Potentiometric, equilibrium studies and thermodynamics of novel thiosemicarbazones and their bivalent transition metal(II) complexes

Article abstract of DOI:10.1016/j.molliq.2016.04.005

Protonation constants of novel thiosemicarbazones including 2-(1-(2-phenyl-hydrazono)-propan-2-ylidene) hydrazine-carbothioamide (TPHP) and N-methyl-2-(1-(2-phenyl-hydrazono)-propan-2-ylidene)hydrazinecarbo-thioamide (MTPHP) ligands and their corresponding metal-ligand formation constants with Cu(II), Ni(II), Mn(II) and Co(II) ions were determined at 15 °C, 25 °C and 35 °C in 50% DMSO solution at I = 0.1 mol·dm^{- 3} NaNO₃. The stability order of complexes with reference to the metal ions has been followed this order Cu(II) > Ni(II) > Co(II) > Mn(II) in concord with the Irving-Williams stability order. Also, chemical equilibrium studies indicated that substitution of hydrogen in thiosemicarbazone (TPHP) by methyl group causes lowering of acidity as a result of the electron-releasing effect of the methyl group in the substituted amino group of the thiosemicarbazide moiety. The speciation of different species in solution has been evaluated as a function of pH. Additionally, the effect of temperature on protonation of thiosemicarbazone ligands and formation of their M(II)-thiosemicarbazone complexes was investigated. The thermodynamic parameters (ΔH, ΔS and ΔG) were calculated and discussed. It was found that both log K₁ and -ΔH₁, for M(II)-thiosemicarbazone complexes are somewhat larger than log₁₀ K₂ and -ΔH₂, indicating a change in the dentate character of these ligands from tridentate (SNN-donors) in 1:1 chelates to bidentate (SN-donors) in 1:2; M:L chelates. Also, the lower values of log₁₀ K₂ and -ΔH₂ than log K₁ and -ΔH₁ may be attributed to steric hindrance produced by the entrance of a second molecule.

Full text of DOI:10.1016/j.molliq.2016.04.005

Journal of Molecular Liquids 219 (2016) 914–922
Contents lists available at ScienceDirect
Journal of Molecular Liquids
journal homepage: www.elsevier.com/locate/molliq
Potentiometric, equilibrium studies and thermodynamics of novel
thiosemicarbazones and their bivalent transition metal(II) complexes
Abeer Taha Abd El-Karim ⁎, Ahmed A. El-Sherif
Department of Chemistry, Faculty of Science, Cairo University, Cairo, Egypt
a r t i c l e i n f o
a b s t r a c t
Article history:
Protonation constants of novel thiosemicarbazones including 2-(1-(2-phenyl-hydrazono)-propan-2-ylidene)
hydrazine-carbothioamide (TPHP) and N-methyl-2-(1-(2-phenyl-hydrazono)-propan-2-ylidene)hydrazinecarbo-
thioamide (MTPHP) ligands and their corresponding metal-ligand formation constants with Cu(II), Ni(II), Mn(II)
Received 13 December 2015
Received in revised form 20 February 2016
Accepted 1 April 2016
and Co(II) ions were determined at 15 °C, 25 °C and 35 °C in 50% DMSO solution at I = 0.1 mol·dm−3 NaNO
3
.
Available online 26 April 2016
The stability order of complexes with reference to the metal ions has been followed this order
Cu(II) N Ni(II) N Co(II) N Mn(II) in concord with the Irving-Williams stability order. Also, chemical equilibrium stud-
ies indicated that substitution of hydrogen in thiosemicarbazone (TPHP) by methyl group causes lowering of acidity
as a result of the electron-releasing effect of the methyl group in the substituted amino group of the
thiosemicarbazide moiety. The speciation of different species in solution has been evaluated as a function of pH.
Additionally, the effect of temperature on protonation of thiosemicarbazone ligands and formation of their M(II)-
thiosemicarbazone complexes was investigated. The thermodynamic parameters (ΔH, ΔS and ΔG) were calculated
Keywords:
Thiosemicarbazone
Protonation
Thermodynamics
Copper, potentiometry
and discussed. It was found that both log K
than log10 and –ΔH , indicating a change in the dentate character of these ligands from tridentate (SNN-donors)
in 1:1 chelates to bidentate (SN-donors) in 1:2; M:L chelates. Also, the lower values of log10 and –ΔH than log K
and –ΔH may be attributed to steric hindrance produced by the entrance of a second molecule.
2016 Elsevier B.V. All rights reserved.
1 1
and –ΔH , for M(II)-thiosemicarbazone complexes are somewhat larger
K
2
2
K
2
2
1
1
©
1
. Introduction
Protonation constants are important physicochemical parameters,
Thiosemicarbazones and their complexes have been extensively stud-
ied because they have a wide range of actual or potential medical applica-
tions [13–17] which include notably antiparasital [18], antibacterial [19]
antitumor activities [20], antiviral [21], fungicidal [22] and antineoplastic
[23]. In spite of the large interest always shown in the coordination prop-
erties of thiosemicarbazone ligands and the attention paid in recent years
to the possible variable biological activities of their metal complexes,
studies dealing with protonation and complex formation equilibria of
thiosemicarbazones and their complexes are still very rare. These find-
ings stimulated our interest to study the solution equilibria of novel
thiosemicarbazones and their metal complexes. In conjunction with our
research program [24–33] directed to study ligands of biological signifi-
cance and their complexes, the present investigation aims to study the
solution equilibria of thiosemicarbazone-ligands and their M(II)–com-
plexes in 50% DMSO–water mixture; the effects of temperature, thermo-
dynamics, nature of the central metal ion and structure of the ligands on
the protonation of novel thiosemicarbazones and their metal complexes.
which can provide critical information about drug properties such as
solubility, lipophilicity, acidity, basicity [1,2], transport behavior,
bonding to receptors [2–4], and permeability [5]. Hence, the relation-
ship between the protonation constants and structure in drug design
studies is important [2,3,6]. Protonation constants are also important
parameters for the selection of the optimum conditions in the develop-
ment of analytical methods [5,7] and choosing a suitable pH value for
carrying out spectrophotometric quantitative analyses. Additionally,
knowledge of the protonation constants of some compounds is
necessary for the calculation of the concentration of each ionized
species at any pH, which is important for the complete understanding
of the physiochemical behavior of such molecules [8] and provides
also information about the stereochemical and conformational struc-
tures of active centers of enzymes [9]. There are various techniques
such as potentiometry, conductometry and spectrophotometry that
are used in the determination of protonation constants. In this study a
potentiometric technique was employed because it has the widest
area of applicability and reliability [10–12].
