All the potentiometric measurements were carried out on pH-meter model ELE international, using combined glass electrode (accurate total 0.01 pH units). Conductometric titration measurements were carried out using conductivity meter model 4320, Jenway, using an immersion cell. Before and after each titration, the electrode was calibrated using standard buffer pH ≈ 4.01, pH ≈ 7.00 and pH ≈ 9.00.

All the reagents were of analytical grade. Pure ethanol and bidistilled water were used for the preparation of the solutions. Sulphathiazole (4-amino-N-(1,3-thiazol-2-yl) benzenesulfonamide, STZ) was purchased from Sigma Chemical Company (USA). The salts of metallic nitrates (BDH, U.K, GENEVA or INDIA) will be prepared. The aqueous metal ion solution was standardized according to a well-known method [22] . Standard solutions of 0.1 M sodium hydroxide, 0.01 M perchloric acid and 0.09 M sodium perchlorate in aqueous solutions were prepared as usual. An STZ stock solution (0. 1 M) was prepared by weighing directly in pure ethanol because it is insoluble in water, slightly soluble in ethanol (96%), in chloroform and in ether. The working solutions (0.001 M) were prepared by accurate dilution using the same solvent.

All pH measurements were performed at 25˚C ± 0.1˚C. The medium was aqueous acid and ionic strength 0.1 M NaClO₄. For binary complexes studies, three types of mixtures (with total volume 50 mL for each) were used; (a) free acid, (b) STZ (as ligand) and (c) chelate produced from individual reaction of STZ with different metal ions; Ti(II), Zr(IV), Sr(II), Al(III), Cr(III), Fe(III), Th(IV), Pb(II), La(III) and Co(II). As shown in Figure 1, mixtures a, b and c

were individually titrated against standard alkali and the plots of pH versus volume of alkali gave the titration curves.

n ¯ H = Y + ( V 1 − V 2 ) ( N ∘ + E ∘ ) ( V ∘ + V 1 ) T c L ∘ (1)

where V₂ and V₁ are the volumes of alkali required to reach the same pH in acid and ligand titration curves. TcL^˚ The total ligand concentration Y is the total number of protons free attached to the ligand molecule, N^˚ is the normality of the alkali, E^˚ is the initial concentration of the free acid and V₀ is the total volume of the titrated solution. The titration curves were used to evaluate n ¯ H (average number of protons associated with STZ); calculations of proton ligand association constants were carried out by plotting a graph of n ¯ H against pH, as we can see in Figure 2. The concentrations were: TcL^˚ = 0.001 M, TcM^˚ = 0.001 M. From these data, the proton-ligand and metal-ligand stability constants were obtained as in Table 1. The maximum numbers of protons that can be released from STZ are two protons in the titration with diluted base (0.1 M) in the range of pH 2.20 - 11.20. STZ behaves as dibasic acid [H₂-STZ]. The acid-base properties of STZ in 25% (v/v) EtOH medium at different ionic strengths of NaClO₄ (I =0.1, 0.2, 0.3, 0.4 and 0.5 M) indicate that one proton from the protonated amino group (NH₂ → NH 3 + ) was deprotonated in the pH range (3.63 - 4.95).

The ionic strength has an important effect in the electrolytic solutions. It also offers ionic atmosphere, which impacts the stability constant of metal complexes. No precipitates produced, indicating that there was no propensity to form hydroxo complexes as the number of moles of NaOH consumed was equivalent to the number of moles of NaClO₄.

It is observed that logk_a values inversely proportional with the concentration of ionic strength.

The second site is the dissociation of proton in the amino group (NH) in pH range (7.10 - 9.40). The values of log K 1 H and log K 2 H (the first and second

proton formation constants of STZ, respectively) are the pH values corresponding to n ¯ H = 0.5 and 1.5, respectively. The values of log K 1 H (7.20) and log K 2 H (3.80) are tabulated in Table 1. However, the reaction mechanism is show as follow:

*These ratios are from potentiometric and conductometric methods.

