Potentiometry and geometrical structure of some azodye compounds and their metal complexes

Article abstract of DOI:10.1016/j.molliq.2015.05.059

A series of ligands of quinoline azodye derivatives (HL_I-HL_VI) have been prepared and characterized by elemental analysis, X-ray diffraction analysis, ¹H NMR and IR spectra. The proton-ligand dissociation constant of the l

Full text of DOI:10.1016/j.molliq.2015.05.059

Journal of Molecular Liquids 209 (2015) 258–266
Contents lists available at ScienceDirect
Journal of Molecular Liquids
journal homepage: www.elsevier.com/locate/molliq
Potentiometry and geometrical structure of some azodye compounds
and their metal complexes
⁎
Amina A. Soayed , Amel F. El-Husseiny
Department of Chemistry, Faculty of Science, Alexandria University, P.O. Box 426, Ibrahimia, Alexandria 21321, Egypt
a r t i c l e i n f o
a b s t r a c t
Article history:
A series of ligands of quinoline azodye derivatives (HLI–HLVI) have been prepared and characterized by elemental
analysis, X-ray diffraction analysis, 1H NMR and IR spectra. The proton–ligand dissociation constant of the ligands
(HLI–HLVI) and the metal stability constant of their Mn2+, Co2+, Ni2+, Cu2+, La3+, UO22+ and Th4+ complexes
have been determined potentiometrically in 1 M KC1 and 40% (V:V) EtOH–water mixture. The molecular struc-
ture of each ligand was optimized theoretically and the quantum chemical parameters were calculated. The in-
fluence of derivatives on the dissociation and stability constants was examined on the basis of the electron
repelling property of the derivatives. Order of the stability constant of the formed complexes was found to be
Mn2+ Co2+ Ni2+ Cu2+ La3+ UO22+ Th4+. The effect of temperature was studied and the corresponding ther-
modynamic parameters (ΔG, ΔH and ΔS) were derived and discussed. The dissociation process was found to be
non-spontaneous, endothermic and entropically unfavorable. On the other hand the formation of the metal com-
plexes was found to be spontaneous, endothermic and entropically favorable.
Received 15 April 2015
Received in revised form 25 May 2015
Accepted 29 May 2015
Available online 9 June 2015
Keywords:
Quinoline azodyes
Potentiometry
Stability constants
Thermodynamics parameters
Quantum chemical parameters
© 2015 Elsevier B.V. All rights reserved.
1. Introduction
In the present paper the potentiometric studies were used to
determine the dissociation constants of quinoline azo dye derivatives
Azo compounds containing heterocyclic moieties have drawn the at-
tention of many researchers. The chemistry of quinoline and its deriva-
tives has attracted special interest due to their electrical and optical
properties [1–5]. The electrical conductivities of quinoline compounds
have received considerable attention [6] because of their physiological
activity, analytical and industrial applications. Some of 8-hydro-
xyquinoline and azo derivatives found numerous applications in analyt-
ical chemistry as chromophoric and metallochromic indicators [7].
Moreover, azo compounds based on quinolone, 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 complexation with the metal ion
[8–10]. The azo compounds are used in dyeing processes and some of
them are used in the analytical separation of many metal ions present
in a mixture [11]. The effect of temperature on the proton–ligand disso-
ciation constants of a series of azo dyes was previously studied potenti-
ometrically [12–15] and the stability constants of their complexes were
found to be in the order Mn+2 b Co+2 b Ni2+ b Cu2+. The corresponding
thermodynamic parameters (ΔG, ΔH, ΔS) were also derived. The stoi-
chiometries of the complexes were determined conductometrically
and indicated the formation of 1:1 and 1:2 (metal:ligand) complexes.
(HLI–HLVI) and the stability constants of their Mn2+, Co2+, Ni2+
,
Cu2+, La3+, UO22+ and Th4+ complexes at different temperatures.
The molecular structures of the investigated ligands were optimized
theoretically and the quantum chemical parameters were calculated.
Moreover, the corresponding thermodynamic functions were also
calculated and discussed.
2. Materials and methods
2.1. Preparation of quinoline azodye derivatives (HLI–HLVI)
Ligands of quinoline azo dye derivatives (HLI–HLVI) were prepared
according to the following procedure. In a typical preparation, 25 ml
of distilled water containing 0.01 mol hydrochloric acid was added to
aniline (0.01 mol) or aniline derivatives. A solution of 0.01 mol sodium
nitrite in 20 ml of water was added drop-wise to the resulting mixture
then stirred and cooled to 0 °C. The formed diazonium chloride was con-
secutively coupled with an alkaline solution of 0.01 mol quinolin-8-ol,
in 10 ml of pyridine. The preparation of ligands (HLI–HLVI) is summa-
rized in Scheme 1. The colored precipitates, which formed immediately,
were filtered through sintered glass crucible, washed several times with
water. The crude products were purified by recrystallization from hot
ethanol, (yield ~75%) then dried in a vacuum desiccator over P2O5.
