Journal of Chemistry

Volume 2013 (2013), Article ID 858374, 10 pages

http://dx.doi.org/10.1155/2013/858374

Research Article

The Audio Frequency Conductance Study of Some Metal Succinate Salts in Aqueous Medium at Different Temperatures (Part I: Magnesium, Manganese (II), Barium and Copper Succinates)

Kosrat N. Kaka,¹ Anis A. Al-Najar,² and Wali M. Hamad³

¹Department of Chemistry, College of Science, University of Koya, Koya Town, Iraq

²Department of Chemistry, College of Science, University of Basrah, Basrah City, Iraq

³Koya Technical Institute, Department of oil Refinery, Koya Town, Iraq

Received 5 January 2012; Revised 25 June 2012; Accepted 26 June 2012

Academic Editor: Jose Corchado

Copyright © 2013 Kosrat N. Kaka et al. This is an open access article distributed under the Creative Commons Attribution License, which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

Abstract.

The audio electrical conductances of aqueous solutions of magnesium, manganese II, barium, and copper succinates have been measured at various temperatures in the range of 298.15 K to 313.15 K, using an audio frequency conductance bridge. The evaluation of conductance data was carried out by minimisation technique using the theoretical equations of the complete and modified forms of Pitts (P) and Fuoss-Hsia (F-H), each a three-parameter equation, association constant (), molar conductance (), and distance parameter (a). Quantitative results showed that these salts do not behave as “strong” electrolytes, and that their dissociations are far from complete. The abnormally low conductances of these electrolytes are not due to the presence of electrically neutral molecules but to the ion-pair formation. The Walden product values, as well as the standard thermodynamics functions (, , ) for the association reaction at the four temperatures studied, have been evaluated.

1. Introduction

The conductance of any electrolyte at infinite dilution is generally recognised to be dependent upon temperature, viscosity, and dielectric constant. Also, it is well established that the extent of association of a series of electrolytes in different solvents is specific for any ion-solvent combination, rather than being dielectric constant dependent. Information on ion-ion and ion-solvent and interactions can be obtained from conductivity measurement [1–4]. In a series of conductance studies, there have been numerous investigations of association behavior of 1 : 1 electrolytes in aqueous and binary mixed solvents.

On the other hand, the conductance behavior of 2 : 2 electrolytes (specially salts of dicarboxylic acids) has received relatively little attention. This is largely due to the difficulty encountered in analysing such data since the usual methods require an arbitrary choice for some of the parameters needed in the analysis. However, there were no available cited recent studies in the literature concerning the association of the bivalent cations with bivalent organic ligands. In part I of this paper, we report conductance measurement of magnesium, manganese II, barium, and copper succinates in aqueous medium at various temperatures in the range 298.15–313.15 K. Metal succinates play an important role in biological and industrial processes [5–8]. Four currently used conductance equations, the complete and expanded forms of both Fuoss-Hsia and Pitts equations, have been used to analyse conductance data for these succinate salts. Also, the thermodynamic parameters , , and were evaluated by investigating the process over a certain temperature range. In addition, Walden product was derived and discussed.

2. Experimental

Measurements were made at a range of temperatures between 298.15 and 313.15 K with Oakton-con510 audio frequency bridge of low impedance. It combines speed and simplicity of operation with 0.05% accuracy. The cell used for measuring conductances has been described elsewhere [9]. The cell constant (as determined by standard solutions of potassium chloride) was 1.04079 cm⁻¹. The method used for measuring conductances has been described [9].