2. Experimental
2.1. Materials and reagents
All chemicals utilized in this investigation were of the analytical
⁎
Corresponding author.
E-mail address: tabdelkarima@yahoo.com (A.T.A. El-Karim).
reagent grade (AR) quality and used without further purification. They
http://dx.doi.org/10.1016/j.molliq.2016.04.005
167-7322/© 2016 Elsevier B.V. All rights reserved.
0

A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
915
Scheme 1. Molecular structure of TPHP and MTPHP thiosemicarbazone compounds.
included HCl and KOH provided by BDH. Aniline, ethyl acetoacetate and
sodium nitrite were obtained from Sigma. Thiosemicarbazide and 4-(N-
methyl)-thiosemicarbazide were purchased from Merck. Metal salts in-
delivery tube and a salt bridge connected with the reference cell filled
with 0.1 M KCl solution in which saturated calomel electrode was dipped.
Temperature was maintained constant inside the cell at 25.0 ± 0.01 °C, by
the circulating water by a thermostated bath. All potentiometric measure-
ments in this study were carried out in water-DMSO mixtures containing
50% DMSO because of low solubility of the synthesized thiosemicarbazone
compounds and possible hydrolysis in aqueous solution.
cluding CuCl
8%), NiCl ·6H
drich, ≥ 99%). DMSO was provided by Aldrich Chemicals Company.
2 2 2 2
·2H O (Aldrich, ≥ 99.99%), CoCl ·6H O (Sigma-Aldrich,
9
2
2
O (Aldrich, 99.9%) and MnCl
2
·4H
2
O (Sigma-Al-
2
2
.2. Synthesis
2.4. Potentiometric titrations
.2.1. Synthesis of 1-(phenyl-hydrazono)-propan-2-one (PHP)
It is prepared as reported in the literature [34,35] as follows: in a 4-l
The protonation constants of the ligands and stability of complex for-
beaker equipped with a mechanical stirrer, 65 g (64 ml, 0.5 mol) of ethyl
acetoacetate was added to 35 g (0.53 mol) of 85% potassium hydroxide in
mation were measured potentiometrically using earlier described meth-
od [36]. pH-metric titrations were carried out by using of Metrohm 686
titroprocessor equipped with a 665 Dosimat. Double-wall glass titration
cell equipped with a magnetic stirring system was used. The cell solution
was stirred continuously at constant speed during the titration using
magnetic stirring system. The glass electrode was calibrated with stan-
dard buffer solutions, potassium hydrogen phthalate (pH = 4.008) and
1
120 ml of water. The mixture is allowed to stand at room temperature
for 24 h. Forty-seven grams (48 ml, 0.5 mol) of aniline is dissolved in
00 ml of aqueous HCl (prepared from equal volumes of concentrated
2
acid and water) in a 2-l beaker. The beaker is equipped with a mechanical
stirrer and immersed in an ice-salt bath. After the solution has cooled to
0
–5 °C, 36 g. (0.52 mol) of sodium nitrite dissolved in 1 l of water is
a mixture of KH
cording to NBS specifications [37]. The titration reaction was investigated
in presence of purified N atmosphere using standard solution of
2 2 4
PO4 and Na HPO (pH = 6.865) at 25.0 °C, prepared ac-
added slowly, with stirring, from a separating funnel. The tip of the
stem of the separating funnel dipped well below the surface of the liquid.
The rate of addition is adjusted to maintain the temperature between 0
and 5 °C. A drop of the reaction mixture is tested from time to time
with starch-iodide paper until nitrous acid persists in the solution during
a 5-min interval. The solution of potassium acetoacetate is cooled to 0 °C,
and 45 ml of concentrated HCl in 150 ml of ice water is added slowly with
stirring. The diazonium salt solution is then added over a period of 20 min,
and the mixture is made basic by the addition of 82 g of sodium acetate
dissolved in 300 ml of water. The temperature of the reaction mixture is
raised slowly to 50 °C and maintained at this temperature for 2 h; the sep-
arated solid is collected on a filter and dried. The yield of crude product is
2
−
3
0.05 mol dm sodium hydroxide free from carbon dioxide. The titration
cell was cleaned with distilled water and dried with a tissue before and
after the experiment. Covered cell calibration lid contains four holes for
Metrohm glass electrode, glass tubing for nitrogen injection, thermomet-
ric probe and plastic tube for alkali solution. Before filling of a tube with
alkali solution, the tube was washed several times with distilled water
and then washed with alkali solution at least 4 times. Also, the air bubbles
were avoided to leak in the tube in order to get accurate results for the
measured volumes. Stock solution of metal salts was prepared and stan-
dardized using complexometric EDTA titrations [38]. In order to avoid
probable hydrolysis of TPHP and MTPHP ligands in aqueous solution, po-
tentiometric titration was carried out in 50% DMSO-water mixture.
To keep the ionic strength (I) constant during the titration process,
supporting electrolyte of sodium nitrate was used i.e., the ionic strength
was maintained constant at 0.1 M sodium nitrate with the addition of
77 g (95%). Purification can be effected by recrystallization from 200 ml of
toluene. The purified product weighs 66 g (82%); m.p. 148–150 °C.
2
.2.2. Synthesis of TPHP and MTPHP thiosemicarbazone compounds
The general route of synthesis is shown as follow: Equimolar
amounts of (PHP) (0.1620 g, 1 mmol) in 25 ml ethanol with an ethanolic
solution (25 ml) of thiosemicarbazide (0.0911 g, 1 mmol) and methyl-
3
appropriate amount of 0.8 M NaNO solution.
As is known, pH-meters read -log aH+ (pH), whereas the potentio-
metric method we used for the calculation of stability constants requires
(
0.1051 g, 1 mmol) thiosemicarbazides were mixed and then refluxed
+
on a hot plate for 4–5 h. The obtained precipitates were separated out,
filtered off, washed with diethyl ether and dried overnight under silica
gel (See Scheme 1).