⁎
Corresponding author.
E-mail addresses: amsoayed@yahoo.com, aminasoayed@yahoo.com (A.A. Soayed).
http://dx.doi.org/10.1016/j.molliq.2015.05.059
0167-7322/© 2015 Elsevier B.V. All rights reserved.

A.A. Soayed, A.F. El-Husseiny / Journal of Molecular Liquids 209 (2015) 258–266
259
Where R =
HLI
X = H
HLII
HLIII
HLIV
HLV
HLVI
X = o-Cl
X = o-CH3 and p-OCH3
X = m-NO2 and p-CH3
X = o-OCH3 and p-NO2
X = o-NO2 and p-OCH3
Scheme 1. Synthesis of azo 8-hydroxyquinoline dyes HLI–HLVI
.
The ligands were characterized by elemental analysis (Table 1) and IR
spectroscopy.
NaOH solution in 40% (by volume) EtOH–water mixture was used as ti-
trant and standardized against oxalic acid.
The apparatus, general conditions and methods of calculations were
the same as in previous works [17,18]. The following mixtures (i)–(iii)
were prepared and titrated potentiometrically at 298 K against standard
0.004 M NaOH in a 40% (by volume) DMF–water mixture:
2.2. Potentiometric studies
The ligand solution (0.001 M) was prepared by dissolving an accu-
rately weighted amount of the solid in DMF. Metal ion solutions
(0.001 M) were prepared from metal chlorides in doubly distilled
water and standardized with EDTA [16]. Solutions of 0.001 M HCl and
1 M KCl were also prepared in doubly distilled water. A carbonate-free
(i) 5 ml 0.001 M HC1 + 5 ml 1.0 M KC1 + 20 ml EtOH
(ii) 5 ml 0.001 M HC1 + 5 ml 1.0 M KC1 + 15 ml EtOH + 5 ml
0.001 M ligand
Table 1
Analytical data of quinoline azodyesa.
4000
3000
2000
1000
0
Compoundb Empirical
formula
M.p.(°C)
Formula Calc. (exp.) %
weight
(g/mol)
C
H
N
HLI
C15H11N3O
267
249.27
72.30
4.50
16.90
(72.20) (4.40) (16.60)
63.50 3.60 14.80
(63.40) (3.40) (14.70)
69.61 5.15 14.33
(69.53) (5.05) (14.20)
59.26 3.73 17.28
(59.15) (3.64) (17.15)
62.33 3.92 18.17
(63.19) (3.80) (18.15)
62.33 3.92 18.17
(63.15) (3.82) (18.07)
HLII
HLIII
HLIV
HLV
HLVI
C15H10ClN3O 273
283.05
293.12
324.29
308.29
308.29
C17H15N3O2
C16H12N4O4
C16H12N4O3
C16H12N4O3
276
287
281
279
a
Further studies with title ligands, using different metals, are in progress and will be
published in due coarse.
0
20
40
θ (ο)
60
80
b
The analytical data agrees satisfactory with the expected formulae represented as
2
given in Scheme 1 (see experimental). HLI–HLVII are the ligand as given in Scheme 1. Air-
stable, high melting temperature, colored, insoluble in water, but soluble in hot ethanol,
and coordinate solvent. Decomposition near 275 °C.
Fig. 1. X-ray diffraction pattern of HLI powder form.

260
A.A. Soayed, A.F. El-Husseiny / Journal of Molecular Liquids 209 (2015) 258–266
HLI
HLII
Carbon atom
Oxygen atom
Hydrogen atom
Nitrogen atom
Chlorine atom
Fig. 2. The optimized structures of ligands (HLI and HLII) within numbering of atoms.
(iii) 5 ml 0.001 M HC1 + 5 ml 1.0 M KC1 + 15 ml EtOH + 5 ml
0.001 M ligand + 1 ml 0.001 M metal chloride
2.3. Measurements
All the compounds and solvents used were purchased from Aldrich
and Sigma and were used as received without further purification. The
IR spectra were recorded as KBr discs using a Perkin-Elmer 1340 spec-
trophotometer. The 1H-NMR spectra was performed using a JEOL FX90
Fourier transform spectrometer with DMSO-d6 as the solvent and TMS
as an internal reference. X-ray diffraction analysis of the ligand (HL)
powder form was recorded on X-ray diffractometer in the range of
diffraction angle 2θ° = 4–70°. This analysis was carried out using Cu
For each mixture, the volume was made up to 50 ml with doubly
distilled water before titration. These titrations were repeated at tem-
peratures 308 and 318 K. A constant temperature was maintained at
0.05 K by using an ultra-thermostat (Neslab 2 RTE 220). The
pH-meter readings in 40% (v/v) EtOH–water mixture were corrected
according to the Van Uitert and Hass relation [19].
