. 5 x 10l6 photon s-l at 350 K. The choice of 200 mg of oxide as the reference amount on which to base any differences between activity in the dark and under illumination was indicated by previous results3 showing that the U.V. photons had access only to that amount of sample in the photocatalytic reactor.Unequivocal evidence for significant photoassisted production of acetaldehyde as a (C,-C,) photoproduct from propan-2-01 was obtained only over the diamagnetic oxides TiO, and ZnO, whose photoactivity at 350 K yielded acetone as a (-H,) photoproduct as well as acetaldehyde. These observations were reminiscent of the occurrence of parallel pathways to (- H,) and (C,-CB) products identified in the literature as electron-transfer-initiated processes in reactions of organometallic complexes with secondary alcohols.12 Results with butan-2-01 were qualitatively very similar to those with propan-2-01 and again gave unequivocal evidence of photoactivity at 350 K only for ZnO and TiO, samples. This contrasted sharply with the complete lack of any type of photoactivity detectable by this procedure at 350 K over V,O,, Cr203, Fe,03, Co30,, NiO or CuO towards either propan-2-01 or butan-2-01 in the presence of 380 Torr of 0,.Manganese oxide emerged (cf. table 1) as the only 3d oxide which exhibited both cations with partially filled d levels and a susceptibility towards the photoenhancement of dehydrogenation of propan-2-01 or butan-2-01 at significant levels. This interesting observation should be qualified at this stage by our feeling that increases of the magnitude shown for A*( - H,) in table 1 may not be sufficiently enhanced relative to the strongly temperature-dependent thermally assisted activity of this interface in the dark as to exclude minor surface heating by the flux of photons (300 < l/nm < 800) or photodesorption of poisons such as CO, as possible sources of A*( - H,) on manganese oxides. For these reasons the value of A*( - H2) for MnO, is bracketted in table 1, Similar reservations are attached to additional bracketted values, shown for A*( - H,) for Cr,O, and Fe203 in table 1, not only because of their rather small absolute magnitude (see later results on photoreaction in the surface monolayer) but also because of their failure to disappear immediately when illumina- tion is ended.Details of the significant thermally activated dark activities of the metal oxides in table 1 towards propan-2-01 at 475 K in the absence of oxygen were recently presented elsewhere,13 together with activation energies for the (- H,) and/or (- H,O) processes. In view of indications thereby received that the relative importance of differing pathways for thermally activated elimination reactions of the alcohol could vary strongly with temperature, it became apparent that reliable values of any differences between photoactivated and thermally activated catalytic activities at 475 K [i.e.A*( - H,), A*( - H,O) or A*(C,-C,)] ideally would require: first, the determination at 475 K of the steady-state dark activity under selected pressures of alcohol and oxygen (cf. columns 5 and 6 of table 1); secondly, a determination of the extent to which illumination at 475 K altered this; and thirdly, a test of whether illumination changed the interface irreversibly by photolysis or photodesorption by checking whether the activity returned to its preillumination steady-state value upon ceasing illumination.Results from such a sequence of measurements with p(R0H) = 20 Torr and p ( 0 , ) = 380 Torr are summarised in the last three columns of table 1 . Dark activities for acetaldehyde production were furthermore below detection for all theCUNNINGHAM, HODNETT, ILYAS, LEAHY A N D TOBIN oxides studied, except for V,O,, MnO, and Cr,O,. For the latter pair of oxides at 475 K, the thermally assisted conversions corresponded to 3.5 x 1014 and 3.5 x 1015 (molecules of CH,CHO produced) rn-, s-l, respectively, and were not detectably enhanced by illumination. On the other hand, the significant enhancements of acetaldehyde product detected upon illumination of ZnO or TiO, at 475 K at the levels listed in the seventh column of table 1 confirm the existence of a photoinitiated pathway to CH,CHO, which had likewise been evident on these diamagnetic oxides at 350 K in the presence of 380 Torr of 0,.Suggestions by other workers14 that such pathways involve alkene formation, followed by oxidative scission of the newly formed C=C double bond, were examined by comparing dark and photoassisted rates of propene production at 475 K in the presence or absence of oxygen. For TiO, and ZnO in the dark, the introduction ofO, at 380 Torr increased rather than decreased the steady-state conversion to propene and did not cause the appearance of detectable acetaldehyde product (thereby arguing against thermally assisted oxidative scission involving gas-phase 0,). For illuminated (CH,),CHOH-TiO, interfaces the extent of photo- assisted conversions to (- H,), (C,-C,), and (- H,O) products were first measured in the presence of 380 Torr of O,, which was then replaced by 380 Torr of argon carrier gas.Such removal of 0, caused the (C,-C,) conversion to decline rapidly almost to zero, whilst the small conversion to propene increased slightly. Upon reintroducing 380 Torr of 0, into the gas flow over the illuminated interface the conversions to (C,-Cp) and (- H,) products increased to their original values, whilst propene declined abruptly at first, but then re-approached the earlier value with oxygenated reactant. In addition to demonstrating the oxygen-reversible nature of the photo- activity, the results established that the changes in (- H,) and (C,-C,) were an order of magnitude greater than the small and opposite changes in (- H,O) product, thereby indicating that the photoassisted oxidation of propene by gas-phase 0, could not be responsible for all the (C,-C,) yield.The results do not, however, exclude photo- activation of surface oxygen. Finally, a comparison was made of the magnitude of any changes in (C,-C,) and (- H,) and (- H,O) yields in the presence of a continuous flow of 0, whenever illumination was commenced. For zinc oxide at 475 K the observed increase in A*(C,-C,) was an order of magnitude greater than the decrease, A*( - H,O), in propene caused by commencing illumination with 0, present. The cumulative effect of these observations provides evidence for an alternative photo- assisted pathway to (C,-C,) involving neither gaseous 0, nor gas-phase propene as an intermediate.One final point to be noted from the values in column 7 is the confirmation they provide that only for the diamagnetic oxides TiO, and ZnO did the data provide unequivocal evidence for (C,-C,)-type photoactivity, there being no detectable photoenhancement of acetaldehyde product over any of the gxides featuring cations with partially filled 3d levels. A working hypothesis capable of explaining the observed restriction to ZnO and TiO, of (C,-C,) photoactivity among the 3d metal oxides investigated here can be developed from the view put forward in our earlier publications1 that photoactivity requires the close coupling of two processes, viz. (i) photoinitiated localization of a hole upon a surface oxide anion and (ii) reaction of the resultant 0--type surface site with alcohol species arriving thereon either from the gas phase or by migration across the adsorbate layer.Those earlier interpretations recognised not only the selective activity of 0--type species towards secondary and primary alcohols, but also the difficulty of distinguishing this from reactivity of the protonated analogues, i.e. surface hydroxy radicals. Similar uncertainties persist in the present work concerning 0- and OH, since the latter may result either from photoinitiated localization of holes on pre-existing OH- surface groups or by proton-transfer to 0- from alcohol. For convenience the 33013302 ELECTRON TRANSFER AT SEMICONDUCTOR INTERFACES following interpretation is given in terms of 0--type intermediates, but OH is not excluded. Earlier workers reported that the substitution of Ti4+ cations in oxides by cations with incomplete 3d shells inhibited photoinduced charge transfer at the surface.15 In attempting to explain the similar sharp decline in photocatalytic activity here reported when switching from ZnO or TiO, to metal oxides with incomplete 3d shells, the following aspects of points (i) and (ii) merit particular attention: the relative ease with which photons may generate 0--type surface species, and the relative susceptibility of any such photogenerated 0--species to rapid degradation into other surface species lacking the specific reactivity of 0- towards alcohols.Since recent studies on Ti0,16 and ZnOl7? l8 single crystals show an absence of states in the band-gap, photogenerated holes in TiO, and ZnO relate predominantly to energy bands derived from 2p orbitals of the oxygen-anion sublattice, and correspond to itinerant 0--type electronic arrangements.(This can apply whether the holes are free or are coupled as hole-electron pairs.) However, for the other oxides examined in this study the wavelengths present in the photon flux used (300 2 R/nm 6 800) were more likely to result in the photogeneration of excited states relating to orbitals derived from the cationslg than to 0--type species (e.g. in Fe,O, photons of energy ca. 2.0eV can produce Fe4+-Fe2+ cation pairs, but charge-transfer from 02- to Fe3+ requires 5.7 eV). Even if photoexcitation of some of the other oxides initially produced 0--type configurations, rapid trapping of such holes by occupied states lying in the band-gap and derived from 3d levels of the variable-valency cations should rapidly convert the holes to electronic configurations differing from 0- and lacking its selective reactivity towards alcohols. PHOTOEFFECTS AT LOW T.A.P.s*.In an effort to achieve higher sensitivity for detection of photoinitiated oxidation of propan-2-01 over the various metal oxides, some experiments were carried out under low pressures (0.1 Torr) of 1802 and (CH,),CHOH and with mass-spectrometric (m.s.) rather than gas-chromatographic (g.c.) detection. Such experiments demonstrated an unmistakable photoassisted rate of oxidation of alcohol at room temperature whenever Cr,O,(red), Fe,O,(red), ZnO and TiO, were simultaneously exposed to U.V.illumination, alcohol vapour and 180,. Furthermore, the acetone photoproduct from these systems was found to be enriched to 30+10% in oxygen-18. The apparent contradiction between this unequivocal m.s. evidence of photoassisted ( - H,) reaction over Cr203 and Fe,O, and the absence of any such evidence from g.c. studies is believed to have its origins in the differing experimental conditions for the two sets of experiments. Thus the continuous-flow g.c. experiments were weighted in favour of the detection of photoproducts arising from truly photocatalytic processes i.e. those capable of yielding photoproducts equivalent to a high turnover per surface site. The m.s. experiments under low pressures in a static reactor were, on the other hand, weighted in favour of the detection of photoproducts arising from photoinitiated reactions with surface monolayers of the metal oxides.Similar observations of photoassisted ( - H,) reaction to sub-monolayer equivalence have previously been observed by infrared spectroscopy for propan-2-01 on ZnO and Al,03.20 Such obser- vations need not imply a photocatalytic process continuing to high t.a.p.s*. Thus the former could proceed via direct photoactivation of surface ions and a once-only reaction with alcohol + 0,, whereas the latter imply hole formation in sub-surface regions and repeated activation of surface sites following their migration to the surface. Blank experiments made by introducing an admixture of laO, + (CH3),C0 over the metal-oxide surfaces followed by illumination established that the observed extents of incorporation of oxyen-18 into the acetone product from propan-2-01 could notCUNNINGHAM, HODNETT, ILYAS, LEAHY AND TOBIN 3303 have arisen by either dark or photoassisted oxygen isotope exchange between 1802 and acetone after its production.Furthermore, the yield of acetone appeared too large in some cases (such as Fe,03, which yielded no detectable propene in the dark or under illumination in the absence of 0,) to be explicable in full in terms of any photoassisted attack by l80, up-on an alkene intermediate formed by a ( - H,O) process. Consequently, the probable existence of some other photoassisted pathway to (CH3),C180 in the presence of l80,(g) could be inferred from these low-turnover results.Experiments were made to test the possibility that this alternative mechanism might involve, initially, the photoassisted formation of H,180 as in eqn (1 a) (which itself is a multistep process involving 0-) hv MO (CH3),CHZ60H ++1802 + (CH3)2C160 + H2180 (1 a) and that oxygen-1 8 subsequently became scrambled between acetone and water via ( 1 b ) in the photoactivated reaction monolayer. In order to test this alternative mechanism, rates of oxygen- 18 exchange were monitored when (CH3),C160 plus water enriched to 38 % in oxygen- 18 were introduced over metal oxide samples at room temperature. The samples had been pretreated in vacuo or in oxygen in a fashion identical to those employed to study alcohol reactions, but they had not been exposed to alcohol.Upon illumination of these [(CH3),C160 + H,180]-metal-oxide interfaces by the same photon flux as employed to study alcohol+180, systems, the fast scrambling of oxygen-18 between acetone and H,O was observed over ZnO and Cr,03. Our experimental observations that photodehydrogenation of propan-2-01 to acetone was always accompanied by 180-incorporation in low-turnover conditions can thus be accounted for by an important role being played by a mechanism involving reactions ( l a ) and (1 b). Furthermore, the importance of reactions ( 1 a) plus ( 1 b) also allows a reinterpretation of our earlier observations2 on the complete loss of the --180H label from 2-methylpropan-2-01 when converted to acetone upon illumination over preoxidised TiO, in the presence of 180,.Process (1 a) in those conditions would lead to the loss of the l80 label from acetone after its formation from 2-methylpropan-2-01 by photoassisted (C,-C,) bond cleavage. In the particular instance of zinc oxide, some support for the model just summarised for photoactivity was derived earlier from demonstrations that creation of additional 0--type surface species (from N,O + e- -+ N, + 0-) enhanced elimination reactions at (alcohol + N,O)-ZnO interfaces under illumination.21 Conversely, inhibiting effects upon ZnO photoactivity were to be expected from incorporation of hole-trapping species capable either of reacting more rapidly than alcohol with 0--type holes or of directly capturing photogenerated holes more efficiently than surface 0,- species.Surface species capable of promoting hole-electron recombination may also lower [O-I. The following data describe inhibition of ZnO photoactivity by species selected for their likely ability to affect [O-] by one-electron-transfer events at the surface. In this context pyrene and tetracyanoethylene (TCNE) were selected as adsorbates likely to interfere with photoactivation of 0,- via one-hole-transfer interactions. The g.c. technique was modified3 to allow for: (i) the introduction of individual pulses of alcohol vapour plus oxygen over a clean, illuminated ZnO surface in order to characterise its photoactivity towards several successive pulses; and (ii) the adsorption of pyrene or TCNE or other adsorbates onto the ZnO surfaces followed by characterization of the photoactivity of these ‘poisoned ’ ZnO surfaces towards successive pulses of [(CH,),CHOH +O,].Fig. 1 (a) illustrates the strongly inhibiting (CH3),C160 + H,180 e (CH,),C180 + H2160 ( 1 b)3 304 ELECTRON TRANSFER AT SEMICONDUCTOR INTERFACES 0.5 n -@ g)----QD--@- - g G * -. -8- -8 I 1 I I // I / / 2 4 9 pulse no. O t @ 0- ------o--- influence of preadsorbed pyrene or TCNE upon the rate of production of acetone over the illuminated ZnO surfaces, with comparable large decreases resulting from the preadsorption of pyrene from its vapour and of TCNE from its vapour or from benzene solution. Table 2 summarises, in the form of percentage activities of the poisoned surfaces relative to activity over an illuminated clean ZnO control, theseCUNNINGHAM, HODNETT, ILYAS, LEAHY A N D TOBIN 3305 TABLE 2.-RELATIVE ACTIVITIES OF CLEAN AND POISONED (CH,),CHOH-ZnO INTERFACES FOR THE INDICATED *PHOTOASSISTED PROCESSES AT 350 Ka AND FOR ALCOHOL ADSORPTION IN THE DARK^ [(CH,)zCHOH(ads-)] (%) [-(Hz)*lC [ -(Co,-C,)*ld [-(HZ0)*le poison (%) (%) (%> pulse 1 other none PYrene(", TCNE(", TCNE 6 (ex.C6H6) CH,COOH,,, C6H50H anisole C6H5NH2 100 10 < 5 < 5 0 0 20 20 100 32 5 5 0 14 39 47 100 5 0 0 0 8 30 10 - 100 95 80& 10 40 40 - 75 0 0 0 0 100 60 & 20 100 60 & 20 a Photoactivities assessed from the extent of conversion averaged over the first four reactant pulses of [(CH,),CHOH+OJ passed over the illuminated ZnO surface at 350 K; extent of adsorption assessed from amount of alcohol retained by ZnO from pulse 1 of alcohol or from other pulses (see text) ; photodehydrogenation activity expressed as percentage of the higher continuing activity noted over a clean ZnO sample in similar conditions, viz.photoassisted (C,-C,) bond cleavage expressed relative to rate of 4.4 x mol CH,CHO per min over clean ZnO or, photodehydration expressed relative to rate of 4 x lo-'" mol C,H6 per min over clean ZnO. inhibiting effects upon (- HJ-type activity and also upon (- H,O)* and (C,-C,) activities at the illuminated interfaces. Activities in each case were the average of the first four reactant pulses. Whilst those data for pyrene and TCNE adsorbates do show inhibiting effects which could be consistent with the interference of these adsorbates in one-electron-transfer processes, other data in table 2 for the effects of other classes of adsorbates demonstrate that other modes of action of the adsorbates should also be taken into account.Thus acidic preadsorbates such as acetic acid or phenol, which may interact with the surface as Bronsted or Lewis acids, may be seen from table 2 to have caused strong inhibition. However, the probable method of interference of those acidic preadsorbates could be shown, by appropriate use of the reactant pulse technique,13 to be via inhibition of irreversible propan-2-01 adsorption [cf. fig. 1 (b)]. Such effectively complete inhibition of alcohol adsorption by these acidic preadsorbates may be contrasted with effects of the one-electron-transfer preadsorbates upon alcohol adsorption. The levels to which irreversible propan-2-01 adsorption was reduced by preadsorption of TCNE (40 %), or of pyrene (95 %), would be quite insufficient (if only Langmuir-type adsorption processes occurred) to explain the much larger percentage inhibitions those preadsorbates exerted on photoactivity (cf.table 2). Basic pread- sorbates, such as pyridine or anisole, which may interact with the ZnO surface as Bronsted or Lewis bases, were notably less effective than acidic adsorbates or TCNE in inhibiting photoactivity towards propan-2-01 or adsorptive capacity (cf. table 2). However, our additional observations of an absence of any significant inhibiting effect upon alcohol adsorption in theJirst pulse following exposure to the basic adsorbates, and the contrasting evidence for some inhibition of alcohol adsorption in subsequent pulses, points towards more complicated surface interactions in these cases (e.g.possibly a base-catalysed conversion of alcohol with an associated induction period). 107 F A R 783306 ELECTRON TRANSFER A T SEMICONDUCTOR INTERFACES Further work will be necessary if these more complex interactions involving bases are to be understood. However, the present results with likely one-electron-transfer adsorbates provide some new indirect support for the involvement of one-electron- transfer processes in the photoactivity of [(CH,),CHOH + 0,) - ZnO*. This in turn is consistent with the view put forward here that the especially high and often observed photoactivity of ZnO (and TiO,) may be understood in terms of persistent 0--type character of photogenerated holes made possible by the absence of occupied states related to d orbitals in the band-gap.Maintenance Grants from the Irish Government (to E. M. L., J. P. T. and B. K. H.), and University College Cork (M. I.) are gratefully acknowledged. J. Cunningham, B. K. Hodnett and A. Walker, Proc. R. i r . Acad. (Centenary Issue), 1977, 41 1 . J. Cunningham, B. Doyle, and E. M. Leahy, J . Chem. SOC., Faraday Trans. I , 1970, 75, 2009. J. Cunningham and B. K. Hodnett, J . Chem. Soc., Faraday Trans. 1, 1981, 77, 2777. D. Rehorek, Z . Anorg. Allg. Chem., 1978,443,255; D. Rehorek, M. Benedix and Ph. Thomas, Znorg. Chim. Acta, 1977, 25, L100. D. Greatorex and T. J. Kemp, Trans. Faraday SOC., 1971,67,56; D. Greatorex, R. J. Hull, T. J. Kemp and T. J. Stone, J . Chem. SOC., Faraday Trans I , 1974, 70, 216; 1972, 63, 2059. H. Noller and K. Thomke, J . Mol. Catal., 1979, 6, 375. H. Vinek, 2. Phys. Chem. (N.F.), 1980, 120, 119. J. Cunningham, E. Finn and N. Samman, Faraday Discuss. Chem. SOC., 1974,58,160; J. Cunningham and P. Meriadeau, J. Chem. SOC., Faraday Trans. I , 1976, 72, 1499. J. Cunningham, E. L. Goold and J. L. G. Fierro, J. Chem. SOC., Faraday Trans I , 1982, 78, 785. lo W. P. Gomes, T. Freund and S. R. Morrison, J. Electrochem. SOC., 1968, 115, 818. l1 L. P. Childs and D. F. Ollis, J. Catal., 1981, 67, 35; 1980, 66, 383. I 2 D. Kochi, Organometallic Mechanisms and Catalysis (Academic Press, New York, 1978), p. 107. l3 J. Cunningham, B. K. Hodnett, M. Ilyas, J. Tobin, E. L. Leahy and J. L. G. Fierro, Faraday Discuss. l4 A. Walker, M. Formenti, P. Merideau and S. J. Teicher, J. Catal., 1977, 50, 237. l5 R. D. Rauh, J. M. Buzby, T. F. Reise and S. A. Alkaitis, J. Phys. Chem., 1979, 83, 2221. l6 R. H. Tait and R. V. Kasowski, Phys. Rev. B, 1970, 20, 5168 and 5178. l8 I. Ivanov and J. Pollman, J . Vac. Sci. Technol., 1981, 19, 344; Solid State Commun., 1980, 36, 361. 2o V. N. Filimonov, Dokl. Akad. Nauk SSSR, 1964, 158, 1408. 21 J. Cunningham, E. L. Goold and D. J. Morrissey, J . Catal., 1978, 53, 68. Chem. SOC., 1981, 72, 283. W. Gopel, Ber. Bunsenges. Phys. Chem., 1978, 82, 745, J. B. Goodenough, Prog. Solid State Chem., 1971, 5. (PAPER 2/147)

