Complexation of 2-Mercaptoimidazol with Some

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Nov 13, 2015 - Also it used in treatment of thyrotoxicosis, 1-methyl-2-mercapto-imidazole ... 2 mo l-1. ) [M]/[L]. Fig. 1: The relation between molar conductance ( m ... in 50% MeOH-H2O mixed solvent at different temperatures. ... Δ Gf (kJ mol-1).

International Journal of Modern Chemistry, 2015, 7(2): 141-155 International Journal of Modern Chemistry ISSN: 21650128 Journal homepage: www.ModernScientificPress.com/Journals/IJMChem.aspx Florida, USA Article

Complexation of 2-Mercaptoimidazol with Some Barium Salts Conductometrically

in

Various

Solvent

at

Different

Temperatures Elsayed T. Helmy*1,2, Esam A. Gomaa2 and Elsayed M. Abou Eleef 3 1

Marine Pollution Laboratory, National Institute of Oceanography and Fisheries, Ministry of

Scientific Research, Alexandria, Egypt 2

Chemistry Department, Faculty of Science, Mansoura University, 35516-Mansoura, Egypt

3

Basic Science Department, Delta Higher Institute for Engineering & Technology, 35681- Dakhlia,

Mansoura, Egypt * Author to whom correspondence should be addressed; E-Mail: [email protected] Article history: Received 21 August 2015, Revised 5 November 2015, Accepted 11 November 2015, Published 13 November 2015.

Abstract: The complexation reaction between 2-meracptoimidazol (MI) ligand with BaCl2 and Ba(NO3)2 salts were studied conductometrically in methanol, 50% methanol-water and water at four different temperatures (293.15K, 298.15K, 303.15K, 308.15K ). On drawing the relation between molar conductance and the ratio of metal to ligand concentrations, different lines are obtained indicating formation of 1:1 and 2:1 [M]/[L] stoichiometric complexes. Limiting molar conductance, Activation energy, Formation constants and thermodynamic parameters of different complexes were determined. Negative values of Gibbs free energy of complexation indicate that the reaction is spontaneous . Keywords: Conductometric studies, Formation constant, Activation energy, Limiting molar conductance, Thermodynamic parameters, 2-mercaptoimidazol.

1. Introduction Conductance measurement is an important and simplest tool to get information in regard to ionsolvation [1-3]. The study of conductance of an electrolyte system with different solvents of varying Copyright © 2012 by Modern Scientific Press Company, Florida, USA

Int. J. Modern Chem. 2015, 7(2): 141-155

142

viscosity and dielectric constant provides satisfactory information regarding ion-solvation [4-7]. In recent years there has been an increasing interest in the study of behavior of electrolytes in partial and non-aqueous solvents with a view of investigating the role of partial or non- aqueous solvents in the solvation of ion [8-12]. 2-mercapto-imidazole derivatives substituted in position 4 (or5) used as antioxidizing agents, in preparation and applications in the pharmaceutical, cosmetic or food industries [13]. Also it used in treatment of thyrotoxicosis, 1-methyl-2-mercapto-imidazole used in treatment of hyperthyroidism. Because there are lake of thermodynamic data for 2-mercaptoimidazol [21], we study complexation of 2-mercaptoimidazol with barium nitrate and barium chloride[12, 21]. All thermodynamic parameter were investigated.

2. Materials and Methods 2.1. Materials All manipulation was performed under aerobic conditions. The barium chloride, barium nitrate and 2-mercaptoimidazol (MI) were purchased from Merck pure and used without any further purification. 2.2. Conductometric Titrations The Conductometric titration of the BaCl2 or Ba(NO3)2 (1x10-3) mole/L against MI (1x10-4) mole/L in water and methanol for BaCl2 while for Ba(NO3)2 measured in water and 50% MeOH-H2O was performed with 0.5 ml interval additions from BaCl2 or Ba(NO3)2 solution. The specific conductance values were recorded using conductivity bridge HANNA, H1 8819N with a cell constant equal to 1 cm1

. The conductometer was connected to the type Kottermann 4130 ultra thermostat. The temperature was

adjusted at (293.15K, 298.15K, 303.15K, and 308.15K).