Table 2
Calculated quantum chemical parameters for HLI and HLII ligands.
Compound
EHOMO (eV)
ELUMO (eV)
ΔE (eV)
Χ (eV)
η (eV)
σ (eV)−1
Pi (eV)
S (eV)−1
ω (eV)
ΔNmax
HLI
HLII
−5.021
−4.643
−2.030
−1.136
2.991
3.507
3.525
2.889
1.495
1.754
0.669
0.570
−3.525
−2.889
0.334
0.285
4.156
2.381
2.358
1.648

A.A. Soayed, A.F. El-Husseiny / Journal of Molecular Liquids 209 (2015) 258–266
261
Table 3
value favors the formation of an intramolecular hydrogen bond between
C8–OH and the N in quinoline moiety. Electron-withdrawing substitu-
ents reduce considerably the intramolecular hydrogen bond as indicated
by the marked shift of the hydroxyl signal to a higher field in the NO2
compound. Electron-donating substituents give the opposite effect,
arising from the increasing basicity of the azo-nitrogen.
It is known that 8-hydroxyquinoline exists in solution in a mono-
mer–dimer equilibrium [23]. In the monomeric form, intramolecular
H-bond is present whereas the strong 950 cm−1 band indicates the
existence of this ligand in a dimeric structure through intermolecular
hydrogen bonding. This is in agreement with previous results [24].
The 2800–3550 cm−1 region is characteristic for the vibrational fre-
quency of the –OH group. The bands in the 3400–3500 cm−1 region
are assigned to intramolecular hydrogen bonding. Thus, HLn ligands
can exist as a five-membered chelate skeleton with hydrogen bonding
classified into two types:
Selected geometrical parameters for ligand (HLI).
Bond lengths (Å)
Bond angles (°)
Bond angles (°)
C(18)–H(30) 1.103 H(29)–C(17)–C(18) 120.031 O(12)–C(3)–C(4) 134.668
C(17)–H(29) 1.103 H(29)–C(17)–C(16) 119.644 N(10)–C(2)–C(3) 101.141
C(16)–H(28) 1.103 C(18)–C(17)–C(16) 120.325 N(10)–C(2)–C(1) 130.758
C(15)–H(27) 1.103 H(28)–C(16)–C(17) 120.321 C(3)–C(2)–C(1) 127.655
C(14)–H(26) 1.105 H(28)–C(16)–C(15) 120.291 C(2)–C(1)–C(6) 114.414
C(9)–H(25)
C(8)–H(24)
C(7)–H(23)
C(5)–H(22)
C(4)–H(21)
C(13)–C(18) 1.348 H(30)–C(18)–C(17) 117.031
C(17)–C(18) 1.343 C(13)–C(18)–C(17) 121.088
C(16)–C(17) 1.341 H(26)–C(14)–C(15) 118.158
C(15)–C(16) 1.341 H(26)–C(14)–C(13) 120.362
C(14)–C(15) 1.342 C(15)–C(14)–C(13) 121.48
C(13)–C(14) 1.348 C(18)–C(13)–C(14) 117.734
1.102 C(17)–C(16)–C(15) 119.388 C(2)–C(1)–C(7) 110.416
1.103 H(27)–C(15)–C(16) 119.897 C(6)–C(1)–C(7) 135.144
1.099 H(27)–C(15)–C(14) 120.117
1.106 C(16)–C(15)–C(14) 119.985
1.101 H(30)–C(18)–C(13) 121.88
N(10)–C(2)
C(3)–C(2)
C(1)–C(2)
C(6)–C(1)
C(7)–C(1)
C(8)–C(7)
C(9)–C(8)
N(10)–C(9)
H(20)–N(10) 1.035 H(24)–C(8)–C(9)
O(12)–H(20) 1.003 C(7)–C(8)–C(9)
C(3)–O(12)
C(4)–C(3)
C(5)–C(4)
C(6)–C(5)
1.256 C(18)–C(13)–N(19) 125.816
1.338 C(14)–C(13)–N(19) 116.45
1.332 C(13)–N(19)–N(11) 119.971
1.355 N(10)–H(20)–O(12) 147.701
(a) An intramolecular hydrogen bond (O–H……N) formed between
hydrogen of (C8–OH) and nitrogen of quinoline ring.
