Stability constants and thermodynamic functions of indium(III) complexes with some organic acids from potentiometric data

Stability constants and thermodynamic functions of indium(III) complexes with some organic acids from potentiometric data

,I, inorg, nut/, Chem.. 1973, Vol. 35, pp. 201-207. Pergamon Press. Printed in Great Britain STABILITY CONSTANTS AND THERMODYNAMIC FUNCTIONS OF IND...

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,I, inorg, nut/, Chem.. 1973, Vol. 35, pp. 201-207.

Pergamon Press.

Printed in Great Britain


Abstraet-Potentiometric studies of the complexes of indium(llI) with several types of organic compound are reported. The order of decreasing stability for the ligands is thiolactic, thioglycolic, lactic, glycolic, a-alanine and glycine. The stepwise changes in the free energy, enthalpy and entropy values for the complexation are evaluated in an ionic strength of0.2M (NaCIO4) at 35°C. The results have been correlated with those for the corresponding analagues of succinic and/3-substituted propionic acid. The stability of the complexes with the analogues of acetic, propionic and succinic acids decreases for a given acid as the substituted group changes from-SH to-OH to-NH2. For acids which contain the same substituent group decreasing stability is observed for the analogues of succinic, propionic and acetic acids, the only exception being/3-mercapto propionic acid. The thermodynamic values reveal that the complexes are stabilised by positive entropy changes. INTRODUCTION

IN THE PRECEDINGcommunications [1-3] the stepwise stability constants and the thermodynamic parameters for the formation of complexes between indium(III) and mercapto, hydroxy and amino analogues of succinic,/3-substituted propionic and benzoic acid have been reported. We now report the extension of this study to the corresponding analogues of o~-substituted propionic and acetic acid. The proton-ligand and metal-ligand stability constants (stepwise) have been obtained by the Bjerrum-Calvin titration technique [4, 5] as used by Irving and Rossotti [6]. The stepwise free energy changes (AG), enthalpy changes (AH) and entropy changes (AS) associated with complexation have been evaluated by the temperature coefficient equations [7]. Indium complexes of glycolic [8, 9] and lactic acid [9] have been studied by e.m.f, and solvent extraction methods. However, titration studies involving pH measurements for determination of the stepwise stability constants and thermoI. 2. 3. 4. 5. 6. 7.

R. Sarin and K. N. Munshi,J. inorg, nucl. Chem. 34, 581 (1972). R. Sarin and K. N. Munshi, A ust. J. Chem. 25,929 (1972). R. Sarin and K. N. Munshi. To be published. J. Bjerrum, Metal Ammine Formation in Aqueous Solution. Haase, Copenhagen (1941). M. Calvin and K. W. Wilson, J. Am. chem. Soc. 67, 2003 (1945). H. Irving and H. S. Rossotti,J. chem. Soe. 2904 (1954). K. B. Yatsimirskii and V. P. Vasi'l'Ev, Instability Constants o f Complex Compounds. Pergamon Press, New York (1960). 8. N. Sund6n, Svensk Kern. Tidskr. 65, 257 (1953). 9. J{mos T6rk/5 and Tamfis Lengyel, Proe. XI Int. Coord. Chem., Jerusalam, Paper no. 1-34 (1968). 201


R. S A R I N and K . N. M U N S H I

dynamic parameters of complexes of indium(III) with thiolactic, lactic, thioglycolic, glycolic, a-alanine and glycine have not been reported. An attempt is also made to assess the effects due to (a) the substitution of - S H , - O H and -NH2 (b) the presence of - S H , - O H , or -NH2 groups at the a or fl position (c) the presence of a methyl group (d) the ring size and (e) the number of tings involved in complexation. EXPERIMENTAL

Materials All chemicals employed were of reagent grade. The ligand solutions were prepared by using thiolactic acid (Schuchardt Miinchen), thioglycolic (E. Merck), lactic (Riedel), glycolic (Riedel), ot-alanine (B.D.H.) and glycine (E. Merck). The metal solution was prepared from 99"9% pure indium nitrate pentahydrate (Schuchardt Miinchen). The sodium hydroxide (carbonate free) solution was prepared by a standard technique [10] using AnalaR material (E. Merck). Stock solutions of sodium perchlorate (neutral) and perehloric acid were prepared by dissolving AnalaR (Riedel) NaCIO4. H20 and AnalaR HC10~ (E. Merck) in double distilled water.

