Comparison of chelating ability of NO-, NS-, ONS-, and ONO-type Schiff base derivatives and their stability constants of Bis-complexes with copper(II)

Article abstract of DOI:10.3906/kim-1303-65

The present study includes important findings relating to the number of donor atoms, species of ligands, and stabilities of complexes. Stabilities of complexes between Cu(II) ion and NO-, NS-, ONS-, and ONO-type Schiff bases were compared. Acid-base properties of the Schiff bases were explained at 25 ± 0.1 ° C and ionic strength (I) of 0.1 M supported by NaCl. The Hyperquad computer program was used for calculation of dissociation and stability constants. The overall stability constants of their Cu(II) complexes were calculated and the various formed complexes between the Schiff bases with Cu(II) ion formulated as CuL₂, CuHL₂, CuH₂L ₂, and CuH_-1L₂(Cu (OH) L₂). The complexes of ONS- and ONO-type tridentate ligands were more stable than those of NO- and NS-type bidentate ligands. TUeBITAK.

Full text of DOI:10.3906/kim-1303-65

Turk J Chem
Turkish Journal of Chemistry
(2014) 38: 109 – 120
¨
http://journals.tubitak.gov.tr/chem/
˙
c
⃝ TUBITAK
doi:10.3906/kim-1303-65
Research Article
Comparison of chelating ability of NO-, NS-, ONS-, and ONO-type Schiff base
derivatives and their stability constants of Bis-complexes with copper(II)
Hasan ATABEY∗, Esra FINDIK, Hayati SARI, Mustafa CEYLAN
Chemistry Department, Science and Arts Faculty, Gaziosmanpa¸sa University, Tokat, Turkey
Received: 21.03.2013
•
Accepted: 09.07.2013
•
Published Online: 16.12.2013
•
Printed: 20.01.2014
Abstract: The present study includes important findings relating to the number of donor atoms, species of ligands, and
stabilities of complexes. Stabilities of complexes between Cu(II) ion and NO-, NS-, ONS-, and ONO-type Schiff bases
were compared. Acid-base properties of the Schiff bases were explained at 25 ± 0.1 ◦ C and ionic strength (I) of 0.1 M
supported by NaCl. The Hyperquad computer program was used for calculation of dissociation and stability constants.
The overall stability constants of their Cu(II) complexes were calculated and the various formed complexes between the
Schiff bases with Cu(II) ion formulated as CuL2 , CuHL2 , CuH2 L2 , and CuH−1 L2 (Cu (OH) L2). The complexes of
ONS- and ONO-type tridentate ligands were more stable than those of NO- and NS-type bidentate ligands.
Key words: Schiff bases, potentiometric titration, Hyperquad, stability constants
1. Introduction
Schiff bases have been used extensively as ligands in the field of coordination chemistry.1−4 By attaching donor
atoms of Schiff bases to transition metal ions very stable complexes are formed in the tetrahedral structures.
Recently, Schiff base complexes have been attracting continuous attention for different applications.5−10 Tran-
sition metals play an important role in the construction of molecular materials that display magnetic properties
and they are used in materials, supramolecular, and biochemistry.11−15 It is well known that the metal com-
plexes of some drugs have higher activity than free ligand forms. In particular, most Cu(II) complexes have
been found to be antibacterial agents.16,17 The predication of acidity constants of organic reagents is important
in estimating their physical and biological activity. They play a fundamental role in many analytical proce-
dures such as acid–base titration, solvent extraction, and complex formation.18−20 The potentiometric titration
method is regarded as a powerful electro-analytical technique21 and is used for the determination of ionic equi-
librium of many ligands and the stability constants of complexes in solutions.22−29 In the present study, the
protonation–deprotonation equilibrium of a series of Schiff bases and the coordination properties of their binary
complexes with Cu(II) ion were investigated using the potentiometric titration method.
2. Results and discussion
2.1. Synthesis of the Schiff bases
The studied Schiff bases were prepared30 from the reactions of 2-aminophenol and 2-aminothiophenol with
related aldehyde derivatives (such as 2-Br-, 2-Cl-, 2-OCH−3 -, 2-OH-benzaldehyde, pyrrol-2-carbaldehyde,
∗Correspondence: hasatabey@gmail.com
109