(b) An intermolecular hydrogen bonding (O–H…..O) formed be-
tween the two C8–OH groups of two different hydroxyl–quinoline
groups or OH of one ring and quinoline N of the other.
1.35
1.354 H(25)–C(9)–N(10)
1.35 C(8)–C(9)–N(10)
1.262 H(24)–C(8)–C(7)
H(25)–C(9)–C(8)
122.571
119.611
117.721
119.654
119.052
121.272
121.4
122.839
116.753
120.015
118.296
128.234
113.431
115.602
119.323
125.001
90.02
The band within the 1100 cm−1 could be due to υC–OH [26]. The
high frequency value of the δOH may account for the existence of
hydrogen bonding.
1.238 N(19)–N(11)–C(6)
1.335 H(23)–C(7)–C(1)
IR spectra exhibited also a medium to strong band in the region
1.35
H(23)–C(7)–C(8)
1540–1570 cm−1 range assigned to υN_N stretching vibrations
1.361 C(1)–C(7)–C(8)
[8,9,25] and a very strong band at 1585
10 cm−1 assignable to
C(13)–N(19) 1.269 C(1)–C(6)–C(5)
N(11)–N(19) 1.249 C(1)–C(6)–N(11)
C(6)–N(11)
υC_C of the phenyl ring.
Analytical data (Table 1) are in agreement with the proposed
formulae.
1.27
C(5)–C(6)–N(11)
H(22)–C(5)–C(4)
H(22)–C(5)–C(6)
C(4)–C(5)–C(6)
H(20)–O(12)–C(3)
H(21)–C(4)–C(3)
H(21)–C(4)–C(5)
C(3)–C(4)–C(5)
C(2)–N(10)–C(9)
C(2)–N(10)–H(20)
C(9)–N(10)–H(20)
C(2)–C(3)–O(12)
C(2)–C(3)–C(4)
121.286
122.441
116.261
118.015
92.722
3.2. X-ray diffraction analysis
Single crystals of the ligands could not be prepared, hence powder
diffraction data were used for structural characterization. Structural de-
termination by X-ray powder diffraction data has gone through a recent
surge so as to become important to get to the structural information of
materials that do not yield good quality single crystals.
139.312
107.584
117.659
The X-ray diffraction (XRD) patterns of HLI ligand in the powder
form is shown in Fig. 1. The XRD pattern shows many diffraction
peaks indicating a polycrystalline phase for this ligand. The average
crystallite size (ξ) can be calculated from the XRD patterns according
to Debye–Scherrer equation:
Kα radiation (λ = 1.541874 Å). The applied voltage and the tube current
are 40 kV and 30 mA, respectively.
The calculations of geometry optimization were performed using
Perkin Elmer ChemBio 3D software by HF method with 3-21G basis
set [20]. Geometry optimization option was employed to obtain the
most stable structure.
Kλ
The pH measurements were carried out using VWR Scientific
ξ ¼
:
β
1=2 cosθ
Instruments Model 8000 pH-meter accurate to
0.01 units. The
pH-meter readings in the non–aqueous medium were corrected [21].
The electrode system was calibrated according to the method of Irving
et al. [22]. Titrations were performed in a double walled glass cell in
an inert atmosphere (nitrogen) at ionic strength of 0.1 M KCl. Potentio-
metric measurements were carried out at different temperatures. The
The equation uses the reference peak width at angle (θ), where λ is
wavelength of X-ray radiation (1.540598 Å), K is constant taken as 0.95
for organic compounds [16] and β1/2 is the width at half maximum of
the reference diffraction peak measured in radians. The dislocation
density (δ) is the number of dislocation lines per unit area of the crystal.
The value of δ is related to the average particle diameter (ξ) by the
relation [27]:
temperatures were controlled to within
0.05 K by circulating
thermostated water (Neslab 2 RTE 220) through the outer jacket of
the vessel.
1
ξ2
3. Results and discussion
δ ¼
:
3.1. Structural determination of the ligands
The 1H NMR spectra of the ligands showed a signal in the region
9.3–9.7 ppm assigned to C8–OH of the hydroxy quinoline ring. This
The values of ξ and δ for HLI ligand were calculated and found to be
24.3 nm and 1.69 × 10−3 nm−2, respectively.

262
A.A. Soayed, A.F. El-Husseiny / Journal of Molecular Liquids 209 (2015) 258–266
LUMO
HOMO
HLI
ΔE = 2.991 eV
HLII
ΔE = 3.507 eV
Fig. 3. HOMO and LUMO molecular orbital of ligands (HLI and HLII).
