1. Introduction

Schiff bases hydrazone derivatives and their metalcomplexes have been studied for their interesting and important properties, e.g., antibacterial[1,2], antifungal[3],antioxidant[4], anticancer[5] and catalytic activity in oxidation of cyclohexene[6]. Moreover, Schiff bases hydrazonederivatives are versatile ligands and they offer the possibility of different modes of coordination towards transition metal ions. Also, some of these derivatives have been applied as iron chelate or drugs in therapy of anaemia[7] and treatment of neuropathic pain[8].

Therefore it prompted us to study Schiff base transition metal complexes. Transition metal ions have a strong role in bio-inorganic chemistry and redox enzyme systems and may provide the basis of models for active sites of biological systems[9]. Copper (II) ion is a biologically active, essentialon, cleating ability and positive redox potential allow participation in biological transport reactions. Cu(II) complexes possess a wide range of biological activity and are among the most potent antiviral, antitumor and anti inflammatory agents[10]. Schiff base transition metal complexes have been extensively studied because of their potential use as catalysts in a wide range of oxidation reactions[11–14]. In recent years many copper, nickel and manganese complexes of Schiff bases were prepared and characterized by several techniques[15, 16].

2. Objectives

This work deals with the determination of solvation free energies (∆Gs) ,enthalpy changes of solvation(∆Hs) and the entropy of solvation (∆Ss) from solubility measurement and identification of coordination behaviour of Schiff base ligand HLtowards CuCl₂.and the determination of thethermodynamic stability constants and thermodynamic functions using the conductometric technique. Thus, thermodynamic studies of complexation reactions of this Schiff base with transition metal ions not only result in important information on the thermodynamics of complexation reaction, but also lead to a better understanding of the high selectivity of this ligand towards different metal cations.

The aim of this work the evaluation the non-covalent behavior of CuCl₂ with2-hydroxyimino-3-(2'-hydazonopyridyl)-butane (HL) in absolute ethanol solutions at 294.15 K. These non-covalent interactions can help us for analysis of salts in bodies and environnement[17].

Scheme. (1). The outline of the synthesis of 2-hydroxyimino-3-(2'-hydrazonopyridyl)-butane (HL)

3. Methods

3.1. Materials

All manipulations were performed under aerobic conditions. The cupper chloride and the used reagents were Merck pure.

3.2. Preparation of HL

2-hydroxyimino-3-(2'-hydrazonopyridyl)-butane (HL) (scheme 1) was prepared by boiling an EtOH solution of 2-hydrazino pyridine (Aldrich) with 2, 3-butanedione monoxime (1:1) under reflux. The product was recrystallised from hot absolute EtOH[ 18]. (M.p: 220℃; yield 80%). The purity of the compound was checked by TLC.

3.3. Conductometric Measurement

The conductometric titration of the CuCl₂ (1x10^-4) mole/L against the ligand (1x10^-3) mole/L in absolute ethanol was performed with 0.2 ml interval additions from HL solution. The specific conductance values were recorded using conductivity bridge ADWA, AD 3000 with a cell constant equal to 1 cm ^-1. The temperature was adjusted at 293.15 K, 298.15 K, 303.15 K and 308.15 K

3.4. Solubility Measurment

Saturated solutions of HL were prepared by dissolving an excess amount of the solid substances in 10 ml. of the corresponding solvent mixtures, using closed test tubes. The solutions were vigorously shaken in a thermostaticwater-bath at the desired temperature. The molal solubilities of the HL were analysed by drying 1ml. of the saturated solutions in small aluminium dishes. Evaporation of the solvent was performed carefully and slowly under a tungsten lamp to prevent any loss in salt weight. Solubility value was taken as an average of three consecutive independent measurements.

4. Results and Discussion

4.1. Association Constants

The specific conductance values (K_s) of different concentrations of CuCl₂ in absolute ethanol were measured experimentally in absence and in the presence of ligand at different temperatures (293.15 K , 298.15 K , 303.15 K and 308.15 K).

The molar conductance (/\_m) values were calculated[19] using equation (1):

(1) ,

where K_s and K_solv are the specific conductance of the solution and the solvent, respectively; K_cell is the cell constant and C is the molar concentration of the CuCl₂ solutions.

