J. Chem. SOC., Faraday Trans. I , 1989, 85(8), 2575-2579 Activity Coefficients for Aqueous Solutions of Potassium Succinate (K,Succ) at 25 *C Miguel A. Esteso,* Luis Fernandez-Merida, Oscar M. Gonzalez-Diaz and Felipe F. Hernandez-Luis Departamento de Quimica Fisica, Uniuersidad de La Laguna, Tenerife, Spain Activity coefficients for aqueous potassium succinate solutions (K,Succ) have been determined at 25 "C from the e.m.f. measurements of cells: ISE-K I K,SUCC (m) I Hg,Succ I Hg. The results obtained were analysed by using various theoretical equations (extended Debye-Huckel, Pitzer etc.) and compared with those reported in a previous paper for the sodium salt. The knowledge of the activity coefficients of electrolytes is essential for an accurate interpretation of some properties of these electrolytes in solution.The literature' provides much data on the coefficients for single electrolytes and mixed electrolytes. Nevertheless, data of 1 : 2 organic electrolytes are not readily available. A series of such 1 : 2 electrolytes (malonates and succinates) is being used in our electrochemistry department as complexing agents in polarographic studies and it is important to know their activity coefficients in aqueous solutions. In a previous paper,2 activity coefficient values for aqueous sodium succinate were presented. Here values for aqueous solutions of potassium succinate (K,Succ) are given and the effect of the change in cation nature is analysed. Because reversible electrodes to both K+ and Succ2- ions are available, it was possible to use the e.m.f.method to determine these activity coefficients. Experimental Chemicals K,SUCC was prepared by neutralizing succinic acid (Merck 'pro analysi') with KOH (Merck 'pro analysi' with Na content less than 0.0002 %). A concentrated stock solution (1.5 mol kg-l) of this salt was prepared and afterwards the working solutions by diluting samples of it with conductivity-grade water ( K = 5 x lo-' L2-l cm-l). All the solutions were prepared by direct weighing of both the stock solution sample and the water. The preparation procedure of Hg2Succ was described in a previous paper from this laboratory. Electrodes The mercurous succinate 1 mercury electrode has also been described previously.2 Its bias potentials were generally less than 0.1 mV and never greater than 0.2 mV.The potassium ion-selective electrode, ISE-K, was Metrohm model 6.0505.0 10 with an inner impedance of 50 ML2 at 20 "C. Owing to small changes with time in its standard potential3 it was necessary to calibrate it before and after each of the measurernent~~-~ by using the reference cell: (1) in which the silver-silver-chloride electrode used, of thermal-electrolytic type, was prepared as described in the literature.6 Before each measurement, a fresh membrane was placed in the ISE-K electrode. ISE-K I KCl(O.1 mol kg-l) I AgCl 1 Ag 25752576 Aqueous Solutions of Potassium Succinate Table 1. Summary of experimental e.m.f. values and activity coefficients for aqueous K,SUCC solutions rn/mol kg-l E/mV Debye-Huckel Pitzer (y&,xp 0.02003 0.03001 0.05002 0.05999 0.08002 0.09002 0.09967 0.2000 0.3000 0.4000 0.6000 0.8000 1 .ooo 1 .500 237.2, 223.8, 207.1 , 200.8, 191.5, 187.8, 184.0, 161.6 148.1 139.4 126.1 115.8 107.7 91.9 237.1, 237.2, 0.672 223.7, 223.8, 0.634 206.9, 206.9, 0.588 201 .o, 200.9, 0.576 191.6, 191.5, 0.550 187.8, 187.6, 0.538 184.5, 184.3, 0.537 161.8 161.6 0.479 148.5 148.4 0.453 139.0 139.1 0.426 125.5 125.7 0.401 115.8 115.9 0.393 108.2 107.9 0.388 - 91.8 0.390 Cells The values of the activity coefficients presented here were obtained from the e.m.f.measurements of the galvanic cell : ISE-K I K,SUCC (m) I Hg,Succ I Hg (2) where m