CHAPTER -VIII. Ion-Solvent Interactions of Sodium Molybdate in Oxalic acid-water Systems at Various Temperatures.

Transcription

1 CHAPTER -III Ion-Solvent Interactions of Sodium Molybdate in Oxalic acid-water Systems at arious Temperatures. 8.. Introduction iscometric, volumetric and acoustic studies of solute in case of aqueous solutions are related to the interactions among the components of a solution -3, knowledge of which is important in solution chemistry. These interactions helps in better understanding of the nature of solute and solvent, i.e., whether the solute modifies or distorts the structure of the solvent. In highly dilute aqueous solutions the two properties, apparent molar volume and viscosity are needed to build up such knowledge of intermolecular interactions. The structural and dynamic nature of water molecules in different environments is also important in aqueous solutions. Moreover carboxylic acids-water solvent system 4-6 and in water systems with other co-solute 7- have been extensively used in volumetric and viscometric studies of salts of carboxylic acids. Recently, studies relating to the measurements of partial molar volume of oxalic acid in aqueous solution and in aqueous solution of fructose have been reported. In both the studies the data indicates the presence of strong ion-solvent interactions. In our present work the effect of addition of sodium molybdate in various mole-fractions of aqueous oxalic acid solution have been studied at various temperatures. Partial molar volumes ( ) and viscosity B- coefficients ( B) of transfer from water to aqueous oxalic acid mixtures have been calculated and discussed in terms of ion-solvent interactions. Partial molar volumes at infinite dilution have been fitted to a second order polynomial equation in terms of temperature and the structure-making or breaking capacity of the electrolyte has φ. been inferred from the sign of ( δ δt ) P φ *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

2 3 8.. Experimental Oxalic acid (of Analytical Reagent Grade) was used after drying over PO5 in a desiccators for more than 4 hours. Sodium molybdate (E. Merck, India) was purified by re-crystallizing twice from conductivity water and then dried in a vacuum desiccators over PO5 for 4 hours before use. Densities (ρ) were measured with an Ostwald-Sprengel type pycnometer having a bulb volume of about 5cm 3 and an internal diameter of about. cm. the measurements were done in a thermostat bath controlled to ±. K. the viscosity (η) was measured by means of suspended Ubbelohde type viscometer, calibrated at 98.5 K with triply distilled water and purified methanol using density and viscosity values from literature. The flow times were accurate to ±.s, and the uncertainty in the viscosity measurements was ± x -4 mpa.s. Adequate precautions were taken to minimize evaporation loses during the actual measurements. The reproducibility in mole fractions was within ±. units. The mass measurements were done on a Mettler AG-85 electronic balance with a precision of ±. mg. The precision of density measurements was ±3x -4 g cm -3. iscosity of the solution, η, is given by the following equation: η= ( Kt L/ t) ρ () where K and L are the viscometer constants and t and ρ are the efflux time of flow in seconds and the [ density of the experimental liquid, respectively. The uncertainty in viscosity measurements is within ±.3 mpa.s. Details of the methods and techniques of density and viscosity measurements have been described elsewhere 3-6.The electrolyte solutions studied here were prepared by mass and the conversion of molality to molarity was accomplished 7 using experimental density values. The experimental values of concentrations(c), densities (ρ ), viscosities (η), and derived parameters at various temperatures are reported in table. *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

3 Results and Discussions The apparent molar volumes ( φ ) were calculated from the observed densities using the following standard expression equation 8 : φ = M ρ( ρ ρ) cρ () Where M is the molar mass of the solute, c is the molarity of the solution; ρ and ρ are the densities of the solvent and the solution respectively. The limiting apparent molar volumes φ were calculated by least square treatment to the plots of φ versus c using the Masson equation 9. φ * =φ + S c (3) where φ is the partial molar volume at infinite dilution and c the experimental slope. The plots of φ against square root of molar concentration ( c) were found to be linear with negative slopes. Calculated values of φand * Sare reported in table 3. φ values can be used to interpret ion-solvent interactions. Table 3 reveals that φ values are positive and increases with increase in temperature and decreases with increasing concentration of oxalic acid in the solvent mixture, suggesting larger electrostriction at higher temperature and in lower amount of oxalic acid in the mixture. However, in case of water, φ are positive but the values decreases with rise of temperature. This may be attributed to a slow desolvation and thermal agitation at higher temperature. From table 3 it is evident that * S values are negative at all temperatures and the values decreases with the increase of experimental temperature which may be attributed to more violent thermal agitation at higher temperatures diminishing the force of ion-ion interactions (ionic-dissociation). However, * S values increases in water at higher temperatures which results in a decrease in hydration of ions i.e., more and more solute molecules are accommodated in the void space left in the *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