3. Materials

The KCl used in this work was purified and used according to the literature [9]. Magnesium succinate was prepared by addition of clean and dried magnesium ribbon to a solution of succinic acid. The PH of solution was continuously tested until the solution became neutral. Then the mixture was filtered to remove the excess of metal added, after which the solvent was evaporated on a water bath. The product was recrystallized twice from conductivity water and dried by air at room temperature. The final product was a white crystalline powder as the tetrahydrate. The composition of the salt was determined by elemental analysis. The salt was further analysed for magnesium by titration with the disodium salt of ethylenediaminetetraacetic acid (EDTA) using eriochrome black T as indicator [10]. Elemental analysis confirmed the composition with experimental/calculated values (mass percent): C 23.4/22.5, H 6.6/5.7, and Mg 11.38/11.24. FT-IR exhibits the following absorption bands (KBr pellet, υ/cm⁻¹): narrow υ(O–H) bands at 3560 and 3435 cm⁻¹; associated with the existence of hydrogen bonds; stretching frequencies corresponding to the carboxylate ion υ() and υ(C–O) at 1560, 1460, 1413, 1332, 1287, and 1245 cm⁻¹; the 1174 cm⁻¹ absorption band is associated with υ(C–C) and the following absorptions with metal oxygen bonding and with υ(C–H) 1037, 972, 884, 804, 716, 671, 579, 524, and 438 cm⁻¹ [11]. The H-NMR was recorded in D₂O solvent  ppm.

Manganese succinate was prepared by the method described by Vogel [12] with some modification. “Analar” grade manganese chloride solution was added to a solution of sodium succinate. The reaction here was continued until the pH of the mixture became neutral. The mixture was heated on a water bath until a quarter of it was gained. The remaining solution was left for 24 hours. Then light pink crystals formed. The mixture was filtered off. The product was washed with chilled conductivity water until it became free from chloride ions; after that the final product was dried by air at room temperature. Elemental analysis confirmed the composition with experimental/calculated values (mass percent): C 20.3/19.8, H 5.7/5.0. FT-IR spectra exhibit the following absorption bands (KBr pellet, υ/cm⁻¹): υ(O–H) bands at 3505 and 3417 cm⁻¹; stretching frequencies corresponding to the carboxylate ion υ(CO₂) and υ(CO) at 1657, 1552, 1462, 1406, 1329, 1285, and 1244 cm⁻¹; the 1172 cm⁻¹ absorption band is associated with υ(C–C) and the following absorptions with metal-oxygen bonding and with υ(C–H) 1037, 970, 856, 798, 671, and 525 cm⁻¹. The H-NMR  ppm.

The same method [12] was carried out for the preparation of barium succinate. “Analar” grade barium chloride dihydrate solution was added to a solution of sodium succinate, and the rest was that as described before for manganese succinate. The product was a white crystalline powder. Elemental analysis confirmed the composition with experimental/calculated values (mass percent): C 15.8/16.6, H 2.3/2.6. FT-IR spectra exhibit the following absorption bands (KBr pellet, υ/cm⁻¹): υ(O–H) 3550, 3435 cm⁻¹; stretching frequencies corresponding to the carboxylate ion υ() and υ(CO) at 1550, 1397, 1328, 1286, and 1239 cm⁻¹; the 1165 cm⁻¹ absorption band is associated with υ(C–C) and the following absorptions with metal-oxygen bonding and with υ(C–H), 1032, 963, 930, 803, 657, 490, and 464. The H-NMR  ppm.

The copper succinate was prepared by reaction stochiometric amounts of both a solution of succinic acid and copper metal. The solution was heated with continuous stirring at 348.15–353.15 K for a period of 3-4 hours. A blue solution appeared which was evaporated on a water bath after filtering and removing the excess metal copper. The product was a blue powder and dried afterwards for 24 hours, as the dihydrate. Elemental analysis confirmed the composition with experimental/calculated values (mass percent): C 22.3/22.3, H 3.2/3.7. The FT-IR spectra exhibit the following absorption bands (KBr pellet, υ/cm⁻¹): υ(O–H) bands at 3579, 3483, and 3380 cm⁻¹ associated with the existence of hydrogen bonds, stretching frequencies corresponding to the carboxylate ion υ() and υ(C–O) at 1625, 1446, 1329, and 1390 cm⁻¹; the 1185 cm⁻¹ absorbance band is associated with υ(C–C) and the following absorptions with metal oxygen bonding and with υ(C–H), 1001, 973, 886, 730, 694, 618, 581, and 520. The H-NMR was recorded in D₂O solvent  ppm.

