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

Abstract

The 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 (Mn²⁺, Co²⁺, Ni²⁺, and Cu²⁺) have been determined potentiometrically in 0.1 mol·dm⁻³ KCl and 10% (by volume) ethanol-water mixture and at 298, 308, and 318 K. The stability constants of the formed complexes increase in the order Mn²⁺, Co²⁺, Ni²⁺, and Cu²⁺. The effect of temperature was studied, and the corresponding thermodynamic parameters (Δ𝐺, Δ𝐻, and Δ𝑆) were derived and discussed. The dissociation process is nonspontaneous, endothermic, and entropically unfavourable. The formation of the metal complexes has been found to be spontaneous, endothermic, and entropically favourable.

1. Introduction

Schiff-base coordination complexes have attracted great attention over the past decades due not only to their facile syntheses, their wide application, and the accessibility of diverse structural modifications, but also to their biological modeling applications, catalysis, design of molecular ferromagnets, and materials chemistry [1–6]. It is well known that N atoms play a key role in the coordination of metals at the active sites of numerous metallobiomolecules [7]. The Schiff bases of 4-aminoantipyrine and their coordination complexes have been extensively investigated because of their biological, clinical, pharmacological, analytical and material applications [8–11]. Also some drugs showed increased activity when administered as metal chelate rather than as free organic compounds [12, 13]. In continuation of earlier work [14–18], we report here the 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 the stability constants of their complexes with Mn²⁺, Co²⁺, Ni²⁺, and Cu²⁺ at different temperatures. Furthermore, the corresponding thermodynamic functions of complexation are evaluated and discussed.

2. Experimental

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) (Figure 1) were prepared by refluxing a mixture of hot solution of 4-aminoantipyrine (0.1 mmol) and 4-hydroxyaniline or 4-carboxyaniline (0.1 mmol) in 50 cm³ ethanol for 2 h [19, 20]. The formed solid product was separated by filtration, purified by crystallization from ethanol, washed with diethyl ether, and dried in a vacuum over anhydrous calcium chloride.

Figure 1 

A ligand solution (0.001 mol·dm⁻³) was prepared by dissolving an accurately weighed amount of the solid in ethanol (Analar). Metal ion solutions (0.0001 mol·dm⁻³) were prepared from Analar metal chlorides in bidistilled water and standardized with EDTA [20]. Solutions of 0.001 mol·dm⁻³ HCl and 1 mol·dm⁻³ KCl were also prepared in bidistilled water. A carbonate-free sodium hydroxide solution in 10% (by volume) ethanol-water mixture was used as titrant and standardized against oxalic acid (Analar).

The apparatus, general conditions, and methods of calculation were the same as in previous work [15–18]. The following mixtures (i)–(iii) were prepared and titrated potentiometrically at 298 K against standard 0.002 mol·dm⁻³ NaOH in a 10% (by volume) ethanol-water mixture:(i)5 cm³ 0.001 mol·dm⁻³ HCl + 5 cm³ 1 mol·dm⁻³ KCl + 5 cm³ ethanol.(ii)5 cm³ 0.001 mol·dm⁻³ HCl + 5 cm³ 1 mol·dm⁻³ KCl + 5 cm³ 0.00l mol·dm⁻³ ligand. (iii)5 cm³ 0.001 mol·dm⁻³ HCl + 5 cm³ l mol·dm⁻³ KCl + 5 cm³ 0.001 mol·dm⁻³ ligand + 10 cm³ 0.0001 mol·dm⁻³ metal chloride.

For each mixture, the volume was made up to 50  cm³ with bidistilled water before the titration. These titrations were repeated for temperatures of 308 K and 318 K. The temperature was controlled to within ± 0.05 K by circulating thermostated water (Neslab 2 RTE 220) through the outer jacket of the vessel. The pH measurements were performed with a Metrohm 836 Titrando (KF & Potentiometric Titrator) equipped with a combined porolyte electrode. The pH-meter readings in the nonaqueous medium were corrected [21]. The 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.

3. Results and Discussion

The average number of the protons associated with ligands (L₁ and L₂) at different pH values, 𝑛_𝐴, was calculated from the titration curves of the acid in the absence and presence of ligands (L₁ and L₂). Apply (1): 𝑛_𝐴𝑉=𝑌+¹−𝑉₂𝑁⁰+𝐸⁰𝑉⁰+𝑉₁𝑇𝐶⁰_𝐿, (1) where 𝑌 is the number of available protons in ligands (L₁ and L₂) (𝑌=1) and 𝑉₁ and 𝑉₂ are the volumes of alkali required to reach the same pH on the titration curve of hydrochloric acid and reagent, respectively, 𝑉⁰ is the initial volume (50 cm³) of the mixture, 𝑇𝐶⁰_𝐿 is the total concentration of the reagent, 𝑁⁰ is the normality of sodium hydroxide solution, and 𝐸⁰ is the initial concentration of the free acid. Thus, the formation curves (𝑛_𝐴 versus pH) for the proton-ligand systems were constructed and found to extend between 0 and 1 in the 𝑛_𝐴 scale. This means that ligands (L₁ and L₂) 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. Three replicate titrations were performed; the average values obtained are listed in Tables 1 and 2. The completely protonated form of ligands (L₁ and L₂) has one dissociable proton that dissociates in the measurable pH range.

