Synthesis and characterization of polydentate schiff-base ligands and their complexes

Article abstract of DOI:10.1246/bcsj.79.867

In this study, we synthesized the amine compounds 2-{(E)-[(2-aminoethyl) imino]methyl}phenol (H₃A) and 2-{(E)-[(3-aminopropyl)imino]methyl}-6- methoxyphenol (H₃B) as the starting materials. From reactions of the starting materials with 2-hydroxy-5-methylisophthalaldehyde, phthalaldehyde, and 2-hydroxy-5-t-butylisophthalaldehyde, we prepared the new ligands H ₃L¹-H₃L³ and H₂L. The Cu(II) and Cd(II) complexes of the ligands have been obtained. Microanalytical data, magnetic moment, ¹H(¹³C) NMR, mass spectra, FT-IR, and conductivity measurements have been used to explain the structures of the ligands and their complexes. The protonation constants of the ligands H ₃L¹-H₃L³ have been studied in a 1:1 molar ratio (M:L). Protonation and stability constants of the Schiff bases and their Cu(II) and Cd(II) complexes have been determined by the potentiometric titration method in 50% dioxane-water media at 25.00 ± 0.02°C under a nitrogen atmosphere and ionic strength of 0.1M NaClO₄. The ligands H₃L1-H₃L³ have seven protonation constants. In other words, the ligand H₂L has six protonation constants. The variation of protonation constants of these compounds were interpreted on the basis of structural effects exposed by the substituents.

Full text of DOI:10.1246/bcsj.79.867

ꢀ 2006 The Chemical Society of Japan
Bull. Chem. Soc. Jpn. Vol. 79, No. 6, 867–875 (2006)
867
Synthesis and Characterization of Polydentate Schiff-Base
Ligands and Their Complexes
ꢀ2
2
Havva Demirelli,1 Mehmet Tu¨mer, and Aysꢀegu¨l Golcu
¨
¨
1Department of Chemistry, Faculty of Education, Gazi University, 06100, Ankara, Turkey
2
˙
¨
Department of Chemistry, Faculty of Science and Arts, K.Marasꢀ Sutc¸u Imam University, 46100, K.Maras, Turkey
¨
Received May 31, 2005; E-mail: mtumer@ksu.edu.tr
In this study, we synthesized the amine compounds 2-{(E)-[(2-aminoethyl)imino]methyl}phenol (H3A) and 2-
{(E)-[(3-aminopropyl)imino]methyl}-6-methoxyphenol (H3B) as the starting materials. From reactions of the starting
materials with 2-hydroxy-5-methylisophthalaldehyde, phthalaldehyde, and 2-hydroxy-5-t-butylisophthalaldehyde, we
prepared the new ligands H3L1–H3L3 and H2L. The Cu(II) and Cd(II) complexes of the ligands have been obtained.
Microanalytical data, magnetic moment, 1H(13C) NMR, mass spectra, FT-IR, and conductivity measurements have been
used to explain the structures of the ligands and their complexes. The protonation constants of the ligands H3L1–H3L3
have been studied in a 1:1 molar ratio (M:L). Protonation and stability constants of the Schiff bases and their Cu(II) and
Cd(II) complexes have been determined by the potentiometric titration method in 50% dioxane–water media at
25:00 ꢁ 0:02 ꢂC under a nitrogen atmosphere and ionic strength of 0.1 M NaClO4. The ligands H3L1–H3L3 have seven
protonation constants. In other words, the ligand H2L has six protonation constants. The variation of protonation con-
stants of these compounds were interpreted on the basis of structural effects exposed by the substituents.
Since proton transfer is known to be crucial for physico-
chemical properties and practical application of Schiff bases,
this process has been widely studied in the literature.1,2 It
has been shown that the position of the tautomeric equilibrium
depends not only on the electronic structure of the molecule,
but also on its conformation. The interactions between the
non-bonded atoms in molecules, leading to the strengthening
of the hydrogen bond, as well as the intramolecular interac-
tions in the solid state, stabilizing the planar structure of the
molecule, shift the equilibrium toward the NH form.3 The in-
tramolecular proton-transfer reaction proceeds comparatively
easily in the ortho-hydroxy Schiff bases. Intramolecular ꢀ-
electron coupling leads to the strengthening of the hydrogen
bonds in these systems. These compounds are potential mate-
rials for molecular memory and optical switch devices4 or fluo-
rescent probes of biological systems.5
Among the different physicochemical properties of organic
compounds, protonation and stability constants determined in
mixed solvents provide an important basis for speculation
about whether substituent effects influence their acidic and ba-
sic properties. It is also accepted that knowledge of the stabil-
ity constants of such Schiff bases and their metal complexes
may eventually help to throw some light on the inactivation
of essential trace metals in biological systems. Therefore, it
was interesting to determine the protonation and stability con-
stants of some Schiff bases–metal complexes in dioxane–water
media of different compositions potentiometrically.6 The study
of the complex formation of Schiff bases can not be carried out
in aqueous solution because the ligands H3L1–H3L3 and their
Cu(II) and Cd(II) complexes are not soluble in water media. A
dioxane–water mixture is a typical binary solution, as the di-
electric constant can be varied from 2.2 to 78 at 25 ꢂC. Another
interesting characteristic of this system is that non-polar di-
oxane is completely miscible with water at every ratio. The di-
oxane–water mixture (50:50%) has been chosen as the solvent
mixture for our study. In such a medium, the Schiff bases and
metal complexes are soluble giving stable solutions.
Numerous electrochemical studies have been made for a
fairly large number of acyclic and macrocyclic Cu(II) com-
plexes derived from Schiff bases. These investigations reveal-
ed that the redox properties of the Cu(II) complexes are mark-
edly influenced by structural and electronic factors.7,8
The present study deals with the synthesis, characterization,
and electrochemical and potentiometric determination of the
Schiff bases and their metal complexes. In addition, we study
the stability constants of their Cu(II) and Cd(II) complexes by
the potentiometric titration method in 50:50% dioxane–water
media at 25:00 ꢁ 0:02 ꢂC and an ionic strength of 0.10 M
NaClO4. The obtained data from titrations were evaluated by
the computer program BEST.9
Experimental
Materials. Stock solutions of the Schiff bases were prepared
in purified dioxane.10 Doubly distilled conductivity (Millipore sys-
tem) water was used as the aqueous medium as well as for the
preparation of dioxane–water mixtures. All the other chemicals
used were of A.R. grade and were used without further purifica-
tion. Stock solutions of 0.03 M Cu(ClO4)2 and Cd(ClO4)2 were
standardized using the appropriate indicator by EDTA titrations.11
Sodium hydroxide solutions were prepared as 50% (v/v) aqueous
dioxane solutions. The concentration of NaOH and absence of car-
bonate were frequently checked by means of Gran plots12 using
potassium hydrogen phthalate (Merck) as the acid. The acid solu-
tions (0.1 M) prepared from Merck pa perchloric acid were titrated
Published on the web June 6, 2006; doi:10.1246/bcsj.79.867

868 Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
Schiff-Base Ligands and Their Complexes
Treatment of Potentiometric Data.
against a standardized 0.1 M sodium hydroxide solution.13 The
ionic strength of each solution was adjusted to 0.10 M by the
addition of NaClO4 as the supporting electrolyte. All organic and
inorganic materials with solvents were purchased from Fluka or
Merck. 2-Hydroxy-5-methylisophthalaldehyde and 2-hydroxy-5-
t-butylisophthalaldehyde were obtained according to the literature
method.14 Physical measurements and 1H(13C) NMR data of the
ligands and their Cd(II) complexes have been given in the Sup-
porting Information.