The limiting molar conductances (/\_M) at infinite dilutions were estimated for CuCl₂ in absolute ethanol alone at different temperatures by extrapolating the relation between /\_m and C_m^½ to zero concentration as shown in Fig.(1) .

The limiting molar conductances (/\_o) at infinite dilutions were estimated for CuCl₂ in the presence of the ligand (HL) by extrapolating the relation between /\_m and C_m^½ to zero concentration Fig. (2)

By drawing the relation between molar conductance (/\_m) and the molar ratio of metal to ligand (M/L) concentrations (Fig. (3), (4), (5), (6)), different lines are obtained with sharp breaks indicating the formation of 1:2, 1:1 and 2:1 (M:L) stoichiometric complexes.

Figure (1). The relation between molar conductance (and\m) and (Cm½) of CuCl2 alone in absolute ethanol at different temperatures (293.15K, 298.15 K, 303.15 K and 308.15 K)

Figure (2). The relation between molar conductance (/\m) and (Cm½) of CuCl2 in presence of HL in absolute ethanol at different temperatures (293.15K, 298.15 K, 303.15 K and 308.15 K)

Figure (3). The relation between /\m and M/L (CuCl2-HL) at 293.15 K

Figure (4). The relation between/\m and M/L (CuCl2-HL) at 298.15 K

Figure (5). The relation between /\m and M/L(CuCl2-HL) at 303.15 K

Figure (6). The relation between /\m and M/L (CuCl2-HL) at 308.15K

The experimental data of (/\_M) and (/\_o) were analyzed for the determination of association and formation constants for each type of the stoichiometric complexes.

The association constants of CuCl₂ in the presence of ligand (HL) in absolute ethanol at different temperatures ( 293.15 K , 298.15 K , 303.15K and 308.15 K) for 2:1 ,1:1 and 1:2 (M:L) were calculated by using equation[20,21]:

(2) ,

where (/\_m , /\₀) are the molar and limiting molar conductance of CuCl₂ in presence of Hl respectively; C_m is molar concentration of CuCl₂ , S(Z) is Fuoss-Shedlovsky factor, equal with unity for strong electrolytes[22]. The calculated association constants are shown in Table (1).

4.2. Gibbs Free Energies of Association

The Gibbs free energies of association (ΔG_A) were calculated from the association constant[23, 24] by applying equation:

(3)

where R is the gas constant (8.341 J) and T is the absolute temperature .The calculated Gibbs free energies were presented in Table (2).

4.3. The Formation Constants for Complexes

The formation constants (K_f) for CuCl₂ complexes were calculated for each type of complexes (1:2), (1:1) and (2:1) (M: L)[25, 26] by using equation:

(4)

where /\_M is the limiting molar conductance of the CuCl₂ alone, /\_obs is the molar conductance of solution during titration and /\_ML is the molar conductance of the complex.

The obtained values (K_f) for CuCl₂-ligand stoichiometric complexes are presented in Table (3)

Table (1). Association constants of CuCl2 with HL at different temperatures (293.15 K, 298.15 K, 303.15 K and 308.15 K) KA Cm 293.15 K 298.15 K 303.15 K 308.15 K 9.52E-05 3.93E+01 4.63E+01 5.33E+01 5.60E+01 9.38E-05 3.66E+01 4.31E+01 4.94E+01 5.19E+01 9.26E-05 3.46E+01 4.09E+01 4.63E+01 4.84E+01 9.09E-05 3.18E+01 3.67E+01 4.20E+01 4.38E+01 8.88E-05 2.89E+01 3.31E+01 3.77E+01 3.88E+01 8.67E-05 2.61E+01 3.00E+01 3.40E+01 3.51E+01 8.47E-05 2.41E+01 2.72E+01 3.07E+01 3.19E+01 8.33E-05 2.27E+01 2.54E+01 2.84E+01 2.96E+01 8.15E-05 2.11E+01 2.32E+01 2.59E+01 2.71E+01 7.98E-05 1.96E+01 2.17E+01 2.40E+01 2.50E+01