Table 1 

Thermodynamic functions for the dissociation of ligand (L₁) in 10% (by volume) ethanol-water mixture and 0.1 mol · dm⁻³ KCl at different temperatures.

Table 2 

Thermodynamic functions for the dissociation of ligand (L₂) in 10% (by volume) ethanol-water mixture and 0.1 mol · dm⁻³ KCl at different temperatures.

The formation curves for the metal complexes were obtained by plotting the average number of ligands attached per metal ion (𝑛) versus the free ligands exponent (pL), according to Irving and Rossotti [24]. The average number of the reagent molecules attached per metal ion, 𝑛, and free ligands exponent, pL, can be calculated using (2) and: 𝑉𝑛=³−𝑉₂𝑁⁰+𝐸⁰𝑉⁰+𝑉₂𝑛_𝐴𝑇𝐶⁰_𝑀,pL=log₁₀∑^𝑛=𝐽_𝑛=0𝛽^H_𝑛1/antilogpH^𝑛𝑇𝐶⁰_𝐿−𝑛𝑇𝐶⁰_𝑀⋅𝑉⁰+𝑉₃𝑉⁰, (2) where 𝑇𝐶⁰_𝑀 is the total concentration of the metal ion present in the solution and 𝛽^H_𝑛 is the overall proton-reagent stability constant. 𝑉₁, 𝑉₂, and 𝑉₃ are the volumes of alkali required to reach the same pH on the titration curves of hydrochloric acid, organic ligand, and complex, respectively. These curves were analyzed, and the successive metal-ligand stability constants were determined using different computational methods [25, 26]. The values of the stability constants (log 𝐾₁ and log 𝐾₂) are given in Tables 3 and 4. The following general remarks can be pointed out.(i)The maximum value of 𝑛 was ~2 indicating the formation of 1 : 1 and 1 : 2 (metal : ligand) complexes only [27].(ii)The 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)The metal titration curves were displaced to the right-hand side of the ligand titration curves along the volume axis, indicating proton release upon complex formation of the metal ion with the ligand. The 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 Mn²⁺, Co²⁺, Ni²⁺, and Cu²⁺ [19, 32, 33]. This order largely reflects that the stability of Cu²⁺ complexes is considerably larger than those of other metals of the 3d series. Under the influence of both the polarizing ability of the metal ion [34] and the ligand field [35] Cu²⁺ will receive some extra stabilization due to tetragonal distortion of octahedral symmetry in its complexes. The greater stability of Cu²⁺ complexes is produced by the well-known Jahn-Teller effect [35].

Table 3 

Stepwise stability constants for ML and ML₂ complexes of ligand (L₁) in 10% (by volume) ethanol-water mixtures and 0.1 mol⋅dm⁻³ KCl at different temperatures.

Table 4 

Stepwise stability constants for ML and ML₂ complexes of ligand (L₂) in 10% (by volume) ethanol-water mixtures and 0.1 mol⋅dm⁻³ KCl at different temperatures.

The dissociation constant (pK^H) for ligands (L₁ and L₂), as well as the stability constants of its complexes with Mn²⁺, Co²⁺, Ni²⁺, and Cu²⁺, has been evaluated at 298 K, 308 K, and 318 K and is given in Tables 1, 2, 5, and 6, respectively. The enthalpy (Δ𝐻) for the dissociation and complexation process was calculated from the slope of the plot pK^H or log 𝐾 versus 1/𝑇 using the graphical representation of van't Hoff equations (3) and (4): Δ𝐺=−2.303𝑅𝑇log𝐾=Δ𝐻−𝑇Δ𝑆, (3) or log𝐾=−Δ𝐻12.303𝑅𝑇+Δ𝑆.2.303𝑅 (4) From the Δ𝐺 and Δ𝐻 values one can deduce the entropy Δ𝑆 using the well-known relationships (3) and (5): ()Δ𝑆=Δ𝐻−Δ𝐺𝑇. (5) All thermodynamic parameters of the dissociation process of ligands (L₁ and L₂) are recorded in Tables 1 and 2. From these results the following conclusions can be made.(i)The pK^H values decrease with increasing temperature; that is, the acidity of the ligand increases [17].(ii)A positive value of Δ𝐻 indicates that the process is endothermic.(iii)A large positive value of Δ𝐺 indicates that the dissociation process is not spontaneous [36].(iv)A negative value of Δ𝑆 is obtained due to the increased order as a result of the solvation process.

Table 5 

Thermodynamic functions for ML and ML₂ complexes of ligand (L₁) in 10% (by volume) ethanol-water mixture and 0.1 mol⋅dm⁻³ KCl.

Table 6 

Thermodynamic functions for ML and ML₂ complexes of ligand (L₂) in 10% (by volume) ethanol-water mixture and 0.1 mol⋅dm⁻³ KCl.

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 Δ𝐻 and Δ𝑆 can then be considered as the sum of two contributions: (a) release of H₂O molecules and (b) metal-ligand bond formation. Examination of these values shows the following.(i)The stability constants (log 𝐾₁ and log 𝐾₂) for ligands (L₁ and L₂) complexes increase with increasing temperature; that is, its stability constants increase with increasing temperature [37].(ii)The negative value of Δ𝐺 for the complexation process suggests the spontaneous nature of such processes [38].(iii)The Δ𝐻 values are positive, meaning that these processes are endothermic and favourable at higher temperature.(iv)The Δ𝑆 values for the ligand complexes are positive, confirming that the complex formation is entropically favourable [16].