Potentiometric and Thermodynamic Studies of Some Schiff-Base Derivatives of 4-Aminoantipyrine and Their Metal Complexes

e proton-ligand dissociation constant of 4-(4-amino-1,5-dimethyl-2-phenyl-1,2-dihydro-pyrazol-3-ylideneamino)-phenol (L1) and 4-(4-amino-1,5-dimethyl-2-phenyl-1,2-dihydro-pyrazol-3-ylideneamino)-benzoic acid (L2) and metal-ligand stability constants of their complexes with metal ions (Mn, Co, Ni, and Cu) have been determined potentiometrically in 0.1mol⋅dm KCl and 10% (by volume) ethanol-water mixture and at 298, 308, and 318 K. e stability constants of the formed complexes increase in the order Mn, Co, Ni, and Cu. e effect of temperature was studied, and the corresponding thermodynamic parameters (ΔGG, ΔHH, and ΔSS) were derived and discussed. e dissociation process is nonspontaneous, endothermic, and entropically unfavourable.e formation of the metal complexes has been found to be spontaneous, endothermic, and entropically favourable.

For each mixture, the volume was made up to 50 cm 3 with bidistilled water before the titration.ese titrations were repeated for temperatures of 308 K and 318 K. e temperature was controlled to within ± 0.05 K by circulating thermostated water (Neslab 2 RTE 220) through the outer jacket of the vessel.e pH measurements were performed with a Metrohm 836 Titrando (KF & Potentiometric Titrator) equipped with a combined porolyte electrode.e pHmeter readings in the nonaqueous medium were corrected [21].e electrode system was calibrated according to the method of Irving et al. [22].All titrations have been carried out between pH 3.0 and 11.0 and under nitrogen atmosphere.

Results and Discussion
e average number of the protons associated with ligands (L 1 and L 2 ) at different pH values,   , was calculated from the titration curves of the acid in the absence and presence of ligands (L 1 and L 2 ).Apply (1): where  is the number of available protons in ligands (L 1 and L 2 ) ( = 1) and  1 and  2 are the volumes of alkali required to reach the same pH on the titration curve of hydrochloric acid and reagent, respectively,  0 is the initial volume (50 cm 3 ) of the mixture,  0  is the total concentration of the reagent,  0 is the normality of sodium hydroxide solution, and  0 is the initial concentration of the free acid.us, the formation curves (  versus pH) for the proton-ligand systems were constructed and found to extend between 0 and 1 in the   scale.is means that ligands (L 1 and L 2 ) have one ionizable proton (the enolized hydrogen ion of the -OH and -COOH group, pK H ), respectively.Different computational methods [23] were applied to evaluate the dissociation constant.ree replicate titrations were performed; the average values obtained are listed in Tables 1 and 2. e completely protonated form of ligands (L 1 and L 2 ) has one dissociable proton that dissociates in the measurable pH range.
e formation curves for the metal complexes were obtained by plotting the average number of ligands attached per metal ion () versus the free ligands exponent (pL), according to Irving and Rossotti [24].e average number of the reagent molecules attached per metal ion, , and free ligands exponent, pL, can be calculated using (2) and: where  0  is the total concentration of the metal ion present in the solution and  H  is the overall proton-reagent stability constant. 1 ,  2 , and  3 are the volumes of alkali required to reach the same pH on the titration curves of hydrochloric acid, organic ligand, and complex, respectively.ese curves were analyzed, and the successive metal-ligand stability constants were determined using different computational methods [25,26].e values of the stability constants (log  1 and log  2 ) are given in Tables 3 and 4. e following general remarks can be pointed out.
(ii) e metal ion solution used in the present study was very dilute (2 × 10 −5 mol⋅dm −3 ); hence there was no possibility of formation of polynuclear complexes [28,29].
(iii) e metal titration curves were displaced to the righthand side of the ligand titration curves along the volume axis, indicating proton release upon complex formation of the metal ion with the ligand.e large decrease in pH for the metal titration curves relative to ligand titration curves points to the formation of strong metal complexes [30,31].
(iv) For the same ligand at constant temperature, the stability of the chelates increases in the order Mn 2 , Co 2 , Ni 2 , and Cu 2 [19,32,33].is order largely re�ects that the stability of Cu 2 complexes is considerably larger than those of other metals of the 3d series.�nder the in�uence of both the polarizing ability of the metal ion [34] and the ligand �eld [35] Cu 2 will receive some extra stabilization due to tetragonal distortion of octahedral symmetry in its complexes.e greater stability of Cu 2 complexes is produced by the well-known Jahn-Teller effect [35].
e dissociation constant (pK H ) for ligands (L 1 and L 2 ), as well as the stability constants of its complexes with Mn 2 , Co 2 , Ni 2 , and Cu 2 , has been evaluated at 298 K, 308 K, and 318 K and is given in Tables 1, 2, 5, and 6, respectively.e enthalpy (Δ) for the dissociation and complexation process was calculated from the slope of the plot pK H or log  versus 1/ using the graphical representation of van't Hoff equations ( 3) and ( 4): or From the Δ and Δ values one can deduce the entropy Δ using the well-known relationships (3) and ( 5): All thermodynamic parameters of the dissociation process of ligands (L 1 and L 2 ) are recorded in Tables 1 and 2. From these results the following conclusions can be made.(i) e pK H values decrease with increasing temperature; that is, the acidity of the ligand increases [17].
(ii) A positive value of Δ indicates that the process is endothermic.
(iii) A large positive value of Δ indicates that the dissociation process is not spontaneous [36].
(iv) A negative value of Δ is obtained due to the increased order as a result of the solvation process.
All the thermodynamic parameters of the stepwise stability constants of complexes are recorded in Tables 5 and 6.It is known that the divalent metal ions exist in solution as octahedrally hydrated species [26], and the obtained values of Δ and Δ can then be considered as the sum of two contributions: (a) release of H 2 O molecules and (b) metalligand bond formation.Examination of these values shows the following.
(i) e stability constants (log  1 and log  2 ) for ligands (L 1 and L 2 ) complexes increase with increasing temperature; that is, its stability constants increase with increasing temperature [37].
(ii) e negative value of Δ for the complexation process suggests the spontaneous nature of such processes [38].
(iii) e Δ values are positive, meaning that these processes are endothermic and favourable at higher temperature.
(iv) e Δ values for the ligand complexes are positive, con�rming that the complex formation is entropically favourable [16].
T 3: Stepwise stability constants for ML and ML 2 complexes of ligand (L 1 ) in 10% (by volume) ethanol-water mixtures and 0.1 mol ⋅ dm −3 KCl at different temperatures.
Standard deviations are given in parentheses.