ISSN:0300-9599    

DOI:10.1039/F19827803297

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 3307-3317 Some Aspects of Stereoselectivities and Kinetics in the Ring-opening Polymerization of Norbornenes using Met at hesis Catalysts The Nature of the Metallacarbene Intermediates BY Ho Huu THOI, BOREDDY S. R. REDDY AND JOHN J. ROONEY* Department of Chemistry, The Queen’s University, Belfast BT9 5AG, Northern Ireland Received 16th February, 1982 Stereoselectivities in the ring-opening polymerization of bicyclo[2 .2. l jhept-2-ene and several 5,5- disubstituted derivatives have been extensively investigated using numerous metathesis catalysts at different temperatures and concentrations of monomer. Some studies of the kinetics of polymerization are also briefly reported. For many Group VI catalysts at moderate to high concentrations of monomer a limiting value of ca.50% for the cis content of the polymers is observed. This feature is related to the observation that under these conditions the kinetics of polymerization are zero order with respect to monomer, and is explained by the postulate that propagation is due solely to two mirror-image forms of metallacarbene which may coordinate monomer with equal facility in either a cis or rrans orientation. However, with Group VI catalysts at high dilution of monomer, or with noble-metal catalysts at moderate concentrations, polynorbornenes having a high trans content are obtained. A head-tail structural bias, which increases with dilution, also becomes evident in polymers of the 5,Sdisubstituted derivatives, especially when the substituents are polar and the catalysts are IrCl, or OsC1,.These results are discussed in terms of Michaelis-Menten kinetic theory with the postulate that the chiral metallacarbenes relax to a more stable symmetrical form. The intermediates described and the kinetic theory developed are used to explain several other important aspects of stereospecificities and selectivities in olefin metathesis and ring-opening polymerization. While selectivities and stereoselectivities in olefin metathesis and ring-opening polymerization of cycloalkenes have been frequently studied, comparatively little attention has been given to the kinetics of these processes. For example, stereoselec- tivities in n-alk-2-ene metathesis are widely discussed2 in terms of the energetics of distinct shapes of puckered metallacycles derived from one type of metallacarbene.The distinction in the metallacycles is supposed to depend on the orientations, axial or equatorial, of substituents in cis or trans 1,3-positions. Recently we have questioned3 this postulate, suggesting instead that there may be different types of metallacarbene with distinct kinetic behaviour, the difference arising if the last-formed double bond is cis, because in that case it may still be coordinated to the metal centre while the next propagation step commences. This theory is strongly supported by evidence obtained from 13C n.m.r. spectra of ring-opened polymers, but is obviously applicable to metathesis of n-alk-2-enes as well. The purpose of the present paper is to show that many other aspects of stereosel- ectivities in ring-opening polymerization are also best discussed in terms of kinetically distinct metallacarbenes and Michaelis-Menten kinetic theory. The possible similarity to enzyme systems in this respect was first alluded to as early as 1970 by hug he^,^ who derived kinetic equations based on the now discarded quasi-cyclobutane mechanism.The kinetics of ring-opening polymerization of norbornene (NBE) 3307 107-23308 POLYMERIZATION OF NORBORNENES catalysed by some Ru complexes have recently been inve~tigated,~ but it was erroneously concluded that they were zero order in monomer, whereas the data clearly show that they are first order. Following hug he^,^ a Michaelis-Menten type equation for the rate of propagation, R, = k, K[M] [cat]/( 1 + K[M]), where [MI is the concen- tration of monomer, was also developed but an equilibrium constant, K , for complexation of monomer was used instead of the correct Michaelis-Menten constant.This is a common error which becomes important if the overall reaction is zero order with respect to monomer. During the past five years we have extensively investigated cis contents ( o , ) , ~ cisltrans blockiness ( k , k,)’ and tacticities3 of ring-opened polymers of cyclic olefins made using numerous metathesis catalysts under a variety of conditions. Tacticities were determined from the 13C n.m.r. spectra of polymers 2a made from racemic and partially resolved 5,5-dimethylbicyclo[2.2. llhept-2-ene (DMNBE), 1 a. Such spectra show four distinct resonances for the tail-head (TH), tail-tail (TT), head-tail (HT) and head-head (HH) C atoms in both the cis and trans junctions. In most cases there 1 ta T H HT TT TT H H HH ( m dyads only shown ) is no bias towards either HT(TH) or HH, TT structures in 2a, but for certain catalysts and conditions which promote the formation of high-trans polymer a slight bias towards HTunits i~clearlyevident.~ We now report further aspects ofthis stereoselective feature together with some significant information about the kinetics of propagation and novel relevant data on oc values for polynorbornenes. The major aim is to delineate further the different types of metallacarbene intermediates involved, the kinetic conditions under which they are important and their stereoselective characteristics.EXPERIMENTAL AND RESULTS Details of polymerization procedures and systems, methods of extraction of polymer samples and 13C n.m.r.analysis have been given el~ewhere.~’ The rates of polymerization of DMNBE at 353 K in CDCl, initiated by Ph(MeO)C=W(CO), were followed by lH n.m.r. using spectral intensities of the monomer relative to those of CH,Cl, which was added to provide an internal standard. The rate was essentially independent of [MI ([MI = 1.78-0.37 mol dm-3) over the first 20-30% conversion (RP = 1.05f0.13 x loM4 mol dm-3 s-l, o, = 0.55), but then suddenly fell off (fig. 1). Zero-order conversion of monomer was also foundHO HUU THOI, B. S. R. REDDY A N D J. J. ROONEY 0.6- 0.4 0, 3309 A n 0 0 A - 3 0 20 4 0 60 80 time/min FIG. 1 .-Concentration of monomer [MI plotted against time for various initial concentrations [MI,: (a) 1.78, (b) 0.81 and (c) 0.37 mol dm-, in degassed CDCI, solution at 353 K.ph(MeO)C=W(CO),] = 0.26 mol dmP3. 0 0 1 2 3 4 5 6 [Ml,/mol dm-3 FIG. 2.-Effect of initial monomer concentration on oE for ring-opened polymers of nor- bornene. Conversion 20-30%. [WCl,] = 5 x rnol dm-, at 293 K in chlorobenzene solvent. [IrCl,] = 4 x lop2 mol dm-, at 348 K in chlorobenzene/ethanol (5/1). 0, WCl,/Ph,Sn (1/2); A, WCl,/Me,Sn ( l / l ) ; 0, IrCl,.3310 POLYMERIZATION OF NORBORNENES TABLE 1.-Cis CONTENTS IN RING-OPENED POLYMERS OF NORBORNENE~ MADE AT 288 K ~ catalyst g c rt rc WCl, WCl,/Me,Sn (1 /2) WC16/Ph4Sn (1 /2) WCl,/EtAlCI, (1 /4) MoCl, MoC1,/Me4Sn (1 /2) MoCl,/Ph,Sn (1 /2) Ph( MeO)C=W(CO),b MesW(CO),/E/EtAlCl, (1 / 1 / l)c MesMo(CO),/E/EtAlCl, ( l / l / l ) MesCr(CO),/E/EtAICl, (1 / 1 / 1) MesW(CO),/AlCl, ( I /20) ReCl,/EtAICl, (1 /4) MesW(CO),/AlCI, (1 /20)d MesW(CO),/AlCl, (1 /20)e 0.55 0.55 0.56 0.5 0.45 0.46 0.47 0.6 0.46 0.42 0.52 0.42 0.55 0.50 0.55 1.7 2.0 2.1 I .4 1.2 1.6 1.6 1.5 - - a [MI was in the range 1.5-4.0 mol dm-,; the temperature was 353 K; Mes is reactant was cyclopentene; reactant was mesitylene and E is norbornene epoxide; cyclooctene.TABLE 2.-STEREOSELECTIVITES IN THE RING-OPENING POLYMERIZATION OF SOME 5,5-DISUBSTITUTED NORBORNENES [MIU monomer /mol dm-, catalyst TH/TT l a l a l a l c l a l a l a l a l a l c l b I d 3.9 1.6 2.0 1.1 3.3 0.54 2.0 0.4 0.4 1.6 1.6 1.5 ReC1, Ph(MeO)C=W(CO), Ph(MeO)C=W(CO), RuC1, RuC1, IrCl, IrCl,b OsCl, RuC1, RuC1, RuCl, WCl,/Me,Sn (1 / 1) 288 373 288 328 323 323 348 348 348 333 333 333 1 .o 1 .o 0.65 1 .o 0.9 0.7 1 .o 0.9 0.86 0.9 1.1 < 0.05 1.1 0.0 1.2 0.26 1.1 1.1 - 3.0 > 3.0 0.22 1.3 1.2 < 0.05 1.5 0.0 1.5 0.0 1.6 - - - - - - - a Using the noble-metal catalysts the solvents for l a and for l b to Id were, respectively, a 1/5 and a 1/1 mixture (vol./vol.) of chlorobenzene and ethanol.Chlorobenzene alone was used for the other catalysts, except for 1 c and Ph(MeO)C=W(CO),, where the solvent was deuterochloroform. The spectrum was complicated by additional signals due to end groups but the TH/TT ratios were clearly of the size indicated.HO HUU THOI, B. S. R. (REDDY A N D J. J. ROONEY 331 1 for the polymerization of cyclopentene ([MI = 5.4 mol dm-3) at 273 K using the very reactive WCl,/Ph,Sn catalyst in chlorobenzene solvent with analysis by a g.1.c.method to be described el~ewhere.~ A similar abrupt termination after 40-50% conversion was observed, but when the reaction was followed at 3 13 K zero-order behaviour was still noted without any sign of a fall-off in rate up to ca. 75% conversion. The reason for these abrupt terminations is not known and is still being investigated. The effects of initial concentration of monomer on ac values for NBE polymerization catalysed by WCl,/Ph,Sn and WCl,/Me Sn at 293 K and IrCl, at 348 K are further shown in fig. 2. In the higher range, 1-4 mol dm-3, oc is independent of [MI but in the lower range, 0.26-0.1 mol dmP3, ac decreases rapidly with dilution for the WCl,-baSed systems.The value of ac N 0.5 is obviously a limit since it also remained essentially constant for WCl,/Ph,Sn in the range 204-403 K.* Similar values (a, 1: 0.5i were observed using a variety of other catalyst systems at 288 K and they are recorded in table 1. The low rtrc values show that the polymers have essentially a random c5ltran.s distribution.' Polymers were made from 1 a, 1 b, 1 c and 1 d using several catalysts, and ac values together with TH/TT ratios for cis and trans junctions are given in table 2. Three features are noteworthy. (i) The TH/TT ratios are unity when ac is high (non- noble-metal catalysts). (ii) A slight bias in favour of HT units is clearly evident with the noble-metal catalyst (low ac values) and becomes more pronounced at high dilution.(iii) The bias is stronger for monomers with polar substituents. DISCUSSION While the interconversion of olefin-metallacarbene and metallacyclobutane com- plexes is now accepted as the mechanism of olefin metathesis, the precise nature of these species in highly active catalysts is still unknown. We will take the unconventional view here, and justify it later, that a metallacyclobutane proper may only be present as a transition state. Furthermore, for highly reactive olefins, e.g. NBE, it seems likely that the appropriate orbitals of Mt=C and C=C are already engaging each other to form a quasi-metallacyclobutane in the monomer-complexation step. These ideas are summarized in scheme 1 , where [Mt] is the metal ion with permanent ligands and I [ M t]=CH' 6' J P,CH 3 4 5 6 3' SCHEME 1 P, is the polymer chain; the essence of scheme 1, as compared with the analogous scheme 1 previously p~stulated,~ is that there are no longer separate and independent monomer-coordination and metallocyclobutane-forming steps, i.e.a complex with discreet carbene and olefin ligands may often never be realized. This may seem surprising but it is worth remembering that metathesis catalysts are very similar to those for Ziegler-Natta polymerization of a-olefins, where the metal ions involved are knownlO to have very little tendency to coordinate ethylene in comparison with metal * For further details see fig. 5 in ref. (3).3312 POLYMERIZATION OF NORBORNENES ions such as Ag+ and those from Group VIII. The difference is ascribed to the importance of dz 4 p K * back-bonding when the d orbitals are extensively occupied; without this bonding component simple olefins are only very weak bases.Two propagating sequences are now envisaged, one of which involves interaction of monomer with the metallacarbene to give the quasi-metallacyclobutane (3 + M -+ 4) followed by bond switching in 4 (4 -+ 5 ) with rapid orbital disengagement (5 -+ 6), and finally decoordination of the newly formed double bond (6 -+ 3’). The metallacyclo- butane is now only a transition state in step 4 -+ 5. It is likely that in many systems 6 is also a transition state in the conversion of 5 to 3’, but we have retained 6 3 6’ in scheme 1 because there is evidence that for some catalysts the alternative sequence which consists only of 6’+ M + 4 -+ 5 --+ 6 is important where displacement of the newly formed junction by monomer is postulated, but only when the double bond is cis.Furthermore, the entry of monomer into 6’ is believed to be such that the orientation which again gives a cis junction is strongly fav~ured.