The ligand protona-
tion and metal ion stability constants were calculated with the
FORTRAN program BEST9 and were obtained through the alge-
braic solution of the mass and charge balance equations evaluated
at each equilibrium point of the formation curves. The input for
the program BEST consists of the components, the concentrations
of each component and initial estimates of the equilibrium con-
stant for each species. The species considered present in the exper-
imental solutions were those that one would expect to form ac-
cording to the principles of coordination chemistry. The program
refines stability constants by the iterative nonlinear least-squares
fit of potentiometric equilibrium curves through a set of simulta-
neous mass balance equations for all the components expressed
in terms of known and unknown equilibrium constants. By suita-
ble use of the program, it is possible to obtain a high degree of
discrimination in the selection of chemical species. Multinuclear
metal complexes and other species, such as MLn, MnLn, HnMLn,
and so on, included at various stages of refinement were rejected
on the basis of increased ꢁ values. The inclusion of these species
did not improve the estimated standard deviation and the value of
their equilibrium constants progressively decreased without con-
vergence. The best set of complexes was: ML, HML, H2ML,
H3ML, and H4ML for Cu(II); ML, HML, H2ML, H3ML, H4ML,
and H5ML for Cd(II). All the models converged at ꢁ < 0:02p[H]
units of the observed p[H] values, which is considered to be an ac-
ceptable fit. The equilibrium constants reported in this paper were
obtained as averaged values of three titrations. Species distribu-
tion curves were calculated with the FORTRAN program SPE.9
Synthesis of the Starting Materials H3A and H3B. The pre-
cursors were prepared according to the modified method described
earlier18,19 by mono condensation of the appropriate diamines with
aldehydes. To the vigorously stirred and cool dilute solution (T ¼
5{10 ꢂC) of the diamine (ethylenediamine and/or 1,3-diaminopro-
pane, 20 mmol) in absolute ethanol (100 mL), was added dropwise
a cooled solution of salicylaldehyde and/or vanilline (15 mmol) in
absolute ethanol (80 mL). After the addition was complete, the
mixture was stirred for 15–30 min and then refluxed for 15–60
min. The resulting solution was evaporated under vacuum to re-
move the solvent, and excess diamine was extracted by benzene.
The obtained compounds were used for the next step without fur-
ther purification.
Potentiometric Apparatus and Procedure. Potentiometric
titrations were carried out in jacketed glass reaction vessels as de-
scribed in Ref. 9. The cell EMF was measured using an Orion EA
940 pH meter (resolution 0.1 mV, accuracy 0.2 mL) equipped with
a Mettler Toledo Inlab 412 combined glass electrode and an Orion
960 automatic titrator containing carbonate-free sodium hydroxide
at a known (ꢃ0:1 M) concentration at 25:00 ꢁ 0:02 ꢂC with an
ionic strength of 0.10 M (NaClO4). The electrode was modified by
replacing its aqueous KCl solution with 0.01 M NaCl + 0.09 M
NaClO4 saturated with AgCl. The temperature was maintained
constant inside the cell at 25:00 ꢁ 0:02 ꢂC by water circulation
from a Haake thermostatted bath (precision ꢁ 0:02). All potentio-
metric measurements were carried out in water–dioxane mixtures
containing 50% dioxane (v/v) because of the low solubility of the
Schiff bases and possible hydrolysis in aqueous solutions. Poten-
tiometric titrations were carried out at constant temperature and in
an inert atmosphere of nitrogen with CO2-free standardized 0.1 M
NaOH in a 50.0 mL solution containing 0.1 M NaClO4 (i) 2:5 ꢄ
10ꢅ3 M HClO4 (for cell calibration), (ii) 6:0 ꢄ 10ꢅ3 M HClO4 +
1:5 ꢄ 10ꢅ3 M Schiff base (for protonation constant of Schiff base),
and (iii) 6:0 ꢄ 10ꢅ3 M HClO4 + 1:5 ꢄ 10ꢅ3 M Schiff base +
1:5 ꢄ 10ꢅ3 M Cu(ClO4)2/Cd(ClO4)2 (for stability constant of
Schiff base–metal complex).
Cell Calibration. The potentiometric cell was calibrated for
use of the combined pH electrode as a hydrogen ion concentration
probe rather than as an activity probe. The ionic strength of the
test solutions used in this study was kept constant; therefore, the
EMF of the cell can be written in the form
0
Ecell ¼ Eꢂ cell þ Ej þ k log½Hþꢆ;
ð1Þ
0
where Eꢂ
represents a quantity independent of [Hþ], but de-
cell
pendent on the activity of Clꢅ in the filling solution of the elec-
trode and the activity coefficient of Hþ in the test solution.15
The activity coefficient of Hþ can be considered to be constant
throughout the titration because the ionic strength of the solution
is almost constant. Ej is the liquid junction potential and the con-
stant k, denoted as electrode calibration 0 slope, represents the
H3A:
calcd): C, 65.80 (65.83); H, 7.40 (7.37); N, 17.10 (17.06). UV–vis
Yield (40%), color: yellow, mp 105 ꢂC. Found (%
(ꢂmax/nm, EtOH as solvent): 407, 342, 315, 290, 268, 255. IR
(KBr, cmꢅ1): 3377 [ꢃ(OH)], 2872 [ꢃ(CH )], 2640 [ꢃ(OH N)],
ꢇꢇꢇ
2
1633 [ꢃ(CH=N)], 1377 [ꢃ(C–OH)]. 1H NMR: (CDCl3 as solvent,
ꢄ in ppm): 10.5 (s, OH), 8.68 (CH=N), 7.99–6.66 (Ar-H, m), 4.52
(t, NH2), 3.23 and 3.22 (t, CH2, t, CH2).
Nernst factor. The calibration constants Eꢂ
and k, were deter-
cell
mined by titration with a 2:5 ꢄ 10ꢅ3 M solution of HClO4 and
0.1 M NaOH solution for the reaction media. All the solutions
for the calibration titrations were made up to an ionic strength
of 0.1 M NaClO4. In all titrations, the experimental points in the
region 2:2 < pH < 3:0 were used for calibrations, where pH rep-
resents ꢅ log½Hþꢆ. Within this range of pH, the Ej is effectively
H3B: Yield (37%), color: dark yellow, mp 58 ꢂC. Found (%
calcd): C, 63.47 (63.44); H, 7.70 (7.74); N, 13.48 (13.45). UV–vis
(ꢂmax/nm, EtOH as solvent): 419, 342, 326, 320, 302, 297, 245.
