Molecular docking, potentiometric and thermodynamic studies of some azo quinoline compounds

Article abstract of DOI:10.1016/j.molliq.2015.12.081

The ligands of 5-(4-derivative phenyl azo)-8-hydroxyquinoline and its derivatives (HL_n) were synthesized from the coupling of the quinoline with diazonium salt derived from aniline and its p-derivatives and characterized by different spectrosco

Full text of DOI:10.1016/j.molliq.2015.12.081

Journal of Molecular Liquids 215 (2016) 740–748
Contents lists available at ScienceDirect
Journal of Molecular Liquids
journal homepage: www.elsevier.com/locate/molliq
Molecular docking, potentiometric and thermodynamic studies of some
azo quinoline compounds
⁎
A.F. Shoair , A.A. El-Bindary, A.Z. El-Sonbati, N.M. Beshry
Chemistry Department, Faculty of Science, Damietta University, Damietta 34517, Egypt
a r t i c l e i n f o
a b s t r a c t
Article history:
The ligands of 5-(4-derivative phenyl azo)-8-hydroxyquinoline and its derivatives (HLn) were synthesized from
the coupling of the quinoline with diazonium salt derived from aniline and its p-derivatives and characterized by
different spectroscopic techniques. Molecular docking was used to predict the binding between azo compounds
with the receptor of breast cancer mutant 3hb5-oxidoreductase. The X-ray diffraction, XRD, pattern of the ligand
(HL3) is polycrystalline nature. The proton-ligand dissociation constant of the azo compounds (HLn) and metal–
ligand stability constants of their complexes with metal ions (Mn2+, Co2+, Ni2+ and Cu2+) have been deter-
mined by potentiometric technique in 0.1 M KCl and 50% (by volume) DMF–water mixture. For the same ligand
at constant temperature, the stability constants of the formed complexes increase in the order Mn2+, Co2+, Ni2+
and Cu2+. The effect of temperature was studied at 298, 308 and 318 K and the corresponding thermodynamic
parameters (ΔG, ΔH and ΔS) were derived and discussed. The dissociation process is non-spontaneous, endo-
thermic and entropically unfavorable. The formation of the metal complexes has been found to be spontaneous,
endothermic and entropically favorable.
Received 22 October 2015
Received in revised form 18 December 2015
Accepted 23 December 2015
Available online xxxx
Keywords:
Quinoline azodye
Potentiometry
Thermodynamic parameters
Molecular docking
X-ray diffraction
© 2016 Elsevier B.V. All rights reserved.
1. Introduction
predict protein–ligand [10,11] and to screen large libraries for mole-
cules that will modulate the activity of a biological receptor.
Quinoline azodye and its derivatives are very important compounds
and have attracted much attention in both academic and applied re-
search used in many applications such as their biological relevance, co-
ordination capacity, their use as metal extracting agent and their
therapeutic properties [1–6]. Also the azo compounds based on quino-
line play a central role as chelating agents for a large number of metal
ions, as they form a stable six and/or five-membered ring after complex-
ation with the metal ion [4,6]. The azo compounds are used in dying
processes; some of them are used in analytical separation of many
metal ions in a mixture. It is well-known that N atoms play a key
role in the coordination of metals at the active sites of numerous
metallobiomolecules [7].
In this paper, the potentiometric studies are used to determine
the dissociation constants of 5-(4-derivative phenyl azo)-8-
hydroxyquinoline (HLn) and the stability constants of its complexes
with some divalent transition metal ions such as Mn2+, Co2+, Ni2+
and Cu2+ at different temperatures. The corresponding thermodynamic
functions are evaluated and discussed. Moreover, the molecular docking
of the ligands (HLn) is studied.
2. Materials and methods
All the compounds and solvents used were purchased from Aldrich
and Sigma and used as received without further purification.
The potentiometric study is one of the most convenient and success-
ful technique employed for metal complex equilibrium measurements.
Such applications are based on studying the influence of the pH on the
equilibrium system components; metal ion, ligand and proton [8,9].