Analysis The stock solutions of organic acids (0.02M), with the exception of glycine and a-alanine, were standardised against sodium hydroxide by potentiometric titrations. Glycine and a-alanine were standardised by adding an equal amount of a 35% solution of formaldehyde to an aqueous solution of the acid [11, 12] and then titrating potentiometricaily. The stock solution of indium nitrate (0.015M) was estimated by the precipitation of indium as the 8-hydroxyquinolinate [ 13] and sodium hydroxide (1.0M) was standardised potentiometrically against standard (0-1M) potassium hydrogen phthalate. The concentration of the perchloric acid (0.1M) stock solution was also determined by potentiometric titration against NaOH.

Procedure The experimental procedure involved potcntiometric titration against sodium hydroxide of the three mixtures prepared as follows: Mixture A-perchloric acid (0.01M); Mixture B-perchloric acid (0.01M) + organic acid (0.005M); Mixture C-perchloric acid (0-01M) + organic acid (0-005M) + metal solution. The concentration of the common ingredients were identical in different cases. The ionic strength of the solution in all cases was adjusted to 0.2M using an appropriate amount of neutral sodium perchlorate solution (1.0M) as a supporting electrolyte. The titrations were run in duplicate using a 5:1 ratio of organic acid to metal ion for mercapto acids and 3:1 and 1:~1ratios for hydroxy and amino acids. The final volume was adjusted to 100 ml using double distilled water. The titrations were carded out in a nitrogen atmosphere at various temperatures with the cell immersed in a thermostated hath (---0.2°(2). The design of the cell was such that it allowed nitrogen to he flushed through and above the solutions; it could also accomodate a glass stirrer, a microburet, a glass electrode and one limb of the agar bridge which was connected to the calomel electrode externally. A pH-meter (Beckmann model H-2) was used for the pH measurements. The equations employed in the calculation of the proton-ligand stability constants and metal-ligand stability constants have been reported earlier [1,2]. Two important factors were considered during the study: (a) the formation of chloro complexes which would lead to low values of stability constants and (h) the hydrolysis of indium(II1) which would lead to high values of stability constants. The formation of chloro complexes was prevented by connecting the calomel electrode externally through a salt bridge (saturated with KNOa). Secondly, it was possible to find a region of negligible hydrolysis by varying the metal ion and ligand concentrations. The agreement between the values of stability constants obtained at different metal ion concentrations indicated that the amount of hydrolysis was negligible below pH values of 10.0, 4-0 and 3.5 for mercapto, hydroxy and amino analogs, respectively. 10. 11. 12. 13.

A. I. Vogel, Quantitative lnorganicAnalysis. Longmans, London (1961). L. Katiznand andJ. Sullivan, J. Phys. Coll. Chem. 45, 346 (1951). S. Sorenson, Biochem. Z. 7, 43 (1907). Ltszl6 Erdey, GravimetricAnalysis, Part II. Pergamanon Press, New York (1965).

Stability constants and thermodynamic functions



The calculated values of proton-ligand stability constants of the ligands and the successive formation constants (log Kn) for the various systems at different temperatures are given in Tables 1 and 2, respectively. The overall order of decreasing stability for the ligands is thiolactic, thioglycolic, lactic, glycolic, a-alanine and glycine. The error limits are ___0.03 for log K1, ± 0-05 for log Kz, ± 0.1 for log K3 and ± 0.03 for log K n. The formation curves for mercapto analogues were studied between h = 0 and 3. N o precipitation was observed throughout the titrations. The metal curves showed significant separation from the corresponding curves in which metal ion was absent. Even at pH - 2.0, h was already 0.5 showing a strong tendency for complexation and liberation of a proton from the sulphydryl group. This conclusion was confirmed by titration of analogous mercapto acids in the presence and absence of In(Ill). The titration curves in the presence of the metal ion showed inflexions corresponding to one, two and three extra moles of base per mole of ligand when the metal-ligand ratios were 1: I, 1:2 and 1:3, respectively. Table 1. Stepwise protonation constants of the ligands at various temperatures and/~ = 0.2M Ligand