IR (KBr, cmꢅ1): 3355 [ꢃ(OH)], 2931 [ꢃ(CH3)], 2847 [ꢃ(CH2)],
2662 [ꢃ(O–H N)], 1632 [ꢃ(CH=N)], 1327 [ꢃ(C–OH)]. 1H NMR:
ꢇꢇꢇ
constant.15 Equation 1 in the form Ecell ¼ þEꢂcell þ k log½Hþꢆ,
(CDCl3 as solvent, ꢄ in ppm): 10.6 (s, OH), 8.40 (CH=N), 7.60–
6.18 (Ar-H, m), 4.03 (t, NH2), 3.96 (OCH3), 3.33 and 3.20 (t,
CH2). 13C NMR (ꢄ in ppm): 162.15 (CH=N), 155.20–110.15
(Ar-C), 60.47 (OCH3), 20.85–20.60 (–CH2–).
0
where Eꢂ ¼ Eꢂ cell þ Ej, was found to reproduce the calibration
cell
data to a precision typically of the order of ꢁ1:0 mV. The stan-
dardization of the combined pH electrode was checked in the al-
kali range too by the addition of an excess of NaOH. By assuming
Synthesis of the Schiff Bases H3L1–H3L3 and H2L. The
unsymmetrical Schiff bases were obtained by condensation of the
half units and the appropriate aldehydes. To the stirred solution of
the precursor (half units, H3A and H3B) in absolute ethanol was
added a solution of 2-hydroxy-5-methylisophthalaldehyde, 2-hy-
the Eꢂ value determined in the acidic range to be reliable and
cell
[OHꢅ] = concentration of base added in excess, we calculated
reproducible values of pKw, for the solvent mixtures examined16,17
and obtained the pKw value of 15.37 in this media.

H. Demirelli et al.
Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
869
Table 1. Analytical and Physical Data for the Schiff-Base Ligands and Their Cu(II) and Cd(II) Complexes
Color
Yield
/%
mp
/ꢂC
Found (calcd)/%
C
Compound
H
N
Cl
Cu
—
H3L1
light brown
brown
brown
yellow
black
brown
light brown
black
brown
red-orange
black
orange
85
72
65
80
71
68
87
69
66
90
73
70
115
>250
>250
97
>250
>250
74
71.09 (71.05)
52.68 (52.63)
45.47 (45.42)
68.42 (68.38)
55.73 (55.68)
48.61 (48.58)
69.65 (69.62)
54.77 (54.72)
48.42 (48.38)
73.29 (73.26)
55.71 (55.69)
51.25 (51.22)
6.18 (6.20)
4.10 (4.06)
3.56 (3.51)
6.67 (6.62)
4.97 (4.94)
4.35 (4.31)
7.23 (7.17)
5.27 (5.23)
4.66 (4.62)
6.15 (6.10)
4.69 (4.64)
4.22 (4.27)
12.31 (12.27)
9.12 (9.09)
7.89 (7.85)
10.32 (10.29)
8.44 (8.38)
7.36 (7.31)
9.60 (9.56)
7.54 (7.51)
6.67 (6.64)
13.18 (13.15)
10.03 (9.99)
9.15 (9.19)
—
[Cu2(L1)]Cl
[Cd2(L1)]Cl
H3L2
5.82 (5.77)
5.04 (4.98)
—
5.37 (5.31)
4.70 (4.64)
—
4.80 (4.76)
4.27 (4.21)
—
12.71 (12.66)
11.70 (11.65)
20.70 (20.64)
31.58 (31.52)
—
19.10 (19.02)
29.43 (29.36)
—
17.11 (17.04)
26.73 (26.66)
—
11.41 (11.34)
18.52 (18.45)
[Cu2(L2)]Cl
[Cd2(L2)]Cl
H3L3
[Cu2(L3)]Cl
[Cd2(L3)]Cl
H2L
215
>250
72-74
154
[Cu(H2L)]Cl2
[Cd(H2L)]Cl2
>250
droxy-5-t-butylisophthalaldehyde, or phthalaldehyde in absolute
ethanol. The mixture was concentrated in vacuum by evaporation
of the solvent until a colored solid precipitated. The product was
collected by filtration, washed with cold ethanol, and recrystal-
lized from the appropriate solvents (ethanol for H2L, methanol/
hexane mixture (1:1) for H3L1–H3L3) to give colored crystals.
Electronic Absorption Spectral Data in Different Solvents
of the Schiff Bases H3L1–H3L3 and H2L. H3L1: (ꢂmax/nm),
hexane: 362, 322, 298, 249; toluene: 476, 364, 350, 324; dioxane–
water: 426, 317, 275; MeOH: 407, 365, 315, 286, 280, 253. H3L2:
(ꢂmax/nm), hexane: 363, 324, 318, 256; toluene: 438, 367, 362,
324; dioxane–water: 429, 403, 362, 322, 278; MeOH: 424, 362,
318, 280, 254. H3L3: (ꢂmax/nm), hexane: 360, 349, 336, 273,
222; toluene: 460, 366, 341, 325, 322, 286; dioxane–water: 423,
288; MeOH: 428, 395, 349, 334, 297, 288. H2L: (ꢂmax/nm),
hexane: 389, 350, 334, 274, 259; toluene: 462, 450, 395, 350,
334, 286; dioxane–water: 346, 336, 301, 276; MeOH: 410, 349,
332, 295, 277.
O
N
N
OH
CH
OH
3
H N
H N
2
2
H3A
H3B
H
H
OH
N
N
N
N
OH
H L
Preparation of the Complexes. The complexes were pre-
pared by similar methods. A solution of the metal salt (M: CuCl2
2
ꢇ
Fig. 1. General formulae of the starting materials and
Schiff-base ligand H2L.
2H2O and CdCl2 H2O) (2 mmol) in absolute EtOH (25 mL) was
ꢇ
added to a solution of the ligands (2 mmol for the ligands H3L1–
H3L3 and 1 mmol for the ligand H2L) in absolute EtOH (20 mL)
and the mixture was boiled under reflux for 6–7 h. At the end of
the reaction, determined by TLC, the precipitate was filtered off,
washed with EtOH and Et2O, and dried in vacuo.
H2L have been synthesized from the reaction of the carbonyl
and amine compounds; they are more stable than the starting
compounds at room temperature. As the ligands have a lot
of polar groups, such as –OH and –CH=N, they are soluble
in polar organic solvents such as, EtOH, MeOH, CHCl3,
(C2H5)2O, etc. It is well known that the oxygen atoms of metal
complexes of salen-type ligands have anionic character, which
sometimes enhances coordination strength to other metal
ions.20 We investigated the synthesis of transition-metal com-
plexes of the ligands and prepared more soluble salen-type
ligands as a host molecule by introducing hydrophobic groups
such as methyl or t-butyl groups into the benzene rings of the
ligands H3L1–H3L3.