The use of protein–ligand docking has become a standard method in po-
tentiometric studies [10]. The protein groups surrounding the ligand
can highly influence the local pH, so that a different protonation could
be favored in the bound state. The molecular docking is widely used to
2.1. Preparation of the ligands
The ligands of 5-(4-derivative phenyl azo)-8-hydroxyquinoline and
its derivatives (HLn) were prepared by dissolving aniline or its p-
substituted derivatives (10 mmol) in hydrochloric acid [12,13]. The
compound was diazotized below −5 °C in an ice–salt bath with a solu-
tion of sodium nitrite (0.8 g, 10 mmol, 30 ml distilled H2O). The diazo-
nium salt was coupled with an alkaline solution of quinoline-8-ol
(1.0 g, 10 mmol) in pyridine. The precipitate was filtered and dried
after thorough washing with water and ethanol. The crude products
were purified by recrystallization from hot ethanol and dried in vacuum
⁎
Corresponding author.
E-mail address: abdel_shoair@yahoo.com (A.F. Shoair).
http://dx.doi.org/10.1016/j.molliq.2015.12.081
0167-7322/© 2016 Elsevier B.V. All rights reserved.

A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
741
Table 1
desiccator over P2O5. Yield percent was 65–81%. The ligands are also
characterized by IR and 1H NMR spectroscopies (Fig. 1).
Analytical data of the ligands (HLn).
Compound
Exp. (Calc.) %
C
HL1 = 5-(4-methoxyphenyl azo)-8-hydroxyquinoline.
H
N
HL2 = 5-(4-methylphenyl azo)-8-hydroxyquinoline.
HL3 = 5-(phenyl azo)-8-hydroxyquinoline.
HL4 = 5-(4-chlorophenyl azo)-8-hydroxyquinoline.
HL1
HL2
HL3
HL4
68.64 (68.82)
72.87 (73.00)
72.12 (72.29)
63.26 (63.49)
4.52 (4.66)
4.78 (4.94)
4.20 (4.42)
3.32 (3.53)
14.84 (15.05)
15.67 (15.97)
16.51 (16.87)
14.44 (14.82)
2.2. Potentiometric studies
2.3. Measurements
A ligand solution (0.001 M) was prepared by dissolving an accurate-
ly weighted amount of the solid in DMF. Metal ion solutions (0.0001 M)
were prepared from metal chlorides in bidistilled water and standard-
ized with EDTA [14]. Solutions of 0.001 M HCl and 1 M KCl were also
prepared in bidistilled water. A carbonate-free NaOH solution in 50%
(by volume) DMF–water mixture was used as titrant and standardized
against oxalic acid.
The apparatus, general conditions and methods of calculation were
the same as in previous work [15,16]. The following mixtures (i)–(iii)
were prepared and titrated potentiometrically at 298 K against standard
0.002 M NaOH in a 50% (by volume) DMF–water mixture:
Elemental microanalyses of the separated ligands for C, H and N
were determined on Automatic Analyzer CHNS Vario ELIII, Germany.
FT-IR spectra (KBr disks, 4000–400 cm−1) by Jasco-4100 spectropho-
tometer. The 1H NMR spectra by Bruker WP 300 MHz using DMSO-d6
as a solvent containing TMS as the internal standard. X-ray diffraction
measurement (XRD) is recorded on X-ray diffractometer in the range
of diffraction angle 2θo = 5–80°. This analysis is carried out using
CuKα1 radiation (λ = 1.540598 Å). The applied voltage and the tube
current are 40 KV and 30 mA, respectively. The pH measurements
were performed with a Metrohm 836 Titrando (KF & Potentiometric Ti-
trator) equipped with a combined porolyte electrode. The temperature
was controlled to within 0.05 K by circulating thermostated water
(Neslab 2 RTE 220) through the outer jacket of the vessel.