3. Results and Discussion The specific conductance values (Ks) of different concentrations of BaCl2 or Ba(NO3)2 in used solvents mentioned above were measured experimentally in the presence of MI at (293.15K, 298.15K, 303.15K, 308.15K ) The molar conductance (Λm) values were calculated [14-28] using equation(1):

m 

(K s  K solv )K cell  1000 C

(1) where Ks and Ksolv are the specific conductance of the solution and the solvent, respectively; Kcell is the cell constant and C is the molar concentration of the BaCl2 or Ba(NO3)2 solution. The limiting molar conductance (Λo) at infinite dilutions were estimated for each salt in the presence of the ligand (MI) by extrapolating the relation between Λm and √Cm to zero concentration .By drawing the relation between Copyright © 2012 by Modern Scientific Press Company, Florida, USA

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molar conductance (Λm) and the molar ratio of metal to ligand [M]/[L] concentrations (Figs 1-4), Different lines are obtained with sharp breaks indicating the formation of 1:1 and 2:1 [M]/[L] stoichiometric complexes. The experimental data of (Λm) and (Λo) were analyzed for the determination of formation constants for each type of the stoichiometric complexes.

293.15 k 298.15 k 303.15 k 308.15 k

480 460 440 420

Molar conductance -1 2 m (S cm mol )

400 380 360 340 320 300 280 260 240 220 200 180 0

1

2

3

4

5

[M]/[L]

Fig. 1: The relation between molar conductance (  m ) and the [M]/[L] molar ratio of BaCl2 to MI in water at different temperatures.

293.15 k 298.15 k 303.15 k 308.15 k

260 240

Molar conductance -1 2 m (S cm mol )

220 200 180 160 140 120 0

1

2

3

4

5

[M]/[L]

Fig. 2: The relation between molar conductance (  m ) and the [M]/[L] molar ratio of BaCl2 to MI in methanol at different temperatures.

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Int. J. Modern Chem. 2015, 7(2): 141-155

144 293.15 K 298.15 K 303.15 K 308.15 K

900 800

Molar conductance -1 2 m (S cm mol )

700 600 500 400 300 200 100 0

1

2

3

4

5

[M]/[L]

Fig. 3: The relation between molar conductance (  m ) and the [M]/[L] molar ratio of Ba(NO3)2 to MI in water at different temperatures.

293.15 k 298.15 k 303.15 k 308.15 k

500 450

Molar conductance -1 2 m (S cm mol )

400 350 300 250 200 150 100 0

1

2

3

4

5

[M]/[L]

Fig. 4: The relation between molar conductance (  m ) and the [M]/[L] molar ratio of Ba(NO3)2 to MI in 50% MeOH-H2O mixed solvent at different temperatures. The formation constants (Kf) for CaCl2 and BaCl2 complexes were calculated for each type of complexes 1:1 and 2:1 [M]/[L] [29-31]by using equation(2): [𝑀𝐿]

𝐾𝑓 = [𝑀][𝐿] =

𝛬𝑀− 𝛬𝑜𝑏𝑠 (𝛬𝑜𝑏𝑠− 𝛬𝑀𝐿) [𝐿]

(2)

where Λm is the limiting molar conductance of the CaCl2 or BaCl2 alone, Λobs. is the molar conductance of solution during titration and ΛML is the molar conductance of the complex. The obtained values (Kf) for BaCl2 -MI and Ba(NO3)2 -MI stoichiometric complexes are presented in Tables (1-8)

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The Gibbs of free energy change of association (  Gf) of different ratios of the metal to ligand formed were calculated [32, 33] from the formation constant (Kf)(tables 1-8) by using equation(3). Gf = - 2.303 RT log Kf

(3)

where R is the gas constant (8.341 J) and T is the absolute temperature.