ATABEY et al./Turk J Chem
furfural, thiophene-2-carbaldehyde, and pyridine-2-carbaldehyde) in ethanol at reflux conditions for 5 h (Figure
1; Table 1).
Table 1. The structure of ligands and their physical properties.
Entry
1
Ligands
N
M.p. (oC)
73-76
Yield (%)
89
Ref.
10
1a
1b
Br
HO
2
3
93-94
95-98
84
91
30,31
N
Cl
HO
N
This work
1c
OCH3
HS
N
4
5
1d
1e
103-105
110-114
90
94
32
33
OCH3 HO
N
OH
HS
N
6
7
182-184
119-121
96
68
32
34
1f
OH
HO
N
2a
NH
HO
N
8
2b
2c
76-78
73
This work
NH
O
HS
N
9
168-170
86
35
HO
N
46-48
78-80
95
92
96
35
10
11
2d
2e
S
S
HO
N
36
HS
N
12
13
3a
3b
172-174
37,38
N
HO
N
119-121
75
40
N
HS
110

ATABEY et al./Turk J Chem
CHO
Y
N
Y
X
Y = -Br, -Cl, -OCH3, -OH
N
CHO
CHO
NH
2
Z
EtOH
Z
Ref./5h.
X
X
Z = NH, O, S
X = -OH, -SH
N
N
N
X
Figure 1. General synthesis route for Schiff bases.
2.2. Dissociation constants
Dissociation constants were potentiometrically obtained from a series of several independent measurements.
Many NO-, NS-, ONS-, and ONO-type Schiff base ligands were investigated and 1e, 1f, 2a, and 3b represent
ONS-, ONO-, NO-, and NS-type Schiff bases, respectively. The distribution curves of ligands having different
coordination properties with respect to the chelating ability of 1e, 1f, 2a, and 3b are shown in Figure 2a–2d.
All the studied ligands have 2 or 3 donor atoms. For example, while 1e, 1f, 3a, and 3b have 3 protonable
donor atoms, other ligands have 2 protonable donor atoms. Consequently, if the fully protonated forms of the
Schiff bases are denoted as LHn , the general notation of its protonation equilibrium is as follows:
LHn + H2O ꢀ Hn−1 + H3O+
(1)
In each stage, one proton dissociates and dissociation constants are given as
[LHn−1] . [H3O]
Kn =
[LHn]
(2)
Dissociation constants of all ligands are calculated using the Hyperquad program under our experimental
conditions and are given in Table 2 in comparison with literature data.
Compound 3b was studied potentiometrically by Issa et al. using the Calvin–Bjerrum titration technique
as modified by Irving and Rossotti at 25 ◦ C and an ionic strength of 0.1 M (NaCl) in 70% (v/v) aqueous ethanol
and dissociation constant of –SH group was determined as 10.12.5 This value is the same as that found in our
study (10.12). Dissociation constants of 3a were studied by Geary et al., Gu¨rkan et al., and Sengupta et
al. in 50% (v/v) aqueous methanol and 50% (v/v) aqueous dioxin, respectively.6−8 Dissociation constants of
azomethine nitrogen of 3a have been determined as about pKa 6. The value is higher than our determined
pKa value (4.36). Additionally, while pKa values reported by Friedrich et al. are 10.46 and 12.46 in 75% (v/v)
aqueous dioxin,9 Gu¨rkan et al.’s values are 9.19 and 10.40 using potentiometric titration at ionic strength of
0.1 M (NaClO4) in 50% (v/v) aqueous methanol for 1f.7 On the other hand, 3 pKa values were obtained as
111