~ The division of 6 into two general types, Pc and pt, where the subscripts define whether the newly formed double bond is cis or trans, respectively, is a necessary distinction, even though pt is never regarded as a propagating species, in contrast to Pc. Also three different forms of each of these, a mirror-irlage pair and a kinetically symmetrical form, have to be considered in order to explain the results of the tacticity s t ~ d i e s . ~ A good example which illustrates this important conclusion is the hexacoordinated W-complex which seems to be the propagating speciesf1 when Casey’s compound, Ph,C=W(CO),, is used as catalyst.The following octahedral geometries seem reasonable for the mirror-image forms and the kinetically symmetrical form, respectively. Several forms of 3 = 3’ also seem to be important. Thus as the newly formed double bond decoordinates in 6 the corresponding pentacoordinated species immediately obtained, P, can also be regarded as essentially octahedral with a vacancy 0. The following three geometries are also envisaged for P H P\$ II/ ’I C -Mt--O P H \ $ II/ ’I C -Mt--O which may then isomerize to more stable forms, P,, where square pyramidal or trigonal bipyramidal structures of the following types seem to be reasonable approximations. H P \ / C II / -yt\ The above discussion by no means exhausts the symmetry possibilities because in the hexacoordinated/pentacoordinated manifold all the permanent ligands, e.g.C1-HO H U U THOI, B. S. R. REDDY A N D J. J. ROONEY 3313 or CO, may not be identical, and corresponding pentacoordinated/tetracoordinated geometries would have to be considered for propagating species such as >C=WOCl,. However, for the purposes of the present paper and for any given catalyst system it is only necessary to consider three kinetically distinct propagating species P,, P and P,, and two relaxation processes, P, + P and P -+ P,, as shown in scheme 2, in order to account for the results over the full range of IT^ values (1 .O-0.0).By definition, all p, species must relax to P. Since each catalytic cycle consists of a series of steps, as detailed in scheme 1, the rate constants in scheme 2 are not those for propagation k2c Pc + M - pc f k1 I + P + M ’ L P, k4 / p, + M--!J%+ PC SCHEME 2 but in subsequent discussion will be regarded as rate constants for the monomer- complexation steps. The kinetics of propagation can be considered in terms of Michaelis-Menten theory which, when applied to scheme 1 (3 + M + 4 -+ 9, leads to three distinct limits as follows. Type (1): Step 4 -+ 3+ M is negligible compared with 4 -+ 5, which is the rate-controlling step, so polymerization is zero order in [MI. Type (2): Step 4 -+ 3+M is also negligible compared with 4 + 5, but with dilution 3 + M -+ 4 becomes slower than 4 -+ 5 and first-order behaviour in [MI is eventually obtained.Type (3): Step 4 -+ 3+M is always significant and in the limit is at equilibrium with the complexation step, such that the rate of polymerization is first order with respect to monomer. One extreme behaviour predicted by scheme 2 is that there is no relaxation of any sort so propagation is confined to the sequence 6’+M -+ 4 + 5 + 6, giving all-cis polymer, as found for NBE or DMNBE at room temperature using ReCI, as ~atalyst.~ In the range 1.0-0.5 the IT, values and cisltrans blockiness’ are then governed by the effect of variables such as temperature and monomer concentration on the ratio k2 [M]/I?,.~ Blockiness demands participation of at least two kinetically distinct propagating species, e.g.P, and P. A second important limiting case, often noted for the most active catalysts, is when IT, 21 0.5, even at 204 K for NBE polymerization using WCl,/Ph,Sn, a fact which supports the view that 6 is now only a transition state and not an intermediate. Furthermore, the essentially random nature of the polymers (table 1) confirms that there is only one type of propagating species which must be assigned as P. The 0, values can then be explained by Michaelis-Menten theory as applied to eqn (1) and (2), where [PM] denotes 4 and P is 3. (2) Steady-state analysis shows that the competition ratio, R,,,, for the formation of cis and trans junctions is given by 4, k2t P+ M - [PM], + Pt. k-It (3) klCk2, 1 kltk2t (k-1, + k 2 d V - l t +k2t) ’ Re,, =3314 POLYMERIZATION OF NORBORNENES When the kinetics are type (1) (cf. fig.l), kWlc and k-,, are negligible compared, respectively, with k,, and k2t, so eqn (3) reduces to R,,, = klCIklt. (4) The temperature independence ofo, and the values (ca. 0.5) are then in accord with the conclusion that for a highly reactive olefin such as NBE monomer complexation with P is non-activated and therefore random, in contrast to the analogous reaction of M with P, using ReCl,. When [MI is decreased and the kinetics are type (2), eqn (4) and therefore o, II 0.5 should still hold if dilution of monomer has no other consequences. However, the data in fig. 2 show that at low [MI o, rapidly decreases. We believe that the reason for this change is that the metallacarbenes P are now isomerizing to the more stable and therefore less reactive form P,.Complexation is now an activated process and because [PM], is lower in energy than [PM],, as a consequence of steric factors associated with substituents on the carbene C atom and on the double bond in M as 4 is formed, P, is trans directing with El, > El, > 0 (cf. El, z El, II 0 for P, but El, < El, > O for P,). El, and El, are the activation energies for monomer complexation in a cis and trans orientation, respectively. A steady-state analysis of the relevant part of scheme 2, neglecting the minor step (5c), gives which predicts that o, may change with dilution from independence of to first-order dependence on [MI, as found for the WC1,-based catalysts (fig. 2). The noble-metal catalysts tend to give polymers with rather low o, values, and the kinetics for NBE polymerization using RuCl, are first order5 even at [MI = 5.0 mol dm-3.Therefore it is quite possible that this is an example of type (3) behaviour, in which case k-l, % k,, and k-,, % k,,, so that eqn (3) becomes Rclt = KCk,,/Ktk,t (6) where K, and K, are, respectively, klc/k-lc and klt/k-lt, the equilibrium constants for monomer complexation. K, < Kt should hold since [PM], is higher in energy than [PM],, so o, < 0.5 is expected by this analysis (k,, e k,, should be true). However, type (2) kinetic behaviour even at high monomer concentrations may be more likely with the noble metals since rapid decomplexation of NBE may never be feasible. This would imply that complexation is activated (3+M + 4 slower than 4 + 5), so El, > El, > 0 is reasonable for P and therefore its trans directing character.In any event, first-order rates in [MI imply that the trans directing P, should be a significant propagating species, and indeed the fall in o, with dilution of NBE noted for the IrCl, catalyst (fig. 2) supports its intervention. The increasing loss of tacticity in polymers of DMNBE with monomer dilution using RuC1, as catalyst had previously led to the same concl~sion.~ We also know from the tacticity studies3 that P, is not involved in Ph(MeO)C=W(CO), catalysed polymerization of DMNBE, where the rates are zero order in [MI (fig. 1) and o, = 0.55. In this case P seems to be the dominant carrier and in agreement with the kinetic theory should not, and does not, relax to P,.HEAD-TAIL STRUCTURAL BIAS I N POLYMERS OF 1 A head-tail structural bias is never observed for ring-opened polymers of 1 a when P, and P are the likely propagating species3 (the substituents act merely as a label), but a slight bias (table 2) is noticed, especially with noble-metal catalysts at highHO HUU THOI, B. S. R. REDDY A N D J. J. ROONEY 3315 dilution of monomer where P, is believed to be a significant carrier. In order to confirm this theory, DMNBE was polymerized using some of the W catalysts described in table I , but with [MI decreased by a factor of ca. 10 in order to bring it within the range (cf. fig. 2) where it is expected that P, should be a significant chain carrier. In agreement with prediction, oc decreased substantially and a head-tail bias in the junctions became evident.12 This is a key observation in conjunction with the type (2) kinetic behaviour expected of such W catalysts at high dilution of M.Thus even cyclopentene polymerization (less strongly complexing monomer than NBEY is zero order in [MI at higher concentrations using the WCl,/Ph,Sn catalyst. Clearly if monomer coordination and metallacycle formation were independent consecutive steps, and the reversibility of the former is negligible, the head-tail bias, which must be due to partial enantiomeric selection,12 would have to be entirely associated with discrimination between the enantiomers in the coordination step (addition is confined to the exo face of norbornenes). This is where scheme 1 has great advantages because all the polar and steric factors required for such discrimination are immediately operative in the complexation step (3+ M -+ 4).Since P, has either a head or tail structure with respect to the positions of the 5,Sgroups and the metal, the formation of the different types of junction, HH, HT, TH and TT, for one type of double bond, e.g. trans, can be considered in terms of the following set of equations, where MH and MT are enantiomers: An analogous set of equations can be written for cis junctions, and a head-tail bias will develop if PH is the dominant carrier with kHT > kHH, or if P, is dominant and kTH > kTT. The reason why P, shows such a bias, but not P, can again be attributed to the fact that P, is the more stable isomeric form and this increases the energy of activation for the step (3+ M + 4) to the point where P, becomes enantioselective as well as trans directing.While the discrimination against the cis orientation is attributed to unfavourable cis- 1,2-~ubstituent interaction in the appropriate complex [PM],, the head-tail bias seems to be determined mainly by the polarities of the Mt=C bond and the C=C bond in the monomer, in agreement with the facts that the bias is more evident for polar substituents in 1 (table 2) and is more pronounced for 0 s and Ir than for the Ru catalyst although the latter is more trans directing at the same dilution. There is good evidence13 that Mt=C is the direction of polarization associated with high metathesis activity, and such polarization is expected to increase down a triad of metals in the transition series as a consequence of the dramatic increase in ligand field effects on the dn orbitals and thus on the dn-pn bonding component in the Mt=C bond.The same theory has been usedl to explain the trends in the ratios of degenerate to productive metathesis of n-alk-1 -enes using Group VI metal catalysts. 6- 6+3316 POLYMERIZATION OF NORBORNENES - If the polarized Mt=C bond is represented by the canonical form Mt-C it is also easy to see by analogy with carbonium-ion chemistry why scheme 1 is also mechanistically realistic (scheme 3). Indeed mechanisms of this type are necessary for olefin metathesis catalysed by All4 and P15 complexes. TC-i c+ c c-c [ M t l - C [ M t l - + C [Mtl-C SCHEME 3 GENERAL ASPECTS For high cis directing catalysts monomer complexation is an activated process.It is not surprising therefore that such catalysts are highly selective in copolymerizationl6? l7 and poor at cross-metathesis with n-alk- 1 -enes,l* because the less strongly complexing olefin will not compete very well. The highly active catalysts (a, = 0.5; cf. table 1) have low selectivities and are good at cross-metathesisl69 l7 since olefin complexation is apparently almost non-activated and therefore rather non- selective. However the high trans directing noble-metal catalysts, such as RuCl,, are again highly selective in copolymerization16 in favour of the more strongly complexing olefin. In view of the first-order kinetics with respect to NBE c~ncentration,~ complexation strengths are again important.The stereoselective feature referred to as ‘formation of cis from cis and trans from trans’ sometimes found in the metathesis of pent-2-enes2* has been much discussed. Here we put forward an entirely novel explanation. It could well be related to cisltrans blockiness in ring-opened polymers and indeed is readily accounted for by two kinetically distinct metallacarbene species, one cis and the other trans directing. Such a pair analogous to P, and P or, better still, to P, and P,, is readily envisaged for the acyclics, cis-pent-2-ene having to displace cis-but-2-ene or cis-hex-3-ene in the complexation step, and therefore being forced to do so largely in the orientation which again gives an all-cis metallacycle (e.g. the behaviour of P, in ring-opening polymerization).For trans-pent-2-ene, however, the propagating metallacarbene is equivalent to P, and the only steric factor operating in the complexation step is the avoidance of an unfavourable cis-l,2-orientation of substituents, so the net result is that ‘trans is made from trans’. Puckering19 of the metallacycles is not therefore of primary importance as to the type of junction, cis or trans, which is ultimately formed, but may instead be of considerable significance in conjunction with group and orbital movements in the acts of formation and fission of the metallacyclobutanes (cf. 4 -+ 5). We thank the S.E.R.C. for financial support and are indebted to Professor K. J. Ivin for useful comments. 1 2 3 4 5 6 7 8 9 10 J. J. Rooney and A. Stewart, in Catalysis, ed. C. Kemball (Specialist Periodical Report, The Chemical Society, London, 1977), vol. 1 , p. 277. M. Leconte and J. M. Basset, J . Am. Chem. SOC., 1979, 101, 7296. Ho Huu Thoi, K. J. Ivin and J. J. Rooney, J . Mol. Curd., 1982, 15, 245. W. B. Hughes, J . Am. Chem. SOC., 1970,92, 532. C. Tanielian, A. Kiennemann and T. Osparpueu, Can. J . Chem., 1980, 58, 2813. K. J. Ivin, D. T. Laverty and J. J. Rooney, Mukromol. Chem., 1977, 178, 1545. K. J. Ivin, D. T. Laverty, J. H. O’Donnell, J . J. Rooney and C. D. Stewart, Mukromol. Chem., 1979, 180, 1989. Ho Huu Thoi, K. J. Ivin and J. J. Rooney, Mukromol. Chem., in press. B. S. R. Reddy, unpublished results. D. G. H. Ballard, J . Polym. Sci., Pofym. Chem. Ed., 1975, 13, 2191.HO HUU THOI, B. S. R. REDDY AND J . J . ROONEY 3317 I 1 C. P. Casey, D. M. Scheck and A. J. Shusterman, Fundamental Research in Homogeneous Catalysis, l 2 G . I. Devine, Ho Huu Shoi, K. J. Ivin, M. A. Mohamed and J. J. Rooney, J . Chem. SOC., Chem. l 3 L. Bencze, K. J. Ivin and J. J. Rooney, J. Chem. SOC., Chem. Commun., ,1980, 834. I5 U. Khlabunde, N. F. Tebbe, C. W. Parshall and R. L. Harlow, J . Mol. Caral., 1980, 8, 57. l6 K. J. Ivin, G. Lapienis, J. J. Rooney and C. D. Stewart, J . Mol. Cataf., 1980, 8, 203. ed. M. Tsutsui (Plenum Press, New York, 1979), vol. 3, p. 141. Commun., 1982, in press. K. J. Ivin, J. J. Rooney and C. D. Stewart, J . Chem. SOC., Chem. Commun., 1978, 603. K. J. Ivin, G. Lapienis and J. J. Rooney, Makromof. Chem., 1982, 183, 9. E. A. Ofstead, J. P. Lawrence, M. L. Senyek and N. Calderon, J. Mol. Cataf., 1980, 8, 227. N. Taghizadeh, F. Quignard, M. Leconte, J. M. Basset, C. Laroche, J. P. Lavaland A. Lattes. J . Mol. Catal., 1982, 15, 219. (PAPER 2/284)