stands for the molality of the electrolytic solution. This cell was built into a Pyrex glass vessel designed in a double H form. The potassium ion-selective electrode was placed in the central compartment, while the other two compartments were occupied by mercury(1) succinate-mercury electrodes.The value of the estimated experimental error of these measurements was k0.15 mV for m < 0.1 mol kg-', although it increased with the increase in the concentration value m, being kO.4 mV for the highest concentration studied. Apparatus The electrometer used to carry out the e.m.f. measurements, was described in a previous paper from this laboratory.' Results and Discussion The values of the e.m.f. for the cell (2) are presented in table 1. Owing to the increase of the experimental error in the value of E when the solution concentration increases, its last decimal figure becomes meaningless from m > 0.1 mol kg-l and was therefore omitted.This e.m.f. can be expressed as: - (3) 3RT 2F E = E*--ln(4li3rny+) where E* is the standard potential of the cell, rn the molality of the solution and y* the stoichiometric mean activity coefficient value. As can be seen in eqn (3), it is necessary to know the value of E* to determine the activity coefficient values. This can be achievedM. A . Esteso et al. 2577 Table 2. Summary of the values obtained from the extended Debye-Hiickel equation E* ./A b,, a/mV mmax 88.8, 5.0 0.077, 0.1, 0.1 88.9, 5.7 0.007, 0.3, 1 Table 3. Summary of the results obtained from Pitzer's equation E* b P 8' C+ o/mV mmax 89.0, 1.2" 0.1670 2.1865 0.0001, 0.2, 1.5 89.0, 1.2" 0.1673 2.1851 0" 0.2, 1.5 89.0, 1.19, 0.1679 2.1902 0" 0.2, 1.5 a Fixed value. by using theoretical and empirical-theoretical equations for y * , which fit the experimental values throughout the whole concentration range studied or, at least, its dilute zone.The equations used can be represented by8,' (4) lny+ - = lny",'lny?+ - i.e. as the result of the sum of two contributions, the electrostatic contribution, In y:, and the short-range contribution, In y", which includes the ion-solvent and ion-ion interactions. The equations used to obtain the value of E* were the extended Debye-Huckel'O and the Pitzer'l equation. In the D-H equation, both terms are given by and and, in the Pitzer equation, by log 7; = b,, I- log (1 + 0.01 802vm) ( 5 b) a and b being 2.0 and 1.2,12 respectively, and the coefficients A, B and A,, 0.5101, 0.3285 and 0.391 5 , respectively.The values obtained for the parameters in eqn ( 5 ) and (6) and those of E* are shown in tables 2 and 3. Moreover, the values of E calculated from the Debye-Huckel and the Pitzer equations are presented in table 1. The comparison between experimental and calculated values shows a better fit by the Pitzer equation. As can be seen, the standard deviation obtained from the Debye-Hiickel equation is greater than that from the Pitzer equation, except when only the data obtained at concentrations below 0.1 mol kg-l are used for the fitting. In this case, the D-H equation would fit better. The fit of the Pitzer2578 0 -0.2 2 - -0.4 Aqueous Solutions of Potassium Succinate limiting l a w t rl 0 0 . 