3.3. Geometrical structures
1
S ¼
ð6Þ
2η
The optimized structures of the ligands (HLI and HLII) are presented
in Fig. 2. Selected geometric parameter bond lengths and bond angles of
ligands (HLI and HLII) are tabulated in Tables 2 and 3. Both the highest
occupied molecular orbital (HOMO) and the lowest unoccupied molec-
ular orbital (LUMO) are the main orbitals that take part in chemical
stability. The HOMO represents the ability to donate an electron,
LUMO as an electron acceptor represents the ability to obtain an elec-
tron. Molecular orbital structures (HOMO & LUMO) for ligands (HLI
and HLII) are presented in Fig. 3. Electronic absorption corresponds to
the transition from the ground state to the first excited state and is
mainly described by one electron excitation from the highest occupied
molecular orbital to the lowest unoccupied molecular orbital. Recently,
the energy gap between HOMO and LUMO has been used to prove the
activity and stability of the compounds [18]. Quantum chemical param-
eters of the ligands (HLI and HLII) were obtained from calculations in-
Pi2
2η
ω ¼
ð7Þ
ð8Þ
Pi
η
ΔNmax ¼ −
:
The value of ΔE for HLI and HLII ligands was found to be 2.991 and
3.507 eV, respectively. The calculations indicated that the HLI ligand is
more stable and highly reactive than HLII ligand.
3.4. Potentiometric studies
3.4.1. Proton–ligand stability constants
cluding the energies of the highest occupied molecular orbital (EHOMO
)
The average number of the protons associated with ligands (HLI–
HLVI) at different pH values, nA, was calculated from the titration curves
of the acid in the absence and presence of ligands (HLI–HLVI) by apply-
ing the following equation:
and the lowest unoccupied molecular orbital (ELUMO) as listed in
⁎
Table 4. Additional parameters such as HOMO–LUMO energy gap, ΔE ,
absolute electronegativities, χ, chemical potentials, Pi, absolute hard-
ness, η, absolute softness, σ, global electrophilicity, ω, global softness,
S, and additional electronic charge, ΔNmax, were calculated using the
following equations [17,28]:
ꢀ
ꢁ
ðV1−V2Þ No þ Eo
ꢀ
ꢁ
nA ¼ Y ꢀ
ð9Þ
Vo−V1 TCoL
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 concen-
tration of the reagent, N° is the normality of sodium hydroxide solu-
tion and E° is the initial concentration of the free acid. Thus, the
formation curves (nA vs. pH) for the proton–ligand systems were
constructed and found to extend between 0 and 1 in the nA scale.
This means that each ligand has one ionizable proton (the enolized
hydrogen ion of the hydroxyl group in the quinoline moiety). It can
be seen that for the same volume of KOH added, the ligand titration
curves had a lower pH value than the acid titration curve. The dis-
placement of a ligand titration curve along the volume axis with
ΔE ¼ ELUMO−EHOMO
ð1Þ
ð2Þ
−ðEHOMO þ ELUMO
Þ
χ ¼
2
ELUMO−EHOMO
η ¼
ð3Þ
2
1
σ ¼
η
ð4Þ
ð5Þ
Pi ¼ −χ

A.A. Soayed, A.F. El-Husseiny / Journal of Molecular Liquids 209 (2015) 258–266
263
Table 4
The phenolic –OH group is known to be weakly acidic, indicating a
stronger bonding between the proton and the oxygen atom. This
means that the proton–ligand stability constant of quinoline azo dyes
should be high due to the dissociation of the –OH group [30].
Selected geometrical parameters for ligand (HLII).
Bond lengths (Å)
Bond angles (°)
Bond angles (°)
C(17)–H(30) 1.104 H(30)–C(17)–C(18) 121.063 N(10)–C(2)–C(3) 115.818
C(16)–H(29) 1.103 H(30)–C(17)–C(16) 117.653 C(1)–C(2)–C(3) 121.274
C(15)–H(28) 1.103 C(18)–C(17)–C(16) 121.285 C(2)–C(1)–C(6) 120.679
C(14)–H(27) 1.106 H(29)–C(16)–C(17) 120.589 C(2)–C(1)–C(7) 116.597
C(9)–H(26)
C(8)–H(25)
C(7)–H(24)
C(5)–H(23)
C(4)–H(22)
C(13)–C(18) 1.354 C(13)–C(18)–C(17) 120.472
C(17)–C(18) 1.346 C(13)–C(18)–Cl(21) 125.001
C(16)–C(17) 1.34
C(15)–C(16) 1.338 H(27)–C(14)–C(15) 116.817
C(14)–C(15) 1.341 H(27)–C(14)–C(13) 120.611
C(13)–C(14) 1.352 C(15)–C(14)–C(13) 122.572
C(9)–N(10)
C(2)–N(10)
C(1)–C(2)
C(3)–C(2)
C(4)–C(3)
C(5)–C(4)
C(6)–C(5)
C(1)–C(6)
C(7)–C(1)
C(8)–C(7)
C(9)–C(8)
C(18)–Cl(21) 1.733 H(24)–C(7)–C(1)
O(12)–H(20) 0.971 H(24)–C(7)–C(8)
C(13)–N(19) 1.271 C(1)–C(7)–C(8)
N(11)–N(19) 1.25
C(6)–N(11)
C(3)–O(12)
An inspection of the results in Table 5 reveals that the pKH values of
HLI and its substituted derivatives are influenced by the inductive or
mesomeric effect of the substituents [31]. The o-OCH3 and o-CH3 deriv-
atives (HLIII and HLV) have a lower acidic character (higher pKH values)
than the o-Cl and o-NO2 derivatives (HLII and HLVI). This is quite reason-
able because the presence of o-OCH3 and p-CH3 groups (i.e. an electron
donating group) will enhance the electron density by its high positive
inductive or mesomeric effect, whereby a stronger O–H bond is formed.