All stock solutions were prepared by weight, and all measurements were done using the weight dilution technique.

Figure 1 illustrates an example of FT-IR run.

Figure 1: FT-IR spectra for Mg succinate.

4. Result and Discussion

The measured molar conductances and the corresponding concentrations in mol·dm⁻³ at four different temperatures are given in Tables 1, 2, 3, and 4. These data were analyzed using the complete and modified forms of both F-H and P equations. For the interpretation of the characteristic parameters of an electrolyte solution from conductance data, a minimisation technique has been used in terms of , , and . For analysis of each set of data, , (), a Fortran computer program has been written. The results of minimisation technique for the best fit values of the three parameters , , and are listed in Tables 5, 6, 7, and 8 together with diffusion coefficient ().

Table 1: Experimental conductance data for Ba succinate.

298.15 K
303.15 K
308.15 K
313.15 K

Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ

0.56411
199.190
0.50910
209.414
0.70798
202.902
0.62824
208.904

0.74531
193.356
0.62278
204.192
0.75117
201.282
0.66546
207.388

0.79218
192.096
0.67767
202.628
0.79106
200.014
0.70088
206.188

0.83645
191.262
0.77676
199.560
0.82796
199.898
0.73761
204.738

0.87756
190.604
0.82113
197.964
0.86275
197.870
0.77281
203.830

0.91808
190.128
0.86265
195.676
0.89685
196.440
0.80525
202.728

0.95837
189.466
0.90303
195.570
0.93330
195.458
0.83474
202.112

1.03271
187.164
0.97511
193.654
0.96168
194.830
0.86236
201.070

1.09135
186.406
1.00758
192.836
0.98779
193.896
0.88292
200.514

1.10711
186.808
1.03740
190.804
1.01067
193.368
0.90937
199.260

1.06599
190.326

Λ (ohm⁻¹cm²mol⁻¹); Conc.: concentration (mol·dm⁻³).

Table 2: Experimental conductance data for Cu succinate.

298.15 K
303.15 K
308.15 K
313.15 K

Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ

0.11875
104.322
0.22592
88.190
0.26685
83.448
0.40075
76.194

0.17218
92.200
0.34929
73.130
0.32137
77.388
0.45082
72.174

0.40553
66.350
0.40233
69.116
0.37259
72.824
0.50755
68.362

0.51577
60.392
0.45097
66.048
0.41829
69.286
0.55875
65.358

0.57000
57.934
0.55605
60.490
0.47100
66.172
0.60713
63.150

0.62195
56.106
0.60216
58.582
0.51965
63.582
0.65040
61.270

0.67296
54.482
0.69118
55.478
0.56405
61.622
0.69433
59.530

0.72053
53.198
0.72973
54.188
0.60434
59.882
0.73652
58.098

0.76804
52.312
0.76627
53.302
0.64214
58.666
0.80993
55.852

0.80697
51.142
0.79876
52.208
0.67746
57.336
0.87104
54.114

0.84202
50.218
0.83003
51.590
0.71418
56.100
0.89970
53.402

0.87450
49.544
0.88436
50.392
0.78025
54.250
0.92202
52.900

0.90625
48.870
0.90694
49.998
0.80909
53.506
0.94071
52.430

0.93728
48.336
0.923890
49.530

Λ (ohm⁻¹cm²mol⁻¹); Conc.: concentration (mol·dm⁻³).

Table 3: Experimental conductance data for Mg succinate.