We studied the keto–enol tautomeric forms of the ligands in
the polar and non-polar organic solvents. The most useful tech-
niques to investigate the tautomeric forms (Fig. 2) of these
ligands are UV–vis and NMR spectroscopy, while IR seems
of limited value here because the location of the ꢃ(C=O) and
ꢃ(C–O) stretches in the spectra is obscured by the abundance
of aromatic skeletal modes. In order to investigate the keto–
enol tautomeric forms of the free ligands, the UV–vis spectra
Results and Discussion
The analytical data of the compounds together with physical
properties are summarized in Table 1. The data obtained from
the analytical and spectroscopic analyses correspond well with
the general formula [M2(Ln)]Cl [n: 1, 2, or 3; M: Cu or Cd] for
the ligands H3L1–H3L3 and [ML]Cl2 for the ligand H2L. First,
we prepared the starting materials 2-(2-aminoethyl)imino-
methylphenol (H3A) and 2-(3-aminopropyl)iminomethyl-6-
methoxyphenol (H3B) (Fig. 1). One of the major difficulties
in preparing the starting materials H3A and H3B is due to
the formation of the bis product. The other problem is very
low yield. Therefore, the reaction conditions must be well ad-
justed. In addition, they are not stable for a long time at room
temperature, and gradually decompose to the oxidized prod-
ucts. Therefore, these starting materials were used immediate-
ly without further purification. The ligands H3L1–H3L3 and

870 Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
Schiff-Base Ligands and Their Complexes
of the same series of compounds were investigated in five sol-
vents (hexane, methanol, ethanol, toluene, and dioxane–water
(50%) mixture). The stronger solute/solvent hydrogen bonding
by the hydroxy or azomethine groups, as well as the increasing
importance of the solvent polarity/polarizability in the stabili-
zation of the electronic excited state, lead to a hypsochromic
shift with both increasing solvent hydrogen-bond acceptor
basicity and solvent polarity/polarizability. This suggests that
most of the solvatochromism in the ligands is due to the sol-
vent polarity and basicity rather than to the solvent acidity.
As the number of CH2 groups in ROH increases, the dielectric
constant decreases. If the molecules contain a polar component
(OH) and a non-polar component (R), then the polarity of a
compound reflects the balance between these two components.
As the relative amount of hydrocarbon character increases, the
polarity decreases. Note that hexane, which is 100% hydrocar-
bon, is the least polar solvent. The UV spectral data of the
ligands are given in the experimental section. For example,
in hexane and toluene, the ligand H3L1 shows eight bands in
the 476–249 nm range. However, in ethanol and methanol, li-
gand H3L1 has bands in the 429–276 nm range. In non-polar
solvents, the bands observed in the 476–322 nm range may
be assigned to the ketoamine tautomeric form.21 Since, in this
media, there are no interactions between the ligands and sol-
vent molecules, the ligands prefer the keto-form. However, in
polar solvents, the polar groups of the ligands interact with the
polar edge of the solvent and, therefore, the hydrogen bonding
forms among the ligands and solvent molecules. The ligands
prefer the enolimine tautomeric forms. In the toluene and
hexane solvents, the free ligands have the ketoamine form.
However, in dioxane–water and MeOH solvents, the enolimine
forms have been observed.22 Other ligands have similar elec-
tronic transitions in polar and non-polar organic solvents.
Electronic spectral data of the Schiff-base ligands and their
Cd(II) and Cu(II) complexes in an EtOH solution are given in
Table 2. The absoption spectra of the ligands are characterized
by several absorption bands in the region 249–461 nm, which
CH
CH
3
3
O
N
N
HN
N
N
OH
N
N
ꢀ
ꢀ
may be assigned to ꢀ–ꢀ , n–ꢀ , and charge-transfer transi-
tion, the longer wavelength bands are assigned to intramolecu-
N
ꢀ
ꢀ
lar charge transfer while the others due to ꢀ–ꢀ and n–ꢀ tran-
sition within the C=N and C–O (double character) bonds are
influenced by charge-transfer interaction. All the complexes
show an intense band in the 324–303 nm range, which is as-
OH
HO
OH
HO
ꢀ
signed to a ꢀ–ꢀ transition associated with the azomethine
linkage.23 The spectra of the complexes show intense bands
in the high-energy region in the 487–358 nm range, which
can be assigned to charge transfer L ! M bands.24 In the spec-
tra of the Cu(II) complexes, the bands observed in the 628–
549 nm region can be attributed to d–d transitions of the metal
ions. These values are of particular importance since they were
highly dependent on the geometry of the molecule. It is known
that the transitions from a square-planar structure to a deformed
tetrahedral structure leads to a red shift of absorption in the
electronic spectra.25 Thus, the smaller value of the wavelength
of the band corresponding to the transitions resembles the
geometry of the complex and that of a square-planar complex.
IR data of the ligands and their complexes are given in
CH
O
3
CH
3
N
HN
HN
NH
NH
O
N
N
N
O
HO
OH
O
Fig. 2. Keto–enol tautomeric forms of the Schiff-base
ligand H3L1.
Table 2. The Electronic Data of the Ligands and Their Complexes
ꢂ
max/nm ("/Mꢅ1 cmꢅ1
)
Compound
ꢅeff (BM) d ! d
ꢀ
350 (1214), 319 (1896) 258 (1720), 52 (1559)
349 (1099), 309 (1347) 274 (1588), 285 (1462), 2.3
253 (1754)
334 (1089) 296 (1590), 249 (1718)
350 (1078), 332 (1155) 296 (1634), 279 (1327), 2.8
281 (1433) 221 (1577)
286 (1528) 256 (1831)
325 (1285), 285 (1773) 278 (1690)
ꢀ
Charge-transfer transitions n ! ꢀ
461 (1093), 411 (1137)
431 (1131), 414 (893)
ꢀ ! ꢀ
ꢀa)
H3L1
H3L2
—
—
—
—
2.5
H3L3
H2L
[Cd2(L1)]Cl
[Cd2(L2)]Cl
[Cd2(L3)]Cl
—
—
Diamag.
Diamag.
Diamag.
—
—
—
—
—
—
428 (1321), 412 (1066)
419 (887), 398 (1285)
361 (1259)
383 (1021)
443 (1102), 361 (1391)
422 (1052)
1.9
25.4
27.1
26.8
57.7
30.5
29.7
32.2
61.3
[Cd(H2L)]Cl2 Diamag.
358 (1252)
281 (1488)
281 (1387)
303 (1398)
284 (1507)
263 (1677)
221 (1533)
220 (1584)
281 (1587)
250 (1788)
[Cu2(L1)]Cl
[Cu2(L2)]Cl
[Cu2(L3)]Cl
1.10
1.11
1.08
612 (378), 581 (299) 361 (1195)
562 (307), 549 (334) 359 (1097)
628 (417)
613 (403)
373 (983)
377 (1123)
[Cu(H2L)]Cl2 1.71
a) ꢀ: ꢁꢅ1 cm2 molꢅ1
.