ATABEY et al./Turk J Chem
100
80
60
40
20
0
100
LH
LH
3
L
L
3
LH
LH
2
LH
LH
2
80
60
40
20
0
4
6
8
10
12
L
2
4
6
8
10
pH
(a)
pH
(b)
100
80
60
40
20
0
100
80
60
40
20
0
LH
LH
2
LH
L
LH
3
LH
2
2
4
6
pH
8
10
2
4
6
8
10
pH
(c)
(d)
Figure 2. The species distribution curves of the ligands (a) 1e, (b) 1f, (c) 2a, and (d) 3b (25.0 ± 0.1 ◦ C, I : 0.1 M
by NaCl, 0.05 mmol HCl).
8.07, 10.76, and 4.47, which were for azomethine nitrogen in 1f in the experimental conditions in this study. In
addition, Demirelli et al. have studied the determination of dissociation constants of 1a, 1d, and 1f in 20%,
40%, and 60% (v/v) aqueous dioxane, respectively, and shown the solvent effect on dissociation constants.10
As a result, different solvent and solvent ratios are shown to have changed the polarity of the solutions.
Therefore, increasing the solvent ratios in solutions causes increasing dipole–dipole interaction among molecules.
Thus, the measured pKa values in an organic solvent–water mix might be different from those in an aqueous
solution. It may also be thought that the high polarity of the solution media causes decreasing electron density
of the azomethine nitrogen. Similarly, high polarity of the solution media increases the electron density of the
phenolic groups. In this case, the pKa values of phenolic groups of the ligands are increased (see Table 2).
2.3. Stability constants
The stability constants of binary complexes between Schiff bases and Cu(II) ion were determined following the
refinement of data by the Hyperquad computer program. The cumulative stability constants (βmlh) are defined
by Eqs. (3) and (4).
112

ATABEY et al./Turk J Chem
Table 2. Dissociation constants of Schiff bases in the literature and this work (25.0 ± 0.1 ◦ C, I : 0.1 M by NaCl, 0.05
mmol HCl).
Ligands pKa1
pKa2
10.14
10.57
11.35
pKa3
logβ2
logβ3
References
Ref.10
Ref.10
1a
4.26
4.21
3.90
-
-
-
-
-
-
-
-
-
Ref.10
3.38 ꢀ 0.03 9.46 ꢀ 0.08
3.29 ꢀ 0.03 9.38 ꢀ 0.09
3.05 ꢀ 0.03 9.75 ꢀ 0.08
-
-
-
12.84 ꢀ 0.03
This work
1b
1c
1d
12.68 ꢀ 0.03
12.80 ꢀ 0.02
-
-
This work
This work
4.45
4.22
3.95
10.22
10.56
11.35
-
-
-
-
-
-
-
-
-
Ref.10
Ref.10
Ref.10
-
-
3.18 ꢀ 0.01 9.21 ꢀ 0.06
12.39 ꢀ 0.03
This work
1e
1f
-
9.32
10.34
-
-
Ref.8
3.94 ꢀ 0.02 9.16 ꢀ 0.03 11.33 ꢀ 0.08 20.45 ꢀ 0.06 24.43 ꢀ 0.07 This work
-
9.19
10.46
8.62
8.73
9.61
10.40
12.46
10.11
10.25
11.30
-
-
-
-
-
-
-
-
-
Ref.7
Ref.9
Ref.10
Ref.10
Ref.10
-
4.28
4.11
4.01
4.47 ꢀ 0.04 8.07 ꢀ 0.01 10.76 ꢀ 0.08 18.83 ꢀ 0.07 23.29 ꢀ 0.07 This work
2a
2b
2c
2d
2e
3a
3.83 ꢀ 0.03
3.75 ꢀ 0.03
-
-
9.15 ꢀ 0.08
12.99 ꢀ 0.03
13.44 ꢀ 0.03
14.58 ꢀ 0.02
13.34 ꢀ 0.03
16.09 ꢀ 0.06
-
-
-
-
-
This work
This work
This work
This work
This work
9.69 ꢀ 0.07
5.43 ꢀ 0.03 9.15 ꢀ 0.02
4.07 ꢀ 0.01 9.28 ꢀ 0.02
6.35 ꢀ 0.02 9.74 ꢀ 0.03
-
-
-
6.41
6.58
6.37
-
-
-
9.25
9.13
9.77
-
-
-
Ref.7
Ref.8
Ref.6
4.36 ꢀ 0.04 6.09 ꢀ 0.09 9.11 ꢀ 0.01
15.19 ꢀ 0.04 19.55 ꢀ 0.03 This work
Ref.5
-
-
10.12
-
3b
3.38 ꢀ 0.04 6.54 ꢀ 0.09 10.12 ꢀ 0.04 16.66 ꢀ 0.04 20.03 ꢀ 0.04 This work
113