Table (2). Gibbs free energies of association of CuCl2 with HL at different temperatures (293.15 K, 298.15 K, 303.15 K and 308.15 K) Δ GA(k J/mol) Cm 293.15 K 298.15 K 303.15 K 308.15 K 9.52E-05 -8.95E+00 -9.51E+00 -1.00E+01 -1.03E+01 9.38E-05 -8.78E+00 -9.33E+00 -9.83E+00 -1.01E+01 9.26E-05 -8.64E+00 -9.20E+00 -9.66E+00 -9.94E+00 9.09E-05 -8.43E+00 -8.93E+00 -9.42E+00 -9.68E+00 8.88E-05 -8.19E+00 -8.68E+00 -9.15E+00 -9.37E+00 8.67E-05 -8.20E+00 -8.43E+00 -8.89E+00 -9.12E+00 8.47E-05 -7.75E+00 -8.19E+00 -8.63E+00 -8.87E+00 8.33E-05 -7.61E+00 -8.02E+00 -8.44E+00 -8.68E+00 8.15E-05 -7.43E+00 -7.79E+00 -8.20E+00 -8.45E+00 7.98E-05 -7.25E+00 -7.62E+00 -8.01E+00 -8.25E+00

Table (3). Formation constants for 1:2, 1:1 and 2:1 (M/L) complexes in absolute ethanol at different temperatures (293.15 K, 298.15 K, 303.15 K and 308.15 K) Kf 1 : 2 (M/L) Cm 293.15 K 298.15 K 303.15 K 308.15 K 8.33E-05 6.78E+05 9.16E+05 2.30E+06 3.50E+06 8.24E-05 3.31E+05 3.74E+05 5.26E+05 6.04E+05 8.06E-05 1.48E+05 1.49E+05 1.60E+05 1.77E+05 7.98E-05 1.15E+05 1.18E+05 1.21E+05 1.27E+05 7.89E-05 8.95E+04 9.34E+04 9.83E+04 9.75E+04 1 : 1 (M/L) Cm 293.15 K 298.15 K 303.15 K 308.15 K 9.09E-05 1.31E+06 2.81E+06 6.92E+06 1.56E+07 8.98E-05 6.00E+05 8.19E+05 9.62E+05 1.02E+06 8.88E-05 3.29E+05 3.90E+05 4.08E+05 4.19E+05 8.77E-05 2.13E+05 2.39E+05 2.40E+05 2.42E+05 8.67E-05 1.56E+05 1.77E+05 1.75E+05 1.79E+05 2 : 1 (M/L) Cm 293.15 K 298.15 K 303.15 K 308.15 K 9.52E-05 7.50E+06 1.49E+07 2.92E+07 1.73E+08 9.51E-05 3.28E+06 5.49E+06 1.54E+07 4.17E+07 9.49E-05 2.56E+06 3.12E+06 8.21E+06 1.63E+07 9.38E-05 1.04E+06 1.22E+06 2.04E+06 1.27E+06 9.26E-05 5.23E+05 6.06E+05 8.88E+05 5.57E+05

4.4. Gibbs Free Energies of Complex Formation

The Gibbs free energies of formation for eachstoichiometric complexes were calculated by using the equation:

(5)

The calculated ΔG_f values are presented in Table (4).

4.5. Enthalpies and Entropies

The enthalpy (ΔH_A) for CuCl₂ complexes were calculated for each type of complexes (1:2) , (1:1) and (2:1) (M:L) by using Van't Hoff equation[26,27] :

(6)

Where R is the gas constant and T is the absolute temperature. By drawing the relation between log K_A and 1/T, different lines are obtained indicating the formation of 1:2,1:1 and 2:1 (M:L) stoichiometric complexes Fig.(7).

4.6. Enthalpies and Entropies of Association

From the relation between log K and 1/T, ΔH_A can be calculated for each type of complexes from the slope of each line (-ΔH/2.303 R).The entropy (ΔS_A) for CuCl₂ complexes were calculated for each type of complexes (1:2), (1:1) and (2:1) (M:L) by using equation :

(7)

Where (S) is the entropy of the system.

The calculated values of (ΔH_A) and (ΔS_A) forCuCl₂-ligand stoichiometric complexes are presented in Table (5):

By drawing the relation between log K_f and 1/T, different lines are obtained indicating the formation of 1:2,1:1 and 2:1 (M:L) stoichiometric complexes Fig.(8).

4.7. Enthapies and Entropies of Complex Formation

The enthalpy (ΔH_f) for CuCl₂ complexes were calculated for each type of complexes (1:2) , (1:1) and (2:1) (M:L) by using van 't Hoff equation .