H. Demirelli et al.
Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
871
Table 3. Infrared Spectral Dataa) for the Schiff-Base Ligands and Their Metal Complexes (cmꢅ1
)
Compound
H3L1
ꢃ(OH)
ꢃ(CH3) ꢃ(ꢆCH2) ꢃ(ꢆ0CH2) ꢃ(O–H N) ꢃ(CH=N)b) ꢃ(CH=N)c) ꢃ(M–O) ꢃ(M–N)
ꢇꢇꢇ
3422 br
3420 br
3408 br
3370 br
—
—
—
3372 br
—
—
2945 s
2959 s
2953 s
—
2899 s
2901 s
2906 s
2901 s
2904 s
2905 s
2909 s
2904 s
2897 s
2898 s
2897 s
2900 s
2862 s
2868 s
2860 s
2835 s
2865 s
2874 s
2868 s
2842 s
2863 s
2866 s
2865 s
2864 s
2741 m
2590 m
2592 m
2590 m
—
1634 s
1633 s
1632 s
1631 s
1610 s
1605 m
1610 m
1614 s
1620 s
1613 s
1617 s
1618 s
1605 m
1612 m
1605 m
1620 m
1590 w
1605 w
1590 s
—
—
—
—
—
—
—
—
H3L2
H3L3
H2L
[Cd2(L1)]Cl
[Cd2(L2)]Cl
[Cd2(L3)]Cl
[Cd(H2L)]Cl2
[Cu2(L1)]Cl
[Cu2(L2)]Cl
[Cu2(L3)]Cl
[Cu(H2L)]Cl2
2948 s
2952 s
2957 s
—
2947 s
2955 s
2951 s
—
532 w
563 w
509 w
532 w
503 w
562 w
534 w
534 w
434 w
428 w
455 w
434 w
444 w
444 w
448 w
442 w
—
—
—
—
—
—
—
1597 w
1595 s
1604 w
1600 w
1600 w
—
3370 br
a) br (broad), s (strong), m (medium), w (weak). b) and c) Salicylidene and diformyl fragments, respectively.
Table 3. In the spectra of the starting compounds H3A and
H3B with the final compounds H3L1–H3L3 and H2L, the broad
bands in the 3370–3422 cmꢅ1 range can be attributed to the
ꢃ(OH) cmꢅ1 vibration. In the Cd(II) and Cu(II) complexes
(except the complexes of the ligand H2L), the bands due to
the OH modes are no longer observed, denoting that all the
hydroxy protons are displaced by Cd(II) and Cu(II) ions, lead-
ing to covalent ꢃ(M–O) bonding with the ligands. In the li-
gands, the broad bands in the range 2741–2590 cmꢅ1 come
from the intramolecular hydrogen bonding along the phenolic
–OH and azomethine –CH=N– groups; these bands disappear
as complexation proceeds between the oxygen and nitrogen
atoms with metals ions. The ꢃ(CH=N) band (1634–1605
cmꢅ1) is shifted to lower wavenumbers denoting that the nitro-
gen atom of the azomethine group coordinates to the Cd(II)
and Cu(II) ions. Because the H2L ligand occurs as mononu-
clear complexes, the O–H group of the o-vanillin fragment
does not coordinate to the metal ions; in the complexes of this
ligand, the ꢃ(OH) vibration is shown at 3372 and 3370 cmꢅ1
for the Cd(II) and Cu(II) complexes, respectively. The bonding
of the Cd(II) and Cu(II) ions to the ligands through the nitro-
gen and oxygen atoms is further supported by the presence of
new bands in the 563–503 and 455–428 cmꢅ1 range due to
ꢃ(M–O) and ꢃ(M–N), respectively.
The molar conductivity data of all compounds in EtOH
(ꢃ10ꢅ3 M solutions) are given in Table 2. The conductivity
data of the Schiff-base ligands are very low and they can be
regarded as non-electrolytes.26 Moreover, all complexes have
cationic form and their conductivity data are quite high (in
the range 25.4–61.3 ꢁꢅ1 cm2 molꢅ1); it may be suggested that
they are weak electrolytes.
The magnetic moments (as BM) of the complexes were
measured at room temperature and the data are given in
Table 2. The Cd(II) complexes of the ligands H3L1–H3L3
and H2L are found in diamagnetic character and tetrahedral
geometry around the metal ion. The structures of the mono-
and binuclear Cu(II) complexes are supported by the magnetic
moment data. The measured magnetic moment value of the
Cu(II) complex of the ligand H2L is 1.71 BM, which is very
close to that of the spin only value (1.73 BM) expected for a
complex having one copper(II) ion with a single unpaired elec-
tron located in an essentially dx2ꢅy2 orbital. This value suggests
that the copper atom is in a tetrahedral environment in its che-
lates.27,28 The observed magnetic moment values of the binu-
clear Cu(II) complexes of the ligands H3L1–H3L3 lie in the
range from 1.08 to 1.11 BM. This confirms the presence of an-
ti-ferromagnetic interaction between the two copper(II) ions.
Binuclear Cu(II) complexes have a square-planar structure
around the central metal ions.
The formulation of the ligands are deduced from analytical
1
data, H and 13C NMR, and further supported by mass spec-
troscopy. The relatively low intensities of the molecular ion
peaks, [M]þ, are indicative of the ease of fragmentation of
the compounds, and this may reflect the number of hetero-
atoms present in each structure. The spectra of the ligands
H2L and H3L1–H3L3 show peaks at m=e 427, 455, 497, and
585, respectively. These peaks can be attributed to the molecu-
lar ion peaks [M]þ. All the ligands decompose in a similar
way. In the mass spectra of the ligands, the highest intensity
peaks are at m=e 130 (100%), 160 (100%), and 202 (100%),
and may be assigned to the [C8H6N2O]2þ, [C9H8N2O]2þ
,
and [C12H14N2O]2þ ions, respectively, which are formed by
the loss of other parts (salicylidene moiety together with di-
amine fragment) of the molecular ions.
1
The H and 13C NMR spectra of the ligands were recorded
in CDCl3 and the data associated with this method are given in
the Supplementary Data section. All the ligands have two dif-
ferent surroundings. These are the 2-hydroxy-5-methyl- or -5-
t-butylisophthalaldehyde, phthalaldehyde, and salicylidene
moiety. The chemical shifts of these are different. In all li-
gands, there are two azomethine groups in the different cen-
ters. In addition, there are also methylene and propylene chains
as bridges. In the methylene chain, there are two CH2 groups.