The presence of o-Cl and o-NO2 groups (i.e. electron withdrawing
groups) will lead to the opposite effect. The high acidic character
found in the HLIV ligand (m-NO2 and p-CH3 groups) is due to the
enhancement of the electron density by the high positive inductive or
mesomeric effect, whereby a stronger O–H bond is formed. The ortho
and para substituents in the quinoline moiety have a direct influence
on the pKH values of the investigated compounds, revealing the copla-
narity of the molecule and thus affording a maximum resonance via
delocalization of its π-system.
1.103 H(29)–C(16)–C(15) 120.329 C(6)–C(1)–C(7) 122.724
1.102 C(17)–C(16)–C(15) 119.082
H(28)–C(15)–C(16) 120.025
1.103 H(28)–C(15)–C(14) 120.418
1.103 C(16)–C(15)–C(14) 119.557
1.1
C(17)–C(18)–Cl(21) 114.527
1.263 C(18)–C(13)–C(14) 117.032
1.269 C(18)–C(13)–N(19) 130.429
1.352 C(14)–C(13)–N(19) 112.539
1.347 C(13)–N(19)–N(11) 123.567
1.341 H(26)–C(9)–N(10) 116.689
1.34
H(26)–C(9)–C(8)
120.338
122.973
121.39
121.039
117.571
1.349 N(10)–C(9)–C(8)
1.355 H(25)–C(8)–C(7)
1.349 H(25)–C(8)–C(9)
C(7)–C(8)–C(9)
1.339 N(19)–N(11)–C(6) 120.172
3.4.2. Metal–ligand stability constants
The formation curves for the metal complexes were obtained by
plotting the average number of ligands attached per metal ion (nA) vs.
the free ligand exponent (pL), according to Irving and Rossotti [32].
These curves were analyzed and the successive stability constants
were determined using different computational methods [33,34]
which agree within 1% error. The average number of the reagent mole-
cules attached per metal ion, n, and free ligand exponent, pL, can be cal-
culated using Eqs. (10) and (11):
1.34
122.951
117.036
120.013
117.326
123.255
119.419
115.82
C(5)–C(6)–C(1)
1.272 C(5)–C(6)–N(11)
1.362 C(1)–C(6)–N(11)
H(23)–C(5)–C(4)
H(23)–C(5)–C(6)
C(4)–C(5)–C(6)
H(20)–O(12)–C(3) 109.344
122.518
121.662
ꢀ
ꢁ
ðV3−V2Þ No þ Eo
ꢀ
ꢁ
n ¼
ð10Þ
Vo−V2 ꢁ nA ꢁ TCo
H(22)–C(4)–C(3)
H(22)–C(4)–C(5)
C(3)–C(4)–C(5)
C(9)–N(10)–C(2)
C(2)–C(3)–C(4)
C(2)–C(3)–O(12)
C(4)–C(3)–O(12)
N(10)–C(2)–C(1)
119.046
119.72
M
121.234
119.937
117.824
121.445
120.731
122.908
and
!
n
n¼J
X
1
βnH
ꢂ
ꢃ
Hþ
Vo þ V3
n¼o
pL ¼ log10
ꢁ
ð11Þ
TCoL −n ꢁ TCoM
Vo
respect to the acid titration curve is an indication of proton dissocia-
tion. The proton–ligand stability constants were calculated using the
method of lrving and Rossotti [29].
where TC°M is the total concentration of the metal ion present in the so-
lution, and 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
Table 5
Thermodynamic function of azoquinoline in 40% (v/v) EtOH–water mixture and 1 M KCl at different temperatures.
