摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 3357-3367 Experimental Measurements of Polymer Unidirectional Fluxes in Polymer + Solvent Systems with Non-zero Chemical-potential Gradients BY MARIE-PAULE I. VAN DAMME, WAYNE D. COMPER AND BARRY N. PRESTON* Biochemistry Department, Monash University, Clayton, Victoria 3 168, Australia Received 1st March, 1982 The diffusion properties of two polymers in saline, namely dextran and albumin, have been followed by measurements of the unidirectional fluxes of their tracer (labelled) counterparts. The relationships between the diffusion coefficients obtained by measurement of the forward and back unidirectional fluxes to the polymer mutual-diffusion and intradiffusion coefficients have been analysed for semi-dilute polymer concentrations. The dynamic behaviour of macromolecules in a concentrated phase is of interest for many diverse aspects in biology and chemistry.Our studies, in particular, have been aimed at obtaining a better understanding of the dynamic transport behaviour of macromolecules in the extracellular matrices of connective tissues. We have previously devoted attention to the diffusional properties of various p o l y m e r ~ . ~ - ~ The standard form of the mutual-diffusion coefficient, which corresponds to the net flux of material resulting from relaxation of a concentration gradient, and the intradiffusion coefficient, which refers to the flux of trace amounts of labelled molecules through a solution of uniform concentration, have been extensively s t ~ d i e d . ~ - ~ The two other forms of diffusion coefficient that can be measured in a polymer + solvent system are those associated with the unidirectional fluxes of trace amounts of labelled material in systems with non-zero chemical-potential gradients as brought about by macroscopic concentration gradients.These types of diffusion coefficients have hitherto received only minor attention. It is the purpose of this paper to present an experimental analysis of the unidirectional-diffusion coefficients of binary systems containing either dextran + water or albumin + water (buffered with phosphate, pH 7.2). This study extends our earlier analysis4 of the measurements of the concentration dependence of the mutual-diffusion and intradiffusion coefficients for these polymers. EXPERIMENTAL MATERIALS - _ FDR7783 dextran (M, = 158 000; M,/M,= 1.32) and fluoresceinylthiocarbamoyl-dextran (FITC-Dx-150, Lot F-108) (z, = 153700; M,/M, = 1.51) were kindly donated by Dr K.Granath, AB Pharmacia (Uppsala, Sweden). These materials have been subject to extensive investigations in terms of their various physicochemical properties.3* 4 3 The two dextrans have similar properties. Bovine serum albumin (BSA, no. A-4503 fraction V, lot 109C-0081) was obtained from Sigma Chemical Company (St Louis, U.S.A.). [14C]Sorbitol (333 Ci rnol-')t was obtained from the Radiochemical Centre, Amersham. 7 1 Ci = 3.7 x 1 O ' O Bq. 33573358 MEASUREMENT OF POLYMER U N I D I R E C T I O N A L FLUXES PREPARATION OF POLYMER SOLUTIONS Stock solutions of polymers of known moisture content were made up by weight in phosphate-buffered saline (PBS) which was 0.14 mol dm-3 NaCl, 2.68 x lop3 mol dmP3 KCl, 1.5 x mol dm-3 KH2P04 and 8.1 x lop3 mol dm-3 Na2HP04 (pH 7.2).PREPARATION OF RADIOACTIVE-LABELLED MATERIALS [1251]BSA was prepared using the lactoperoxidase technique of Mar~halonis.~ [3H]BSA was prepared using the technique of Tack et aL8 The radioactive-labelled BSA was separated from the excess free 12512 or 3H on Sephadex G-25 PD-10 Column (AB Pharmacia, Uppsala, Sweden). The labelled monomeric BSA used in the diffusion studies was freed from dimer and any residual free 1251 or 3H by gel chromatography on Ultrogel AcA 54 (LKB, Bromma, Sweden) or G-150 (AB Pharmacia) at 4OC, using PBS buffer as eluate. The specific activity of the [1251]BSA was ca.2 Ci kg-' and of the [3H]BSA 260 Ci kg-l. Immediately before each experiment the [1251]BSA was further purified from free iodide on a Sephadex G-25 PD-10 column (AB Pharmacia). There was no release of free 12512 during the course of the diffusion experiments. [3H]FDR7783 dextran was prepared using the technique of Preston et al.4 with slight modifications. Dextran (100 mg) was dissolved in 0.1 mol dmP3 NaOH (1 cm3) to which Phenol Red (2 drops) was added; NaB3H4 solution (Radiochemical Centre, Amersham) (0.1 cm3 of 100 mCi in 0.1 mol dmp3 NaOH) was added and the reaction was carried out at 4 "C overnight. At the end of the reaction, acetic acid was added in order to acidify the medium. After 1 h, the medium was neutralized with NaOH. The polymer was separated from tritiated water on a Sephadex G-25 PD-10 column (AB Pharmacia).The preparation was then dialysed extensively against water and then purified on a Sepharose CL6B (AB Pharmacia) column (1.5 x 65 cm) in PBS. It has been shown previously4 that this labelling procedure resulted in minimal change in the molecular size distribution of the dextran sample. The specific activity of the [3H]FDR7783 was in the range 20-30 Ci kg-l. METHODS DETERMINATION OF DIFFUSION COEFFICIENTS Unidirectional-diffusion coefficients of labelled polymers, which were present in trace quantities, were measured in Sundelof diffusion cells. This technique has been described in detail el~ewhere.~ The cells consist of two cylindrical chambers which, by a shearing mechanism, can be moved in two different positions, as shown schematically in fig.1; a filling or emptying position, where the two chambers are separated, and a measuring position, where both chambers are brought over each other, forming thus a horizontal boundary. We define the concentrations in the bottom and top chambers as c b and C,, respectively. The mean concentration across the initial boundary is C[ = ( c b + C,)/2] and the difference in concentration is given as AC( = c b - C,). The total amount (mass) of diffusing solute Q transported across the boundary during the time t is given by Q2 = A2Ci D t / n where C, corresponds to the initial concentration of the labelled material in one of the chambers, A is the cross-sectional area of the diffusion compartments and D is the diffusion coefficient.The forward unidirecti%nal-diffusion coefficient representing the flux from the bottom to the top chamber is given by D while the back unidirectional-diffusion coefficient is given by z. For a system where C, = 0 the forward unidirectional-diffusion coefficient B is identical to the mutual-diffusion coefficient D. In the case where c b = C,, the back and forward unidirectional diffusion coefficients should be identical and are known as the intradiffusion coefficient D+. However, it is very difficult in practice to carry out these experiments owing to the absence of any stabilising gradient. Values of D+ were thus obtained at c b - C, N 5 kg mp3. For non-zero values of C, we shall examine the relationship of the mutual-diffusion coefficient and unidirectional-diffusion coefficients through the use of an additivity equation where D = (DC,,-BCJ/AC.FIG.1.- C, and M-P. I. V A N DAMME, W. D . COMPER A N D B. f i I l i -Schematic diagram of the formation C, represent the concentrations of N. P R ng positi ESTON on m ea sur i n g posit ion emptying posit ion 3359 of the free-liquid boundary in the Sundelof diffusion material in the bottom and top chambers of the respectively. cell. cell, Mutual-diffusion coefficients of BSA were also measured by free diffusion in a Beckman model E analytical ultracentrifuge using Schlieren optics. The initial concentration gradient was ca. 5 kg mP3 for routine measurements. RADIOACTIVE COUNTING PROCEDURES [3H]dextran samples were prepared for liquid scintillation counting with Aquasol (NEF-934 New England Nuclear, Boston, U.S.A.). Studies have shown that the degree of quenching of 3H counts is dependent on the total quantity of dextran present in the vial, as shown in fig.2. It is evident that there is a marked increase in quenching with increasing concentration of dextran, and we have chosen for routine measurements to work in the dextran concentration range 30-50 mg per vial, since although the counting efficiency is decreased by 40%, it remains relatively constant over this range. Furthermore, under these conditions the quenching does not change over a period of 3 days (fig. 2). In addition, the counting efficiency for these [3H]dextran samples is sensitive to the water/Aquasol ratio, as shown in fig. 3. A 22.5% (v/v) aqueous sample in Aquasol was used, conditions in which a stiff gel was formed.Note that preparation of counting samples with < ca. 19% H,O resulted in incomplete gel formation and precipitation of the sample with time. [3H]BSA and [14C]sorbitol were counted under similar conditions to [3H]dextran. 1251 was determined in an Autogamma analyser PW5420 (Philips, Holland).3360 MEASUREMENT OF POLYMER UNIDIRECTIONAL FLUXES 0 10 20 30 40 50 60 FIG. 2.-Effect of the addition of unlabelled dextran on the counting of [3H]dextran. The [3H]dextran samples (volume 1.3 ~ m - ~ ) were prepared for liquid scintillation counting with 4.5 cmP3 Aquasol. The experimental points represent the mean value of five determinations (0) together with their standard deviation (+). Some vials were left at 4OC for 3 days and then recounted, as indicated by the symbol dextran mass/10-6 kg 20 21 22 23 24 25 26 percentage of water in sample FIG.3.-Counting efficiency for [3H]dextran as a function of the percentage of the volume of aqueous sample with respect to the total volume of the solution (which contains Aquasol and aqueous sample) that is counted. The total volume was kept constant at 5.8 ~ m - ~ . The experimental points represent the mean value of three determinations (0) together with their standard derivation (+). Symbols represented by (0) correspond to vials which had been maintained at 4OC and recounted 3 days later. The dashed area corresponds to incomplete gel formation and precipitation of sample.M-P. I. V A N DAMME, W. D.COMPER A N D B. N. PRESTON 336 I FLUORESCENCE MEASUREMENTS Fluorescence measurements were performed on an Hirachi model 101 spectrofluorimeter (Hitachi Ltd, Japan) at 492 nm (excitation) and 515 nm (emission). RESULTS The use of the Sundelof cells allows studies to be made of the transport kinetics of labelled material across the initial boundary. The diffusion studies involved measurements associated with the transfer of < 15% of the total labelled material across the boundary. As shown in fig. 4, the plot of Q2 as a function of time is linear, as predicted by eqn (1); this is in agreement with a diffusional mechanism. In all diffusion studies reported here the regression coefficient of results plotted according to eqn (1) I s > 0.975. 0 1 2 3 4 time/ 1 O5 s FIG. 4.-A plot of Q2/C,2 (where Q is the total amount of material crossing the boundary after time t and C,, corresponds to the initial concentration of the labelled material in one of the chambers) as a function of time for the measurement of the unidirectional-diffusion coefficient of [3H]dextran from the bottom chamber to the top chamber where initially C, = 69 kg m-3 and C,= 64 kg mP3.The concentration dependence of the forward unidirectional-diffusion coefficient (n) of [1251]BSA and E3H]FDR7783 were studied when Ct = 0. The data are presented in fig. 5 and 6, respectively. For [lZ5I]BSA, the values of the a (= D) obtained do not show any concentration dependence and are in good agreement with the coefficients obtained from the analytical ultracentrifuge using Schlieren optics (fig.5), although some differences are apparent at high albumin concentrations. These values are also in good agreement to those obtained by the open-ended capillary technique2 and laser light scattering . lo In contrast to the diffusion of albumin, the diffusion coefficient of dextran shows - significant concentration dependence. When the diffusion values are plotted against C [ =(C,+ Ct)/2] we find that they are in good agreement with the mutual-diffusion coefficients (evaluated from ultracentrifuge and proton correlation spectroscopy) published previ~usly.~ As with the albumin transport, even at high C, values the Q2 value for dextran was linear with time, in accord with eqn (1). Therefore, while transfer of material across the boundary will produce a time-dependent change in AC, we do not detect any significant time-dependent variation in the diffusion coefficient.In general, transport across a free liquid boundary where high polymer concentration3362 MEASUREMENT OF POLYMER UNIDIRECTIONAL FLUXES 0 11' ri [ iI t ii i I 1 I 0 100 200 300 mean albumin concentration/kg m-3 FIG. 5.-Plot of the unidirectional-diffusion coefficient of [1251]albumin obtained from the Sundelof diffusion cell as a function of the mean concentration of albumin across the boundary (C) when C, = 0 (0). The value of the mutual-diffusion coefficient of albumin obtained by the boundary relaxation method in the analytical ultracentrifuge is also shown (0). The variation of the intradiffusion coefficient of [1251]albumin is shown as (m).The error bar corresponds to 95% confidence limits. gradients exist yields a mutual-diffusion coefficient at C which is essentially independent of AC. The values of both the forward-flux and the back-flux unidirectional-diffusion coefficients, a and D, respectively, have been measured when Cb was maintained constant while varying Ct (table 1 for [1251]BSA and table 2 for [3H]FDR7783). When minimal values of c b - Ct are studied the values of 25 and converge to the same value: that of the intradiffusion coefficient D+ whose values are presented in fig. 5 and 6 . Values of Td at intermediate values of Ct appear to correspond to some average of the mutual-diffusion and intradiffusion coefficient measured at the corresponding value of Cb. Values of D for [1251]BSA measured at constant Cb were found to be slightly dependent on Ct, suggesting that is directly related to C, for this particular polymer.We also find thatB is ca. 15% lower than the value of D+ at the concentration c b . In the case of [3H]FDR7783 (table 2), however, t h e n values tend to increase with increasing C, ; again, the values of the are significantly lower than the intradiffusion coefficients at corresponding concentrations. The calculated values of D, using eqn (2), are presented in the last columns of table 1 (for [1251]BSA) and table 2 (for [3H]FDR7783). These calculated values are in good agreement with mutal-diffusion coefficients evaluated experimentally. Estimates of D from eqn (2) when ( c b - C,) is small are subject to large error, which may explain the anomalous values obtained with albumin.Note that when the magnitude of D is low, then the difference between values of D a n d z is apparent even at small values of ACM-P. I. VAN DAMME, W. D . COMPER A N D B. N . PRESTON 3363 i i I t I 0 20 40 60 80 100 mean dextran concentration/kg m-3 FIG. 6.-Plot of the unidirectional-diffusion coefficient of [3H]dextran obtained from the Sundelof diffusion cell as a function of the mean concentration of dextran across the boundary when C, = 0 (0). The corresponding intradiffusion coefficient of [3H]dextran is also shown (m). The error bar represents 95 % confidence limits. TABLE 1 .-UNIDIRECTIONAL-DIFFUSION COEFFICIENTS OF ['251]ALBUMIN WITH c, MAINTAINED AT A CONSTANT VALUE OF (a) 157 kg m-3 AND (b) 327 kg mP3 WHILE VARYING C, IN EACH CASE diffusion coefficienta/ lo-" m2 s-' C,/kg m-3 + D (a) Ob 25.3 51.0 103.3 151.7 50.6 146.3 235.0 321.9 (4 Ob 157.0 131.8 106.2 53.8 5.5 327.0 276.6 181.0 92.3 5.5 6.13 (5.27-7.00) 5.86 (5.24-6.47) 5.38 (4.68-6.08) 4.58 (4.18-4.98) 3.29 (2.91-3.67) 6.01 (5.36-6.66) 5.96 (5.26-6.65) 3.31 (2.81-3.82) 1.63 (1.41-1.85) 1.47 (1.07-1.87) 2.82 (2.40-3.24) 2.80 (2.42-3.18) 2.78 (2.52-3.04) 1.07 (0.95-1.19) 0.83 (0.7 1-0.95) 0.91 (0.72-1.10) 0.63 (0.53-0.73) 0.70 (0.55-0.85) - - 6.2 6.4 6.6 7.9 16.1 6.0 6.8 5.2 4.2 35.1 a The numbers in parentheses indicate 95% confidence limits.For back-flux diffusion coefficient C, was initially 5 kg m-3. (i.e. ca. 5 kg m-3). This points to the necessity to extrapolate to zero AC in order to obtain an accurate value of the intradiffusion coefficient under these conditions.However, we have routinely employed the convenient practice of evaluating D+ values at AC % 5 kg m-3 in spite of the inaccuracy introduced, especially when dealing with low-magnitude diffusion coefficients. A more extensive study was undertaken in the comparison between t h e z and D+ values over a wide range of Cb values (tables 3 and 4). The values of t h e z for [1251]BSA were found to be consistently ca. 20% lower than the values of D+. For the [3H]FDR7783, the magnitude of theD is significantly less than the D+, at corresponding concentrations. These low values of D were also obtained when FITC-dextran was3364 MEASUREMENT OF POLYMER UNIDIRECTIONAL FLUXES TABLE 2.-UNIDIRECTIONAL-DIFFUSlON COEFFICIENTS OF [3H]DEXTRAN WITH c b MAINTAINED AT A CONSTANT VALUE OF (a) 45 kg rnP3 AND (b) 92.5 kg m-3 WHILE VARYING C, IN EACH CASE diffusion coefficienta / 1 O-ll m2 s-' CJkg m-3 D (a> O6 4.5 13.5 26.9 40.5 22.4 45.0 68.5 87.8 (4 O6 45.0 40.5 31.5 18.1 4.5 92.5 70.0 47.5 23.9 4.6 3.25 (2.71-3.75) 3.20 (2.81-3.59) 2.66 (2.52-2.80) 1.93 (1.63-2.23) 1.07 (0.87-1.27) 3.8 7 (3.33-4.40) 3.6 1 (3.57-3.65) 1.57 (1.32-1.81) 0.74 (0.67-0.81) 0.66 (0.49-0.83) - 0.45 (0.40-0.50) 3.51 0.64 (0.55-0.73) 3.51 0.69 (0.62-0.65) 3.65 0.70 (0.66-0.74) 4.41 a The numbers in parentheses indicate 95% confidence limits.For back-flux diffusion Diffusion coefficient the dextran concentration in the upper solution was initially 5 kg m-3. coefficients too low to measure with any accuracy.TABLE 3 .-COMPARISON OF THE BACK-FLUX UNIDIRECTIONAL-DIFFUSION COEFFICIENJ (B) OF ['251]ALBUMIN WITH ITS INTRADIFFUSION COEFFICIENT (D'). FOR THE DETERMINATION OF D, c, WAS 5 kg mP3, WHEREAS FOR D+ MEASUREMENTS C, - C, WAS ca. 5 kg M - ~ . ~~ ~~~ - diffusion coefficientsa/ 1 0-l1 m2 s-' 45.8 4.68 (4.28-5.08) 5.38 (4.70-6.10) 103.3 3.48 (3.08-3.88) 3.72 (3.38-4.06) 164.7 2.61 (1.97-3.24) 3.22 (2.91-3.43) 2 12.4 2.03 ( 1 37-2.19) 2.57 (1.97-3.17) 263.7 1.44 (1.36-1.52) 2.20 (1.88-2.52) 33 1.4 1.07 (0.95-1.19) 1.47 (1.07-1.87) a The numbers in parentheses indicate 95% confidence limits. used as an alternative tracer (table 4). A similarity between the values o f z and D+ has been noted previously for diffusion measurements with hyaluronate for low values for [3H]FDR7783 may be due to a boundary disturbance occurring in this system (for details see Discussion).We have used trace quantities of labelled solvent markers of different sizes, namely [14C]sorbitol and [3H]albumin, to monitor any such disturbance. The results of the unidirectional- diffusion coefficient of these trace materials under different polymer concentration conditions are shown in table 5. For convenience, we describe the direction of unidirectional transport of the solvent marker the same way as its dextran counterpart. It is clear from the studies of transport of [14C]sorbitol in dextran that the D for [14C]sorbitol into a dextran solution of either 50 or 100 kg mP3 is not significantly different from the diffusion coefficient when Cb = C,.Therefore, in using this particular solvent marker no substantial anomalies could be detected in the system. of Cbe5 This marked retardation of theM-P. I. V A N DAMME, W. D . COMPER A N D B. N. PRESTON 3365 In the case of using [3H]albumin as a solvent marker, using essentially the same procedure as for studies with [14C]sorbitol, we find now that t h e 3 is significantly lower as compared with the diffusion coefficient of [3H]albumin when C, z Ct. These results would substantiate the depressed back-flux diffusion coefficients of the [3H]dextran in similar systems (see table 4). TABLE 4.-cOMPARISON OF THE BACK-FLUX UNIDIRECTIONAL-DIFFUSION COEFFICIENTS (B) OF [3H]FDR7783 DEXTRAN AND FITC-DEXTRAN _WITH THEIR CORRESPONDING INTRADIFFUSION COEFFICIENTS (D+). FOR THE DETERMINATION OF D, C, WAS 4.45 kg mP3, WQEREAS FOR INTRADIF- FUSION MEASUREMENTS C, - C, WAS ca.5 kg m-3. diffusion coefficientsa/ lo-" m2 s-' D D+ c ~~ fluorescein- fluorescein- C,/kg m-3 [3H]FDR7783 dextran I3 HI FD R77 8 3 dextran 45.0 0.45 0.59 1.05 1 .08b 68.9 0.21 0.20 1.10 0.90b 92.5 0.10 0.10 0.66 0.67 (0.40-0.50) (0.55-0.63) (0.85-1 -25) (1.05-1.1 1) (0.18-0.24) (0.17-0.23) (0.94-1.26) (0.87-0.93) (0.07-0.13) (0.09-0.12) (0.49-0.8 3) (0.57-0.76) a The numbers in parentheses indicate 95% confidence limits. From ref. (3). TABLE TH THE TRANSPORT OF SOLVENT MARKERS, ['4C]SORBITOL AND [3H]ALBUMIN, IN SYSTEMS WITH ZERO AND NON-ZERO CONCENTRATION GRADIENTS OF FDR7783 DEXTRAN~ dextran FDR7783 concentration/kg mP3 diffusion coefficient of solvent marker (solvent marker, [14C]sorbitol = S*) 50+S* 4 50, 50 +- 5+S* lOO+S* -+ loob 100 +- 5+s* (solvent marker, [3H]albumin = A*) lOO+A* -+ 95 100 t 5+A* 47.3 (42.8-52.0) 5 I .9 (45.6-58.2) 37.7 (33.1-41.9) 41.8 (36.5-47.2) 1.67 (1.33-2.02) 1.02 (0.81-1.21) a Arrow indicates direction in which the flux of solvent marker was measured.Unlabelled sorbitol at a concentration of 5 kg m-3 was added to the bottom compartment solution in order to stabilize the boundary. Further insights into the relationship of and Z7 have been made in studying an 'ideal' system of the diffusion of [14C]sorbitol in sorbitol. We have performed unidirectional diffusion analysis by a similar procedure to that described above for polymer diffusion, with the results shown in table 6.We find that (i) there is a moderate concentration dependence of D for sorbitol_up to a concentration of 200 kg mP3, (ii) there is no major difference in the value of D when C, = 5 kg mP3 as compared with3366 MEASUREMENT OF POLYMER UNIDIRECTIONAL FLUXES TABLE 6.-uNIDIRECTIONAL-DIFFUSION COEFFICIENTS OF [ ''C]SORBITOL IN SORBITOL diffusion coefficient/ 10-l' m2 s-l C,/kg m-3 CJkg rnP3 e D 5 0 52.7 (48.9-56.5) - 104.5 0 48.4 (46.0-50.8) 47.5 (43.7-51.3)" 104.5 99.1 41 .O (37.6-44.4) 43.1 (41.2-45.0) 217.8 0 44.5 (39.9-49.1) 3 8.0 (3 6.6- 39. 4)a 217.8 210.1 30.0 (28.7-3 1.3) 32.0 (29.2-34.8) a For back-flux diffusion coefficient C, was initially 5 kg m-3. its value when AC is low and (iii) D a n d D are approximately the same for low AC values. These results would emphasize the fact that the anomalously low values for dextran are the result of the polymeric nature of this material.DISCUSSION Analysis of the flux of labelled polymers in the Sundelof diffusion cell allows the examination of polymer transport in systems with large polymer concentration gradients across a free liquid boundary. This has only been previously possible with the use of a diaphragm cell1' in which a millipore filter or glass filter is used to stabilise the initial boundary. The use of such a diffusion apparatus to measure polymer diffusion may appear to be unsatisfactory. The measurements of Keller et aZ.12 of the mutual diffusion of albumin or haemoglobin using the diaphragm cell has subsequently come under criticism due to the very low values of the mutual-diffusion coefficient obtained by this technique as compared with values obtained by laser light scattering13 and the open-ended capillary technique.2 Re-analysis of the polymer flux results of Keller et d.12 for albumin and haemoglobin on the basis of diffusion in a column of cross-section A l4 [i.e.eqn (l)] can yield mutual-diffusion coefficients comparable to those obtained by others. This analysis involves arbitrarily choosing one set of their experimental flux data as an internal standard for the evaluation of all diffusion-cell constants as embodied in the term A of eqn (1). Therefore, the anomalous nature of the mutual-diffusion coefficients measured by Keller et a l l 2 may stem either from the mathematical treatment of the results or the effect of interposing a membrane between the two interdiffusing solutions.With regard to dextran diffusion, the agreement between the unidirectional-diffusion coefficient of the dextran from Sundelof diffusion cells with the mutual-diffusion coefficient from the centrifuge (where low values of AC are employed) may indicate that the mutual-diffusion coefficient is not significantly dependent on AC. Caution is introduced here, however, as we have previously shown4 that the mutual-diffusion coefficient of dextran measured in the ultracentrifuge can be dependent on AC when studied over a limited range of AC. With the presence of large concentration gradients of dextran we have observed anomalous behaviour nssociated with the unidirectional back fluxes of this polymer.This behaviour may be regarded as some form of boundary disturbance, being equivalent to an overall displacement in the position of the initial boundary. An approximate calculation of the distance the boundary would have to be dis- placed in order to account for the difference between the back flux and intradiffusion flux indicates that this distance is extremely small and would be difficult to detect byM-P. I. V A N DAMME, W. D. COMPER AND B. N. PRESTON 3367 conventional measurements. It is the relatively lower-magnitude diffusion coefficients which would be more greatly affected by this boundary movement, such as those obtained with E3H]dextran transport. We have previously shown that boundary disturbances may result from certain hydrodynamic instabilities that yield countercurrent volume flows of high ordered and intricate nature in multicomponent polymer Obviously while no macroscopic changes in the volume of the system did occur, microscopic volume fluxes were visualized.Although this type of behaviour has not been visualized in the binary polymer +water systems studied here, it is conceivable that a form of microscopic volume flow may occur. This would be consistent with the results obtained by using [3H]albumin as a trace solvent marker (see table 5). It is informative to consider the case of macroscopic dimensions where a membrane of finite thickness is interposed between the two solutions. Kedem and Katchalsky16 have derived the commonly used expression for the total volume flow Jv = - L, oAII (3) where L, is the membrane filtration coefficient or hydraulic permeability, o is the reflection coefficient and A l l the difference in osmotic pressure across the membrane.It is often claimed that o is a measure of the selectivity of the system to solute and solvent (see Ogston and Michell' for more accurate expressions). A relationship could then be considered, in a binary system with a free-liquid boundary, of the interplay of dynamic polymer-polymer interactions, yielding a pseudo-membrane which is selective to the movement of the polymer itself and an osmotic gradient generated by the polymer concentration gradient. Note, however, that eqn (3) has been derived for discontinuous membrane systems in which absolute differences in thermodynamic parameters across the membrane, such as pi and n, are taken as the driving forces.The situation in a free boundary system is more complex, particularly with respect to an understanding of the effective osmotic pressure at any particular point in the boundary and the microscopic volume flows that may occur. This project was supported by the Australian Research Grants Committee (grant no. D68/16898, D2 73/14137 and DS 79/15252). We thank Gregory Checkley and Geoffrey Wilson for their expert technical assistance. B. N. Preston, T. C. Laurent and W. D. Comper, in Glycosaminoglycan Assemblies in the Extracellular Marrix, ed. D. A. Rees and S. Arnott (to be published by Humana Press). R. G. Kitchen, B. N. Preston and J. D. Wells, J. Polymer Sci., Polym. Symp., 1976, 55, 39. T. C. Laurent, L-0. Sundelof, K-0. Wik and B. Warmegird, Eur. J. Biochem., 1976, 68, 95. B. N. Preston, W. D. Comper, A. E. Hughes, I. Snook and W. van Megen, J. Chem. SOC., Faraday Trans. I , 1982, 78, 1209. K - 0 . Wik and W. D. Comper, Biopolymers, 1982, 21, 583. A. G. Ogston and B. N. Preston, Biochem. J., 1979, 183, 1. J. J. Marchalonis, J. Exp. Med., 1972, 135, 956. B. F. Tack, J. Dean, D. Eilat, P. E. Lorenz and A. N. Schechter, J. Biol. Chem., 1980, 255, 8842. L-0. Sundelof, Anal. Biochem., in press; T. C. Laurent, B. N. Preston, L-0. Sundelofand M-P. I. Van Damme, Anal. Biochem., in press. lo B. D. Fair and A. M. Jamieson, J. Colloid Interface Sci., 1980, 73, 130. l1 A. R. Gordon, Ann. N.Y. Acad. Sci., 1945, 46, 285. l2 K. H. Keller, E. R. Canales and S. I. Yum, J . Phys. Chem., 1971, 75, 379. l 3 R. S. Hall and C. S. Johnson Jr, J. Chem. Phys., 1980, 72, 4251. l4 L-0. Sundelof, Ark. Kemi, 1966, 25, 1. l5 B. N. Preston, T. C. Laurent, W. D. Comper and G. C. Checkley, Nature (London), 1980, 287, 499. l6 0. Kedem and A. Katchalsky, Biochim. Biophys. Acta, 1958, 27, 229. A. G. Ogston and C. C. Michel, Prog. Biophys. Mol. Biol., 1978, 34, 197. (PAPER 2/367)