–log[H ] (p[H]). Hence, the first step in computations was to convert
the pH-meter readings (B) recorded in DMSO-water solutions to
+
hydrogen ion concentration [H ] [39]. This can be achieved by using
the widely used relation given by the Van Uitert and Hass equation,
Eq. (1) [40] as shown below,
2
.3. Instruments
ꢀ
þꢁ
Potentiometric measurements were made using a Metrohm 686
− log10 H ¼ B þ log10UH
ð1Þ
titroprocessor equipped with a 665 Dosimat (Switzerland-Herisau). A
thermostatted glass-cell was used equipped with a magnetic stirring sys-
tem, a Metrohm glass electrode, a thermometric probe, a microburet
where log10
U
H
is the correction factor for the solvent composition and
in DMSO-water mixtures
ionic strength for which B is read. Values of pK
w

916
A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
were determined as described previously [41]. For this purpose, various
amounts of standard NaOH solution were added to 50% aqueous DMSO
2.5. Data processing
−
3
−
solutions containing 0.1 mol·dm NaNO
3
. The [OH ] was calculated
The calculations were obtained from ca. 100 data points in each
titration using the computer program MINIQUAD-75 [45]. The species
distribution diagrams were obtained using the program SPECIES [46]
under the experimental conditions employed.
+
from the amount of base added. The [H ] was calculated from the pH
−
+
value. The product of [OH ] and [H ] was taken. The mean values obtain-
+
−
w
ed in this way for -log10 [H ][OH ] (pK = 15.52 ± 0.2) are in agree-
ment with the literature values [42].
A carbonate-free sodium hydroxide solution in 50% (by volume)
DMSO-water mixture was used as titrant and standardized against
potassium hydrogen phthalate (Analar). The apparatus, general condi-
tions, and method of calculations were the same as in previous work
3. Results and discussion
3.1. Protonation constants of thiosemicarbazone ligands
[
43,44].
The proton association constants of the ligands (TPHP or
MTPHP) were determined potentiometrically by titrating (1.25 ×
The study of protonation equilibria by the studied thiosemicarbazone
ligands cannot be carried out in aqueous solution because of the nature
of the compounds involved. These compounds are insoluble in water.
This solvent has been widely used for potentiometric determination of
protonation and formation equilibria as mentioned in literature
[47–49]. The mixture DMSO-water 50%:50% was the chosen solvent for
our study.
The stoichiometric protonation constants of the investigated
thiosemicarbazone compounds (TPHP and MTPHP) were determined
in 50% DMSO-water mixture at 25 °C and these constants are given in
Table 1. The compounds studied here have two protonation constants.
The log KSH and log KNH are related to the protonation of thiolate sulfur
and hydrazo group, respectively. This is also illustrated in the species
−
3
−3
3
1
0
mol·dm ) of the ligand solution (40 cm ). The stability
constants of the metal(II) complexes were determined using po-
tentiometric data obtained from (40 cm ) mixture containing
3
−
3
−3
MCl
1.25 × 10
The equilibrium constants evaluated from the titration data are
2
·nH
2
O (1.25 × 10
mol·dm ) + (TPHP or MTPHP)
−
3
−3
−3
−3
(
mol·dm
/ 2.5 × 10
mol·dm ).
defined by Eqs. (1) and (2), where M, L and H stand for the metal(II)
ion, thiosemicarbazone ligands (TPHP/MTPHP) and proton, respectively
h
i
pðMÞ þ qðLÞ þ rðHÞ⇌ MpðLÞ ðHÞ
ð2Þ
ð3Þ
q
r
distribution of the TPHP ligand in Fig. 1. By rising of pH, the ligand
h
i
−
(
H
2
L) loses its protons from SH group to form HL , which is the
MpðLÞ ðHÞ
q
r
predominant species in pH range 8.8–9.7. As conditions become more
βpqr
¼
p
q
r
½
Mꢀ ½Lꢀ ½Hꢀ ꢀ
alkaline, the second NH group begins deprotonation to the free ligand
2
−
L
anion which is the predominant species at pH N 11. From Table 1,
log KSH of (MTPHP) ligand is 8.48 and that of (TPHP) is 7.41 and log
NH of (MTPHP) is 11.15 and that of (TPHP) is 11.07. Substitution by
methyl group instead of hydrogen causes a 1.07 unit increase of log
SH and a 0.08 unit increase for log KNH, which means a lowering of acid-
Table 1
Protonation constants of TPHP and MTPHP thiosemicarbazones in 50% DMSO-50% H
2
O (v/v)
K
at different temperatures and I = 0.1 mol dm− NaNO
3
3
.
K
System
TPHP
T (°C)
p
q
r
log10
β
S
log KNH
11.16
log KSH
ity. This can be explained by the electron-releasing effect of the methyl
group in the substituted amino group. Also, it is observed that, the
change in log KNH is smaller than log KSH. This is in accordance with
the fact that, the inductive effect decreases as the distance increases.
The log KNH values ranges from (10.76–11.27) is similar to that
found in literature for the hydrazo moiety (10.98) [50]. Also, the order
of basicity of thiosemicarbazone ligands, is as follows; MTPHP N TPHP.
The increase in basicity or decrease in the acidity is due to
the electron-donating property (+I-effect) of the N-substituted
moiety of the thiosemicarbazide moiety i.e., Methyl N Hydrogen.
This is in accord with the published data in literature for N-
15
0
0
0
0
0
0
0
0
0
0
0
0
1
1
1
1
1
1
1
1
1
1
1
1
1
2
1
2
1
2
1
2
1
2
1
2
11.16 ± 0.03
18.71 ± 0.05
11.07 ± 0.01
18.48 ± 0.04
10.87 ± 0.05
18.21 ± 0.07
11.27 ± 0.02
19.87 ± 0.04
11.15 ± 0.04
19.63 ± 0.07
10.93 ± 0.02
19.29 ± 0.06
2.1 E-8
7.55
7.41
7.34
8.60
8.48
8.36
2
3
5
5
3.7 E-8
6.3E-8
4.3E-8
5.1E-8
4.9E-8
11.07
10.87
11.27
11.15
10.93
MTPHP
15
2
3
5
5
Fig. 1. Concentration distribution of various species as a function of pH in the TPHP system (at concentration of 1.25 mol·dm−3 for TPHP) at 25 °C and I = 0.1 mol·dm−3 NaNO
.
3

A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
917
alkyliminobis(methylphosphonic acid) derivatives whereas the pK
Table 2
Stability constants of M(II)-thiosemicarbazones in 50% DMSO-50% water at different
values of N-ethyliminobis(methylphosphonic acid) (12.20, NH) [51],
and N-methyliminobis(methylphosphonic acid) (11.75,NH) [51] are
considerably greater than that of iminobis(methylphosphonic acid)
temperatures and I = 0.1 mol dm− NaNO
3
3
.