The formation curves of the complexation equilibrium as shown in Figure 3 is obtained by plotting the degree of the complex formation ( n ¯ ) against the negative logarithm of the concentration of non-protonated ligand (pL) using the Irving and Rossotti equations [23] . From these formation curves, the values of stability constants listed in Table 1 were determined using the half-integral method [23] . The obtained results shows that both Al(III) and Zr(IV) ions forms (1:1), (1:2) and (1:3) metal to ligand complexes. Also, some metal ions viz; Zr(IV), Sr(II), Al(III), Fe(III), Th(IV) and Pb(II) forms (1:1) and (1:2) metal to ligand complexes, but in the case of Co(II), Cr(III), Ti(II) and La(III) only one complex (1:1) metal to ligand was formed. This may be due to the nature of metal ion, concentration of ligand and ionic strength.

n ¯ = ( V 3 − V 2 ) ( N ∘ + E ∘ ) ( V ∘ + V 2 ) n ¯ A T c M ∘ (2)

p L = log [ 1 + β 1 [ H + ] + β 2 [ H + ] 2 ( T c L ∘ − n ¯ T c M ∘ ) × V ∘ + V 3 V ∘ ] (3)

where Tcm^˚ denotes the total concentration of metal present in solution, β is the proton ligand stability constant, the other terms have their usual meaning as mentioned before. The stability constants of the complexes formed with STZ were listed in Table 1 in the order of stability of the different binary complexes formed between STZ and the metal ions investigated in this study, is in the expected Irving-Williams order [23] :

Fe ( III ) > Co ( II ) > Ti ( II ) > Zr ( IV ) > Al ( III ) > La ( III ) > Cr ( III ) > Sr ( II ) > Pb ( II ) > Th ( IV )

STZ behavior may be based on the bidentate nature which coordinates through the oxygen atom of sulphonamide group and a sulphur atom in the thiazole ring, forming stable six-membered chelate ring:

Ratio (1:1) metal: ligand

Also, the study of ionic strength effect on the stability constant of STZ with Al(III), Th(IV), Cr(III), Pb(II) and Zr(IV) metal ions in aqueous solution at 25˚C ± 0.1˚C was discussed. The relation between log K 1 and the ionic strength is shown in Figure 4. As we can see, the stability constants of metal-ligand complex (1:1) were decreased as the ionic strength increase.

The conductometric analysis shows a greater use in tracing the complex formation. This method finds its useful application as a sensitive tool to test for decimal variations in ionic radii of transition metal ions investigated. The conductometric analysis is based on changes in the electrical conductivity values of solutions as a result of complex formation. These changes depend on the number of ions present, and their mobility’s. Conductivity measurements are employed to trace the different types of chelate species formed between metal ions and STZ. Figure 5 shows the conductometric titration curve for the binary ligand system containing Ti(II), Zr(IV), Sr(II), Al(III), Cr(III), Fe(III), Th(IV), Pb(II), La(III) and Co(II) ions. As can be seen, an initial decrease in conductance and a minimum at (1:1).This may be due to the neutralization of H⁺ ions resulting from the formation of the (M (STZ)) complex. Furthermore, the conductance

increases slightly between (1:1), (1:2) and/or (1:3), probably due to the formation of the binary ligand complex and the release of H⁺ ions from STZ.

Distribution curves of STZ at I = 0.01M NaClO₄ shown in Figure 6 observe the pH range (2.20 - 4.80) the major species is α̥ (H₂-STZ) and in the pH range (4.80 - 6.00) the major one is α₁ (H-STZ^-1), but α₂ (STZ^-2) is major in the pH range (6.00 - 11.40). The species distribution curves can be obtained by plotting the mole fraction of metal species and pH. The curves are depicted in Figure 7. The analysis of these diagrams reveals that at low pH value, most of the metal ions are often present as free ions. This indicated that no complex occur in the acidic medium. On increasing the concentration of the ligand by increasing the pH of the solution during the titration, the mole fraction of the free metal ion tends to decrease, while that of ML species tend to develop at moderately acidic media. However, the value of log K 1 > log K 2 indicates that there will be an appreciable

concentration of ML species in this pH region. Further increasing the pH of the solution, the essential change is the increase in ML2 concentration with decrease in ML. Above this region almost the entire metal ion remains in the form of ML or ML₂ species on increasing the pH of the solution. Some fraction species at intersection points and maximum pH, complexes are represented.