ISSN:0300-9599    

DOI:10.1039/F19827803307

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chern. Soc., Faruduy Trans. 1, 1982, 78, 3319-3329 Thermodynamics of n-Alkane + Dimethylsiloxane Mixtures Part 4.-Surface Tensions BY BERYL E D M O N D S ~ A N D IAN A . MCLURE* Department of Chemistry, The University, Sheffield S3 7HF Receiued 17th February, 1982 Surface tensions have been measured by the differential capillary rise method under orthobaric conditions for binary mixtures of four n-alkanes (n-pentane, n-heptane, n-decane and n-tetradecane, replaced in the case of dimer by n-hexadecane) with four linear dimethylsiloxanes (dimer, trimer, tetramer and pentamer) at 303.2 K. The sign and magnitude of the excess surface tension depend ultimately upon the chain lengths of the components of the mixtures irrespective of whether a volume-fraction-based or a mole-fraction-based ideality is adopted.The dependence of the volume-fraction-based excess surface tension on chain length is very similar to that previously found for the excess volume. The results are analysed in terms of Prigogine's parallel-layer theory as modified by Gaines and Prigogine's average potential theory. Only the latter gives a useful description of the complexity of the chain-length dependence of the excess surface tension, and it is qualitative at best even then. We have reported previously measurements of gas-liquid critical temperatures and pressures, vapour pressures and excess enthalpies and volumes of mixing for mixtures of the type n-alkane+linear dimethy1siloxane.l Here we report the results of our measurements of surface tensions for sixteen of these systems and show that their complex dependence on the chain length of the components can be interpreted qualitatively in terms of a corresponding-states approach, not unlike that which was successful for the similar chain-length dependence of the excess volume.The sixteen binary mixtures studied under orthobaric conditions, i.e. at the saturation vapour pressure of the mixture at 303.2 K, comprised each of n-pentane, n-heptane, n-decane and either n-tetradecane or (for dimer only) n-hexadecane with each of hexamethyldisiloxane (dimer), octamethyltrisiloxane (trimer), decamethyl- te t rasiloxane (te tramer) and dodecame t hylpen tasil oxane (pen tamer). EXPERIMENTAL The orthobaric surface tension was measured by the differential capillary rise technique using a Pyrex glass cell containing only the degassed mixture.The diameters of the two Veridia precision-bore capillaries were confirmed at several points along the length of each capillary by weighing mercury. The diameters were found to be 1.00, and 0.19, mm (nominaIly 1.00 and 0.20mm, respectively) and the variation along the length of each capillary was 0.5%. The apparatus was cleaned successfully with permanganic acid, hydrogen peroxide and distilled water before oven drying. Mixtures were made up directly in the cell and their compositions were determined by weighing. The filled cell was wholly immersed in a water-filled thermostat whose temperature remained constant within & 5 mK. Equilibrium was deemed to have been reached when the difference in level of the menisci in the two capillaries, Ah, remained constant to within the precision of the cathetometer (kO.03 mm) for t Present address : Institution of Chemical Engineers, 3319 1 h.The surface tension y was Rugby CV21 3HQ.3320 n-A L K A N E + D I M E T H Y LSI LOX A NE MIXTURES where rl and r2 are the radii of the capillaries, g is the acceleration of free fall in our laboratory. and d is the density of the mixture calculated from the known densities of the n-alkanes2 and dimethyl~iloxanes~ and the excess volumes of mixing.' The contact angle was assumed to be zero; this was supported by visual observation, the good agreement of our values and literature values far the surface tension of the pure components and the similar agreement for the mixture benzene+n-hexane.* More details of the technique and of the materials used are available elsewhere.'.RESULTS Table 1 contains a comparison of our results for the surface tension y for the pure dimethylsiloxanes with some literature values; the agreement is within & 0.1 mN m-l in most caws. A fuller comparison will appear with the results of our measurements over a wide range of temperature.6 As a check on our procedure for mixtures we have measured y for benzene + n-hexane at 308.2 K. Fig. 1 shows our results and those of Schmidt et a1.;4 the agreement is satisfactory. The results of our measurements of y for n-alkane + dimethylsiloxane mixtures at 303.2 K are listed in table 2 and illustrated in fig. 2 as a function of the dimethylsiloxane mole fraction x,.The precision of y is believed to be 0.1 mN m-l, although as indicated above the accuracy may be less, and the precision of x, is - +0.001. TABLE 1 .-COMPARISON OF OUR SURFACE TENSIONS, y, FOR PURE SUBSTANCES AT VARIOUS TEMPERATURES substance T / K y/mN m-l y(literature)/rnN m-' Dimer 298 303 308 313 318 323 trimer tetramer pen tamer 298 303 308 313 318 323 298 303 308 313 318 323 328 333 298 303 313 318 323 15.4 15.0 14.7 14.3 13.9 13.5 16.6 16.2 15.8 15.4 15.0 14.6 17.3 16.9 16.5 16.1 15.7 15.4 15.0 14.7 17.7 17.4 16.7 16.4 16.0 15.7 (293 K)' 15.3 (297 K)', 14.82 (298 K)'O 16.96 (293 K)7 16.6 (297 K)9 16.05 (298 K)l0 17.60 (293 K)' 17.08 (298 K)1° 18.10 (293 K)' 17.7 (297 K)87g 17.08 (298 K)'OB. EDMONDS A N D I. A. MCLURE 332 1 FIG. 1.-Surface tension, 'J, for hexane (l)+benzene (2) mixtures at (from top to bottom) 303.2, 308.2 and 313.2 K: 0, this work; 0, according to ref (4).DISCUSSION Prigogine and Sarolea produced in 1950 the first molecular theory for the surface tension of chain-molecule mixtures,ll and successive refinements appeared over the next ten years. The most suitable data for testing the development of the theory have been the surface-tension results of Marechal for benzene-based monomer + dimer mixtures,12 those of Aveyard for binary mixtures of n-alkanes13 and those of LeGrand and Gaines for dimethylsiloxane oligomer + polymer mixturesg Each of these mixtures is of the kind which we shall call monohomologous, in that they contain only substances belonging to a single homologous series. The surface tension of such mixtures has an uncomplicated dependence on composition and chain length, and, not surprisingly, the behaviour can be fairly readily described by the Prigogine theory in either the original form or one of the modified forms.By contrast, just as with the bulk thermodynamics of chain-molecule liquid mixtures, the surface tension of polyhomologous mixtures, i.e. those containing substances from at least two homo- logous series, is more complicated than that of monohomologous mixtures and presents a greater challenge to theory. Thus our measurements of the surface tension of many of the binary dihomologous mixtures formed from the n-alkane, linear dimethylsiloxane and perfluoro-n-alkane series offer the opportunity for a more extended test of the theory of the surface tension of chain-molecule mixtures than has been available hitherto.Here we confine ourselves to the presentation and discussion of our results for n-alkane + dimethylsiloxane mixtures. From our previous work we know that the deviations from bulk ideality of these mixtures are modest, and we find that their surface behaviour is similarly only slightly non-ideal. The dependence on composition of the surface tension of all n-alkane + dimethyl- siloxane mixtures which we have studied is simple and regular. Only for n- pentane + dimer does aneotropy, or surface azeotropy, occur, and this probably owes3322 n-A L K ANE + D I M E THY LSI LOX ANE MIXTURES TABLE 2.-sURFACE TENSIONS, y, AT 303.2 K FOR n-ALKANE + DIMETHYLSILOXANE MIXTURES OF MOLE FRACTION X n-pentane n- hep tane n-decane n-hexadecane x, y/mN rn-l x, y/mNm-l x, y/mNm-l x, 7lmNm-l ~ 0 0.05 0.19 0.34 0.46 0.58 0.75 0.79 1 0 0.075 0.20 0.25 0.375 0.55 0.62 0.80 1 0 0.10 0.24 0.39 0.48 0.55 0.80 1 0 0.03 0.12 0.25 0.36 0.47 0.695 0.87 15.0 15.0 15.05 15.1 15.15 15.1 15.1 15.05 15.0 15.0 15.05 15.4 15.5 15.7 15.95 16.0 16.1 16.2 15.0 15.3 15.75 16.15 16.3 16.45 16.75 16.9 15.0 15.1 15.6 16.15 16.5 16.75 17.1 17.3 n-alkanes (1) + hexamethyldisiloxane (2) 0 19.3 0 22.9 0.08 18.7 0.07 22.1 0.19 18.3 0.16 21.1 0.37 17.3 0.295 19.75 0.55 16.5 0.42 18.7 0.67 16.1 0.61 17.4 0.80 15.65 0.75 16.5 0.88 15.45 0.90 15.6 1 15.0 1 15.0 n-alkanes (1) + octamethyltrisiloxane (2) 0 19.3 0 22.9 0.14 0.28 0.40 0.55 0.6 1 0.82 1 8.6 0.14 21.1 8.0 0.20 20.6 7.6 0.34 19.5 7.15 0.41 18.95 7.0 0.56 18.1 6.5 0.66 17.6 6.2 0.75 17.2 1 16.2 n-alkanes (1) + decamethlytetrasiloxane (2) 0 19.3 0 22.9 0.1 1 18.7 0.13 21.1 0.27 18.15 0.20 20.5 0.44 17.7 0.325 19.65 0.52 17.5 0.46 18.9 0.73 17.2 0.60 18.3 1 16.9 0.81 17.45 1 16.9 - - - - n-alkanes (1) + dodecamethylpentasiloxane (2) 0 19.3 0 22.9 0.05 19.0 0.04 22.2 0.13 18.7 0.12 21.3 0.27 18.3 0.32 19.8 0.40 18.0 0.49 19.0 0.515 17.8 0.67 18.3 0.58 17.7 0.76 18.0 0.77 17.5 1 17.4 0 0.175 0.28 0.38 0.55 0.60 0.76 0.86 1 0 0.20 0.25 0.44 0.55 0.66 0.70 1 0 0.18 0.255 0.33 0.44 0.61 0.71 0.745 1 0 0.14 0.18 0.34 0.52 0.65 0.78 1 26.6 23.2 21.6 20.5 18.7 18.3 17.0 16.2 15.0 25.7 22.1 21.5 19.65 18.75 18.0 17.8 16.2 25.7 22.3 21.5 20.75 19.9 18.8 18.3 18.1 16.9 25.7 22.5 21.9 20.4 19.4 18.8 18.3 17.4 more to the near equality of the surface tensions of the components than to any deeper cause; thus the occurrence here resembles a surface Bancroft point.The deviation of y from linearity in x is nowhere particularly marked; nonetheless we shall demonstrate that the dependence of this deviation from linearity on the chain length of the components is complex and provides the principal interest of the work. All of the foregoing contrasts sharply with the behaviour of mixtures containing perfluoro-B. EDMONDS A N D I. A. MCLURE 3323 0 0.2 0.L 0.6 0.8 1 .Y 2 288 26 I E Z E . - I c % . + 2 6 1 -1 0 0.2 01, 0.6 0.8 1 .Y 2 u I I r t l r l I I I 0.2 0.L 0.6 0.8 1 .Y 2 FIG. 2.-Surface tension, y , for n-alkane (l)+linear demethylsiloxane (2) mixtures at 303.2 K.In each diagram the curves correspond from bottom to top to n-pentane, n-heptane, n-decane and n-hexadecane (with dimer only) or n-tetradecane. Solid curves are drawn through the set of points corresponding to a given siloxane: 0, dimer; A, trimer; tetramer; '7, pentamer. n-alkanes, which is rich in non-linear surface-tension - composition relationships and in which both positive and negative aneotropy abound.14 The discussion of the surface thermodynamics of liquid mixtures is hampered by the lack of a statement of ideal behaviour enjoying the status of, say, Raoult's law or its near equivalents in the bulk thermodynamics of liquid mixtures. At least three3324 n-A LK AN E + D I MET H Y L S I LOX AN E MIX T U RES ways of describing ideality at the gas-liquid interface have been used.Two of these are essentially empirical and share with Raoult’s law the advantages of simplicity. They are ?id = x,Y,+x,Y, (2) and (3) where x i , 4i and yi are the mole fraction, volume fraction and surface tension in the pure state, respectively, of component i of the mixture and yid and yid are the surface tensions of the ideal mixture according to the two conventions. The second of these conventions is more appropriate for polymer mixtures. The third statement of ideality in currency was developed by Guggenheim from the quasi-crystalline treatment of interfaces;15 in this treatment the ideal surface tension of the mixture yg is given by exp(-yga/kT) =x,exp(-y,a/kT)+x,exp ( - y , a / k T ) (4) where a is the area of surface per molecule and T is temperature.This expression is obtained from the general theory by setting the interchange energy w equal to zero. This procedure is the counterpart of that which generates Raoult’s law from the quasi-crystalline theory for bulk mixture thermodynamics. The presence of the surface area a in eqn (4) introduces an undesirable lack of generality into this definition of surface ideality. The main attraction of the Guggenheim treatment is that it is based on a simple model and so it is easy to determine the effect of changes in the interchange energy on the surface tension of the mixture. The first two conventions are based on no model and it is hard to interpret in any simple way deviations from ideality so defined. Objections notwithstanding, each of these conventions possesses certain virtues of which we shall take advantage in turn.Our principle use of the concept of surface ideality is the definition of excess surface tension of surface ideality where the subscript indicates which of the three conventions for surface ideality is in Fig. 3 shows the excess surface tension for the equimolar mixture 7,” (x = 0.5) for all sixteen mixtures as a function of the n-alkane chain length n,. The most striking features are the regular family of smooth curves connecting quartets of points corresponding to mixtures containing a common dimethylsiloxane partner and the crossover point at n, = 6.5, which corresponds to the excess surface tension of an equimolar mixture of the hypothetical n-alkane of chain length 6.5 with any linear dimethylsiloxane. The magnitude of 7,” always increases with dimethylsiloxane chain length but only for mixtures containing n-pentane is 7,” positive.We see no reason to anticipate a cyclic pattern whereby positive y,” would recur at higher alkane chain length. In view of the relationship between mole fraction and volume fraction, it is no surprise to find regularities in the behaviour of the excess surface tension for the mixture of volume fraction 0.5 yf (4 = 0.5) shown on fig. 3 similar to that of 4E (x = 0.5) in fig. 4 although topologically different. Only for n-pentane dimer is yf positive, and the order of increasing magnitude of y,f on the high-chain-length side of the crossover point, in this case at nc z 7, is opposite to that for y f ; i.e.yf decreases with increasing dimethylsiloxane chain length for nc > 7. This feature and the generally smaller magnitude of yf compared with that of 7,” suggest that for chain- Play -B. EDMONDS A N D I. A. MCLURE 3325 1 ~ ~ ~ ' 1 ~ ~ 1 1 1 ' 0 - I E G 5 -1 - h m c I1 8 - - 2 -2 - 5 6 7 8 9 10 11 12 13 1L 15 16 'lc FIG. 3.