ATABEY et al./Turk J Chem
mM + lL + hH ꢀ MmLlHh
[MmLlHh]
(3)
(4)
βmlh
=
,
[M]m [L]l [H]h
where M is Cu(II) ion, L is ligand, and H is proton, and m, l, and h are the respective stoichiometric coefficients.
The potentiometric data for the Cu(II) – L2 systems indicate that there is a significant tendency toward the
formation of ML2 species. Cu(II) ion complexes were formed by releasing 2 of the hydrogen ions from the fully
protonated form of the ligands.11 Compounds 1e and 1f serve as tridentate ligands by the coordinating of
imino, phenolic –OH and –SH groups with Cu(II) ion. The others (1a, 1b, 1c, 1d, 2a, 2b, 2c, 2d, 2e, 3a, and
3b) serve as bidentate ligands. Thus, the stability constants of Cu(II) complexes of 1e and 1f are higher than
those of the others.
Coordination numbers of a central atom can change to 6 or 4 depending on the ligand structures in the
complex formation. For example, the coordination number of Cu(II) ion was observed as 4 against all studied
ligands except 1e and 1f in this study. Therefore, tetrahedral complexes were obtained. On the other hand, the
coordination number of Cu(II) ion was 6 against 1e and 1f, because they are tridentate ligands. The molecular
structures of 4- and 6-coordinated complexes of copper with 1a and 1f are given in Figure 3a and 3b.
Figure 3. Molecular structures of 4- and 6-coordinated complexes of Cu(II) with 1a and 1f (a) Cu(II)–1a2 (b)
Cu(II)–1f 2
.
Complexes having 6 coordination numbers form an octahedral structure and they can be formulated as
MX2 , where MX2 is a structure of bis-complex, M is Cu(II) ion, and X is the ligand 1e (or 1f). Therefore,
the octahedral complexes between 1e and 1f and Cu(II) ion are more stable than the tetrahedral complexes.
This situation was supported by the experimental (Table 3) and the semiempirical molecule orbital (SE-MO)
PM3 method. 3D structures of complex species and their formation heats (Hf ) are calculated by PM3 method.
Accordingly, formation heats (Hf ) of Cu–1e2 /1f 2 /2a2 /3b2 complexes were determined as 608.78 kcal/mol,
613.42 kcal/mol, 728.96 kcal/mol, and 776.13 kcal/mol, respectively, and the findings are given in Figure 4a–4d.
114

ATABEY et al./Turk J Chem
Figure 4. Comparison of 3D structure of complex species and their formation heats (Hf ) (a) Cu(II)–1e2 complex (ONS
type) (Hf : 608.78 kcal/mol) (b) Cu(II)–1f 2 complex (ONO type) (Hf : 613.42 kcal/mol) (c) Cu(II)–2a2 complex (NO
type) (Hf : 728.96 kcal/mol) (d) Cu(II)–3b2 complex (NS type) (Hf : 776.13 kcal/mol).
As a result, different electron densities on the donor atoms are an important factor affecting the stability
of the Cu(II) complexes with the ligands. The various complexes between Cu(II) ion and the Schiff bases were
formulated as CuL2 , CuHL2 , CuH2 L2 , and CuH−1 L2 (Cu (OH) L2), depending on pH. The overall stability
constants of detectable Cu(II)–L2 species are given in Table 3.
115