In other words, in the propylene chain, there are three CH2
groups. These are shown in the different regions. The ligands
show approximately the same peaks. For example, in the
1H NMR spectrum of the ligand H3L1, the triplets in the
3.26–3.18 ppm range can be attributed to the protons of the
CH2 groups. The protons of the methyl group of the aromatic
ring are observed at 2.29 ppm as a singlet. The multiplets in the
6.62–7.33 ppm range may be assigned to the protons of ben-
zene rings. The azomethine protons of the ligand are observed
in the 8.36 and 8.42 ppm range as a singlet. A singlet at 10.45
ppm may be attributed to the OH proton. When D2O is added
to the solution of the ligands, the OH peaks disappear. Other
ligands also show similar spectral properties. In the spectrum

872 Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
Schiff-Base Ligands and Their Complexes
Table 4. Successive Protonation Constants of H3L1, H3L2, and H3L3 Schiff Bases in 50% Dioxane–Water Mixture (ꢅ ¼ 0:100 M
NaClO4, 25:00 ꢁ 0:02 ꢂC)
H
H
H
H
H
H
H
Compound
Log K1
11:21 ꢁ 0:02 10:78 ꢁ 0:02
11:20 ꢁ 0:02 10:75 ꢁ 0:02 10:03 ꢁ 0:02 8:20 ꢁ 0:02 6:98 ꢁ 0:02 5:30 ꢁ 0:02 3:15 ꢁ 0:02
11:18 ꢁ 0:02 10:70 ꢁ 0:02 9:51 ꢁ 0:02 8:18 ꢁ 0:02 6:95 ꢁ 0:02 5:25 ꢁ 0:02 3:11 ꢁ 0:02
Log K2
Log K3
Log K4
Log K5
Log K6
Log K7
Log ꢇH
H3L1
H3L2
H3L3
9:56 ꢁ 0:02 8:20 ꢁ 0:02 6:99 ꢁ 0:02 5:23 ꢁ 0:02 3:08 ꢁ 0:02
55.05
55.61
54.88
of the Cd(II) complex of the H3L1 ligand, signals due to the
hydrogen atom of the azomethine groups of the ligand shifted
downfield according to the free ligand. This shows that the
nitrogen atom of the azomethine group coordinates to the met-
al ions. The OH signals observed for the ligands disappear and
this situation confirms that the phenolic C–OH group coordi-
nates to the transition-metal ions. In the complex, the signals
of the methylene groups also shifted downfield.
CH
3
CH
3
N
OH
N
7
6
3
N
OH
3
N
7
6
5
N
N
4
5
N 4
N
13C NMR spectra of the ligands and their Cd(II) complexes
were recorded in CDCl3 and are a good diagnostic tool for
determining the mode of bonding. The 13C NMR spectra of the
ligands and their complexes demonstrate similar properties. In
the spectrum of the ligand H3L1, signals due to the azomethine
carbon atoms are observed at 168.48 and 163.05 ppm. In the
Cd(II) complex, these signals are shifted downfield, suggesting
the involvement of the nitrogen lone pair in coordination with
the metal ion. The signals of the carbon atoms of the benzene
rings are observed approximately in the 118.95–157.75 ppm
range and the signals observed in the 61.79–62.50 ppm range
may be attributed to the carbon atoms of the CH2 groups.
The signal at 21.65 ppm can be attributed to the methyl carbon
atom of the benzene ring. However, in the spectra of the Cd(II)
complex, the signals due to these groups are observed at 60.05,
62.21, and 21.12 ppm. It may be suggested that these groups
are not affected by complexation.
HO
OH
OH
HO
2
2
1
1
O
O
CH
CH
3
3
H3L1
H3L2
C(CH )
3
3
N
OH
N
7
3
6
5
N
N4
HO
OH
1
2
O
O
CH
CH
3
3
H3L3
Protonation Constants of Schiff Bases. Table 4 gives
the stoichiometric protonation constants for the Schiff bases
H3L1–H3L3 in a 50% dioxane–water mixture (v/v) at 25 ꢁ
0:02 ꢂC. The chemical structures of the Schiff bases are shown
Fig. 3. The protonation regions of the Schiff-bases ligands
H3L1–H3L3.
in Fig. 3 without any information about the protonation re-
H
the same protonation constant values. However, when the val-
ues in Table 5 are examined, the protonation constants in the
related regions are different. The difference in this region
can be explained with intermolecular hydrogen bonds.
gions. As it can also be seen from Fig. 3, there are Log K1
H
,
Log K2H, and Log K3 values. They show the phenolic oxygen
atoms numbered with (1), (2), and (3). In addition, there are
H
Log K4H, Log K5H, Log K6H, and Log K7 belonging to the
Potentiometric equilibrium titration curves of the systems
consisting of the free ligand H3L1 and its Cu(II) and Cd(II)
complexes are presented in Fig. 4. The ‘‘a’’ represents the num-
ber of moles of base added to the experimental solution per
mole of the ligand. In the potentiometric titration curve of the
ligand H3L1, there are four inflections at a ¼ 1, a ¼ 2, a ¼ 3,
and a ¼ 5 (where a = moles of base added per mole of ligand).
From a ¼ 0 to 1, a ¼ 1 to 2, a ¼ 2 to 3, a ¼ 3 to 5, and a ¼ 5 to
7, there are three narrow and two extensive buffer regions. The
first three narrow buffer regions correspond to the completion
of neutralization of the first two most acidic and one neutral
ammonium ions. In extensive buffer regions, the other ammo-
nium group and phenolic groups are neutralized. The buffer re-
gions at higher pH are related to the association of the ammo-
nium and phenolic groups in the sequential overlapping steps.
This is also illustrated in the species distribution of the li-
gand H3L1 in Fig. 5. At pH < 4, the ligand exists in the fully
protonated form as H7L4þ. As the pH increases, the ligand los-
es its protons from the imine nitrogens in order to form the
numbers (4), (5), (6), and (7), showing the nitrogen atoms.
The protonation regions of the phenolic oxygens in the Schiff
bases H3L2 and H3L3 are different from the Schiff base H3L1.
This situation may be due to the resemblance of the chemical
structure of the ligands H3L2 and H3L3. Moreover, there are
methoxy groups on these ligands. The methoxy groups on
the Schiff-base ligands H3L3 and H3L2 are more electron-
offering, and the ꢀ-electron density of the methoxy groups can
delocalize on the aromatic ring with resonance.
When the structures of the Schiff bases are examined, the
phenolic oxygens with numbers (2) and (3) (in the ligand
H3L1), with numbers (1) and (2) phenolic oxygens (in the
ligands H3L2 and H3L3), the imine nitrogens with numbers
(4) and (5) (in the ligand H3L1), and the numbers (6) and (7)
(in the ligands H3L2 and H3L3) imine nitrogens are identical.