ISSN:0300-9599    

DOI:10.1039/F19827803357

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Furuday Trans. I , 1982, 78, 3369-3378 Diffusion of Tritiated Water (HTO) in Dextran + Water Mixtures BY WAYNE D. COMPER, MARIE-PAULE I. VAN DAMME AND BARRY N. PRESTON* Biochemistry Department, Monash University, Clayton, Victoria 3 168, Australia Received 1st March, 1982 The diffusion of HTO has been measured in dextran solutions using an open-ended capillary technique and a newly developed Sundelof diffusion cell. HTO diffusion has been examined as a function of dextran concentration and molecular weight. These results, together with our previous results on the intradiffusion and mutal-diffusion coefficients of dextrans, now provide a complete set of conventional translational diffusion coefficients for both components in this binary system. Various assumptions associated with the theoretical description of polymer translational motion can now be examined.In order to understand the molecular motion ofpolymers in binary polymer + solvent systems knowledge of the relationship between the diffusional behaviour of both components is important. It is apparent, however, that in studies on aqueous polymer systems this relationship has been overlooked. Previous on the diffusion of dextran and other water-soluble polymers have been interpreted on the basis of several gross assumptions associated with the nature of the solutions and the relative movement of both polymer and water. In this study, these assumptions are tested for dextran + water mixtures. We have experimentally determined all the conventional diffusion coefficients for each component.The intradiffusion and mutual-diffusion coefficients of dextran, which have been reported ear lie^,^ are now complemented in this study with intra- diffusion measurements of HTO in dextran solutions. The relationship between the various motions of each component can now be represented and analysed in a unified and comprehensive manner. HTO diffusion is probably a good representation of molecular water motion rather than that of a special rapid H+ transfer process since HTO and H,lsO move at the same rate.4 We have studied HTO diffusion in two dextrans of different molecular weights, namely dextran TI0 (M, = 10400) and dextran FDR7783 (M, = 158200). Some measurements of the diffusion of water in dextran solutions by a proton magnetic resonance-pulse method have been reported previously5 and will be compared with values obtained in this study.EXPERIMENTAL MATERIALS Tritiated water (lot no. 1275-133, 0.25 mCi g-l)? was from - - New England Nuclear (Boston, U.S.A.). The polymer dextran samples T10 (M, = 10400; Mw/Mn = 1.68) and FDR7783 (M, = 158200; Mw/Mn = 1.32) were either supplied or kindly donated by AB Pharmacia (Uppsala, Sweden). Dextrans were labelled with tritium as described by Preston et aL3 Dextran solutions were prepared by weight in distilled water either from the dextran as supplied t 1 Ci = 3.7 x 1O1O Bq. 109 3369 FAR 783370 DIFFUSION OF kITO IN DEXTRAN+WATER or from dried samples. Conversion of polymer concentration into mass!volume units was carried out on the basis of the dry weight of the solid and the partial specific volume of the polymer.METHODS The intradiffusion coefficient of HTO in dextran solutions was measured by two techniques, namely a modified3 open-ended capillary technique of Anderson and Saddington6 and in a newly developed diffusion cell.'? Equilibrium dialysis of HTO in dextran solutions was performed at 4OC using Visking tubing (size 8/32 in, Medicell International, London).? The tubing was pre-treated with 1 % acetic acid and washed in sodium carbonate prior to use. Each dialysis bag contained dextran (at - the required concentration) and fluorescein-labelled dextran (FITC-Dx- 150, lot F- 108, M , 7 153 700 from AB Pharmacia, Sweden) at a low concentration of 1.2 kg mV3 and of known specific activity. The fluorescein-labelled dextran was used to monitor any changes in volume of the sample during dialysis.The total dextran concentration was varied up to 133 kg m-3. The dialysis bags (there were five for each dextran concentration studied) were placed in a cylinder containing HTO (ca. 70000 cpm ~ m - ~ ) . $ The solutions were gently rocked for 4 days. The final concentration of dextran in the bag was estimated by measuring the concentration of fluorescein-labelled dextran with an Hitachi spectrofluorimeter (Hitachi Ltd, Japan) using exciting light of wavelength 493 nm and reading the emission at 515 nm. The ratio of HTO inside the bag to HTO in the dialysate (as measured by radioactivity) gives the partition coefficient Kav. Radioactive counting procedures have been described elsewhere.8 HTO samples containing dextran were prepared for liquid scintillation counting with Aquasol (NEF-934 New England Nuclear, Boston, U.S.A.) (22.5% aqueous sample in Aquasol to form a stiff gel).There was no change in the counting efficiency of HTO over a wide range of dextran concentration (varied from 1 to 30 mg per counting vial). The partial specific volume of dextran T10 was determined by measuring the pyknometric density a t 20 *C of dextran solutions at concentrations of 104.8 and 217.4 kg m-3. A value of 0.60 cm3 g-l was found at both concentrations and agrees with values obtained at lower dextran concentrations by Edmond et al.9 NOMENCLATURE We designate dextran as component 1 and water as component 2 and their corresponding trace-labelled counterparts by the subscript*.We shall be dealing with three diffusion coefficients, namely the mutal-diffusion coefficient of component 1, D,, and the intradiffusion coefficients 0;' and Dl. Diffusion coefficients at infinite dilution will be further designated by the superscript O. RESULTS INTRADIFFUSION OF HTO IN DEXTRAN Fig. 1 shows the dependence of the reduced diffusion coefficient of water Di/(Di)O on the concentration of dextran. Several features of the results may be noted: (i) the reduced diffusion coefficient of water decreases with increasing dextran concentration, (ii) the magnitude of the decrease of the diffusion rate appears to be independent of the molar mass of the dextrans used in this study and (iii) the values of Dt/(Dt)" obtained from the two different methods appear to be similar. Other measurements of water diffusion in dextrans of various molecular weights5 showed property (ii) for dextrans with molecular weights up to 150000.However, the magnitude of their reduced diffusion coefficients at high dextran concentrations is considerably lower than that obtained in this study. We can offer no explanation for this difference. Note t 1 in = 2.54 x 10-2 m. $ cpm = counts per minute.W. D. COMPER, M-P. I. VAN DAMME AND B. N. PRESTON 3371 0 0.3 0 50 100 150 200 250 mean dextran concentration/kg m-3 FIG. 1.-variation of the reduced diffusion coefficient D:/(D:)O for HTO diffusion in dextran T10 (0) and dextran FDR7783 (A) as obtained from Sundelof diffusion cells. Values of D:/(D:)O obtained by the open-ended capillary technique are shown by the corresponding open symbols.that our values have been corroborated through the use of two independent methods and are also in accord with magnitude of reduced diffusion coefficients of low- molecular-weight solutes in dextran solutions.1° DISCUSSION OBSTRUCTION AND EX C LUDED-VO LU ME MODELS An approach which has proved successful in explaining the hindered transport of compact macromolecules in chain polymer solutions is the stochastic model of Ogston et This model takes into account the molecular size of the interacting species on the basis of the excluded-volume concept. The reduced diffusion coefficient is given in terms of the equivalent Stokes radius r2 of the migrating species, the effective cylindrical radius of the fibrous network molecules r1 and their effective specific volume (1) 6 so that D;/<D:>~ = A exp [ - BC!] where B = (r1+r2) Vt/r1 (2) and A is a constant close to unity and C, is the concentration of component 1 in mass/volume units.An independent estimate of B may be obtained from equilibrium partition of the migrating species between the free solution and a compartment containing dextran such that the partition coefficient (KaV) is given by1' K,, = exp [ - B2C1]. (3) Using the experimental values of K,,, the calculated values of B in eqn (1) and (3) may be compared. Linear regression analysis of the variation of In [D;f/(D:)"] against Ci yields values of A = 1.15 (kO.09) and B = 1.333 (k0.17). This value of B from diffusion data compares favourably with the values of B obtained from eqn (3) and K,, data only at dextran concentrations c ca.70 kg m-3 (fig. 2). The high error bar on B values from equilibrium partition experiments, which is due to the relatively high 109-23372 DIFFUSION OF HTO I N DEXTRAN+WATER 1 . 8 1 1 1.6 1.4 B 1 2 0.6 0.4 LL 0.15 0.20 0.25 0.30 0.35 &-/(kg rn-3)* FIG. 2.-Variation of B calculated from eqn (3) using experimental determination of K,, from equilibrium partition experiments, with C,b (0): The solid line represents the value of B obtained from linear regression analysis of HTO diffusion data in dextran using eqn (1). The filled area represents the 95% confidence limits of this line. values of Kav owing to the small effect of the dextran on water distribution for the dextran concentrations studied, precludes any further assessment of the stochastic model, particularly in estimating a sensible value of rl from eqn (2).Another treatment, which has been derived for dilute solutions, is the simple obstruction model of Wang.12 In this case, if it is assumed that exchange between bound (within the hydration layer of the polymer) and unbound water is rapid then the reduced diffusion coefficient may be expressed as where a is a geometric factor, pz is the density of water, w, is the mass fraction of the polymer and His the mass of water bound per unit mass of polymer. A calculation of H by a graphical procedure gives a best-fit value of H = 0.4 (i.e. 0.4 g of water bound per g of dextran or ca. 4.4 moles of H,O per mole of disaccharide unit).This value of H is similar to that found for DNA by Wang.13 PHENOMENOLOGICAL FRICTIONAL COEFFICIENT INTERPRETATION HTO INTERACTION WITH DEXTRAN CHAIN SEGMENTS The mutual-diffusion coefficient of dextran D, may be given a314 where pl is the chemical potential of component 1, ci and Oi are the molar concentration and partial molar volume, respectively, of component i a n d f l is the frictional coefficient15 obtained from mutual diffusion such that (6) W l -- =fE(u1-uz) ax where ui represents the velocity of component i. Note that the thermodynamic termW. D. COMPER, M-P. I. VAN DAMME A N D B. N. PRESTON 3373 14 ‘i 12 5 10 0 E 13 x 2 6 y \ 4 2 0 I I dextran concentrationlkg m-3 100 200 0 100 200 dextran concentration/kg m-3 FIG. 3.-Calculated values of fly from eqn (5) as a function of C, for (a) dextran TI0 and (b) dextran FDR7783 [values of D, were obtained from ref.(3)]. c1apl/dc1 in eqn ( 5 ) may be evaluated through a standard virial expansion (see Appendix). An estimate of the type of interaction that occurs between HTO and dextran may be made through a comparison of the flz values obtained from dextran diffusion and HTO diffusion. Evaluation off’ in eqn ( 5 ) from mutal diffusion and thermodynamic data of dextran3 is shown in fig. 3 . For both dextrans, the calculated values of fE are seen to sharply increase with dextran concentration (fig. 3 (a) and (b)]. The higher the molecular weight of the dextran the higher the magnitude of ffi. An alternative estimate of flz may be made through values of HTO diffusion in dextran.The intradiffusion coefficient of water may be described in terms of frictional coefficients, when c2* -g cz, such that14 (7) D$ = RWfZ *I +JZ *A’ Derivation of eqn (7) requires the frictional coefficients to obey the reciprocal relation and the relation c2f21 = C l f 1 2 f - 2 =f1*2 = f 1 z * The substitution of eqn (8) and (9) into eqn (7) and rearrangement yields an expression for f12 such that where ffi is defined as the f12 coefficient obtained from intradiffusion measurements through this equation as distinct fromfg in eqn (5). The~retically,f{~ andfg should be identical. Derivation of eqn (10) requires that ( D l ) O = RT/f2 * 2 (1 1) which is assumed to be constant over the concentration range of dextran studied. This assumption is probably quite reasonable in view of experimental data which suggest3374 DIFFUSION OF HTO I N DEXTRAN+WATER 0 a a a .0 0 a a I 1 I I I 0 100 200 300 dextran concentration/kg m-3 FIG. 4.-Calculated values of F{2 from eqn (12) as a function of C, for dextran T10 (0) and dextran FDR7783 (0). that polymers and polysaccharides similar to dextran have little or no effect on the molecular properties or structure of water at the polymer concentrations used in this study.16 Evaluation from eqn (10) through experimental data of D i reveals that this quantity is considerably higher (not shown) in magnitude than values o f f g evaluated from eqn ( 5 ) and given in fig. 3. However, it is of interest to convertf:, in eqn (10) into a frictional coefficient, F[2, evaluated in terms of mass/volume concentration units such that (12) The values of F,I, are found to be independent of dextran concentration and molecular weight (fig.4). The magnitude of I;;i is also seen to be considerably less than the values of ffi calculated from mutual-diffusion data [fig. 3(a) and (b)]. These results give rise to the concept that a water molecule is undergoing a frictional interaction with only a segment of the dextran molecule. This would also suggest that in order forhy [obtained from eqn (5)] to be equal tof:, [defined in eqn (lo)], the adjustable parameter should be an effective molecular weight of component 1, (M,),,,, to be embodied in the c, term in eqn (10). In this case we define I - zM2 4 2 -f12--* Ml The calculated values of (Ml)eff are given in fig.5 . It is clear that these values are dependent on dextran concentration and overall molecular weight of the dextran. The latter probably reflects the dependence of (M,),,, on the relative mobility [as embodied in eqn (6)] of the water molecule and chain segment of the dextran [which is equivalent to (A41)eff]. The data would suggest that while some form of entanglement is envisaged at higher dextran concentrations, which may reduce the molecular-weight dependence in certain dynamic parameters, the segmental interaction of dextran with water depends on the size of the dextran forming the transient network structure. In p.m.r.-echo studies, Tsitsishvili et aL5 arrived at a value of (M,),,, = 1300 g mol-l. In contrast to our findings they claim that (Ml),,, was essentially concentration independent in the polymer entanglement region.They also used a different method of calculating (M,),,, which relied on an obscure relation between the degree of hydration and the probability of centres of collision between dextran and water.W. D. COMPER, M-P. I. VAN DAMME AND B. N. PRESTON 3375 t o4 L e, h .. 3 1 o2 t I I 100 200 dextran concentration/kg m-3 FIG. 5.-Calculated values of (M,),ff from eqn (13) as a function of C, for dextran TI0 (-) and dextran FDR7783 (----). Values ofhy and 0: for a particular concentration have been obtained from the graph of their concentration dependence. RELATIONSHIP OF DYNAMIC DEXTRAN-DEXTRAN INTERACTIONS AND D E X T R A N-W A T E R I N T E R A C T I 0 N S T 0 T R AN S LA T I 0 N D I F F U S I 0 N COEFFICIENTS In order to establish the relationship of the various binary frictional interactions that can occur in the dextran+water system we also require an estimate of dextran-dextran interactions.This may be performed through a study of the intra- diffusion of dextran. The intradiffusion coefficient of dextran may be described in terms of frictional coefficients, when c, * 6 c,, such that14 Using the values off,? and DT we may calculate fi *, from eqn (14) (see table 1). The calculated values off, *, are found to be consistently negative over almost the whole dextran concentration range for both dextrans studied. The magnitude of this parameter is relatively lower for TI0 than it is for FDR7783. It is unlikely that experimental errors in the estimate of D, and D,t could account for the consistent behaviour of the value of f, *,.It could be argued in this study, however, that difficulties with theoretical analysis of the experimental results is due to the polydis- persity associated with the dextran fractions used. A serious error may then be introduced through an incorrect assignment of the value of M , in calculatingfly from eqn (5) (see also the Appendix). For example, if a & 50% error, at worst, was made in the assignment of M , (due to effects of polydispersity etc.) then calculated values of fly would also be expected to be ca. 50% in error at high dextran concentrations. Even so, the negative value off, *, would still be observed for FDR7783. In fact, due to the low degree of polydispersity for FDR7783 (with Mw/Mn = 1.32), incorrect assignments of M , inf,, calculations would be far less serious than estimated in this hypothetical case.For T10 dextran, introduction of f 50% error in M , will result in a positive or negative value off, *,. At this stage we emphasise that the values off, *,3376 DIFFUSION OF HTO I N DEXTRAN+WATER TABLE 1 .-EVALUATION OF THE FRICTIONAL COEFFICIENT f, *1 frictional coefficients/ 1 0l6 dyn s mol-l cm-' f1*1 v12 +fi *1> f 1 2 (subtract column 3 C/kg m-3 [from eqn (14)la [from eqn (5)] from column 2) 20 28 36 54 62 68 98 104 136 150 160 198 7 25 43 64 70 2.6 3 .O 3.2 3.7 4.1 4.2 6.1 5.3 6.1 8.9 7.5 17.7 dextran T10 3.0 3.2 3.5 4.2 4.6 4.9 6.4 6.7 8.6 9.5 10.2 13.3 dextran FDR7783 12 14 15 26 23 40 22 61 36 93 - 0.4 - 0.2 - 0.3 - 0.5 -0.5 - 0.7 -0.3 - 1.4 -2.5 - 0.6 - 2.7 + 4.4 -2 -11 - 17 - 39 - 57 a Experimental determination of Dr.are either negative or low in magnitude as compared with fly. This does realise the unexpected finding that in concentration regimes where solute-solute interactions occur in some form (as manifested by the molecular-weight-independent dynamic behaviour of transient statistical network structures in concentrated polymer ~olutions)~~ the concentration-dependent positive frictional factor resides primarily in solvent-solute interactions, as embodied in f12, rather than in dynamic solute-solute interactions. It appears that the latter type of interactions may dominate the thermodynamic term cli3pl/ac, in eqn (5). VALIDITY OF EQUATIONS DESCRIBING POLYMER MUTUAL DIFFUSION In practice, untested assumptions are frequently made to yield a relationship between D, and Dt.A degree of inconsistency and ambiguity has been associated with these expressions.18 A number of these assumptions can now be properly assessed as we have evaluated all three conventional diffusion coefficients in the dextran + water system. Eqn ( 5 ) may be expressed in the following f0n1-P which requires the assumption that f 1 2 f 2 1 = f 2 *2f1*1*W. D. COMPER, M-P. I. VAN DAMME AND B. N. PRESTON 3377 Eqn (15) may be reduced to a more familiar form as our studies on the magnitude of DZ and D l have established that for dextran concentrations up to 200 kg m-3 the condition D;cl < DZc2 holds so that As all frictional coefficients other than possibly f l have positive values, therefore, in assuming f 2 *2 to be constant, low magnitude or negative values offl *1 as found in table 1 will invalidate the assumption embodied in eqn (16) and therefore invalidate eqn (1 5) and (17).A sufficient but not necessary condition for eqn (1 5) and (1 7) to hold is the property of the solution which Bearman14 has called regular, that - D: - - 3 D; ul so that eqn (15) may be reduced to the form12 This equation has been used previously in the analysis of polymer diffu~ion.~~ 2o The results of dextran and water intradiffusion reported here together with measurements of the partial specific volume of dextran do not bear this relation out. The factor D t / D ; decreases with dextral1 concentration whereas fi2/fi1 remains constant.At this stage, it appears that expressions for mutual-diffusion coefficients of the form expressed in eqn (1 5 ) and (1 9) are not valid for dextrans and most probably not for other water-soluble polymers. APPENDIX EVALUATION OF THE THERMODYNAMIC TERM Clapl/acl I N EQN ( 5 ) Algebraic expression for the chemical potential of component 1 as a function of composition has been given by OgstonZ1 in the form of p,-& = RT(lnrn,+a2rnl+a,rn~+. . . ) (A 1) where rn, is the molality of component 1 (moles per gram of solvent) and a2, a 3 . . . are the coefficients expressing thermodynamic non-ideality. Conversion of the concentration units of eqn (A 1) to a molar scale, where rn, = c1/c2M2, and then differentiation with respect to c, gives the form of cli3pl/acl required such that A problem arises, however, in the determination of the coefficients in the virial expansion of eqn (A 2).It is customary to relate these to a standard virial coefficient form through the osmotic pressure equation. Since the osmotic pressure ll of a polymer solution can be written as then with the use of the Gibbs-Duhem equation and eqn (A 1) we have3378 DIFFUSION OF HTO IN DEXTRAN-kWATER or we may rewrite eqn (A 4) in terms of standard virial coefficients (which are related to the virial expansion of concentration terms in mass/volume units) such that IIc,8, = RT[c, + A,(Ml)2 cf + A3(Ml), c:+ . . .I. (A 5 ) We will now show that the relationship between the Ai in eqn (A 5 ) to the coefficients in eqn (A 2) may take different forms, which ultimately leads to different expressions for c, i3p,/'?cl.If the values of A,, A , etc. have been obtained by fitting the experimenta! osmotic pressure data to eqn (A 5) then by correspondence to eqn (A 4) we also state the relationships 2a3 and A - a2 A - - 2c, M , ( M , ) , - 3(C2M2), (M,? which on substitution into eqn (A 2) gives, with no assumptions, This equation has been used in the main text and e l ~ e w h e r e . ~ ? ~ ~ If, as is commonly assumed for dilute solutions, c2tj2 * 1 then in eqn (A 4) and (A 5) values of A , and A , take the form defined in eqn (A 6). Substitution of these values into eqn (A 2) will give c, ["I = RT[ 1 + 2A,(M,), c, + 3A3(M1)3 cg + . . . ] (A 8) acl T , p when c,~, N 1 is consistently Employed.Eqn (A 8) has been employed in various treatments of polymerl9 diffusion but is strictly only valid for dilute polymer solutions. This project was supported by the Australian Research Grants Committee (grant no. D68/16898, D2 73/14127 and DS 79/15252). We thank Gregory Checkley and Robert Kitchen for their expert assistance. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 T. C. Laurent, L-0. Sundelof, K-0. Wik and B. Warmegird, Eur. J . Biochem., 1976, 68, 95. R. G. Kitchen, B. N. Preston and J. D. Wells, J . Polym. Sci., Pofym. Symp., 1976, 55, 39. B. N. Preston, W. D. Comper, A. E. Hughes, I. Snook and W. van Megen, J . Chem. SOC., Faraday Trans. 1, 1982, 78, 1209. D. Eisenberg and W. Kauzmann, The Structure and Properties of Water (Oxford University Press, London, 1969). V. G. Tsitsishvili, V. Ya Grinberg, E. I. Fedin and V. B. Tolstoguzov, Pofym. Sci. USSR, 1979, 20, 2888. J. S. Anderson and K. Saddington, J. Chem. SOC., 1949, 5381. L-0. Sundelof, Anal. Biochem., in press; T. C. Laurent, B. N. Preston, L-0. Sundelof and M-P. 1. Van Damme, Anal. Biochem., in press. M-P. I. Van Damme, W. D. Comper and B. N. Preston, J . Chem. SOC., Faraday Trans. I , 1982, 78, 3357. E. Edmond, S. Farquhar, J. R. Dunstone and A. G . Ogston, Biochem. J., 1968, 108, 755. B. N. Preston, T. C. Laurent and W. D. Comper, in Gfycosaminogfycan Assemblies in the Extracellular Matrix, ed. D. A. Rees and S. Arnott (to be published by Humana Press). A. G. Ogston, B. N. Preston and J. D. Wells, Proc. R. SOC. London, Ser. A , 1973, 333, 297. J. H. Wang, J. Am. Chem. Soc., 1954, 76, 4755. J. H. Wang, J. Am. Chem. SOC., 1955, 77, 258. R. J. Bearman, J. Phys. Chern., 1961, 65, 1961. K. S. Spiegler, Trans. Faraday SOC., 1958, 54, 1409. W. D. Comper and T. C. Laurent, Physiol. Rev. 1978, 58, 255. P. G. de Gennes, Nature (London), 1979, 282, 367. J. S. Vrentas and J. L. Duda, J . Appf. Polym. Sci., 1976, 20, 2569. T. Loflin and E. McLaughlin, J. Phys. Chem., 1969, 73, 186. K-0. Wik and W. D. Comper, Biopolymers, 1982, 21, 583. A. G. Ogston, Arch. Biochem. Biophys., 1962, Suppl. 1 , 39. (PAPER 2/366)

