摘要:

J. Chem. SOC., Faraday Trans. 1, 1988, 84(2), 647-656 Differential Thermo-osmotic Permeability in Water-Cellophane Systems Crist6bal Fernandez-Pineda and M. Isabel Vazquez-Gonzalez Department of Physics, Faculty of Sciences, University of Malaga, Malaga, Spain The thermo-osmosis of pure water through two different cellophane membranes at temperatures ranging between 33 and 47 "C has been studied. Two types of experiments have been carried out: one to determine the dependence of the phenomena on stirring rate, with the average temperature and the temperature difference between the two bulk phases being kept constant while varying the stirring rate; the other varying the temperature difference between the two bulk phases, with the temperature of the cold bulk phase being kept as close to a fixed value as possible and employing a constant stirring rate.In all experiments the thermo-osmotic flow was from the hot to the cold side, and thermo-osmotic permeability was found to decrease with the mean temperature. The experimental values of the global thermo-osmotic permeability were corrected by taking into account the temperature polarization, and from these corrected values the differential thermo-osmotic permeabilities were calculated for the two membranes employed. The relationship between the differential thermo-osmotic permeability and the temperature was found to be different for the two membranes, being linear for the 600P membrane and quadratic for the 500P membrane. However, within the range of temperatures used in the present study, in each case, the differential thermo-osmotic permeability decreased with temperature.For comparison, the differential thermo-osmotic per- meabilities of two analogous membranes mentioned in the literature were calculated, and their behaviour was shown to be very similar to those of the membranes used in the present study within the same temperature range. The phenomena of mass transport through a membrane due exclusively to a temperature difference between the two membrane surfaces is known as thermo-osmosis and has been widely studied.'Y2 A comparison of the published results of a large number of authors, the majority of whom employed cellophane or cellulose acetate membranes with different acetyl contents, reveals qualitative and quantitative differences between the transport coefficients.The recorded differences in mass transport behaviour arise from the different nature of the membranes, and some authors have demonstrated that thenno-osmotic phenomena can only occur in suitably dense membrane^.^-' However, recent reports show that thenno-osmotic flows across porous membranes can Thermo-osmosis is also reported in systems employing 'charged ' or ' activated ' membranes and electrolyte solutions. 13-16 Inspection of the literature also reveals contradictory results for analogous membranes. Carr and S01lner~~ state that thermo-osmotic flow across cellophane membranes is undetectable when water and electrolyte solutions are employed, while Alexander and Wirtz,l* Rastogi et al.,374 Hasse et al.,19-22 and more recently Vink and C h i ~ t h i ~ ~ have obtained marked thermo-osmotic effects in experiments with similar cellophane membranes.Rastogi et al.,374 working with Du Pont 600 membranes, that had first been washed in progressive dilutions of NaOH, then in dilute HCl, and finally repeatedly in conductivity water, found that the thermo-osmotic permeabilities for water 647648 Thermo-osmosis in Water-Cellophane Systems are independent of AT, up to AT = 19 "C, when measurements are made at a constant mean temperature. In ref. (3) the reported constant mean temperature was 40 "C. When the measurements are made at a constant temperature difference, AT, thermo-osmotic permeability decreases as a function of mean temperature up to a minimum value and thereafter increases with temperature.In ref. (3) again, AT = 14 "C, the mean temperature ranged from 40 to 55 "C and the minimum value of the thermo-osmotic permeability was near 51 "C. This behaviour is contrary to that reported by Haase et al.,19-22 who used cellophane membranes with water as the permeant. They employed two types of membranes in ref. (19), 0.3 and 0.6 kg m-2, manufactured by Kalle and Co. (Wiesbaden). For some experiments, these were pretreated by washing in water and subsequently impregnated with a deposit of copper ferrocyanide. In ref. (20)-(22) the membranes were manufactured by Kalle AG Wiesbaden-Biebrich. In all these published studies the thermo-osmotic permeabilities measured at constant AT values were found to decrease with average temperature, with the flux going from the hot to the cold side of the membrane.Moreover, these authors found20-22 that above t,, x 55 "C the direction of the flux changes, i.e. water goes from the cold to the warm compartment. The measurements were made within two temperature ranges, between 11 and 77 "C20 and between 10 and 90 0C.21v22 In ref. (20)-(22), AT x 1.3 "C. Most of the reported differences have been explained by Belluc~i~~ using the concept of temperature polarization, first suggested by Vink and C h i ~ t h i ~ ~ and used implicitly by Dariel and Kedem.5 However, because a number of contradictory observations of thermo-osmotic phenomena taking place in cellophane membranes have not been completely explained or eliminated, we have designed new experiments to measure thermo-osmotic phenomena in the water-cellophane systems, to provide more information on these conflicting results, and these are reported in the present paper.The first set of experiments attempted to study the influence of stirring rate on thermo- osmosis in order to obtain the correction factor introduced by Belluc~i~~ and to calculate the true global thermo-osmotic permeability coefficient, which could then be used to interpret experimental results for all values of the temperature difference between the bulk components and all average temperatures. In the second set of experiments, using a constant stirring rate, the bulk temperature in one of the half-cells was varied while maintaining the bulk temperature in the other almost constant. Values of the thermo-osmotic permeabilities obtained for each experimental run were then corrected by the method explained above, and from these the differential thermo-osmotic permeability, b( t), was calculated by a procedure analogous to that used to obtain the differential diffusion c o e f f i ~ i e n t ~ ~ ! ~ ~ and previously used for other types of measurements in thermo-osmotic experiments.' Experimental Membranes Two different commercial membranes were used, 600P and 500P, supplied by Cellophane Espaiiola S.A. Both membranes were pretreated before the experiments by being washed in twice-distilled and deionized water for 72 h. The air remaining in the membrane matrix was eliminated by immersing the samples in water in an Erlenmeyer flask and submitting them to a partial vacuum provided by a Buchner funnel attached to a water tap. The thickness, 6, of the wet membranes was determined by a Millitron-Compact measuring instrument (1202-IC) to an accuracy of f 2 pm.The fractional void volume, E , was determined by the procedure described by Fernindez-Pineda and Serrano;27 the results are shown in table 1.C. Fernandez- Pineda and M . I. Vazquez-Gonzakz Table 1. Thickness (S), density @) and fractional void volume ( E ) of the membranes employed membrane 6/pm p/103 kg m-3 & 600P 6 2 f 4 1.49 k 0.23 0.75 +O. 14 500P 51 _+4 1.26 _+ 0.15 0.74 2 0.10 649 Permeant Pure water (twice-distilled and deionized) was used for all experiments. To eliminate dissolved air and avoid the possible formation of air bubbles in the thermo-osmotic cell, the water was boiled and then filtered through a Millipore filter of 0.45 pm nominal pore size before being introduced into the measuring apparatus.Apparatus All the measurements were made with an experimental device similar to those employed by Mengual et a1.' and Fernandez-Pineda and Ser~ano.