-Volume-fraction-based excess surface tension, $, for n-alkane + linear dimethylsiloxane mixtures at 303.2 K and 4 = 0.5 plotted against the length, n,, of the alkane component of the mixture. Solid curves are drawn through the set of points corresponding to a given siloxane: 0, dimer; A, trimer; 0, tetramer; V, pentamer. ~ " " " " 1 ~ FIG. 4.-Moie-fraction-based excess surface tension, y:, for n-alkane + linear dimethylsiloxane mixtures at 303.2 K and x = 0.5 plotted against the chain length, n,, of the alkane component of the mixtures.Solid curves are drawn through the set of point corresponding to a given siloxane: 0, dimer; A, trimcr, 0, tetramer; V, pentamer. molecule mixtures yy is a more appropriate convention for surface ideality than ?id, at least in the predictive sense, and especially so when all components of the mixture are of large molecular weight. Both fig. 3 and 4 are very reminiscent of the diagram showing the fractional volume change at 4 = 0.5 for n-alkane+dimethylsiloxane mixtures as a function of nc [fig. 1 of ref. (lc)], in which the crossover points occurs at n, z 6.3326 n-A LK A N E + DIM ET HY LS I LOXANE MIXTURES Any successful theory of the surface tension of chain-molecule mixtures can reasonably be expected to describe and account for the following features of our experimental results and their counterparts in volume fraction terms.(1) The small positive 7,” of mixtures with n, < 6.5 and the increasing magnitude of 7,” of such mixtures with dimethylsiloxane size. (2) The existence of the crossover point in the neighbourhood of nc z 6.5. (3) The increasingly negative y,” with increasing dimethylsiloxane oligomer size at nc > 6.5. We describe next our attempts to achieve these aims in terms of two treatments of increasing sophistication. PR IGOG INE ’ s PAR A LLE L-L A Y ER MODEL TREATMENT A more realistic description of the surface tension of mixtures of molecules of different sizes was given by Prigogine and others, first for monomer+dimer and monomer + trimer, and then generally for monomer + oligomer of chain length r, called rmers by the Brussels school.A simplified model was later developed by Prigogine and Marichal for the special case in which rigid molecules only were considered and only those configurations in which the rmer lay parallel to the surface were taken into account.16 This treatment, based on the parallel-layer model, was developed chiefly for the athermal case, which we would expect to be a reasonable representation of n-alkane + dimethylsiloxane mixtures. The equations which yield the surface tension of the mixture are Y = Y1 + (W4 [In W 4 1 ) + (1 - r-l) (4; - 4211 where r is the chain length of component 2 of the mixture. A modification of these equations was suggested by Gains on the basis of a model in which only the surface layer behaves at her mall^.^ Non-ideality is allowed for by including a non-athermal interaction term in the chemical potential for the bulk mixture.The resulting generalised expressions for a mixture of oligomers of chain length rl and r2 are Y = Y 1 + ( W r 1 4 {In W 4 1 ) + (4; - 42) [I - Wr2)II - ( P / 4 4; = Y2 + ( W r 2 a) {In (4242) + (4; - 41) [(T2/Yl) - 11) -@/a) 4;. The following useful expressions for rl and r2 have been developed previously,lc where as before component 1 is the n-alkane and component 2 is the dimethylsiloxane: rl = (nc+2)/2 and Y, = 3nSi/2. The expressions are similar but not identical to those given by Simha and Havlik.” The surface area a was calculated from the volume per segment as a = (&/Nri)2/3 where Vi is the molar volume of component i and N is Avogadro’s number.The quantity p is in a sense an interaction energy similar to Guggenheim’s exchange energy w, but in Gaines’ treatment it assumes the role of a differential or excess interaction energy between the surface and the bulk of the mixture. In practice, p is an adjustable parameter which absorbs the consequences of effects neglected in the main theory. The computational procedure was to calculate 4’, the surface volume fraction, from the experimental y at 4 = 0.5, and then to calculate p at 4 = 0.5. Knowing B, 4’ was calcuated at each 4 and thus the entire theoretical isotherm was generated. The calculated and experimental isotherms are in good agreement.Typical resultsB. EDMONDS A N D I. A. MCLURE 2L * 3327 0 0.2 0.L 0.6 0.8 @2 FIG. 5. Predictions of the Prigogine-Gains parallel-layer model for surface tensions for the mixtures n-heptane +dimer (lower curve) and n-decane + pentamer (upper curve) at 303.2 K. The solid line represents the theory and the points are experimental. The dashed line corresponds to volume-fraction-based ideality. TABLE 3.-vALUES OF PARAMETER p REQUIRED TO DESCRIBE SURFACE TENSION OF n-ALKANE+ LINEAR DIMETHYLSILOXANE MIXTURES AT 303.2 K USING THE PRIGOGINE-GAINES PARALLEL- LAYER MODEL n-a1 kane dimer trimer tetramer pentamer n-pen tane - 5.8 + 9.4 6.4 5.8 n-heptane - 24.6 - 14.3 -9.3 -6.1 n-decane - 52.9 - 38 - 39.6 -38.1 n-tetradecane - 72.4a - 60.8 - 64.9 - 55.4 a n-Hexadecane + dimer.are shown in fig. 5 for n-heptane+dimer and n-decane+pentamer. Table 3 lists the values of p; they show some regularity with n-alkane chain length but they are independent of dimethylsiloxane chain length. This pattern can be interpreted loosely as reflecting the widely suspected difference between end- and middle-group interactions for n-alkanes, which is thought to be unimportant for dimethylsiloxanes and perfluoroalkanes. In Gaines' work on mixtures of toluene or tetrachloroethylene with dimethysiloxanes the values of p were considerably scattered, and little obvious correlation could be made with dimethylsiloxane chain length.'* The success of the model in describing our results is certainly to some extent a result of fortuitous compensation of errors.For example, deficiencies in evaluating the bulk-mixture chemical potentials are cancelled by errors in the surface-layer chemical potential. Only where the relative adsorption is weak, as in n-alkane + dimethylsiloxane mixtures, is the theory trustworthy, and even here its success is marred by the need to guess or otherwise put a value on p. Another objectionable feature of the method is that in no way does it account for the influence of chain flexibility on the surface tension of mixtures. In the next section we move to a discussion of a treatment which does attempt to introduce this effect.3328 n-A L K A N E -k D I METHY LSI L O X A N E MIXTURES PR I GOG I N E'S A VE RAG E-PO TE N T I A L THEORY In parallel with the extension to chain-molecule mixtures of the average-potential model (APM) for the bulk behaviour of mixtures of small molecules, Bellemans extended Prigogine's average-potential theory for the surface tension of mixtures of small molecules to mixtures of chain molecule^.^^ The essential features of the theory are that the surface tension of both components of the mixtures and of the mixture itself obey the same reduced surface equation of state, i.e.that the reduced surface tension yo2/&, where a and E are characteristic size and energy parameters, respectively, is a universal function of the reduced temperature_F= kT/E. The replacement of E , CT and by average values ( E ) , ( 0 ) and thus (7') generates the dynamic surface tension Ydyn of the mixture corresponding to a freshly formed surface of the same compcpition as the bulk of the mixture.If E and a for the components are not too dissimilar then ydyn and y of the mixture, sometimes termed Ystat, are related by In the form of the theory used here the usual APM rules for ( E ) and (a) were used. The reduced temperature was taken as ckT/qE* where c and q are the structural quantities related to entropy and energy in the usual corresponding-states treatment of chain molecules. The previous expressions were used for r, and r2 and the expressions of Simha and Havlik were used for c c, = ( n c + 7 ) / 6 and c2 = n,,+2. For q the following quasi-crystalline expression was used : Y = Ydyn - 4 2 ri r z a ? / 8 ( r ~ 41 -k r2 4 2 ) k T ] (&d~n/dd)~. qz = r ( z - 2 ) + 2 . At 303.2 K all our substances are far from their critical point, and so z = 12 was assumed.The surface area per segment of component 1, a;, was again estimated from a: = Vm(~egment)/~,N2/3. . 1 1 1 1 1 I I I ~ l 8 5 6 7 8 9 10 11 12 13 1G 15 16 "( FIG. &Predicted volume-fraction-based excess surface tension, y 8, for n-alkane + linear dimethylsiloxane mixtures at 303.2 K and 4 = 0.5 using Prigogine's average-potential model as extended to oligomer mixtures by Bellemans; nc is the chain length of the alkane component. The calculations incorporate the assumptions 6 = p = 8. From top to bottom along the left-hand side of the diagrams of the lines represent mixtures with second component in the order dimer, timer, tetramer and pentamer.B. E D M O N D S A N D I.A. MCLURE 3329 In the calculation the Progogine interaction parameters p, 6 and 8 were all set to zero. Patterson and Rastogi have demonstrated by phenomenological analysis that the n-alkanes and the dimethylsiloxanes obey the same surface principle of corresponding states20 and we have extended the certainty of this using new measurements of the surface tensions of the dimethylsil~xanes.~ The predicted surface tensions are in relatively pcor agreement with experiment, as can be seen from fig. 6 at least as far as magnitude is concerned. However, and we believe that this is more important, the variation of yf with chain length is qualitatively well-reproduced. Better agreement could undoubtedly be forced by adjustment of the various parameters in the equations, but in view of the well-known deficiencies of the average-potential model there seems little purpose in so doing.Note that the essential features of this model - the establishing of the various r, c and q average-parameter quantities - proceeded in the same way as for the similarly successful prediction of A V / V for the same mixtures. An alternative analysis of our results, free of the problems of the average-potential model, has been published by Dickinson.21 B.E. gratefully acknowledges receipt of a CASE studentship from the S.R.C. We both record our gratitude to Dr M. La1 of the Unilever Research Laboratory, Port Sunlight, for useful discussions during the course of this work. I E. Dickinson and I . A. McLure, J . Chmi. Sol,., Furudqj' Truns. I , 1974, 70. (a) 2313, ( h ) 2321 and (c) 2328. R. A. Orwoll and P. J. Flory, J . Am. Chem. Soc.. 1967. 55, 68 14. 1. A. McLure, A. J. Pretty and P. A. Sadler, J . Chem. Eng. Datu. 1977. 22. 372. R. L. Schmidt, J. C. Randall and H. L. Clever, J . Phys. Chem.. 1966, 70. 3912. B. Edmonds, Ph.D. Thesis (University of Sheffield, 1972). J. F. Neville, B. Edmonds and I . A. McLure. Polymer. in preparation. H. W. Fox, P. W. Taylor and W. A. Zisman, tnd. Eng. Chrm., 1947, 39. 1401. G. L. Gaines, J . Phys. Chem., 1969, 73, 3143. D. G. LeGrand and G . L. Gaines Jr, J . Polvn. Sci.. Purt C. 1971, 34, 45. I . Prigogine and L. Sarolea, J . Chim. Phj,.s., 1950, 47. 807. *IJ M. J. Hunter, E. L. Warwick, J. F. Hyde and G. C. Currie. J . Am. Chtwi. Soc .. 1946. 68, 2334. l 2 J . Marechal, Bull. Soc. Chim. BeIg., 1952, 61, 149. ':' R. Aveyard. Truns. Furudliy Soc., 1967, 63, 2778. I . A. McLure, B. Edmonds and M. Lal, Nuture ( P h j ~ . Sci.), 1973, 241, 7 I . E. A. Guggenheim, Truns. Furuduj, Soc., 1945, 41, 150; Mi.\-turc.s (Clarendon Press, Oxford. 1952). chap. IX. I . Prigogine and J. Marechal, J . Colloid Sci., 1952, 7, 122. l 7 R. Simha and A. J. Havlik, J . Am. Chtm Soc., 1964. 86, 197. IN G . L. Gaines. J . Phys. Chtwi.. 1969, 73. 3150. IY A. Bellemans, J . Chim. Phys., 1960, 57, 40. I" D. Patterson and A. K . Rastogi, J . PJ1y.s. Chcvn., 1970. 74, 1067. E. Dickinson. J . Colloid 1nrerfuc.e Sci.. 1975. 53. 467. (PAPER 2/301)