ISSN:0300-9599    

DOI:10.1039/F19827803369

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 3379-3382 Catalytic Deamination by Reversed-flow Gas Chromatography BY MICHAEL KOTINOPOULOS, GEORGE KARAISKAKIS AND NICHOLAS A. KATSANOS* Physical Chemistry Laboratory, University of Patras, Patras, Greece Received 22nd March, 1982 Rate constants and activation parameters for the deamination of 1 -aminopropane, 2-aminopropane and aminocyclohexane over 13X molecular sieve have been determined using the technique of reversed-flow gas chromatography. In aminocyclohexane the fraction of the surface which is catalytically active has been estimated, and this is found to increase with increasing temperature. Reversed-flow gas chromatography (r.f.g.c.) is a new differential method for studying the detailed kinetics of surface catalysed reactionsly and other slow physical processes, such as diffu~ion.~ The method has been used successfully to study the dehydration of alcohols over 13X molecular sieve and y-aluminium ~ x i d e .~ The activation parameters determined were found to agree with those determined by other techniques. Moreover, the fraction of catalytically active surface sites was determined and found equal to the fractional conversion of the reactant to products. In the present paper we used the r.f.g.c. technique to study another class of organic reactions, namely catalytic deaminations to alkenes. The catalyst was molecular sieve 1 3X and the reactants 1 -aminopropane, 2-aminopropane and aminocyclohexane, to form unsaturated hydrocarbons as main products. The r.f.g.c. method is very simple. It uses a conventional gas chromatograph and a column consisting of two lengths I' and 1 in series containing the catalyst.The reactant is introduced as a pulse between these two lengths, and after a certain time the direction of the carrier-gas flow is repeatedly reversed. This gives rise to extra peaks in the chromatographic trace, whose height or area under the curve depends on the exact time of each flow reversal. The analytical mathematical form of this dependence, which is characteristic of the mechanistic scheme of the reaction, permits the calculation of rate constants from the experimental data. EXPERIMENTAL MATERIALS Molecular sieve 13X, 80-100 mesh (Applied Science Laboratories) was used as the catalytic surface. 1 -aminopropane (puriss grade), 2-aminopropane (puriss grade), aminocyclohexane (purum grade) and cyclohexene (puriss grade) were obtained from Fluka AG.Propene was from Matheson (G. P. grade, 99.7% purity) and benzene was from Merck AG (Uvasol). With 1 -aminopropane as reactant nitrogen was used as carrier gas, while with 2-aminopropane and aminocyclohexane the carrier gas was helium. Both gases were products of Linde, Athens, (99.99% purity). 33793380 CATALYTIC DEAMINATIONS APPARATUS AND PROCEDURE The experimental set-up and the procedure followed in the r.f.g.c. method have already been reported.2 The lengths l' + I of the chromatographic column (glass, I.D. 4 mm) containing the catalyst were 3 + 108, 0.5 + 8 and 1.3 +46 cm for reactants 1-aminopropane, 2-aminopropane and aminocyclohexane, respectively. The conditioning of the columns was conducted in situ by holding them at 693 K for 19 h, under carrier-gas flow (0.76 cm3 s-l, corrected at column temperature). During the kinetic runs the flow rate was in the region 0.65-0.87 cm3 s-l.For each amine, 1 mm3 of liquid, using a microsyringe, was injected onto the column length 1 through an injector placed at the junction of the two columns, and with the carrier gas flowing in direction F (forward). Plots and calculations were made on a Hewlett-Packard 9825 A desk-top computer connected to a 9872 B plotter. RESULTS AND DISCUSSION The main products of the deamination reactions studied were identified experi- mentally by comparing their retention times with those of pure substances, under the same conditions.These main products were propene in the deaminations of 1-aminopropane and 2-aminopropane, and cyclohexene in the case of aminocyclohexane. The kinetics of the deamination reactions studied conform to a simple first-order decomposition of the adsorbed amine to give the adsorbed product(s). Therefore, the same equation as for the dehydration of alcohols4 describes the kinetic law f = W b P (ktR) - 11 exp (- k4,J (1) wherefis the area under peaks F and R (where F indicates forward and R reversed with respect to the direction of the gas flow), rn the mass of amine injected and g its 5 - 4 Y .- c 0) .3 Y - 2 3 s E - 2 1 0 1 2 3 4 t t o t l 1 0 3 s FIG. 1.-Plots of eqn ( 1 ) for the deamination of aminocyclohexane to cyclohexene over 13X molecular sieve at 597 K: 0, R peaks; A, F peaks.M. KOTINOPOULOS, G.KARAISKAKIS AND N. C. KATSANOS 338 1 TABLE RATE CONSTANTS FROM R PEAKS (kR) AND FROM F PEAKS (kF) FOR THE The fvalues in the rate constants are standard errors. DEAMINATION OF THREE AMINES OVER 13x MOLECULAR SIEVE AT VARIOUS TEMPERATURES 1 -aminopropane 624 644 65 1 657 2-aminopropane 580 592 61 1 623 aminocyclohexane 574 586 597 604 626 3.9 f 0.6 5.8 f0.5 7 f l 7.8 f 0.3 1.54f0.02 3.0 fO.l 7.0f0.1 12.4 f 0.3 2.3 f 0.3 3.3 f0.3 7.3 f 0.3 10.9k0.8 35f1 4.3 +0.3 6.6f0.3 8.0 f0.5 7.9 f 0.2 1.64 f 0.04 3.0k0.2 6.73 f 0.06 12.1 f0.2 2.22 & 0.08 3.7 f0.3 7.4 f 0.3 10.3 k0.5 32+ 1 - 10 - 14 - 14 - 1.3 - 6.5 0.0 3.9 2.4 3.5 -1.4 5.5 8.6 - 12 TABLE 2.-ACTIVATION ENERGIES (E,) AND ENTROPIES (AS’) FOR THE DEAMINATION OF THREE AMINES OVER 13x MOLECULAR SIEVE EJkJ mol-l -ASz/J K-l mol-l amine R peaks F peaks R peaks F peaks 1 -aminopropane 71 f 2 66f9 212f3 218f 17 2-aminopropane 144f3 138f3 84+6 94&6 aminocyclohexane 162 f 1 1 156f5 51 f 17 59f8 fraction held on reactive surface sites, k the rate constant, ttot the total length of time until the last reversal of the gas flow, and t , the retention time of the product in the direction opposite to the flow.An example of plotting lnfagainst ttot, according to eqn(l), is shown in fig. 1 for both R and F peaks. The deviation from linearity in all plots was small, as judged from t-tests of significance on the relevant coefficients of regression. The probability of the various t values being exceeded was < 1 % in all cases.From the slopes of the above-mentioned plots the values of k were calculated and are collected in table 1. They are denoted as k , (from R peaks) or kF (from F peaks). The last column of this table gives the percentage difference between the two rate constants, showing a fairly good agreement between the k , and k , values in most cases, as predicted by the theory. Some relatively large differencies may be due to accidental errors. Eqn (1) has been derived2 on the basis that the mechanistic scheme of the reaction is fast k fast A+S $ A-S -+ D-S + D+S3382 CATALYTIC DEAMINATIONS where A is the gaseous reactant, D the gaseous product, and A-S and D-S the respective adsorbed species on active centres S. Deamination of aminocyclohexane over y-A120, at 525 K has been shown2 to proceed via the formation of an adsorbed intermediate B-S between A-S and D-S.The fact that the present results conform to a simple first-order law is most probably due to very fast formation or disappearance of the intermediate B-S, caused either by the different catalytic surface and/or by the higher temperatures used here. The energies and entropies of activation for the deamination reactions, calculated from conventional Arrhenius plots, are given in table 2. The differences between activation parameters determined from R peaks and from F peaks lie Yjithin the limits of experimental error, showing that secondary reactions of the detected product or irreversible adsorption of it are negligible. This is because the lengths Z and I' of the column responsible for the R and F peaks, respectively, are very different (by a factor of 16-36, cf.the Experimental section). No comparison of our results with literature values can be made, since to the best of our knowledge no deaminations over zeolites have been studied previously. As in the dehydration of alcohol^,^ one can find the absolute value for the pre-exponential factor in eqn (I), mg[exp (ktR)- 13, if the response of the detecting system is known. Then, knowing m and tR from experiment, and k from the gradient of eqn (l), the fraction g can be calculated. The mean values determined from F and R peaks for the deamination of aminocyclohexane are 0.24, 0.34, 0.46 and 0.76 at temperatures 574, 586, 604 and 626 K, respectively. In aminopropanes the absolute value of g could not be calculated, because the response of the flame ionization detector to the product propene could not be found accurately. It was possible, however, to ascertain that the g value in this case does not change significantly with temperature, as in the case of aminocyclohexane above.It was shown previously4 that the g values give the fractional conversion of the reactant to product. Thus, an increasing conversion of aminocyclohexane to cyclohexene with increasing working temperature is observed here. This is not due to irreversible adsorption of the product on the solid catalyst, since the g values calculated from the F and R peaks are not significantly different, in spite of the fact that these two types of peaks are due to two different column lengths (Z' and I ) containing very different amounts of catalyst. It is hoped that the present paper, together with previous studies1* 2 * will help to introduce workers in various fields of heterogeneous catalysis to the technique of reversed-flow gas chromatography, which is a new tool for studying the kinetics of surface reactions. We thank Mrs Margaret Barkoula for assistance. N. A. Katsanos and I. Georgiadou, J. Chem. SOC., Chem. Commun., 1980, 242 and 640. N. A. Katsanos, J. Chem. SOC., Faraday Trans. I , 1982, 78, 1051. N. A. Katsanos and G. Karaiskakis, J. Chromatogr., 1982, 237, I . G. Karaiskakis, N. A. Katsanos, 1. Georgiadou and A. Lycoughiotis, J. Chem. SOC., Faraday Trans. I , 1982, 78, 2017. (PAPER 2/496)