~' It is a cell which consists of two equal 0.1 m long cylindrical chambers with a diameter of 0.06 m. They are surrounded by concentric, cylindrical walls of the same length but with a diameter of 0.09 m. The arrangement was temperature-stabilized by circulating water between the two concentric cylinders, and each water jacket was connected to a different thermostat. The water was stirred by a chain-drive magnetic-cell stirrer assembly2' to ensure the uniformity of water temperatures in each chamber.Each chamber was connected to a glass tube placed vertically and inserted in the superior part of the half-cell. Each chamber, with its corresponding glass tube, could be filled separately by means of a two-way stopcock on the rubber tube connecting the chamber with a water reservoir located at a higher level. Each arrangement, consisting of the half-cell, glass tube, rubber tube and water reservoir was enclosed in its own large air-bath which provided a temperature-controlled environment. The membrane was placed between the two chambers in a methacrylate holder with two stainless-steel grids between which the membrane was fixed. The membrane surface area exposed to the flow was q = (20.4k0.6) x m2, and the cross-sections of the glass tubes were (20k 1) x lo-' and (1 9 3) x m2 (determined by the mercury-drop technique).The glass tubes were pretreated with Rhodorsil-240 obtained from Rh6ne-Poulenc to prevent meniscus formation and to obtain a flat interface which would facilitate accurate measurement of the differences in pressure. Measurement of Thermo-osmotic Permeability The evolution with time of the difference in hydrostatic pressures, AP, in the thermo-osmotic experiments towards steady state is determined by an equation of the where and AP AP APm B ATb '- ATb ATb A y =--, y ---A, y =--=- where t is the time, ATb is the temperature difference measured between the bulk phases, APo and APm are the differences in hydrostatic pressure at t = 0 and t = GO, respectively,650 Thermo-osmosis in Water-Cellophane Systems A and B are integral, global or average phenomenological coefficients termed permeability and thermo-osmotic permeability, respectively, qo is the cross-section of the glass tube used for the pressure measurements, 6 is the membrane thickness, g is the acceleration of gravity, M is the molar mass and q is the membrane surface area exposed to the flow.In each experimental period of 7-10 days, the meniscus heights of the water in the glass tubes were measured, using an Ealing cathetometer with 1 x lo-' m accuracy, in runs comprising at least 15 different periods. The flows, in all cases, were from the hot to the cold side of the membrane. The sequences of height differences obtained were transformed into units of pressure which were subsequently divided by the corresponding temperature difference and fitted to a curve with the shape of eqn (1) following the method proposed by Lybanon?' From these fits values of z and yW were obtained.The hydraulic permeability, A, was obtained from z through eqn (2b), then the value of B was obtained from A and ym using eqn (2a). To eliminate errors originating from a bulging of the membrane that occurred at the beginning of each experiment, the initial pressure difference, APo(t = 0), used in the calculations was always greater than the initial experimental value. The values of AP, ranged between 1.5 x lo-, and 2 x lo-, rnH,O.T Results and Discussion In a previous phase of this experimental work, the relationship between the temperatures at which the thermostats were set, the chamber water temperatures and the stirring rates were investigated.'* 239 30 The chamber water temperatures were measured by platinum resistance thermometers (100 Q at 0 "C).The stirring rate was measured with a digital tachometer (On0 Sokki, model HT-430). Two types of experiments were carried out. In one we kept constant the temperature values of the thermostats (44.9 and 34.3 "C for the 600P membrane; 47.2 and 3 1.7 "C for the 500P membrane) and varied the stirring rate (between 0 and 346 r.p.m. for the 600P membrane and between 0 and 330 r.p.m. for the 500P membrane). In the other we kept constant the stirring rate (277 r.p.m. for the 600P membrane and 220 r.p.m. for the 500P membrane) and varied the temperature values of the thermostats.The results obtained for both membranes are analogous. Inspection of the data reveals the following. (a) The values of the differences in temperature between the water in the two bulk phases correspond closely to the temperature differences between the two thermostats. (The literature reports a greater divergence between the thermostat and chamber temperatures.'* 23) (b) For stirring rates > 80 r.p.m., ATb and Tb, the average temperature in the bulk, are constant for a given AT,(AT, = 10.6f0.2 "c, AT, = 8.5k0.2 "c and i+b = 39.4f0.2 "c for the 600P mem- brane; AT, = 15.5k0.2 "C, AT, = 12.2k0.2 "C and T b = 39.4k0.2 "C for the 500P membrane) within experimental error for the two membranes employed. At 0 r.p.m., for the 600P membrane ATb = 10.4k0.2 "c and Tb = 39.9k0.2 "c, while for the 500P membrane AT, = 12.8f0.2 "C and Tb = 41.0k0.2 "C.(c) For a given stirring rate, ATb depends on AT,. Two other groups of experiments were carried out. In the first, the measurements were made at different stirring rates for each membrane while maintaining constant the water temperatures in both chambers. For the 600P membrane these temperatures were t, = 43.7 "C and t, = 35.2 "C; for the 500P membrane they were C, = 45.5 "C and t, = 33.2 "C. In fig. 1 plots of I API us. time at different stirring rates are shown for a few illustrative cases. Tables 2 and 3 show the calculated values of A and B at different stirring rates for the 600P and 500P membranes, respectively. The values of the hydraulic permeability, A, oscillate about an average value in both cases.In each case a linear fit by the least- squares method of the data for A vs. stirring rate was obtained which gave correlation t 1 mH,O = 980.665 Pa.C. Fernandez- Pineda and M. I. Vazquez-Gonzalez 65 1 L I I I I 50 100 150 tlh Fig. 1. Evolution with time of the difference in hydrostatic pressure, I AP I, for different stirring rates: A, 198; 0, 314; x , 340 r.p.m. The solid lines represent the curves fitted to eqn (1) for membrane 600P. The water temperatures in each chamber were tl = 43.7 "C and f , = 35.2 "C. Table 2. Values of A and B at different stirring rates for membrane 600P (tl = 43.7 "C, f , = 35.2 "C and AT, = 8.5 "C) stirring B/ 10-lo mol m-' rate/r.p.m. A/10-12 mol s kg-' K-1 s-1 346 1.8f0.3 314 2.7 f 0.4 277 3.6 f 0.5 230 1.7 f 0.3 198 2.9 & 0.4 157 4.1 k0.8 109 2.8 f 0.4 3.7 k0.4 3.6 f0.4 3.4 f0.3 3.4 k0.3 3.2 k0.3 2.68 _+ 0.25 2.27 f 0.21 Table 3.Values of A and B at different stirring rates for membrane 500P (tl = 45.5 "C, t, = 33.2 "C and AT, = 12.3 "C) stirring B/ rnol m-' rate/r.p.m. A/10-12 mol s kg-' K-1 s-1 330 280 270 220 185 140 120 85 1.3k0.3 4.1 f 0.9 3.0 & 0.6 4.1 f 0.9 5.4k 1.1 2.3 f0.5 2.8 f 0.6 4.8 f 1.2 4.7 k 0.7 4.2 k 0.6 3.9 f 0.6 3.5 f0.5 3.3 & 0.4 3.1 f 0.4 2.7 k 0.4 2.2 f 0.3652 Thermo-osmosis in Water-Cellophane Systems coefficient values of -0.42 for the 600P membrane and -0.37 for the 500P membrane; their respective slopes were -4.12 x mol s kg-l. Student’s t- test was applied to determine whether these slopes were significantly different from zero.In both cases a significance level of 0.05 was set, so that the t-test gave strong evidence of zero slopes. This confirms that, within the range of stirring rates studied, the hydraulic permeability was independent of stirring rate7 and so may be considered to have a constant value. In each case the mean value was chosen : for the 600P membrane it was (2.8 0.4) x mol s kg-l. The values of the thermo-osmotic permeability, B, shown in tables 2 and 3 were calculated using the average values of A given above and those of I AP,/ATb I obtained from the fits of IAP/ATbI us. time data at each stirring rate, for both membranes. In the second group of experiments the water temperature in one of the bulk phases was varied while keeping almost fixed, in most cases, the water temperature in the other, and employing a constant stirring rate; for membrane 600P this was 277 r.p.m.and for membrane 500P it was 220 r.p.m. The evolutions of I AP I with time are analogous to those given in fig. 1 . From the values of z obtained from the fits, the values of the hydraulic permeability were obtained in the same way as in the previous set of experiments. Again it was found that the values of the hydraulic permeability, A , oscillated about an average value. A linear fit by the least-squares method was made to determine A as a function of the average temperature for each membrane. A correlation coefficient of - 0.38 was obtained for membrane 600P, and of -0.36 for membrane 500P; the values of their slopes were -4.65 x mol s kg-l, respectively.Student’s t-test was again applied: in both-cases a significance level of 0.05 was set, so that the t-test gave strong evidence of zero slopes. These results confirm that the hydraulic permeability was