298.15 K
303.15 K
308.15 K
313.15 K

Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ

0.32106
152.146
0.53293
167.682
0.56479
171.312
0.06619
222.736

0.38543
150.700
0.62320
165.106
0.65514
166.354
0.23327
200.850

0.44357
147.958
0.73614
159.570
0.77615
163.214
0.52183
181.330

0.49967
144.364
0.95818
153.276
0.88264
159.736
0.72094
172.758

0.56965
141.704
1.07099
150.494
1.01288
157.180
0.81690
169.664

0.63635
139.118
1.19035
147.426
1.13065
153.924
0.93341
165.490

0.69845
136.062
1.32226
145.708
1.25809
151.156
1.16002
160.302

0.77167
133.942
1.44684
143.592
1.40034
148.436
1.41143
155.162

0.83990
131.736
1.57115
142.498
1.53116
146.800
1.53523
152.818

0.91505
129.448
1.83767
137.972
1.66882
144.824
1.66980
149.540

Λ (ohm⁻¹cm²mol⁻¹); Conc.: concentration (mol·dm⁻³).

Table 4: Experimental conductance data for Mn-succinate.

298.15 K
303.15 K
308.15 K
313.15 K

Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ
Conc. *10³
Λ

0.43244
162.542
0.19989
182.826
0.10976
197.114
0.34023
176.396

0.51084
158.992
0.31941
171.114
0.14567
191.570
0.39961
169.720

0.60785
153.734
0.37725
167.640
0.33335
170.980
0.47590
167.116

0.69413
149.622
0.46203
162.784
0.39655
167.354
0.53986
163.222

1.01124
139.240
0.54652
159.042
0.47351
160.484
0.61789
159.032

1.12663
135.140
1.08379
138.060
0.54065
159.326
0.69363
155.548

1.24575
132.034
1.20277
134.786
0.69278
148.376
0.77113
151.386

1.37319
129.446
1.32895
131.778
0.77883
146.682
0.85864
147.472

1.49389
127.348
1.56919
125.532
0.87155
141.826
0.93962
146.116

1.59976
125.100
1.69673
124.530
1.25557
131.398
1.02234
143.202

Λ (ohm⁻¹cm²mol⁻¹); Conc.: concentration (mol·dm⁻³).

Table 5: Best fit results for Ba-succinate.

Temps.
Parameter
F-H complete*
F-H modified*
Pitts complete
Pitts modified

298.15 K
Λ₀
240.00
238.80
231.40–232.60
236.40

135.00
114.00
275.00–285.00
78.00

a
0.99
1.01
0.30
0.69

*10⁻⁵
1.597
1.589
1.544
1.573

σ
0.1859
0.2110
0.3086
0.2727

ηΛ₀
2.13816
2.12747
2.06689
2.10609

303.15 K
Λ₀
261.20
259.20
250.80–251.00
254.20

182.0
216.0
380.0–382.0
81.0

a
0.70
1.39
0.30
0.70

*10⁻⁵
1.767
1.754
1.698
1.720

σ
0.2410
0.2339
0.2867
0.2267

ηΛ₀
2.08490
2.06893
2.00268
2.02902

308.15 K
Λ₀
279.80
270.00
261.60–261.80
263.60

420.0
263.0
430.0
110.0

a
1.41
1.39
0.70
0.30

*10⁻⁵
1.916
1.857
1.800
1.813

σ
0.1866
0.2308
0.2889
0.2190

ηΛ₀
2.02239
1.95156
1.89157
1.90530

313.15 K
Λ₀
288.80
276.00
275.20
272.00

455.0
264.7
540.0
150.0

a
1.40
1.39
0.35
0.41

*10⁻⁵
2.019
1.929
1.924
1.901

σ
0.2509
0.3156
0.4114
0.3189

ηΛ₀
1.88875
1.80504
1.79981
1.77888

referred values.

Units of Λ₀: ohm⁻¹cm²mol⁻¹, : dm³mol⁻¹, a: nm, : cm²s⁻¹, σ: unitless, and ηΛ₀: ohm⁻¹cm²mol−1 Cp.

Table 6: Best fit results for Cu-succinate.