ISSN:0300-9599    

DOI:10.1039/F19827803379

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 3383-3392 Ion-Water Geometry and the Tammann-Tait-Gibson Effective Pressure and Radius BY JEAN V. LEYENDEKKERS Building A1 2, School of Biological Sciences, University of Sydney, Sydney, N.S.W. 2006, Australia Received 23rd March, 1982 An analysis based on X-ray and neutron diffraction data shows that the ion-water geometry can be summarised in terms of the Tammann-Tait-Gibson (T.T.G.) radius of the ion and two angles describing the relative orientations of the ion and water molecule. One angle only depends on Zi (Zis the ionic strength) whilst the other angle, a, depends also on the size of the ion. The relative values of a [(a+-aK+) for the cations and (a--aC1-) for the anions] are directly proportional to [d(l/Z&,tra/dm]m-ro where is the proton magnetic relaxation rate.The relationship of the T.T.G. effective pressure to the geometry is demonstrated as well. 1. INTRODUCTION The most detailed information about the microscopic static structure of liquids is obtained from neutron and X-ray diffraction measurements.1 Enderby and Neilsonl, have reviewed recent advances in this field, as applied to the structure of ionic liquids. These authors have also given a general survey of X-ray and neutron scattering by aqueous solutions of e1ectrolytes.l They summarised most of the published X-ray results, covering fifty electrolytes (ion-oxygen distances and hydration numbers for 19 cations and 6 anions). These studies covered a range of concentrations [0.056-27.5 mol (kg H20)-'] but most of the measurements were at concentrations above 2 mol (kg H,O)-'.Recently, Ohtorno's group has made neutron diffraction measurements on alkali3v4 and acid5 halide solutions at ca. 1 mol (kg H,O)-'. Theoretical studies of ion-water geometries have been reviewed recently.', Predom- inant in such studies are the simulation methods (mathematical and numerical). Numerical simulations use either the Monte Carlo (MC) or molecular dynamics (MD) techniquess? '. The structure of a single-ion-single-water molecule8 or a single ion or ion-pair in a cluster of water moleculesg has been analysed using quantum-chemical methods. At present it appears that the cation-oxygen distances1* and possibly the anion- deuterium distances are relatively insensitive to concentration. However, the orienta- tions of the water-molecules near the ions are a function of concentration. In the present paper the ion-water geometry is analysed, especially in relation to concen- tration effects, using the recently derived T.T.G.(Tammann-Tait-Gibson) radius (r+A) of an ion in water.I0 The relationship of the T.T.G. effective pressurelo to this geometry will also be considered. 33833384 ION-WATER GEOMETRY 2. RELATIONSHIP BETWEEN THE T.T.G. AND CRYSTAL RADII Values of (r+A) have been estimated from the density of the solutions, using the T.T.G. model.l0 The term A is a length (in A) resulting from packing and electrical deformation effects1° and r is the crystal radius (Pauling) for monatomic ions or the interatomic distance for other ions.lo The length A is related to r uia the following equations.For the cations, oxyanions and ions such as SCN- (r+A) = 1.304+0.643r (1 a) (r+A) = -0.85+ 1.43r (1 b) with an average deviation for all ions (excluding Zn2+ and Be2+) of _+ 0.07 A. For the halides with an average deviation of kO.02. The ions H+ and OH- do not fit either of these equations, probably because of unreliable ( r + A). 3. CATION-WATER GEOMETRY Recent accurate neutron diffraction re~ultsl-~ indicate that cation-oxygen distances r+o are insensitive to concentration changes. The following equation was derived using the neutron diffraction data of Enderby and Neilson1g2: or r+o = 1.67(r++A+)-0.95 r+o + 0.95 = 2 sin (8/2) (r+ +A+) where 8 = 112.9O. This value of 8 is consistent with the cation-hydrogen distances from neutron diffraction data as illustrated below.Values of r+o from eqn (2) are compared with the corresponding X-ray, neutron diffraction and theoretical values in table 1. The calculated value of r+o for Na+ appears low but is in reasonable agreement with the ab initio quantum calculations. (3) If eqn (1) and ( 2 ) are combined r+o = 1.07r+ + 1.23 which is to be compared with the well known relationship r+o 'v (r+ + 1.38) (I .38 being the radius of the water molecule). The length A can be split into the geometric component (0.55 A for the cations in table 1) and the electronic component Ae1,lo so that eqn (2) may be written r+o = 2 sin (8/2) (r+ + Ael+) ( 2 4 neglecting the residual of 0.03 on the left-hand side. The value of the cation-hydrogen (or deuterium) distance, T + ~ , is a function of concentration (and indicates the orientation of the water molecule).For a given concentration it is found that r+H is directly proportional to (r+ +A+), which may be expressed r+H = 2 sin ( 4 / 2 ) (Y+ +A+). The neutron diffraction data for NiCI,'? (giving accurate values of r+H as a function of concentration) were used to derive the following equation (4) 4 / 2 - 40/2 = - 6.83314 + I . I 71 ( 5 ) where do equals 8 from eqn (2b); eqn (5) apparently is independent of the ion as itJ. V. LEYENDEKKERS 3385 TABLE 1 .-COMPARISON OF CALCULATED AND EXPERIMENTAL CATION-OXYGEN DISTANCES (A) ( r + A) r+0 ion ra b C expt. eqn (2)b eqn (2)" theoreticald H+ Li+ Na+ K+ Rb+ cs+ NH,+ Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Zn2+ Cd2+ Las+ Ce3+ co2+ Ni2+ CU2+ ~ 1 3 + 1n3+ Cr3+ Er3+ 0.36e 0.68 0.97 1.33 1.47 1.67 1.03 0.35 0.66 0.99 1.12 1.34 0.74 0.97 1.14 1.07 0.74 0.72 0.72 0.5 0.8 1 0.69 0.96 1.27 1.74 1.85 2.09 2.23 2.38 2.18 (1 .06)h 1.76 1.99 2.04 2.23 1.69 2.09 2.04 1.81* (1.44)h - - - - - - 1.54 1.74 1.93 2.16 2.25 2.38 1.97 1.53 1.73 1.94 2.03 2.17 1.78 1.93 2.04 1.99 1.78 1.77 1.77 1.63 1.83 1.75 1.92 - 1.95f 1.909 2.08-2.25 2.4-2.69 2.37-2.43 2.6-2.89 2.8-2.9 - 2.8 5-3.05' 2.8-3.0 2.0-2.1 2.4f 2.26-2.4 2.61 2.9 2.26 - - - - 2.1 2.07f 2.05-2.1 1.93-2.3 1.9 2.35 2.3 1.90- 1.98 1.17 1.98 - - 2.14 - 2.53 2.77 3.02 2.69 (0.82)h 1.99 2.37 2.46 2.77 (1 .46)h 1.87 2.54 2.46 2.07 - - - - - - - - - 1.62 1.96 - - 2.27 - 2.66 2.8 1 3.02 2.34 1.61 1.94 2.29 2.44 2.67 2.02 2.29 2.46 2.41 2.02 2.00 2.00 1.77 2.11 1.97 2.26 - - - - - 1.81-1.89 Sch 2.0 Cls 2.06-2.10 MD 2.20-2.36 Sch 2.3 Cls 2.3 1 MD 2.69-2.90 Sch 2.8 Cls 3.10 MD - - 1.95 Sch 2.40 Sch * Pauling crystal radius (interatomic distance for NH,+), ref. (10); derived from density of solution and independently of r10 (* from neutron diffraction data); from eqn (1); from the compilation in ref.(2) : Sch, single-ion-single-water molecule; Cls, cluster theory; MD, molecular dynamics; ref. (10); f neutron diffraction (k0.05 A) from ref. (2); 9 neutron diffraction from ref (3)-(9, other values from X-ray diffraction data listed in ref. (1); bracketed values are less reliable. gives good predictions for the other cations (Li+, K+, Ca2+, Na+ and Cs'); see table 2. The ionic strength I is given by fz, z- V?, with z representing the ionic charge, v the number of moles of ions per mole of salt and rn the molality [mol(kg H,O)-l].The value of T + ~ decreases with concentration-up to I,,, (8.53rn) and then increases. Fig. 1 illustrates the ion-water configuration as the concentrason changes. M represents the cation, H the hydrogen and 0 the oxygen atoms. In dilute solutions (as rn --+ 0) the triangles MH,P, and MHP coincide so that 4 = 4, = 8 = 112.9'. As the concentration increases up to I,,,, MHP swings away, MH shortens and the angles a and 4 change. 4 has a minimum values of 93' at Imax, and assuming the 0-H value3386 ION-WATER GEOMETRY TABLE 2.-cOMPARISON OF CALCULATED AND EXPERIMENTAL VALUES OF CATION-DEUTERIUM DISTANCES ~ ~~ solute amax “ P expt. eqn (4) and (5) theoretical type, year NiC1, Ni2+ NiC1, LiCl Li+ LiCl LiI LiF Li+ CaC1, Ca2+ Na+ NaCl NaF Na+ K+ KF KCl K+ c s + CSCl CsF 18.9 19.9 17.9 18.8 16.9 15.1 13.23 9.15 8.53 4.38 2.55 1.38 0.258 0.0 9.95 8.53 3.57 2.2 2.2 0.555 0.278 0.1 0.0 13.47 8.53 0.0 8.53 2.2 0.555 0.555 0.278 8.53 2.22 0.555 0.555 0.278 0.0 8.53 2.2 2.2 0.0 - 2.67 f 0.02 2.67 2.67 2.76 2.80 2.80 f 0.03 2.50f0.02 2.55 f 0.02 - - - - - - - - - 2.93 & 0.05 - - - - - - -_ - - - - - - - - - - - 2.65 2.63 2.625 2.66 2.71 2.78 2.90 3.02 2.53 2.523 2.57 2.62 2.62 2.74 2.79 2.83 2.90 2.92 2.886 3.35 2.683 2.79 2.92 2.92 2.96 3.031 3.15 3.29 3.29 3.35 3.49 3.451 3.59 3.59 3.97 - - - - - _ _ _ _ _ _ __ _- - - 3.0 2.67 k 0.04 2.5 2.6-2.7 - - - - - - 2.8 2.8 3.0 2.8 3.0 3.8 3.2 3.5 3.4 - - - 3.6 3.6 - - - - - - - - - - __ - Cls, 1976 MD, 1981 CIS, 1976 Cls, 1978 - __ - - - MD, 1976 MD, 1976 CIS, 1976 CIS, 1976 Cls, 1978 Cls, 1976 Cls, 1976 Cls, 1976 CIS, 1978 - - - MD, 1976 MD, 1976 - a At I = 8 .5 3 ~ ~ ; ionic strength, tz+z-vm, where v is the number of moles of ions per mole of solute, z is the electronic charge, m is the molality [mol (kg H,O)-’]; experimental and theoretical values from ref. ( 1 ) and (2), MD, 1981 from ref. (7); values of (r++A+) from table 1, 3rd column from left. of 0.95 A remains constant, a reaches a maximum value of 15-20 O, depending on the ion (table 2). Above Imax, either a remains fixed and the 0-H distance changes or a decreases and 0-H remains fixed. On the other hand both 0-H and a could change. If a is fixed, the value of 0-H is given by rOH = (Y:O + r$H -2r+Or+H cos amax)+.( 6 )J. V. LEYENDEKKERS 3387 FIG. MPO M p/#' 'mi" 5 1 .-Cation-water geometry as a function of concentration. do = B = 1 1 3 O , dmin = 93O, scale for Li+ = PoHo = MP, = P,H, = (r++A+) OH,:]= 0, OH,:I= 1.0, OH,:I= 8.53, OH,:]= 1 8 ~ . amax values listed in table 2. Using eqn (6), with I = 18m, for a NiC1, solution rOH is 1.01 A whereas for a LiCl solution rOH is 1.02 A. From the above equations the values of r+H as rn 4 0 are given by r+O+0.95 A. 4. ANION-WATER GEOMETRY Neutron and X-ray data are avialable for the halides. A preliminary fit of all the data indicated a linear relationship between the anion-oxygen distance, rPO and (r-+A-). However, only the most accurate data, the neutron diffraction results for CI- of Enderby and Neilson,l? were used to derive the following equations (double weight was given to the CaCI, data, which are the most accurate): and rPO - 1.1 = 2 sin 0112) (r- + A-) x/2-xo/2 = 5.74fi- 1.071 with x 0 / 2 = 30°, so that, as rn -+ 0 r-O = 1.1 +(r-+A-).(7 b) Values of rw0 calculated from eqn (7a) and (8) are compared with the experimental and theoretical values in fig. 2. Since eqn (7a) and (8) were derived from the data of Enderby and Neilson it is interesting that the results for DCl around l g obtained by Ohtomo et ~ 1 . ~ are consistent with the calculated values. As yet, there is no definitive method for analysis of neutron diffraction data on aqueous ionic solutions. The Enderby group used first- and second-order difference spectroscopy, whereas Ohtomo's group used a subtraction m e t h ~ d .~ The latter method is apparently more accurate for dilute solutions, so that the good agreement with the extrapolated value [from eqn (7a) and (S)] indicates that the form of eqn (8) is correct. The values of r-o do not seem to be very sensitive to the cation (1 2 different cations, fig. 2) and eqn (8) appears adequate for all the halides. Only theoretical values of rp0 are available for F-. The calculated values from eqn (7a) and (8) [with (r+A) = 1.081 range from 2.2 to 2.4 for I ranging from 0 to 7 . 1 9 ~ . At I = 2 . 2 ~ the calculated value is 2.38 whereas the3388 I ON-W A TER GEOMETRY 0 2 4 6 8 10 12 14 16 18 20 Ilmol (kg H,O)-' FIG. 2.-Anion-oxygen distances as a function of concentration [(a) I-, (b) C1-, ( c ) Br-1: 0, neutron diffraction data with error bars' values at lm are from ref. ( 5 ) ; 0 , X-ray data from the compilation in ref.(1); x , MD; +,Clsaslistedinref. (2). (r+A)valuesinA: C1-, 1.8;Br-, 1.97;1-, 2.22. Curvescalculated from eqn (8) and (9). FIG. 3.-Anion-water geometry as a function of concentration. x-,, = 60°,xmax = 75.4O. H,,, corresponds to Qmax. AH = scale for CI-, O H = 1.1 A. O,, I = 0 ; 0,, I = 3 or 13m; O,,,, I = 7 . 2 ~ . A P = Q P = (r+A). a- = ( O A H (see text). theoretical predictions are 2.2 (MD)2 and 3.0 (Cls);2 at 0 . 5 5 5 ~ ~ the calculated value is 2.30 compared with 2.3-2.7 (Cls).2 For OH- [with ( r + A) g1.241, the calclulated values at 17.5m, 4.48~1 and 2m are 2.53, 2.61 and 2.56, whereas the X-ray values are all 2.9 A ( t h e n l y data available').From eqn (8) x/2 reaches a maximum (37.7O) at I = 7 . 1 9 ~ . If the anion-deuteriumJ. V. LEYENDEKKERS 3389 distance, T - ~ , remains fairly independent of concentration, the changes of the anion-water geometry with concentration can be interpreted via fig. 3. A , 0 and H represent the positions of the anion, oxygen and hydrogen, 0 Q is fixed at 1 . 1 A [eqn (7a)], AH remains constant in length but moves along the arc SQ,,, as indicated, whilst A 0 increased until Imax is reached. kc very dilute concentrations (m + 0), assuming rOH remains constant at 1 . 1 . (a value indicated from the LiCl results at higher concentrations), xo is 60°, A 0 equals [(r+A)+ 1.11, and AH, is inclined to A 0 by ca. 20'. As the concentration increases up to ca.3m, H , moves closer to A 0 and finally, at Imax, em,, and Hmax coincide and A , Hand o a r e linear. Further increases in the concentration result in 0 moving back towards 0,. Of course, the real situation might be more complex than this since A H might change in length. This does not seem to be the case for LiC1, but the value for NiCl, is slightly larger than those for the alkali-halide solutions. For constant A H ( Y - ~ ) , it can be seen that r-H = AQmax = 2 sin ( xrnax/2) (r- + A-) (9) withxmax/2 = 37.7'. This gives values of r - H of 2.20,2.41,2.72 and 1.32 for Cl-, Br-, I- and F-, respectively. The experimental value for C1- is ca. 2.23 (k0.04) for Li+, Na+, Rb+, Ca2+ and Ni2+, for I ranging from 3.57 to 13.47m. The ion-water distances calculated above are within thepredicted (theoretical) ranges.The MD values for C1- and the cations are remarkably close to the T.T.G. values at 2m - (table 2 and fig. 2). 5. EFECTIVE PRESSURE AND ION-WATER GEOMETRY The T.T.G. effective pressure, p e , is obtained from the density of the solution via the equation', Pe = ( 1) (BT+ 1) (10) where f ( m ) = a,m + a,m3/2 + a2m2. The a coefficients are related to the coefficients of the Masson equation for the apparent molal volume, dv, given by 4v = 4: -I- Slmt + SLm (1 1) with a, = (My/,-&)/J = - K O (BT+ l ) / J ; a, = - Sk/J, and a, = - S;l/J. K' is the limiting value of the partial molal compressibility, B, is the Tait parameter (3005 bar at 25 'C), J = 315v, (u, is the specific volume of water). The T.T.G.volume My/, is given by and the distance ( r + A) has been discussed in section 2. My/, = 4 ~ N x (r+A)3 = 2.52 (r+A)3 The partial molal quantity pe is given byll pe = 2pe/8rn = In 10(p,+~,+ I ) (a,+ 1.~a,m~+2a,m) p z = ( d p J 2 ~ ~ ) ~ = In 10(BT+ l ) ~ , . = In I O(B, + 1 ( I . 5a,mi) ( 1 2 4 (12b) so that as rn + 0 (indicated by L) In dilute solutions where p e + B, p,, - p where3390 I ON-W A TER GEOMETRY with irepresenting an ion in the solution; z and v are defined in section 3 . For a single-salt solution w = $z+z-v. The effective pressure depends on the attractive forces between the ions and the induced or residual charges on the water molecules, and has been discussed in detail.l01 l1 The problem of interest here is how p e and pe are related to the ion-water geometry derived above.From eqn (lo), (1 2a) and (1 3) it can be seen that this problem efiectively reduces to the problem of the relationship between the geometry and S, [S&/wi, eqn (1 3)], since S, is usually small.loyll The relationship of a, and the microstructure has been discussed previously. The ionic contributions of S, can be analysed on the basis of equations given previously,1° viz. S,+I4 = ($z$S,, + 1.2 - C+)Ii (14) where SDH is the Debye-Huckel limiting slope (1.865 cm3 kgb mor3l2 at 25 "C), C+ = 6.2Bn,, and C-= 10.5Bn,,, with where is the proton magnetic relaxation ratelo and the superscript O indicates pure water. As shown in sections 3 and 4, the ion-water geometry can be summarised in terms of (r + A) and the two angles 4 and a, for the cation and x and a- for the anion (fig.1 and 3). The angles 4 and x only depend on I t (and the sign of the charge), as do the terms, 1.214 and -0.6114 which were previously interpreted as arising from different orientations of anion versus cation.1° The angles a, and a_, however, are specific for each ion, like the terms C+ and C-. These comparisons suggest that a and the C terms of eqn (14) and (15) might be correlated. Values of a, were calculated using cos a, = (r$o + r:H - r&,)/2r+or+H (16) and similarly for a_. For the cations the values of r+H were calculated from eqn (4) and ( 5 ) and rOH was taken as 0.95 A (section 3). For the anions r-o values were calculated from eqn (7a) and (8) and rOH was taken as 1.1 A (section 4). The results of these calculations for I = l m - are shown in table 3.The C terms were calculated from (iztS,, + 1.2 - S,+) and ( -$z2SDH + 0.6 + S,-), using the ionic S , values from density datalo (table 3). The value of C for K+ is practically zero, and C for C1- is also small. This is because at moderate concentrations these two ions have little or no influence on the mean correlation time z, of water1, (B,,, is proportional to T~). Because of this, and the fact that a and C were found to be linearly correlated, we have for cations C+ = 0.62 [$z$ (a+ -a,+)] with an average deviation of 0.08. Eqn (1 7) is based on the a values from table 3, but (a+ -aK+) is nearly independent of concentration (within the errors of estimate of a, ca. f0.3'), except for H+. C- = 1.16 [iz? (a- - acl-)] (18) For the halides with an average deviation of kO.1.The angle a depends on the ionic strength and on the size of the ion [eqn (2), (4),J . V. LEYENDEKKERS 339 1 TABLE 3.-T.T.G. PARAMETERS AND ION-WATER GEOMETRY AT I = 1 mol (kg H20)-la H+ Li+ Na+ K+ Rb+ cs+ Mg2+ Ca2+ Sr2+ Ba2+ Zn2+ Cd2+ Ce3+ F- c1- Br- I- 1.54 1.74 1.85 2.09 2.23 2.38 1.76 1.99 2.04 2.15 1.78 1.93 1.99 1.08 1.80 1.97 2.22 1.62 1.98 2.14 2.53 2.77 3.02 1.99 2.40 2.49 2.70 2.02 2.29 2.41 2.33 3.15 3.34 3.63 2.39 2.70 2.87 3.24 3.46 3.69 2.73 3.08 3.16 3.33 2.76 2.99 3.08 1.32 2.22 2.41 2.72 16.2 15.5 14.2 12.8 12.2 11.6 14.8 14.0 13.8 13.5 14.6 14.2 14.2 14.4 12.7 11.9 11.4 1.03 1.18 1.67 2.12 2.35 2.44 2.58 3.24 3.5 4.0 2.70 3.18 5.67 1.18 0.2 1 -0.17 -0.59 1.10 0.95 0.46 0.01 - 0.22 -0.31 2.35 1.69 1.43 0.93 2.23 1.75 3.92 0.85 -0.12 -0.50 - 0.92 a Units: distances in A, S , and C in cm3 kgi m ~ l - ~ ' ~ , a in degrees.(5), (7a), (8) and (16)]. For a given ionic strength the angle a K + apparently represents a configuration that enables the adjacent water molecules to maintain their normal z, (and hence the normal intramolecular relaxation rate of pure water). The deviations of a, from a,+ thus reflect the magnitude of the change in the relaxation rate due to the ion. Similar considerations apply for the anions, with C1- having the minimal effect on z,. 6. DISCUSSION At present, computer simulations can predict the gross features of ion-water coordination, whilst the experiementally derived structural picture is more detailed. Recent progress in this regard has been thoroughly reviewed.' However, ion-water geometry as a function of concentration has not been elucidated and the foregoing correlations should be of interest in this regard.A detailed knowledge of the static structure as a function of concentration will greatly reduce most of the problems of interpretation that beset spectroscopic studies of all kinds.' In addition, a second problem, the relationship between the macroscopic and microscopic properties of aqueous electrolyte solutions, is not yet resolved. Most of the chemical thermodynamic properties of such solutions can be predicted via the T.T.G. model using density data. Since p e and My/, are the two basic quantities of this model the links between them and the microstructure certainly contribute to the solution of this second problem.The results given here support the following picture. The ion holds the adjacent water molecules tightly, consistent with the solvation-sheath concept. For the cation, the adjacent oxygen atoms remain at an average distance that is practically independent of the concentration, whereas the hydrogen atoms adopt relative positions according to the concentration. Apparently, these positions provide the minimum disruption to the microdynamic behaviour of the water molecule. A similar situation prevails for the anion, with the roles of oxygen and hydrogen roughly interchanged.3392 ION-WA TER GEOMETRY The value of S, can be interpreted as being made up of the Debye-Huckel term plus an ion-size related term, and a term due to an additional effect of the charge’s sign on the relative orientation of ion and water molecule (charge distribution effects), viz.Sv+ = ~z:[S,, + 0.62 (a+ - a,+)] + S,(@) S,-= $22 [SDH+ 1.16 (a--acl-)]+S,(x>. With the exception of the term S, the contributions tope can be interpreted in terms of the geometry and microdynamics of the ion and water molecule. The functions for 4 and x involve a term linear in I [eqn (5) and (S)] which will affect the orientations, particularly at higher concentrations. Probably S, is related to ion-water orientation effects incorporated in this linear term. The anticipated increase in the accuracy of experimentally derived neutron data’ should enable this problem to be resolved. J. E. Enderby and G. W. Neilson, in Water: A Comprehensive Treatise, ed F. Franks (Plenum, New York, 1979), vol. 6, chap. 1; and Rep. Prog. Phys., 1981, 44, 593. J. E. Enderby and G. W. Neilson, Adv. Phys., 1980, 29, 323. N. Ohtomo and K. Arakawa, Bull. Chem. SOC. Jpn, 1979, 52, 2755. N. Ohtomo and K. Arakawa, Bull. Chem. SOC. Jpn, 1989, 53, 1789. N. Ohtomo, K. Arakawa, M. Takeuchi, T. Yamaguchi and H. Ohtaki, Bull. Chem. SOC. Jpn, 1981, 54, 1314. D. W. Wood, in Water: A Comprehensive Treatise, ed. F. Franks (Plenum, New York, 1979), vol. 6, chap. 6. ’ Gy. I. Szasz, K. Heinzinger and G. Palinkas, Chem. Phys. Lett., 1981, 78, 194. * P. Schuster, W. Jakubutz and W. Marius, Top. Curr. Chem., 1975, 60, 1. J. E. Clementi and R. Barsotti, Chem. Phys. Lett., 1978, 59, 21. lo J. V. Leyendekkers, J. Chem. SOC., Faraday Trans. I , 1982, 78, 357. l1 J. V. Leyendekkers, J. Chem. SOC., Faraday Trans. I , 1981, 77, 1529. l2 H. G. Hertz, in Water: A Comprehensive Treatise, ed F. Franks (Plenum, New York, 1973), vol. 3, chap. 7. (PAPER 2/498)

ISSN:0300-9599    

DOI:10.1039/F19827803383

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982,78, 3393-3396 Formation of Manganese@) Porphyrin Derivatives from Manganese(II1) Derivatives by Ionizing Radiation BY RAMAKRISHNA RAO AND MARTYN C. R. SYMONS* Department of Chemistry, The University, Leicester LE 1 7RH AND ANTHONY HARRIMAN Davy Faraday Research Laboratory, The Royal Institution, London W 1X 4BS Received 26th March, 1982 Exposure of manganese(Ir1) porphyrin solutions to 6oCo prays at 77 K gave the manganese(I1) derivatives, with gI (apparent) z 6, A(55Mn) = 77 G and ID( 2 0.3 cm-I. The spectra are very similar to those for Mn" derivatives prepared chemically. Weak features in the g = 2 region for these complexes were obscured by features from a symmetrical Mn" derivative of unknown origin. These features grew in intensity on melting and re-freezing.These results suggest that low-temperature irradiation coupled with e.s.r. spectroscopy may be a suitable technique for detecting Mn"' derivatives in biological systems. Although manganese occurs quite widely in biological systems, many centres give no e.s.r. signals in the pure state. An important example is that of chloroplasts. We have established that the use of ionizing radiation at low temperatures (usually 77 K) is a powerful method for inducing specific electron addition,' and that this technique can usefully be applied to biological systems.2-6 It seemed possible that biological systems containing Mn1I1 ions, and therefore giving no detectable e.s.r. signals, might be conveniently converted into MnlI derivatives by simple electron addition under low-temperature conditions that should inhibit further reactions. The aim of this study was to test this hypothesis using MnlIr porphyrins as good model systems.The results show that the method is indeed feasible. Previous work7 has shown that MnlI1 porphyrins undergo very inefficient photoreduction in outgassed solution to form the corresponding MnIl porphyrin. Similar reductions can be achieved with chemical reductants,s such as dithionite, although both techniques have little application for biological samples. EXPERIMENTAL PREPARATION MnTSPP was prepared as described by Harriman and Porterg and purified by repeated chromatography on Dowex 50-wX8 cation-exchange resin followed by exhaustive dialysis against deionized water. MnTPP was prepared from tetraphenylporphyrin (chlorine-free, Aldrich) (1 g), dissolved in glacial acetic acid (80 cm3) and Mn(OAc), (4 g) was added.The mixture was refluxed for 4 h and then evaporated to dryness. The solid product was washed with water and dried overnight at 90 OC under vacuum. It was purified by repeated chromatography on alumina using CHCl, as eluent. 33933394 FORMATION OF MnI1 PORPHYRINS Dilute solutions (10-50 mmol dm-3) in methanol (CD,OD was used to minimise e.s.r. absorption from the solvent) were degassed and frozen as small beads in liquid nitrogen. They were exposed to 6oCo y-rays at 77 K in a Vickrad source to doses of 0.5-3.0 Mrad.* The e.s.r. signals assigned to Mn" grew steadily during this period. c 1000 G (9.0977 GHzI MI (55Mn) -"Z I I - 3/2 + 3200 G (9.0957 GHzI 50 G W H A I I I I I I MI ( Mn) -72 -% -% *% *3/2 +% FIG.1.-First derivative X-band e.s.r. spectra for a solution of Mn"' TSPP in CD,OD after exposure to 6oCo y-rays at 77 K. (a) At 77 K, showing the g = 6 feature assigned to species A (MnIITSPP) and (6) after annealing to the glass-point and re-cooling, showing features assigned to species B. (The main hyperfine features are indicated with MI values; the intermediate lines are forbidden transitions and the central feature is due to organic radicals.) E.s.r. spectra were measured on a Varian E-109 X-band spectrometer calibrated with a Hewlett-Packard 52461, frequency counter and a Bruker BH 12 E field probe standardised with a sample of diphenylpicrylhydrazyl. Samples were annealed by decanting the liquid nitrogen from the insert Dewar flask and monitoring the spectra continuously. Samples were re-cooled whenever significant spectral changes were observed.RESULTS AND DISCUSSION The solutions showed no e.s.r. features at 77 K before irradiation. After exposure, in addition to intense signals in the g = 2 region from *CD, and *CD,OD radicals, a broad resonance appeared in the g = 6 region which was just resolved into six hyperfine components from coupling to 55Mn nuclei ( I = $) [species A, fig. 1 (a)]. * 1 rad = J kg-I.R. RAO, M. C. R. SYMONS A N D A. H A R R I M A N 3395 After annealing to remove solvent features, two sets of lines from 55Mn in the g = 2 region were detected. One set, characteristic of Mn’I with only a small zero-field splitting, grew in intensity on annealing [species B, fig.l(b)]. At the same time, the other set, assigned to species A, decayed slightly and became less well resolved. TABLE 1 .- E.s.R. PARAMETERS FOR SOME MnI1 PORPHYRINS apparent values of ga 55Mn hyperfine couplingb complex g /I g, A / / A , Mn(II)TSPPC ca. 2.00 5.9 & 0.1 ca. 78 78-L- 1 Mn(I I)(TPP)(PY Id 2.00 5.96 74 74 a These are not true values of g (see text). G = lop4 T. D > 0.3 cm-’. D = 1.2 cm-1.8 SPECIES A Because of the great width of the e.s.r. features, the data given in table 1 have large error limits. Nevertheless, comparison of these results with results reported elsewhere for similar complexes7v8 show that simple electron addition must have occurred to give the MnII TSPP derivative.These results are typical of high-spin (S = 5) d5 systems with large tetragonal zero-field splitting parameters (D). The results require that D be greater than the microwave energy (ca. 0.3 ~ m - l ) . ~ Because of the absorption from species B, it was difficult to obtain details of the expected ‘parallel’ feature in the g = 2 region, but extra features, giving “gI1 ’’ z 2.00 and A,, z 78 G could be picked out. The g = 2 features are usually far weaker than the g = 6 features for such complexes and the hyperfine splitting is usually isotropic. We stress that the quoted values of g , , and gl are nothing to do with the true g-tensor components, but simply serve to identify transitions within the S = f manifold that are intense at X-band frequenciesg In fact, the values of g must be nearly isotropic at g = 2.00. We conclude that simple electron addition occurred on irradiation.The broadening observed on annealing presumably reflects some change in the environment as the Mnl* system relaxed to its equilibrium dimensions. It may be significant that the MnlI ion resides some 0.56 A above the plane of the porphyrin ring whereas the MnlI1 ion is much nearer to being in-the-plane. Consequently, upon reduction there must be a large geometry change and this may contribute towards the annealing effect. SPECIES B The e.s.r. spectrum for this species is typical of high-spin (S = g) MnI1 complexes with near cubic symmetry. In addition to the six main hyperfine components, intermediate lines due to formally forbidden transitions are apparent.The broadening of the hyperfine lines on going from low to high field indicates a small D component. Although this species is clearly a product of irradiation, it seems to us unlikely that it is formed directly from MnlI1 TSPP, since this would require the ejection of manganese from the porphyrin ring and resolvation. This is extremely unlikely in dilute solutions for which the only expected process is electron additi0n.l It is more likely that some MnlI1 impurity with a higher electron affinity than that of MnlI1 TSPP is involved, in which case species B is of no special significance.3396 FORMATION OF MnII PORPHYRINS CONCLUSION These results clearly establish that MnlI1 complexes which do not involve manganese clusters can be readily reduced by ionizing radiation, and the resulting MnIL complexes can be detected by e.s.r.spectroscopy in low concentration. Preliminary results with chloroplasts have not been successful. This may be because manganese pairs are present, as suggested by the recent work of Dismukes and Siderer.lo The success of the method with simple Mn porphyrins together with the failure to detect anything with chloroplasts suggests that more model systems are required. It should be possible to replace the haem in a natural protein, such as myoglobin, with a Mn porphyrin and so obtain a system intermediate between the simple porphyrins and the chloroplast. We thank the S.E.R.C. for a grant to D.N.R.R. M. C. R. Symons, Pure Appl. Chem., 1981, 53, 223. M. C. R. Symons and R. L. Petersen, Biochim. Biophys. Acta, 1978, 535, 241. M. C. R. Symons and R. L. Petersen, Biochim. Biophys. Acta, 1978,537, 70. M. C. R. Symons and R. L. Petersen, Biochim. Biophys. Acta, 1978, 535, 247. M. C. R. Symons and R. L. Petersen, J . Chem. Res. (S), 1978, 382; (M), 1978, 4572. 'I A. Harriman and G. Porter, J. Chem. Soc., Faraday Trans. 2, 1979, 75, 1543. I. A. Duncan, A. Harriman and G. Porter, J. Chem. Soc., Faraday Trans. 2, 1980, 76, 1415. A. Harriman and G. Porter, J. Chem. SOC., Faraday Trans. 2, 1979, 75, 1532. * M. C. R. Symons and R. L. Petersen, Proc. R. SOC. London, Ser. B, 1978, 201, 285. lo G. C. Dismukes and Y. Siderer, FEBS Lett., 1980,121,78; Proc. Natl Acad. Sci. USA, 1981, 78, 274. (PAPER 2/524)