independent of the mean temperature and may be considered as constant within the temperature range studied. The average values for the 600P and 500P membranes were (2.04 _+ 0.1 1) x 10-l2 and (3.6 f 0.4) x The values of the thermo-osmotic permeability, B, were calculated from the average values of hydraulic permeability given above and values of I AP,ATb I obtained from the fits of I AP/ATb 1 us. time data at each value of ATb for both membranes. The values of A and B are shown in tables 4 and 5. and -6.09 x mol s kg-l, and for the 500P membrane it was (3.5 f 0.5) x and - 1.79 x mol s kg-’, respectively.The Correction Factor Following the procedure proposed by Belluc~i,’~ the ratio between the flow which would occur at an infinite stirring rate and the flow at a given stirring rate, R, for the same average temperature in the bulk fluids, i“,, and with the same temperature difference between them, ATb, is given by where (ATm)R is the actual temperature difference between the membrane sides at a stirring rate R, K is the thermal conductivity of the membrane-permeant system, 6 is the membrane thickness, h is the convection coefficient in the half-cells andfis a correction factor for the temperature polarization. Note thatfis independent of Tb and ATb, and thus the same factor may be used to interpret experimental results for all the values of ATb and Tb, with the sole restriction that the stirring rate, R, remains constant.Considering the empirical description given by Haase,28 eqn (3) is transformed (in the present case, the hydraulic permeability, A , is independent of the stirring rate) in theC. Fernandez- Pineda and M . I. Vazquez-Gonzalez 653 Table 4. Values of hydraulic permeability (in 10-l2 mol s kg-') and thermo-osmotic permeability (in 1 O-'O mol m-' s-l K-l) (experimental, corrected and calculated) for different temperature dif- ferences, AT,,, at a stirring rate of 277 r.p.m. for membrane 600P. AT,/"C AT," = ATb f '/"C A Bexptl B = f'exptl Bcalc, 36.8 - 33.4 38.3 - 33.6 39.8 - 33.9 41.8 - 34.3 43.6 - 34.7 45.3 - 34.9 40.8 - 37.2 42.3 - 37.6 43.9 - 37.9 45.8 - 38.2 42.7 - 39.3 44.4 - 39.7 46.0 -40.0 44.5 - 41.1 46.4 - 41.5 36.19 - 34.01 37.46 - 34.44 38.74- 34.96 40.45 - 35.65 42.00 - 36.30 43.43 - 36.78 40.15 - 37.85 41.46 - 38.44 42.82 - 38.98 44.44- 39.56 42.09 - 39.91 43.56 - 40.54 44.92 - 41.08 43.89 - 41.7 1 45.52 - 42.38 2.1 f0.5 2.6 f 0.4 2.2 f 0.3 1.6 f 0.2 2.3 f0.3 2.1 f 0.3 1.8f0.3 2.0 f 0.3 3.1 f0.3 2.3 f 0.3 1.5f0.2 2.4 f 0.4 1.9 f0.3 1.7f0.3 1.8f0.3 3.17 f 0.25 3.14f0.25 3.13 & 0.24 2.90 f 0.23 2.83 f0.22 2.62 f 0.21 2.61 f 0.24 2.64 f 0.2 1 2.44 f 0.18 2.27 f 0.18 2.42 f 0.19 2.30 & 0.19 2.17 f0.18 2.28 f 0.28 2.01 f0.16 4.9 f 0.8 4.9 f 0.8 4.9 f 0.8 4.5 fO.8 4.4 f 0.8 4.1 f0.7 4.1 f0.7 4.1 f0.7 3.8 f 0.6 3.5 f 0.6 3.8 f0.6 3.6f0.6 3.4k0.6 3.6 f 0.6 3.1 f0.5 5.1 4.9 4.7 4.5 4.2 4.0 4.3 4.1 3.9 3.6 3.8 3.6 3.4 3.4 3.2 ~~ a AT, is the temperature difference between the sides of the membrane.f is the correction fact or. Table 5. Values of hydraulic permeability (in 10-l2 mol s kg-') and thermo-osmotic permeability (in 1O-lo mol m-' s-' K-' ) (experimental, corrected and calculated) for different temperature dif- ferences ATb, at a stirring rate of 220 r.p.m. for membrane 500P. 35.9 - 32.3 37.4 - 32.6 39.1 - 32.8 40.8-33.1 42.6 - 33.4 44.3 - 33.9 45.5 - 34.1 39.7 - 36.3 41.6-38.1 43.5 -40.2 45.5-42.1 47.4 -44.2 35.05 - 33.14 36.28 - 33.72 37.63 - 34.27 39.00 - 34.90 40.45 - 35.55 41.87 - 36.33 42.83 - 36.77 38.90 - 37.10 40.78 - 38.92 42.73 -40.97 44.70 - 42.90 46.65 - 44.95 2.3 & 0.5 3.0 f0.7 6.2f 1.3 5.1 & 1.1 4.0 f 1 .O 5.3f 1.1 4.6f 1.1 3.5 f0.7 1.7 f 0.4 1.8 f 0.4 2.2 f 0.5 2.7 f 0.6 5.5 & 0.7 5.5 k0.7 4.8 f 0.6 4.7 k 0.6 4.3 f 0.5 4.1 f 0.5 4.0 f 0.5 4.6 & 0.6 4.0 f 0.5 4.1 f0.5 3.8 k0.5 3.8 f0.5 10.3 & 2.5 10.3 f 2.5 9.1 f2.2 8.8 f 2.1 8.1 f 1.9 7.8 f 1.9 7.6 f 1.9 8.6 f 2.0 7.6 f 1.8 7.8 f 1.9 7.1 f 1.7 7.2 f 1.7 10.4 9.8 9.3 8.9 8.5 8.1 8.0 8.3 7.6 7.2 7.1 7.2 To obtain (lAPm/ATbl)R=m a linear fit was made of l/(lAPm/ATbl)R against reciprocal stirring rate.The value of 0 r.p.m. was not used, as it would have modified the conditions of the same T,, or AT,,. The value of the intercept of the linear fit give the value of (AP,/ATb I)R=oo, which corresponds to an infinite stirring rate. From this value the correction factors for temperature polarization are obtained. For membrane 600P (stirring rate 277 r.p.m.) f= 1.56f0.09, while for membrane 500P (stirring rate 220 r.p.m.) f= 1.88 0.12.The values of the 'corrected' thermo-osmotic permeability, P O r r (= B), are calculated by multiplying the values of the experimental thermo-osmotic permeability, (BexptJR, by the correction factor. The results for each membrane are shown in tables 4 and 5. The ratio [AT,/(AT,),] (= the correction factor) was calculated in a similar way by 22 FAR 1654 Thermo-osmosis in Water-Cellophane Systems Mengual et al.,' while Tasaka and Futamura3' used a different method of estimation. In the case of cellulose acetate membrane^,^ the factors for the two membranes used were 1.41 and 1.45. For various types of membranes ranging in thickness from 3 x to 7 x m the values of the correction factor ranged from 1.45 to 1 .35.30 These values are of the same order of magnitude as those presented in this paper.Differential Thermo-osmotic Permeability To establish the relationship between the values of P O r r , obtained for different temperature differences between the two sides of the membrane, with the corresponding differential coefficient, b(t), the latter is expressed in a similar way to the differential diffusion coefficient : 7 9 25 1 I't, J - b( t ) d t p o r r - t2--1 t , where t, and t, are the Celsius temperatures on each side of the membrane. As the form of b(t) is unknown, it was assumed that a development in the form of power series may be admitted, following the method proposed by Hammond and Stokes:26 (6) where a,,, a,, a,, ...are constants. Then, substituting b(t) in eqn ( 5 ) by eqn (6), the following result is obtained : P O r r = a, + (a,/2) ( t , + t,) + (a,/3) ( t i + t, t, + tt) + (a3/4) (ti + ti t, + t, tt + t;) + (a4/5)(t; + ti t, + tf ti + t , t! + tt) + .-- (7) To determine the type of dependence that b(t) has on temperature, the data contained in tables 4 and 5 were fitted to an equation with the form of eqn (7) by multiple regression analy~is,~, in which the number of terms was varied from 2 to 5. Once the fits were made, to select the optimum, the following inspection procedure was used. (i) Observe the corrected determination coefficient for each fit, R2. (ii) Test, by applying the standard Student's t-test at the significance level a = 0.05, whether the individual coefficients may be assumed equal to zero for every fit.(iii) Test, by using the F-test at the same significance level, a = 0.05, whether all coefficients in the model are zero for each fit. When these foregoing tests had been applied, the best fit for the 500P membrane was found to be quadratic and the best fit for the 600 P membrane was linear. Table 6 shows the corresponding functions b(t)/6 for the two membranes. With the object of comparing the results of the present work with those reported in the literature for similar cellophane membranes, the differential thermo-osmotic permeabilities were calculated using the data of Rastogi et aL3 and of Haase et al.,l Note that the experimental values of thermo-osmotic permeability used in these two papers are not corrected for temperature polarization. All the values in the above published papers, and in the present work, are expressed as b(t)/6 in mol rn-, s-' K-l t o permit a comparative analysis between the results obtained from previous data and those of the present work. After carrying out the same type of fits and tests as for the membranes in the present work, the optimum fit for the membrane used by Haase et aL21 was found to be cubic [four terms in eqn (7)]; for the membrane used by Rastogi et aL3 it was quadratic [(three terms in eqn (7)].The results of the above optimal fits are shown in table 6. Inspection of table 6 (and the corresponding graphical representations) suggests the following. (1) Within the temperature range studied in the present work, all the differential coefficients decreased with increasing Celsius temperature.