ATABEY et al./Turk J Chem
Table 3. Overall stability constants in Cu(II)–L2 binary system (25.0 ± 0.1 ◦ C, I : 0.1 M by NaCl, 0.05 mmol HCl).
Ligands
m
1
1
h
0
1
l
2
2
logβ
14.07 ꢀ 0.05
Ligands
m
1
1
1
1
1
1
1
1
h
0
1
–1
0
1
–1
0
1
l
logβ
2
2
2
2
2
2
2
2
15.88 ꢀ 0.03
22.90 ꢀ 0.05
6.68 ꢀ 0.07
16.18 ꢀ 0.03
22.96 ꢀ 0.03
5.99 ꢀ 0.07
11.59 ꢀ 0.08
18.94 ꢀ 0.15
1a
20.38 ꢀ 0.08
2a
1
1
1
1
1
0
1
–1
0
–1
2
2
2
2
2
13.43 ꢀ 0.05
19.97 ꢀ 0.08
5.99 ꢀ 0.09
13.65 ꢀ 0.15
6.56 ꢀ 0.10
1b
1c
2b
2c
1d
1e
1f
1
1
0
–1
2
2
11.24 ꢀ 0.04
1.67 ꢀ 0.09
2d
2e
3a
1
1
0
–1
2
2
13.71 ꢀ 0.05
5.45 ꢀ 0.11
1
1
1
0
1
2
2
2
2
18.85 ꢀ 0.09
28.64 ꢀ 0.05
35.55 ꢀ 0.07
1
1
0
–1
2
2
13.29 ꢀ 0.04
3.84 ꢀ 0.07
1
1
1
0
1
–1
2
2
2
13.60 ꢀ 0.06
18.91 ꢀ 0.04
4.76 ꢀ 0.01
1
1
1
0
1
2
2
2
2
17.51 ꢀ 0.05
26.33 ꢀ 0.04
34.11 ꢀ 0.02
3b
1
1
1
0
1
–1
2
2
2
14.65 ꢀ 0.06
21.52 ꢀ 0.08
6.93 ꢀ 0.08
Electron pairs on donor atoms play a critical role for complex formation. Mobility of the electron pairs
facilitates participation in coordination. However, electron-withdrawing groups on the ligands cause decreasing
stability in the complexes because of the limitation of electron mobility. This situation is clearly seen in Table
3. Differences in electronegativity of Br and Cl atoms cause different stability constants in the Cu(II)–1a2 and
Cu(II)–1b2 complexes. The same case can be said for the –OH and –SH groups.
The species distribution curves of the complexes between Cu(II) ion and 1e, 1f, 2a, and 3a ligands are
given in Figure 5a–5d.
In Figure 5, the species distribution curves of 1e differ from those of 1f because of the different electron
density of the –OH and –SH groups. In the Cu(II)–1e2 system, 3 main complexes (CuL2, CuHL2 , and CuH2 L2)
were obtained at between pH 5 and 11. The CuHL2 species start occurring at pH 5 and reach the maximum
at pH 8–9 by 90%; and the CuL2 species start to form at pH 8 and reach the maximum at pH 11 by 99%. In
the Cu(II)–1f 2 system, similarly, CuL2 and CuH2 L2 complex species were observed in the acidic and basic
region at 99%, the same as in the Cu(II)–1e2 system. However, CuHL2 species reach the maximum at pH 8–9
and approx. 60%.
In Cu(II)–2a and Cu(II)–3a systems, the main complexes (CuL2 and CuHL2) were obtained in neutral
and acidic regions. The CuL2 and CuHL2 species exist above pH 7 at 90% and 98%, respectively. For both
complexes (Cu(II)–2a and Cu(II)–3a), hydrolysis species (CuH−1 L2) were also observed at pH 11 and at 99%.
The logβCuL2 values are shown in Figure 6.
116