ISSN:0300-9599    

DOI:10.1039/F19827803393

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 3397-3408 Solvation and Hydrophobic Hydration of Alkyl-substituted Ureas and Amides in NN-Dimethylformamide + Water Mixtures BY AART Rouw AND Gus SOMSEN* Department of Chemistry, Free University, De Boelelaan 1083, 1081 HV Amsterdam, The Netherlands Received 31st March, 1982 Enthalpies of solution of five alkyl-substituted ureas and seven different amides have been determined at 298.15 K in mixtures of NN-dimethylformamide (DMF) and H,O. Methyl substitution of the ureas causes changes in the enthalpies of transfer from H,O to DMF which show that either side of these molecules is solvated independently. From the measurements on the amides it is concluded that methyl substitution at the N atom gives changes in the enthalpies of transfer from H,O to DMF which are different from those caused by methyl substitution at the C atom.Analysis of the data in the mixed solvent shows that introduction of more or longer alkyl groups into the molecules makes both ureas and amides considerably more hydrophobic. After accounting for the influence of the NH protons, in both ureas and amides, the enthalpic effect of hydrophobic hydration of the solutes was calculated by application of a clathrate-like hydration model. Enthalpic effect of N-substituted methyl groups in ureas and amides prove to be virtually equal. The variation in the enthalpic effects of hydrophobic hydration with the number of C atoms in the n-alkyl group is comparable to that found for alcohols and amines. This paper is part of a project in which we are investigating the solvation of hydrophobic solutes in aqueous mixed solvents.After an earlier series of reports on tetra-alkylammonium salts [for references see ref. (l)] we have recently focussed our attention on non-electrolytes. Thus far we have studied monohydric alcohols' and alkylamines, as two representative types of model compounds. In our investigations we determine the enthalpies of solution of a (hydrophobic) compound in mixtures of water and an organic solvent. Since aprotic NN- dimethylformamide (DMF) has proved to be a suitable reference solvent with regard to water and since it is completely miscible with water, we confine ourselves mainly to DMF + H,O mixtures. Due to hydrophobic hydration the enthalpy of solution of a hydrophobic solute shows a strong endothermic shift upon addition of small amounts of DMF (or other cosolvents) to water.This is caused by the collapse of the hydrophobic hydration sphere around the solute. Analysis of this effect in terms of a model description3 yields information on the hydrophobic properties of the compound. In this paper we turn our attention to two other types of organic model compounds, i.e. substituted ureas and amides. Because of their biochemical importance aqueous solutions of these compounds have been studied extensively.*-15 Among the thermo- dynamic data enthalpies of solution in water have been reported for both alkylureaslO and amides.I1-l4 Enthalpies of solution in pure non-aqueous solvents have been published for amides only.14 As far as we know enthalpies of solution of these compounds in aqueous mixtures, covering the whole composition range from pure water to pure organic solvent, have never been determined, with the exception of urea, for which such measurements were reported in an earlier study from our 1ab0ratory.l~ These data provide a useful background to this work.33973398 SOLVATION + HYDRATION OF UREAS + AMIDES In this paper we discuss the solution properties of ureas and amides substituted with alkylgroups in various ways. The comparison of the hydrophobic effects of these compounds can provide information on the properties of alkyl groups at different positions in the solute molecule. To this aim we determined calorimetrically enthalpies of solution of twelve compounds in the mixed solvent DMF+H,O over the whole composition range.We investigated the following compounds : methylurea (MU); 1,l -dimethylurea (1,l -DMU); 1,3-dimethylurea (1,3-DMU); tetramethylurea (TMU); ethylurea (EU); formamide (FA); N-methylformamide (NMF); NN-dimethyl- formamide (DMF); acetamide (AA); NN-dimethylacetamide (DMA); Nn-butyl- acetamide (NBA) and butyramide (BA). For comparison we also determined the enthalpies of solution of TMU in the non-aqueous mixtures of NN-dime t h yl formamide + N-met hy 1 formamide. EXPERIMENTAL The enthalpies of solution were measured with an LKB 8700 precision calorimetry system, equipped with a 100cm3 reaction vessel. The experimental procedure and a test of the calorimeter system have been reported before.16 Solid compounds were transferred into glass ampoules of 1 cm3, which were closed with silicone rubber stoppers and sealed with wax.Liquid solutes were transferred with a syringe into another type of 1 cm3 ampoule. These ampoules had a narrow neck which could be sealed with the aid of a microburner. NN-dimethylforrnamide (Baker, Analyzed Reagent) and N-methylformamide (Merck, zur Synthese) were purified and dried as before.' The solvent mixtures were prepared by mass. For the aqueous mixtures we used distilled deionized water. MU (EGA, Steinheim) and I ,3-DMU (Koch-Light, Purissimum) were recrystallized from absolute ethanol. 1,l -DMU (Merck, zur Synthese) was recrystallized from a mixture of ethanol and chloroform. EU (Merck, zur Synthese), AA (Baker, Analyzed Reagent) and BA (Fluka, Purum) were recrystallized from a mixture of ethanol and diethylether.These solid compounds were dried under vacuum over P,O, for at least 48 h before use. TMU (Fluka, Purum) was distilled at reduced pressure. FA (Baker, Analyzed Reagent), NMF (Merck, zur Synthese), DMA (Baker, Analyzed Reagent) and NBA (Eastman) were purified by distillation from NaOH under reduced pressure. In all cases only the middle fraction was used. DMF (Baker, Analyzed Reagent) was used as such. All these liquid solutes were dried over molecular sieve (Baker, 4A), except for FA, where we used 3A molecular sieve. The water content of these liquids, determined with a modified Karl Fisher titration," was always below 0.02 mass %. The purity of the liquids was determined by gas-liquid chromatography and found to be > 99.8 mol %.RESULTS The amount of solute used during the calorimetric experiments led to final solute concentrations between 0.005 and 0.02 mol dm-3. In this range we did not observe any concentration dependence of the enthalpies of solution, so we considered our measured enthalpy values as those at infinite dilution and hence as standard enthalpies of solution, AH?. Correction for the presence of solute vapour in incompletely filled ampoulesls proved negligible because of the low volatility of the solutes. The values of the enthalpies of solution of the urea compounds in mixtures of DMF+H,O are listed in table 1, together with the average deviations. The values represent the mean result of two to four measurements agreeing within 0.15 kJ mol-1 and they refer to 298.15 K.Corresponding results on the amides are listed in table 2. In addition table 3 gives the results for TMU in mixtures of DMF+NMF. Some of the enthalpies of solution in pure water can be compared with results from the literature. Values of 1,3-DMU and TMU have also been determined by Ahluwalia and coworkers.1° Their results differ from ours by 0.3 and 0.6 kJ mol-l, respectively.A. ROUW AND G . SOMSEN 3399 TABLE 1 .-ENTHALPIES OF SOLUTION OF ALKYL-SUBSTITUTED UREAS IN DMF + H20 MIXTURES AT 298.15 K MU 1,l -DMU 1,3-DMU 0.000 0.060 0.189 0.325 0.450 0.550 0.650 0.770 0.850 0.900 0.950 1 .ooo 1 1.43 f 0.04 11.13f0.02 1 1 .OO f 0.03 11.19+0.01 11.75 f 0.01 12.26 & 0.0 1 12.84f 0.01 13.07 fO.O1 12.68 f 0.03 12.24 & 0.01 11.73 f 0.01 11.19 f O .O 1 19.56 f 0.05 19.35f0.08 19.26 f 0.03 19.26 f 0.05 19.65 f 0.03 19.70&0.02 19.22 fO.01 17.76f0.01 16.09 f 0.01 14.90 f 0.03 13.47 f 0.02 12.01 fO.05 11.17 f 0.02 10.68 f0.04 10.40 f 0.05 10.30_+0.04 10.45 k0.02 10.45 & 0.0 1 10.12 f 0.01 8.58 & 0.05 6.60 & 0.03 4.91 f 0.06 3.06 f 0.07 1.05 & 0.03 x w TMU EU 0.000 0.060 0.189 0.325 0.450 0.550 0.650 0.770 0.850 0.900 0.950 1 .ooo 0.17 & 0.02 0.10 f 0.01 0.05 f 0.03 - 0.21 f 0.02 - 0.96 f 0.04 - 2.60 & 0.04 - 6.54 k 0.04 - 10.80f0.01 - 14.24 f 0.06 - 18.42&0.01 -24.53k0.01 - 13.35 f 0.03 13.19f0.01 13.16 f 0.05 13.48 f 0.04 14.20f0.01 15.06 f 0.01 15.82f0.01 15.87 & 0.01 14.91 f0.02 13.84 f 0.02 12.40 & 0.0 1 10.69 f 0.05 On the other hand the difference between our values for FA, NMF, AA and NBA and those determined by Wadso and coworkers12~13 is always smaller than 0.07 kJ mo1-l. The agreement with results for FA, NMF, DMF, AA and DMA reported by Stimson and Schrierll is also very good.In a recent study Spencer et d . 1 4 have determined enthalpies of solution of FA, NMF, DMF, AA and DMA in both water and DMF. Considering the limited precision of their results the agreement with our values is fair. No literature data are available for the enthalpies of solution of MU, 1,l -DMU, EU and BA in water and for those of the substituted ureas, BA and NBA, in DMF. DISCUSSION The results for the enthalpies of solution of the alkylureas and the amides, listed in tables 1 and 2, are visualized in fig. 1 and 2.To facilitate the comparison we have plotted the enthalpies of transfer from H20 to the DMF+H,O mixtures, AHtr3400 SOLVATION -k HYDRATION OF UREAS + AMIDES TABLE 2.-ENTHALPIES OF SOLUTION OF AMIDES IN DMF 4- H 2 0 MIXTURES AT 298.15 K x w FA NMF DMF AA 0.000 0.060 0.189 0.325 0.450 0.550 0.650 0.770 0.850 0.900 0.950 1 .ooo - 3.88 f 0.03 -3.71 f0.05 - 3.28 f 0.03 - 2.44 f 0.02 - 1.32f0.01 -0.17 f 0.02 1.05 f 0.02 2.31 fO.01 2.71 fO.01 2.76 f 0.02 2.61 f 0.05 1.97 50.01 - 0.24 f 0.0 1 - 0.22 f 0.01 - 0.07 f 0.0 1 0.21 fO.01 0.50 f 0.03 0.65f0.01 0.48f0.01 -0.52 0.01 - 2.08 f 0.05 -3.36k0.01 - 4.95 f 0.02 -7.1 1 fO.01 0.00 0.06 f 0.01 0.04 f 0.03 - -0.19f0.03 - 0.83 f 0.05 - 2.08 f 0.06 -4.89f0.01 - 7.62 f 0.04 -9.69f0.01 - 12.02 Ifr 0.01 - 15.27 f 0.01 11.56f0.01 11.51 kO.01 1 1.49 f 0.02 1 1.83 f 0.02 12.40 f 0.03 12.94f0.01 13.31 f0.02 1 3.02 f 0.04 12.25 f 0.04 1 1.49 f 0.02 10.75 f 0.04 9.73 f 0.02 DMA NBA BA 0.000 0.060 0.189 0.325 0.450 0.550 0.650 0.770 0.850 0.900 0.950 1 .ooo -0.08f0.01 - 0.34 f 0.01 - 0.66 f 0.02 - 1.05f0.01 - 1.72 f 0.03 - 2.70 f 0.02 -4.58k0.04 - 8.29 & 0.04 - 11.84_+0.02 - 14.48 f 0.03 - 17.54f0.01 - 21.46 f 0.04 3.42 f 0.01 3.35 fO.01 3.43 f 0.01 3.76 & 0.04 4.29 & 0.02 4.76 & 0.02 4.81 f0.03 3.12 f 0.01 -0.31 fO.01 -3.71 f0.03 - 8.45 f 0.01 - 14.79 f 0.03 14.31 f0.03 14.40 & 0.05 14.83 f0.02 15.64f0.01 16.82 f 0.01 17.75 f 0.03 1 8.64 f 0.05 18.56f0.01 16.84 f 0.05 14.98 fO.O1 12.33 kO.01 9.07 f 0.01 TABLE 3.-ENTHALPIES OF SOLUTION OF TETRAMETHYLUREA IN MIXTURES OF DMF + NMF AT 298.15 K X,,, AHp/kJ mol-l 0.000 - 2.43 & 0.03 0.250 - 1.39 f 0.02 0.500 -0.73 fO.01 0.750 - 0.23 f 0.02 1 .ooo 0.17 f 0.02 (H,O -+ DMF+H,O), against X,, the mole fraction of water in the mixture.Enthalpies of transfer show the same variations as enthalpies of solution, but they do not contain contributions from interactions in the pure solid or the liquid solutes. Hence they reflect directly the enthalpic differences in solvation between the different mixtures and water. For comparison we have included in fig. 1 the results for urea (U) in the earlier study of De Visser et aL.l5A. ROUW AND G . SOMSEN 340 1 s l \ . X + LL 2 t 0, E I 1 I I I 0.2 0.4 0.6 0.8 FIG. 1 .-Enthalpies of transfer from H,O to mixtures of DMF + H,O for various alkyl-substituted ureas.xw +24 - I - 0 E 2 +16 3: + LL 2 +8 t . s 0, z I I I I 0.2 0.4 0.6 0.8 FIG. 2.-Enthalpies of transfer from H,O to mixtures of DMF+H,O for various amides. *w In both figures we see that, upon introduction of either more or longer alkyl groups into the ‘parent’ compounds U and FA, the curves of AH,,(H,O+ DMF+H,O) gradually come to show the features of hydrophobic hydration which we have described in our previous work.1*2 The most striking feature is the strong endothermic shift in the enthalpy of transfer in mixtures with a small mole fraction of organic 110 FAR 783402 SOLVATION + HYDRATION OF UREAS + AMIDES TABLE 4.-ENTHALPIES OF TRANSFER OF SOME ALKYLUREAS AND AMIDES FROM H20 TO DMF AT 298.15 K compound AH,,(H,O -+ DMF)/kJ mol-l U MU 1,1-DMU 1,3-DMU TMU FA NMF DMF AA NMA DMA ~~ - 9.43 0.24 7.55 10.12 24.70 - 5.85 6.87 15.27 1.83 14.43a 21.38 a Calculated +731 t 9.67 /Ii1 I3 /Ii1 +9.88 \ from data in ref.(1 1) and (14). DMU \9.88, I + 34.13 FIG. 3.-Changes in the enthalpies of transfer from H,O to DMF for ureas, caused by methyl substitution. Values in parentheses were estimated. (Energies are in kJ mol-l.) solvent. This is most pronounced for TMU, the compound with the largest number of methyl groups, and for NBA, which bears the longest alkyl chain. Differences between the curves of compounds containing an equal number of methyl groups, such as MU, NMF and AA or 1,I-DMU and 1,3-DMU, show that the position of the substituents also determines the shape of these curves. From the results we have calculated enthalpies of transfer from water to pure DMF, AH,,.(H,O -+ DMF), for the parent compounds U, FA and AA and their N-methyl- substituted derivatives.The results are listed in table 4. In urea the introduction of one methyl group, to form MU, causes a shift in AH,,(H,O + DMF) of 9.67 kJ mol-l. This is distinctly higher than the effect of the introduction of a second methyl group on the same N atom, to form l,l-DMU, which amounts to 7.31 kJ mol-l. However, introduction of a second methyl group on the other N atom, to form 1,3-DMU, results in a shift (9.88 kJ mol-l) which is very close to that of the first methyl group. This is a strong indication that the NH, and NHCH, groups on either side of the molecules are solvated independently, which should imply that the enthalpy of transfer of TMU will be larger than that of U by 9.67+ 9.88 + 7.31 + 7.31 = 34.17 kJ mol-1 (see fig.3). Experimentally the difference in the enthalpies of transfer of U and TMU is found to be 34.13 kJ mol-l, in excellent agreement with the predicted value. Inspection of the data in the solvent mixtures shows that this additivity scheme also applies to aA. ROUW AND G . SOMSEN 3403 good approximation to the enthalpies of transfer from water to the mixtures. Consequently the assumption of independent solvation of both sides of the urea molecules is substantiated by the experimental results in both the pure solvents and the mixtures. Although the magnitudes of the changes depicted in fig. 3 are determined by different effects, e.g. changes in solute-solvent hydrogen bonding, we believe them to be due largely to hydrophobic hydration of the methyl groups. For the amides the enthalpies of transfer in table 4 show that in this case also the change in AH,,(H,O + DMF) caused by a first methyl group introduced into FA and AA to form NMF and NMA (12.72 and 12.60 kJ mol-l, respectively) appears to be higher than the effect of a second N-methyl substitution (8.40 and 6.95 kJ mol-l, respectively, see fig.4). Because the amides are not symmetrical, introduction of a methyl group at the other side of the molecule will give different results. Indeed the difference in AH,,(H,O -+ DMF) between AA and FA amounts to 7.68 kJ mol-l, and such a lower value is also found for the differences between NMA and NMF and between DMA and DMF (see fig.4). Hence the effect of C-methyl substitution is definitely smaller than the effect of N-methyl substitution. This is not unexpected, since on C-substitution an aprotic CH group is replaced by a CCH, group, whereas on N-methyl substitution a protic NH group is transferred into an aprotic NCH, group. + Z68 +Z56 + 6.11 In our previous study on amines2 we have shown that, besides the enthalpies of transfer, a good and even better description of the enthalpies of solution in the mixtures is provided by the quantity 6 given by 6 = AH,,(H20 -+ DMF+H,O)-(1 -XW)AH,,(H,O -+ DMF). (1) 6 represents the deviation of the enthalpies of transfer in the mixtures from a change proportional to X,. We have shown that the values of 6 for a given solute are related to its hydrophobic properties.This relation can be described quantitatively by a model3 which involves the formation of clathrate-like cages of water molecules around each alkyl group in the water. These cages break down when cosolvent is added. According to this model 6 can be written as 6 = Hb(W) (XG-X,). (2) In this expression the two parameters Hb(W) and n represent the enthalpic effect of hydrophobic hydration of the solute in pure water and the number of hydration sites of one alkyl group, respectively. Although this model approach is a first approximation3+ l9 it has proved to be a useful means of comparing the hydrophobic properties of a wide range of solutes. However, one should realize that in this model eqn (2) assigns 6 values exclusively to the hydrophobic hydration of the alkyl groups in a solute. Any contributions from other effects are disregarded.This implies that in the absence of hydrophobic effects 6 should be zero at each mixture composition, i.e. the enthalpy of solution should change proportionally to the mole-fraction composition. Such behaviour has indeed I 10-23404 SOLVATION + HYDRATION OF UREAS + AMIDES TABLE 5.-vALUES OF 6 IN kJ m0l-l OF UREA [INTERPOLATED DATA OF REF. (1 5)] AND DIFFERENCES BETWEEN 6 VALUES OF ALKYL-SUBSTITUTED UREAS AND AMIDES AND THE RESPECTIVE PARENT COMPOUNDS x w W ) d(MU) - d(U) G(NMF) - A(FA) 0.000 0.060 0.189 0.325 0.450 0.550 0.650 0.770 0.850 0.900 0.950 1 .ooo 0.00 -0.84 - 2.20 - 3.06 - 3.33 - 3.08 -2.53 - 1.72 - 1.14 -0.77 - 0.39 0.00 0.00 0.55 1.82 2.90 3.76 4.04 4.10 3.54 2.59 1.80 0.92 0.00 0.00 0.61 1.98 3.14 3.90 4.18 4.06 3.32 2.38 1.68 0.89 0.00 x w d(1,l -DMU)-d(U) d(DMF)-&FA) d(DMA)-d(AA) 0.000 0.060 0.189 0.325 0.450 0.550 0.650 0.770 0.850 0.900 0.950 1 .ooo 0.00 1.08 3.34 5.22 6.82 7.38 7.10 5.74 4.10 2.92 1.48 0.00 0.00 1.10 3.46 5.44 6.76 7.08 7.26 5.18 3.74 2.66 1.52 0.00 0.00 0.96 3.18 5.12 6.32 6.76 6.46 5.38 4.16 3.26 1.92 0.00 been found for various non-hydrophobic solutes in DMF + H,O m i x t ~ r e s ~ 9 ~ ~ and also for hydrophobic solutes in non-aqueous mixtures such as DMF + NMF.lv2y20 This suggests that for these compounds the solvation changes gradually from one solvent to another.Fig. 1 shows that for the non-hydrophobic U in DMF+H,O the enthalpies of transfer deviate in a negative sense from linear behaviour.Values of 6(U), which were calculated by interpolation of the data in ref. (1 5), are listed in table 5. One might expect that these (negative) values for the parent compound cannot be ignored when considering the results for substituted ureas, because the latter contain similar polar groups. In this respect it is interesting to look at the behaviour of TMU. For X , < 0.5, where hydrophobic effects can be excluded, the enthalpies of transfer of TMU change linearly with X,. Consequently we think that the deviation from linear behaviour of U is mainly due to its NH protons. Support for this view is provided by our results for TMU in DMF + NMF. Inspection of table 3 shows that the enthalpies of solution of TMU in this mixture change almost linearly with composition.On the other hand the previous study of De Visser et aZ.15 has shown that the enthalpies of solution of U in DMF + NMF show negative deviations from linearity, comparable to those in DMF+H,O.A. ROUW AND G. SOMSEN 3405 Hence in order to obtain significant information from our 6 values on the hydrophobic hydration of alkyl groups in the substituted ureas we should apply a ‘correction’ which accounts for the influence of the NH protons. Because the substituted ureas contain a different number of NH protons we should apply corrections to 6 proportional to this number. For simplicity we now assume that the d(U) values are composed of four additive 6(NH) contributions : 6(NH) = 1/4 6(U) (3) and that these contributions can also be applied to the alkyl-substituted ureas.The corrected 6 values obtained in this way refer to the alkyl parts of the substituted ureas only. Regression analysis with respect to eqn (2) of a set of these values for a particular compound now yields values for the model parameter Hb(W). The results are listed in table 6. For the ureas table 6 also gives the contribution per methyl group to Hb(W). TABLE 6.-ENTHALPIC EFFECTS OF HYDROPHOBIC HYDRATION OF N-METHYL-SUBSTITUTED UREAS - Hb(W)/kJ mol-l solute ‘correction’ total per methyl group ~ MU - id( U) 7.1 7.1 1 , l - D M U -$6(U) 1 3 . 0 6.5 1 , 3 - D M U -@(U) 1 3 . 2 6.6 TMU - 24.6 6.2 We note that they are very similar, which indicates that to a good approximation the hydrophobic hydration of the methyl groups is independent of the solvation of the neighbouring group in the solute molecule.A similar independence has been found for other types of ~ 0 1 u t e s . ~ ~ ~ ~ Table 6 also shows a small trend in the Hb(W) values per methyl group (from MU to TMU). In our view this trend bears limited significance, because of the approximations made in the correction procedure. Some of the amides contain NH protons comparable to those in the ureas, so we can expect that in the analysis of the results for these compounds corrections will again be necessary. However, the situation is more complicated here, because amides other than FA contain alkyl groups substituted in different ways, i.e. at the C atom and/or at the N atom. These may contribute to 6 differently. Consequently the data cannot be treated in the same manner as for the ureas.For both ureas and amides, however, the effects of N-methyl substitution alone can be estimated by elimination of the contribution to 6 of the CH or C-substituted alkyl groups. This is done by comparing the differences 6(MU) - 6(U) and G(NMF) -&FA) and the differences 6( 1 , 1 -DMU) - 6(U), G(DMF) - 6(FA) and S(DMA) - 6(AA). The first are due to one methyl group and one NH group, the second to two methyl groups and one NH, group. The results given in table 5 show a close agreement between these differences. This strongly suggests that ureas, formamides and acetamides possess equal contributions of the methyl and NH groups to 6. This implies that the 6(NH) values of the ureas [eqn (3)] can also be used as a correction for the amides.Corrected 6 values obtained in this way were subjected to a curve-fitting program with respect to eqn (2). The resulting values for H6(W) are listed in table 7. It appears that in FA the hydrophobic effect of the CH group is equal to - 3.7 kJ mo1-l. Hb(W) of the CH3406 SOLVATION + HYDRATION OF UREAS + AMIDES group in the formamides can also be obtained from the difference between $Hb(W) for TMU and Hb(W) for DMF (see table 6). The result, -3.8 kJ mol-', is in remarkable agreement with the value from table 7. Since no corrections have been applied to the experimental results for TMU and DMF, this close agreement corroborates the reliability of our correction procedure for the amides. These considerations show that FA is slightly hydrophobic.Indeed the curve for FA in fig. 2 shows a slight maximum at high X , and is distinctly different from the curve for (non-hydophobic) U in fig. 1. The Hb(W) values in table 7 for AA and BA refer to a methyl and an n-propyl group, respectively, as C-substituents. For NMF and NBA the Hb(W) values are the sum of contributions of two different groups. In TABLE 7.-ENTHALPIC EFFECTS OF HYDROPHOBIC HYDRATION OF AMIDES ACCOUNTING FOR NH CONTRIBUTIONS solute ' corrections ' - Hb(W)/kJ mol-l FA - is( U) NMF - @(U) AA - id( U) NBA - as( U) BA - $3(U) DMF - DMA - 3.7 10.9 16.1 8.0 19.3 22.6 14.9 NMF the effect is due to a CH and a methyl group. Adoption of a contribution from CH as -3.7 kJ mol-l leaves a value of -7.2 kJ mol-1 for the methyl group, which compares very well with the Hb(W) value for the methyl group in AA.In the same way the contribution of the n-butyl group in NBA can be estimated at - 14.6 kJ mol-l. The procedures applied above give us the opportunity to calculate Hb(W) and n values for the separate alkyl groups of all compounds used in this study. For example, subtracting the 6 values of DMF from the contributions of the CH group [equal to 6(F)-@(U)], gives the sum of the contribution of two methyl groups. The resulting 6 values can be analysed by means of eqn (2). For the substituted acetamides DMA and NBA we have eliminated the contribution of the C-substituted methyl group by substraction of the values of 6(AA). In this way we obtained values for Hb(W) and n of methyl groups in various compounds and of ethyl, n-propyl and n-butyl groups in one solute.They are given in table 8, together with A, the mean deviation of the fit to eqn (2). Table 8 also indicates the different subtraction procedures used. From table 8 it is clear that Hb(W) values of methyl groups in ureas and amides are comparable. In our view any differences are mainly caused by deficiencies in the calculation procedures. Table 8 also shows that the values of n increase with the size of the alkyl group. The same trend has been observed with other compounds.' We have shown earlier3 that absolute values of n do not bear any physical significance and that only trends are important. The change in Hb(W) with the number of C atoms in the alkyl groups is shown in fig. 5, in which we used an average value for the methyl groups.Fig. 5 includes earlier results for alcohols' and amines.2 Although the amount of data from ureas and amides is limited, a distinct parallelism between the three classes of compounds can be observed. The value of Hb(W) for an n-propyl group in the curve for the amides seems to deviate. It originates from BA. This is not surprising since BA has a backbone of four carbon atoms. A similar high value of Hb(W) can be observed for AA.A. ROUW AND G . SOMSEN 3407 TABLE 8.-ENTHALPIC EFFECTS OF HYDROPHOBIC HYDRATION AND THE NUMBER OF HYDRATION SITES PER ALKYL GROUP FOR ALKYL GROUPS IN VARIOUS COMPOUNDS alkyl group subtraction procedure -Hb(W)/kJ mol-' n A/kJ mol-I Me Me Me Me Me Me Me Me Et n-Pr n-Bu d(MU) -;d(U) t[d( 1 1 -DMU) - @(U)] &[d( 173-DMU) -+d(U)] G(NMF) - [d(F) - td(U)] +(d(DMF)-[d(F)-&d(U)]) 6(AA) - @(U) d(EU) -id(U) d(BA) - &S(U) G(NBA) - [d(AA) -$d(U)] $d(TMU) t@(DMA) - [ W A ) - mJ>l> 7.1 6.5 6.6 6.2 8.1 7.0 5.6 8.0 10.7 14.9 15.1 4.0 0.07 3.5 0.08 4.3 0.08 5.1 0.02 3.4 0.06 3.0 0.08 4.0 0.03 4.4 0.09 5.5 0.09 6.6 0.14 8.7 0.16 I I I I I I I 1 2 3 4 5 6 number of carbonatoms FIG.5.-Variation of Hb(W) with the number of C atoms in alcohols (x), amines (0) and ureas + amides (0). Recently Spencer et aZ.14 discussed the enthalpic effects caused by the transfer of amides from non-aqueous solvents to water. From an analysis of the transfer process they calculated the difference in the enthalpic effect caused by the structuring of the water molecules around DMF and DMA, and found a value of 6.8 kJ mol-l.From our results we calculate a difference of only 3.2 kJ mol-l. In our view this discrepancy may be caused by the fact that the value of Spencer et al. was obtained as a remainder, after calculation of several other contributions to the transfer enthalpies. This remainder is exclusively assigned to the water-structuring effect, without accounting for differences between the transfer enthalpies of the CON groups in DMF and DMA. We are grateful to Dr M. Booij for helpful discussions.3408 SOLVATION + HYDRATION OF UREAS + AMIDES A. C. Rouw and G. Somsen, J. Chem. Thermodyn., 1981, 13, 67. A. C. Rouw and G. Somsen, J. Solution Chem., 1981, 10, 533. W. J. M. Heuvelsland, M. Bloemendal, C. de Visser and G. Somsen, J. Phys. Chem., 1980,84,2391. I. M. Klotz and J. S. Franzen, J. Am. Chem. Soc., 1962,84, 3461. D. W. James, R. F. Armishaw and R. L. Frost, J. Phys. Chem., 1976, 80, 1346. 0. D. Bonner, J. M. Bednarek and R. K. Arisman, J. Am. Chem. SOC., 1977,99, 2898. R. Kummel and H. Hesse, 2. Phys. Chem. (Leipzig), 1981, 262, 705. ' P. R. Philip, G. Perron and J. E. Desnoyers, Can. J. Chem., 1974, 52, 1709. ' E. G. Finer, F. Franks and M. J. Tait, J. Am. Chem. Soc., 1972,94, 4424. lo S. Subramanian, T. S. Sarma, D. Balasubramanian and J. C. Ahluwalia, J. Phys. Chem., 1971, 75, l1 E. R. Stimson and E. E. Schrier, J. Chem. Eng. Data, 1974, 19, 354. l2 J. Konicek and I. Wadso, Acta Chem. Scand., 1971, 25, 1541. l3 R. Skold, J. Suurkuusk and I. Wadso, J. Chem. Thermodyn., 1976, 8, 1075. l4 J. N. Spencer, S. K. Berger, C. R. Powell, B. D. Henning, G. S. Furman, W. M. Loffredo, E. M. Rydberg, R. A. Neubert, C. E. Shoop and D. N. Blauch, J. Phys. Chem., 1981, 85, 1236. l5 C. de Visser, H. J. M. Griinbauer and G. Somsen, 2. Phys. Chem (N.F.), 1975, 97, 69. l6 W. J. M. Heuvelsland, C. de Visser and G. Somsen, J. Phys. Chem., 1978, 82, 29. 815. J. C. Verhoef and E. Barendrecht, Anal. Chim. Acta, 1977,94, 395. I. Wadso, Sci. Tools, 1966, 13, 33. l8 W. J. M. Heuvelsland, C. de Visser and G. Somsen. J. Phys. Chem., 1979, 82, 29. 2o W. J. M. Heuvelsland, C. de Visser, G. Somsen, A. LoSurdo and W-Y. Wen, J. Solution Chem., 1979, 8, 21. (PAPER 2/553)