Martell et al. obtained similar protonation regions in their
study dealing with macrocyclic ligands like our ligands.29,30
In fact, the identical protonation regions are supposed to have

H. Demirelli et al.
Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
873
Table 5. Logarithms of the Stability Constants of the Complexes Cu(II) of the Schiff Bases, in 50% Dioxane–Water Mixture
(ꢅ ¼ 0:100 M NaClO4, 25:00 ꢁ 0:02 ꢂC)
Equilibrium quotient
[Cu2(L1)Cl]Cl
[Cu2(L2)Cl]Cl
[Cu2(L3)Cl]Cl
[Cd2(L1)Cl]Cl
[Cd2(L2)Cl]Cl
[Cd2(L3)Cl]Cl
Log ꢇ
Log ꢇ
Log ꢇ
Log ꢇ
Log ꢇ
Log ꢇ
[H4ML]/[M][L][H]4
[H3ML]/[M][L][H]3
[H2ML]/[M][L][H]2
[HML]/[M][L][H]
[ML]/[M][L]
54:78 ꢁ 0:04
47:06 ꢁ 0:04
38:25 ꢁ 0:03
28:42 ꢁ 0:03
17:99 ꢁ 0:02
55:65 ꢁ 0:04
47:46 ꢁ 0:03
38:50 ꢁ 0:03
28:25 ꢁ 0:02
17:80 ꢁ 0:02
55:57 ꢁ 0:06
47:43 ꢁ 0:05
38:41 ꢁ 0:04
28:31 ꢁ 0:04
17:75 ꢁ 0:03
49:5 ꢁ 0:1
42:28 ꢁ 0:08
34:63 ꢁ 0:08
25:68 ꢁ 0:06
15:71 ꢁ 0:06
49:03 ꢁ 0:09
42:50 ꢁ 0:07
34:67 ꢁ 0:07
25:73 ꢁ 0:05
15:75 ꢁ 0:04
48:95 ꢁ 0:10
42:41 ꢁ 0:08
34:54 ꢁ 0:08
25:58 ꢁ 0:07
15:61 ꢁ 0:07
100
90
CuH4L
CuL
80
70
CuH3L
Cd+2
CuH2L
CuHL
+4
H7L1
60
%
50
40
30
20
10
Cu+2
Cu
2.0
3.0
4.0
5.0
6.0
7.0
8.0
9.0
10.0
11.0
-Log[H+]
12.0
Fig. 6. Species distribution diagram for the Cu–L1 Schiff-
base system in 1:1 molar ratio as a function of the pH.
ꢅ ¼ 0:100 M NaClO4, t ¼ 25:0 ꢂC, TL ¼ 1:5 ꢄ 10ꢅ3 M,
TCu ¼ 1:5 ꢄ 10ꢅ3 M, % = percentage concentration of
species.
Fig. 4. Potentiometric titration curves for the Schiff base
H3L1 and 1:1 stoichiometries of Cu(II) and Cd(II) to the
Schiff base H3L1 (ML) as a function of added NaOH (a =
moles of base added per mole of metal ion (or ligand)).
tonation constant, log ꢇH, for the Schiff base H3L2 is 55.61,
whereas for the corresponding smaller Schiff base H3L3 con-
taining the t-butyl group is 54.88. Thus, the Schiff base H3L2
displays a much stronger overall basic character than the li-
gand H3L3. These results can be explained in that the t-butyl
group in H3L3 Schiff base gives more electron density than
the ligand H3L2 with the inductive effect.
100
H7L+4
H6L+3
90
H5L+2
L-3
80
70
60
50
40
30
20
10
0
H3L
H4L+
H2L-
Stability Constants of the Schiff-Bases Complexes. The
potentiometric titration curves of the ligand H3L1 with equiv-
alents of ligand to metal ion for Cd(II) and Cu(II) are shown in
Fig. 6. The metal ions depress the titration curve of the free li-
gand by the release of protons according to the abilities of the
metal ions to bind to the ligand Schiff bases. As the titration
curves of the complexes formed by cadmium and copper to-
gether with the Schiff base H3L1 are examined, two inflection
points can be observed at a ¼ 2 and a ¼ 6 for Cd(II), and at
a ¼ 3 and a ¼ 6 for Cu(II). The presence of an inflection point
on the titration curves at a ¼ 3 and a ¼ 2 for the Cu(II) and
Cd(II), respectively, can be explained by the formation of pro-
tonated complexes. The ligand H3L1, because of its geometric
shape, can not form bonds with the metal ion at every proton-
ation region; as a result, the protonated complexes are formed
in an acidic medium. For example, the oxygen atoms of the ba-
sic phonolate anion and one of the nitrogen atoms of the imine
group are protonated in the Cu–L1 system and the CuH4L
complex forms, and oxygens of the phonolate anion and two
of the nitrogen atoms of the imine group are protonated in
the Cd–L1 system and CdH5L complex forms. The Cu(II)
-2
HL
2
4
6
8
10
-Log H+
12
Fig. 5. Species distribution diagram for the Schiff-base
(H3L1) systems as a function of the pH. ꢅ ¼ 0:100 M
NaClO4, t ¼ 25:0 ꢂC, TL ¼ 1:5 ꢄ 10ꢅ3 M, % = percent-
age concentration of species.
H6L3þ, H5L2þ, and H4Lþ species, respectively. H6L3þ is the
predominant species in the pH range 2–7 and its maximum
concentration reaches 85.5% at pH ¼ 4:2. When the pH is
above 6, one of the three phenolic hydroxy protons begins to
deprotonated to form H3L, which reaches its maximum con-
centration (72.0%) at pH ¼ 9:0. The log values of overall pro-

874 Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
Schiff-Base Ligands and Their Complexes
Table 6. Logarithms of the Stability Constants of the Com-
plexes Cd(II) of the Schiff Bases, in 50% Dioxane–Water
Mixture (ꢅ ¼ 0:100 M NaClO4, t ¼ 25:00 ꢁ 0:02 ꢂC)
ation of the cadmium ion (Cd(II)) (its radius is larger than the
Cu(II) ion). Therefore, the stability constants of the Cd(II)
complexes are much smaller than the stability constants of
the Cu(II) complexes.
The species distribution curve of the Cd–L1 system is very
different from the Cu–L1 system because the hydrolysis of
cadmium is more dominant than the formation of Cd–L1 com-
plexes. In the Cd–L1 system, the complex CdH2L is the most
predominate species (Fig. 7) and its maximum, 46.6%, is
reached at pH ¼ 9:5. The next most predominate species are
the complexes CdH5L, CdH3L, CdHL, and CdL. These com-
plexes have almost the same percentages. The percentage of
the species CdH4L is the smallest in this system.
Equilibrium quotient
H3L1
H3L2
H3L3
Log ꢇ
Log ꢇ
Log ꢇ
[H5ML]/[M][L][H]5 49:5 ꢁ 0:1 49:03 ꢁ 0:09 48:95 ꢁ 0:10
[H4ML]/[M][L][H]4 42:28 ꢁ 0:08 42:50 ꢁ 0:07 42:41 ꢁ 0:08
[H3ML]/[M][L][H]3 34:63 ꢁ 0:08 34:67 ꢁ 0:07 34:54 ꢁ 0:08
[H2ML]/[M][L][H]2 25:68 ꢁ 0:06 25:73 ꢁ 0:05 25:58 ꢁ 0:07
[HML]/[M][L][H]
[ML]/[M][L]
15:71 ꢁ 0:06 15:75 ꢁ 0:04 15:61 ꢁ 0:07
5:15 ꢁ 0:04 5:11 ꢁ 0:03 5:07 ꢁ 0:05
100
80
Supporting Information
Supplementary data associated with this article can be found in
the online version of this journal. This material is available free of
charge on the web at http://www.csj.jp/journals/bcsj/.