(2) The behaviour of membrane 500P, within the measured temperature range and when extrapolated, is similar to that reported by Rastogi et aL3 Both membranes show a decrease in the differential coefficient with increasing temperature for a given range of temperatures, b(t) = a,+a, t + a , t 2 + ...C. Fernandez- Pineda and M. I. Vazquez-Gonzalez 655 Table 6. Optimal differential thermo-osmotic permeability divided by membrane thickness, b(t)/6, for the different membranes as a function of Celsius temperature membrane b(t)/S/mol m-2 s-' K-' -~ ~ 600P 2.04 x -0.35 x 10-9 500P 6OO2l 6003 39.5 x 15.22 x 10-6t+ 149.2 x 10-'t2 14.3 x 10-5-5.90x 10-6t+67.5 x 1.75 x 10-5+0.11 x 10-6t-8.32 x 10-9t2- 14.57 x 10-12t3 after which they exhibit an increase with temperature.The values of the function b(t)/d for the 600P membrane resemble, within the measured temperature range and when extrapolated, the results of Haase et aL2' since they exhibit a decrease with increasing temperature (up to ca. 50 "C); they reach zero at between 50 and 60 "C, after which they change sign and continue to decrease with higher temperatures. (3) Above a temperature close to 50 "C, the behaviour of the b(t)/d functions for the membranes of Rastogi et aL3 and Haase et aL21 differs significantly. Those of Rastogi et aL3 exhibit a minimum and thereafter increase with temperature, while those of Haase et decrease to zero and continue decreasing with increasing temperature.In our opinion, these differences may result from the different membrane pre-treatment procedures employed by the various investigators. Note that it was not possible to obtain experimental results for temperatures outside the range 33-47 "C because the experimental periods required are so great that the membranes deteriorate. However, investigations using higher temperatures are planned for the future. References 1 2 3 4 5 6 7 8 9 10 11 12 13 14 I5 16 17 18 19 20 21 22 23 24 25 G. Lippmann, C. R. Acad. Sci. (Paris), 1907, 145, 104. M. Aubert, Ann. Chim. Phys., 1912, 26, 145. R. P. Rastogi, R. L. B. Blokhra and R. K. Agarwal, J. Electrochem. Soc., 1962, 109, 616. R. P. Rastogi and K. Sing, Trans. Faraduy SOC., 1966, 62, 1754, M. S. Dariel and 0. Kedem, J. Phys.Chem., 1971, 79, 1773. J. I. Mengual, J. Aguilar and C. Fernandez-Pineda, J. Membr. Sci., 1978, 4, 209. J. 1. Mengual, F. Garcia-Ldpez and C. Fernandez-Pineda, J. Membr. Sci., 1986, 26, 211. F. S. Gaeta and D. G. Mita, J. Membr. Sci., 1978, 3, 191. F. S. Gaeta and D. G. Mita, J. Phys. Chem., 1979,83, 2276. D. G. Mita, U. Asprino, A. D'Acunto, F. S. Gaeta, F. Bellucci and E. Drioli, Gazz. Chim. Ital., 1979, F. Bellucci, E. Drioli, F. G. Summa, F. S. Gaeta, D. G. Mita and N. Pagliuca, J. Chem. Soc., Faraday Trans. 2, 1979, 75, 247. F. Bellucci, E. Drioli, F. S. Gaeta, D. G. Mita, N. Pagliuca and D. Tomadacis, J. Membr. Sci., 1980, 7 , 169. J. W. Lorimer and S. H. Chan, International Symposium on Macromolecules, Helsinki, 1972, abstract J. W. Lorimer, in Charged Gels and Membranes II, ed. E. Stltgny (Reidel, Dordrecht, 1976), p. 76. M. Tasaka, S. Abe, S. Sugiura and M. Nagasawa, Biophys. Chem., 1977, 6, 271. M. Tasaka and M. Nagasawa, Biophys. Chem., 1978, 8, 111. C. W. Carr and K. Sollner, J. Electrochem. Soc., 1962, 109, 616. K. F. Alexander and K. Wirtz, Z . Phys. Chem., 1950, 195, 165. R. Haase and C. Steinert, 2. Phys. Chem. N.F., 1959, 21, 270. R. Haase and H. J. D. De Greiff, 2. Phys. Chem. N.F., 1965, 44, 301. R. Haase, H. J. De Greiff and H. J. Buchner, Z . Naturforsch., Teil A , 1970, 25, 1080. R. Haase and H. J. De Greiff, Z. Naturforsch., Teil A , 1971, 26, 1773. H. Vink and S. A. A. Chisthi, J. Membr. Sci., 1976, 1, 149. F. Bellucci, J. Membr. Sci., 1981, 9, 285. R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 2nd edn, 1970). 109,475. 11-39. 22-2656 Thermo- osmosis in Wa ter-Cellophane Systems 26 B. R. Hammond and R. H. Stokes, Trans. Faraday SOC., 1953, 49, 890. 27 C. Fernandez-Pineda and F. Serrano, J. Membr. Sci., 1984, 19, 309. 28 R. Haase, Thermodynamics of Irreversible Processes (Addison-Wesley, London, 1969). 29 M. Lybanon, An. J. Phys., 1984, 52, 22. 30 M. Tasaka and H. Futamura, J. Membr. Sci., 1986, 28, 183. 31 J. Kmenta, Elements of Econometrics (MacMillan, New York, 1971). Paper 71867; Received 18th May, 1987

ISSN:0300-9599    

DOI:10.1039/F19888400647

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1988, 84(2), 657-664 Spectrochemistry of Solutions Part 20.l-The Infrared, Near Infrared and Visible Spectra of Liquid Ammonia Joyce C. Dougal, Peter Gans* and J. Bernard Gill* Department of Inorganic and Structural Chemistry, The University, Leeds LS2 9JT The infrared, near-infrared and visible spectra of liquid ammonia at 293 K have been recorded. The fundamental and first six overtone regions have been assigned on the basis of the existence of four overtone progressions. The spectral data have been interpreted, together with other published spectral data, in terms of a model in which the local environment of the ammonia molecules is the predominating influence. Evidence is presented for the existence, in the liquid, of ammonia molecules with 0, I, 2 or 3 hydrogen bonds.In 1936 Costeanu and Barchewitz reported the spectrum between 600 and 950 nm of liquid ammonia at 198 K. They found two series of bands which could be assigned to overtones in anharmonic progressions, pointing to fundamental vibrations at 3333 and 3407 cm-1.2 Although many studies of the vibrational spectra of liquid ammonia have been reported ~ubsequently,~-ll this remarkable observation appears to have been neglected. Roberts and Lagowski examined one of the bands with the aid of curve resolution and difference spectroscopy, but they admitted that their analysis was not completely satisfactory, l2 Having already built a cell in order to examine the infrared spectra of electrolytes dissolved in liquid ammonia, we have now constructed a cell suitable for obtaining spectra in the visible and ultraviolet regions of the spectrum.By using both these cells, we have been able to measure the spectrum of liquid ammonia at room temperature from 200 to 8000 nm (50000 to 1250 cm-l). The infrared spectrum between 3000 and 3600cm-l has been analysed by curve resolution, but in the other region the band maxima were identified in the second or fourth derivative of the spectrum with respect to wavelength. We have found at least four overtone progressions. In the discussion we have attempted to collate all vibrational spectral observations by means of a model of liquid-ammonia structure which shows a significant and variegated degree of hydrogen bonding. Experiment a1 The cell used for recording infrared spectra has been described previo~s1y.l~ The cell used for recording the near-infrared and visible spectra was constructed to the same design as that of the infrared cell, the main differences being that the path length is fixed at 10 mm, and the windows are of quartz.A path length of 1 mm was obtained by inserting a 9mm quartz spacer into the cell. All spectra were recorded at ambient temperature. Infrared spectra were recorded as before.13 Below 4000 cm-l a path length of 12.8 pm was used. For the region 3048-3600 cm-l 25 independent scans were run at 15 min intervals, the infrared beam being blocked off between measurements to minimise sample heating. The 25 scans were co-added. The infrared spectrum was also recorded on a Nicolet 5-MX interferometer from 1200 to 4800 cm-' with a 50 pm path length.Between 4500 and 12 500 cm-I the spectra were recorded in sections on a Cary 14H 657658 Spectrochemistry of Liquid Ammonia spectrophotometer equipped with digital paper-tape output. The sample path length was 1 mm below 7200 cm-l and 10 mm above 7200 cm-'. The paper tapes were read into a BBC microcomputer recently programmed by us for use as an infrared data station with the SP3 infrared spectrophotometer. The spectra were co-added using the software facilities provided by that program. Between 2 and 6 scans were co-added for each group of bands. Spectra in the region 12500-50000 cm-l were recorded on a Pye-Unicam PU8800 spectrophotometer, connected to a BBC microcomputer for which we have written software with facilities similar to those available for the analysis of infrared spectra.Results The infrared spectrum between 3048 and 3600 cm-' is shown in fig. 1, together with the fit obtained with our VIPER ~r0gram.l~ The bands all have Lorentzian shape, and the second derivative of the fitted spectrum was in good agreement with the second derivatives obtained from the observed spectrum by the convolution method, except that the band at 3369 cm-' appeared as a shoulder in the convolution but as a resolved band in the derivative of the sum of Lorentzians. The band at 3339 cm-l is more than twice as broad as any other band in the