ФИЗИЧЕСКАЯ ХИМИЯ РАСТВОРОВ

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1 ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ, 29, том 83, 1, с УДК ФИЗИЧЕСКАЯ ХИМИЯ РАСТВОРОВ APPARENT MOLAR VOLUME, VISCOSITY AND ADIABATIC COMPRESSIBILITY OF SOME MINERAL SULPHATES IN AQUEOUS BINARY MIXTURES OF FORMAMIDE AT , AND K 29 M. N. Roy, R. Chanda and B. K. Sarkar Department of Chemistry, University of North Bengal, Darjeeling-73413, India Received 28 July, 28 Abstract The densities and viscosities of several sulphates, viz., ammonium sulphate, sodium sulphate, potassium sulphate, magnesium sulphate, zinc sulphate and cadmium sulphate in aqueous binary mixtures of formamide (FA) have been determined at , and K and at atmospheric pressure. The ultrasonic speeds of the electrolytic solutions have also been measured at K. Apparent molar volumes (φ V ), viscosity B-coefficients and adiabatic compressibilities (K S ) of these electrolytic solutions were calculated from the experimental densitiy, viscosity and acoustic data. The density and viscosity data were evaluated by using Masson s and Jones Dole equation respectively; the derived parameters have been analyzed in terms of ion ion and ion solvent interactions. The structure making/breaking capacities of the electrolytes have been inferred from the sign of 2 2 ( φ V/ T ) P. The results showed that all the electrolytes act as structuremakers in these media. Also the compressibility data indicated electrostriction of the solvent molecules around the cations. The activation parameters of viscous flow were also determined and discussed by the application of transition state theory. INTRODUCTION Studies on density, viscosity and acoustic properties of ionic solutions are of great help in characterising the structure and properties of solutions. Various types of interactions exist between the ions in the solutions and of these, ion ion and ion solvent interactions are of current interest. These interactions [1 4] help in better understanding of the nature of solute and solvent, i.e. whether the added solute modifies or distorts the structure of the solvent. Recently we have undertaken a comprehensive program to study the solvation and association behaviour of some electrolytes [4 8] in different aqueous and non-aqueous solvent media from the measurement of transport and thermodynamic properties. Formamide (FA) is choosen for the study as it is a simplest amide that contains a peptide linkage, the fundamental building block of proteins [9]. It along with some of its derivatives serves as good solvents for many organic and inorganic compounds and is also used as a plasticizer. Hence, in the present study, we reported density, viscosity and ultrasonic speed of ammonium sulphate, sodium sulphate, potassium sulphate, magnesium sulphate, zinc sulphate, and cadmium sulphate in 1, 2 and 3 mass % formamide + water mixture at , and K and the derived parameters were discussed in terms of ion ion and ion solvent interactions. EXPERIMENTAL Extrapure grade Formamide procured from S.d.fine-Chem. Limited, India was purified by the standard procedures [1]. After purification its purity was ascertained by GLC and also by comparing experimental values of densities and viscosities with their literature values [11 14]. The physical properties of pure and aqueous binary mixtures of formamide at different temperatures are listed in Table 1. All the sulphate salts (Anal R grade) were procured from