Table 1: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for Ba(NO3)2 and MI complex 1:1 [M]/[L] formation in water at different temperatures Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

922.37

320.91

2.666E+04

-24.8375

298.15 K

1277.48

327.32

6.368E+04

-27.4199

303.15 K

1372.98

362.40

8.336E+04

-28.5584

308.15 K

1391.24

368.97

6.137E+04

-28.2446

Table 2: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for Ba(NO3)2 and MI complex 2:1 [M]/[L] formation in water at different temperatures. Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

922.37

221.64

3.073E+04

-25.1842

298.15 K

1277.48

229.63

4.571E+04

-26.598

303.15 K

1372.98

240.12

4.815E+04

-27.1752

308.15 K

1391.24

247.61

1.835E+05

-31.0509

It was observed that complexation of Ba(NO3)2 and MI in water, 303.15 K is the best condition for formation of 1:1 [M]/[L] , while 303.15K is the best for formation of 2:1 [M]/[L], Also formation constant values increases by increasing temperatures. Table 3: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for Ba(NO3)2 and MI complex 1:1 [M]/[L] formation in 50% MeOH-H2O mixed solvent at different temperatures Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

730.75

237.45

5.430E+04

-26.5714

298.15 K

748.17

241.76

4.202E+04

-26.389

303.15 K

785.32

246.82

4.000E+04

-26.7075

308.15 K 843.87 253.55 4.942E+06 -39.4884 Table 4: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for Copyright © 2012 by Modern Scientific Press Company, Florida, USA

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Ba(NO3)2 and MI complex 2:1 [M]/[L] formation in 50% MeOH-H2O mixed solvent at different temperatures Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

730.75

163.71

3.046E+04

-25.1625

298.15 K

748.17

173.74

3.120E+04

-25.6513

303.15 K

785.32

179.14

3.090E+04

-26.0573

308.15 K

843.87

182.33

3.354E+04

-26.6969

It was observed that complexation of Ba(NO3)2 and MI in 50% MeOH-H2O, 293.15 K is the best condition for formation of 1:1 [M]/[L]and the formation constant values decreases by increasing temperatures. While 308.15K is the best for formation of 2:1 [M]/[L]. Also formation constant values increases by increasing temperatures. Using Ba(NO3)2 complex formation is favoured in water than in 50% MeOH-H2O and 1:1 [M]/[L] is more stable. Table 5: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for BaCl2 and MI complex 1:1 [M]/[L] formation in water at different temperatures Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

247.53

131.84

4.143E+04

-25.9121

298.15 K

207.59

142.45

4.740E+04

-26.6881

303.15 K

314.11

153.69

6.708E+04

-28.0106

308.15 K

395.81

163.13

5.012E+04

-27.7258

Table 6: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for BaCl2 and MI complex 2:1 [M]/[L] formation in water at different temperatures Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

424.83

222.51

3.429E+04

-25.4509

298.15 K

207.59

122.18

3.307E+04

-25.7957

303.15 K

314.11

128.39

3.010E+04

-25.9906

308.15 K

395.81

133.24

2.661E+04

-26.1036

It was observed that complexation of BaCl2 and MI in water, 303.15 K is the best condition for formation of 1:1 [M]/[L], While 293.15K is the best for formation of 2:1 [M]/[L] and the formation constant values decreases by increasing temperatures. Table 7: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for

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BaCl2 and MI complex 1:1 [M]/[L] formation in Methanol at different temperatures Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

187.91

136.97

2.266E+04

-24.4414

298.15 K

235.35

149.07

5.860E+04

-27.2138

303.15 K

278.25

175.68

4.357E+04

-26.9233

308.15 K

339.28

186.77

2.486E+04

-25.9297

Table 8: Limiting molar conductance (Λo), Formation constant (Kf), Gibbs free energy change (ΔGf) for BaCl2 and MI complex 2:1 [M]/[L] formation in Methanol at different temperatures Temp. K

Λo (S cm2.mol-1)

Λobs. (S cm2.mol-1)

Kf

Δ Gf (kJ mol-1)