4 33 packing of water molecules. A quantitative comparison of the magnitude of and * S values shows that ion-solvent interactions dominate over ion-ion interactions in all the solutions and at all experimental temperatures. Partial molar volumes φ φ at infinite dilution were fitted to a second order polynomial in terms of absolute temperature (T) using the expression φ = a+ at + over the temperature range under study where T is the temperature in K. at alues of coefficients of the above equation for sodium molybdate in aqueous oxalic acid mixtures are reported in table 4. The partial molar expansibilities ( φ ) at infinite dilution can be obtained by differentiating equation (4) with respect to temperature. φ E (4) ( δ ) = a at (5) E = δφ T P + The values φeof different solutions of the studied electrolytes at 33.5, 33.5 and 33.5 K are reported in table 5. Table 5 reveals that φevalue increases with temperature in water but in ternary solutions φedecreases slightly with increasing temperature. This may be attributed to gradual disappearance of caging or packing effect, in the ternary solutions. According to Helper 3 the sign of ( δφ δ is a E T) P better criterion in characterizing the long range structure-making and breaking ability of the solutes in solution. The general thermodynamic expression used is as follows ( φ δt) = ( δ φ δt ) a (6) δ E = P P δ E δt) P If the sign of ( φ is positive the solute is a structure maker otherwise it is a structure breaker. As is evident from table 5, sodium molybdate predominately acts as a structure breaker in all the solvent mixture. Partial molar volume φ of transfer from water to different aqueous oxalic acid solution has been determined using the following relation 4, 5 : *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

5 34 The φ = φ [Aqueous catechol solution] φ [Water] (7) φ value is independent from solute-solute interaction and provides information regarding solute and co-solute interaction 4. As can be seen from table 7, the values of φ are negative at 33.5K for.5(m) and.5(m) oxalic acid solution but becomes positive at higher temperatures and increases monotonically with concentration of oxalic acid in the remaining solutions. These results further confirm the presence of strong ion-solvent interactions in the chosen solution for sodium molybdate. The experimental results of relative viscosities of sodium molybdate in water and binary mixtures of aqueous oxalic acid have been analyzed using Jones-Dole 6 equation: (ηη ) c = (η r ) c= A+ Bc (8) Where η and η are the viscosities of the solvent/solvent mixtures and solution respectively. A and B are the constants estimated by a least-squares method and are reported in table 6. From the table it is evident that the values of the A coefficient are negative for all the solutions under investigation at all experimental temperatures. These results indicate the presence of very weak ion-ion interactions and these interactions further decrease with the rise of experimental temperatures and increase with an increase of oxalic acid in the mixture. The effects of ion-solvent interactions on the solution viscosity can be inferred from the B-coefficient 7, 8. The viscosity B-coefficient is a valuable tool to provide information concerning the solvation of the solutes and their effects on the structure of the solvent. From table 6 it is evident that the values of the B-coefficient are negative in water but becomes positive just on addition of small amount of oxalic acid and the B-coefficient values decreases as the concentration of oxalic acid increases. This indicates that the ion-solvent interactions which were quite weak in water have been enhanced by the addition of oxalic acid in water. In other words, sodium molybdate mix more ideally with water+oxalic acid as compared to water and there is perfect solvation of sodium molybdate in water+oxalic acid mixtures *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

6 35 resulting in strong ion-solvent interaction. These conclusions are in excellent agreement with those drawn from φ values discussed earlier. iscosity B-coefficient of transfer ( B) from water to different aqueous oxalic acid solutions have been determined using the relation 4, 5. B= B [Aqueous oxalic acid solution] B [Water] (9) The B values as shown in table 7 ( B vs. molarity of oxalic acid in solution) as a function of molarity of oxalic acid supplements the result obtained from φ as discussed above. A number of studies 9, 3 suggest that db/dt is a better criterion for determining the structure-making/breaking nature of any solute rather than B-coefficient. Table 6 indicates that the values of B-coefficient increases with rise of temperature (positive db/dt) suggesting the structure-breaking tendency of sodium molybdate in the studied solvent systems. The adiabatic compressibility (β) was evaluated from the following equation: β = u ρ where ρ is the solution density and u is the sound speed in the solution. The apparent molal adiabatic compressibility ( φ k ) of the solutions was determined from the relation, φ = β / ρ + ( βρ β ρ) / mρρ k M Where β, β are the adiabatic compressibility of the solvent and solution respectively and m is the molality of the solution. Limiting partial molal adiabatic compressibilities φkand experimental slopes * SKwere obtained by fitting φkagainst the square root of molality of the electrolyte musing the method of least squares. φ = φ + S m * k k k *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