E. Merck (India) Ltd. They were used after drying over P 2 O 5 in a desiccator for few hours. The reagents were always placed in the desiccator over P 2 O 5 to keep them in dry atmosphere. Freshly distilled conductivity water (specific conductance <1 6 Ω 1 cm 1 ) was used for preparing aqueous mixtures of formamide. The densities (ρ) were measured with an Ostwald Sprengel type pycnometer having a bulb volume of 25 cm 3 and an internal diameter of the capillary of about.1 cm. The pycnometer was calibrated at , and K with doubly distilled water and benzene. The pycnometer with the test solution was equilibrated in a water-bath maintained at ±.1 K of the desired temperatures. The pycnometer was then removed from the thermostatic bath, properly dried and weighed. Adequate precautions were taken to avoid evaporation losses during the time of actual measurements. Averages of triplicate measurements were taken into account. Mass measurements accurate 1922

2 APPARENT MOLAR VOLUME 1923 Table 1. Physical properties of pure formamide (FA) and different mass% of formamide (FA) + H 2 O mixtures at different temperatures T, K ρ 1 3, kg m 3 η, mpa s u, m s 1 Present work Literature Present work Literature Present work Literature 1 mass% of formamide + water mass% of formamide + water mass% of formamide + water Pure formamide [11] [11] [12] [13] [13] [14] [14] to ±.1 mg were made on a digital electronic analytical balance (Mettler, AG 285, Switzerland). The viscosities (η) were measured by means of a suspended Ubbelobde type viscometer [15], calibrated at , and K with triply-distilled water and purified methanol. The flow times were accurate to ±.1 s. The ultrasonic speeds of sound for the solutions were determined by a multi-frequency ultrasonic interferometer (Mittal Enterprises, New Delhi, India) working at 2 MHz, calibrated with purified water and benzene at K. The temperature stability was maintained within ±.1 K by circulating thermostatic water around the cell with a circulating pump. The binary aqueous solutions of formamide as well as the solutions of sulphates were made by mass and conversion of molality into molarity was done [16] using experimental density values. The uncertainties in the density, viscosity and speed of sound measurements were estimated to be ±.1 g cm 3, ±.3 mpa s and ±.2 m s 1 respectively. The details of the methods and techniques have been described elsewhere [17 23]. RESULTS AND DISSCUSSIONS The apparent molar volumes (φ V ) have been calculated using the following equation [24]: φ V = M --- ρ 1( ρ ρ ) , cρ (1) where c is the molar concentration of the solution, M is the molecular weight of the solute; ρ and ρ are the densities of the solvent and solution respectively. The values of limiting apparent molar volumes ( φ V ) and experimental slopes ( S V * ) at different temperatures have been obtained by using a least-square treatment to the plots of φ V against c 1/2 using the Masson s equation [25]: φ V = φ V + S V *c 1/2. (2) The values of limiting apparent molar volumes ( ) and experimental slopes ( S V * ) are listed in Table 2. φ V The intercept of equation [2] is a measure of ion solvent interactions [26]. Table 2 shows that the values of φ V are generally small for most of the electrolytes and decrease with an increase in temperature; this indicates the presence of weak ion solvent interaction and such interaction decreases as the temperature of the mixtures increases. This may be due to large electrostriction at higher temperatures. The decrease in φ V may also be attributed to the decrease in solvation. The slope S V * of equation [2] may be attributed to the measure of ion ion interactions [27 29]. A perusal of Table 2 reveals that the S V * values for most of the examined sulphates, are large positive at different temperatures and also increase in magnitude with a rise in φ V ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