Temps.
Parameter
F-H complete*
F-H modified*
Pitts complete
Pitts modified

298.15 K
Λ₀
226.20
221.40
214.60–215.00
217.00

11600.0
10510.0
11100.0–11108.0
10050.0

a
1.01
0.71
0.10
0.31

*10⁻⁵
1.505
1.473
1.429
1.444

σ
0.1723
0.2198
0.2417
0.1998

ηΛ₀
2.01522
1.97245
1.91365
1.93325

303.15 K
Λ₀
271.40
267.60
267.80 and 268.00
262.20

16300.0
15020.0
17000.0
14550.0

a
0.69
0.31
1.00
0.30

*10⁻⁵
1.836
1.811
1.812
1.774

σ
0.1800
0.1971
0.2513
0.1869

ηΛ₀
2.16632
2.13598
2.13838
2.09288

308.15 K
Λ₀
287.80
279.60
277.80–278.10
280.40

18100.0
15800.0
17950.0–18050.0
16200.0

a
1.01
0.31
0.30
0.31

*10⁻⁵
1.980
1.923
1.912
1.929

σ
0.1130
0.2095
0.2569
0.1770

ηΛ₀
2.08022
2.02095
2.00902
2.02673

313.15 K
Λ₀
359.20
345.00
331.70–332.30
344.80

25300.0
22100.0
22700.0–22900.0
22450.0

a
0.29
0.29
1.40
0.29

*10⁻⁵
2.511
2.411
2.321
2.410

σ
0.1083
0.1027
0.2589
0.1071

ηΛ₀
2.34917
2.25630
2.17128
2.25500

referred values.

Units of Λ₀: ohm⁻¹cm²mol⁻¹, : dm³mol⁻¹, a: nm, : cm²s⁻¹, σ: unitless, and ηΛ₀: ohm⁻¹cm²mol−1 Cp.

Table 7: Best fit results for Mg-succinate.

Temps.
Parameter
F-H complete*
F-H modified*
Pitts complete
Pitts modified

298.15 K
Λ₀
225.20
222.60
220.00
223.00

120.0
170.0
320.0
110.0

a
0.59
1.40
0.70
0.60

*10⁻⁵
1.499
1.481
1.464
1.484

σ
0.5177
0.2390
0.4397
0.3709

ηΛ₀
2.00631
1.98314
1.95998
1.98671

303.15 K
Λ₀
230.00
229.20
228.00
228.20

175.0
200.0
460.0
160.0

a
0.65
1.00
0.30
0.60

*10⁻⁵
1.556
1.551
1.543
1.544

σ
0.0975
0.0961
11.4499
0.2885

ηΛ₀
1.83586
1.82947
1.81990
1.82149

308.15 K
Λ₀
231.80
232.00
231.60
231.40

370.0
295.0
560.0
280.0

a
1.00
1.00
1.40
0.70

*10⁻⁵
1.594
1.596
1.593
1.592

σ
0.3931
0.9841
1.6379
1.3726

ηΛ₀
1.67545
1.67690
1.67401
1.67256

313.15 K
Λ₀
260.00
253.20
251.20
253.80

480.0
329.0
592.0
305.0

a
1.39
1.41
1.00
1.01

*10⁻⁵
1.817
1.770
1.756
1.774

σ
0.8112
0.6646
0.7072
0.6792

ηΛ₀
1.70040
1.65593
1.64285
1.65985

referred values.

Units of Λ₀: ohm⁻¹cm²mol⁻¹, : dm³mol⁻¹, a: nm, : cm²s⁻¹, σ: unitless, and ηΛ₀: ohm⁻¹cm²mol−1 Cp.

Table 8: Best fit results for Mn-succinate.