7 36 alues of m, u,β, φ k, φkand * SKare presented in table 9. Since the values of φ K and * SKare measures of ion-solvent and ion-ion interactions respectively, a perusal of table 9 shows that the values are in excellent agreement with those drawn from the * values of φ and Sdiscussed earlier. The viscosity data have also been analyzed on the basis of transition state theory of relative viscosity of solutes as suggested by Feakins. et. al 3. The B parameter in terms of this theory is given by the following equation. B= ( φ φ ) / + φ ( µ µ ) /RT 3 o # # v, v, v, o o Where φv,and φv,are the partial molar volumes of the solvent and the solute respectively. The free energy of activation per mole of the pure solvent ( µ ) and free energy of activation per mole of the pure solute µ was calculated with the help of equation 4 and 5 respectively µ = G = RT ln( ηφ hn A ) (4) and µ = µ + RT ( B+ φ φ ) / φ (5) where N is Avogadro s number and the other symbols have their usual significance. The values of µ and µ are reported in table 8. From table 8 it is evident that µ is practically constant at all the solvent composition and temperature, implying that µ is mainly dependent on the viscosity B-coefficients and ( φ φ ) terms. It is also clear from table 8 that the values of,, µ are positive and larger than µ at all experimental temperatures which suggests that the formation of the transition state is less favored 3 in presence of sodium molybdate, thereby suggesting that the formation of transition state is accompanied by the rupture and distortion of the intermolecular bonds between oxalic acid and water i.e., solvent 3. According to Feakins model, µ increases with temperature for solutes having positive values of db/dt. This is nicely shown by sodium molybdate *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

8 37 as recorded in table 8. The entropy of activation for solutions has been calculated using the following relation 3. Where S = d ( µ ) / dt (6) # # S has been determined from the negative slope of the plots of µ against T by using a least square treatment. The activation enthalpy The value of S and H has been calculated using the relation: H = µ + T S (7) H are listed in table8 and they are found to be negative for all the solutions and at all experimental temperatures suggesting that the transition state is associated with bond breaking and increase in order. Although a detailed mechanism for this cannot be easily advanced, it may be suggested that the slipplane is in the disordered state 3, Conclusion The values of apparent molar volume ( φ ) and viscosity B- coefficients for sodium molybdate indicate the presence of strong ion-solvent interactions and these interactions are further strengthened at higher temperature. Also sodium molybdate acts as a water-structure promoter due to hydrophobic hydration in the presence of oxalic acid. *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

9 38 References. R. Mathpal, B. K. Joshi, S. Joshi, N. Kandpal, Monahshefte fur Chemie. 6,37, M. B. Roga, T. M. Sidd, J. Baethel, R.Neueder, A. Apelblat, J. Soln. Chem.,, C. Zhao, P. Ma, J. Li. J. Chem. Thermodyn. 5, 37, R. K. Chowdoji, R. S. Brahmji. Indian J. Chem. 983, A, P. K. Mandal, B. K. Seal, D. K. Chatterji, A. S. Basu. J. Soln. Chem. 978, 7, M. C. S. Shubha, R. K. Chowdji, R. S. Brahmji. J. Indian Chem. Soc. 99, 67, U. N. Das, B. K. Mohanty. Indian J. Chem. 996, 35A, R. L. Blokhara, S. Kumar, S. Kant. Indian J. Chem. 99, 9A, P. S. Nikam, A. R. Hiray. J. Indian Chem. Soc. 989, 66, M. L. Parmer, R. K. Awasthi, M. N. Guleria. J. Chem Sci. 4, 6(), 33.. R. R. Gupta, M. Singh. Indian J. Chem. 7, 46A, M. L. Parmer, M. N. Guleria. J. Indian Chem. Soc. 5, 8, M. N. Roy, A. Sinha, B. Sinha. J. Solution. Chem. 5, 34, M. N. Roy, B. Sinha,. K. Dakua, J. Chem. Eng. Data. 6, 5, M. N. Roy, A. Sinha, Fluid Phase Equilib. 6, 43, M. N. Roy, M. Das. Russian. J. Phys.Chem. 6, 8, S D. P. Shoemaker, C. W. S. Garland, Experiments in Physical Chemistry (Mc Graw Hill, New York 967). 8. M. N. Roy, B. Sinha, R. Dey, A. Sinha. Int. J. Thermophys. 5, 6, D. O. Masson, Phil. Mag. 99, 8, 8.. C. Guha, J. M. Chakraborty, S. Karanjai, B. Das. J. Phys. Chem. B. 3, 7, 84.. F. J. Millero, Structure and Transport Process in Water and Aqueous Solutions (R. A. Horne, New York 97).. M. L. Parmar, D. S. Banyal. Indian J. Chem. 5, 44A, 58. *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