3 1924 ROY et al. Table 2. Limiting apparent molar volumes ( φ V ) and experimental slopes ( S V ) of different sulphate salts in different aqueous binary mixtures of formamide (FA) at different temperatures (standard errors are given in parenthesis) FA, mass % φ 1 6, m 3 mol 1 1 6, m 3 l 1/2 mol 3/2 V S V K K K K K K (NH 4 ) 2 SO (±.27) (±.64) (±.5) (±.5) (±.2) (±.3) (±.4) (±.56) (±.9) (±.4) (±.3) (±.5) (±.2) (±.62) (±.67) (±.3) (±.1) (±.1) Na 2 SO (±.9) (±.6) (±.72) (±.1) (±.2) (±.3) (±.7) (±.66) (±.93) (±.1) (±.1) (±.4) (±.59) (±.83) (±.39) (±.5) (±.1) (±.1) K 2 SO (±.44) (±.36) (±.19) (±.5) (±.2) (±.2) (±.27) (±.75) (±.91) (±.2) (±.2) (±.3) (±.88) (±.89) (±.73) (±.4) (±.2) (±.1) MgSO (±.68) (±.7) (±.94) (±.9) (±.3) (±.1) (±.88) (±.95) (±.6) (±.5) (±.2) (±.3) (±.1) (±.59) (±.53) (±.3) (±.4) (±.2) ZnSO (±.4) (±.72) (±.45) (±.2) (±.3) (±.3) (±.68) (±.17) (±.35) (±.2) (±.5) (±.2) (±.71) (±.9) (±.14) (±.2) (±.4) (±.6) CdSO (±.9) (±.34) (±.62) (±.1) (±.1) (±.1) (±.59) (±.75) (±.18) (±.1) (±.3) (±.2) (±.67) (±.52) (±.91) (±.9) (±.3) (±.3) ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

4 APPARENT MOLAR VOLUME 1925 temperature. This suggests that ion ion interactions further strengthen with a rise in temperature. This may be attributed to the desolvation of the electrolytes [3] in these systems due to more thermal agitation at higher temperatures. Further, it is found that S V * increases on going from 1 to 3 mass % of formamide in the mixed solvents; thereby indicating increasing trend of ion ion interactions. This in turn supports the behaviour of φ V indicating decreased ion solvent interac- tions as the content of formamide in the solutions increases. Partial molar volumes ( φ V ) at infinite dilution were fitted to a second order polynomials [31] in terms of absolute temperature (T): φ V = a + a 1 T+ a 2 T 2. (3) The values of coefficients a, a 1 and a 2 for the studied sulphates at different temperatures in different binary mixtures of formamide are recorded in Table 3. The limiting apparent molar expansibilities ( φ E ) at infinite dilution can be obtained [31] by differentiating equation [3] with respect to temperature, φ E = ( φ V / T) P = a 1 + 2a 2 T. (4) The values of along with the sign of ( 2 φ V / T 2 ) P for the studied electrolytes at different temperatures in different binary mixtures of formamide are recorded in Table 4. It is evident from Table 4 that the values of φ E increases as the temperature increases for all electrolytic solutions. This may be attributed to the presence of caging or packing effect [32, 33]. Hepler [34] has developed a technique of examining the sign of ( 2 φ V / T 2 ) P for various solutes in terms of long range structure making or breaking capacities of the solutes in mixed solvent systems using the following thermodynamic expression [35]: ( C P / P) T = T( 2 φ V / T 2 ) P. (5) According to this, the left hand side of the equation (5) should be positive for all structure breaking solutes and therefore, structure breaking