System
T (°C)
p
q
r
log10β
S
(
11.20, NH) [52]. This means that, substitution with an alkyl group has
Cu-TPHP
15
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
1
2
1
2
1
2
1
2
1
2
1
2
1
2
1
2
1
2
1
2
1
2
1
2
1
2
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
1
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
11.84 ± 0.04
18.17 ± 0.06
22.62 ± 0.07
11.70 ± 0.05
17.93 ± 0.06
22.32 ± 0.05
11.53 ± 0.04
17.67 ± 0.08
22.02 ± 0.07
12.30 ± 0.05
18.80 ± 0.06
23.64 ± 0.06
12.14 ± 0.06
18.53 ± 0.07
23.31 ± 0.04
11.98 ± 0.05
18.29 ± 0.08
23.02 ± 0.07
11.08 ± 0.06
16.94 ± 0.09
21.20 ± 0.05
10.64 ± 0.05
16.72 ± 0.07
20.93 ± 0.07
10.49 ± 0.04
16.48 ± 0.08
20.63 ± 0.06
11.21 ± 0.03
17.40 ± 0.06
21.40 ± 0.08
11.02 ± 0.04
17.14 ± 0.06
21.12 ± 0.09
10.91 ± 0.03
16.91 ± 0.06
20.82 ± 0.08
10.18 ± 0.04
19.95 ± 0.06
10.02 ± 0.03
19.66 ± 0.07
9.90 ± 0.05
19.40 ± 0.06
10.59 ± 0.05
20.21 ± 0.07
10.47 ± 0.04
19.96 ± 0.08
10.30 ± 0.06
19.65 ± 0.06
9.87 ± 0.04
19.18 ± 0.08
9.72 ± 0.06
18.92 ± 0.09
9.60 ± 0.05
18.65 ± 0.08
10.05 ± 0.06
19.45 ± 0.09
9.91 ± 0.07
19.20 ± 0.08
9.77 ± 0.08
18.90 ± 0.09
3.8E−7
a great influence on the protonation equilibria of the neighboring
groups.
2
3
5
5
4.2E−7
5.1E−7
5.6E−7
6.2E−7
6.8E−7
7.5E−8
6.3E−8
7.7E−8
4.8E−8
5.2E−8
4.8E−8
3
.2. Stability constants of the thiosemicarbazone complexes
The stability constants of ML complexes of thiosemicarbazone
ligands with some divalent metal ions in DMSO-water solution were de-
termined by use of the Miniquad-75 computer program. Comparing the
titration curve of the free thiosemicarbaazone ligand with those of the
complexed ligand shows that addition of the copper ion to the free
thiosemicarbazone ligand solution of the complexes shifts the curve to
lower pHs. In other words, the curves of the complexes are situated at
lower pHs than the free ligand curve as they required more alkali to
have the same pH as the free ligand. This can be explained simply as a
result of proton release from the coordinated ligand, which implies
complex formation. The stoichiometric stability constants of M(II)
complexes of the investigated thiosemicarbazone ligands were deter-
mined in 50% DMSO-water mixture at different temperatures and
these constants are given in Table 2. The data also showed the formation
of the metal complexes with stoichiometric coefficients 110, 111 and
Cu-MTPHP
15
25
35
Ni-TPHP
15
2
3
5
5
1
20.
The high stability of M(II)-MTPHP complexes than M(II)-TPHP com-
plexes can be attributed to the presence of —CH group in the
Ni-MTPHP
15
3
thiosemicarbazide moiety which is in agreement with the basicity of
these ligands. This is quite reasonable because the presence of the
above —CH group (i.e. an electron-donating group) will enhance the
3
electron density by their high positive inductive effect, whereby
stronger chelation was formed. This study indicated that alkylation of
the N-terminal of thiosemicaarbazone has a highest effect on the
stability of the formed complexes. This is in accord with the published
data in literature [53].
2
3
5
5
Co-TPHP
15
3.4E−8
5.9E−8
4.2E−8
3.9E−8
3.5E−8
4.6E−8
5.1E−8
4.9E−8
7.2E−8
8.1E−8
3.4E−8
9.5E−8
25
a
The pK of the protonated complex can be calculated using Eq. (4)
[54,55].
35
15
Co-MTPHP
Mn-TPHP
Mn-MTPHP
pK ¼ logβ111− logβ110
ð4Þ
a
25
This value is in fair agreement with pK
a
of SH group (7.41) taking
into consideration the acidification upon complexation i.e. the lower
value (6.25) than that of free TPHP ligand (7.36) indicates acidification
upon coordination to Cu(II) as a representative example of protonated
metal(II)-thiosemicarbazone complexes by 1.16 pH units (7.41 to
35
15
25
6
.25). This may indicate that the NN donors of TPHP are the binding
35
sites in the protonated (CuHL) complex.
In order to investigate change in the concentration of the copper(II)
complexes with pH, the species distribution diagram for Cu-TPHP
system is examined (Fig. 2) as a representative example of Cu(II)
complexes. In the Cu-TPHP distribution diagram, it will be seen that
the complex CuHL is formed with maximum percent of 80% at pH 5.0.
The deprotonated Cu-TPHP complex is formed with maximum percent
of 63% at pH 8.0 forms. Cu(TPHP) complex is formed with maximum
2
percent of 18% at pH 8.0. Copper(II) ion was found to combine easily
with TPHP-thiosemicarbazone ligand to form deprotonated (CuL),
15
25
35
Definitions of stability constants.
ML = [ML]/[M][L].
MHL = [MHL]/[ML][H].
L = thiosemicarbazone ligands); (Charges are omitted for simplicity).
K
K
(
protonated (CuHL) and CuL
solution.
2
complexes depending on the pH of the
3
.3. The relationship between the properties of central metal ion and
stability of complexes
properties of the metal ions, such as the atomic number, ionic radius,
ionization potential and electronegativity. Here, we have discussed
the relationship between the periodic table properties [56] and the
stability constants of complexes. In this investigation, it was found
In order to explain why a given ligand prefers binding to one
metal instead of another metal, it is important to link the relation-
ship between the stability constants of metal complexes and
II
that, the formation constants of M -complexes of some transition

918
A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
Fig. 2. Concentration distribution of various species as a function of pH in the Cu-TPHP system (at concentrations of 1.25 mol·dm for Cu(II) and 1.25 mol·dm−3 for TPHP) at 25 °C and
−3
I = 0.1 mol·dm− NaNO
3
3
.
metal ions with thiosemicarbazones obeyed this arrangement:
3.4. Effect of temperature and thermodynamics
2
+
2+
2+
Mn(II) b Co
Williams' order [57].