fit. The spectrum observed can be classified into seven regions, corresponding to the fundamental and the first six overtones of the N-H stretching vibrations.In each region there is a fundamental or overtone complex, and a complex of combination bands of those vibrations and each of the two fundamental bending vibrations. Also, there are ternary combinations involving a stretching mode and both bending modes. The presence of underlying band maxima was ascertained by means of our MULTISMOOTH program,15 using the fact that the derivative of the spectrum with respect to wavelength is apparently more well resolved than the spectrum itself. Where possible the positions of the band maxima were determined by locating, using a least-squares method, the zero crossing points in first or third derivatives of the spectra. These peak positions are given in table 1. The spectrum in the first overtone region and its fourth derivative is shown in fig.2; Roberts and Lagoswki reported a band of similar general appearance.12 The spectrum and fourth derivative of the region involving the combinations of first overtone and bending modes is shown in fig. 3. The intensity of absorption decreases by a factor of ca. 20 for each step up the overtone sequence. Thus, whilst a 1 mm path length gave excellent spectra in the first overtone region, a 10 mm path length did not result in such a high signal-to-noise ratio in the second overtone region, and only the second derivative could be interpreted with confidence. Because the spectra obtained in the third and fourth overtone regions yielded rather noisy second derivatives, these regions could not be examined in detail. Whilst the higher overtones were only just detectable, as shown in fig.4, their positions could be established. Four series of overtone and combination bands have been assigned, designated by the subscripts a, b, c and d. 13 bands in the (a) series fundamental, overtones and combination bands were fitted by the method of unweighted least-squares to eqn (l), yielding the results shown in table 2. v = v, 0, +v: x,, +v2 v2 +v2 v, XZ3 +v, v, + v, v, X,,. The fundamental and overtone wavenumbers of the (b), (c) and ( d ) series were fitted to (2) eqn (2) : The values for the (a) series agree within 3 cm-l with those reported by Costeanu and Barchewitz,2 whilst for the (b) series the 198 K harmonic wavenumber is some 43 cm-l lower than the room-temperature value. v = vcu,+v2x33.J .C. Dougal, P. Gans and J . B. Gill 659 3600 3400 32 00 wavenum ber/ cm -' Fig. 1. Infrared spectrum of the fundamental N-H stretching region of liquid ammonia at 293 K, together with five Lorentzian bands used to fit the spectrum. For the combination bands with v, in the (b) series the two equations v = v v ~ ~ s + v , + v x , , (v = 2,3) (3) were solved for the two unknowns v, and A',,; v, and X34 were calculated in a similar manner. The band wavenumbers derived from these parameters are therefore equal to the observed wavenumbers. Bands at 6414 and 9404 cm-l have not been assigned; they could either belong to a fifth overtone progression or they could be combination bands derived from symmetric and antisymmetric stretching vibrations. Also, the combination bands at 4458 and 4914 cm-l do not fit into the progressions of combination bands with overtones, implying that they derive from different N-H stretching fundamentals.Discussion Very little is known about the structure of liquid ammonia. In crystalline amrnonial6* '' each molecule forms three N-H ... N bonds and three H - - . N-H bonds. However, although the hydrogen bonds are approximately linear, they are unusual in that the nitrogen atom's lone electron pair cannot lie on the bond axis. The structure presumably arises because the hydrogen bonding is relatively weak, so that packing requirements determine the overall structure. The structure is in marked contrast to that of crystalline water, in which there are an equal number of protons and lone pairs on each water molecule, resulting in a fully hydrogen-bonded structure.The ammonia molecule, with three protons and one lone pair, cannot form a regular hydrogen-bonded structure in three dimensions. X-Ray diffraction data obtained from the liquid have been interpreted'' in a way that indicates that each nitrogen atom has two distinct nearest nitrogen neighbours. Unfortunately, molecular-dynamics simulations have not yet yielded a model which reproduces the diffraction data adeq~ate1y.l~~ 2o All that could be saidlg is that the models examined agree in predicting a significant degree of association in the liquid. This is consistent with the quantum-mechanical calculations on the ammonia dimer,lg which suggest that the configuration with the linear N-H ... N bond is the most energetically favourable configuration, with a dimerisation energy of ca. 15 kJ mol-l.We propose a new interpretation of the vibrational spectroscopic data, not in terms of the overall liquid structure, but in terms of the immediate environment of the ammonia molecule. There is, of course, a link between these two descriptions, but we do not think it can be derived from the vibrational data.660 Spec t rochem is t ry of Liquid Ammonia Table 1. Observed and calculated peak maxima (in cm-l) in the spectrum of liquid ammonia at room temperature obs calc assignment notes 1640 1750 3 203 3 247 3 339 3 369 3 386 4 360 4458 4914 6071 6 309 6 384 6414 6 541 6 637 6704 677 1 7 622 7 704 7 757 7 843 8 122 8 230 8316 8 379 9 194 9404 9581 9 864 9 978 10111 10 702 10943 11 171 11 450 12515 12 889 13635 14037 15 299 15 822 16431 17 934 'Id 3337 v3, 3373 VSc - - 6535 2v3a 6663 2v3, 6700 2~~~ 6764 2v3, 7624 v,+2v3, exact v, + 2v3, v, + 2v3c - '2 + 2v3d - 8135 v,+2v3, exact vp + 2v3, v4 + 2v3c v, + 2v,, + v, - '4 + 2v3d - - - - 9594 3v3ar 9979 3V3c 9823 3v3, 10113 3v3d 10 699 v, + 3v3, exact v, + 3v3, 11 156 v,+3v3, exact v, + 3v3, 12513 4v3, 12888 4v3, 13 635 v, + 4v3, 14044 v,+4v3, 15294 5v3, 15847 5v3, 16423 v, + 5v3, 17935 6,, 1046 (Raman)3 broad 32 17 (Raman)3 3260 (Raman)3 3303 (Raman)3 c - - - broad - - (1046 + 3386 + 1640 = 6072) - - (1046 + 6704 = 7750) (1046+6771 = 7871) - (1640+6704 = 8344) (1640+6771 = 8411) (1046+ 1640+6541 = 9227)J .C. Dougal, P . Gans and J . B. Gill 661 1 I 1500 1600 w avelengh/nm Fig. 2.Absorption spectrum of liquid ammonia in the first overtone region (a) and fourth derivative of the absorbance with respect to wavelength (b). w avenum ber/ cm - I I I I 1 1 1 1 wavelength/nm Fig. 3. Absorption spectrum of liquid ammonia in the region of the combination bands with the first overtones (a) and fourth derivatives of the absorbance with respect to wavelength (b).662 Spectrochemistry of Liquid Ammonia 0.3 w avelength/nm Fig. 4. Absorption spectrum of liquid ammonia (10 mm cuvette) in the third, fourth and fifth overtone regions. The same spectrum is shown on two different absorbance scales, (a) and (b). Table 2. Calculated harmonic wavenumber and anharmonicity constants for overtone band progressions ; fundamental wavenumbers and anharmonicity constants for combination band progressions" band o,/cm-' X,,/cm-' v,/cm-' XZ3/cm-l v,/cm-' X,,/cm-' a 34O7( 4) -70(1) 1056(17) 16 (4) 1658 (21) -32(6) b - 3441 (15) -55(4) 1043 12 1607 -7 C 3397 ( 7) -23 (3) - d 3404(ll) -11 (4) - - - - - - - " Least-squares values of the error in the last digits are given in parentheses.The isolated ammonia molecular has Csv symmetry, giving rise to stretching vibrations of A; and E' symmetry, both of which are infrared and Raman active. There is also the possibility of Fermi resonance between 2v, and vl. The ammonia molecule in the liquid may have various environments. For example it may donate to 0, 1, 2 or 3 N-H ..- N bonds, and it may form H N-H bonds of the conventional kind or of the kind that exists in the solid. We suggest that the effect of these relatively weak interactions can, as a first approximation, be considered as a perturbation.Corset and Lascombe have calculated the effects of a perturbation consisting of a change in the N-H stretching force-constant. 