293.15 K

187.91

131.11

2.112E+04

-24.2706

298.15 K

235.35

135.22

2.367E+04

-24.9661

303.15 K

278.25

152.44

2.137E+04

-25.1278

308.15 K

339.28

156.02

2.287E+04

-25.7158

It was observed that complexation of BaCl2 and MI in MeOH, 298.15 K is the best condition for formation of both complex ratio and it was found that complex formation is favoured using water than MeOH. The previous tables shows that decreasing in Gibbs free energy of complexaion by increasing in temperatures indicating more spontaneous process also the values is case of 1:1 [M]/[L] is smaller than in case of 2:1 [M]/[L] for each salt used . While limiting molar conductance increases by increasing temperatures and values for 1:1 [M]/[L] is higher than that in case of 2:1 [M]/[L] . The enthalpy (ΔHf) for the complexes formed were calculated for each type of complexes, 1:1 and 2:1 [M]/[L] by using van 't Hoff equation (4): 𝑑𝑙𝑛𝐾 𝑑𝑇

=

o Δ𝐻A

𝑅𝑇 2

(4)

where R is the gas constant (8.341 J) and T is the absolute temperature. By drawing the relation between log Kf and 1/T different lines are obtained for the formation of 1:1 and 2:1 [M]/[L] stoichiometric complexes for, BaCl2 and Ba(NO3)2 with 2-mercaptoimidazol. From the relation between log Kf and 1/T, figs.(5-8), ΔHf can be calculated for each type of complexes from the slope of each line (-ΔHf/2.303R).The entropy (ΔSf) for complexes were calculated for each type of complexes 1:1 and 2:1[M]/[L] tables (9-16)by using equation (5): ΔGf = ΔHf – TΔSf

(5)

where (S) is the entropy of system.

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1:1 M:L 2:1 M:L

5.3 5.2 5.1 5.0

log Kf

4.9 4.8 4.7 4.6 4.5 4.4 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 5: The relation between (log Kf) and (1/T) for 1:1 and 2:1 [M]/[L] stoichiometric complexes of Ba(NO3)2 and MI in water.

1:1 M:L 2:1 M:L 4.70 4.65

log Kf

4.60 4.55 4.50 4.45 4.40 4.35 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 6: The relation between (log Kf) and (1/T) for 1:1 and 2:1[M]/[L] stoichiometric complexes of Ba(NO3)2 and MI in 50% MeOH-H2O mixed solvent

1:1 M:L 2:1 M:L

4.9

4.8

log Kf

4.7

4.6

4.5

4.4

4.3 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 7: The relation between (log Kf) and (1/T) for1:1 and 2:1 [M]/[L] stoichiometric complexes of BaCl2 and MI in water. Copyright © 2012 by Modern Scientific Press Company, Florida, USA

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1:1 M:L 2:1 M:L

5.1 5.0 4.9

log Kf

4.8 4.7 4.6 4.5 4.4 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 8: The relation between (log Kf) and (1/T) for 1:1 and 2:1 [M]/[L] stoichiometric complexes of BaCl2 and MI in methanol.

Table 9: Enthalpy change (ΔHf ) and entropy change (ΔSf) for Ba(NO3)2 and MI complex 1:1 [M]/[L] formation in water at different temperatures Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

69.3247

94.1622

3.21E-01

298.15K

69.3247

96.7446

3.24E-01

303.15K

69.3247

97.8831

3.23E-01

308.15K

69.3247

97.5693

3.17E-01

Table 10: Enthalpy change (ΔHf) and entropy change (ΔSf) for Ba(NO3)2 and MI complex 2:1 [M]/[L] formation in water at different temperatures Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

42.2055

67.3898

2.29E-01

298.15K

42.2055

68.8035

2.31E-01

303.15K

42.2055

69.3807

2.29E-01

308.15K

42.2055

73.2565

2.38E-01

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Table 11: Enthalpy change (ΔHf) and entropy change (ΔSf) for Ba(NO3)2 and MI complex 1:1 [M]/[L] formation in 50% MeOH-H2O mixed solvent at different temperatures Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