ISSN:0300-9599    

DOI:10.1039/F19827803319

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 3331-3340 Ruthenium Dioxide: A Redox Catalyst for the Generation of Hydrogen from Water BY PATRICK KELLER AND ALEC MORADPOUR* Laboratoire de Physique des Solides, Universite de Paris-Sud, 91 405 Orsay, France AND EDMOND AMOUYAL Laboratoire des Processus Photophysiques et Photochimiques, Universite de Paris-Sud, 91 405 Orsay, France Received 19th February, 1982 Ruthenium dioxide, already known to be a catalyst for the oxidation of water to oxygen, has been shown to mediate effectively the generation of hydrogen from water in Ru(bipy)g+, methyl viologen and EDTA model systems, with efficiencies even higher (pH > 5) than those of previously investigated platinum catalysts. The main factor limiting the evolution of hydrogen is found to be the destruction of the organic electron-transfer relay, and this side-reaction is attributed to H&,,, species formed on the particles of catalyst. Ru0,-coated electrodes are widely used as so-called dimensionally stable anodes in the large-scale production of chlorine.In addition to this wide-spread industrial application, which is due to the high corrosion resistance and low overvoltage for chlorine evolution, other electrocatalytic properties as well as its basic electrochemical features have been widely studied1 during the last decade. Thus an exceptionally low overvoltage was also observed for oxygen evolution with this material;, however, the electrocatalytic activity decreased with time and the long-term stability of this type of electrode when generating oxygen was not sati~factory.~ In fact, in this case electrocatalysis has been attributed to a metastable substoichiometric RuO,(x < 2) oxidation state, and the fall-off of the oxygen- generation efficiency with time has been assigned to the possible formation of the more stable (less active) stoichiometric R U O , ~ ~ and/or to the oxidation to RuO,, which subsequently passes into solution causing degradation of these electrode^.^ On the other hand, RuO, exhibits hydrogen overvoltages similar to those on Pt electrodes,6 although hydrogen discharge produces a marked modification of the RuO, electrode and persistent H, evolution causes the RuO, layer to collapse, presumably by its reduction to metallic R U .~ Heterogeneous catalysts prepared from electrode materials, such as platinum or graphite, have long been known to promote various redox reactions.' The use of platinum as a catalyst for the reduction of water to hydrogen by vanadium(I1) was observed in 19028 and the oxidation of water to oxygen by cerium(rv) (also catalysed by platinum) has been studied more recently in a series of redox reaction^.^ These redox processes (catalysed by PtO,, Ir0,lo or RuO,ll) have recently been rediscovered, but the catalytic nature of the RuO, powders used in these experiments involving Ce4+ has been questioned.12 The renewed interest attracted by such redox catalysts has been stimulated by studies of the development of water-splitting processes to store solar energy.l 9 Thus, 333 13332 RuO, AS A REDOX CATALYST light-induced hydrogen formation catalysed by colloidal Pt149 l5 and oxygen generation promoted by Ru0,16 have been proposed in ' sacrificial ' model systems, where H, and 0, are produced at the expense of irreversible consumption of externally added compounds. Curiously, the catalytic activity of RuO,, which could have been used to promote both hydrogen and oxygen formation in these experiments because of its very low overvoltage for the generation of both species, was generally restricted to studies of oxygen production.l1?l6 Moreover, the necessity of the presence of Pt and RuO,, combined on semiconductor particles, was emphasized in recent studies of cyclic (non-sacrificial) water-splitting to hydrogen and oxygen by visible light.17 We have recently reported18 that RuO, is an effective redox catalyst for the generation of hydrogen with the Ru(bipy)i+/methyl viologen (MV2+)/EDTA sacrificial system.We now report detailed studies of the light-induced, Ru0,-catalysed formation of H, and examine the unavoidable limitations of this catalysis encountered in the present sacrificial photosystem and which may be anticipated for other photocatalytic processes. EXPERIMENTAL MATERIALS The sources and methods of preparation and purification of Ru(bipy);+, methyl viologen (MV2+), EDTA and platinum hydrosols have already been described.15 The ruthenium dioxide catalyst was used either as the commercially available powder (Alfa Ventron, soluble form) or as a more finely dispersed sample obtained by loading solid supports [Linde molecular sieve type LZ-Y52 (Alfa Ventron, catalyst A) or titanium dioxide (Alfa Ventron, catalyst B)] as described elsewhere.16 The mixed catalyst (RuO,/IrO,) loaded on the above mentioned zeolite was also prepared by this method.lG Aqueous solutions were prepared from distilled water and the pH was adjusted to the desired value using acetate (pH 4 and 5 ) , phosphate (pH 6, 7 and 8) or borate (pH 9) buffers.METHODS PHOTOCATALYTIC H, PRODUCTION The photochemical procedures and apparatus have been described in detail.'*, l5 30 cm3 solutions containing Ru(bipy)i+ (2 x lop4 mol dm-3), MV2+ (5 x lop4 mol dmp3) and EDTA (0.2 mol dm-3) were first thoroughly purged by argon and then the required amount of catalyst powder was added. The continuously stirred mixtures were irradiated with a 250 W halogen slide-projector lamp.The production rates and the total amounts of hydrogen were measured, after the evolved gases had been bubbled through a 50% potassium hydroxide solution, with a 6cm3 scale gas volumeter. The amount of MV2+ present during irradiation (or the electrochemical experiments) was measured by high-pressure liquid chromatography (h.p.1.c.) analysis. H, FORMATION MEDIATED B Y THE ELECTROCHEMICALLY REDUCED MV2+ A specially designed two-compartment electrochemical gas-tight cell was used. Aqueous solutions containing EDTA (0.12 rnol dmP3) and MV2+ (5 x rnol dmP3), buffered to pH 6, 8 or 9 were introduced respectively into the anodic (50 cm3, graphite auxiliary electrode) and the cathodic (60 cm3, Hg pool working electrode and saturated calomel electrode) compartments. The two solutions, linked by an agar/KCl bridge, were purged with argon before the addition of the catalyst.Th,: required potential to generate MV'+ (-0.7 V us. SCE) was then applied (Tacussel PRT 100- IX), the cathodic compartment being connected to the above mentioned gas volumeter. The rates of formation of hydrogen (if produced) and yields, as well as the amounts of MV2+ present in these electrolytic runs, were measured as a function of time.P . KELLER, A. MORADPOUR A N D E. AMOUYAL 3333 CATALYTIC HYDROGENATIONS The possibility of the hydrogenation of MV2+ catalysed by RuO, (1 atm* H,) was also examined with aqueous mixtures (pH 5, 6 and 8) containing MV2+ and the catalyst, as previously effected with platinum ~ata1ysts.l~ RESULTS AND DISCUSSION Ru0,-c A TA L Y SED H, PHOTO PRODUCTION : pH E F F E c TS Visible-light irradiation of outgassed solutions of Ru(bipy)i+, MV2+ and EDTA results in appreciable evolution of hydrogen when RuO, is added to the aqueous mixtures; the catalytic efficiency of this oxide was first examined as a function of pH (table 1).TABLE 1 .--HETEROGENEOUS CATALYSIS OF H, PHOTOPRODUCTION FROM IRRADIATED AQUEOUS SOLUTIONS OF Ru(bipy)i+ + MV2+ + EDTAU H, formationd - MV2+ amount of rate yield (turnover entry pH catalyst/pmolbp /cm3 h-l /mmol number)e 1 2 3 4 5 6 7 8 9 10 1 1 12 4 5 6 7 8 9 5 5 5 5 5 6 I :30 I:30 I:30 I:30 I:30 I:30 I:60 I :90 I : 120 I : 180 I1 : 0.7 I1 :0.4 1.5 2.6 2.6 1.2 0.2 0.0 2.7 2.6 2.6 2.3 4.5 0.2 0.63 1.03 0.96 0.41 0.13 0.00 1.12 1.12 1.12 0.80 1.35 0.04 42 69 64 27 9 0 74 74 74 54 90 3 a See Experimental section for concentrations; I = RuO,, Alfa Ventron, soluble form; I1 = colloidal Pt, catalyst A, ref.(1 5 ) ; determined within lo%, ref. (1 5 ) ; defined as the ratio of the total amount of H, to the initial amount of MV2+. As irradiation proceeded, hydrogen was first produced at a constant rate, but then gas formation decreased and spontaneously stopped at every investigated pH. The rates and total yields, measured in the usual way, varied markedly as a function of pH (table 1, entries 1-7), maximum values being observed for pH 5 and 6. The irradiated solutions were also analysed by h.p.1.c. as a function of the irradiation time with respect to the amounts of MV2+ present in the mixtures.These analyses allowed us t o assign the termination of gas production to the almost total disappearance of the organic relay; the rates of this destruction were also observed to be dependent on the pH values (fig. 1). Several aspects of the Ru0,-catalysed formation of hydrogen are reminiscent of * 1 atm = 101 325 Pa.3334 RuO, AS A REDOX CATALYST 100 75 h E A 50 > z 25 0 L I 1 200 LOO 600 time/min FIG. 1 .-Analysis of MVZ+ during hydrogen photoproduction by Ru(bipy)$+ + MVz+ + EDTA + RuO,, using h.p. 1.c. Irradiations correspond to the experiments reported in table 1 : curve A (entry 2, pH 5), curve B (entry 4, pH 7). features typical of the results obtained previously with a colloidal-platinum catalyst :15 (i) the hydrogen-formation rates and the dependence of the overall yields on the pH and (ii) the limitation of gas production by the destructipn of the MV2+ organic relay.However, significant differences distinguish the two processes. With the platinum catalyst, an MV2+-hydrogenation reaction limits the catalytic efficiency of hydrogen photogeneration. This side-reaction (see also fig. 2) accounts for the observed variations in the hydrogen formation rates (quantum yieldslS) and the overall yields as a function of the amount of Pt catalyst added: both factors increased and then sharply decreased as the amount of Pt was increased.l49 1 5 9 l9 On the other hand, when the amount of catalyst added corresponded to that required to obtain the optimum efficiency of the photosystem drastic differences were observed, this time as a function of the pH of the irradiated solution, in the levels of hydrogen production.Thus, the yields at pH 6 were lower, by more than one order of magnitude, than those obtained at pH 5 (compare entries 1 1 and 12 in table 1); these profound differences were explained by the increasing efficiency of Pt to hydrogenate MV2+ as the pH value was raised.15 Turning now to RuO,, the possible hydrogenation of the MV2+ relay was also envisaged in the present photoproduction of hydrogen. In fact the use of an Ru0,-basedP . KELLER, A. MORADPOUR A N D E. AMOUYAL 3335 t o o 75 h E + pl > 2 50 25 C g B I I 50 100 time/min FIG. 2.-Analysis of the stability of MV2+ in the presence of Pt and RuO, under hydrogen, using h.p. 1.c. Curves A and B, Pt catalyst, pH 5 and 7, re~pectively;'~ curve C , RuO, catalyst, pH 5, 6 and 8 (see text for further details).catalyst in heterogeneous hydrogenation is not uncommon.2o This possibility was first examined (as with the Pt catalyst)15 in the absence of any extraneously formed photochemical intermediates, by the simple combination of MV2+ and RuO, under hydrogen (fig. 2). Under these conditions, the stability of MV2+ in H,-purged solutions in the presence of RuO, was higher in comparison with colloidal Pt. Whereas the latter promoted large decreases in the initial amounts of MV2+, with strongly pH-dependent hydrogenation rates (fig. 2, curves A and B), the former was inactive catalytically for hydrogenation over the photochemical timescale. The hydrogenation of MV2+ was only noticeable after extremely long induction periods, (as sometimes observed with Ru0,-based catalysts before ruthenium metal, effectively active in the hydrogenation, is p-oduced in situ20)* 8 h at pH 5 and > 24 h at pH 6 and 8 (fig. 2, curve C ) .Moreover, in Pt-catalysed processes the electron-exchange reaction between hydrogen and MV2+ [equilibrium (l)] MVZ++iH, MV+'+H* catalyst is readily identified (i.e. at pH 7) by the intense blue colour of MV'+. When RuO, is added to aqueous solutions of MV2+ under hydrogen, the formation of the radical cation is no longer detected. These facts indicate the inability of this oxide to * Ruthenium metal as a catalytic species for the formation of hydrogen in the present photosystem may, however, be excluded because of its great ability to catalyse the hydrogenation of MV2+ with an efficiency even higher than that of colloidal Pt.3336 RUO, AS A REDOX CATALYST chemisorb hydrogen dissociatively21 (to hydrogenate MV2+ to undesired side-products) within the photochemical timescale, although the evolution of hydrogen from photochemically reduced MV'+ [through the reverse step of reaction (l)] is mediated efficiently by this catalyst.The striking difference in the observed effects of pH on the rates of hydrogen photoproduction catalysed by RuO, (table 1, entries 1-5) as compared with the Pt-mediated process (table 1, entries 1 1 and 12), is at first surprising if only the overpotential for H, evolution is considered (similar values have been determined for both catalysts6).This simply reflects the large reaction-rate differences for the catalysis by Pt and RuO, of the two competing processes, hydrogen formation [reaction (l)] and undesired hydrogen consumption. Considering the respective Nernst potentials associated with MV2+ oxidation (Emed) and water reduction (EH,) half-reactions [reactions (2) and (3)], MV2+ + e- g MV' + (2) H+ +e- $H2 (3) E H , = - (RT/nF) pH the catalysed formation of H, through the MV2+-mediated redox equilibrium [the sum of reactions (2) and (3)] requires, on a simple thermodynamic basis, the amount of MV'+ to exceed the equilibrium value at a given pH. Steady-state gas production then results from an excess of MV'+ generated upon illumination: at pH 8 the calculated I value of MV2+/MV'+ was 0.286 (MV'+ = 77%) and H, was effectively observed with RuO, as a catalyst (table 1, entry 5), whereas Pt failed to produce any detectable amount of H, under these conditions.However, an analysis of the amount of MV2+ in the irradiated aqueous mixture of Ru(bipy)t+ + EDTA + PtO,* at pH 8 showed a fast decrease in MV2+ (fig. 3, curve A) as compared with the rates of the uncatalysed (presumably radical-initiated) MV2+ destruction processL5 (fig. 3, curve B) or with the Ru0,-catalysed hydrogen production at this pH (fig. 3, curve C). Thus the Pt-based catalysts even prevented the production of H, in high-pH media, due to their high efficiency to hydrogenate the MV2+ relay. In this'case the H, produced was used exclusively in situ (at pH 8) to reduce the unsaturated organic compound.Note that a quantitative comparison of H, formation rates, corresponding to rather different H, yields, as a function of pH may be rather meaningless for the photosystems under study. Consequently, due to the interference of the side-reactions inherent in the MV2+/Pt couple (which depending on conditions may even be the main reaction, vide supra results at pH 8) these rate values may not bz appropriate as supporting data to test the otherwise interesting electrochemical predictive model of colloidal catalysis recently 23 RU0,/MV2+ COUPLE: THE LIMITING FACTOR I N H, GENERATION The hydrogenation of MV2+ has been ruled out for Ru0,-catalysed photoprocess by the independent investigation of the simple RuO, + MV2+ mixtures under hydrogen (fig. 2, curve C).Nevertheless the H, evolution obtained for visible-light irradiation of the Ru(bipy);+ + MV2+ + EDTA + RuO, solution stopped spontaneously, and the total hydrogen yields (table 1) were thus limited, as with Pt at every investigated pH, by the disappearance of the organic relay. The occurrence of other side-reactions, involving for example the EDTA radical-type oxidation by-products,15 seemed * PtO, was used here instead of colloidal Pt, which was unstable at this pH.P. 100 75 A 5 0 > E: 25 KELLER, A. MORADPOUR AND E. AMOUYAL 3337 I I-, 200 L 00 600 time/min FIG. 3.-H.p. 1.c. analysis of the amount of MV2+ present in the photochemical experiments at pH 8. Irradiated solution corresponds to the composition reported in table 1. Curve A, PtO,; curve B, without catalyst; curve C, RuO, (3 x lops mol).reasonable. Therefore hydrogen formation, mediated by electrochemically reduced MV'+ and catalysed by RuO,, was investigated in the absence of any extraneous photochemical electron-transfer process. The stability of MV2+ was examined, by h.p.1.c. analysis, in a simple two-component RuO,/MV'+ system. Hydrogen formation was in fact observed when the radical cation was continuously generated at pH 6 and 8, and the corresponding amounts of MV2+ were found to decrease with time (fig. 4, curves A and B). At pH 9 however, the organic compound was found to be very stable (fig. 4, curve C) and hydrogen formation was no longer detected in this medium. Thus whenever H, was generated, a side-reaction inherent to the simple RuO,/MV'+ couple took place and destroyed the organic relay.The possibility of Ru0,-catalysed hydrogenation of MV'+ was first considered. However, this hypothesis (which was in contradiction with the inability of RuO, to activate hydrogen21) was rejected, as MV'+ was also found to be stable (pH 9) when H,-purged solutions (instead of argon) were used in the electrochemical cell. The 'reactivity' of H,, produced by the Ru0,-catalysed process [reaction (4)] MV'+ + H+ + MV2+ + BH, (4) RuO, on MV'+ was thus completely different from that of the added hydrogen gas. I on FAR 783338 100 75 E + N > E 50 25 RuO, AS A REDOX CATALYST \ 200 400 600 time/min FIG. 4.-Analysis of the stability of MV2+ in the Ru0,-catalysed formation of hydrogen mediated by electrochemically reduced MV'+, using h.p.1.c.Curves A and B, pH 6 and 8, respectively; curve C, pH 9 (see Experimental section for the details). These results clearly show that Hiads) species, spilled over the catalyst surface by the electron-transfer relay,24 react efficiently with the organic salt. With electrodes made of RuO, these adsorbed species are assumed to lead to hydrogen through an ion-atom recombination step2' [reactions (5) and (6)] H++e-,H;,d,, ( 5 ) (6) and are also envisaged to be the catalytically active intermediates in heterogeneous hydrogenations. RuO, did indeed catalyse effectively the ' self-destruction ' of MV' + by these Hiads) species, although it was not, strictly speaking, a hydrogenation catalyst under the present conditions. Hiads) + H,O+ + e- * H2 + H,O MODIFIED RUO,-BASED CATALYSTS The maximum efficiencies in the photoproduction of hydrogen obtained in the Ru0,-catalysed system were closely comparable to the values obtained in the colloidal-Pt-catalysed photosystem (cf. entries 7 and 11 in table l), but the amount of RuO, macrodispersed powder required in these irradiations (entry 7 = 60 pmol) was roughly two orders of magnitude higher than the corresponding amount of the more finely dispersed Pt hydrosols (entry 11 = 0.7 pmol).In order to increase the catalyst efficiency (turnover number) solid-support deposited oxides were investigated. Very efficient zeolite-supported metal oxides have been proposed for the photoinduced generation of oxygen from water,16 and now several oxides have been prepared by a similar procedure and examined as catalysts in hydrogen photogeneration.TheP. KELLER, A. MORADPOUR AND E. AMOUYAL 3339 hydrogen-formation rates and yields promoted by these supported catalysts (table 2) are close to those previously measured using unsupported RuO, powders (entry 1, table 2), but these values were obtained with significantly lower amounts of added oxides (cf. entries 1 and 3, table 2). Moreover, the most efficient sample for hydrogen production (RuO, + 11-0, mixed oxides, entry 4) was also found to be the best catalyst for oxygen formation in the Ru(bipy):+ + Co(NH,),Cl* Cl, sacrificial mode1,16 although this is probably not related to any similarity between the catalytic sites (not yet precisely defined) involved in hydrogen and oxygen formation. TABLE 2.-FORMATION OF HYDROGEN BY THE VISIBLE-LIGHT IRRADIATION OF AQUEOUS SOLUTIONS OF Ru(bipy)i+ + MV2+ + EDTAU PROMOTED BY VARIOUS HETEROGENEOUS METAL-OXIDE CATALYSTS catalyst H, formationC amount added rate yield entry typeb /pmol /cm3 h-l /mmol 1 RuO, (powder) 60 2.7 1.12 2 RuO,/zeolite 15 1.8 0 .9 2 3 RuO,/TiO, 7.5 2.4 0.67 (catalyst A) (catalyst B) zeolite (catalyst C) 4 RuO,+IrO,/ 11.4 3.0 1.12 a , See Experimental section for concentrations and for catalyst preparation method; see note (d) of table 1. CONCLUSIONS The efficiency of Ru0,-based catalysts for the generation of hydrogen from aqueous mixtures of Ru(bipy)i+ + MV2+ + EDTA irradiated by visible light was found to be at least comparable to (or even higher than, depending on the pH values) the efficiency of previously investigated colloidal-Pt catalysts.The higher efficiency of RuO, is related to an inability to chemisorb hydrogen dissociatively and to catalyse the undesired hydrogenation of the organic relay, which limits the formation of hydrogen promoted by Pt. Nevertheless, MV2+ was found to be unstable in this medium and its destruction is attributed to the side-reactions initiated by Hiads) species confined to the catalyst surface. More generally, these H a adsorbed intermediates, unavoidable in any catalytic hydrogen-formation process, may obviously contribute to other undesired short-circuit processes (e.g. if oxygen was produced at a neighbouring site). These reactions must be considered in the absence of any specific catalyst (such as Pt if H, and 0, are considered) for the recombination steps.The design of a new method of H, and 0, production using separate c~mpartments~~ is therefore highly attractive and will certainly be developed in the future. 108-23 340 RUO, AS A REDOX CATALYST S. Trasatti and W. E. O'Grady, in Advances in Electrochemistry and Electrochemical Engineering, ed. H. Gerischer and C. W. Tobias (John Wiley, New York, 1981), p. 177. ' L. D. Burke, 0. J. Murphy, F. F. O'Neill and S . Venkatesan, J . Chem. Soc., Faraday Trans. 1 , 1977, 73, 1659, and references cited therein. R. S. Yeo, J. Orehotsky, W. Wissher and S . Srinivasan, J. Electrochem. Soc., 1981, 128, 1900. S. Trasatti, J . Electroanal. Chem., 1980, 111, 125. L. D. Burke and J. F. Healy, J . Electroanal.Chem., 1981, 124, 327, and references cited therein. D. Galizzioli, F. Tantardini and S. Trasatti, J . Appl. Electrochem., 1974, 4, 57. M. Spiro, J . Chem. Soc., Faraday Trans. I , 1979, 75, 1507. A. Piccini and L. Marino, Z . Anorg. Allg. Chem., 1972, 32, 55. M. Spiro and A. B. Ravno, J . Chem. Soc., 1965, 78. lo J. Kiwi and M. Gratzel, Angew. Chem., Int. Ed. Engl., 1978, 17, 860. l 1 J. Kiwi and M. Gratzel, Angew. Chem., Int. Ed. Engl., 1979, 18, 624. l 2 A. Mills and M. L. Zeeman, J . Chem. SOC., Chem. Commun., 1981, 948. l 3 N. Sutin and C. Creutz, Pure Appl. Chem., 1980, 52, 2717. l 4 P. Keller and A. Moradpour, J . Am. Chem. Soc., 1980, 102, 7193 and references cited therein. l5 P. Keller, A. Moradpour, E. Amouyal and H. B. Kagan, Nouv. J . Chim., 1980, 4, 377. l6 J. M. Lehn, J. P. Sauvage and R. Ziessel, NOUD. J . Chim., 1980, 4, 355. l7 E. Borgarello, J. Kiwi, E. Pelizzetti, M. Visca and M. Gratzel, J . Am. Chem. Soc., 1981, 103, 6324. l8 E. Amouyal, P. Keller and A. Moradpour, J . Chem. Soc., Chem. Commun., 1980, 1019. l9 E. Amouyal, D. Grand, A. Moradpour and P. Keller, Nouu. J. Chim., 1982, 6, 241. * O See, for example, L. F. Fieser and M. Fieser, Reagents for Organic Synthesis (John Wiley, New York, 21 D. Galizzioli, F. Tantardini and S . Trasatti, J . Appl. Electrochem., 1975, 5, 203. 22 D. S. Miller, A. J. Bard, G . McLendon and J. Ferguson, J . Am. Chem. Soc., 1981, 103, 5336. 23 D. S. Miller and G. McLendon, J . Am. Chem. SOC., 1981, 103, 6791. 24 K. Kopple, D. Meyerstein and D. Meisel, J. Phys. Chem., 1980, 84, 870, and references cited therein 25 F. Chojnowski, P. Clechet, J. R. Martin, J. M. Herrmann and P. Pichat, Chem. Phys. Lett., 1981, 1967), p. 983. for similar processes. 84, 555. (PAPER 2/3 1 1)