t = 25°C

t = 35°C

t = 45°C

log K~ n log K2 n








log K1H log Kz n

9.92 3.53

9.82 3.56

9-72 3.60

Lactic Glycolic ct-Alanine Glycine

log K H

3.57 3-53 2.43 2.38

3.59 3.56 2.45 2.40

3.61 3.59 2.45 2.41


iogK n log K u log K n

Table 2. Stepwise stability constants of the indium(Ill) complexes at various temperatures and tz = 0-2M Ligands


t = 25°C

t = 35°C

t = 45°C

log Kt log Kz log K3 log K~ log K2 log Ka

12-28 10-72 6.55 12.10 10.33 6"34

12"15 10.56 6.37 11.98 10-19 6"16

12"01 10.41 6.20 11.87 10-07 6-00

log K1 log K2

3" 14







log KI log Kz

2.91 2-53

3 -00 2.58

3 -07 2.63

tx-Alanine Glycine

log K~ log K~

2.51 2-39

2-57 2-46

2.63 2-54





R. S A R I N and K. N. M U N S H I

The log K values summarised in Table 2 are the mean values obtained by the various computational methods reported in the literature[14]. The formation curves for hydroxy analogues are unsymmetrical above h = 1.4. This may be due to the onset of hydrolysis at higher pH values. However, this does not affect the accuracy of log K1K2 as in the region ~ = 1, no evidence of hydrolysis is found. In the cases of amino analogues of the two acids the log K2 values obtained were not of good precision due to the early precipitation around h = 1 and, therefore, only the log K1 values are reported. The values of free energy change (AG), enthalpy change (AH) and entropy change (AS) calculated from the temperature coefficient equations [7] are summarised in Table 3. The AG, AH and AS values are not the standard AG °, AH ° and AS° which are obtained from the extrapolated values of the concentration stability constants to zero ionic strength. For comparison purpose only, relative values will suffice since under the same conditions of ionic strength the values of AG,

Table 3. Stepwise thermodynamic parameters of complexes of indium (III) at 35°C and/z = 0"2M AG1 (kcal/mole)

AG2 (kcal/mole)

AGa (kcal/mole)

--17.1 -16.9 - 4.5 - 4.2 - 3.6 - 3"4

--14.9 -14.4 - 3-7 - 3.6

-9"0 --8-7


AH1 * (kcal/mole)

AH2 (kcal/mole)

AH3 (kcal/mole)

Thiolactic Thioglycolic Lactic Glycolic a-Alanine Glycine

-5.8 -6-1 +3.2 +3.1 +3.5 +3.0

-6-5 -6-3 +2.4 +2-2

-7-4 -7.0


AS1t (cal/deg/mole)

AS~ (cal/deg/mole)

ASa (cal/deg/mole)

Thiolactic Thioglycolic Lactic Glycolic cz-Alanine Glycine

+37 +35 +25 +24 +23 +21

+27 +26 +20 + 19

+5 +5

Ligand Thiolactic Thioglycolic Lactic Glycolic a-Alanine Glycine

*_ 0.5 kcal/mole. t ± 2.0 cal/deg/mole. 14. H. Irving and H. S. Rossotti, J. chem. Soc. 3397 (1953).

Stability constants and thermodynamic functions


AH and AS for a series of related complexes will show the same variations as the values of standard thermodynamic functions. DISCUSSION

