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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148
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
Int. J. Modern Chem. 2015, 7(2): 141-155
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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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Int. J. Modern Chem. 2015, 7(2): 141-155
151
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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Int. J. Modern Chem. 2015, 7(2): 141-155
153
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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