ISSN:0300-9599    

DOI:10.1039/F19827803397

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 3409-3415 Osmotic and Activity Coefficients of Aqueous Solutions of Sodium Tetraphenylboron at 0 and 25 O C BY THELMA M. HERRINGTON* AND CHRISTINE MARY TAYLOR Department of Chemistry, The University of Reading, Whiteknights, Reading RG6 2AD Received 1st April, 1982 The osmotic coefficient of the solvent and the activity coefficient of the solute for aqueous solutions of sodium tetraphenylboron have been determined at 0 and 25 OC. The solute-solute interaction left after subtracting the Debye-Huckel contribution for the electrostatic forces indicates ion pairing. The large tetraphenylborate ion shows singular behaviour in its physico-chemical properties compared with other quaternary ions of similar size (e.g. 44A~+, 44P+ and Bu4N+).Skinner and FUOSS~ found that the limiting conductances of B& and Am,BuN+ are nearly equal but that there is an unexplained difference in behaviour at dilutions less than infinite; the former ion has a positive coefficient in the conductance function while the latter has a negative one. The n.m.r. spectra of B4h in water of Coetze and Sharpe2 indicate specific solute-solvent effects; the chemical shift of the water hydroxy proton showed that the orientation of water molecules by B4; is very different from that of I$~As+ and 44P+. Krishnan and Friedman3 found that in both water and methanol the solvation enthalpy of 44A~+ is close to that of four benzene molecules, whereas the solvation enthalpy of B4; shows a large extra negative contribution. Jolicoeur and Philip4 found that the temperature dependence of the heat capacity of the tetraphenylborate ion was considerably greater than that of d4P+ and Bu4N+.However, Millero5 determined the apparent molar volumes of aqueous sodium tetraphenylboron solutions between 0 and 60 OC and found that the deviations of the apparent molar volumes from the limiting law were large and negative; this behaviour is similar to that of tetra-alkylammonium salts. There has long been speculation as to whether the unusual properties of the tetraphenylborate anion could be the result of ion pairing. Covington and Tait6 analysed their conductivity data at 25 OC in terms of two models, one of which indicated extensive ion pairing. However, further spectroscopic studies7 failed to confirm this.Millero5 has suggested that because of the similarities of the deviations of the apparent molar volumes of NaB4, and Bu4NBr aqueous solutions, similar types of ion-ion interactions are responsible in both systems. He points out that as it is usually considered that some type of cationsation interaction is responsible for the large negative deviations of the Bu4NX solutions, by analogy the negative deviations of the NaB4, solutions might be attributed to anion-anion interactions. This work is an attempt to explain why the large B4; anion confers similar behaviour to the large R4N+ cation for some physico-chemical properties but not others, and to what extent considerations of ion pairing are relevant. It was decided to obtain equilibrium thermodynamic data for aqueous solutions of sodium tetra- phenylboron at 0 and 25 OC and to analyse them in terms of an electrolyte plus a non-electrolyte contribution.Sodium tetraphenylboron is soluble up to 0.15 molal at 34093410 OSMOTIC A N D ACTIVITY COEFFICIENTS OF AQUEOUS NaB4, 25 OC. In these dilute solutions, if the Debye-Huckel electrostatic contribution is subtracted from In y, then the remainder is linear in the molality as for a non-electrolyte. The recent theory of Pitzer8 was also used to calculate the electrostatic contribution. The non-electrolyte contribution can then be analysed using rigorous statistical- mechanical t h e ~ r i e s . ~ EXPERIMENTAL FREEZING-POINT MEASUREMENTS The freezing-point depressions were measured with a model 3R osmometer (Advanced Instruments Inc.).The solution is supercooled and the formation of ice crystals initiated by a vibrating probe. Heat of fusion is liberated and the sample temperature rises to its freezing point. To obtain reproducible results, care must be taken to exclude dust, centre the thermistor probe and keep the temperature of the freezing bath constant. The freezing point could be read to 0.1 x K; at least ten readings were taken at each molality and the standard deviation varied from 0.2 to 0.5 at the highest molality. The instrument was calibrated with sodium chloride solutions over the same freezing-point range and the calibration was checked before and after each run. The overall accuracy in the freezing point depression of the sodium tetraphenylboron solutions is & lop4 K.VAPOUR-PRESSURE MEASUREMENTS If a drop of solution and a drop of solvent are suspended side-by-side in aconstant-temperature enclosure saturated with solvent vapour, a differential mass transfer occurs between the two drops and the vapour phase. This transfer cmses a temperature difference between the two drops which is proportional to the vapour-pressure lowering.1o. l1 The drops are formed on two thermistor beads mounted in opposite arms of a resistance bridge, so that the temperature difference is measured as a change in resistance. The calibration constant for the instrument was determined using sodium chloride solutions of known molality against water as a reference. For the sodium tetraphenylboron solutions a solution of sodium chloride of known molality and closely matching vapour pressure was used, so that only a small correction to an accurately known osmotic coefficient needs to be determined.MATERIALS The sodium tetraphenylboron (B.D.H. Ltd) was purified by the method of Skinner and FUOSS.~ All solutions were prepared using once distilled but previously deionized water; it had a conductance of 1 x lop6 Rpl cm-l. Solutions were made up by weight and buoyancy corrections applied to give a precision of kO.1 mg. RESULTS OSMOTIC COEFFICIENT AT 0°C The freezing-point apparatus was calibrated using aqueous sodium chloride solutions of accurately known molality. The freezing-point data of Scatchard and Prentiss12 were smoothed using Debye-Hiickel theory13 [eqn (Al) and (A3) of the Appendix]: these smoothed values of the osmotic coefficient were used for calibration (table 1).Values for the osmotic coefficient of aqueous sodium tetraphenylboron solutions were calculated at 0 O C using the equation 4 = 812 x 1.8606 m. These values for the osmotic coefficient were smoothed by both the Debye-Hiickel and Pitzer methods (see the Appendix) to give the smoothed values for the osmotic * A glossary of symbols is given in the Appendix.T. M. HERRINGTON AND C. M. TAYLOR 341 1 TABLE 1 .-SMOOTHED VALUES OF THE OSMOTIC COEFFICIENT FOR AQUEOUS SODIUM CHLORIDE SOLUTIONS AT 0 AND 25°C rnlmol kg-* 4 (0 "C) 4 (25 "C) 0.0100 0.0300 0.0500 0.0700 0.0900 0.1 100 0.1300 0.1500 0.969 0.952 0.944 0.938 0.935 0.932 0.930 0.929 0.968 0.95 1 0.943 0.938 0.935 0.932 0.93 1 0.930 TABLE 2.-sMOOTHED VALUES OF THE OSMOTIC COEFFICIENT OF THE SOLVENT AND ACTIVITY COEFFICIENT OF THE SOLUTE FOR AQUEOUS SOLUTIONS OF SODIUM TETRAPHENYLBORON AT 0 AND 25 O C 0 O C 25 "C m/mol kg-l In Y Y 4 In Y Y 0.0100 0.0300 0.0500 0.0700 0.0900 0.1100 0.1300 0.1500 0.975 0.967 0.967 0.969 0.974 0.979 0.985 0.992 -0.093 -0.136 -0.155 -0.164 -0.168 -0.168 -0.166 -0.161 0.912 0.873 0.856 0.848 0.845 0.845 0.847 0.851 0.918 - 0.228 0.796 0.888 - 0.306 0.736 0.862 -0.374 0.688 0.646 0.813 - 0.496 0.610 0.791 -0.550 0.577 0.769 - 0.603 0.547 0.836 - 0.437 coefficient of the solvent and the activity coefficient of the solute, y , recorded in table 2.(The smoothed values of 4 and y were the same by both methods to the number of figures given.) The values of cu found were cu(Debye-Huckel) = 1.03 kg mol-l and w(Pitzer) = 1.18 kg mol-l.OSMOTIC COEFFICIENTS AT 25 O C The Mechrolab 30 1 vapour-pressure osmometer was calibrated using the same sodium chloride solutions as for the freezing-point apparatus. The osmotic coefficients for aqueous sodium chloride solutions were required at molalities of < 0.15 mol kg-l: the activity coefficient of the solute is accurately known at these low molalities from electromotive-force measurements. The data of Brown and McInnes14 and of Longworth15 were smoothed using eqn (A 2) and (A 4) of the Appendix. The values of the osmotic coefficient used for calibration are given in table 1. The experimental values for the osmotic coefficient of aqueous sodium tetraphenyl- boron solutions were smoothed in the same manner as the results at 0 O C .Smoothed values of the osmotic coefficient of the solvent and the activity coefficient of the solute are given in table 2. It was found that co(Debye-Hiickel) = - 1.84 kg mol-l and w(Pitzer) = - 1.69 kg mol-l.3412 OSMOTIC AND ACTIVITY COEFFICIENTS OF AQUEOUS NaB4, DISCUSSION From theoretical considerations16 the Gibbs energy of a solution of mole ratio of solute to solvent r i may be written G / N l k T = ,u:/kT+rn,u~/kT-m+inln Eii+$A22m2+&22m3+ .... (2) According to the theory of McMillan and Mayer,17 for a solution of a solute in a solvent the osmotic pressure, n, is given by n/kT= n+B,*nn2+B,*,,n3+ . . . . ( 3 ) Hill1* has shown that the coefficients A,, etc. may be related to the coefficients I?,*, etc. For example (4) A,, V: = 2B:: - VF + by1 where B,*,O = - b:,, the solute-solute cluster integral.Now b:l, the solute-solvent cluster integral, is related to the partial molecular volume of solute at infinite dilution by16 For aqueous solutions of sodium tetraphenylboron let us denote the non-electrolyte contribution to the solute activity coefficient by In y*, then from eqn (2) in dilute by, = -@+kTu. ( 5 ) solution 2In y* = A , , i ~ + B , , , i ? i ~ f .... Thus from eqn (A 2) and (6) o = A,, M1/2, and from eqn (4) and ( 5 ) NB,*,O = A,, V,O/2+ V?-RTu/2. (7) Values for the solute-solute virial coefficient B,*,O were calculated from the values obtained for o. The critical compilation of Bradley and Pitzerlg was used for the compressibility of water. For water Vy is 18.02 cm3 mo1-1 at 0 OC and 18.07 cm3 mol-l at 25 OC.,O The apparent molar volume of sodium tetraphenylboron at O°C is 267.2 cm3 mol-l and at 25 O C 276.4 cm3 ~ 0 I - l .~ Values for NB,*,O are given in table 3 using both the Debye-Huckel and Pitzer values for o. Different values are obtained for the non-electrolyte contribution depending on which theory is used for the electrostatic forces. As the values are of the same order of magnitude only the Debye-Huckel method was used for comparison with other electrolytes. Data for aqueous solutions of sodium chloride and aqueous solutions of sucrose in table 3 are from ref. (9) and (21). The values for NB,*,O for sodium tetraphenylboron are very different from those of sucrose and sodium chloride; at 0 OC NB,*,O is large and positive and at 25 O C large and negative for sodium tetraphenylboron, whereas it is not markedly temperature dependent for either sucrose or sodium chloride.BZ: can be considered to be composed of an attractive and a repulsive contribution from the intermolecular forces, thus B,*,O = S+@* ( 8 ) where S is the repulsive and (DA the attractive contribution. If a hard-sphere model is assumed for the repulsive contribution, then the temperature dependence of B:: is that of the attractive. contribution. In table 3 values of N(a@*/aT),, obtained from heat-of-dilution datag for sodium tetraphenylboron are compared with those for sodium chloride and sucrose. If the repulsive contribution to NB,*,O is constant then the negative value at 25 OC implies that the attractive contribution increases markedly with increasing temperature, and this is supported by the large negative value forT.M. HERRINGTON AND C. M. TAYLOR 341 3 TABLE 3.-sOLUTE-SOLUTE VIRIAL COEFFICIENT B::, SOLUTE-SOLVENT VIRIAL COEFFICIENT B::, AND TEMPERATURE DEPENDENCE OF THE SOLUTE-SOLUTE INTERACTION COEFFICIENT NB,*,0/cm3 mol-l NB,*,O - N(i3DA/aT), NB,*p species T/K Debye-Hiickel Pitzer /cm3 mol-l /cm3 mol-l K-' /cm3 mol-l NaB4, NaB44 NaCl NaCP sucroseb sucrosea Bu,NCl B u N C 1 273.15 298.15 - 2 7 3 . 1 5 2 9 8 . 1 5 2 7 3 . 1 5 298.15 273.15 2 9 8 . 1 5 - - 266 1 3 0 0 1 4 5 0 1570 - 1420 72.51" 275 1 2 232 307 - 2.20 1 6 3 3 0 - 205 - - 28 5 0.56 2 1 0 - 287 262 7.86 2 9 3 - - - - - - - - - - - - - a Data from ref. (9); * Data from ref. (21). N(a@*/a T ) p .Thus on this model ion pairing in sodium tetraphenylboron solutions should be appreciable at 25 O C . The small negative value of N(a@*/aT), for sucrose is consistent with the slight decrease in NB,*,O with increasing temperature and shows that the attractive force between two sucrose molecules is not highly temperature dependent. From our analysis for sodium chloride w(0 "C) = 0.22 kg mol-1 and w(25 "C) = 0.29 kg mol-l: V p values for sodium chloride were taken from Dunn22 and the values for NB,*,O are given in table 3. The increase in NB,*,O with temperature for sodium chloride is consistent with the small positive value for N(&DA/dT),, indicating that the solute-solute attraction decreases with increasing temperature. It has been suggested5 that the behaviour of the apparent molar volume of sodium tetraphenylboron solutions indicates that this electrolyte behaves like the tetra- alkylammonium chlorides in its ion-ion interactions and like sodium chloride in its ion-water interactions.From eqn ( 5 ) values for the solute-solvent virial coefficient B;r,O = --byl can be calculated. Data are given for tetrabutylammonium chloride in table 3 . 9 3 23 Sodium tetraphenylboron and tetrabutylammonium chloride have similar values for NB,*P, which implies similar solute-solvent interaction, if we assume comparable hard-sphere molar volumes, but very different values for their solute-solute interactions. Sodium chloride has considerably smaller values for NBT,O consistent with a smaller hard-sphere repulsive force. For all three electrolytes the solute-solvent attractive forces decrease with increasing temperature.Friedman and co-w~rkers~~ have analysed the thermodynamic properties of aqueous solutions of tetra- alkylammonium halides using ion-pair correlation functions. The models that best fit the data show that + - ion pairing is of considerably greater importance than + + or - - ion pairing. However, all types of ion pairing contribute to thermodynamic and conductance data and our analysis makes no attempt to distinguish between them. CONCLUSION Analysis of the non-electrostatic interaction between solute molecules in aqueous solutions of sodium tetraphenylboron using the McMillan-Mayer interaction coeffi- cients for the solute-solute cluster integral indicates considerable ion pairing at 25 O C , decreasing rapidly with decreasing temperature.The trend of the osmotic-coefficient data obtained at 0 and 25 *C is in agreement with our measurements of the heat of3414 OSMOTIC AND ACTIVITY COEFFICIENTS OF AQUEOUS NaB4, dilution at 25 0C.9 In the same temperature range the solute-solvent interaction decreases slightly. The solute-solute behaviour of the large tetraphenylboron anion is quite different from that of the large tetrabutylammonium cation, whereas their solute-solvent interactions are similar. The similar negative deviations of the apparent molar volumes of NaB4, and Bu,NBr aqueous solutions come in fact from the non-electrolyte term linear in the molality and reflect the pressure derivative (i3A22/i3p)T. From eqn (4) and (5)’ A,, reflects both solute-solute and solute-solvent interaction : it seems plausible that the pressure dependence of the similar solute-solvent interaction dominates this term.There has been considerable speculation6* as to whether or not ion pairing occurs in sodium tetraphenylboron aqueous solutions at 25 OC. A recent coilductance study by Schiavo el al.25 concludes that ‘ion pairing’ is not negligible for sodium tetraphenylboron solutions at 25 OC. Ion-pairing constants have also been found for a range of other solvents. APPENDIX DATA SMOOTHING Standard values for the osmotic coefficient of solvent tjSt: and the activity coefficient of solute, ySt, are defined using electrolyte theory and the experimental values of the osmotic or activity coefficient are fitted to the equations tj-tjSt = om12 (A 1) (A 2) In y-ln ySt = om.Two expressions for the standard values were used. (a) Based on Debye-Huckel theory:I3 1 - @t = abP3 m-l[ 1 + bm: - (1 + bm;)-l- 2 In (1 + bmj)] In ySt = -am:( 1 +bm$)-l. (A 3) (A 4) At 0 OC, a = 1.132 kgi mol-a and at 25 OC, a = 1.173 kg’ mol-i; b = 1 .O kgi mol-4 (b) Based on the theory of Pitzer? 1 - qPt = A, mi( 1 + bm;)-’ In ySt = -A, mi( 1 + brni)-l - 2A, In (1 + bmi)/b (A 5 ) (A 6) where A , = a13 and b = 1.2 k& mol-4. G k iTi m MI n N GLOSSARY OF SYMBOLS Pitzer electrostatic parameter solute-solute virial coefficient (Bzf = - b;,) solute-solvent virial coefficient (B1*p = - byl) cluster integral for two molecules of solute in pure solvent cluster integral for one molecule of solute and one of solvent in pure solvent Gibbs energy Bol tzmann’s constant mole ratio of solute to solvent (N2/N1) molaiity of solute in mol kg-l molar mass of solvent in kg mol-1 number density of solute Avogadro’s constant number of molecules of solvent number of molecules of soluteT.M. HERRINGTON AND C. M. TAYLOR 3415 gas constant repulsive contribution to the cluster integral absolute temperature partial molecular volume of solvent molecular volume of pure solvent molar volume of pure solvent partial molecular volume of solute at infinite dilution partial molar volume of solute at infinite dilution De bye-Hiickel electrostatic parameter activity coefficient of solute on the molality scale non-electrolyte contribution to the solute activity coefficient osmotic pressure osmotic coefficient attractive contribution to the configuration integral isothermal compressibility of solvent molecular chemical potential of pure solvent molecular chemical potential of the solute at infinite dilution non-electrolyte solute-solute interaction parameter freezing point depression J.F. Skinner and R. M. Fuoss, J. Phys. Chem., 1964, 68, 1882. * J. F. Coetze and W. R. Sharpe, J . Phys. Chem., 1971, 75, 3141. C. V. Krishnan and H. L. Friedman, J. Phys. Chem., 1971, 75, 3606. C. Jolicoeur and P. R. Philip, J. Solution Chem., 1975, 4, 3. F. J. Millero, J. Chem. Eng. Data, 1970, 15, 562. A. K. Covington and M. J. Tait, Electrochim. Acta, 1967, 12, 113. A. K. Covington and M. J. Tait, Electrochim. Acta, 1967, 12, 123. T. M. Herrington and E. L. Mole, J. Chem. SOC., Faraday Trans. I , 1982, 78, 2095. * K. S. Pitzer, J. Phys. Chem., 1973, 77, 268. lo E. J. Baldes, J. Sci. Instrum., 1934, 1 1 , 223. l 1 A. P. Brady, H. Huffand and J. W. McBain, J. Phys. Chem., 1951, 55, 304. l 2 G. Scatchard and S. S. Prentiss, J. Am. Chem. SOC., 1933, 55, 4355. l 3 P. Debye and E. Hiickel, Phys. Z., 1923, 24, 185. l 4 A. S. Brown and D. A. McInnes, J . Am. Chem. SOC., 1935, 57, 1356. l6 J. E. Garrod and T. M. Herrington, J. Phys. Chem., 1969, 73, 1877. l 7 W. G. McMillan and J. E. Mayer, J. Chem. Phys., 1945, 13, 276. I s T. L. Hill, J . Am. Chem. SOC., 1957, 79, 4885. l 9 D. J. Bradley and K. S. Pitzer, J. Phys. Chem., 1979, 83, 1599. *O G. S. Kell and E. Whalley, Philos. Trans. R . SOC. London, Ser. A, 1965, 258, 565. 22 L. A. Dunn, Trans. Faraday SOC., 1968, 64, 2951. 23 W-Y. Wen, in Structure and Transport Processes in Water and Aqueous Solutions, ed. R. A. Horne 24 P. S. Ramanathan, C. V. Krishnan and H. L. Friedman, J . Solution Chem., 1972, 1 , 237. 25 S. Schiavo, R. M. Fuoss and G. Marrosu, J. Solution Chem., 1980, 9, 563. L. G. Longsworth, J. Am. Chem. SOC., 1932, 54, 2741. T. M. Herrington and C. P. Meunier, J. Chem. SOC., Furaday Trans. 1, 1982, 78, 225. (Wiley, New York, 1970), chap. 15. (PAPER 2/557)