Temps.
Parameter
F-H complete*
F-H modified*
Pitts complete
Pitts modified

298.15 K
Λ₀
224.80
220.00
220.40–220.80
223.40

353.0
490.0
850.0
550.0

a
0.65
1.10
0.30
0.65

*10⁻⁵
1.496
1.464
1.468
1.487

σ
0.2028
0.1772
0.3644
0.3928

ηΛ₀
2.00274
1.95998
1.96533
1.99027

303.15 K
Λ₀
226.30
225.80
228.00
226.00

500.0
525.0
950.0
553.0

a
0.60
1.20
0.30
0.75

*10⁻⁵
1.531
1.528
1.543
1.530

σ
0.4441
0.5212
0.7283
0.8156

ηΛ₀
1.80633
1.80234
1.81990
1.80393

308.15 K
Λ₀
232.80
229.80
230.60
230.00

680.0
580.0
990.0
555.0

a
1.00
1.10
0.30
0.80

*10⁻⁵
1.601
1.581
1.586
1.582

σ
0.5795
0.4829
0.7623
0.5641

ηΛ₀
1.68268
1.66099
1.66678
1.66244

313.15 K
Λ₀
246.00
239.60
240.00
242.60

820.0
590.0
1010.0
585.0

a
1.40
1.41
0.30
0.71

*10⁻⁵
1.720
1.675
1.678
1.696

σ
0.4200
0.5529
0.5570
0.7734

ηΛ₀
1.60884
1.56698
1.56960
1.58660

referred values.

Units of Λ₀: ohm⁻¹cm²mol⁻¹, : dm³mol⁻¹, a: nm, : cm²s⁻¹, σ: unitless, and ηΛ₀: ohm⁻¹cm²mol−1 Cp.

Walden product and the corresponding values of the standard deviation are given by

where is defined by

Standard thermodynamic quantities for the association reaction are obtained from the temperature dependence of the association constant . The standard enthalpy change () was determined from the slope of versus [13] as follows:

and the standard Gibbs energy () and entropy () changes were calculated using the well-known relations

Also the diffusion coefficients of the salts () can be founded from the equation [14]

where  j K⁻¹ mole⁻¹, , and , while the Walden’s product can be calculated as a function of temperature by using

It was noted that the minimisation and values for the complete and modified Fuoss-Hsia equations were almost greater than those values obtained using the complete and modified forms of the Pitts. The reason for this has been discussed by Fernandez-Prini and Prue [15, 16] in outline, and by Pitts et al. [17] in detail. The Brownian terms in the velocity of the ion were neglected by Pitts. Also, Pitts equation does not allow for the kinetic (osmotic) terms which contributed to the increase in velocity of the ion.

It was obvious from the minimisation technique that the parameter had the greatest effect on the values of during the variation of the three parameters , , and . This is to be expected, in so far as is the leading term in all the conductance equations. However, was relatively insensitive to the values of and at the corresponding minima of .

Form the association constants given in Tables 5–8, at different temperatures, it can therefore be anticipated that these salts do not behave as “strong” electrolytes, and that their dissociations are far from complete. It can also be seen that copper succinate undergoes more ion-pair formation according to the quantitative conductance data. In conclusion, it appears that the copper succinate is capable of showing a marked tendency to autocomplex formation, similar to the same tendency as for copper malonate [18, 19]. The high-association constant for the Cu(II) salt, with the ready availability of d-levels in the cation, illustrates this tendency. At present, however, it is not possible to make a calculation for such complex formation in our computer programs. So, no further analysis was done for this salt. On the above views, it can be generally predicated, for all studied salts, that the degree of dissociation is highly changed from one to another.

The succinate ion possesses the donor property to a less degree than both oxalate and malonate ions [20, 21], so it can also be emphasized that succinate salts are more highly dissociated than the oxalates [22] and malonates [23, 24].

In Figures 2, 3, 4, and 5, curves of molar conductance versus square root of concentration at four temperatures have been obtained for the four succinate salts. For these electrolytes, the measured conductivity was abnormally small. Abnormally weak salts usually form autocomplexes readily, as are shown by Ives and Riley’s conductivity measurements [25] for copper malonate. In this work, with the exception of copper succinate, however, the experimental conductances give no evidence of this. The same conclusion was found for other systems by other authors [9, 18–24].

Figure 2: Molar conductance versus square root of concentration for Ba-succinate at different temperatures.