Docking calculations were carried out using a Docking Server [17–
19]. The MMFF94 force field [20] was used for energy minimization of
ligand molecule using Docking Server. Gasteiger partial charges were
added to the ligand atoms. Non-polar hydrogen atoms were merged
and rotatable bonds were defined. Docking calculations were carried
out on 3hb5-oxidoreductase protein model. Essential hydrogen atoms,
Kollman united atom type charges, and solvation parameters were
added with the aid of AutoDock tools [21]. Affinity (grid) maps of
20 × 20 × 20 Å grid points and 0.375 Å spacing were generated using
the Autogrid program [21]. AutoDock parameter set- and distance-
dependent dielectric functions were used in the calculation of the van
der Waals and the electrostatic terms, respectively. Docking simulations
were performed using the Lamarckian genetic algorithm (LGA) and the
Solis & Wets local search method [22]. Initial position, orientation and
torsions of the ligand molecules were set randomly. Each docking ex-
periment was derived from 10 different runs that were set to terminate
after a maximum of 250,000 energy evaluations. The population size
was set to 150. During the search, a translational step of 0.2 Å, and qua-
ternion and torsion steps of 5 were applied.
i) 5 cm3 0.001 M HCl + 5 cm3 1 M KCl + 25 cm3 DMF.
ii) 5 cm3 0.001 M HCl + 5 cm3 1 M KCl + 20 cm3 DMF + 5 cm3
0.001 M ligand.
iii) 5 cm3 0.001 M HCl + 5 cm3 l M KCl + 20 cm3 DMF + 5 cm3
0.001 M ligand +10 cm3 0.0001 M metal chloride.
For each mixture, the volume was made up to 50 cm3 with bidistilled
water before the titration. These titrations were repeated for the tem-
peratures of 308 and 318 K. All titrations have been carried out between
pH 3.5 and 13.0 and under nitrogen atmosphere.
3. Results and discussion
3.1. Characterization of the ligands (HLn)
The chemical structures of the ligands were elucidated by elemental
analyses Table 1. The infrared spectra of ligands (HLn) shows two bands
in the range of 3266–3315 and 1570–1590 cm−1 for stretching OH of
quinoline at C8-position and CNquin. (nitrogen atom of azomethine of
quinoline group), respectively. The aromatic C–H bands was observed
at 3000–3120 cm−1 as used and methyl C–H vibration of methoxy
group was observed at 2990–2850 cm−1 and exhibit band in the
range of 1500–1504 cm−1 which could be assigned to υN_N stretching
vibration [23–27].
1H NMR spectra of ligands (HLn) were recorded in
dimethylsulphoxide (DMSO-d6) solution using tetramethylsilane
(TMS) as internal standard. The 1H NMR spectra of quinoline and ben-
zene rings appeared in the range of 7.01–8.25 ppm. For the HL1 has a
singlet observed at 3.88 ppm is assigned to OCH3 protons (the integra-
tion curve shows three protons). Also HL2 has a singlet at 3.76 ppm
which is assigned to the CH3 protons. The 1H NMR spectra show two
Fig. 1. The formation mechanism of azo quinoline derivatives (HLn).

742
A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
for a, b, c, α, β and γ, respectively. The inter-planar spacing (d) and Mill-
er indices (hkl) which are estimated by CRYSFIRE are listed in Table 2.
3.3. Molecular docking
The docking study showed a favorable interaction between the re-
ceptor breast cancer (3hb5) and the ligands of 5-(4-derivative phenyl
azo)-8-hydroxyquinoline and its derivatives (HLn) and the calculated
energy is listed in Table 3 and Fig. 3. According to the results obtained
in this study, HB plot curve indicated that, the ligands (HLn) bind to
the protein with hydrogen bond interactions and decomposed interac-
tion energies in kcal/mol were existing between the ligands (HLn)
with 3hb5 receptor as shown in Fig. 4. The calculated efficiency is favor-
able where Ki values estimated by AutoDock were compared with ex-
perimental Ki values, when available, and the Gibbs free energy is
negative. Also, based on this data, we can propose that interaction be-
tween the 3hb5 receptor and the ligands (HLn) is possible. 2D plot
curves of docking with ligands (HLn) are shown in Fig. 5. This interac-
tion could activate apoptosis in cancer cells energy of interactions
with ligands (HLn). Binding energies are most widely used as mode of
measuring binding affinity of ligands. Thus, decrease in binding energy
due to mutation will increase the binding affinity of the ligands towards
the receptor [10]. The characteristic feature of ligands was represented
in the presence of several active sites available for hydrogen bonding.
Fig. 2. X-ray diffraction pattern of HL3 in powder form.
Table 2
Crystallographic data of HL3.
Peak no.