ATABEY et al./Turk J Chem
100
100
80
60
40
20
0
CuH2L2
Cu2L
CuL2
CuH2L2
CuHL2
80
60
40
20
0
CuHL2
5
6
7
8
9
10
11
6
7
8
9
10
11
pH
pH
(a) Cu (II) - 1e2 system
(b) Cu (II) - 1f2 system
100
80
60
40
20
0
100
80
60
40
20
0
CuHL2
CuH-1L2
CuHL2
CuH-1L2
CuL2
CuL2
5
6
7
8
9
10
11
4
6
8
10
pH
pH
(c) Cu (II) - 2a2 system
(d) Cu (II) – 3b2 system
Figure 5. The species distribution curves of complexes between Cu(II) ion and 1a, 1e, 2a, and 3a ligands (a) Cu
(II)–1e2 system (b) Cu(II)–1f 2 system (c) Cu(II)–2a2 system (d) Cu(II)–3b2 system (25.0 ± 0.1 ◦ C, I : 0.1 M by
NaCl, 0.05 mmol HCl).
19
18
17
16
15
14
13
12
11
1a 1b 1c 1d 1e 1f 2a 2b 2c 2d 2e 3a 3b
Ligands
Figure 6. Changing of the logβCuL2 values for the ligands.
117

ATABEY et al./Turk J Chem
3. Experimental procedure and methods
3.1. Preparation of the Schiff bases
The studied Schiff bases were synthesized by a procedure reported by Perumal et al.30 To a stirred solution of
2-methoxybenzaldehyde (1.36 g, 10 mmol) in ethanol (10 mL) was added a solution of 2-aminothiophenol (1.87
g, 15 mmol) in ethanol (10 mL). The mixture was refluxed for 5 h. After cooling the reaction mixture, the
precipitated substance was filtered and recrystallized in ethanol. The other Schiff bases were prepared by the
above-mentioned procedure. The physical data of unknown compounds:
(E)-2-(2-methoxybenzylideneamino)benzenethiol (1c): (yield 91%; mp 95–98 ◦ C); 1 H NMR (400 MHz,
CDCl3 ) δg= 8.64 (d, J = 7.6 Hz, 1H), 8.19 (d, J = 8.0 Hz, 1H), 7.98 (d, J = 8.0, Hz, 1H), 7.55 (t, J =
7.6 Hz, 1H), 7.48 (t, J = 7.2 Hz, 1H), 7.42 (s, 1H, HC = N), 7.28 (t, J = 8.0 Hz, 1H), 7.19 (t, J = 7.6 Hz.
1H9, 7.05 (d, J = 8.4 Hz, 1H), 4.04 (s, 3H, -OCH3), 3.88 (brs, 1H, -SH). 13 C NMR (100 MHz, CDCl3 ):gδ =
163.21, 159.16, 157.29, 136.49, 131.87, 129.57, 125.99, 124.68, 122.85, 121.31, 121.20, 116.42, 111.74, 55.89. IR
(Liquid): 3544, 3475, 3413, 3226, 1612, 1563, 1415, 1138, 1041, 884, 863, 747, 605, 482. Elemental Anal. Cald:
C, 69.11; H, 5.39; N, 5.76; S, 13.18. Found: C, 68.91; H, 5.27; N, 5.72; S, 13.28.
(E)-2-((1H-pyrrol-2-yl)methyleneamino)benzenethiol (2b): (yield, 73%; mp 76–78 ◦ C); 1 H NMR (400
MHz, CDCl3 ) δg= 11.74 (s, -NH), 8.36 (s, 1H, HC = N), 7.44 (d, J = 8.0 Hz, 1H), 7.24–7.21 (d, J = 8.0,
Hz, 1H), 7.16–7.11 (m, 3H), 6.80 (bs, 1H), 6.24 (m, 1H), 4.56 (s, 1H, SH). 13 C NMR (100 MHz, CDCl3 ): δ =
151.22, 149.74, 130.80, 130.61, 127.62, 126.30, 126.02, 125.28, 118.22, 117.89, 110.54. IR (Liquid): 3554, 3482,
3415, 3235, 1616, 1567, 1413, 1132, 1037, 881, 867, 744, 601, 480. Elemental Anal. Calcud: C, 65.32; H, 4.98;
N, 13.85; S, 15.85. Found: C, 65.28; H, 5.18; N, 13.78; S, 15.98.
3.2. Apparatus and materials
Firstly, Schiff bases were dissolved in sufficient ethanol and diluted at a ratio of 1/10. Next, 1 × 10−3
M
stock solution was prepared for each ligand. Ethanol, NaCl, and CuCl2 were purchased from Merck, potassium
hydrogen phthalate (KHP) and borax (Na2 B4 O7) from Fluka, and 0.1 M NaOH and 0.1 M HCl as standard
from Aldrich. All reagents were of analytical quality and were used without further purification. A solution of
metal ion (1 × 10−3 M) was prepared from CuCl2 as received and standardized with ethylenediaminetetraacetic
acid (EDTA).40 Next, 1.0 M NaCl stock solution was prepared from the original bottle. For all solutions, CO2 -
free double-distilled deionized water was obtained with an aquaMAX-Ultra water purification system (Young
Lin Inst.). Its resistivity was 18.2 MΩ cm−1
.
3.3. Potentiometric measurements
All potentiometric pH measurements were carried out on solutions in a 100-mL double-walled glass vessel using
the Molspin pH meter with Orion 8102BNUWP ROSS ultra combination pH electrode and the temperature
was controlled at 25.0 ꢀ 0.1 ◦ C by circulating water through the double-walled glass vessel, from a constant-
temperature bath (DIGITERM 100, SELECTA). The electrode was calibrated according to the instructions in
the Molspin Manual.41 An automatic burette was connected to the Molspin pH-mV-meter. The pH electrode
was calibrated with a buffer solution of pH 4.005 (KHP) and pH 9.180 (borax)42 at 25.0 (ꢀ0.1) ◦ C. During
the titration, nitrogen (99.9%) was purged through the cell. The Hyperquad43 computer program was used for
the calculation of both dissociation and stability constants.
118

