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

Authors: A. A. El-Bindary; A. Z. El-Sonbati; M. A. Diab; M. K. Abd El-Kader
Journal: Journal of Chemistry
 (2013)

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## 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.


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## Body

## 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 1A 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):
(1)n-A=Y+(V1-V2)(N0+E0)(V0+V1)TCL0,
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, TCL0 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, pKH), 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.
Dissociation
Gibbs energy
Enthalpy
Entropy
T/K
constant
kJ·mol−1
kJ·mol−1
J·mol−1·K−1
pKH
Δ
G
1
Δ
H
1
-
Δ
S
1
298
9.25 (0.10)
52.77
70.56
308
9.07 (0.09)
53.48
31.75
70.57
318
8.90 (0.09)
54.19
70.56
Standard deviations are given in parentheses.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.
T/K
Dissociationconstant
Gibbs energykJ·mol−1
EnthalpykJ·mol−1
EntropyJ·mol−1·K−1
pKH
Δ
G
1
Δ
H
1
−ΔS1
298
4.26 (0.09)
24.31
24.92
308
4.09 (0.09)
24.12
31.73
24.72
318
3.91 (0.08)
23.81
24.92
Standard deviations are given in parentheses.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:
(2)n-=(V3-V2)(N0+E0)(V0+V2)n-ATCM0,pL=log10∑n=0n=JβnH[1/antilogpH]nTCL0-n-TCM0·V0+V3V0,
where TCM0 is the total concentration of the metal ion present in the solution and βnH 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 ofn- 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-Tellereffect [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.
M
n
+
298 K
308 K
318 K
logK1
logK2
logK1
logK2
logK1
logK2
Mn2+
6.84 (0.08)
6.04 (0.11)
7.02 (0.09)
6.22 (0.10)
7.19 (0.10)
6.40 (0.11)
Co2+
6.95 (0.09)
6.13 (0.10)
7.13 (0.09)
6.31 (0.10)
7.31 (0.11)
6.48 (0.11)
Ni2+
7.02 (0.09)
6.19 (0.09)
7.20 (0.10)
6.37 (0.09)
7.38 (0.11)
6.54 (0.10)
Cu2+
7.15 (0.09)
6.34 (0.09)
7.34 (0.09)
6.52 (0.10)
7.53 (0.11)
6.70 (0.10)
Standard deviations are given in parentheses.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.
M
n
+
298 K
308 K
318 K
logK1
logK2
logK1
logK2
logK1
logK2
Mn2+
7.33 (0.08)
6.74 (0.09)
7.50 (0.09)
6.90 (0.10)
7.67 (0.10)
7.07 (0.10)
Co2+
7.54 (0.09)
6.85 (0.10)
7.72 (0.10)
7.02 (0.10)
7.90 (0.10)
7.20 (0.11)
Ni2+
7.63 (0.09)
6.92 (0.10)
7.80 (0.11)
7.10 (0.09)
7.98 (0.11)
7.27 (0.10)
Cu2+
7.78 (0.09)
7.07 (0.10)
7.97 (0.10)
7.26 (0.10)
8.16 (0.11)
7.44 (0.10)
Standard deviations are given in parentheses.The dissociation constant (pKH) 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 pKH or log K versus 1/T using the graphical representation of van't Hoff equations (3) and (4):
(3)ΔG=-2.303RTlogK=ΔH-TΔS,
or
(4)logK=(-ΔH2.303R)(1T)+(ΔS2.303R).
From the ΔG and ΔH values one can deduce the entropy ΔS using the well-known relationships (3) and (5):
(5)ΔS=(ΔH-ΔG)T.
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)
ThepKH values decrease with increasing temperature; that is, the acidity of the ligand increases [17].(ii)
A positive value ofΔH indicates that the process is endothermic.(iii)
A large positive value ofΔG indicates that the dissociation process is not spontaneous [36].(iv)
A negative value ofΔS is obtained due to the increased order as a result of the solvation process.Table 5
Thermodynamic functions for ML and ML2complexes of ligand (L1) in 10% (by volume) ethanol-water mixture and 0.1 mol·dm−3 KCl.
M
n
+
T/K
Gibbs energy/kJ·mol−1
Enthalpy/kJ·mol−1
Entropy/J·mol−1·K−1
-
Δ
G
1
-
Δ
G
2
Δ
H
1
Δ
H
2
Δ
S
1
Δ
S
2
298
39.03
34.46
237.52
225.21
Mn2+
308
41.39
36.68
31.75
32.65
237.50
225.10
318
43.78
38.97
237.52
225.21
298
39.66
34.98
242.64
223.92
Co2+
308
42.05
37.21
32.64
31.75
242.52
223.91
318
44.51
39.46
242.64
223.92
298
40.06
35.32
243.98
225.07
Ni2+
308
42.46
37.57
32.65
31.75
243.87
225.06
318
44.94
39.82
243.98
225.07
298
40.79
36.18
252.55
230.96
Cu2+
308
43.29
38.45
34.46
32.65
252.43
230.84
318
45.85
40.79
252.55
230.96Table 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.
M
n
+
T/K
Gibbs energy/kJ·mol−1
Enthalpy/kJ·mol−1
Entropy/J·mol−1·K−1
-
Δ
G
1
-
Δ
G
2
Δ
H
1
Δ
H
2
Δ
S
1
Δ
S
2
298
41.82
38.46
243.82
229.45
Mn2+
308
44.23
40.69
30.83
29.92
243.72
229.25
318
46.70
43.05
243.82
229.45
298
43.02
39.09
253.93
237.64
Co2+
308
45.53
41.39
32.65
31.73
253.82
237.44
318
48.01
43.83
253.93
237.65
298
43.54
39.48
252.58
239.05
Ni2+
308
45.99
41.87
31.73
31.75
252.38
239.04
318
48.59
44.27
252.58
239.05
298
44.39
40.34
264.61
248.00
Cu2+
308
47.48
42.81
34.46
33.56
264.50
247.99
318
49.68
45.30
264.62
248.05All the thermodynamic parameters of the stepwise stability constants of complexes are recorded in Tables5 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 (logK1 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  ofΔG for  the  complexation  process  suggests  the spontaneous nature of such processes [38].(iii)
TheΔH values are positive, meaning that these processes are endothermic and favourable at higher temperature.(iv)
TheΔS values for the ligand complexes are positive, confirming that the complex formation is entropically favourable [16].



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