2θobs. (°)
dobs. (Å)
dcal. (Å)
(h k l)
3.4. Potentiometric studies
1
2
3
6.514
7.917
9.788
13.56162
11.16377
9.030481
13.70641
11.21252
8.988331
011.
111.
The hydrogen ion concentration, [H+], was determined
potentiometerically in the usual manner [10]. The average number of
the protons associated with ligands (HLn) at different pH values,‾nA,
was calculated from the titration curves of the acid in the absence and
presence of ligands (HLn) by applying the following equation:
210.
4
14.586
6.070134
6.070134
113.
222.
033.
412.
323.
5
6
7
8
9
15.772
19.490
20.388
21.288
23.871
5.614431
4.5526
4.354438
4.170547
3.724855
5.606259
4.563389
4.344614
4.170547
3.730266
ꢀ
ꢁ
ðV1−V2Þ N0 þ E0
432.
234.
522.
10
11
12
24.110
25.052
26.359
3.689176
3.553044
3.378509
3.692713
3.562905
3.385934
ꢀ
ꢁ
nA ¼ Y ꢀ
ð1Þ
V0−V1 TC0L
343.
053.
325.
344.
13
14
15
16
27.205
28.082
29.584
44.022
3.276478
3.175185
3.017235
2.055778
3.275085
3.176495
3.0196
where Y is the number of available protons in ligands (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, V° is the
initial volume (50 cm3) of the mixture, TC°L is the total concentration
of the reagent, N° is the normality of sodium hydroxide solution and
E° is the initial concentration of the free acid. Thus, the formation curves
(n‾A vs. pH) for the proton-ligand systems were constructed and found to
extend between 0 and 1 in the n‾A scale. This means that quinoline azo
dyes have one ionizable proton (the enolized hydrogen ion of the phe-
nolic –OH group, pKH as shown in Table 4). The completely protonated
form of ligands (HLn) has one dissociable proton, that dissociates in the
measurable pH range. The deprotonation of the hydroxyl group most
probably results in the formation of stable intramolecular H-bonding
with nitrogen of the azo group. Such an interaction decreases the disso-
ciation process of phenolic –OH group, i.e. increases the pKH value [28,
29].
2.055778
763.
584.
17
47.760
1.902915
1.902915
singlets for C8–OH at ~9.55–10.30 ppm and HCN at ~9.09–9.30 ppm
[12].
3.2. X-ray diffraction analysis
The X-ray diffraction (XRD) pattern of HL3 ligand in powder form is
shown in Fig. 2. The XRD pattern shows that the ligand (HL3) has a poly-
crystalline nature. The calculated crystal system of HL3 is found to be
monoclinic with space group P21/A. The estimated lattice parameters
are found to be 20.4710 Å, 18.8150 Å, 19.9590 Å, 90.0°, 92.7° and 90.0°
The formation curves for the metal complexes were obtained by
plotting the average number of ligands attached per metal ion (n‾A) vs.
Table 3
Energy values obtained in docking calculations of ligands (HLn) with receptor breast cancer mutant 3hb5.
Receptor
Est. free energy of binding
(kcal/mol)
Est. inhibition constant
(Ki) (μM)
vdW+ bond + desolve
energy (kcal/mol)
Electrostatic
energy (kcal/mol)
Total intercooled
energy (kcal/mol)
Interact surface
HL1
HL2
HL3
HL4
−6.23
−6.83
−6.72
−7.46
27.09
9.88
11.90
3.41
−7.64
−7.67
−7.64
−8.14
−0.10
−0.08
−0.06
−0.05
−7.74
−7.75
−7.71
−8.20
732.707
710.141
659.553
716.271

A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
743
Fig. 3. The ligands (HLn) (green in (A) and blue in (B)) in interaction with receptor breast cancer mutant 3hb5. (For interpretation of the references to color in this figure legend, the reader
is referred to the web version of this article).

744
A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
Fig. 4. HB plot of interaction between ligands (HLn) and receptor breast cancer mutant 3hb5.
the free ligands exponent (pL). The average number of the reagent mol-
ecules attached per metal ion, n‾, and free ligands exponent, pL, can be
calculated using Eqs. (2) and (3):
curves relative to ligand titration curves points to the formation
of strong metal complexes [10].