The data in.Table 4 provide a summary of results obtained in the present work and those reported earlier [ 1,2]. The various aspects which have been considered for the correlation of these results are, (a) the ratio of successive stability constants; (b) the effect of substitution of - S H , - O H and - N H 2 at a and/3 position; (c) the presence of the methyl group; (d) the ring size and (e) the number of rings involved in complexation. Unfortunately the values of AH2, AH3, AS2 and AS:~ are not very accurate, owing to the magnitude of the error limits for the subsequent complexes. Hence only the parameters for the mono-ligand complexes are included for discussion. The values of log K,JKn+I are positive in all cases. Similar results have been reported in the literature [4] and can be assumed to be due to a combination of factors such as statistics, steric hinderance and, for charged ligands, coulombic interactions. The high positive values of log K1/K2 may be explained in terms of M ---> L, 7r bonding. Indium has a d 1° configuration and due to the availability of vacant dzr orbitals on the sulphur atom, there can be interaction between it and the d-orbitals of indium(III), leading to the synergic stabilisation of a L -+ M, tr bond. Hence, if the donor dxu and d~z orbitals of the metal ion form a 7r bond with the empty d orbitals of a ligand on the X coordinate they will be less available for dTr bonding on the - X coordinate. Consequently w bonding will contribute less to the second M - L bond[15]. It is seen from Table 3 that AH2 is is almost equal to AH1 while AS2 is appreciably smaller than AS1. This reflects the Table 4. Summary of the values of indium(llI) complexes for correlations ASz (e.u.) Ref.

log K n

Thiomalic Aspartic

17.51 7.64 5-52

14-46 4.60 3.28

-19.9 -6-5 -4.6

-6.1 +4-2 +4.9

+45 +35 +31

* *















3-58 13-50 3-59 2.45 13.38 3.56 2.40

2.77 12.15 3.21 2.57 11.98 3.00 2.46

--3.9 -17.1 --4-5 -3.6 -16.9 --4;2 - 3 -4

+3.0 -5.8 +3.2 +3.5 -6. I +3.1 +3 -0

+22 ÷37 +25 +23 +35 +24 +21

$ $ ~:



Thiolactic Lactic a-Alanine Thioglycolic Glycolic Glycine

log K

AG~ AHI (kcal/mole) (kcal/mole)


*.?Data from previous work, Refs. [ 12].

*Data from this work. 15. J. Lewis and R. G. Wilkins, Modern Coordination Chemistry. lnterscience, New York (1964).


R. S A R I N and K. N. M U N S H I

decreasing importance of disruption of the hydration layer of the cation to the overall entropies of complexation. This decrease in the hydration sphere results from steric effects of the anions immediately adjacent to the cation, as well as the decrease in the net charge attracting the hydration sphere. Thus, it would seem reasonable to conclude that the greater tendency for the formation of first-step complexes is essentially due to an entropy effect. It is reported that the stability of the complexes of lanthanides [16] and some transition elements[17] decreases as the substituted group changes from -NH2 to - O H to -SH. In the present case the stability of the complexes of acetic, propionic and succinic acids decreases for any given acid as the substituted group changes from - S H to - O H to -NH~. This decrease parallels the order of decreasing basicity in any given series (Table 4). Considering those acids which contain the same substituted group, decreasing stability is observed in the analogues of succinic, ~-suhstituted propionic, a-substituted propionic and acetic acids, except in the case of the/~-mercapto-propionate complex, which is less stable than amercapto-propionate and mercaptoacetate. This can attributed to the fact that in the first case the molecule contains a less stable six-membered ring whereas in the other cases five-membered rings are present. The high stability of the mercapto analogues is also reflected by the highly negative values of enthalpy and greater positive values of entropy. The similar positive values of enthalpy for the complexes of analogous hydroxy and amino acids show that the bonding is similar in these cases. Thus, it is concluded that although enthalpy changes are unfavorable for the amino and hydroxy complexes they are stabilised by the relatively large positive entropy changes due to the liberation of water molecules from the ions accompanying complex formation, whereas for the mercapto analogues complex formation is favoured by both entropy and enthalpy. In case of a-substituted analogues the movement o f the hydrocarbon chain portion of the complex is hindered by the position of the amino and hydroxy group whereas in the g-substituted series less steric hinderance occurs. This will result in a larger disruption of the water structure around the complexes, which will account for the higher entropy and consequently greater stability for the ~substituted complexes. The general conclusion that the higher entropy changes are favorable for complexation when the reaction is endothermic is supported by the results of McAuley and Nancollas [18]. Since the above explanation is consistent with our results reported in Table 4, it is concluded that the greater stability of //-amino and fl-hydroxypropionate complexes is due to the more positive entropy changes than those for the complexes of a-amino and a-hydroxypropionate, respectively. The data also show that the complexes of hydroxy and amino analogues of propionic acid are more stable than those of the corresponding analogues of acetic acid. This may again be explained on the basis of the more positive entropy changes in the former cases. Thus, on complexation the net entropy change is positive because the disruption of the large hydration zone around indium(Ill) is more important than the energy release on complex formation. The larger 16. M. Cefola, A. S. Tompa, P. S. Gentile and A. V. Celiano, lnorg. Chem. 1,290 (1962). 17. M. Cefola, R. C. Taylor, P. S. Gentile and A. V. Celiano, J. phys. Chem. 66, 790 (1962). 18. A. McAuley and G. J. Nancollas, J. chem. Soc. 989 (1963).