The stability constants for the ternary systems were computed from titrations in which the concentrations of metal ions: STZ: Gly were kept in the ratio 1:1:1, listed in Table 2. The experimental data shown in Figure 8 reveal that the formation of the ternary complex M (STZ) (Gly) shifts the buffer region of the ligands to lower pH values, which indicates that the ternary complex is more stable than the binary complex. The horizontal distance between curves e and f

were measured and used for the calculation of n ¯ m i x (average number of secondary ligand (L) molecule attach edper (M-STZ) binary complex using Equation 4:

n ¯ m i x = ( V 4 − V 3 ) − ( V 2 − V 1 ) [ ( N ∘ + E ∘ ) + T c L ∘ ( Y − n ¯ H ) ] ( V ∘ + V 3 ) n ¯ H   T c M ∘ (4)

where V₁, V₂, V₃, and V₄ are the volumes of NaOH required to reach the same pH values for solutions of (free acid), (free acid + STZ), (free acid + STZ + metal ion) and (free acid + STZ + metal ion + Gly), respectively. The difference ( V 4 – V 3 ) – ( V 2 – V 1 ) can be used for the calculation of n ¯ m i x (average number of secondary ligand molecules associated with one [ M ( Gly ) ] n + ion. The free secondary ligand exponent, pL_mix was calculated using Equation (5):

p L m i x = log [ ∑ n = 0 i β n H ( 1 10 B ) n T c L ∘ − n ¯ m i x T c M ∘ . V ∘ + V 4 V ∘ ] (5)

According to the obtained results, the complex equilibria of M-STZ-Gly can be represented by the following scheme:

M + STZ ↔ M ( STZ ) (6)

M ( STZ ) + Gly ↔ M ( STZ ) ( Gly ) (7)

K M ( STZ ) ( Gly ) M ( STZ ) = [ M ( STZ ) ( Gly ) ] [ M ( STZ ) ] [ Gly ] (8)

In order to compare the stabilities of the ternary complex species with those of the parent binary complexes, the ∆logK value (the difference between the stabilities of the binary and the ternary complexes) were determined. The parameter ∆logK is determined by Equation (9):

Δ log K = log K M ( STZ ) ( Gly ) M ( STZ ) − log K M ( Gly ) M (9)

It was found that the difference was positive in terms of stability and showed a statistical increase in the value of stability constants of the mixed ligand complex. As shown in Table 2, the values of ternary stability constants are found to decrease in this order: Al(III) > Ti(II) > Pb(II) > Th(IV) > Zr(IV) > Sr(II) > La(III) > Cr(III) > Co(II) > Fe(III)

Formation curves corresponding to the various metal ions-STZ-Gly systems were obtained by plotting n ¯ m i x vs. pL_mix, the results are shown in Figure 9. Values of the formation constant of the STZ ternary complexes (tabulated in Table 2) shows that, the metal ions such as Al(III) and Th(IV) used in M-STZ-Gly ternary complexes are more stable than the (1:1) M-STZ binary complex and

the (1:1) M-Gly binary complexes. Thus, the (1:1) M-STZ complex has a greater tendency toward combination with Gly molecule to form ternary complexes. But some metal ions viz; Fe(III), Cr(III), La(III) and Co(II) forms less stable ternary complexes with STZ and Gly than the other corresponding binary complexes of them. Therefore, the binary complexes of the primary ligand (STZ) with these metal ions may be at odds with the combination with the Gly molecule to form ternary complexes. On the other hand, some metal ions like, Pb(II), Sr(II), Ti(II) and Zr(IV) have up-normal values. This behavior may be interpreted on the basis of statistical considerations and the nature of species formed in the solution.