ISSN:0300-9599    

DOI:10.1039/F19827803331

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans, I, 1982, 78, 3341-3356 Radiation Chemistry of Dilute Aqueous Solutions of Thallous Ion Formation of Colloidal Thallium and its Catalysis of the Reduction of Water by (CH,),cOH and CH$HOH Radicals BY GEORGE V. BUXTON* A N D TREVOR RHODES University of Leeds, Cookridge Radiation Research Centre, Cookridge Hospital, Leeds LS16 6QB A N D ROBIN M. SELLERS Central Electricity Generating Board, Berkeley Nuclear Laboratories, Berkeley, Gloucestershire GL 1 3 9PB Received 26th February, 1982 In the absence of 0, relatively stable (several hours to several days) thallium metal colloids are formed when dilute aqueous solutions of thallous ion ([Tl+l0 z 1.2 x lop4 mol drn-,) are irradiated (dose rate z 10 Gy min-I) under reducing conditions (lo-' mol dm-3 propan-2-01 or ethanol) in the pH range 6-12 in the presence of mol dm-, surfactant (sodium dodecyl sulphate or Triton-X-100).The colloid is also stable at pH 3, but it cannot be formed at this pH because H,O+ competes with TI+ for e&. Once nucleation has occurred, T1+ is reduced by (CH,)$OH and CH,cHOH at the particle surface. Electrophoresis measurements showed that the particles are negatively charged, and kinetic analysis indicated that their mean diameter ranges from 30 nm at pH 3.4 to 24 nm at pH 11.8 under the experimental conditions specified above. Increasing the dose rate or [Tl+] resulted in smaller particles being formed. Colloidal thallium catalyses the reduction of water by (CH,),cOH and (CH,)cHOH, and also catalyses the disproportionation of these radicals.A mechanism is proposed for these processes in which the rate-determining step for hydrogen production is the discharge of H,O+ (low pH) or H,O (pH 2 7) at the metal surface, and radical disproportionation involving electron transfer to and from the particle. The rate constant for the discharge of H,O is estimated to be 820 s-I. The radiation chemistry of aqueous colloidal systems naturally involves the reactions of free radicals at phase interfaces and is a topic of burgeoning interest from several standpoints. These range from the effects of a radiation field on physical properties, the subject of much of the early work,l and the radiation-induced dissolution of metal oxides,, which are of importance in water-cooled nuclear reactor^,^ to catalysis of the cleavage of water into hydrogen and oxygen by free radicals4 and the kinetics of electrode processes.It has been shown recently that ~ i l v e r , ~ ~ 6gold 7 , and cadmium9 metal sols catalyse the reduction of water by reducing radicals such as (CH,),cOH, CH$HOH and CO; -, and that silver sols catalyse multielectron reductions of inorganiclo and organic species.ll The noble metals behave as microelectrodes, accumulating electrons from the reducing radicals until they acquire a negative potential sufficient to cause evolution of hydrogen from water. In this paper we report a study of the preparation and properties of colloidal thallium which, like cadmium, is an electronegative base metal ( E e = -0.336 V). 33413342 RADIATION CHEMISTRY OF AQUEOUS TI+ EXPERIMENTAL All chemicals were AnalaR grade and were used as received.Solutions to be irradiated were made up in triply distilled water. 0, was removed by evacuation. Irradiations were carried out in conventional Pyrex cells fitted with an optical cell on a sidearm. A 6oCo pray source providing a maximum dose rate of 10 Gy min-' was used in the majority of the experiments; in a few cases higher dose rates were achieved using a d.c. beam of electrons from a 3 MV Van de Graaff accelerator. Gaseous products were collected by Toepler pump and analysed by gas chromatography. Solution products were determined by conventional g.1.c. methods. Electrophoresis measurements were made with a Rank Bros. mark I1 instrument using a cylindrical cell configuration. Changes of light transmittance by the colloidal solutions were measured as absorbance using a Pye Unicam SP 8-100 u.v.-visible spectrophotometer.The observed absorbance is due to light absorption and light scattering, but the relative contributions of these two processes to the observed absorbance spectrum have not been established in the present work. RESULTS FORMATION A N D SPECTRA OF THALLIUM SOLS Fig. 1 shows the development of the spectrum due to colloidal thallium formed by the radiolysis of oxygen-free 1.25 x lop4 mol dm-3 T1+ (as Tl,SO,) solution buffered at pH 7 with 4 x mol dm-3 phosphate and containing 10-1 mol dm-3 propan-2-01 and mol dm-3 sodium dodecyl sulphate (SDS). The surfactant is necessary to prevent coagulation and precipitation of the colloid.Similar results were obtained when SDS was replaced by Triton-X-100. No colloid was formed when the solution initially contained N 2 0 or acetone. The developing spectrum shows a gradual red shift with dose up to a dose of ca. 350 Gy followed by a sharper blue shift between 350 and 400 Gy, and no appreciable further change at larger doses. Fig. 2 summarises the spectral changes observed at various pH including neutral unbuffered solution. The colour of these sols was generally purple and the particles were too large to pass through a 25 nm millipore filter. The intensity of the weak absorption band at ca. 300 nm (fig. 1) was roughly proportional to dose at all pH, but the proportionality constant increased abruptly between pH 10 and 11. This is shown in the inset to fig.1. The sharp absorption peak at 216 nm in fig. 1 is due to T1+ and was used to measure the change in concentration of the ion during colloid formation. Different results were obtained depending on whether the absorbance at 216 nm was measured in the presence of the colloid or after precipitating it by repeatedly freezing and thawing the solutions. These are shown in fig. 3, where it is seen that there was a residual [Tl+] of ca. mol dmP3 after colloid formation. The smaller residual concentration observed in the presence of the colloid suggests that most of the ions are on or near the surface of the particles. Electrophoretic measurements showed that the particles were negatively charged whether the surfactant was SDS or Triton-X-100 when the anion was sulphate o r chloride.Electrophoretic mobilities were measured for the thallous chloride/Triton- X-100 system at natural pH and pH 1 1.8, and were 1.8 x lops and 8 x m2 s-l V-l, respectively. The < potential was calculated for these mobilities using the method of Wiersema et a1.12 was found to be 33 mV at natural pH and 83 mV at pH 11.8, corresponding toG . V. BUXTON, T. RHODES AND R. M. SELLERS 3 343 2.c 1 . 8 1.6 1 . 4 1 . 2 aJ K -e 1.0 s: D m 0.8 0.6 0 . 4 0.2 0 wavelength/nm FIG. 1 .-Development of the absorption spectrum of colloidal thallium with radiation dose. Initial conditions: 1.25 x lop4 mol dmp3 T1+, lo-' mol dmF3 propan-2-01, mol dm-3 C,,H,5S0,Na, buffered at pH 7 with 4 x mol dm-3 phosphate. Dose rate 10 Gy min-*. The dose is shown for each spectrum.Inset: Dependence on pH of the rate of increase of absorbance at 300 nm with dose. one negative charge per 70 and 7 nm2, respectively. These results are consistent with T1+ being incorporated into the metal lattice at the particle surface, any excess of anions being held at the surface by electrostatic and chemisorptive effects, and cations being concentrated in the diffuse region of the double layer. The effects of dose rate and the concentrations of TI+ and SDS on colloid formation were examined briefly. The effect of dose rate is shown in fig. 4 and, since smaller particles scatter light of shorter wavelengths, the blue shift in the absorption spectrum with increasing dose rate indicated the formation of smaller particles. The effect of [T1+] is shown in fig.5. By decreasing [SDS] in 1.25 x mol dm-3 T1+ solution it was found that as little as mol dm-3 SDS was sufficient to stablise the colloid. Because of these variations in colloid properties with experimental conditions, the standard conditions, except for pH, chosen for colloid preparation were as detailed for fig. 1 . Colloidal solutions at pH > 7 were obtained by adding NaOH to unbuffered solutions before irradiation, and at pH < 7 by adding HC10, to unbuffered solutions after irradiation.3344 RADIATION CHEMISTRY OF AQUEOUS T1+ FIG. 2.-Dependence on dose of A,,, of the absorption spectrum of colloidal thallium at natural pH (A), pH 7 (O), pH 10.8 (O), pH 11.8 (+) and pH 10.8 with ethanol instead of propan-2-01 (---). 1.25 1.00 m E 5 0.75 E P 0, +, k, 0.5 U --.- 0.25 I I I I I I I '0 d ose/G y FIG. 3.-Decrease in thallous-ion concentration with dose at pH 7 (A) and pH 10.8 (x), and at pH 7 after precipitation of the colloid by freezing and thawing (0). The solid circles represent [TI+], - 2[N,] (see text).G. V. BUXTON, T. RHODES AND R. M. SELLERS 3345 1 .o 0.9 0.8 0.7 0.6 0.5 0 . 4 0.3 0.2 0.1 0 a, E s1 D m I I I I I I I 2 09 400 600 8 00 1000 X/nm Fic. 4.-Effect of dose rate on the spectrum of colloidal thallium at natural pH. [Tl+], = 1.25 x lop4 mol dm--3, dose = 700 Gy. The dose rate (Gy s-') is shown for each spectrum. STABILITY O F COLLOIDAL THALLIUM In air-free solution the colloid was stable for several days at pH 7- 12, and for several hours at pH 3; in more acidic solution the metal dissolved more rapidly.The colloid could be formed at pH 13 but it precipitated as soon as the dose exceeded the formation dose. The colloid dissolved immediately when exposed to air and over a period of hours in the presence of N,O. In the latter case N, and T1+ were produced in the ratio of 1 :2 and the N, yield provided an accurate measure of the amount of T1+ that had been reduced to the metal. mol dm-3 H,O, was added to a sol containing 7 x lop5 mol dmU3 T1 under air-free conditions, the colloid dissolved completely at pH 7 but not at pH 1 1.8 where the colour of the sol changed from purple to yellow. These results indicate that H,O, reacts by different mechanisms in neutral and alkaline solutions and it seems likely that this is the cause of the higher rate of colloid formation at pH > 10 (see fig.1 inset). When 5 x FORMATION OF H, AND ORGANIC PRODUCTS Fig. 6 shows the dependence of the formation of H, and acetone on dose under the same conditions as for fig. 1. The data are characterised by an abrupt increase in G(H,) at the dose where colloid formation is complete from 1.0 to 2.5, indicating catalytic production of H,. Table 1 summarises data for propan-2-01 and ethanol solutions at various pH ; table 2 shows the effect of dose rate on G(HJ for propan-2-01 solutions. Both sets of data relate to yields obtained from linear yield against dose plots after colloid formation (see fig. 6).3346 RADIATION CHEMISTRY OF AQUEOUS TI+ 1 .c 0.9 0.8 0.7 0.6 2 0.5 B % 0.4 0.3 0.2 0.1 0 a, c L I I I I I I 1 1 0 LOO 600 800 1000 X/nm FIG.5.-Effect of thallous ion concentration on the spectrum of colloidal thallium at natural pH for a dose of 200Gy at a dose rate of 10Gymin-'. The concentration of T1+ (moldm-3) is shown for each spectrum. H, was generally measured within 15 min after irradiation, but an experiment at pH 1 1.8 showed that a small amount of the gas (cu. 2.5 x lop6 mol dm-3) was formed at longer times in a thermal reaction (see fig. 6 inset.) SILVER COLLOID For comparison purposes the yield of H, was measured for a colloidal silver solution at natural pH which initially contained 2 x mol dm-3 Ag+, lop3 mol dm-3 SDS, 3 x lo-, mol dm-3 acetone and 10-l mol dmP3 propan-2-01. For a dose rate of 9.7 Gy min-l, G(H,) = 3.38 0.08 after colloid formation was complete and is in good agreement with data from previous SILVER COLLOID COATED WITH THALLIUM It has been shown" that when a solution containing Ag+ and Tl+ is irradiated under conditions similar to the standard conditions described above, silver particles are formed first and subsequently become coated with thallium.Since silver particles under these conditions are much smaller than thallium particles, this experiment provides a method of generating effectively small thallium particles for the purpose of obtaining a relationship between Amax of the absorption spectrum and particle size. Fig. 7 summarises the results obtained on irradiating a solution containing 1.25 x mol dmP3 TI+, 1.6 x lop4 mol dm-3 Ag+, 10-l mol dmP3 propan-2-01, lo-, mol dmP3 acetone and The results show clearly the formation of the silver sol first with Amax at 390 nm, mol dm-3 SDS at natural pH.G .V. BUXTON, T. RHODES AND R. M. SELLERS 3347 250 200 m E E i -0 150 c OJ U 2 100 0 N 50 1 I I I I I I P / / Y. y I 2 '2; 6;. 110 1; time/min / 0 400 800 1200 1 dose/Gy 00 FIG. 6.-Yields of H, (a) and acetone (0) from colloidal thallium solution at pH 7 containing lo-' mol dm-3 propan-2-01. Dose rate = 10 Gy min-l. The broken line represents the theoretical yield of acetone (see text). Inset: Post-irradiation formation of H,. TABLE YIELDS OF PRODUCTS FROM THE RADIOLYSIS OF COLLOIDAL SOLUTIONS OF THALLIUM CONTAINING 0.1 mol dm-3 PROPAN-~-OL OR ETHANOL (DOSE RATE = 10 Gy min-l) colloid PH G(H2) G(organic product) propan-2-01 absenta presentb presentb presentb presentb ethanol absenta presentb acetone natural 1 .OO f 0.05 3.5 f 0.1 natural 4.99 & 0.02 4.90 + 0.02 7 2.50 & 0.04 2.5 + 0.1 10.7 2.1 fO.l 2.06 f 0.03 11.8 1.20 & 0.05 1.24& 0.10 10.8 1 .OO f 0.05 1.5 f 0.2 1.9 +O.1 10.8 1.39f0.05 0.25k0.05 1.10+0.05 butane-2,3-diol acetaldehyde a Solution saturated with N,O. Yields obtained from the slopes of linear yield against dose plots after formation of colloid was complete. followed by the coating with T1 which shifted Amax progressively to the blue until it reached 350 nm and remained constant when no further reduction of Tl+ occurred. When N,O was added to the colloidal system, the thallium metal dissolved as before and the silver sol remained with almost the same spectrum as the initial one.3348 RADIATION CHEMISTRY OF AQUEOUS TI+ TABLE 2.-EFFECT OF DOSE RATE ON G(H,) FROM COLLOIDAL SOLUTIONS OF THALLIUM CONTAINING 0.1 mOl dm-3 PROPAN-2-OL pH dose rate/Gy min-la G(H,) 3.4b7C 10 naturalb 10 7 10 2.0 2.0 1.9 0.7 10.8 9.7 2.0 2.9 & 0.1 3.25 f 0.03 2.28 & 0.02 2.76 0.02 2.50 L 0.04 3.15 & 0.05 3.28 & 0.02 2.1 kO.1 2.85 & 0.05 a Dose rate was varied after the formation of colloid was complete; 0.03-0.05 mol dm-3 acetone added after colloid formation (see Discussion section); pH adjusted with HClO, after 2 .o 0.5 a, 5 -E s: -2 1.0 0.5 0 I I I I I 1 1 200 300 400 500 600 h/nm FIG.7.--Spectrum of thallium-coated silver colloid (-), silver colloid (. . .), and silver colloid after dissolution of the thallium coating by N,O (---). [Ag+], = 1.6 x mol dm-3, [TI+],, = 1.25 x loe4 mol dm-3, natural pH, dose rate = 10 Gy min-I.G.V. BUXTON, T. RHODES A N D R. M. SELLERS 3 349 DISCUSSION MECHANISM OF THALLIUM COLLOID FORMATION The radiolysis of water is described by reaction (1) H20*ecq, H, OH, H,, H20,, H+ ( 1 ) and Geiq = Go, = GH+ = 2.7, G, = 0.6, GH2 = 0.4 and GH2O2 = 0.6 are the yields (molecules per 100 eV) of the primary species appropriate to our experimental ~0nditions.l~ The radicals formed in reaction (1) can react as follows ecq + T1+ -+ T1° (2) eiq + H+ -+ H (3) (4) ( 5 ) H+RHOH -+ROH+H2 (6) (7) (8) H2O ecq+RO-+RO- -+ ROH+OH- H,O eLq + N20 + N, + 6- -+ OH + OH- OH + RHOH + ROH + H 2 0 ROH + ROH + RO + RHOH ROH + ROH -+ HORROH (9) where RHOH represents propan-2-01 or ethanol, ROH is (CH),),oH or CH$HOH etc.Although OH is known not to abstract a hydrogen atom exclusively from the carbon atom carrying the hydroxy group,14 the product yields from N,O-saturated solutions of ethanol and propan-2-01 (table 1) are consistent with all other radicals being transformed into (CH,),eOH or CH,cHOH, through a radical transfer reaction with the parent alcoho1,15 before disproportionation or combination can occur. - E* for Tlo/T1+ is estimated to be - I .9 V,16 so that eGq ( E e = - 2.7 V) is the only species capable of reducing T1+ since E* for (CH,),cOH and CH$HOH are - 1.5 and - 1.1 V, respectively.li This is confirmed by our observation that T1+ was not reduced when the solution initially contained acetone or N,O, i.e. when reaction (2) was suppressed by reactions (4) or (5).However, fig. 3 shows that G( - TI+) increased after some colloid had formed, and this may be explained by a lowering of the reduction potential of T1+ on or near the metal surface so that reaction (10) can occur ROH + TI&rface -+ T1:urface + RO + H+. (10) The larger initial G( - T1+) at pH 10.8 (fig. 3) is probably due to ionisation of a fraction (ca. 4%) of the (CH,),COH radicals (pK = 12.29 since ES for (CH,),cO- is estimated to be - 2 . 