ISSN:0300-9599    

DOI:10.1039/F19827803409

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 3417-3429 Electrochemistry of Pol yace t ylene, ( CH)z Characteristics of Polyacetylene Cathodes BY KEIICHI KANETO, MACRAE MAXFIELD, DAVID P. NAIRNS AND ALAN G. MACDIARMID* Department of Chemistry, University of Pennsylvania, Philadelphia, Pennsylvania 19104, U.S.A. AND ALAN J. HEEGER Department of Physics, University of Pennsylvania, Philadelphia, Pennsylvania 19104, U.S.A. Received 1st April, 1982 The relationship of cell potential to degree of oxidation, coulombic and energy efficiencies, constant-current discharge characteristics, energy density and maximum power density of a partly oxidized polyacetylene, [CH(ClO,),] 0, < 0.07), cathode in a cell of the type [CH(ClO,),],ILiClO,ILi are discussed. Coulombic efficiencies ranging from 100 to 86 % and energy efficiencies ranging from 8 1 to 68 %during a chargedischarge cycle are found at oxidation levels ranging from 1.54 to 6.0%.Energy densities of ca. 255 W h kg-’ (based on the weights of the electroactive materials involved in the discharge process) are obtained for 7.0% oxidized polyacetylene cathodes under constant-current discharge conditions. Maximum power densities of ca. 30 kW kg-’ are observed. We have shown previously that polyacetylene, (CH),, may be controllably oxidized or reduced electrochemically with the incorporation of a variety of counter anions or cations, and that as oxidation or reduction proceeds its conductivity increases through the semiconducting to the metallic regime to give ultimately a series of ‘organic metals’.l We have also demonstrated that ca.0.1 mm thick films of partly oxidized or reduced (CH), may be used as the cathode and anode materials, respectively, in lightweight, rechargeable storage b a t t e r i e ~ . ~ ? ~ Polyacetylene is the first example of a covalent polymer whose conductivity can be increased, by chemical or electrochemical oxidation or reduction, into the metallic regime. Increases in electrical conductivity of over twelve orders of magnitude have been 0bserved.l The partly oxidized material exists as a stabilized polycarbonium ion, (CHY+),, in combination with the corresponding number of monovalent counter anions, A-, such that the overall composition is [CHu+A;],. Similarly, the partly reduced material exists as a polycarbanion, (CHY-)$, in combination with the corresponding number of monovalent counter cations, M+, such that the overall composition is [M,+CHY-],.Extensive studies of the partly oxidized material have shown that a semiconductor-metal transition occurs at oxidation concentrations > 7 mol% with a highly conducting transitional regime extending from ca. 0.1 mol% up to 7 mol%. In the concentration range > 1 mol%, the conductivity increases upon further oxidation at a relatively slow rate up to a value of ca. lo3 Q-l cm-l at room temperature. We summarize here selected important electrochemical properties of the more extensively studied partly oxidized (CH),, such as voltage against degree of oxidation curves, coulombic and energy efficiencies, constant-current discharge characteristics, energy density and maximum power density, which demonstrate its potential for use as a cathode material in rechargeable batteries.34173418 ELECTROCHEMISTRY OF POLYACETYLENE EXPERIMENTAL All investigations were performed using cells constructed by sandwiching a separator between a (CH), film electrode and a lithium electrode, squeezing the assembly into rectangular glass tubing (3 mm x 10 mm) adding the electrolyte together with ca. 200 mg of Woelm B Super 1 basic alumina and sealing the tube under vacuum across the protruding electrode leads. Using this method of construction the minimum amount of electrolyte was employed, air was excluded and the electrodes were held firmly in place, ca. 0.5 mm apart from each other. The (CH), electrodes were constructed of a single sheet of cis-rich (CH), (ca.85% cis isomer; 1-2.5 cm2; 0.1 mm thick; 3-10 mg). On oxidation to the levels involved in the present study, spontaneous isomerization to the trans isomer occurs. Cis film was used in construction of the cells since it is more flexible and is therefore more easily folded into the current collectors, as described below. One current collector was constructed by spot welding a 10 mil* platinum wire lead to a piece of 52 mesh platinum gauze of similar area to the (CH), film. The other current collector was constructed by spot welding a 10 mil nickel wire lead to a piece of 189 mesh nickel grid (Delker Corp.) of similar area to the (CH), film. This nickel grid was then folded tightly around a piece of 0.13 mm lithium ribbon (Alfa Ventron Corp.) whose area was approximately half that of the nickel grid.This composite was then placed on top of the separator [of similar size to the (CH), film] which was placed on top of the (CH), film which was itself placed on top of the platinum gauze. The whole assembly was then folded tightly in half so that the Li-Ni electrode was at the inside centre and the platinum grid was on the outside of the whole assembly. The whole assembly was then inserted into the rectangular glass tubing. Hydrophobic polypropylene (Celanese Celgard K-442) was used as the separator material in the cell potential and the energy density studies; freshly kiln dried (600 "C) glass filter-paper (Reeve Angel 934AH) was employed in the power density studies. The electrolyte solution was 1 .O mol dm-3 LiClO, (Alfa Ventron Corp.) in propylene carbonate.The solvent, electrolyte and lithium were purified and handled as described previ~usly.~ Charging of cells was performed with a Princeton Applied Research potentiostat/galvanostat model 173 by oxidizing the (CH), electrode either at constant current [ca. 0.1 mA per mg of (CH),], or at a series of stepped up constant potentials. Discharge studies were carried out in a similar manner, either at a constant-current discharge or at a series of stepped down constant applied potentials. Coulombs passed during charge or discharge cycles were recorded using a PAR coulometer, model 179. Voltages, resistances and currents were measured with a Keithley 177 microvolt digital multimeter.Voltage-charge, voltage-time and current-time graphs were recorded with a Houston Instruments Corp. Omnigraphics XY recorder. All voltages are given with respect to a lithium electrode. It must be stressed that the electrochemical characteristics of (CH), electrodes are extremely sensitive to the method of cell construction, presence of impurities (especially oxygen), relative ratio of electrolyte to (CH),, charging conditions etc. The results given in this report were obtained by following the described procedures exactly. RESULTS AND DISCUSSION CELL POTENTIAL (v) AND DEGREE OF OXIDATION (@ The positive and negative terminals of a d.c. power supply were attached to the (CH), and Li electrodes, respectively, immersed in a solution of 1 .O mol dm-3 LiClO, in propylene carbonate, and the cell was charged at a constant applied voltage [initial current ca.0.1 mA per mg of (CH),]. When the charging current fell to ca. 0.0 14 mA per mg of (CH),, the applied voltage was increased by steps of ca. 0.02 to 0.3 V. This step charging procedure was repeated at increasingly higher applied voltages until the desired final voltage was reached. The overall charging process in which the (CH), is oxidized and the Li+ is reduced is represented by (CH), + xyLi+(ClO,)- -+ [CHy+(ClO,);], + xyLi. (1) * 1 mil = 0.001 in = 2.54 x 10-5 m.KANETO, MAXFIELD, NAIRNS, MACDIARMID A N D HEEGER 3419 The charging process was interrupted at periodic intervals and V (o.c., immediate), the open-circuit voltage, was measured within 2-3 s of the interruption of the charging process.An open-circuit voltage is actually a voltage measured while the cell is being discharged at a very low rate, in these experiments at ca. 0.3 PA. Hence, during the few seconds needed to perform an open-circuit voltage measurement, negligible discharge of the cell occurs. The percentage oxidation of the polyacetylene electrode at each V(0.c.) measurement was calculated from the amount of charge passed and from the weight of (CH), (3.5 mg) employed. The total time taken to oxidize the (CH), to 6.00/,, i.e. to composition [CH(C104)o.06],, was ca. 3 h. The curve showing the relationship between V(o.c., immediate) and the percentage oxidation is given in fig. 1. 4.0 3.5 5 c - v L. 3.0 1 I I I I I A similar procedure was carried out with a different but similarly constructed cell [using 8.1 mg of (CH),., except that the cell was permitted to stand for 5 min after the charging process was interrupted before the open-circuit voltage, V (o.c., 5 min) was measured. The time taken for the whole experiment was 7 h.The curve showing the relationship between V(o.c., 5 min) and the percentage oxidation is given in fig. 1. In another study, using the same cell employed in the V(o.c., immediate) experiment described above, the charging process was interrupted at periodic intervals and the cell was permitted to stand for ca. 24 h to permit diffusion of the (C10,)- ions from the exterior to the interior of the ca. 200 I$ diameter (CH), fibrils. The open-circuit voltage, V (o.c., 24 h), was then measured. The cell was then discharged [reverse reaction to that given by eqn (I)] to 2.50 V, a voltage characteristic of parent, neutral (CH),, during 16 h using the method described below.The percentage oxidation of the polyacetylene cathode at each point was then calculated from the amount of charge (in coulombs) released during discharge. The curve showing the relationship between V(o.c., 24 h) and the percentage oxidation of the (CH), is given in fig. 1. The data for this experiment were collected over a period of three weeks. The empirical equation representing the V(o.c., 24 h) against percentage oxidation curve is given by the relationship (2) V(o.c., 24 h) = 3.43 + 0.14 In y3420 ELECTROCHEMISTRY OF POLYACETYLENE where y is the percentage oxidation. The excellent agreement between the curve (for values of y > 0.05%) given by this equation and the experimental points is shown in fig.1. In all the studies described above, significant oxidation occurred only at an applied potential > ca. 3.1 V. After the onset of oxidation, the V(0.c.) rises rapidly with increasing oxidation up to oxidation levels of ca. 1 % and then increases more slowly. Thus the electrochemical potential of a (CH), cathode can be varied over a relatively large range depending on the degree of oxidation. From fig. 1 it is apparent that for a given level of oxidation the observed open-circuit voltage decreases in the order V(o.c., immediate) > V(o.c., 5 min) > V(0.c. 24 h). This is interpreted as indicating that diffusion of the (C10,)- ions from the surface of the fibrils to the interior is relatively slow.Since the open-circuit voltage will be determined by that portion of the (CH), in contact with the electrolyte, the open-circuit voltages will decrease with time on standing as the extent of oxidation on the exterior of the fibrils decreases and the extent of oxidation in the interior increases by this diffusion process. The important role of diffusion is confirmed by the observation that the open-circuit voltage increases with time after rapid partial discharge. In this case, the (C10,)- ions and the associated positive charges on the polymer diffuse from the more highly oxidized interior of fibrils to the (ClO,)--depleted fibril exterior. Increases of almost 1 V were observed after cells were permitted to stand for 20 min after a relatively rapid discharge to 2.50 V (see below).Studies are in progress to determine the time taken to attain diffusion equilibrium within a fibril. For a 200 A (CH), fibril the diffusion time constant, 7,* is ca. 2 days;, for a 50 A fibril it is ca. 8 ha5 These values appear to vary significantly with the value of the applied voltage and extent of oxidation. The relationship between the open-circuit voltage and the degree of oxidation of the (CH), is believed to be given most accurately by the V(o.c., 24 h) curve because: (i) equilibration of C10, ions will be more complete than in the experiments from which the V(o.c., immediate) and V(o.c., 5 min) curves are generated and (ii) the amount of charge released during discharge is believed to give a more reliable measure of the extent of oxidation than the amount injected during charging, since at higher levels of oxidation some charge is lost during a charge-discharge cycle.Possible reasons for the loss of charge are discussed below. COULOMBIC EFFICIENCY A N D ENERGY EFFICIENCY A cell containing 3.5 mg of (CH), was charged at a constant applied potential of a magnitude such that the initial current was ca. 0.1 mA per mg of (CH),. When the charging current fell to ca. 10% of the initial value, the applied voltage was increased to give again a charging current of ca. 0.1 mA per mg of (CH),. When the charging current fell to ca. 10% of the initial value, the applied potential was again increased and the process was repeated as many times as necessary in a given experiment. This procedure results in a charging rate corresponding to ca.2% oxidation of the (CH), film per hour. The charging current was periodically discontinued and ~(o.c., immediate) was measured. The amount of charge involved in charging [Q(in)] to this value of V(o.c., immediate) was recorded and used to calculate the apparent extent of oxidation of the (CH)x at this value of V(o.c., immediate). The term ' apparent extent * 7 is related to the time taken to reach diffusion equilibrium in the following way: consider a solid cylinder surrounded by a constant concentration, C,,, of a substance, A, which is diffusing into the material of the cylinder. At any given time the average concentration of A in the cylinder is gven by c.At time = 2, c/Co = 0.74; at time = 27, C/Co = 0.91; at time = 32, c/Co = 0.97; at time = 45, C/Co = 0.99. Hence, diffusion equilrbrium is virtually established after time = 4 ~ . ~4.0 TABLE 1 .-COULOMBIC AND ENERGY EFFICIENCIES OF A [CH(ClO,),],lLiClO,(P.C.)ILi CELL* - - 2 coulombic efficiency energy efficiency oxidation [Q(out, total)/Q(in, total)] x 100 [E(out, total)/E(in, total)] x 100 (%> (%I (%) 4.5 I I I I 1 I -,B -0 -0-0-0-4 B - A0’ 3.5 .+’ I - v a 0 - 1 I I 1 I I 2.0 * 1.54 2.01U 2.17 2.51 4.0 6.0 100.0 99.2 100.1 95.8 89.5 85.7 80.8 (88.7)c 79.7b (89.5) 81.5 (89.2) 78.2 (87.7) 72.8 (78.8) 68.2 (74.8) a The discharge curves for 2.01,2.17 and 2.5 1 % oxidation were similar; for ease in examining the curves in fig. 2, the discharge curves for 2.01 and 2.51% oxidation are, therefore, not shown. Slower rates of charging and discharging results in higher energy efficiencies, e.g.85.8% at 2.0% oxidation. Values in parentheses are obtained from the area under V(0.c. immediate) curves during charging and discharging. * Nofe added inproqf: Studies performed since this manuscript was submitted for publication show that higher coulombic efficiencies at higher levels of oxidation are obtained consistently if the (CH), is oxidized at the rate of 1 % per hour at constant current charging conditions and is then discharged first at a constant applied current of ca. 0.2 mA mg-I to 2.50 V and then at a constant applied voltage of 2.50 V for 16 h. Coulombic efficiencies in excess of 92% have been obtained for oxidation levels of 8.0%, and of 87% for oxidation levels of 10.0%. These high coulombic efficiencies have also been confirmed in an independent laboratory (D.J. Frydrych and G. C. Farrington, personal communication; J . Electrochem. SOC. 1982, to be published).3422 ELECTROCHEMISTRY OF POLYACETYLENE of oxidation’ is employed since the system is not at diffusion equilibrium and also because the results of this study show that at higher levels of oxidation, complete charge recovery is not obtained. Further points were obtained by restarting the charging process with a higher applied voltage and repeating the above operations until the maximum value of ~ ( o . c . , immediate) desired in a given cycle was obtained. The total number of coulombs [Q (in total)] used in the charging operation in this cycle was then recorded and the degree of oxidation of the (CH), was calculated.An example of the stepped-up charging potential (solid-line steps) and the corresponding ~ ( o . c . , immediate) curve is shown in fig. 2 for the 6.0% oxidation experiment. An identical procedure was used in the studies at lower levels of oxidation shown in fig. 2 and summarized in table 1. The constant applied charging potentials used in these studies were significantly greater than the near diffusion equilibrium potentials, ~(o.c., 24 h), for a given degree of oxidation. For example, the final constant applied potential step for the 6% oxidized film given in fig. 2 was 3.95 V, implying a local surface oxidation level considerably in excess of 6%. This 3.95 V value may be compared with the V(o.c., 24 h) value [calculated from eqn (2)] of 3.68 V for 6.0% oxidized film.A similar procedure was used when discharging. The cell was discharged during 1 h to ca. 3 V in a series of constant applied potential steps, as illustrated for the 6.0% oxidation cycle in fig. 2. It was then finally discharged at a fixed potential of 2.50 V during ca. 16 h and the total amount of charge (in coulombs) liberated, Q(out, total) was recorded. The extent of apparent oxidation of the (CH), at any selected ~(o.c., immediate) value in the discharge cycle was calculated from Q(in, total) - Q(out), where Q(out) was the number of coulombs released on discharging to a given V(o.c., immediate) value in a given discharge cycle. From fig. 2 it can be seen that hysteresis is present in the charge-discharge curves.From the data obtained in these experiments, the coulombic efficiency of a charge-discharge cycle, Q(out, total)/Q(in, total), may be calculated, where Q(out, total) is the total number of coulombs released in the discharge process to 2.50 V in a given charge-discharge cycle. The energy efficiency of a given charge-discharge cycle, E(out, total)/&, total), may also be calculated. The total energy (amount of charge x voltage) expended in a given charge cycle, E(in, total), is given by the area under the applied stepped-up voltage curve (shown in fig. 2 for the 6.0% oxidized cycle). Similarly, the total energy released, E(out, total), in the discharge portion of a cycle is given by the area under the corresponding applied stepped-down voltage curve.Energy efficiencies calculated from the areas under the respective stepped-up charge and stepped-down discharge curves using the applied V(charge) and V(dis- charge) potentials at each step are given in table I for a number of different (CH), oxidation levels. The energy efficiencies listed will be dependent on the arbitrarily chosen values of the stepped-up charge and stepped-down discharge voltages used in a given experiment. Energy efficiencies obtained under different charge and discharge conditions will, therefore, be higher or lower than the values listed in table 1. If the areas under the V(o.c., immediate) charge and discharge curves are used, then the energy efficiencies given in parentheses in table I are obtained.These refer to the percentage of the stored energy in a given charge cycle which could be released if the cell were discharged so that the discharge potential followed the V(o.c., immediate) discharge curve all the way down to 2.50 V. These data suggest that cells containing (CH), electrodes may have excellent potential for use in storage batteries. Slower rates of charging and discharging will result in higher energy efficiencies due to reduction in 12R loss. As shown above, higher charging rates result in a greater degree of oxidation of the (CH), on the exterior of the fibrils since the dissipationKANETO, MAXFIELD, NAIRNS, MACDIARMID A N D HEEGER 3423 of charge to the interior of a fibril is limited by diffusion. Thus at higher charging rates a higher voltage is required to produce a given average level of oxidation.Conversely, at higher discharge rates the discharge potential will be lower because of depletion of charge on the exterior of the fibrils. Hence, the energy efficiencies under near equilibrium conditions should be greater than the values given in table 1. For example, charging to 2.0%, [V(charge) = 3.69 V], at an average current of ca. 0.015 mA per mg of (CH), and discharging in a series of decreasing constant potential steps to 2.50 V, in each of which the discharge current was similar to the above, gave an energy efficiency of 85.8%. The value corresponding to that given in parentheses in table 1 is 89.7%. The average charging and discharging currents in this experiment were ca.3-4 times smaller than the average currents used in obtaining the data in table 1 . Studies are in progress to determine whether the loss of charge over a complete charge-discharge cycle is real or only apparent. The loss of charge could be caused by chemical reaction of the electrolyte or by reaction of impurities in the electrolyte with [CHV+(ClO,);],. Preliminary studies4 suggest that the charging potential, which will determine the local extent of oxidation on the fibril surface, should not exceed ca. 3.75 V if chemical reaction of the film with the propylene carbonate/LiClO, electrolyte is to be avoided. Higher values may be possible with other electrolyte systems. The 100% coulombic efficiency studies given in table 1 were carried out at charging potentials < 3.75 V.Films consisting of smaller diameter fibrils, having smaller z values, can be oxidized more rapidly without exceeding this critical charging potential. However, at least a significant