(iv) For the same ligand at constant temperature, the stability of the
chelates increases in the order Mn2+, Co2+, Ni2+ and Cu2+
[16]. This order largely reflects that the stability of Cu2+ com-
plexes is considerably larger than those of other metals of the
3d series. Under the influence of both the polarizing ability of
the metal ion and the ligands field [10] Cu2+ will receive some
extra stabilization due to tetragonal distortion of octahedral sym-
metry in its complexes. The greater stability of Cu2+ complexes is
produced by the well-known Jahn–Teller effect [30].
ðV3−V2ÞðN∘ þ E∘Þ
nA
¼
ð2Þ
ðV∘−V2Þ ꢁ nAꢁTC∘
M
and
ꢂ
ꢃ
n¼J
n
X
βnH
1
Hþ
V0 þ V3
½
ꢂ
n¼0
pL ¼ log10
ꢁ
ð3Þ
TC0L −n:TC0M
V0
Stepwise dissociation constants for ligands (HLn) and the stepwise
stability constants of their complexes with Mn2+, Co2+, Ni2+ and
Cu2+ have been calculated at 298, 308 and 318 K. The corresponding
thermodynamic parameters (ΔG, ΔH and ΔS) were evaluated.
The dissociation constants (pKH) for ligands (HLn), as well as the sta-
bility constants of their complexes with Mn2+, Co2+, Ni2+ and Cu2+
have been evaluated at 298, 308 and 318 K, and are given in Tables 4
and 5, respectively. The enthalpy (ΔH) for the dissociation and com-
plexation process was calculated from the slope of the plot pKH or log
K vs. 1/T as shown in Fig. 6 using the graphical representation of van't
Hoff Eqs. (4) and (5):
where TC°M is the total concentration of the metal ion present in the so-
lution, βHn 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 [18]. The values of the stability constants
(log K1 and log K2) are given in Table 5. The following general remarks
can be pointed out:
(i) The maximum value of n‾ was ~2 indicating the formation of 1:1
and 1:2 (metal:ligand) complexes only [10].
(ii) The metal ion solution used in the present study was very dilute
(2 × 10−5 M), hence there was no possibility of formation of
polynuclear complexes [15].
(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 ligands. The large decrease in pH for the metal titration
ΔG ¼ −2:303 RT log K ¼ ΔH–TΔS
ð4Þ
or
ꢂ
ꢃꢂ ꢃ
−ΔH
2:303R
1
T
ΔS
log K ¼
þ
2:303R
ð5Þ

A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
745
Fig. 5. 2D plot of interaction between ligands (HLn) and receptor breast cancer mutant 3hb5.

746
A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
Table 4
Thermodynamic functions for the dissociation of the ligands (HLn) in 50% (by volume) DMF–water mixtures and 0.1 M KCl at different temperatures.
Gibbs energy
Enthalpy change
Entropy change
Dissociation constant
pKH1
(kJ mol−1
)
(kJ mol−1
)
(J mol−1 K−1
)
Temperature
(K)
Compound
HL1
ΔG1
ΔH1
- ΔS1
298
308
318
298
308
318
298
308
318
298
308
318
9.48
9.35
9.24
9.25
9.10
8.95
9.00
8.85
8.71
8.82
8.69
8.56
54.09
55.14
56.26
52.78
53.67
54.49
51.35
52.19
53.03
50.33
51.25
52.12
21.79
108.41
108.29
108.41
85.81
85.90
85.81
84.03
84.03
84.04
89.75
89.83
89.75
HL2
HL3
HL4
27.21
26.31
23.58
Table 5
Stepwise stability constants for complexes of the ligands (HLn) in 50% (by volume) DMF–
water mixtures and 0.1 M KCl at different temperatures.
Table 6
Thermodynamic functions for ML and ML2 complexes of the ligands (HLn) in 50% (by vol-
ume) DMF–water mixtures and 0.1 M KCl.
Comp.
HL1
Mn+
298 K
log k1
308 K
log k1
318 K
log k1
Gibbs energy
(kJ mol−1
Enthalpy change
(kJ mol−1
Entropy change
(J mol−1 K−1
log k2
log k2
log k2
Comp.