solutes posses negative value of ( 2 φ V / T 2 ) P. On the other hand, positive value of ( 2 φ V / T 2 ) P should be associated with structure making solutes. In the present study, Table 4 revealed that the values of ( 2 φ V / T 2 ) P for all the electrolytic solutions are positive, thereby suggesting that these electrolytes act as structure makers in these solvent mixtures. The viscosity data of all the electrolytic solutions at different temperatures in different aqueous binary mixtures of formamide have been analyzed by Jones Dole [36] equation: φ E Table 3. Values of various coefficients of Eq. (3) for different sulphate salts in 1, 2 and 3 mass% of FA + H 2 O mixtures FA, mass % a 1 6, a 1 1 6, a 2 1 6, m 3 mol 1 m 3 mol 1 K 1 m 3 mol 1 K 2 (NH 4 ) 2 SO Na 2 SO K 2 SO MgSO ZnSO CdSO ( η/η 1)/c 1/2 = A+ Bc 1/2, (6) where η and η are the viscosities of solvent-mixtures and solutions respectively. The values of A and B parameters have been determined from the intercept and slope of linear plots of (η/η 1)/c 1/2 versus c 1/2 and recorded in Table 5. Table 5 shows that the values of A- coefficients, for most of the sulphates under investigation, are positive and increase with the rise in temperature, thereby suggesting strong ion ion interactions [37] and these interactions further strengthen with the increase of temperature. Also it is found that A-coefficient values increase with the amount of formamide in the mixtures, suggesting the predominance of ion ion interactions in higher amount of formamide. A perusal of Table 5 shows that the values of B-coefficient are positive and decreases with a rise in both the temperatures and the amount of formamide in the mixtures. This indicates the decreasing trend in ion solvent interactions and justifies the predominance of ion ion interactions at higher temperatures and increased amount of formamide in the mixtures. ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

5 1926 ROY et al. Table 4. Limiting apparent molar expansibilities ( φ V Also the temperature derivative of B, i.e., db/dt is found to be negative, indicating that these electrolytes act as structure makers in these solvent mixtures [38]. These conclusions are in excellent agreement with our earlier view drawn from magnitude of ( 2 φ V / T 2 ) P illustrated earlier. The viscosity data have also been analyzed on the basis of transition state theory of relative viscosity of the electrolytes as suggested by Feakings et. al [39] using the following equation: RT = µ ( 1B+ V 1 ), (7) V 1 where is the contribution per mole of the solute to the free energy of activation of viscous flow of the solutions; V 1 and are the partial molar volumes of the solvent and solute, respectively; ( ) of the so- ) and ( 2 / T 2 ) P for various sulphate compounds in 1, 2 and 3 mass% of FA + H 2 O mixtures at different temperatures FA, mass% 1 6, m 3 mol 1 K 1 ( 2 φ / T 2 ) P 1 12 V, K K K m 6 mol 2 K 2 φ E (NH 4 ) 2 SO Na 2 SO K 2 SO MgSO ZnSO CdSO φ E lutions was determined from the above relation. The free energy of activation of viscous flow per mole of the pure solvent ( µ 1 ) is given by the relation [39.4]: = = RTln( η V 1 /hn A ), (8) where N A is the Avogadro s number, h is the Planck constant, η is the viscosity of the solvent, R is the gas constant and T is the absolute temperature. The values of the parameters µ 1 and are reported in Table 6. Table 6 shows that µ 1 is practically constant at all the solvent compositions and