portion of the apparent loss of coulombs may be due simply to the fact that an insufficient amount of time had been allowed for complete diffusion of (C104)- ions from the interior of the fibrils to the surface during discharge. Factors which increase the amount of time elapsing between the initiation of a charge cycle and termination of a charge cycle will favour more complete diffusion of (C10,)- ions into the interior of a fibril. Such factors would therefore imply that a correspondingly longer time would be needed for (ClO,)--ion diffusion to the exterior of a fibril during complete discharge.If the coulombic loss is real and is caused by reaction of the (CHY+), species with the electrolyte or with traces of impurities, then the real percentage oxidation values in fig. 2 will be slightly less than shown, to the extent indicated by the coulombic efficiencies at various oxidation levels as given in table I . In situ optical studies5 during the oxidation-reduction cycle have recently been carried out. The optical cell utilized a thin (CH), film (ca. 2000 A thick by 2 cm2 in area; fibril diameter ca. 50 A) in 2 cm3 of electrolyte so that the electrolyte to polymer ratio was ca. lo5. Under these conditions, even trace impurities in the electrolyte might lead to harmful side reactions. Nevertheless, these studies show that the (CH), is not degraded either chemically or electrochemically during cycling at quasi-equilibrium diffusion conditions to a maximum charging voltage of 3.73 V (7.8% oxidation).Hence, the (CH), electrodes can be stable in the electrochemical environment under the experimental conditions employed in these optical studies. Other related investigations show, quite conclusively, that traces of air and unknown impurities in the system cause the voltage of a cell to fall sharply on standing. The stability is increased after several charge-discharge cycles, possibly because initially the (CHy+), species acts as a scavenger for impurities. The V(0.c.) of a carefully prepared, previously cycled cell after 1 day was 3.67 V. This fell to 3.61 V in one week and during the next two weeks fell at a rate of 0.003 V per day to a value of 3.56 V.Detailed long-term stability studies are in progress.3424 ELECTROCHEMISTRY OF POLYACETYLENE ENERGY DENSITY The change in voltage of a cell as it is discharged at a constant current is an important characteristic of the cell. Ideally, it should show very little change in voltage until it is almost completely discharged when the voltage would then drop sharply. Fig. 3 shows the excellent characteristics of a cell constructed from 3.3 mg of (CH), oxidized to 7.0% when discharged at constant currents of 0.1, 0.55 and 1.0 mA to 2.50 V immediately after charging. Greater weights of film will naturally provide a longer 'plateau' region for any given discharge current. These discharge rates in A per kg of [CH(ClO,),~,,], are 19.5, 107 and 195, respectively. The total energy released upon discharge (charge x voltage) is given by the area under each curve.To obtain the energy density (W h kg-l) the mass of the electroactive material involved was calculated using the weight of [CH(ClO,),~,,], employed and the weight of Li consumed in the discharge reaction [the reverse reaction to that given in eqn (l)]. The energy density values obtained are: 0.1 mA, 258 W h kg-l; 0.55 mA, 255 W h kg-'; 1.0 mA, 254 W h kg-l when the cell is discharged to 2.50 V, the voltage characteristic of parent, neutral (CH),. The corresponding energy density values obtained from the curves in fig. 3 on discharge to 3.0 V, the approximate potential'at which the discharge voltage begins to drop rapidly, are 2 19,2 18 and 21 6 W h kg-l, respectively. These values may be compared with the previously reported energy density of 176 W h kg-l obtained in a 3 min short-circuit discharge of a cell using 6.0% oxidized (CH),.3 Empirical rules may be applied to obtain a rough estimate of the expected energy density of a packaged battery, including the weight of electrolyte, solvent and casing from the experimental energy density of 218 W h kg-l.A reduction factor of seven gives a reasonably conservative estimate and results in a value of 31 W h kg-' for the completely packaged battery. This is approximately the same value as that found for t/min FIG. 3.-Relationship between the discharge voltage, Y, and discharge time, t, during a constant current discharge at 0.1 mA ( x ), 0.55 mA (O), and 1.0 mA (0) of a (CH),Il mol dmP3 LiClO, in propylene carbonatelLi cell.KANETO, MAXFIELD, NAIRNS, MACDIARMID A N D HEEGER 3425 the average lead/acid automobile battery.It should be noted that the above energy- density value is for 7% oxidized (CH),. Higher levels of oxidation will result in larger energy densities. The coulombic efficiencies for the 0.1, 0.55 and 1 .O mA discharges shown in fig. 3 are given in table 2. After completion of a discharge given in fig. 3, the cell was permitted to rest for 20 min, during which time diffusion of (C10,)- from the interior to the exterior of a fibril continued; an increase in V(0.c.) was observed. An additional constant potential discharge at 2.90 V for 16 h resulted in the release of further charge, as shown in table 2.Note that the 86.7% coulombic efficiency value in table 2 for a 7.0% oxidized film is in complete agreement with the 85.7% value in table 1 for 6.0% oxidized film obtained in a different, although related, type of experiment. The release of additional charge is caused by diffusion of (C10,)- ions within a fibril, not to diffusion of Li+ ions in the electrolyte to the surface of the fibrils. This is obviously the case since if the total quantity of charge released were dictated by the rate of diffusion of Li+ ions within the electrolyte to the fibril surface, then more charge should have been liberated in the slower 0.1 mA discharge than in the faster 1.0 mA discharge. As shown in table 2, the reverse is the case for this 7.07; oxidized film, more charge being released in the faster discharge.The mechanism for this unexpected observation is presently being studied. It is believed to be related to the charging times used in each of the three discharge studies. These were 12, 7 and 5.5 h for the 0.1, 0.55 and 1.0 mA studies, respectively, suggesting that the smaller the time allowed for diffusion of (C10,)- ions into the (CH), during charging the greater is the coulombic recovery at these relatively high discharge rates. TABLE 2.-cOULOMBIC EFFICIENCIES, [Q(out, total)/Q(in, total)] X 100, FOR SELECTED CONSTANT CURRENT DISCHARGES OF A [CH(C104),~,,],~LiC104(P.C.)(Li CELL coulombic efficiency increase of total coulombic after discharge cycle V(0.c.) during efficiency after a 16 h I(discharge) shown in fig.3 20 min rest constant voltage discharge /mA to 2 . 5 0 V (%) period to 2.90 V (%) 0.1 0.55 1 .o 74.0 2.50 + 3.42 80.3 79.3 2.50 + 3.40 84.4 86.7 2.50 + 3.36 87.2 POWER DENSITY The power delivered by a cell is given by the product of the voltage and current during a discharge. Power is dependent on a number of factors related to the nature of the electroactive materials, the packaging of the electrodes and the internal resistance of the cell. In the simplest case, when the resistance of the external load, R,, through which the cell is discharged is equal to the internal resistance, Ri, the power delivered by the cell is at its maximum. The maximum power, PmaX, was measured in three ways by one or the other of the following relationships:w P h, rn TABLE 3.-TYPICAL MAXIMUM POWER DENSITIES OF [CH(ClO,),],lLiClO,lLi CELLS M M r CI el 0 0 weight 56 is 28f5 eqn (4) ; 1 4.2 f 0.1 4.9 & 0.2 2.1 3.52 22.4 22.5 f 0.2 1.8 f 0.2 - 2 4.2 4.9 2.1 3.54 21.3 22.0 86f 1 33+2 eqn ( 5 ) 4 37f7 eqn (4) g 3 7.0 8.2 2.0 3.52 13.0 11.0 1.83 - 4 7.0 8.4 2.5 3.55 11.0 10.5 1.80 36+7 eqn (4) 5 12.0 17.4 & 0.2 5.5 3.66 8.0 10.0 2.08 f 0.04 208 f 5 252 l b eqn (3) g of electro- maximuma weight of active oxidation ~(o.c., 15 h) power density, experiment (CH),/mg material/mg (%) /v R i m RlIQ V m P Im/mA Pm,,/kW kg-l method m - - 4 a The calculated error in the P,,, values arises from the uncertainties in measuring the mass of the (CH),, the discharge current and the discharge voltage, and in adjusting the variable resistor.The errors associated with each measurement are given for experiment 1. They are of similar magnitude m in the other experiments. The small error in experiment 5 reflects the greater accuracy in reading the discharge voltage, which fell less rapidly due to the greater extent of oxidation of the (CH),. r MKANETO, MAXFIELD, NAIRNS, MACDIARMID A N D HEEGER 3427 where V, and I , are the voltage and current, respectively, measured at the very beginning of a discharge cycle under matched load conditions. The maximum power density was calculated in kW kg-l. The internal resistance of the cell was first measured by charging the cell as described above at a rate of ca. 2% oxidation of the (CH), per hour. Charging was terminated at a level of oxidation required for a given experiment listed in table 3.After standing for ca. 5 h to permit partial equilibration of (C10,)- ions within the fibrils, V(o.c., 5 h) was measured and the cell was discharged for ca. 1 s through an ammeter and coulometer in series whose total resistance, R,, (ca. 2 a), had been determined previously by an ohmmeter. The discharge current, I, was measured and was used to calculate Ri, values of which are listed in table 3, by means of the relationship: The coulombs lost in this measurement were also recorded. The cell was recharged using the same number of coulombs lost in the above operation. To measure P,,,, the cell was permitted to stand for ca. 15 h and V(o.c., 15 h) was recorded. It was then discharged through one of the three circuits shown in fig.4 in which the total external load, R,, was made equal to Ri. The adjustable external load includes the resistance of any other instruments such as coulometers etc. V(O.C., 5 h) = (Ri + &,)I. ( 6 ) R I 1 3 R Q R P which might have been used in a circuit in certain instances. The circuits are based on the use of eqn (3), (4) and (5), respectively. Values of V, were measured ca. 2 s after a discharge was commenced. During the initial 2 s period, the voltage fell very rapidly from the V(0.c.) value and then decreased more slowly. Values of I, were taken to be the peak discharge current which rises from zero to a maximum within 1-2 s then declines steadily. The results of five typical experiments using different cells containing differing amounts of (CH), doped to different extents are given in table 3.With the exception of experiment 5, it can be seen that V, z V(o.c., 15 h)/2, indicating that maximum power discharge conditions are operative. In experiment 5 , the Ri and R, values are not exactly matched and hence V , is greater than V(o.c., 15 h)/2. In this case, the product of I , and R,, 2.08 V, is identical to the experimentally determined value of VIn. The results of a number of experiments show that the maximum power density values at oxidation levels > 2% do not vary greatly either with the percentage3428 ELECTROCHEMISTRY OF POLYACETYLENE oxidation or with the total mass of film employed. Maximum power densities in the range of 30 kW kg-l are obtained (see table 3). These are based on the weight of [CH(ClO,),], employed and the weight of lithium which would be oxidized in the discharge reaction if it went to completion. Application of the empirical conversion factor of seven discussed previously gives maximum power densities of ca.4-5 kW kg-l. This may be compared with that of ca. 0.2 kW kg-l for a typical lead/acid automobile battery . The extremely high maximum power density values are probably related primarily to the fact that the effective surface area of ca. 4 mg of (CH), film (1 cm2 x 0.01 cm) is ca. 2.5 x lo3 cm2. Since the fibrils composing the film are only ca. 200 A in diameter, no portion of the (CH), can ever be more than ca. 100 from the electrolyte. This, coupled with the large surface area of the film, permits good accessibility of the (CH), to the (ClO,)- ions with consequent extremely high current densities during discharge (and charge) processes.The high power density is also related to the fact that because of the high electronic conductivity of partly oxidized (CH),, the removal of positive charge can readily occur and is limited primarily by the rate at which the C10, ion can diffuse from the interior to the exterior of a fibril during discharge. Any partial removal of positive charge (i.e. addition of electronic charge) from the interior of a [CHY+(ClO,)-], fibril will greatly increase the electrostatic repulsion between the remaining (ClO,)- ions and will hence increase their rate of diffusion to the fibril exterior. Recent preliminary studies4 show that the diffusion constant for the diffusion of charge into (CH), fibrils depends on the potential applied to the (CH), electrode.This may, in part, be the origin of the higher coulombic efficiencies at higher rates of discharge discussed above. In general, thinner films result in larger power densities. This is probably related to the observation that in thin (CH), films, formed during shorter polymerization times, the fibril diameter is smaller than that found for fibrils in films produced during longer polymerization times.' The larger surface area per unit weight of the thin fibrils contributes greatly to the production of larger discharge currents and correspondingly larger power densities. Since the power delivered at any given time during discharge is the product of the discharge voltage and discharge current, the power against time relationships for constant discharge currents of 0.1,0.55 and 1 .O mA may be obtained simply from the data given in fig. 3. The power against time curves obviously have exactly the same shape as the voltage against time curves. The average power densities are 70, 354 and 591 W kg-l for the 0.1, 0.55 and 1.0 mA discharges, respectively. The present results suggest that electrochemical studies not only of (CH), but also of other conducting organic polymers represent an extensive area for further research, not only of fundamental scientific interest but of possible potential technological value. We thank Dr Shahab Etemad for suggesting the logarithmic relationship between the potential of partly oxidized (CH), and its degree of oxidation and Mr James H. Kaufman for many helpful discussions. This study was supported by the U.S. Department of Energy, Contract no. DE-AC02-8 1 -ERlOS32. A. G. MacDiarmid and A. J. Heeger, Synth. Met., 1979/80, 1, 101 ; A. J. Heeger and A. G. MacDiarmid, in The Physics and Chemistry of Low Dimensional Solids, ed. L. Alcacer (D. Reidel, Dordrecht, Holland, 1979), p. 353. P. J. Nigrey, A. G. MacDiarmid and A. J. Heeger, J . Chem. Soc., Chem. Commun., 1979, 594; P . J. Nigrey, D. Macinnes Jr, D. P. Nairns, A. G. MacDiarmid and A. J . Heeger, in Conductive Polymers, ed. R. B. Seymour (Plenum Press, New York, 1981), p. 227.KANETO, MAXFIELD, N A I R N S , MACDIARMID A N D HEEGER 3429 P. J. Nigrey, D. MacInnes Jr, D. P. Nairns, A. G. MacDiarmid and A. J. Heeger, J. Electrochem. Soc,, 1981, 128, 1651. K. Kaneto, A. G. MacDiarmid and A. J. Heeger, to be published. A. Feldblum, J. H. Kaufman, S. Etemad, A. J. Heeger, T-C. Chung and A. G. MacDiarmid, Phys. Rev. B, 1982, in press. W. Jost, Dzflusion (Academic Press, New York, 1960), p. 45. ' M. Aldissi, Ph.D. Thesis (Universite des Sciences et Techniques du Languedoc, Montpellier, 1981), pp. 48-60; M. Aldissi and F. Schue (U.S.T.L., Place E. Bataillon, 34060 Montpellier Cedex, France), unpublished observations. (PAPER 2/558)

ISSN:0300-9599    

DOI:10.1039/F19827803417

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

Reviews of Books The Thermodynamics of Soil Solutions. By G. SPOSITO (Clarendon Press : Oxford University As the title implies, this book is directed primarily towards the application of thermodynamics to the highly complicated aqueous solutions which are in equilibrium with various types of soil. However, unlike some other books on specialised applications of thermodynamics, there is much in this volume which will be of interest and value to physical chemists interested in similar, but less complicated, systems and solutions. Throughout the book the author not only stresses the applications of thermodynamics to real systems, but also emphasises the many limitations to its use. The first two chapters review some fundamental concepts of chemical thermodynamics and care has been taken to explain these and their consequences in greater detail than is common in many textbooks of physical chemistry.These chapters would be useful reading for anyone wanting to refresh his or her ideas about the nature of some basic thermodynamic assumptions and definitions, such as the nature of the standard state. However, the author has adopted the unusual approach of defining activities in such a way as to make them dimensionless, with the unfortunate consequence that the usually dimensionless activity coefficients must be assigned the dimension [concentration]-l or [molalityl-l. One error in this section (p. 25): the implication that the conditions for equilibrium when temperature and pressure are held constant is dU = 0 arises from a slightly lax definition of chemical potential a few pages earlier.The remainder of the book deals with detailed applications, chap. 3, 4 and 5 being devoted to solubility equilibria, electrochemical equilibria and ion exchange, respectively. In each case the author shows how the many complex interactions in soil-solution systems can be broken down into the contributing parts. Both soil chemists and classical thermodynamicists should find these sections of considerable interest. Most of the final third of the book is devoted to a consideration of ion exchange, one chapter being devoted to the thermodynamic background and a second to molecular considerations. The final chapter considers the problem from a soil physicist’s point of view and discusses the properties of the water rather than of the solutes and solid phases present.At no stage does the author shirk from presenting his work in a properly quantitative way and these last chapters particularly are mathematically demanding. Certainly they are such as to be useful to only the most dedicated undergraduate and much of the book is really directed towards postgraduate students and research workers. However, one excellent feature of the book will make some of it useful reading for undergraduates. At the end of each chapter the author includes both detailed notes and references to specific points raised in the main text and also a comprehensive, annotated further-reading list for more general points. Many references are given to recent work and the whole presentation seems to be free of typographic errors. The book concludes with three indexes including a comprehensive one on the symbols used throughout the text .This book should certainly be purchased for the library of any institution or individual interested in the physical chemistry of soils. Its price and relatively narrow area of specialization mean that it is unlikely to be widely purchased by undergraduates, but could be useful in showing how many concepts from classical thermodynamics and electrochemistry can be extended to apply to the highly complex interactions encountered in soil-water systems. A. D. PETHYBRIDGE Press, Oxford, 1981). Pp. xii + 223. Price E24. Received 1st April, 1982 343 13432 REVIEWS OF BOOKS Mass Spectrometry. Principles and Applications. By I. HOWE, D.E. WILLIAMS and R. D. BOWEN (McGraw-Hill, New York, 2nd edn, 1981). Pp. xii+276. Price E24.95. This is a revised version of the book written by Drs Howe and Williams in 1972 which was a first-class introduction to mass spectrometry. Such has been the development of the subject over the last ten years that it was now out-of-date in some respects. This revised edition remedies this and although the authors state that the book is not intended for experts, I am sure that many practising mass spectroscopists as well as newcomers to the field will find much to interest them. After a short opening chapter on instrumentation, there follow four chapters on the underlying principles of ionisation processes and the kinetics and energetics of unimolecular and bimolecular reactions of ions.The reader is given a good qualitative description of the role played by kinetics in determining the observed mass spectrum and the chapter on energetics of unimolecular reactions gives a concise account of the present position in a branch of the subject in which the authors have made a number of important contributions. Personally, I should have liked to have seen a slightly more extended account of bimolecular reactions but there is sufficient here to give the reader an indication of what may be found in the literature. The remainder of the book is devoted to applications, major chapters being devoted to the interpretation of electron impact spectra and the analysis of mixtures. In the last decade, on-line data systems have become increasingly available and a short chapter on computer techniques provides a useful introduction to what is now possible with such a system. A selection of analytical applications and a brief description of isotopic labelling methods lead to an unusual final chapter in which the questions to be considered in choosing a mass spectrometer system are discussed. Overall, this book can be recommended as a readable and well-balanced introduction to the subject. K. R. JENNINGS Received 22nd March, 1982

ISSN:0300-9599    

DOI:10.1039/F19827803431

出版商:RSC    

年代:1982

数据来源: RSC