45.171

71.7423

2.447E-01

298.15K

45.171

71.5623

2.400E-01

303.15K

45.171

71.8784

2.371E-01

308.15K

45.171

84.6594

2.747E-01

Table 12: Enthalpy change (ΔHf) and entropy change (ΔSf) for Ba(NO3)2 and MI complex 2:1 [M]/[L] formation in 50% MeOH-H2O mixed solvent at different temperatures. Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

65.6611

90.8235

3.098E-01

298.15K

65.6611

91.3123

3.063E-01

303.15K

65.6611

91.7184

3.026E-01

308.15K

65.6611

92.3580

2.997E-01

Table 13: Enthalpy change (ΔHf) and entropy change (ΔSf) for BaCl2 and MI complex 1:1 [M]/[L] formation in water at different temperatures. Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

15.7582

41.6704

1.421E-01

298.15K

15.7582

42.4463

1.424E-01

303.15K

15.7582

43.7688

1.444E-01

308.15K

15.7582

43.4840

1.411E-01

Table 14: Enthalpy change (ΔHf) and entropy change (ΔSf) for BaCl2 and MI complex 2:1 [M]/[L] formation in water at different temperatures. Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

29.7445

55.1954

1.883E-01

298.15K

29.7445

55.5402

1.863E-01

303.15K

29.7445

55.7351

1.839E-01

308.15K

29.7445

55.8481

1.812E-01

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Table 15: Enthalpy change (ΔHf) and entropy change (ΔSf) for BaCl2 and MI complex 1:1 [M]/[L] formation in methanol at different temperatures. Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

44.2969

68.7383

2.345E-01

298.15K

44.2969

71.5107

2.398E-01

303.15K

44.2969

71.2202

2.349E-01

308.15K

44.2969

70.2266

2.279E-01

Table 16: Enthalpy change (ΔHf) and entropy change (ΔSf) for BaCl2 and MI complex 2:1 [M]/[L] formation in methanol at different temperatures. Temperature

ΔHa kJ.mol-1

TΔSa kJ.mol-1

Δ Sa kJ.mol-1

293.15K

38.5389

62.8094

2.143E-01

298.15K

38.5389

63.5050

2.130E-01

303.15K

38.5389

63.6667

2.100E-01

308.15K

38.5389

64.2547

2.085E-01

Since the conductance of an ion depends mainly on its mobility, It’s quite reasonable to treat the conductance data similar to the one that employed for the rate process taking place with the change of temperature on the basis of equation (6). Λ0 =A e-Ea/RT

(6)

where A is the frequency factor, (R) is the gas constant and (Ea) is the Arrhenius activation energy of transport process. Therefore, on plotting of log ( o ) versus 1/T figs. (9-12) straight lines are obtained from their slopes the (Ea) values are evaluated and represented in Tables (17-20) 3.15

3.14

3.13

log 

3.12

3.11

3.10

3.09

3.08 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 9: The relation between (log Λo) and (1/T) for 1:1[M]/[L] stoichiometric complexes of Ba(NO3)2 and MI in water. Copyright © 2012 by Modern Scientific Press Company, Florida, USA

Int. J. Modern Chem. 2015, 7(2): 141-155

152

2.93 2.92 2.91

log 

2.90 2.89 2.88 2.87 2.86 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 10: The relation between (log Λo) and (1/T) for 1:1[M]/[L] stoichiometric complexes of Ba(NO3)2 and MI in 50% MeOH-H2O mixed solvent.

2.60

log 

2.55

2.50

2.45

2.40

2.35 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 11: The relation between (log Λo) and (1/T) for 2:1 [M]/[L] stoichiometric complexes of BaCl2 and MI in water

2.45 2.40

log 

2.35 2.30 2.25 2.20 2.15 2.10 0.00325

0.00330

0.00335

0.00340

-1

1/T, K

Fig. 12: The relation between (log Λo) and (1/T) for 2:1 [M]/[L] stoichiometric complexes of BaCl2 and MI in methanol.