temperatures, implying that is dependent mainly on the viscosity B-coefficients and ( V 1 ) terms. The values are positive and larger than that of µ 1 at all the experimental temperatures; this suggests that the formation of the transition state is less favourable in the presence of the studied sulphate salts. According to Feakings et al. [39], > µ 1 for electrolytes having positive B-coefficients and indicates a stronger ion solvent interactions, thereby suggesting that the formation of transition state is accompanied by the rupture and distortion of the intermolecular forces in solvent structure [39]. The greater the value of, the greater is the structure making tendency of the electrolyte. Table 6 also shows that the values of decreases with the increase of temperature and this suggest that all the experimental solutes are act as structure makers [41]. The entropy of activation for electrolytic solutions has been calculated using the relation: [39] S 2 = d( )/dt, (9) µ 1 G 1 where has been obtained from the negative slope of the plots of versus T by using a least squares treatment. The activation enthalpy ( ) has been calculated using the relation [39]: = + T S 2. (1) S 2 The values of and are reported in Table 6 and they are found to be positive for all the electrolytic solutions at all the temperatures suggesting that the transition state is associated with bond breaking and decrease in order. Although a detailed mechanism for this cannot be easily advanced, it may be suggested that the slip-plane is some where in the centrosymmetric region [39, 4]. Adiabatic compressibility (K S ) has been calculated at K using the relation [42]: K S = (u 2 ρ) 1, (11) ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

6 APPARENT MOLAR VOLUME 1927 Table 5. Values of A- and B-coefficients of different sulphate salts in different aqueous binary mixtures of formamide (FA) at different temperatures (standard errors are given in parenthesis) FA, mass. % A 1 3, l 1/2 mol 1/2 B 1 3, l mol K K K K K K (NH 4 ) 2 SO (±.5) (±.2) (±.1) (±.16) (±.1) (±.8) (±.2) (±.3) (±.1) (±.8) (±.16) (±.1) (±.2) (±.1) (±.4) (±.1) (±.8) (±.3) Na 2 SO (±.3) (±.2) (±.3) (±.3) (±.4) (±.12) (±.2) (±.2) (±.3) (±.2) (±.3) (±.13) (±.2) (±.4) (±.1) (±.3) (±.1) (±.8) K 2 SO (±.1) (±.3) (±.2) (±.3) (±.9) (±.8) (±.1) (±.2) (±.2) (±.3) (±.6) (±.12) (±.2) (±.1) (±.1) (±.9) (±.5) (±.1) MgSO (±.2) (±.3) (±.3) (±.3) (±.5) (±.7) (±.2) (±.6) (±.3) (±.4) (±.3) (±.8) (±.2) (±.2) (±.3) (±.5) (±.5) (±.8) ZnSO (±.1) (±.1) (±.3) (±.2) (±.2) (±.8) (±.2) (±.2) (±.2) (±.3) (±.4) (±.7) (±.5) (±.2) (±.4) (±.9) (±.4) (±.14) CdSO (±.1) (±.3) (±.4) (±.1) (±.1) (±.2) (±.3) (±.3) (±.2) (±.1) (±.1) (±.1) (±.5) (±.1) (±.3) (±.2) (±.1) (±.2) ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

7 1928 ROY et al. Table 6. Values of,,,, and for (NH 4 ) 2 SO 4, Na 2 SO 4, K 2 SO 4, MgSO 4, ZnSO 4 and CdSO 4 in different aqueous binary mixtures of formamide (FA) at different temperatures V 1 µ 1 Parameter 1 mass % of FA + water 2 mass% of FA + water 3 mass% of FA + water K K K K K K K K K V 1 µ 1 1 6, m 3 mol , kj mol Amminium sulphate 1 6, m 3 mol , kj mol , kj mol , kj mol Sodium sulphate 1 6, m 3 mol , kj mol , kj mol , kj mol Potassium sulphate 1 6, m 3 mol , kj mol , kj mol , kj mol Magnesium sulphate 1 6, m 3 mol , kj mol , kj mol , kj mol Zinc sulphate 1 6, m 3 mol , kj mol , kj mol , kj mol Cadmium sulphate 1 6, m 3 mol , kj mol , kj mol , kj mol ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