CdH-2
60
%
Cd
CdH5L
L
References
CdH2L
CdHL
1
2
3
E. Hadjoudis, Mol. Eng. 1995, 5, 301.
T. Dziembowska, Pol. J. Chem. 1998, 72, 193.
A. Filarowski, T. Glowiak, A. Koll, J. Mol. Struct. 1995,
40
20
0
CdH3L
CdL
CdH4L
484, 75.
4 T. Sekikawa, T. Kobayashi, T. Inabe, J. Phys. Chem. A
1997, 101, 644.
CdH-1
5 B. Zheng, S. Brett, J. P. Tite, T. A. Brodie, J. Rhodes,
Science 1992, 256, 1560.
2.00
4.00
6.00
8.00
10.00
12.00
- Log [H+]
˘
E. Kꢂlꢂc¸, F. Koseoglu, E. Canel, Indian J. Chem. 1997, 36,
6
115.
7
¨
Fig. 7. Species distribution diagram for the Cd–L1 Schiff-
base system in 1:1 molar ratio as a function of the pH.
ꢅ ¼ 0:100 M NaClO4, t ¼ 25:0 ꢂC, TL ¼ 1:5 ꢄ 10ꢅ3 M,
TCd ¼ 1:5 ꢄ 10ꢅ3 M, % = percentage concentration of
species.
G. S. Patterson, R. H. Holm, Bioinorg. Chem. 1975, 4, 257.
D. B. Rorabacher, M. J. Martin, M. J. Koenigbauer, M.
Malik, P. R. Schroeder, J. F. Eudicott, L. A. Ochrymowyez, in
Copper Coordination Chemistry: Biochemical and Inorganic
Perspectives, ed. by K. D. Karlin, J. Zubieta, Academic Press,
New York, 1983.
8
ion, which is known to form complexes of tetragonal shape,
does not show an inflection point at a ¼ 4; this is because a
proton is removed from the protonated complex and the pro-
tonation constants of the ligand intersects each other.
9
A. E. Martell, R. J. Motekaitis, The Determination and Use
of Stability Constants, VCH, New York, 1988.
10 D. D. Perrin, W. L. F. Armarego, Purification of Labora-
tory Chemicals, Pergamon, 1992.
The stoichiometric stability constants of the metal com-
plexes (Cu(II) and Cd(II)) of the investigated Schiff bases were
determined in the 50% dioxane–water mixture at 25 ꢂC and
these constants are given in Tables 5 and 6. The metal ions
Cu(II) and Cd(II) were found to combine easily with the Schiff
bases H3L1, H3L2, and H3L3 to form deprotonated (ML) and
multiprotonated (MHnL) in concentrations depending on the
pH value of the solution. When the values in Tables 5 and 6
are examined, the following row is obtained for stability con-
stants of the mononuclear complexes (ML) of the Schiff bases:
H3L1 > H3L2 > H3L3. When these rows, which are obtained
for stability constants of the Schiff bases, are compared with
the row that is obtained for protonation constants, it is seen
that there is no correlation between the stability constants and
protonation constants. The order of the stability constants for
all ligands was found to be Cu(II) > Cd(II), in agreement with
the increasing acidity of the metal ion. Comparison of the
Cu(II) and Cd(II) complexes of these three ligands indicate
that the size of the cavity in the ligands and metal ion radius
plays an important factor in the stability of mononuclear com-
plex formation.29,31 The cavity is too small to allow complex-
11 G. Schwarzenbach, M. Flascha, Complexometric Titra-
tions, Methuen, London, 1957.
12 F. J. C. Rossotti, H. Rossotti, J. Chem. Educ. 1965, 7, 375.
13 G. Gran, Acta Chem. Scand. 1950, 4, 559.
14 T. N. Sorrell, C. O’Connor, O. P. Anderson, J. H.
Reibenspies, J. Am. Chem. Soc. 1985, 107, 4199.
15 H. Rossotti, The Study of Ionic Equilibria: An Introduction,
Longman, London, 1978.
16 P. May, D. R. Williams, P. W. Linder, R. G. Torrington,
Talanta 1982, 29, 249.
17 P. Serjant, Potentiometry and Potentiometric Titrations,
Wiley, New York, 1984.
18 D. M. Boghaci, M. Lashanizadegan, J. Sci., Islamic Repub.
Iran 2000, 11, 301.
19 D. M. Boghaci, M. Lashanizadegan, Synth. React. Inorg.
Met.-Org. Chem. 2000, 30, 1535.
20 D. Cunningham, P. McArdle, M. Mitchell, N. N.
Chonchubhair, M. O’Gara, F. Franceschi, C. Floriani, Inorg.
Chem. 2000, 39, 1639.
21 M. Tumer, C. C¸ elik, H. Koksal, S. Serin, Transition Met.
¨
Chem. 1999, 24, 525.
¨

H. Demirelli et al.
22 a) M. Tumer, H. Koksal, S. Serin, M. Dꢂgrak, Transition
Bull. Chem. Soc. Jpn. Vol. 79, No. 6 (2006)
875
˘
Met. Chem. 1999, 24, 13. b) A. Golcu, M. Tumer, H. Demirelli,
516. b) J. Mukherjee, R. Gupta, T. Mallah, R. Mukherjee, Inorg.
Chim. Acta 2005, 358, 2711.
¨
¨
¨
¨
¨
˘
28 M. Dolaz, M. Tumer, M. Dꢂgrak, Transition Met. Chem.
R. A. Wheatley, Inorg. Chim. Acta 2005, 358, 1785.
23 H. Koksal, M. Dolaz, M. Tumer, S. Serin, Synth. React.
Inorg. Met.-Org. Chem. 2001, 31, 1141.
24 R. Atkins, G. Brewer, E. Kokot, G. M. Mockler, E. Sinn,
Inorg. Chem. 1985, 24, 127.
25 A. L. Nivorozhkin, H. Toflund, P. L. Jorgensen, L. E.
Nivorozhkin, J. Chem. Soc., Dalton Trans. 1996, 1215.
26 W. J. Geary, Coord. Chem. Rev. 1971, 7, 81.
¨
2004, 29, 528.
29 D. Kong, J. Mao, A. E. Martell, A. Clearfield, Inorg. Chim.
Acta 2002, 338, 78.
30 G. Shangguan, A. E. Martell, Z. Zhang, J. H. Reibenspies,
Inorg. Chim. Acta 2000, 299, 47.
¨
¨
31 A. E. Martell, R. D. Hancock, Metal Complexes in
Aqueous Solutions, Plenum Press, New York, 1996.
27 a) M. Dolaz, M. Tumer, Transition Met. Chem. 2004, 29,
¨