5 1 .o 1.5 2 .o I /(mol kg-')'/* Fig. 1. Plot of logy, - against P2 for aqueous Na,Succ and K,SUCC solutions. equation is better with the data obtained for the whole concentration range studied, even when 0" = 0, as occurs in the case of other electrolytes at concentrations up to 2 mol kg-'.'' The values obtained for E* from both equations agree with each other, the mean value of E* being 89.0 mV.By replacing this value in eqn (3), the experimental stoichiometric activity coefficient values shown in table 1, were calculated. Another attempt to obtain the value of E* was made by using the Scatchard equation13~'* to fit the experimental e.m.f. values, but the fit did not improve the Pitzer model and moreover it needed the contribution of five adjustable parameters, a(i). So, it was discarded. In fig. 1 the values of logy, for K,SUCC are plotted against Ill2, together with those for Na,Succ presented in a previous paper of our laboratory., As can be seen, within the experimental error no effect of the cation nature appears up to 0.15 mol kg-l, but from this concentration onwards Y ~ ~ , ~ ~ ~ ~ > yKzSucc. A similar behaviour has also been observed in the case of the acid salts, NaHSucc and KHSUCC.'~? l6 This behaviour is most probably due to changes in the ion-ion and ion-solvent interactions which take place in these electrolytic systems. In the first part of the plot (low concentration zone), the ion-ion interaction is predominant and therefore no differences appear because of the presence of either Na or K.In the second part, the ion-solvent interactions increase with increasing concentration and eventually predominate over the ion-ion interactions. Since this ion-solvent interaction is greater for the sodium salts than for the potassium (hNa > hK,17 h being the hydration number), the curve of Na,Succ rises much more than that of K,Succ.In fact, if we fit the experimental data to the Glueckauf hydration equation in the mole fraction statistics," for which the electrostatic term, log y y , - is given by eqn (5a) and: logy: - = --log(l-0.018hm)+-log[1-0.018m(h-v)] h - v (7) h V V a value of hKZSucc = 1.5 (for m 1 mol kg-') is obtained, whereas hNazSucc = 8.1 (for m 5 1.4 mol kg-l). These results confirm that hNaZSucc > hKZSucc. This is true even though the absolute values obtained for such hydration parameters can only be considered apparent values, due to the fact that the ionic association effects which can take place in these salts" (more marked in K,SUCC than in Na,Succ) have not been taken into account.M.A . Esteso et al. 2579 References 1 K. S. Pitzer, in Activity Coefficients in Electrolyte Solutions, ed. R. M. Pytkowicz (C.R.C. Press, Boca 2 M. A. Esteso, L. Fernandez-Merida and F. F. Hernandez-Luis, J. Electroanal. Chem., 1987, 230, 69. 3 C. C. Briggs and T. H. Lilley, J. Chem. Thermodyn., 1974, 6, 599. 4 J. Torrent, F. Sanz and J. Virgili, J. Solution Chem., 1986, 15, 363. 5 M. A. Esteso, L. Fernandez-Merida and F. F. Hernandez-Luis, J. Electroanal. Chem., 1987, 230, 77. 6 G. J. Janz, in Reference Electrodes. Theory and Practice, ed. D. J. G. Ives and G. J. Janz (Academic 7 M. A. Esteso, L. Fernandez-Merida, F. F. Hernandez-Luis and 0. M. Gonzalez-Diaz, J. Electroanal. 8 H. S . Harned and R. A. Robinson, in Multicomponent Electrolyte Solutions (Pergamon Press, Oxford, 9 M. Whitfield, in Activity Coefficients in Electrolyte Solutions, ed. R. M. Pytkowicz (C.R.C. Press, Boca Raton, Florida, 1979), vol. I, chap. 7. Press, New York, 1961), p. 209. Chem., 1988, 255, 71. 1968), chap. 2. Raton, Florida, 1979), vol. 11, chap. 3. 10 E. Hiickel, Phys. Z., 1925, 26, 93. 11 K. S. Pitzer, J. Phys. Chem., 1973, 77, 268. 12 K. S. Pitzer and G. Mayorga, J. Phys. Chem., 1973, 77, 2300. 13 G. Scatchard, J. Am. Chem. SOC., 1961, 83, 2636. 14 R. M. Rush and J. S. Johnson, J. Phys. Chem., 1968, 72, 767. 15 J. M. Stokes, J. Am. Chem. SOC., 1948, 70, 1944. 16 W. J. Hamer and Y. C. Wu, J. Phys. Chem. Ref. Data, 1972, 1, 1047. 17 E. Glueckauf, Trans. Faraday SOC., 1955, 51, 1235. 18 R. G. Bates, Pure Appl. Chem., 1973, 36,407. 19 H. S. Harned and B. B. Owen, Physical Chemistry of Electrolyte Solutions (Reinhold, New York, 3rd edn, 1958), chap. 13. Paper 8/04165B; Received 20th October, 1988