The calculated values of (ΔH_f) and (ΔS_f) for CuCl₂-ligand stoichiometric complexes are presented in Table (6):

4.8. Acivation Energies

Since the conductance of an ion depends mainly on its mobility, it is quite reasonable to treat the rate process taking place with the change of temperature on the basis of equation (8) :

(8)

where A is the frequency factor, R is the gas constant and Ea is the Arrhenius activation energy of the transfer process. Consequently, from the plot of log /\₀ vs. 1/T, the E_a values can be evaluated[ 27] as shown in Fig (9) .

E_a value is 14.6996 KJ/mol.

4.9. Soubility Measurement

The solubility (S) of2-hydroxyimino-3-(2'-hydazonopyrridyl)-butane (HL) in (EtOH-H₂O) mixtures at different temperatures (293.15, 298.15, 303.15 and 308.15 K) was determined by gravimetric technique. The results are illustrated in Table 1. Solubility was calculated as an average of the two experimental data. The molal solubility is calculated by using equation (9):

Molal solubility (S) = W.1000/do.M g.mole /1000 g .

(9)

where (W) is the weight of one ml. of saturated solution, after its complete evaporation in the aluminum dish under the effect of tungsten lamp,(M) is the molecular weight of HL and (d_o) is the density of pure solvent used as it shown in Table (7) ; Fig.(10) the molal solubility was increased with the increase of the content of the organic solvent used (EtOH).This can be explained on the basis of the fact that like dissolve like as well as the lower and higher ion-solvent interactions. The molal solubility of HL was increased with the increase of temperatures

Table (4). Gibbs free energies of formation of CuCl2 with HL at different temperatures (293.15 K, 298.15 K, 303.15 K and 308.15 K) ΔGf(k J/mol) 1 : 2 (M/L) Cm 293.15 K 298.15 K 303.15 K 308.15 K 8.33E-05 -32.7243 -34.0278 -36.9243 -38.6072 8.24E-05 -30.9752 -31.8059 -33.2027 -34.1017 8.06E-05 -29.0121 -29.5208 -30.1944 -30.9637 7.98E-05 -28.3917 -28.9411 -29.4929 -30.1157 7.89E-05 -27.7900 -28.3702 -28.9740 -29.4305 1 : 1 (M/L) Cm 293.15 K 298.15 K 303.15 K 308.15 K 9.09E-05 -35.2606 -36.8031 -39.6971 -42.4394 8.98E-05 -33.1527 -33.7523 -34.7234 -35.4580 8.88E-05 -31.5954 -31.9121 -32.5605 -33.1690 8.77E-05 -30.4886 -30.6951 -31.2246 -31.7554 8.67E-05 -29.6959 -29.9470 -30.4239 -30.9824 2 : 1 (M/L) Cm 293.15 K 298.15 K 303.15 K 308.15 K 9.52E-05 -38.5812 -40.9373 -43.3258 -48.5997 9.51E-05 -36.5700 -38.4688 -41.7124 -44.9520 9.49E-05 -35.9583 -37.0661 -40.1272 -42.5507 9.38E-05 -33.7636 -34.7418 -36.6154 -36.0169 9.26E-05 -32.0900 -33.0054 -34.5210 -33.8937

Figure (7). The relation between (log KA) and (1/T)

Table (5). The enthalpies and entropies of association of complexes at different temperatures (293.15 K, 298.15 K, 303.15 K and 308.15 K) ΔHA(KJ/mol) and ΔSA(KJ/mol.K) 1 : 2 (M/L) Temp 293.15 K 298.15 K 303.15 K 308.15 K ΔHA 13.8326 13.8326 13.8326 13.8326 ΔSA 0.07316 0.07329 0.07347 0.07305 1 : 1 (M/L) Temp 293.15 K 298.15 K 303.15 K 308.15 K ΔHA 16.5043 16.5043 16.5043 16.5043 ΔSA 0.08505 0.08530 0.08551 0.08497 2 : 1 (M/L) Temp 293.15 K 298.15 K 303.15 K 308.15 K ΔHA 18.1299 18.1299 18.1299 18.1299 ΔSA 0.09236 0.09270 0.09279 0.09138

Figure (8). The relation between (log Kf) and (1/T)