10 39 3. L. G. Hepler, Can. J. Chem. 969, 47, K. Belibagli, E. Agranci. J. Solution. Chem. 99, 9, R. K. Wadi, P. Ramasami, J. Chem.Soc.Faraday Trans, 997, 93, G. Jones, M. Dole, J. Am. Chem. Soc. 99, 5, F. J. Millero, Chem. Rev. 97, 7, F. J. Millero, A. Losurdo, C. Shin. J. Phys. Chem. 978, 8, R. Gopal, M. A. Siddique. J. Phys. Chem. 969, 7, N. Saha, B. Das. J. Chem. Eng. Data997, 4, D. Feakins, D. J. Freemantle, K. G. Lawrence, J. Chem, Soc. Faraday Trans. I, 974, 7, B. Samantaray, S. Mishra, U. N. Dash, J. Teach Res Chem. 5,, Mithlesh, M. Singh, J. Indian. Chem. Soc. 6, 83, 83. *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

11 4 Table Density (ρ) and viscosity (η) of aqueous oxalic acid mixtures at different temperatures Temp. (K) ρ /kg m η /mpa s U /m. sec -.5(M) (M) (M) *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

12 4 Table. Concentration (c), density (ρ), viscosity (η), apparent molar volume ( φ v ), ( ηr ) / cof sodium molybdate in different aqueous oxalic acid mixtures at different temperatures. c /mol dm 3 ρ /kg m η/ mpa s.(m) φ 6 /m mol 3 ( ηr ) c 33.5 K K K *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

13 4.5(M) 33.5 K K K (M) 33.5K *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

14 K K (M) 33.5K K *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

15 K *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

16 45 Table 3 Limiting apparent molar volumes ( φ ) and experimental slopes ( molybdate at different temperatures. * S ) for sodium Aqueous oxalic acid solution/(mol dm -3 ) / m mol 6 3 φ / (m mol ) * / S 33.5K 33.5K 33.5K 33.5K 33.5K 33.5K Table 4 alues of the coefficients of equation (4) for sodium molybdate in different aqueous oxalic acid mixtures. Aqueous oxalic acid 3 solution /(mol dm ) a 6 / 3 m mol a 6 / a 6 / (m mol K ) (m mol K ) *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

17 46 Table 5 Limiting partial molar expansibilities for sodium molybdate in different aqueous oxalic acid mixtures at different temperatures. Aqueous oxalic acid solution/(mol dm -3 ) 3 - φ / (m mol K ) 6 E 33.5K 33.5K 33.5K δφ E / (m mol K ) δt p Table 6 alues of A and B coefficients sodium molybdate in different aqueous oxalic acid mixtures at different temperatures Aqueous oxalic acid solution/(mol 3/ / A / (m mol ) 3 B/ (m mol ) dm -3 ) 33.5K 33.5K 33.5K 33.5K 33.5K 33.5K *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

18 47 Table 7 Partial molar volumes ( φ ) and viscosity B-coefficients ( B) of transfer from water to different aqueous oxalic acid solutions for sodium molybdate at different temperatures. Aqueous oxalic acid solutionmol dm 3 6 φ /m mol 3 6 φ /m mol 3 3 B/ (m mol ) B 3 / (m mol ) 33.5K K K *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

19 48 Table 8 alues of ( φ φ ),,,,, µ µ S T and H for sodium molybdate in different aqueous oxalic acid mixtures at different temperatures. Parameter 33.5K 33.5K 33.5K. mole.dm -3 ( φ φ ) / m mol ,, µ /(kj mol ) µ /(kj mol ) T S /(kj mol ) H /(kj mol ) ,,.5 mole.dm -3 ( φ φ ) / m mol µ /(kj mol ) µ /(kj mol ) T S /(kj mol ) H /(kj mol ) ,,. mole.dm -3 ( φ φ ) / m mol µ /(kj mol ) µ /(kj mol ) T S /(kj mol ) H /(kj mol ) *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

20 49 6 3,,.5 mole.dm -3 ( φ φ ) / m mol / kj mol µ µ /(kj mol ) T S /(kj mol ) H /(kj mol ) Table 9 Molality (m), density(ρ), sound speed(u), adiabatic compressibility(β), partial molal adiabatic compressibility( φ K ), limiting partial adiabatic compressibility( φ K), and experimental slope( S * K) of sodium molybdate in different aqueous oxalic acid mixtures at 33.5K. m / mol kg u / m s β. /Pa φ. K m mol Pa 3 φ. K m mol Pa 3 S K. * 3 3 m mol Pa kg mol dm mol dm *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3

21 mol dm mol dm *Published in Journal of Teaching and Research in Chemistry 8, 5(), 3