HL1
Mn+
T/K
)
)
)
Mn2+
Co2+
Ni2+
Cu2+
Mn2+
Co2+
Ni2+
Cu2+
Mn2+
Co2+
Ni2+
Cu2+
Mn2+
Co2+
Ni2+
Cu2+
5.82
5.95
6.10
6.21
6.00
6.12
6.25
6.39
6.21
6.35
6.43
6.55
6.38
6.52
6.66
6.80
4.59
4.64
4.78
4.90
4.80
4.93
5.09
5.22
4.98
5.12
5.26
5.37
5.16
5.30
5.42
5.55
5.95
6.10
6.21
6.35
6.14
6.26
6.40
6.54
6.35
6.49
6.57
6.67
6.53
6.64
6.78
6.92
4.71
4.76
4.90
5.03
4.93
5.05
5.22
5.37
5.12
5.25
5.39
5.50
5.31
5.43
5.55
5.67
6.09
6.22
6.35
6.47
6.27
6.38
6.53
6.68
6.47
6.61
6.70
6.80
6.67
6.78
6.92
7.05
4.85
4.91
5.03
5.15
5.07
5.19
5.35
5.49
5.24
5.40
5.55
5.65
5.43
5.57
5.68
5.79
- ΔG1
- ΔG2
ΔH1
ΔH2
ΔS1
ΔS2
Mn2+
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
26.19
27.78
29.53
26.48
28.07
29.90
27.27
28.90
30.63
27.96
29.66
31.36
27.39
29.07
30.87
28.13
29.78
31.60
29.04
30.78
32.58
29.78
31.67
33.43
28.42
30.19
31.91
29.21
30.96
32.88
30.01
31.79
33.79
30.64
32.44
34.40
29.44
31.31
33.06
30.24
32.02
33.91
30.93
32.73
34.58
31.67
33.44
35.25
33.21
35.09
37.08
33.95
35.97
37.87
34.81
36.62
38.66
35.43
37.45
39.39
34.24
36.21
38.18
34.92
36.92
38.85
35.66
37.74
39.76
36.46
38.57
40.67
35.43
37.45
39.39
36.23
38.27
40.25
36.69
38.75
40.79
37.37
39.34
41.40
36.40
38.51
40.61
37.20
39.16
41.28
38.00
39.98
42.13
38.80
40.81
42.93
23.56
24.48
166.95
166.68
166.95
170.91
170.55
170.92
167.57
167.40
167.58
169.94
169.95
169.94
174.04
173.87
174.05
173.46
173.19
173.46
176.59
176.51
176.59
182.22
182.42
182.21
174.55
174.65
174.54
183.18
182.91
183.19
188.87
188.50
188.88
187.97
187.69
187.97
181.07
181.27
181.06
183.62
183.44
183.62
182.90
182.82
182.91
179.31
179.23
179.31
193.57
193.40
193.58
196.19
196.40
196.19
192.78
192.42
192.79
198.10
198.21
198.09
197.09
197.10
197.09
196.37
196.48
196.37
204.95
205.05
204.95
210.64
210.65
210.64
198.10
198.21
198.09
200.78
200.89
200.77
205.32
205.33
205.32
201.46
201.29
201.47
210.45
210.45
210.45
203.90
203.63
203.91
206.58
206.31
206.59
206.25
206.08
206.25
HL2
HL3
HL4
Co2+
Ni2+
Cu2+
Mn2+
Co2+
Ni2+
Cu2+
Mn2+
Co2+
Ni2+
Cu2+
Mn2+
Co2+
Ni2+
Cu2+
24.46
22.66
22.68
24.48
23.56
23.58
24.52
23.60
25.37
26.27
25.37
24.52
24.48
23.58
21.77
24.52
22.64
23.60
24.50
23.60
25.41
26.31
23.60
23.60
24.50
22.66
26.31
23.56
23.56
22.66
HL2
HL3
HL4
where R is the gas constant =8.314 J mol−1 K−1, K is the dissociation
constant for the ligand stability and T is the temperature (K).
Fig. 6. Van't Hoff plot pKH1 of ligands (HLn) against 1/T.