Figure 3: Molar conductance versus square root of concentration for Cu-succinate at different temperatures.

Figure 4: Molar conductance versus square root of concentration for Mg-succinate at different temperatures.

Figure 5: Molar conductance versus square root of concentration for Mn-succinate at different temperatures.

Considering the difficulties resulting from the limited solubilities of the succinate salts, the values obtained in this work for , , and are satisfactory, providing a particularly good example for typical incompletely dissociated electrolytes, the thermodynamic association constants of which conform to the law of mass action. The abnormally low conductances of these electrolytes are not due to the presence of electrically neutral molecules but to the ion-pair formation. This is to be indicated by further measurements at high different frequencies which will be taken into consideration in the future studies. Such evidence has been shown for both oxalate [22] and malonate salts [9], as well as for Mg and Mn sulphates [26].

It is obvious from Tables (5–8) that some values are small, while others are large depending on the type of the four conductance equations used. The explanation for such discrepancies in these results is that the small values indicate that the linkage is largely covalent [18, 19, 27] and not merely due to “coulomb” ion association. The large values were explained by the fact that the assumption of spherical symmetry and a central charge is far from truth for the succinate ion just like that in barium malonate [18, 19].

Recently, an investigation of Mg, Ca, and Ba complexes in aqueous solution was carried out [28]. The result showed that the conductance of Mg⁺² ion is much lower due to greater interaction between the charge on the ion and the dipoles of the adjacent solvent molecules, which leads to a reduction in mobility. A good agreement between their values of for Mg and Ba succinates and our present values was obtained for both salts.

Unfortunately, no recent determinations of at 298.15 K for all salts studied have been found. So, in comparison with ionic mobilities at infinite dilution at 298.15 K taken from [14]: , , , and together with the value of 58.8 [14] or 56.99 [29] of succinate ion, good agreement has been shown in Tables 5–8.

In order to obtain a better understanding of the thermodynamics of the association reactions for the studied salts, it is useful to consider the enthalpic and entropic contributions to these associations. The standard enthalpy, free energy and entropy changes were determined by using (3) and (4), respectively, at different temperatures. Summarizing at this point, give the conductometric association constant , the standard free energy change for the postulated equilibrium can be determined. This opens the way to a search for correlations between , , and on the one hand and basic properties of solutes and solvents on the other hand. According to the minimisation technique used here, each of the seven systems gave a unique best set of parameters at each temperature. While the best and showed an expected trend with temperature, this trend for was quite irregular and covered a broader range of values. Tables 5–8 show clearly the temperature dependence of the of particular systems. The increase of with an increase of temperature is due to the decrease of solvent viscosity. Bearing in mind that dielectric constant decreases as temperature increases, an increase in temperature should play important role in stabilizing ion-pairs of reaction. So, increases as shown in the same mentioned tables. Furthermore, when the dielectric constant of the solvent becomes smaller as temperature increases, the electrostatic contribution dominates. As a consequence, many complexes become significantly stable at high temperatures [30]. The same trend was found by Franchini et al. [31] for a conductometric study of dissociation of picric acid in two different organic solvents at different temperatures, as well as by Sokol et al. [13] in their study for the thermodynamics of the association reaction conductance measurements.

Standard thermodynamic quantities for the association reaction of metal ion (M⁺²) and succinate ions are obtained from the temperature dependence of , (3), as given in Tables 9, 10, 11, and 12.

Table 9: Thermodynamic data for Ba-succinate.

(kJ mol⁻¹)
(kJ mol⁻¹)
(J K⁻¹ mol⁻¹)

298.15 K
303.15 K
308.15 K
313.15 K
298.15 K
303.15 K
308.15 K
313.15 K

F-H complete
69.681
−12.159
−13.116
−15.475
−15.934
274.493
273.122
276.346
273.400

F-H modified
42.614
−11.740
−13.548
−14.276
−14.524
182.303
185.260
184.617
182.461

Pitts complete
32.510
−13.968
−14.978
−15.535
−16.380
155.887
156.647
155.913
156.122

Pitts modified
35.069
−10.800
−11.076
−12.042
−13.045
153.846
152.219
152.884
153.646

Table 10: Thermodynamic data for Cu-succinate.