Comp.
Temp. (K)
Dissociation constant pKH
Enthalpy change
Free energy change
Entropy change
(kJ mol−1) ΔG
(kJ mol−1) ΔH
(J mol−1 K−1) −ΔS
HLI
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
298
308
318
9.40
9.25
9.10
8.30
8.16
8.00
9.60
9.45
9.32
9.80
9.66
9.51
9.70
9.55
9.40
8.50
8.36
8.22
53.63
54.55
55.41
47.36
48.12
48.71
54.78
55.73
56.75
55.92
56.97
57.90
55.27
56.25
57.16
48.50
49.30
50.05
27.21
27.19
25.41
26.29
27.21
25.39
88.658
88.766
88.679
67.684
67.955
67.673
98.557
98.442
98.553
99.430
99.610
99.403
94.161
94.286
94.182
77.550
77.630
77.547
HLII
HLIII
HLIV
HLV
HLVI

264
A.A. Soayed, A.F. El-Husseiny / Journal of Molecular Liquids 209 (2015) 258–266
(b)
(a)
(b)
(a)
2+
Cu
2+
Cu
11.9
9.7
HL
HL
HL
I
11.1
V
9.2
11.8
11.7
11.6
11.5
11.4
11.3
11.2
11.1
11.0
9.6
9.5
9.4
9.3
9.2
9.1
9.0
8.9
HL
II
11.0
10.9
10.8
10.7
10.6
10.5
10.4
10.3
2+
Ni
9.1
9.0
8.9
8.8
8.7
8.6
8.5
VI
2+
Ni
2+
Co
2+
Co
2+
Mn
2+
Mn
T=308 K
T=308 K
25
26
27
28
29
25
26
27
28
29
Atomic number
Atomic number
Fig. 4. The relation between stability constant (log β) and atomic number of metal com-
plexes for ligands a) HLI and b) HLII.
Fig. 6. The relation between stability constant (log β) and atomic number of metal com-
plexes for a) HLV and b) HLVI
.
titration curves of hydrochloric acid, organic ligand and complex, re-
spectively. The following general remarks may be pointed out:
(a)
(b)
2+
Cu
HL
I
51
62
HL
II
(i) The maximum value of n was ≈ 2 indicating the formation of 1:1
and 1:2 (metal:ligand) complexes only.
(ii) The metal ion solution used in the present study was very dilute
(2 × 10−4 M); hence there was no possibility for the formation of
polynuclear complexes [35].
(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 ligand. The large decrease in pH for the metal titration curves
relative to ligand titration curves points to the formation of
strong metal complexes [36].
(iv) In most cases, the color of the solution after complex formation
was observed to be different from the color of the ligand at the
same pH.
50
49
48
47
46
2+
61
60
59
58
57
Ni
2+
Co
2+
Mn
T=298 K
25
26
27
28
29
Atomic number
(v) For the same ligand at constant temperature, the stability of the
chelates increases in the order Mn2+ Co2+ Ni2+ Cu2+ La3+
UO22+ Th4+ [14,37]. This order largely reflects the changes in
the heat of complex formation across the series from a combina-
tion of the influence of both the polarizing ability of the metal
Fig. 7. The relation between free energy change (ΔGβ) and atomic number of metal com-
plexes for ligands a) HLI and b) HLII.
ion [38] and the crystal field stabilization energies [39]. The great-
er stability of Cu2+ complexes is produced by the well-known
Jahn–Teller effect.
(b)
(b)
(a)
(a)
65
64
63
62
61
60
59
58
2+
Cu
2+
Cu
HL
HL
11.6
11.5
11.4
11.3
11.2
11.1
11.0
10.9
10.8
10.7
12.1
12.0
11.9
11.8
11.7
11.6
11.5
11.4
11.3
11.2
HL
HL
III
51
50
49
48
47
46
III
IV
IV
2+
Ni
2+
Ni
2+
Co
2+
Co
2+
Mn
2+
Mn
T=298 K
T=308 K
25
26
27
28
29
25
26
27
28
29
Atomic number
Atomic number
Fig. 5. The relation between stability constant (log β) and atomic number of metal com-
plexes for ligands a) HLIII and b) HLIV
Fig. 8. The relation between free energy change (ΔGβ) and atomic number of metal com-
plexes for ligands a) HLIII and b) HLIV
.
.