A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
747
Fig. 7. Van't Hoff plot of log K1 of Mn+ complexes with ligands (HLn) against 1/T.
From the ΔG and ΔH values, one can deduce the entropy ΔS using
the well-known relationships (4) and (6):
(iii) The positive values of ΔH mean that the complex formation pro-
cesses are endothermic and favored at higher temperature.
(iv) The positive values of ΔS are confirming that the complex forma-
tion processes are thermodynamically favorable [16].
ΔS ¼ ðΔH−ΔGÞ=T
ð6Þ
The thermodynamic parameters of the dissociation process of li-
gands (HLn) are recorded in Table 4. From these results, the following
can be made:
Inspection of the results in Table 6 reveals that the pKH values of 5-
(4-derivative phenyl azo)-8-hydroxyquinoline and its derivatives are
influenced by the inductive or mesomeric effect of the substituents.
HL1 has a lower acidic character (higher pKH values) than HL2. This is
quite reasonable because the presence of p-CH3 group (i.e. an
electron-donating effect) will enhance the electron density by their
high positive inductive or mesomeric effect, whereby a stronger O\\H
bond is formed. The presence of p-Cl group (i.e. an electron-
withdrawing effect) will lead to the opposite effect [18].
(i) The pKH values decrease with increasing temperature, i.e. the
acidity of ligands increases [18].
(ii) Positive values of ΔH indicate that dissociation is accompanied
by absorption of heat and the process is endothermic.
(iii) Positive values of ΔG indicate that the dissociation process is not
spontaneous [16].
(iv) Negative values of ΔS are due to increased order as a result of the
solvation processes.
4. Conclusion
5-(4-Derivative phenyl azo)-8-hydroxyquinoline and its derivatives
(HLn) have been synthesized and characterized by different spectro-
scopic techniques. The proton-ligand dissociation constant of ligands
(HLn) and metal–ligand stability constants of their complexes with
metal ions (Mn2+, Co2+, Ni2+ and Cu2+) at different temperatures
were determined. The stability constants of the formed complexes in-
crease in the order Mn2+, Co2+, Ni2+ and Cu2+. The dissociation pro-
cess is non-spontaneous, endothermic and entropically unfavorable.
The formation of the metal complexes has been found to be spontane-
ous, endothermic and entropically favorable. The stability constants
(log K1 and log K2) for the complexes increase with increasing temper-
ature. The pKH1 value of the ligands increases in the order of
HL1 N HL2 N HL3 N HL4.
All the thermodynamic parameters of stepwise stability con-
stants for the complexes of ligands (HLn) are recorded in Table 6.
The obtained values ΔH and ΔS can then be considered as the
sum of two contributions: (a) Release of H2O molecules and
(b) Metal-ligands bond formation. Examination of these values
shows that:
(i) The stability constants (log K1 and log K2) for quinoline azo dyes
complexes increase with increasing temperature, i.e. its stability
constants increase with increasing temperature as shown in
Figs. 7 and 8.
(ii) The negative values of ΔG for the complexes formation suggest a
spontaneous nature of such process [9].

748
A.F. Shoair et al. / Journal of Molecular Liquids 215 (2016) 740–748
Fig. 8. Van't Hoff plot of log K2 of Mn+ complexes with ligands (HLn) against 1/T.
[15] A.A. El-Bindary, A.Z. El-Sonbati, M.A. Diab, E.E. El-Katori, H.A. Seyam, Int. J. Adv. Res.
References
2 (2014) 493–502.
[16] A.A. El-Bindary, A.Z. El-Sonbati, M.A. Diab, Sh.M. Morgan, J. Mol. Liq. 201 (2015)
[1] M.A. Diab, A.A. El-Bindary, A.Z. El-Sonbati, O.L. Salem, J. Mol. Struct. 1018 (2012)
176–184.
[2] N.A. El-Ghamaz, E.M. El-Menyawy, M.A. Diab, A.A. El-Bindary, A.Z. El-Sonbati, S.G.
Nozha, Solid State Sci. 30 (2014) 44–54.
[3] N.A. El-Ghamaz, M.A. Diab, A.A. El-Bindary, A.Z. El-Sonbati, S.G. Nozha, Spectrochim.
Acta A 143 (2015) 200–212.