(kJ mol⁻¹)
(kJ mol⁻¹)
(J K⁻¹ mol⁻¹)

298.15 K
303.15 K
308.15 K
313.15 K
298.15 K
303.15 K
308.15 K
313.15 K

F-H complete
37.950
−23.199
−24.445
−25.117
−26.396
168.065
168.450
173.372
172.070

F-H modified
35.414
−22.954
−24.239
−24.768
−26.044
158.155
161.510
161.252
159.469

Pitts complete
34.494
−23.090
−24.551
−25.102
−26.125
162.542
163.193
162.353
162.459

Pitts modified
39.125
−22.843
−24.159
−24.832
−26.085
167.448
165.597
166.045
166.597

Table 11: Thermodynamic data for Mg-succinate.

(kJ mol⁻¹)
(kJ mol⁻¹)
(J K⁻¹ mol⁻¹)

298.15 K
303.15 K
308.15 K
313.15 K
298.15 K
303.15 K
308.15 K
313.15 K

F-H complete
76.249
−11.867
−13.017
−15.150
−16.074
296.520
294.787
297.659
294.372

F-H modified
36.812
−12.731
−13.354
−14.570
−15.090
162.843
166.122
165.789
163.933

Pitts complete
31.863
−14.299
−15.453
−16.212
−16.620
153.717
154.514
153.814
154.056

Pitts modified
56.333
−11.652
−12.791
−14.436
−14.893
225.165
222.361
221.888
221.548

Table 12: Thermodynamic data for Mn-succinate.

(kJ mol⁻¹)
(kJ mol⁻¹)
(J K⁻¹ mol⁻¹)

298.15 K
303.15 K
308.15 K
313.15 K
298.15 K
303.15 K
308.15 K
313.15 K

F-H complete
44.111
−14.542
−15.663
−16.709
−17.468
188.730
188.774
193.366
191.745

F-H modified
10.223
−15.355
−15.786
−16.302
−16.611
73.663
78.412
79.503
79.025

Pitts complete
8.720
−16.720
−17.281
−17.672
−18.010
76.095
78.171
78.710
80.152

Pitts modified
2.906
−15.641
−15.903
−16.189
−16.589
45.969
46.121
48.508
50.936

The positive entropy values of the associations indicate a change from more order states towards a less order states. This might reflect the role of solvent in the association process.

The above-mentioned results show that the term is the most important factor in determining the stability of the association, and the differences in the magnitude of values reflected the differences in the attractive forces between opposite charges of the ions of ion-pair.

Many difficulties and uncertainties are associated with these calculations which are derived from the inevitable experimental errors in values and partly from the calculator’s choice of the function which represents the dependence on the temperature. However, the values of have proved useful in yielding structural information about solute species and solute-solvent interactions. So, such values indicate a spontaneous association of ions (i.e., the yield of ion-pairs increases) despite of that they differ from one equation to the other depending on their mathematical derivations.

From the dependence of the Walden product on the temperature, information can be obtained on ion-solvent interactions. Data from Tables 5–8 show that the Walden product dependence on the temperature is substantially obeyed. For more consistent results, a consideration should be taken for the electrostatic interaction between the fields of the moving ions and those of the solvent dipoles in the surrounding solvent. The gradual decrease of with temperature has been attributed to a decrease in dielectric constant, causing a disturbance in the hydrodynamic radii of ions followed by a change in their mobility.

Generally, it is concluded that either Fuoss-Hsia complete or modified is better than those equations of Pitts.

Acknowledgment

With affection and deep appreciation the authors acknowledge their indebtedness to: Mr. Salam G. Tahir for doing CHN and HNMR measurements for the prepared compounds in University of Wales-Bangor (UWB), UK.

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