Stability constants and thermodynamic functions


positive values of entropy for the propionate complexes relative to the acetate complexes is due to the presence of a bulky methyl group in the former. Our results show that the five-membered chelate rings formed with a-mercaptopropionate and mercaptoacetate are more stable than the six-membered chelate rings formed by /3-mercapto propionate. This agrees with other results [15, 19] and it may be due to the greater proton affinity of the ligands which give rise to the weaker six-membered ring complexes [20]. The thermodynamic data in Table 4 show that it is primarily entropy change which determines the relative stability of five and six-membered ring chelates. We have recently reported [3] that ortho-mercapto benzoic acid also forms six-membered rings with indium (III), but it is difficult to correlate the stabilities of the complexes with the analogues of propionic acid because as the donor group changes many other changes also occur, e.g. bond strength, ring size, steric factors and resonance[21]. The stability of fl-mercaptopropionic acid is less than that of ortho-mercapto benzoic acid. An analogous situation was reported by Schwarzenbach [22, 23] who found that a six-membered aliphatic ring binds calcium ions less effectively than a six-membered aromatic ring. Thiomalic acid can act both as a bidentate and a tridentate ligand since aspartic acid is reported to act as a tridentate ligand with first row transition metal ions [24] and as a bidentate ligand with alkaline earth ions [25]. For the alkaline earth chelates, the stabilities of the aspartates were only 0.2 log units higher than those of glycinates, and this small increase was attributed to the inductive effect of the negative carboxylate group which increases the basicity of the donor group towards the metal ion. Similarly, comparison of the stability constants for the thiomalate complexes of indium with those of the corresponding analogues of acetic and propionic acids show larger differences (approximately 2.5 log units in K t values). This difference is too large to be attributed solely to the inductive effect of an additional carboxylate ion in either a five or six-membered bidentate, and hence it is reasonable to assume that thiomalic acid forms a tridentate chelate. Similarly, it is reasonable to assume from the values of the stability constants of the malate and aspartate complexes that they are chelates in which the ligands are bidentate whereas in the corresponding complexes of propionic and acetic acids the ligands are monodentate. This is also apparent from the thermodynamic parameters, since the entropy changes are more positive for the succinic acid complexes than those of the corresponding complexes with acetic and propionic acids. Acknowledgements-We thank Prof. R. H. Sahasrabudhey, Head of the Chemistry Department, Nagpur University, Nagpur for provision of facilities and the University Grants Commission for the award of a Junior Research Fellowship (R.S). 19. L. Birkofel and L. Hartwig, Chem. Ber. 90, 260 (1957). 20. H. Irving, R. J. P. Williams, D. J. Ferrette and A. E. Williams, J. chem. Soc. 3494 (1954). 21. A. E. Martell and M. Calvin, Chemistry of the Metal Chelate Compounds. Prentice-Hall, New York (1952). 22. G. Schwarzenbach, A. Willi and R. D. Bech, Helv. chirn. Acta 30, 1303 (1947). 23. G. Schwarzenbach, H. Ackermann and P. Ruckstuhl, Heir. Chim. Acta 32, 1175 (1949). 24. S. Chaberek, Jr. and A. E. Martell, J. Am. chem. Soc. 74, 5052 (1952). 25. R. F. Lumb and A. E. Martell, J. phys. Chem. 57,690 (1953).