ATABEY et al./Turk J Chem
The cell was equipped with a magnetic stirrer. Atmospheric CO2 was excluded from the titration cell
with a purging steam of purified N2 . The system was maintained at an ionic strength of 0.1 M by NaCl as a
supporting electrolyte. A solution containing about 0.01 mmol of the ligands, and the required amount of 1.0
M NaCl and 0.1 M HCl were put into the titration cell. Finally, doubly distilled deionized water was added to
the cell to a total volume of 50 mL and titration was started. The pH data points were collected after each
addition of 0.03 mL of the standardized NaOH solution. The second solution contained the same amounts of
components plus 0.005 mmol of Cu(II) solution and doubly distilled deionized water was added to the same
total volume. The potentiometric studies were carried out at the metal:L molar ratios of 1:2 and each titration
was repeated 3 times.
4. Conclusion
In this work, the effect of substituents on the Cu(II)–L2 complexes was discussed. Different electron densities
on the donor atoms are an important factor affecting the stability of the Cu(II) complexes with the ligands.
The various complexes between Cu(II) ion and Schiff bases were formulated as CuL2 , CuHL2 , CuH2 L2 , and
CuH−1 L2 (Cu (OH) L2) depending on pH. Stability constants of binary complexes between Cu(II) and Schiff
bases were determined in 0.1 M ionic strength (NaCl) and at 25.0 ꢀ 0.1 ◦ C, using the combined glass electrode,
potentiometrically. The logβCuL2 values for the ligands are shown in Figure 4. The dissociation constants and
overall stability constants were calculated using Hyperquad and the results are given in Tables 1 and 2. The
coordination number of Cu(II) ion was 4 against all studied ligands except 1e and 1f in this study. Therefore,
tetrahedral complexes were obtained. However, the coordination number of Cu(II) ion was 6 against 1e and 1f
because they are tridentate ligands. Cu(II)–1e and –1f complexes are in octahedral structure. As a result, the
complexes of ONS- and ONO-type tridentate ligands are more stable than those of NO- and NS-type bidentate
ligands. The logβCuL2 values are changed as 1e > 1f > 2b > 2a > 3b > 1a > 2d > 1c > 3a > 2c >
1b > 2e > 1d.
Acknowledgment
The authors gratefully acknowledge the financial support of this work from the Scientific Research Council of
Gaziosmanpa¸sa University.
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