The relation between the log10
1/r) of the studied transition metal ions represents nearly a linear
relationship. Additionally, a linear relation has been noticed between
log10 ML and the electronegativities of the metal ions under investiga-
tion (Fig. 3). This is consistent with the fact that the increase in the elec-
b Ni
b Cu
which is consistent with Irving-
The values of thermodynamic parameters that are related to the pro-
tonation of thiosemicarbazones and their metal(II) complexes have
been calculated from the temperature dependent data given in
Tables 3 and 4. Values of ΔH and ΔS were obtained by drawing the rela-
tionship between the values of equilibrium constants (logK) versus re-
ciprocal of temperature (1/T) (logK = − ΔH/RT + ΔS/R) leading to an
intercept ΔS/R and a slope –ΔH/R (Figs. 4 and 5). Main conclusions
from the data can be summarized as follows:
KML and reciprocal of the ionic radii
(
K
2
+
2+
tronegativity of metals (Mn(II) (1.55)
(
b
Co (1.88)
b Ni
2
+
1.91) b Cu (2.0)) will reduce the difference in electronegativity be-
tween the metal atom and the donor atom of the ligand. Therefore,
the metal-ligand bond should have more covalent character, leading
to increase the degree of stability of complexes.
(I) The protonation reaction of the thiosemicarbazones is exothermic
with a net negative ΔG (Table 5).
A good linear correlation has been obtained between the stability
constants of metal complexes and the second ionization potential of
the metal ions (Fig. 3) under investigation. Generally, it is observed
(II) Often, the color of the solution after formation of the complex dif-
fers from the color associated with the free ligand at the same pH.
(III) Formation constants of metal complexes at different tempera-
tures have been calculated and discussed as follows:
[1] These values decrease with rising of temperatures, proposing that
the process of complex formation is favored at low temperature.
2
+
that the stability constant of the Cu complex is larger in comparison
to the other metals. The ligand field had given Cu2 further stability
as a result of tetragonal distortion of the octahedral symmetry [58].
+
Fig. 3. Variation of the stability constants for the M(II)-TPHP complexes with properties of the metal(II) ions.

A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
919
Table 3
the 3d series. Under the influence of both the polarizing ability
of the metal ion [63] and the ligand field [64], Cu will receive
Stepwise stability constants for ML and ML
ligands in 50% (v/v) DMSO-water mixtures and 0.1 mol·dm
temperatures.
2
complexes of TPHP-thiosemicarbazone
2+
at different[
-
−
3
NaNO
3
2
-
some extra stabilization due to tetragonal distortion of octahedral
symmetry in its complexes. The greater stability of Cu2 com-
plexes is produced by the well-known Jahn-Teller effect [64].
[5] All negative values of the Gibb's free energy associated with the
process of formation of complexes illustrate the spontaneous na-
ture of the complex formation reactions.
+
Mn+
15 °C
logK
25 °C
logK
35 °C
1
logK
2
1
logK
2
logK
1
logK
2
2
+
Mn
9.87
10.18
11.08
11.84
9.31
9.77
10.12
10.78
9.72
10.02
10.94
11.70
9.21
9.64
9.60
9.90
9.05
9.50
2
+
Co
2
+
Ni
9.99
10.62
10.79
11.53
9.84
10.49
[6] The negative values of the heat content (ΔH) showed that the
process of formation of complexes is exothermic demonstrating
that the process of chelation is preferable at low temperatures.
2
+
Cu
[
7] The values (ΔS) for the ligand complexes are positive, confirming
Table 4
that the complex formation is entropically favourable [65] and the
Stepwise stability constants for ML and ML
2
complexes of MTPHP-thiosemicarbazone
+
mechanism of complexation is based on hydrogen ion (H ) liber-
ligands in 50% (v/v) DMSO-water mixtures and 0.1 mol.dm− NaNO
3
at different
3
ation and the release of water molecules [66]. During formation of
metal chelates, water molecules from the primary hydration
sphere of the metal ion are displaced by the chelating ligand.
Thus there is an increase in the number of particles in the system
i.e., randomness of the system increases as shown in the following
equation
temperatures.
Mn+
15 °C
logK
25 °C
logK
35 °C
logK
1
logK
2
logK
1
2
1
logK
2
2
+
Mn
10.05
10.59
11.21
12.30
9.40
9.62
10.19
11.14
9.91
10.47
11.03
12.14
9.29
9.49
10.09
11.04
9.77
10.30
10.91
11.98
9.13
9.35
9.91
2
+
Co
2
+
Ni
2
+
ꢀ
ꢁþ2
þ L−ðaqÞ⇌MLþðaqÞ þ nH2O:
Cu
10.84
MðH2OÞn
ð5Þ
ðaqÞ
]
It is known that divalent metal ions exist in solution in the form of oc-
tahedral hydrated species [59]. So it can be considered that the values
of entropy and enthalpy, which was obtained as the sum of the contri-
butions of both liberation of water molecules and formation of bonds
between the metal and ligand. From these results the following conclu-
sions can be derived:
[
8] The negative values of both ΔG and ΔH of the complexation
process indicate that the complexation process proceeds sponta-
neously and exothermically, respectively, i.e., the complex for-
mation is enthalpically favourable.
In general, the abnormal higher positive values of ΔS for all the
complex systems are consistent with the hypothesis that a large
number of water molecules are released upon complexation
with probability of change for the coordination number [67].
This was supported by the values of ΔH, where it was found
[
3] Of the data (Tables 3 and 4) reveals that logK
1
for TPHP and
·This is due to
MTPHP complexes is somewhat larger than logK
2
the fact that the interaction of a second bulky ligand molecule is
usually weaker than the first ligand, i.e., the ML (1:2) species is
2
1 2
that –ΔH N −ΔH for the TPHP and PTPHP-thiosemicarbazone
not formed until complete formation of the ML(1:1) species.