8 APPARENT MOLAR VOLUME 1929 Table 7. Molal concentration (m, mol kg 1 ), ultrasonic speed of sound (u, m s 1 ), adiabatic compressibility (K S, Pa 1 ), and apparent molal adiabatic compressibility (φ k, m 3 mol 1 Pa 1 ), of some sulphates in different aqueous binary mixtures of formamide (FA) at K m u K S 1 1 φ k 1 1 m u K S 1 1 φ k 1 1 m u K S 1 1 φ k 1 1 (NH 4 ) 2 SO 4, 1 mass% of FA K 2 SO 4, 1 mass% of FA ZnSO 4, 1 mass% of FA (NH 4 ) 2 SO 4, 2 mass% of FA K 2 SO 4, 2 mass% of FA ZnSO 4, 2 mass% of FA (NH 4 ) 2 SO 4, 3 mass% of FA K 2 SO 4, 3 mass% of FA ZnSO 4, 3 mass% of FA Na 2 SO 4, 1 mass% of FA MgSO 4, 1 mass% of FA CdSO 4, 1 mass% of FA Na 2 SO 4, 2 mass% of FA MgSO 4, 2 mass% of FA CdSO 4, 2 mass% of FA Na 2 SO 4, 3 mass% of FA MgSO 4, 3 mass% of FA CdSO 4, 3 mass% of FA ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

9 193 ROY et al. Table 8. Limiting apparent molar adiabatic compressibility ( φ, m 3 mol 1 Pa 1 ) and experimental slope (, m 3 mol 3/2 К S К kg 1/2 Pa 1 ) of some sulfates at K in mixtures formamide (FA) water FA, mass% φ К S К φ К S К φ К S К (NH 4 ) 2 SO 4 K 2 SO 4 ZnSO (±.58) (±.7) (±.83) (±.4) (±.28) (±.7) (±.25) (±.3) (±.7) (±.1) (±.1) (±.2) (±.1) (±.2) (±.17) (±.1) (±.7) (±.4) Na 2 SO 4 MgSO 4 CdSO (±.46) (±.3) (±.9) (±.1) (±.99) (±.15) (±.11) (±.1) (±.1) (±.1) (±.1) (±.1) (±.14) (±.1) (±.9) (±.7) (±.16) (±.3) where u is the ultrasonic velocity and ρ is the density of the solution. The apparent molal adiabatic compressibility (φ K ) of the solution was computed from the relation [43], K φ S M 1( K K S ρ K S, ρ) = , (12) ρ mρρ where K S, K S, are the adiabatic compressibility of solution and solvent respectively and m is the molal concentration of the solution. In equation (12) molality (m) has been used rather than concentration (c), because molality is independent of temperature and we performed acoustic calculations only at K. The limiting apparent molal adiabatic compressibilities ( φ K ) were obtained by extrapolating the plots of φ K versus m 1/2 of the solution to zero concentration by the computerized least-squares method [44, 45], φ K = φ K + S K *m 1/2, (13) where S K * is the experimental slope. The values of m, u, K S, φ K, φ K and S K * for the studied electrolytes at K in different aqueous binary mixtures of formamide are determined and recorded in Tables 7, 8. A perusal of Table 8 shows that φ K values are negative and S K * values are positive for all the ternary solutions. Since the values of φ K and S K * are measures of ion solvent and ion ion interactions respectively [46], the results are in good agreement with those drawn from the values of φ V and S V * discussed earlier. CONCLUSION In summary, the study reveals that ion ion interactions are predominant over ion solvent interactions for (NH 4 ) 2 SO 4, Na 2 SO 4, K 2 SO 4, MgSO 4, ZnSO 4 and CdSO 4 in different aqueous binary mixtures of formamide at all experimental temperatures. Also, the sulphates under investigation were found to act as structure makers in the solvent mixtures studied. ACKNOWLEDGMENTS The authors are thankful to the Departmental Special Assistance Scheme under the University Grants Commission, New Delhi (No. 54/6/DRS/22, SAP-1) for the instrumental and financial assistance. REFERENCES 1. D. Das, B. Das and D. K. Hazra, J. Solution Chem. 32, 85 (23). 2. J. M. McDowali and C. A. Vincent, J. Chem. Soc. Faraday Trans. 1, 1862 (1974). 3. M. R. J. Deck, K. J. Bird and A. J. Parker, Aust. J. Chem. 28, 955 (1975). 4. M. N. Roy, B. Sinha, R. Dey and A. Sinha, Int. J. Thermophys. 26, 1549 (25). 5. M. N. Roy, B. Sinha, V. K. Dakua and A. Sinha, Pak. J. Sci. Ind. Res. 49, 153 (26). 6. M. N. Roy and A. Sinha., J. Indian. Chem. Soc. 83, 16 (26). 7. M. N. Roy and A. Sinha, Phys. Chem. Liq. 45, 67 (27). 8. A. Choudhury and M. N. Roy, Pak. J. Sci. Ind. Res. 48(3), 162 (25). ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том