36–42.
[17] A.Z. El-Sonbati, G.G. Mohamed, A.A. El-Bindary, W.M.I. Hassan, M.A. Diab, Sh.M.
Morgan, A.K. Elkholy, J. Mol. Liq. 212 (2015) 487–502.
[18] A.A. El-Bindary, M.A. Hussein, R.A. El-Boz, J. Mol. Liq. 211 (2015) 256–267.
[19] Z. Bikadi, E. Hazai, J. Chem. Inf. 11 (2009) 1–15.
[20] T.A. Halgren, J. Comput. Chem. 17 (1998) 490–519.
[21] G.M. Morris, D.S. Goodsell, J. Comput. Chem. 19 (1998) 1639–1662.
[22] R.J.B. Wets, Mathematics of Operations Research 6 (1981) 19–30.
[23] A.Z. El-Sonbati, M.A. Diab, A.A. El-Bindary, A.M. Eldesoky, Sh.M. Morgan,
Spectrochim. Acta A 135 (2015) 774–791.
[4] M.A. Diab, A.Z. El-Sonbati, A.A. El-Bindary, A.M. Barakat, Spectrochim. Acta A 116
(2013) 428–439.
[5] C. Karadaş, D. Kara, J. Food Compos. Anal. 32 (2013) 90–98.
[6] M.A. Diab, A.A. El-Bindary, A.Z. El-Sonbati, O.L. Salem, J. Mol. Struct. 1007 (2012)
11–19.
[7] A. Nezhadali, G. Taslimi, Alex. Eng. J. 52 (2013) 797–800.
[8] A.F. Shoair, A.R. El-Shobaky, E.A. Azab, J. Mol. Liq. 203 (2015) 59–65.
[9] A.A. El-Bindary, A.Z. El-Sonbati, M.A. Diab, A.M. Barakat, J. App. Sol. Chem. Mode. 2
(2013) 191–196.
[24] A.Z. El-Sonbati, M.A. Diab, A.A. El-Bindary, G.G. Mohamed, Sh.M. Morgan, Inorg.
Chim. Acta 430 (2015) 96–107.
[25] N.A. El-Ghamaz, A.Z. El-Sonbati, M.A. Diab, A.A. El-Bindary, G.G. Mohamed, Sh.M.
Morgan, Spectrochim. Acta A 147 (2015) 200–211.
[26] N.A. El-Ghamaz, A.Z. El-Sonbati, M.A. Diab, A.A. El-Bindary, M.K. Awad, Sh.M.
Morgan, Mater. Sci. Semicond. Process. 19 (2014) 150–162.
[27] M.M. Ghoneim, A.Z. El-Sonbati, A.A. El-Bindary, M.A. Diab, L.S. Serag, Spectrochim.
Acta A 140 (2015) 111–131.
[28] A.Z. El-Sonbati, A.A. El-Bindary, R.M. Ahmed, J. Sol. Chem. 32 (2003) 617–623.
[29] A.A. El-Bindary, A.F. Shoair, A.Z. El-Sonbati, M.A. Diab, E.E. Abdo, J. Mol. Liq. 212
(2015) 570–584.
[10] A.Z. El-Sonbati, G.G. Mohamed, A.A. El-Bindary, W.M.I. Hassan, A.K. Elkholy, J. Mol.
Liq. 209 (2015) 625–634.
[11] T. Cheng, Q. Li, Z. Zhou, Y. Wang, S. Bryant, AAPS J. 14 (2012) 133–141.
[12] N.A. El-Ghamaz, A.A. El-Bindary, A.Z. El-Sonbati, N.M. Beshry, J. Mol. Liq. 211 (2015)
628–639.
[13] A.Z. El-Sonbati, A.A. El-Bindary, A.F. Shoair, R.M. Younes, Chem. Pharm. Bull. 49
(2001) 1308–1313.
[30] L.E. Orgel, An Introduction to Transition Metal Chemistry Ligand Field Theory, Me-
[14] G.H. Jeffery, J. Bassett, J. Mendham, R.C. Deney, Vogel's Textbook of Quantitative
Chemical Analysis, fifth ed. Longman, London, 1989.
thuen, London, 1966, p. 55.