Table (6). The enthalpies and entropies of formation of complexes at different temperatures (293.15 K, 298.15 K, 303.15 K and 308.15 K) ΔHf (KJ/mol) and ΔSf(KJ/mol.K) 1 : 2 (M/L) Temp 293.15K 298.15 K 303.15 K 308.15 K ΔHf 87.7317 87.7317 87.7317 87.7317 ΔSf 0.4109 0.4083 0.4112 0.4099 1 : 1 (M/L) Temp 293.15 K 298.15 K 303.15 K 308.15 K ΔHf 123.0207 123.0207 123.0207 123.0207 ΔSf 0.5399 0.5360 0.5367 0.0536 2 : 1 (M/L) Temp 293.15 K 298.15 K 303.15 K 308.15 K ΔHf 162.4927 162.4927 162.4927 162.4927 ΔSf 0.6859 0.6823 0.6789 0.6850

Figure (9). The relation of (log /\0) and 1/T

Table (7). The Molal solubility (S) of HL at different temperatures (293.15 K, 298.15 K, 303.15K and 308.15 K) Vol. % ofEtOH (S) of HL 293.15 K 298.15 K 303.15 K 308.15 K 0 0.0005 0.0007 0.0009 0.0011 10 0.0011 0.0015 0.0019 0.0021 20 0.0016 0.0019 0.0025 0.0030 30 0.0018 0.0024 0.0029 0.0033 40 0.0032 0.0044 0.0051 0.0057 50 0.0047 0.0063 0.0072 0.0081 60 0.0069 0.0076 0.0089 0.0115 70 0.0114 0.0125 0.0148 0.0179 80 0.0179 0.0199 0.0221 0.0246 90 0.0201 0.0229 0.0248 0.0275 100 0.0257 0.0278 0.0311 0.0339

Figure (10). Variation of the molal solubility (S) of HL with the mole fraction (Xs) of EtOH at different temperatures

Table (8). The solvation free energies (∆G) S of HL in EtOH-H2O mixture at different temperatures (293.15 K, 298.15 K, 303.15Kand 308.15 K) Vol. % ofEtOH (ΔG)S(KJ/mol) (ΔH)s(KJ/mol) 293.15 K 298.15K 303.15K 308.15K 0 18.5286 18.0104 17.6789 17.7005 35.1255 10 16.6066 16.1209 15.7953 15.7994 32.7945 20 15.6932 15.5348 15.1035 14.8854 32.4435 30 15.3393 14.9045 14.7294 14.6412 28.6884 40 13.9657 13.4528 13.3063 13.2407 27.6306 50 13.0406 12.5629 12.4370 12.3403 26.1393 60 12.1305 12.0978 11.9026 11.4422 25.3114 70 10.9065 10.8642 10.6206 10.3085 22.8197 80 9.8067 9.7113 9.6099 9.4937 15.9020 90 9.5241 9.3632 9.3193 9.2082 15.3366 100 8.9250 8.8825 8.7487 8.6721 14.1599

4.10. Thermodynamics of Solvation

The solvation free energies ∆G_S of HL in EtOH-H₂O mixture at different temperatures(293.15 K,298.15 K, 303.15 K and 308.15 K) were calculated from the solubility measurements by using the following equation (10):

(10) .

The value of (log K_sp) depends mainly on the solvation of the solute in the solvent under investigation .In case of neutral compound (the activity coefficient is close to one), the values of (log K_sp) can be equal to log (S).

The enthalpy changes of solvation(∆H_s) of HL in EtOH-H₂O mixtures were calculated from the plots of (log K_sp) versus (1/T) ,where the slope equals (-∆H_s/2.303 R) using the following equation(11):

(11)

5. Conclusions

The stability constants for the complexation of copper (II) ion with 2-hydroxyimino-3-(2'-hydazonopyridyl)-butane (HL) were determined conductometrically at different temperatures. Thermodynamic parameters of complexation were determined from the temperature dependence of the formation constant. The negative values of ∆G show the ability of the studied ligand to form stable complexes and the process trend to proceed spontaneously. However, the obtained positive values of ∆H means that enthalpy is not the driving force for the formation of the complexes. Furthermore, the positive values of ∆S indicate that entropy is responsible for the complexing process. The formation constants and Gibbs free energies of different complexes follow that order: K_f (2:1) › K_f (1:1) › K_f (1:2) for (M:L), and ∆G_f (2:1) › ∆G_f (1:1) › ∆G_f (1:2) for (M:L)