A.A. Soayed, A.F. El-Husseiny / Journal of Molecular Liquids 209 (2015) 258–266
265
(a)
(b)
From the values of free energy change (ΔG) and enthalpy change
(ΔH), we can deduce the entropy (ΔS) using the well-known relation-
ships (12) and (14):
2+
Cu
68
67
HL
V
67
66
65
64
63
62
61
66
65
64
63
62
61
60
HL
VI
2+
Ni
ΔS ¼ ðΔH–ΔGÞ=T:
ð14Þ
2+
Co
All thermodynamic parameters of the dissociation process of ligands
(HLI–HLVI) are recorded in Table 5. From these results the following con-
clusions can be made:
2+
Mn
(i) The pKH values decrease with increasing temperature, i.e., the
acidity of the ligands increases, independent of the nature of
the substituent [17,36].
(ii) A positive value of ΔH indicates that the process is endothermic.
(iii) A large positive value of ΔG indicates that the dissociation pro-
cess is not spontaneous.
(iv) The dissociation processes for the compounds have negative
values of ΔS due to the increased order as a result of the solvation
processes.
T=298 K
25
26
27
28
29
Atomic number
Fig. 9. The relation between free energy change (ΔGβ) and atomic number of metal com-
plexes for ligands a) HLV and b) HLVI
.
An inspection of the results reveals that the stability constant values
of the complexes of HLIV, HLV and HLVI are influenced by the inductive or
mesomeric effect of the substituents. This behavior correlates with the
effect of substitution on the phenyl ring as follows:
It is known that the divalent metal ions exist in solution as octahe-
dral hydrated species and the obtained values of ΔH and ΔS can then
be considered as the sum of two contributions: (a) release of H2O mol-
ecules and (b) metal–ligand bond formation. Examination of these
values shows that:
(i) The high stability of HLIV complexes can be attributed to the pres-
ence of the p-CH3 group relative to the azo group. This is quite
reasonable because of the presence p-CH3 group (i.e. an electron
donating effect), whereby stronger chelation was formed and
therefore the stability of the complexes.
(ii) The stability of HLV complexes can be attributed to the presence
of the p-NO2 and o-OCH3 groups relative to the azo group.
(iii) The low stability of HLVI complexes can be attributed to the pres-
ence of the NO2 group in the o-position relative to the azo group.
This is caused by the negative indicative effect of the NO2 group,
which decreases its ability for chelation, and therefore the stabil-
ity of the complexes.
(i) The stepwise stability constant (log K1 and log K2) for ligand
complexes increases with increasing temperature, i.e., its stabili-
ty constants increase with increasing the temperature.
(ii) The negative value of ΔG for the complexation process suggests
the spontaneous nature of such process.
(iii) The values of free energy change of formed complexes were in-
creased with increasing atomic number as shown in Figs. 7–9.
(iv) The ΔH values are positive, meaning that these processes are en-
dothermic and favorable at higher temperature,
(v) The ΔS values for the ligand complexes are positive, confirming
that the complex formation is entropically favorable.
(iv) For the ligands with the same metal ion at constant temperature,
the stability of the chelates increases in the order HLI N HLII N HLIII
[36].
4. Conclusion
(v) The stability constants of Mn2+, Co2+, Ni2+ and Cu2+ complexes
were increased with increasing atomic number in the order
Cu2+ N Ni2+ N Co2+ N Mn2+ at constant temperature as shown
in Figs. 4–6.
The proton–ligand dissociation constant of quinoline azo-dye deriv-
atives (HLI–HLVI) and metal–ligand stability constants of their com-
plexes with metal ions (Mn2+, Co2+, Ni2+, Cu2+, La3+, UO22+ and
Th4+) at different temperatures were determined. At constant temper-
ature the stability constants of the formed complexes decreases in the
order Cu2+, Ni2+, Co2+ and Mn2+. The dissociation process is non-
spontaneous, endothermic and entropically unfavorable. The formation
of the metal complexes has been found to be spontaneous, endothermic
and entropically favorable. The values of stability constants of quinoline
azodye derivatives (HLI–HLVI) complexes increase with increasing
temperature.
3.4.3. Effect of temperature
The dissociation constant (pKH) for azoquinoline and its substituted
derivatives, as well as the stability constants of their complexes with
Mn2+, Co2+, Ni2+, Cu2+, La3+, UO22+ and Th4+ have been evaluated
at 298, 308 and 318 K. The enthalpy (ΔH) for the dissociation and com-
plexation process was calculated from the slope of the plot pKH or log K
vs. 1
(13):
/
T using the graphical representation of Van't Hoff Eqs. (12) and
Appendix A. Supplementary data
ΔG ¼ −2:303 RT log K ¼ ΔH−TΔS
ð12Þ
Supplementary data to this article can be found online at http://dx.
doi.org/10.1016/j.molliq.2015.05.059.
or
ꢄ
ꢅꢄ ꢅ
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−ΔH
2:303R
1
T
ΔS
log K ¼
þ
2:303R
ð13Þ
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