This can be ascribed to: (i) The increase in the Lewis acidity of
the free metal ion (M ) as compared to the 1:1 chelated ion
complexes (Tables 6 and 7). The lower negative values of ΔH
2
were taken as good evidence for a change in the dentate character
of the TPHP and PTPHP ligands from tridentate (SNN-donors) in
the 1:1 complexes to bidentate (SN-donors) in the 1:2, M:L,
complexes, whereby the steric hindrance in the 1:2 species is
relieved. Similar observations were obtained by Evans et al. [68].
The higher values of the thermodynamic functions for the
+
n
+
n–1
(
ML
of a second bulky ligand molecule on the ML
4] For the same ligand at constant temperature, the stability of the
) and (ii) The steric hindrance caused by the addition
+
n–1
chelated ion.
[
chelates increases in the order Cu2 N Ni
60–62]. This order largely reflects that the stability of Cu
complexes is considerably larger than those of other metals of
+
2+
N Co
2+
N Mn
2+
2
+
[
9
Cu(II)-complexes is attributed to the 3d -configuration of
Cu(II) which undergoes Jahn-Teller distortion. On the other
Fig. 4. Effect of temperature on the protonation constant of TPHP and MTPHP .

920
A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
Fig. 5. Effect of temperature on the formation constant of M-TPHP complexes.
Table 5
Thermodynamics for the association of TPHP, MTPHP and PTPHP thiosemicarbazones in
Table 7
Thermodynamic functions for ML and ML
2
complexes of MTPHP, thiosemicarbazones in
2
50% DMSO-50% H O(v/v).
−3
50% (by volume) DMSO-water mixture and 0.1 mol·dm NaNO .
3
System
ΔH°
kJmol−1)
ΔS°
ΔG°
Mn+
Enthalpy/kJ·mol−1
Entropy/J·mol−·K−1
T/K
Gibbs
energy/kJ·mol
(JK mol−1
−1
)
(kJmol−1)
(
−1
TPHP
L
HL + H ⇌ H
ΔG
Mn2+ 288 −55.45 −51.86 −46.05
98 −55.62 −52.09
1
ΔG
2
ΔH
1
ΔH
2
ΔS
1
ΔS
2
−
+ H+ ⇌ LH
−47.60
−34.61
48.82
24.12
−61.97
−41.67
+
L+
−44.34
−44.40
−46.04
−49.38
32.55
37.53
43.08
52.85
26.04
30.06
35.42
45.57
2
2
MTPHP
308 −55.77 −52.12
288 −58.43 −56.22 −47.65
298 −58.77 −56.42
−
+
Co2+
L
+ H ⇌ LH
−55.83
−39.48
22.21
27.57
−62.34
−47.55
+
L+
HL + H ⇌ H
2
3
08 −58.80 −56.57
288 −61.85 −61.46 −49.40
98 −61.96 −61.97
Ni2+
2
hand, the lower values for the Mn(II)-complexes may be
attributed to the 3d – configuration of Mn(II) which probably
exists as tetrahedral [Mn(OH
5
308 −62.28 −61.88
288 −67.86 −62.57 −52.63
Cu2+
2
+
2
)
4
]
in aqueous solutions [69].
2
3
98 −68.14 −62.70
08 −68.39 −63.02
3
.5. Effect of ionic strength on the protonation constants of
where m
i
is the molar concentrations of the ions and Z
i
are their charges.
thiosemicarbazone compounds
This is related to the activity coefficients of the ions in solution by the
following Eq. (7)
The ionic strength of the medium (μ) is given by Eq. (6)
2
logγꢁ ¼ −αZþjZ−jμ 1=2
ð7Þ
μ ¼ 1=2ΣZ m
ð6Þ
i
i
where γ
of dissociation, and Z
When the concentration of free ions in solution is very low, γ
close to unity and the ions are not affected by each other. At higher
concentrations, γ is not close to unity and the ions of opposite charges
±
is the activity coefficient of cations and anions, α is the degree
|Z | is the positive product of the ionic charges.
becomes
Table 6
Thermodynamic functions for ML and ML
+
−
2
complexes of TPHP, thiosemicarbazones in 50%
−
3
(by volume) DMSO-water mixture and 0.1 mol·dm NaNO
3
.
±
Mn+
Enthalpy/kJ·mol−1
Entropy/J·mol−1·K−1
T/K Gibbs
energy/kJ·mol−1
±
attract each other [70]. Therefore, the effect of variation of ionic strength
ΔG
1
ΔG
2
ΔH
1
ΔH
2
ΔS
1
ΔS
2
Mn2+ 288 −54.46 −51.37 −44.48
−42.77
−44.42
−46.04
−47.73
34.46
34.84
46.52
49.97
30.06
32.93
34.08
40.59
2
3
98 −54.50 −51.53
08 −54.80 −51.60
Table 8
Co2+
Proton-ligand association constants of TPHP at different ionic strengths at 25 °C.
288 −56.17 −53.90 −46.07
2
3
98 −56.24 −54.11
08 −56.52 −54.23
log KNHa
System
Ionic strength
log KSH
Ni2+
288 −61.13 −55.83 −47.69
TPHP
TPHP
TPHP
TPHP
TPHP
0.05
0.10
0.15
0.20
0.25
11.21 ± 0.05
11.07 ± 0.01
10.97 ± 0.06
10.88 ± 0.01
10.79 ± 0.03
7.54 ± 0.08
7.41 ± 0.04
7.29 ± 0.06
7.18 ± 0.05
7.09 ± 0.08
2
3
98 −61.41 −56.07
08 −61.60 −56.17
Cu2+
288 −65.32 −59.48 −50.96
2
3
98 −65.67 −59.61
08 −65.82 −59.88
a
Protonation constants ± standard deviations.

A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
921
pffiffiffi
Fig. 6. Plot of log K vs μ at 25 °C (the values of correlation coefficients are 0.991–0.993).
on the protonation constants of the TPHP as a representative example of
References
thiosemicarbazone compounds has been investigated. For this purpose,
the proton ligand protonation constants of the TPHP ligand have been
evaluated at five different ionic strengths (0.05, 0.10, 0.15, 0.20 and
[
[
1] M. Meloun, S. Bordovská, A. Vrána, Anal. Chim. Acta 584 (2) (2007) 419–432.
2] G. Roda, C. Dallanoce, G. Grazioso, V. Liberti, M. deAmici, Anal. Sci. 26 (1) (2010)
51–54.