21 The various environments cause different degrees of shift in the A; and E' wavenumbers, and in the cases of 1 and 2 N-H --. N bond structures, separation of the components of the E' vibration with the lifting of degeneracy. Thus, although the species with 0, 1, 2 or 3 N-H -.. N bonds should give rise to eight polarised and four depolarised Raman bands (all infrared active), some bands will be at almost the same wavenumber. Changes in the relative absorption intensities may also occur. In the region of the fundamental stretching vibrations the infrared spectrum showsJ.C. Dougal, P. Gans and J. B. Gill 663 remarkably few coincidences with the Raman spectrum, in which it is now generally agreed that there are three polarized bands.3* The infrared band at 3247 cm-l does not coincide with any of the polarised Raman lines, but it has been classed as a symmetric vibration v,, because of its low wavenumber. It therefore appears that there are four distinct symmetric vibrations ; ignoring Fermi resonance, that could mean that there are four ammonia environments. We believe that Fenni resonance should be ignored, following the interpretation by Gardiner et aL3 of data obtained from solutions of ammonia in carbon tetrachloride and acetonitrile. In dilute solutions they observed a strong band at ca.3314 cm-' and a very weak band at ca. 3206 cm-l (both polarised). As the ammonia concentration is increased the strong band shifts to lower wavenumber, the weak band shifts to higher frequency and grows in intensity relative to the strong band, new bands grow in relative intensity at ca. 3387 and 3217 cm-l (pure-liquid values), and v, shifts from 989 to 1046 cm-l. These changes are consistent with changes in the relative proportions of various ammonia environments, i.e. increasing aggregation, with concentration. Since both CH3CN and CCl, are weak hydrogen-bond acceptors it is reasonable to suppose that the strong band observed in the dilute solutions is predominantly that of monomeric ammonia molecules hydrogen-bonded to the solvent. In that case the weak band belongs to an ammonia aggregate such as a dimer, and is not the result of Fermi resonance.Lemley et aL4 have suggested a two species model for the solvent structure. Their most convincing evidence is the observation of two polarised N-D stretching vibrations in the Raman spectrum of NH,D. However, two bands would be predicted with the local environment model. One environment would have a N-D N bond and one would not; perturbations due to other hydrogen bonds in the molecule would be small. From the intensities of the two bands we would conclude that there are roughly equal proportions of the species with and without the deuterium bond. Our observation of four symmetric stretching fundamental bands together with four overtone progressions is highly suggestive of the existence of four local environments. These could well be the environments (IWIV): I I1 I11 IV N N N N I I I I I In the local-environment model hydrogen bonding involving an adjacent ammonia molecule and the nitrogen atom would be expected to perturb the spectrum less than hydrogen bonding involving the nitrogen atom of another ammonia molecule.The packing constraints that result in unusual hydrogen bonding in solid ammonia are not present in the liquid. It is therefore likely that the normal hydrogen bond with the lone pair on the bond axis will predominate. Only one environment can exist in an infinite structure; the species which has one N-H . - - N bond and one H ... N-H bond can exist in an infinite chain. The other environments can only exist in small aggregates.Various workers have studied the effect of temperature on the Raman-active N-H stretching vibration^.^. 5, 6 q lo All have observed large changes in the relative intensity of the component bands. Perhaps the most significant observation is growth of the depolarised band at ca. 3385 cm-l at low temperatures. This band, which we assign as664 Spectrochemistry of Liquid Ammonia v3d, is not found in the spectrum of a dilute solution of ammonia in carbon tetrachloride or a~etonitrile.~ It is thus a band characteristic of aggregates, and would be expected on that basis to be more prominent at low temperatures, where the liquid is expected to be more structured. Kruh and Petz have reported radial distribution functions calculated from X-ray diffraction data obtained at 277, 228 and 199 K.22 They concluded that the structure is more highly ordered at low temperatures.Buback has reported infrared spectra of the fundamental regions as a function of density, between 303 and 573 K; above the critical temperature, 405.5 K, there is of course only one phase. At 423 K the wavenumber of the band maximum in the v3 region increases from ca. 3385 to ca. 3405 cm-l at the lowest den~ity.~ This suggests to us a change in the direction of less aggregation, with the smaller aggregates having a higher absorption wavenumber. We have reported that the v4 bending region in the Raman spectrum is more complex than a single band.'' In addition to the peak at 1640 cm-' another band was found at 1750 cm-l, with five times the integrated intensity at 293 K, and approximately equal intensity at 196 K.We have now also found a band at 1750 cm-l in the infrared spectrum (table 1). This is yet another clear indication of a temperature-dependent aggregation process existing in liquid ammonia. The precise form of the aggregation is as yet unknown, but the new data presented here on the fundamental and overtone absorptions, taken together with the published data on Raman scattering, are consistent with a model in which any one ammonia molecule may form between 0 and 3 N-H .-- N bonds with neighbouring ammonia molecules. We are grateful for the support of the S.E.R.C. both for the grant for infrared and computer equipment (GR/B/00817), and for the provision of a CASE studentship for J.C. D. (cooperating body Johnson Matthey Technology Centre). We also thank Prof. B. L. Shaw F.R.S. for the use of the interferometer, and Dr T. R. Griffiths for the use of the Cary 14H spectrophotometer. References 1 Part 19. P. Gans, J. B. Gill and L. H. Johnson, J. Chem. SOC., Dalton Trans., 1987, 673. 2 G. Costeanu and P. Barchewitz, C. R. Acud. Sci., 1936, 203, 1499. 3 D. J. Gardiner, R. E. Hester and W. E. L. Grossman, J . Raman Spectrosc., 1973, 1, 87. 4 A. T. Lemley, J. H. Roberts, K. R. Plowman and J. J. Lagowski, J. Phys. Chem., 1973, 18, 2185. 5 M. Buback, Ber. Bunsenges. Phys. Chem., 1974, 78, 1230. 6 J. W. Lundeen and W. H. Koehler, J. Chem. Phys., 1975, 79, 2957. 7 C. A. Plint, R. M. Small and H. L. Welsh, Can. J. Phys., 1954, 32, 653. 8 T. Birchall and I. Drummond, J. Chem. SOC. A , 1970, 1859. 9 B. Bettignies and F. Wallart, C. R. Acad. Sci., 1970, 271, 640. 10 M. Schwartz and C. H. Wang, J . Chem. Phys., 1973, 59, 5258. 11 P. Gans and J. B. Gill, J. Chem. Soc., Dalton Trans., 1976, 779. 12 J. H. Roberts and J. J. Lagowski, Electrons, Fluids, Nat. Met. Ammonia Solutions, 3rd Colloq. Weyl, ed. 13 P. Gans, J. B. Gill, Y. M. MacInnes and C. Reyner, Spectrochim. Acta, Part A , 1986, 42, 1349. 14 P. Gans, Comput. Chem., 1977, 1, 291. 15 P. Gans and J. B. Gill, Appl. Spectrosc., 1983, 37, 515. 16 I. Olovsson and D. H. Templeton, Acta Crystallogr., 1959, 12, 832. 17 J. W. Reed and P. M. Harris, J . Chem. Phys., 1961, 35, 1730. 18 A. H. Narten, J . Chem. Phys., 1977, 66, 3117. 19 A. Hinchcliffe, D. G. Bounds, M. L. Klein, I. R. McDonald and R. Rhigini, J. Chem. Phys., 1981,74, 20 M. L. Klein and I. R. McDonald, J . Chem. Phys., 1981, 74,4214. 21 J. Corset and J. Lascombe, J. Chim. Phys., 1967, 64, 665. 22 R. F. Kruh and J. L. Petz, J . Chem. Phys., 1964, 41, 890. J. Jortner (Springer Verlag, New York, 1973), p. 39. 121 1. Paper 71901 ; Received 19th May, 1987

ISSN:0300-9599    

DOI:10.1039/F19888400657