2 V.17 CATALYTIC ACTION OF COLLOIDAL THALLIUM Once the colloid has been formed, all the radicals produced in reaction (1) form ROH through reactions (4), (6) and (7). The data in table 1 show that at this stage the overall reaction taking place is RHOH + RO+H,3350 RADIATION CHEMISTRY OF AQUEOUS TIS i.e.the thallium colloid catalyses the dehydrogenation of the alcohols. In reality, reactions (8) and (9), which occur in the bulk solution, are replaced to some extent by reaction (I 1) T1 ROH + ROH + 2R0 + H, (1 1) but half of the ROH are generated from the product RO through reaction (4). In unbuffered solution at natural pH, H+ increases through reaction (10) and reaction (3) becomes faster than (4). Thus RO is protected but an equivalent amount of H, is formed via reactions (3) and (6) (see table 1). Nevertheless, the yield of catalytic H, can still be measured in acidic solution by adding sufficient acetone to the colloidal solution to ensure that reaction (3) does not occur (see table 2). The extent of reaction ( I I ) is given by for which the maximum value will be G(H2)excess = .fG(ROH) = 3.0.In practice the largest measured G(H2)excesS for thallium colloid was 2.3 (see table 2) compared with 2.4 for the silver colloid. In the latter case it was proposed5 that the shortfall in G(H,) is due to reoxidation of the metal by H,O, in a two-electron process. Such a reaction would explain the H, yields at pH 7 in the thallium case and also account for the complete dissolution of 7 x mol dm-3 T1 by 5 x mol dm-3 H,O, that was observed. At high pH, however, both the effect of added H,O, and the increased rate of colloid formation are consistent with H,O, acting as a one-electron oxidant above pH 10 (see fig. 1 inset). This change in mechanism with pH might be associated with the ionisation of H,O, (pK = 1 1.75) at the particle surface, since the electrophoresis measurements are consistent with the hydroxide-ion concentration being higher at the surface than in the bulk solution. The overall result of one-electron oxidation by H,O, is its catalytic decomposition through reaction (12) followed by (7) and (10) (12) With increasing pH G(H2)excess becomes quite small, indicating that reaction (1 1) is no longer efficient.Similar observations have been reported for silver5 and gold8 sols, and have been ascribed5* to the catalysed disproportionation of ROH througn one radical donating an electron to, and the second radical accepting an electron from, the metal particle. The data for propan-2-01 in the absence of colloid (see table 1) show that (CH,),COH only disproportionates in the bulk solution [reaction (S)], whereas for ethanol 43% of CH3cHOH combine [reaction (9)]. In the latter case, therefore, catalysed disproportionation will show up as an increase in the yield of acetaldehyde relative to butane-2,3-diol.Thus the data for ethanol (table 1) provide clear evidence for the catalysed disproportionation of ROH at high pH. H,O, + T1+ Tl+ + OH + OH-. SIZE OF THE COLLOID PARTICLES In order to elucidate the reactions of ROH at the colloid surface it is necessary to know what fraction of the radicals actually reach the surface, which in turn requires a knowledge of the size and shape of the particles. Because of the instability of colloidal thallium in air we have been unable to examine the sol by electron microscopy, and attempts to use light-scattering methods were unsuccessful.We believe, however, thatG. V. BUXTON, T. RHODES A N D R. M. SELLERS 3351 meaningful average particle sizes can be obtained using a kinetic method as outlined below. We begin with the propan-2-01 system at pH 3.4 and assume that (i) all (CH,),COH reaching the particle surface form H,; (ii) the radicals donate an electron on the first encounter (i.e. the reaction is diffusion controlled); (iii) the particles are uniform spheres and have the same density as thallium metal (p = 11.85 g cm-,); (iv) the competing reactions involving (CH,),t)OH are 2 (CH,),COH -+ disproportionation in the bulk solution (CH,),t)OH -+ diffusion to the particle surface. (1 3) (14) Then under steady-state conditions for (CH,),COH (R) d[R1 = 0 = G(R) d - 2kl3[RI2 - k14[R] [PI (1) dt where G(R)d is the rate of formation of R [G(R) is the yield of R and d the dose rate in appropriate units], and P represents a thallium particle. When reaction (14) is diffusion controlled 4nrP D, N 1000 N - since rp % rH.and DR % D,, where r and D are radius and diffusion coefficients, respectively, and N is Avogadro's number. The concentration of particles is given by MTlO1 [PI = ~ Nm where A4 is the molecular weight of T1 and rn = g r i p is the mass of a particle. Hence 3M[T1°] DR 1OOOrbp ' k14PI = The fraction,f, of radicals that reach the particle surface is given by and eqn (I) and (11) can be solved for kI4[P], and hence rp. Taking D, = cm2 s-l, 2k1, = 1.5 x lo9 dm3 mol-l s-l,19 GHtO, = 0.6 and G(R) = 6.0 gives rp = 15.4f 1.1 nm and 15.5 kO.9 for dose rates of 10.0 and 2.2 Gy min-l, respectively.The agreement between these two values indicates that the assumptions made in their calculation are valid. The data for the ethanol system can be treated in an analogous way. In this case not all the radicals reaching the surface form H,, but the alternative assumption made is that butane-2,3-diol is only formed in the bulk solution and is therefore a measure of the number of radicals that combine or disproportionate in the bulk solution.3352 RADIATION CHEMISTRY OF AQUEOUS TI+ Taking D, = cm2 s-l and 2k, = 2k, = 2.3 x lo9 dm3 mol-l s-' gives rp = 14.4 0.9 nm, which is in good agreement with the values of rp calculated from the propan-2-01 data. In fact, this slightly smaller value for the ethanol system is consistent with the blue shift in A, (see fig.2). The data for the silver sol coated with thallium (see fig. 7) provide further information on the relationship between A, and particle size for colloidal thallium. Since the exterior of these particles consists of a shell of thallium they are expected to have the optical properties of a small thallium particle. The radius of the silver particles is taken to be 7 nm6 and the thallium coated particles are calculated to have a radius of 8.7 nm. Fig. 8 shows that the limited amount of data available conform to a linear relationship between A, and particle size. 81 I I I I 300 400 500 600 700 ~rnaxlnm FIG. 8.-Dependence of A,,, of the spectrum of colloidal thallium on particle size (see text).OXIDATION A N D REDUCTION OF R o H AT THE PARTICLE SURFACE Knowing the size of the particles, it is possible to calculate the yield of radicals that donate, G(R),,, or accept, G(R)red, electrons at the surface using the following relationships for pH < 10: and the same relationships without the term GHZo2 for higher pH. The results obtained, using the data in fig. 8 to estimate particle size, are listed in table 3. KINETICS OF SURFACE REACTIONS It is clear from the data in tables 2 and 3 that colloidal thallium catalyses both the reduction of water by ROH and their disproportionation, the latter reaction becoming more important in alkaline solution. These features relate to the mechanism of hydrogen evolution on a thallium surface.G . V. BUXTON, T.R H 0 4 E S AND R. M. SELLERS 3353 TABLE 3.-CALCULATED YIELDS OF (CH,),t'OH RADICALS WHICH ARE OXIDISED OH REDUCED AT THE COLLOID SURFACE 3.4 15.5 10.0 2.0 natural 15.5 10.0 2.0 7 13.8 10.0 1.9 0.7 10.8 12.2 10.0 2.0 11.8 11.8 10.0 10.8" 14.4 10.0 0.83 0.95 0.83 0.95 0.88 0.97 0.99 0.92 0.97 0.93 0.8 1 5.0 5.7 4.35 5 . 4 4.75 5.67 5.85 3.91 5.72 3.0 2.82 0 0 0.60 0.35 0.55 0.17 0.09 1.60 1.05 2.58 2.04 " Data for CH,t'HOH. The process of hydrogen disLharge at a cathodically charged metal is generally assumed to be as follows: Volmer or discharge reaction M+H++e- + M-H M + H,O + e- --+ M-H +OH- (16) followed by Heyrovsky or electrochemical reaction M-H+H++e-+M+H, M-H + H,O+e- + M + H, + OH- (18) M-H+M-H +2M+H, (19) where M-H represents a hydrogen atom adsorbed on the surface of the metal.Theories put forward to account for hydrogen overvoltage are based on one of these three reactions being rate limiting. We show below that our data for colloidal thallium are consistent with reaction (1 6) being the slow step in neutral and alkaline solution. If we represent a thallium particle in its equilibrium state as Tl,, having T1+ on or near the surface, then the reactions in which electrons are removed from the particles can be written as or Tafel or catalytic reaction -TI+ kOH + Tl,,, + T1, + RHOH +OH- (20) HzO -,TI+ 2H,O + TI,,, -+ Tl,-l + 20H-+ H, -ZTP H,O, + Tl,,, -, Tl,-l + 20H-3354 RADIATION CHEMISTRY OF AQUEOUS TI+ and reactions in which electrons are donated to the particles as T1+ ROH + Tl,-, + T1, + RO + H+ Tlf ROH +TI, + Tl,,, + RO + Ht.T1, and Tl,-, may be regarded as oxidised, and Tl,,, as reduced, forms of the colloid. Under catalytic conditions where the rates of the charging and discharging reactions are equal, the steady-state approximation may be applied to ROH, Tl,-,, Tl,,, and H,O,. This treatment leads to the following expression for k21: and the same expression without the term 2GHzO1 at high pH when H,O, is a one- electron oxidant. The assumptions made in deriving eqn (111) are that k,, = k,, = k,, (all diffusion controlled) and The mechanism of reaction (21) may be written as follows: Volmer reaction -T1+ H,O + Tl,,, + TI,--H +OH- T1+ OH-+ TI, - H + Tl,,, + H,O Heyrovsky reaction -T1+ H,O + Tl,-H + Tl,-, + H, + OH- Application of the steady-state approximation to TI,-H gives eqn (IV) for k,, k25 k - 21 - 1 +k,,[OH-]/k,,’ Values of k,, can be determinted from experimental quantities using eqn (111) and hence k,, and k,,/k,, can be calculated using eqn (IV).Fig. 9 shows a comparison of the values of k,, determined experimentally with calculated values obtained using eqn (IV) with k,, = 8.2 x 10, s-l and k26/k2, = 8.2 x lo3 dm3 mol-l. The two sets of data agree reasonably well over the pH range 7-1 2 and for dose rates 0.7- 10 Gy min-’, showing that the proposed mechanism of H, evolution is consistent with the experimental data. We conclude, therefore, that the Volmer reaction (16) is the rate-determining step for thallium. This is not unexpected, since thallium has a relatively high hydrogen overvoltage and low adsorption energy for hydrogen atoms.20 In the kinetic analysis outlined above it is implicitly assumed that [T1,-HI 6 [Tl,,,], and therefore that reactions between ROH and TI,-H can be neglected.This assumption seems to be justified by the data in fig. 9, and is easily rationalised if reaction (25) is the rate-determining step.G. V. BUXTON. T. RHODES AND R. M. SELLERS 3355 7 8 9 10 11 12 PH FIG. 9.-Dependence of k,, on pH for dose rates of 10 Gy min-I (O), 2 Gy min-I (A) and 0.7 Gy min-I ( x ). The solid circle represents k,, when propan-2-01 is replaced by ethanol (dose rate 10 Gy min-I). The solid line is calculated with k,, = 820 s-I and k,,/k,, = 8.2 x lo3 dm3 mol-' (see text). In acidic solution it is expected that reactions (28) and (29), equivalent tc reactions (15) and (17), will occur, so that eqn (IV) simplifies to eqn (V) -Tlf H++Tl,+, -, Tl,-H -TI+ H++Tl,-H -P Tl,-,+H, This accounts for the fact that G(H,) is larger at pH 3.4 than at pH 7 (see table 2) since reaction (20) will be less able to compete with reaction (21) under these conditions.In calculating the particle size at pH 3.4 we,assumed that reaction (20) did not occur at all, which implies that the diffusion of ROH to the particle surface is the rate-determining step. The colloidal thallium particles may be treated as microelectrodes by analogy with other colloidal systems.,, 9 9 21 It has been pointed out2, that the microelectrode will acquire a mixed potential under steady-state conditions corresponding to equal anodic and cathodic currents. In the present case the anodic current is generated by reaction (24) and the cathodic current by reactions (20) and (21), which are in competition.Thus the mixed potential attained by the particle will be independent of pH below the pK of the radicals as long as reaction (20) takes place. This differs from the usual electrochemical situation where the hydrogen evolution reaction is the sole cathodic reaction and is pH dependent.22 Henglein23 has discussed redox reactions of free radicals at electrodes in terms of the distributions of their electronic energy levels. Catalytic disproportionation of (CH,),cOH and CH,cHOH and H, formation at a thallium particle requires that the energy distributions of occupied and unoccupied electronic levels in the radicals overlap with one another at a potential where reduction of water can occur.This3356 RADIATION CHEMISTRY OF AQUEOUS TIS potential is -0.7 V at pH 11.8, so the thallium particles must acquire at least this value. As noted earlier, the values of Ee for (CH,),cOH and CH,cHOH are - 1.5 and - 1.1 V, respectively,16 so that the particles will acquire a negative potential which is lower in the ethanol system than in the propan-2-01 system. This explains the observation (see tables 1 and 3) that reaction (21) competes less with reaction (20) at pH 10.8 when ROH is CH,cHOH than when it is (CH,),cOH. An implication of this interpretation of the surface disproportionation of ROH is that these radicals also disproportionate in the bulk solution by an electron-transfer mechanism, as has been suggested before.,, The slow post-irradiation formation of H, (see fig.6 inset) at pH 1 1.8 indicates that the thallium particles store an excess of electrons during irradiation, as has been observed for silver,6 gold8 and cadmium sols. From the amount of H, formed (2.5 x lod6 mol dm-,) the quantity of stored charge, Q,, at the steady state is ca. 0.5 C dmP3 when the charging rate is 0.6 C dmP3 min-l. This value of Q, for thallium can be compared with the values for silver6 (0.8 C dm-3), golds (0.3 C drn-,) and cadmiumg (1.4 C dm-3). However, as in the case of ~admium,~ it is more appropriate to consider stored charge on thallium particles as stored reduction equivalents since stored electrons will be neutralised by residual thallous ions. T. R. thanks the S.E.R.C. for a CASE studentship, and R. M. S. thanks the C.E.G.B. for permission to publish this paper. 1 2 a 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 M. Haissinsky, Nuclear Chemistry and Its Applications (Addison-Wesley, Reading, Mass., 1964), p. 396. G . V. Buxton, T. Rhodes and R. M. Sellers, Nature (London), 1982, 295, 583. R. M. Sellers, Central Electricity Generating Board Report no. RD/B/N3707 (1976). K. Kalyanasundaram and M. Gratzel, Angew. Chem., Int. Ed. Engl., 1979, 18, 701, and references therein. A. Henglein, J . Phys. Chem., 1979, 83, 2209. A. Henglein and J. Lilie, J . Am. Chem. Soc., 1981, 103, 1059. D. Meisel, J . Am. Chem. Soc., 1979, 101, 6133. J. Westerhausen, A. Henglein and J. Lilie, Ber. Bunsenges. Phys. Chem., 198 1, 85, 182. A. Henglein and J. Lilie, J . Phys. Chem., 1981, 85, 1246. A. Henglein, Ber. Bunsenges. Phys. Chem., 1980, 84, 253. A. Henglein, J. Phys. Chem., 1979, 83, 2858. P. H. Wiersema, A. L. Loeb and J. T. G. Overbeek, J. Colloid Interface Sci., 1966, 22, 78. I. G. Draganic and Z. D. Draganic, The Radiation Chemistry of Water (Academic Press, London, 1971). K-D. Asmus, J. Mockel and A. Henglein, J . Phys. Chem., 1973, 77, 1218. C . E. Burchill and I . S . Ginns, Can. J . Chem., 1970,48, 1232; 2628. J. Butler and A. Henglein, Radiat. Phys. Chem., 1980, 15, 603. M. Breitenkamp, A. Henglein and J. Lilie, Ber. Bunsenges. Phys. Chem., 1977, 81, 556. K-D. Asmus, A. Henglein, A. Wigger and G. Beck, Ber. Runsenges Phys. Chem., 1966, 70, 756. M. Simic, P. Neta and E. Hayon, J . Phys. Chem., 1969, 73, 3794. B. E. Conway and J. O’M. Bockris, J . Chem. Phjx, 1957, 26, 532. D. S. Miller, A. J. Bard, G. McLendon and J. Ferguson, J Am. Chem. SOC., 1981, 103, 5336. D. S. Miller and G. McLendon, J . Am. Chem. Soc., 1981, 103, 6791. A. Henglein, in Electroanalytical Chemistry, ed. A. J. Bard (Marcel Dekker, New York, 1976), vol. 9, p. 163. (PAPER 2/352)

ISSN:0300-9599    

DOI:10.1039/F19827803341

出版商:RSC    

年代:1982

数据来源: RSC