3] M. Sanchooli, J. Chemistry vol.2013 (2013), 989362 8 pp.
4] A.E. Pütün, G. Bereket, E. Keskin, J. Chem. Eng. Data 40 (1) (1995) 221–224.
5] I. Narin, S. Sarioglan, B. Anilanmert, H. Sari, J. Solut. Chem. 39 (10) (2010)
1582–1588.
[
[
[
0
.25 M) using sodium nitrate as a supporting electrolyte at constant
temperature (298 K). It was observed that the protonation constants
of the TPHP thiosemicarbazone decrease with increase in the ionic
strength of the medium (Table 8) which is in good agreement with
the Debye-Hückel equation [71]. logKProtonation values were plotted vs.
[
6] S.S¸. Anli, Y. Altun, N. Sanli, G. Alsancak, J.L. Beltran, J. Chemical and Engineering Data,
54 (11) (2009) 3014–3021.
[
7] O. Hakli, K. Ertekin, M.S. Ozer, S. Aycan, J. Anal. Chem. 63 (11) (2008) 1051–1056.
square root of μ as per the Debye-Huckel equation (Fig. 6). The plots
[8] H. Sigel, R.B. Martin, Chem. Rev. 82 (1982) 385–426.
9] C. Öğretir, S. Yarligan, S. Demirayak, T. Arslan, J. Mol. Struct. (THEOCHEM) 666-667
2003) 609–615.
pffiffiffi
[
of log K vs μ for all systems were found to be linear with correlation
(
coefficient ranges from 0.98–0.99.
[
[
[
10] F.C. Rossotti, H. Rossotti, The Determination of Stability Constants, McGraw-Hill Inc.,
New York, 1961.
11] A.A. El-Sherif, M.R. Shehata, M.M. Shoukry, N. Mahmoud, J. Solution Chem. 45
4
. Conclusions
(2016) 410–430.
12] A.E. Martell, M. Calvin, Chemistry of Metal Chelate Compounds, Prentice-Hall Inc.,
Complex formation equilibria of M(II), with thiosemicarbazones
New York, 1952.
were investigated in 50% DMSO-H
studies indicated that substitution of hydrogen in thiosemicarbazone
TPHP) by methyl group causes lowering of acidity as a result of the
2
O mixture. Chemical equilibrium
[13] K.J. Duffy, A.N. Shaw, E. Delorme, S.B. Dillon, C. Erickson-Miller, L. Giampa, Y. Huang,
R.M. Keenan, P. Lamb, N. Liu, S.G. Miller, A.T. Price, J. Rosen, H. Smith, K.J. Wiggall, L.
Zhang, J.I. Luengo, J. Med. Chem. 45 (2002) 3573.
(
[
14] J.R. Dilworth, R. Hueting, Inorg. Chim. Acta 389 (2012) 3.
electron-releasing effect of the methyl group in the substituted amino
group. The protonation constants of the TPHP thiosemicarbazone com-
pound decrease with increase in the ionic strength of the medium.
M(II)-MTPHP complexes is more stable than M(II)-TPHP complexes as
a result of electron-releasing effect of N-methyl group complexes. It is
hoped that the obtained data will be a significant contribution to
workers carrying out mechanistic studies in biological media. For the
[15] J.A. Lessa, I.C. Mendes, P.R.O. da Silva, M.A. Soares, R.G. dos Santos, N.L. Spezialic, N.C.
Romeiro, E.J. Barreiro, H. Beraldo, Eur. J. Med. Chem. 45 (2010) 5671.
[
16] A. Karakucuk-Iyidogăn, D. Tasdemir, E.E. Oruc-Emre, J. Balzarini, Eur. J. Med. Chem.
46 (2011) 5616.
[17] R.J. Glisoni, M.L. Cuestas, V.L. Mathet, J.R. Oubina, A.G. Moglioni, A. Sosnik, Eur. J.
Pharm. Sci. 47 (2012) 596.
[
18] X. Du, C. Guo, E. Hansel, P.S. Doyle, C.R. Caffrey, T.P. Holler, J.H. McKerrow, F.E. Cohen,
J. Med. Chem. 45 (2002) 2695.
[19] D. Kovala-Demertzi, M.A. Demertzis, E. Filiou, A.A. Pantazaki, P.N. Yadav, J.R. Miller,
Y. Zheng, D.A. Kyriakidis, Biometals 16 (2003) 411.
1 2
TPHP and MTPHP complexes, log K N log K indicating that the vacant
[
[
20] J.P. Scovill, D.L. Klayman, D.G. Franchino, J. Med. Chem. 25 (1982) 1261.
21] L. Klayman, J.P. Scovill, J.F. Bartosevich, J. Bruce, J. Med. Chem. 26 (1983) 35.
sites of the metal ions are more freely available for binding of the first
ligand than for the second one. The complexation (chelation) process
is spontaneous, exothermic and entropically favourable. More stable
complexes will be formed with: i) Hard-hard or soft-soft interactions
of the metal ions and the ligands; ii) Small sizes of the metal ions; iii)
Higher basicity of the ligands and iv) Lower temperatures which is con-
sistent with exothermic complexation.
[22] D.K. Demertzi, M.A. Demertzis, J.R. Miller, C. Papadopoulou, C. Dodorou, G. Filousis, J.
Inorg. Biochem. 86 (2001) 555.
[
[
23] P.K. Singh, D.N. Kumar, Spectrochim. Acta A 64 (2006) 853.
24] A.A. El-Sherif, J. Solut. Chem. 41 (2010) 249–260.
[25] A.A. El-Sherif, J. Coord. Chem. 64 (12) (2011) 2035–2055.
[25] A.A. El-Sherif, J. Coord. Chem. 264 (7) (2011) 1240–1253.
[
[
[
26] A.A. El-Sherif, J. Solut. Chem. 39 (2010) 131–150.
27] A.A. El-Sherif, M.M. Shoukry, J. Main Group Metal Chem. 29 (4) (2006) 189–200.
28] A.A. El-Sherif, M.M. Shoukry, Spectrochim. Acta A 66 (2007) 691–700.
Abbreviations
[29] A.A. El-Sherif, M.M. Shoukry, J. Coord. Chem. 59 (14) (2006) 1541–1556.
[
[
30] A.A. El-Sherif, M.M. Shoukry, J. Coord. Chem. 58 (16) (2005) 1401–1415.