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1988, 84(2), 665-674 Dynamic Studies of the Interaction between Diols and Water by Ultrasonic Methods Part 4. - 3-Methylbutane- 1,3-diol and 2,2-Dimethylpropane- 1,3-diol Solutions Sadakatsu Nishikawa," Naohiko Nakayama and Nobuyoshi Nakao Department of Chemistry, Faculty of Science and Engineering, Saga University, Saga 840, Japan In order to justify the relationship between the structures of diols in aqueous media and their ultrasonic properties, measurements of ultrasonic ab- sorption, sound velocity, density and viscosity have been made in aqueous solutions of 3-methylbutane- 1,3-diol and 2,2-dimethylpropane- 1,3-diol, which are isomers of each other. In both solutions, a single relaxational ultrasonic absorption has been observed in the frequency range 15-220 MHz.The absorption mechanism has been interpreted in terms of reaction kinetics associated with the interaction between the solute and solvent. As a result, the effect of the solute on the solvent (water) structure has been estimated, and it has been found that these diols act as water-structure promoters. Furthermore, the greater the hydrophobicity of the solute molecule, the more effectively it promotes the water structure. The trend in hydrophobicity determined from sound absorption has been confirmed from the concentration dependences of the compressibility. The correlation between the solvent structural parameters and the compressibility has been examined. The apparent molar volume has also been determined and is discussed with regard to the ultrasonic parameters.In previous studies1 of aqueous solutions of various diols by ultrasonic methods, the conformations or the structures of the diol molecules in an aqueous medium have been predicted to affect the ultrasonic characteristics ; i.e. relaxational absorption associated with an interaction between solute and solvent has sometimes been observed, depending upon the structure of the solute. In addition, relaxational absorption due to the conformational changes in some diols would be expected both in aqueous and organic solutions. In order to justify such predictions, ultrasonic absorption and velocity results are required in other solutions. For this purpose, we have chosen two diols, 3- methylbutane- 1,3-diol and 2,2-dimethylpropane- 1,3-diol, which are isomers.In this paper we report ultrasonic absorption, ultrasonic velocity, density and viscosity results in these solutions, and these are compared with results reported previously. Experimental The chemicals used in this study were purchased from the Tokyo Kasei Co. Ltd and were the purest grades obtainable. 3-Methylbutane- 1,3-diol was distilled under reduced pressure. 2,2-Dimethylpropane- 1,3-diol was a solid at room temperature and was used without further purification. The sample solutions were prepared with doubly distilled water by weighing at their desired concentrations. The concentration of the 2,2- dimethylpropane-1,3-diol solution was limited to 5 mol dm-3 because of its solubility in water. The ultrasonic absorption coefficient was measured by an improved pulse method in the frequency range 15-220 MHz : details have been given elsewhere.2 An interferometer 665666 Interaction between Diols and Water 2 3 4 5 CJmol dm-3 Fig.1. Concentration dependence of the ultrasonic absorption in an aqueous solution of 2,2- dimethylpropane-1,3-diol at various frequencies at 25 "C: 0, 14.54; 0, 45.52; a; 92.63 and 8, 222.1 MHz. at 2.5 MHz and a 'sing-around' meter at 1.92 MHz were used to determine the sound velocity. By both methods the value of the sound velocity was obtained to within an accuracy of better than +1 m s-l. The solution density was measured using a pyknometer of volume ca. 4cm3. The viscosity coefficient was measured using an Ubbelode viscometer. All items of apparatus were immersed in a water bath whose temperature was controlled to within k0.002 "C.Results and Discussion Fig. 1 shows the ultrasonic absorption coefficient divided by the square of the measurement frequency, ( a / f 2 ) , at various frequencies for an aqueous solution of 2,2- dimethylpropane- 1,3-diol at 25 "C. For concentrations < 2.0 mol dm-3 the (a/f2) values are independent of the frequency, and it was not possible to measure the absorption at concentrations > 5 mol dm-3 because of the diol's solubility in water. However, the values of (a/f2) in the range 2.7-5.0mol dm-3 have been found to depend considerably on both the frequency and the concentration, as is seen in fig. 1. In a solution of 3-methylbutane- 1,3-diol the absorption measurements were carried out in the concentration range up to 6.99 mol dm-3 at 20 "C and the absorption is dependent on frequency in the range above 2.02 mol dm-3.When the absorption mechanism is5000 I000 " ' 500 E v) P I 2 1 h N i a S. Nishikawa, N. Nakayama and N. Nakao I I I - - I I I 667 . - 10 5c 100 500 f/M& Fig. 2. Ultrasonic absorption plots in aqueous solution of 2,2-dimethylpropane- 1,3-diol: @,2.75; a, 3.00; 0, 3.60 and a, 5.00 mol d ~ n - ~ . f/MHz Fig. 3. Ultrasonic absorption plots in aqueous solution of 3-methylbutane-l,3-diol at 20 "C: 8, 3.02; 0, 4.03; a, 5.02; 0, 5.97 and (>, 6.99 mol dm-3.668 Interaction between Diols and Water Table 1. Ultrasonic and thermodynamic parameters in aqueous solutions of 3-methylbutane- 1,3- diol and 2,2-dimethylpropane- 1,3-diol C,/mol dm-3 f,/MHz A/lO-" s2cm-' B/10-" s2 cm-' v/m s-' p / g q/cP 2.02 3.02 3.50 4.03 4.50 5.02 5.58 5.76 5.97 6.58 6.99 2.75 2.90 3 .OO 3.10 3.20 3.40 3.50 3.60 3.75 3.80 4.00 4.25 4.50 5.00 3-methylbutane- 1,3-diol solution at 25 "C 210 22.6 25.9 1635 127 43.5 73.7 1678 116 105 91.1 1689 114 163 118 1693 114 210 157 1690 114 253 198 1682 127 288 234 1672 135 357 217 1664 131 329 285 1660 165 355 347 1640 139 436 434 1624 2,2-dimethylpropane- 1,3-diol solution at 25 "C 97.6 82.6 55.2 1642 96.4 85.4 73.1 1645 94.8 97.6 83.6 1646 98.8 128 72.5 1648 92.9 111 92.4 1649 88.3 165 91.5 1651 87.9 160 106 1651 85.0 195 94.8 1650 89.4 201 111 1650 91.2 210 112 1650 95.9 213 126 1648 99.6 242 130 1645 101 226 132 1642 103 267 267 1630 1.0042 1.0078 1.0095 1.0101 1.0104 1.0096 1.0089 1.0076 1.0072 1.0939 1.0018 1.0065 1.0068 1.0075 1.0072 1.0075 1.0077 1.0078 1.0080 1.0080 1.0083 1.0084 1.0082 1.008 1 1.0071 - - - - - - - - - - - 2.7 18 2.909 3.074 3.21 1 3.364 3.635 3.848 3.872 4.222 4.400 4.882 5.398 6.103 8.27 1 interpreted, it may be appropriate to analyse the frequency dependence of the absorption coefficient.If the absorption is associated with a relaxation process, it may be expressed by the following equation : wheref, is the relaxation frequency and A and B are constants. Fig. 2 and 3 show the frequency dependence of the absorption for both solutions. All the spectra are well represented by eqn (1). The ultrasonic parameters have been determined by a least- squares computer program, and the solid curves in the figures indicate the calculated values.Ultrasonic results determined thus are listed in table 1, along with other ultrasonic and thermodynamic parameters. In order to illustrate the concentration dependence of the excess absorption amplitude, A, and the background absorption, B, the results are shown in fig. 4 for an aqueous solution of 2,2-dimethylpropane- 1,3-diol. They increase monotonically with the analytical concentration. These trends are contrasted by those observed in many non-electrolyte aqueous solutions ;3 i.e. the peak sound absorption vs. concentration phenomena are usually observed in solutions of non- electrolytes that have a hydroxy group. In order to see if the observed relaxational absorption is due to a viscosity effect, the viscosity coefficient, 7, has been measured in the solutions and the classical absorption (denoted by the subscript cl) has been calculated using the equation (a/. 2)c1 = 8Z27/3PV3 (2)" I E N S.Nishikawa, N . Nakayama and N . Nakao 300b 200 too 669 2 3 4 5 CJmol dm-3 Fig. 4. Concentration dependences of the excess absorption amplitude, A , (0) the background absorption, B, (0) and the classical absorption, ( C C / ~ ~ ) ~ , , (0) for an aqueous solution of 2,2-dimethylpropane- 1,3-diol at 25 "C. where p is the solution density and v is the sound velocity. The results are also shown in fig. 4, and it is seen that the observed absorption is not associated with the viscous one. The background absorption, B, is still higher than that of the classical absorption, which may indicate that another relaxation process may exist in the higher frequency range. On the other hand, characteristic trends in the relaxation frequencies were found in both solutions, as is seen in fig.5 : they show a minimum. In the solution of 3-methylbutane- 1,3-diol, the relaxation frequency exists at a high frequency range, so that the absorption measurements were carried out at 20 "C in order to obtain the ultrasonic parameters as accurately as possible. From the dependence of the ultrasonic parameter, fr, on the concentration, we consider that a plausible absorption mechanism may be one due to a reaction associated with interactions between the solute and solvent. Following previous interpretations, the model may be simply expressed by k f AB$A+B (3) kb where k, and k, are the forward and backward rate constants, respectively, and AB is the complex formed by the solute A and solvent B.To a good approximation we can say that the solvent water molecules consist of non-hydrogen- bonded and hydrogen-bonded molecules, and the former may participate in the reaction under consideration. Further, it is assumed that the reaction associated with the formation and breakage of water hydrogen bonds is so fast4 that it is expected not to affect the solute-solvent interactions. Under these assumptions the relationship between the