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Table 17: Activation energy (Ea), Gibbs free energy change (ΔG*) and entropy change (ΔS*) of activation for Ba(NO3)2 and MI complex 1:1 [M]/[L] formation in water at different temperatures. Temp.

Ea (kJ mol-1)

ΔG* (kJ mol-1)

TΔS* (kJ mol-1)

ΔS* (kJ mol-1K-1)

293.15 K

6.6949

-16.6390

23.3339

7.960E-02

298.15 K

6.6949

-17.7301

24.4251

8.192E-02

303.15 K

6.6949

-18.2092

24.9041

8.215E-02

308.15 K

6.6949

-18.5433

25.2383

8.190E-02

Table 18: Activation energy (Ea), Gibbs free energy change (ΔG*) and entropy change (ΔS*) of activation for Ba(NO3)2 and MI complex 1:1 [M]/[L] formation in 50% MeOH-H2O mixed solvent at different temperatures. Temp.

Ea (kJ mol-1)

ΔG* (kJ mol-1)

TΔS* (kJ mol-1)

ΔS* (kJ mol-1K-1)

293.15 K

7.1292

-16.0714

23.2006

7.914E-02

298.15 K

7.1292

-16.4039

23.5331

7.893E-02

303.15 K

7.1292

-16.8011

23.9303

7.894E-02

308.15 K

7.1292

-17.2625

24.3917

7.916E-02

Table 19: Activation energy (Ea), Gibbs free energy change (ΔG*) and entropy change (ΔS*) of activation for BaCl2 and MI complex 1:1 [M]/[L] formation in water at different temperatures. Temp.

Ea (kJ mol-1)

ΔG* (kJ mol-1)

TΔS* (kJ mol-1)

ΔS* (kJ mol-1K-1)

293.15 K

-14.7495

24.0322

-14.7495

38.7817

298.15 K

-15.5037

24.0322

-15.5037

39.5359

303.15 K

-16.0876

24.0322

-16.0876

40.1198

308.15 K

-16.5788

24.0322

-16.5788

40.6110

Table 20: Activation energy (Ea), Gibbs free energy change (ΔG*) and entropy change (ΔS*) of activation for BaCl2 and MI complex 1:1 [M]/[L] formation in methanol at different temperatures. Temp.

Ea (kJ mol-1)

ΔG* (kJ mol-1)

TΔS* (kJ mol-1)

ΔS* (kJ mol-1K-1)

293.15 K

17.7946

-12.7613

30.5560

1.042E-01

298.15 K

17.7946

-13.537

31.3316

1.051E-01

303.15 K

17.7946

-14.1861

31.9807

1.055E-01

308.15 K

17.7946

-14.9281

32.7227

1.062E-01

A perusal of tables shows that the values of Ea are positive. This indicates higher mobilites of ions in solutions and hence Λo increases with increasing temperatures.

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Int. J. Modern Chem. 2015, 7(2): 141-155

154

4. Conclusions We study the complexation reaction between 2-mercaptoimidazol

with BaCl2 and Ba(NO3)2

condctometrically in water ,methanol and 50% MeOH-H2O at four different temperatures (293.15K, 298.15K, 303.15K, 308.15K ). On drawing the relation between molar conductance and the ratio of metal to ligand concentrations, Different lines are obtained indicating the formation of 1:1 and 2:1 [M]/[L] stoichiometric complexes. The formation constants and Gibbs free energies of different complexes were determined. Negative values of Gibbs free energy of complexation indicate that the reaction is spontaneous. It’s shown that formation constant increases by increasing temperature. The decreasing in Gibbs free energy of complexaion by increasing in temperatures indicating more spontaneous process also this values is case of 1:1 [M]/[L] is smaller than in case of 2:1 [M]/[L] for each salt used . Also limiting molar conductance increases by increasing temperatures and its values for 1:1 [M]/[L] is higher than that in case of 2:1 [M]/[L]. The positive values of activation energy indicates higher mobility of ions in solutions and hence Λo increases with increasing temperatures.

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