31] N.A. Al-Awadi, N.M. Shuaib, A.A. El-Sherif, A. El-Dissouky, E. Al-Saleh, Bioinorg.
Chem. Appl., 2008, Art. No. 479897.
PHP
TPHP
1-(phenyl-hydrazono)-propan-2-one
2-(1-(2-phenyl-hydrazono)-propan-2-ylidene) hydrazine-
carbothioamide
[32] A.A. El-Sherif, J. Solut. Chem. 39 (2010) 1562–1581.
[
33] A.A. El-Sherif, M.M. Shoukry, J. Progress in Reaction Kinetics and Mechanism 36
2011) 215–226.
34] N. Rabjohn, Organic Synthesis, Collective Volume 4, John Wiley and Sons Inc., 1963
(
MTPHP N-methyl-2-(1-(2-phenyl-hydrazono)-propan-2-
[
ylidene)hydrazinecarbo-thioamide
[35] A.A. El-Sherif, Inorg. Chim. Acta 362 (2009) 4991–5000.

9
22
A.T.A. El-Karim, A.A. El-Sherif / Journal of Molecular Liquids 219 (2016) 914–922
[54] M. Aljahdali, A.A. El-Sherif, M.M. Shoukry, S.E. Mohamed, J. Solut. Chem. 42 (2013)
[36] A.A. El-Sherif, J. Solut. Chem. 39 (2010) 131–151.
[37] R.G. Bates, Determination of pH, Theory and Practice, 2nd edn. John Wiley and Sons,
1028–1050.
New York, 1975.
[55] M.S. Aljahdali, A.T. Abdelkarim, A.A. El-Sherif, J. Solut. Chem. 42 (2013) 2240–2266.
[56] J.E. Huheey, Inorganic Chemistry-Principles of Structure and Reactivity, Harper SI
Edn, New York, 1983.
[
38] F.J. Welcher, The Analytical Uses of Ethylenediamine Tetraacetic Acid, Princeton, Van
Nostand, 1965.
[
[
[
[
39] M. Woolley, D.G. Hurkot, L.G. Hepler, J. Phys. Chem. 74 (1970) 3908–3913.
40] G.L. Van Uitert, C.G. Hass, J American Chemical Society 75 (1971) 451.
41] E.P. Serjeant, Potentiometry and Potentiometric Titrations, Wiley, New York, 1984.
42] A. Golcu, M. Tumer, H. Demirelli, R.A. Wheatley, Inorg. Chim. Acta 358 (2005)
[57] H.M. Irving, H.S. Rossotti, J. Chem. Soc. 2904 (1954).
[58] F.A. Cotton, G. Wilkinson, Advanced Inorganic Chemistry, Wiley, London, 1962.
[59] H.M.N.H. Irving, R.J.P. Williams, The stability of transition-metal complexes, J. Chem.
Soc. (1952) 3192–3210.
1
785–1797.
[60] A.A. El-Sherif, M.R. Shehata, M.M. Shoukry, M.H. Barakat, Bioinorg. Chem. Appl.
Volume 2012, Article ID 984291, 10 pp. http://dx.doi.org/10.1155/2012/984291.
[61] W.U. Malik, G.D. Tuli, R.D. Madan, Selected Topics in Inorganic Chemistry, third ed.
S. Chand & Company LTD, New Delhi, 1984.
[
[
43] A.E. Martell, R.J. Motekaitis, The Determination and Use of Stability Constants, VCH,
Weinheim, 1988.
44] M. Meloun, J. Havel, J.H. Högfelt, Computation of Solution Equilibria, Wiley, N. Y,
1
988.
[62] G.A. Ibañez, G.M. Escandar, Polyhedron 17 (25–26) (1998) 4433–4441.
[63] F.R. Harlly, R.M. Burgess, R.M. Alcock, Solution Equilibria, Ellis Harwood, Chichester,
UK, 1980.
[
[
[
45] P. Gans, A. Sabatini, A. Vacca, Inorg. Chim. Acta 18 (1976) 237–239.
46] L. Pettit, University of Leeds Personal Communication.
47] A.A. El-Sherif, M.M. Shoukry, M.M.A. Abd-Elgawad, J. Solut. Chem. 42 (2013)
[64] L.E. Orgel, An Introduction to Transition Metal Chemistry Ligand Field Theory,
Methuen, London, 1966.
[65] A.A. El-Sherif, T.M.A. Eldebss, J. Spectrochimica Acta (A) 79 (2011) 1803–1814.
[66] B. Jeragh, D. El-Wahaib, A.A. El-Sherif, A. El-Dissouky, J. Chem. Eng. Data 52 (5)
(2007) 1609–1614.
[67] A. McAuley, G.H. Nancollas, K. Torrance, Inorg. Chem. 6 (1967) 136.
[68] A.G. Evans, J.C. Evans, B.A. El-Shetary, C.C. Rowlands, P.H. Morgan, J. Coord. Chem. 19
(1979) 19–29.
4
12–427.
48] A.A. El-Sherif, M.R. Shehata, M.M. Shoukry, M.H. Barakat, Spectrochimica. Acta (A)
6 (2012) 889–897.
49] D. Martin, H.G. Hauthal, Dimethylsulphoxide, Van Nostrand Reinhold, Wokingham,
975.
[
[
[
[
[
[
9
1
50] E. Labisbal, K.D. Haslow, A. Sousa-Pedrares, J. Valdes-Martinez, S. Hernandez-Ortega,
D.X. West, Polyhedron 22 (2002) 2831–2843.
51] B. Kurzak, A. Kamecka, K. Kurzak, J. Jezierska, P. Kafarski, Polyhedron 19 (2000)
[69] P. George, D.S. McClure, Progress in Inorganic Chemistry, Vol. IInterscience,
2083.
NewYork, 1959.
52] M. Aljahdali, C. Foti, A.A. El-Sherif, M.M.A. Mohamed, A.A. Soliman, O.S. Al Ruqi,
Monatsh. Chem. 144 (2013) (2013) 1467–1480.
53] É.A. Enyedy, M.F. Primik, C.R. Kowol, V.B. Arion, T. Kiss, B.K. Keppler, Dalton Trans. 40
[70] K. Denbigh, Principles of Chemical Equilibrium, Cambridge Univ. Press, London,
1955.
[71] J. Bjerrum, Metal-Amine Formation in Aqueous Solution, Haase, Copenhagen, 1941.
(22) (2011) 5895–5905.