relaxation frequency and the analytical concentrations for the process expressed by eqn (2) is derived as5670 Interaction between Diols and Water 2 4 6 C,/mol dm-3 Fig. 5. Concentration dependences of the relaxation frequency for aqueous solutions of 2,2-dimethylpropane- 1,3-diol (0) and 3-methylbutane- 1,3-diol (0).Table 2. Rate and thermodynamic constants for aqueous solutions of diols ~~ solute k,/ 1 O8 s-' k,/ 1 O8 dm3 mo1-1 s-l P ref. a a a (1 6) pentane- 1,5-diol - - - 3-methylbutane- 1,3-diol 1.1 kO.1 2.1 fO.l 0.155 f 0.004 this work 2,2-dimethylpropane- 1,3-diol 0.99 If: 0.07 2.2 & 0.2 0.104 f 0.002 this work a No chemical excess absorption. where Ce and Cw are the analytical concentrations of solute and solvent, respectively, p is the fraction of non-hydrogen-bonded water molecules and K,, is defined as K,, = k,/k,. The rate and thermodynamic parameters in eqn (4) were determined so as to obtain the best fit of the experimental data, f,, to the equation by means of a non-linear least-squares program.The determined values are listed in fable 2. The value of p is considered to be the fraction of less structured water molecules, and is smaller than that in pure water.6 This means that the solute acts as a promoter of water structure. In this case p is expected not to be so dependent on temperature. For example, the enthalpy change and entropy change between the two states of water have been estimated to be 8 kJ mol-1 and 15 J mol-1 K-l for an aqueous solution of ally1 Cello~olve.~ Although the measurement temperatures for the two solutions in this study are different, it may be appropriate to compare the results : the value of B for the 3-methylbutane- 1,3-diol solution is smaller than that for the 2,2-dimethylpropane- 1,3- diol solution.From the structure of these two molecules, it is expected that the two C(2) methyl groups in the latter solute may be bulky and hydrogen-bonded water may form around the hydrophobic part of the solute molecule. This is also expected from the differences in solubility of the two diols in water. Note that no excess absorption is observed in the solution of pentane- 1,5-diol, another isomer.1bS. Nishikawa, N . Nakayama and N. Nakao 67 1 I I I I I I 2 4 6 C,/mol dm-3 Fig. 6. Plots of the maximum excess absorption per unit wavelength and the calculated pv2T as a function of concentration: (0) and (1) for a solution of 3-methylbutane-1,3-diol; (0) and (2) for a solution of 2,2-dimethylpropane- 1,3 diol. Another parameter obtained from measurements of the sound absorption and velocity is the maximum excess absorption per wavelength, ,urnax, which is related to the standard volume change of the reaction, AV, and the standard enthalpy change, AH.It is derived as follows: ( 5 ) where R is the gas constant, Tis the temperature, a, is the thermal expansion coefficient, C , is the specific heat at the low-frequency limit, C," is that at the high-frequency limit and I' is a concentration term expressed by ,urn,, = zpv2T ( A V - a, AH/PC,")~ (C,"/CP)/2RT r = [ 1 / c A + ~ / c B + 1 / c A B - ~ / ( c A + c ~ + c A B ) ] - ' (6) where Ci indicates the equilibrium concentration of reactant i. In order to obtain the changes in volume and enthalpy of the reaction, values of the two specific heats and the thermal expansion coefficient are necessary.However, these are not readily available. In a previous study1' we determined these volume and enthalpy changes on the assumption that the heat capacity was close to that of a 2-butoxyethanol solution. However, in the case of the present solutions, the contribution of specific-heat terms may be very important. Despite this uncertainty, we considered that the most effective term in determining the concentration dependence of ,urnax is pv2T. Fig. 6 shows the experimental values of the maximum excess absorption per wavelength and the calculated value of p v 2 r In the low concentration range the dependences resemble each other, while at higher concentrations they do not. This type of trend has also been observed in other diol solutions, and has been attributed to a relaxation process associated with rotational isomerization in the concentrated solution.' There may also be relaxation due to conformational changes of the solutes in concentrated solutions.Desnoyers' group8 has studied thermodynamic properties of various non-electrolytes672 Interaction between Diols and Water I I I 4 2 4 6 CJmol dm-3 Fig. 7. Concentration dependences of the apparent molar volume for solutions of 3-methylbutane- 1,3-diol (0) and 2,2-dimethylpropane- 1,3-diol (0). in aqueous media. They have reported that when the hydrophobicity of the solute increases, characteristic concentration dependences of the apparent molar volume and compressibility are observed in the solution. In order to see how the present solutes behave, the concentration dependences of these parameters were determined using the following relations : where q5v is the apparent molar volume, po is the solvent density, Me is the molecular weight of the solute and m is the molality, which is obtained from the molarity of the solutions.The adiabatic compressibility, K,, and the apparent molar compressibility, r ~ 5 ~ , of the solution are obtained from the results for the sound velocity and the solution (8) density using the equations lc, = 1/@v2) and A = K , A + 10OO(lcs - Kw)/mPo (9) #v = M e / P - - Po)/(mPPrJ) (7) where IC, is the compressibility of the solvent water. Fig. 7 shows the concentration dependence of the apparent molar volume for the two solutions. The curve for the 2,2- dimethylpropane- 1.3diol solution is not so sharp as that of the 3-methylbutane- 1.3diol solution, although the hydrophobicity of the former is considered to be larger from the absorption results ; i.e.the value of /3 for the 2,2-dimethylpropane- 1,3-diol solution isOO S. Nishikawa, N. Nakayama and N . Nakao 673 1 I I - 0. I 0.2 0.3 mole fraction, x Fig. 8. Concentration dependences of the apparent molar compressibility for solutions of 3-methylbutane- 1,3-diol (0) and 2,2-dimethylpropane- 1,3-diol (@). 0.2 P 0.1 C I / I 1 0.05 0.1 0 0.15 X min Fig. 9. Plots of j3 and the mole fraction where the compressibility shows a minimum, x,,,: 0 3-methylbutane- 1,3-diol; a, 2,2-dimethylpropane- 1,3-diol; 0, 2-butoxyethanol ; 0, 2-isobu- toxyethanol; (>, 2-t-butoxyethanol ; a, 3-methoxy-3-methylbutan- 1-01.674 Interaction between Diols and Water smaller than that for the 3-methylbutane- 1,3-diol solution. The trends in the apparent molar volume and the hydrophobicity of the solute might not hold for solutions of the present isomers.Fig. 8 shows the dependence of the apparent molar compressibility for the two solutions. The curve for the 2,2-dimethylpropane- 1,3-diol solution is seen to increase more sharply than that of the 3-methylbutane- 1,3-diol solution. This result is consistent with the prediction from the work of Desnoyers and coworkers.' The calculated compressibility passes through a minimum. In order to illustrate the correlation between compressibility and absorption, plots of B and the mole fraction at which the compressibility shows a minimum are shown in fig. 9, and are seen to pass through the origin. In solutions of some ethers which are isomers each other the same dependence has been found,3b although the slope is different, depending on the type of isomeric group. If the minimum-compressibility concentration were very low, the solute would be expected to act as a very effective promoter of water structure. This work was partly supported by The Naito Foundation. References 1 (a) S. Nishikawa and M. Mashima, J. Chem. SOC., Faraday Trans I , 1982,78, 1294; (b) S. Nishikawa, J. Chem. SOC., Faraday Trans. I , 1983,79,2651; (c) S. Nishikawa and N. Nakao, J. Chem. SOC., Faraday Trans. I , 1985, 81, 1931. 2 S. Nishikwa and K. Kotegawa, J. Phys. Chem., 1985, 89, 2896. 3 (a) M. Blandamer, Introduction to Chemical Ultrasonics (Academic Press, London, 1983); (b) S. 4 L. Hall, Phys. Rev., 1948, 73, 775. 5 S. Nishikawa, M. Mashima and T. Yasunaga, Bull. Chem. SOC. Jpn, 1975, 48, 661. 6 C. M. Davis Jr and J. Jarzynski, A&. Mol. Relax. Process, 1968, 1, 155. 7 S. Nishikawa and T. Yamaguchi, Bull. Chem. SOC. Jpn, 1983,56, 1585. 8 (a) G. ROUX, G. Perron and J. E. Desnoyers, J. Solution Chem., 1978, 7, 639. (b) J. Lara and J. E. Nishikawa, R. Shinohara and G. Tanaka, Bull. Chem. SOC. Jpn, 1986, 59, 827. Desnoyers, J. Solution Chem., 1981, 10, 465. Paper 7/1042; Received 15th June, 1987

ISSN:0300-9599    

DOI:10.1039/F19888400665

出版商:RSC    

年代:1988

数据来源: RSC