10 APPARENT MOLAR VOLUME G. T. Fraser, R. D. Suenram and F. J. Lovas, J. Mol. Liq. 189, 165 (1988). 1. D. D. Perrin and W. L. F. Armarego, Purification of Laboratory Chemicals, 3rd edition, Pergamon Press, New York, (1988). 11. A. K. Covington and T. Dickinson, Physical chemistry of Organic Solvent systems, Plenum, London and New York, (1973) M. N. Roy, B. K. Sarkar and R. Chanda, J. Chem. Eng. Data, 52, 163 (27). 13. A. M. Cases, A. C. G. Marigliano, C. M. Bonatti and H. N. Solimo, J. Chem. Eng. Data, 46, 712 (21). 14. A. C. G. Marigliano and H. N. J. Solimo, Chem. Eng. Data, 47(4), 796 (22). 15. J. R. Suindells and T. B. Godfray, J. Res. Natd. Bur. Stand, 48, 1 (1952). 16. D. P. Shoemaker and C. W. Garland, Experiment s in Physical Chemistry, McGraw-Hill, New York, 131 (1967). 17. V. K. Dakua, B. Sinha and M. N. Roy, Indian J. Chem. 45A, 1381 (26). 18. B. Sinha, B. K. Sarkar and M. N. Roy, J. Chem. Thermodynamics, 4, 394 (28). 19. M. N. Roy, A. Sinha and B. Sinha, J. Solution. Chem. 34, 1311 (25). 2. M. N. Roy, B. Sinha and V. K. Dakua, J. Chem. Eng. Data, 51, 59 (26). 21. M. N. Roy and A. Sinha, Fluid Phase Equilibria, 243, 133 (26). 22. M. N. Roy and M. Das, Russian J. Phys. Chem. 8, S163 (26). 23. B. K. Sarkar, B. Sinha and M. N. Roy, Russian J. Phys. Chem. 82 (6), c1 (28). 24. O. Redllich and D. M. Meyer, Chem Rev. 64, 221 (1964). 25. D. O. Masson, Phil. Mag. 8, 218 (1929). 26. A. N. Kannappan and R. Palani, Indian J. Chem. 46A, 54 (27). 27. R. Gopal and P. Ramanand, Indian J. Chem. 16A, 25 (1978). 28. N. Saha, B. Das and D. K. Hazra, J. Chem. Eng. Data. 4, 1264 (1995). 29. F. J. Millero, J. Chem. Eng. Data, 18, 47 (1973). 3. C. Guha, J. M. Chakraborty, S. Karanjai and B. Das, J. Phys. Chem. 17 B, (23). 31. P. S. Nikam, A. S. Sawant, J. S. Aher and R. S. Khainer, J. Indian Chem. Soc. 77, 197 (2). 32. F. J. Millero, Structure and Transport Processes in Water and Aqueous Solution, edited by R. A. Horne, Wiley, New York, 15, 622 (1971). 33. R. Dey, A. Jha and M. N. Roy, J.Chem. Eng. Data, 46, 1327 (21). 34. L. G. Hepler, Can. J. Chem, 47, 4613 (1969). 35. R. R. Gupta and M. Singh, Indian J. Chem. 46A, 455 (27). 36. G. Jones and M. Dole, J. Amer. Chem. Soc. 51, 295 (1929). 37. H. Falkenhagen and M. Dole, Phys. Z. 31, 611 (1929). 38. T. S. Sharma and J. C. Ahluwalia, Rev. Chem. Soc. London, 2, 217 (1973). 39. D. Feakins, D. J. Freemantle and K. G. Lawrence, J. Chem. Soc. Faraday Trans I, 7, 795 (1974). 4. S. Glasston, K. Laidler and H. Eyring, The Theory of Rate Processes, McGraw-Hill, New York, 477 (1941). 41. M. L. Parmar and R. C. Thakur, Indian J. Chem. 45A, 1631 (26). 42. B. R. Reddy and D. L. Reddy, Indian J. Pure Appl. Phys. 37, 13 (1999). 43. D. P. Gulwade, M. L. Norwade and K. N. Wadodkar, Indian J. Chem. 43A, 212 (24). 44. F. Gucker, Chem. Rev. 13, 111 (1933). 45. Debye and Huckel, Z Phys. 24, 185 (1923). 46. K. N. Mehrotra and M. Anis, J. Indian Chem. Soc. 74, 72 (1997). ЖУРНАЛ ФИЗИЧЕСКОЙ ХИМИИ том *

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