Download Potintiometric and Thermodynamic Studies of Some Divalent Metal

Potintiometric and Thermodynamic Studies of Some Divalent Metal Ion
Complexes with a New Schiff Base Hydrazone Containing Quinoline Moiety
Noura H. AL-Sha'alan1, Mostafa El-Behery2* and Gaozaa El-Nawmosy1
ABSTRACT
Metal complexes of some divalent metal ions (Co, Ni, Cu, Mn, Cd, and Zn) with a new
hydrazone Schiff base consisting of 7-chloro-4-(o-hydroxybenzilidenehydrazo)-quinoline,
were investigated potentiometrically and spectrophotometrically and found to have
stoichiometric formulae 1:1 and 1:2 (M:L). Formation constants of the proton-ligand and
metal-ligand complexes were determined potentiometrically at different temperatures (10, 20,
30 and 40 oC) and 0.05 mol dm-3, ionic strength (KNO3) in 75 % (v/v) ethanol-water mixture.
The standard thermodynamic parameters viz. G o , H o , and S o , for the proton-ligand and
the stepwise metal-ligand complexes were evaluated. Thermodynamic functions were
analyzed in terms of the electrostatic (el) and non-electrostatic (cratic, c) components. H co
was found to be linearly correlated with the acceptor number of the metal ion ( ANM ),
whereas H elo was linearly correlated with the ionic radii of the metal ion. Electronic, IR, and
1
H NMR spectra were used to characterize the free ligand, HL. Formation constants, and
analytical applications of the ligand complexes with Co(II), and Ni(II) ions were also
evaluated and studied. The constant values obtained from spectrophotometric studies are
found to be consistent with those obtained from potentiometric studies. The study was
performed also for Fe(III)-complex.
Keywords:
dissociation; stability constants; thermodynamic parameters; Schiff base
hydrazone
1- INTRODUCTION
*
Corresponding author, e-mail: [email protected]
Numerous hydrazone Schiff base compounds containing quinoline moiety are well known in
natural materials and show interesting biological and antiviral activities1-3. Many derivatives
of hydrazon compounds form coloured complexes with different metal ions and can be used
as analytical reagents for their determinations4-9. The coordination compounds of
aroylhydrazones have been reported to act as enzyme inhibitors10 and are useful due to their
pharmacological applications11-13. Hydrazones derived from condensation of isonicotinic acid
hydrazide (INH) with pyridine aldehydes have been found to show better antitubercular
activity than INH14,15. It is believed that ligands containg different donor atoms like O – N
type form more stable complexes than either O – O or N – N types16. We have previously
investigated and reported the physicochemical properties of the solid complexes of the Schiff
base hydrazone ligand, 7-chloro-4-(o-hydroxybenzilidenehydrazo)-quinoline, which contains
an O –N ligand system and forms stable solid coloured complexes with transition metal
ions17.
This aim of this work is to study the complex formation of such a Schiff base hydrazone
ligand with Co, Ni, Cu, Mn, Cd, Zn and Fe(III) in an ethanol–water mixture using
potentiometric and spectroscopic methods to throw light on their composition, structures and
analytical applications. Thermodynamics parameters are computed and analyzed in order to
investigate the bond character between the metal ligand.
2. EXPERIMENTAL
2.1.
Reagents and Materials
All chemicals used were of Anal. Grade and the solutions of metal nitrates of Co(II), Ni(II),
Cu(II), Mn(II), Cd(II), Zn(II) and Fe(III) ions were prepared in carbonate free double distilled
water and they were standardized using EDTA titrations18. The HL ligand was prepared as
described previously17. The structure of the ligand was elucidated by IR, Mass, 1H-NMR and
electronic (UV-VIS) spectroscopy, as well as micro analytical analyses17. The results have
been previously published17and weighed quantity of the ligand was dissolved in 75% (v/v)
ethanol-water medium. Ethanol was freshly refluxed and distilled over magnesium powder
and iodine18.
2.2.
Potentiometric measurements
2
A Hanna pH-meter Model 302 digital with conventional pH-electrode assembly was used for
pH measurements at 10, 20, 30 and 40oC. The ionic strength of the medium was kept virtually
constant at 0.05 mol dm-3 with KNO3 as background electrolyte. The temperature was
maintained constant by use of double-jacket cells with water circulated from a constanttemperature bath. Purified nitrogen gas was bubbled through the solution before and during
the titrations. Multiple titrations were carried out for each system.. The pH meter readings
were recorded6-9. Electronic absorption spectra of the free ligand and its complexes in 75 %
(v/v) ethanol–water mixture were recorded using a Shimadzu Model Uv-probe
spectrophotometer. The compositions of the complexes in solutions were determined by the
mole ratio method19,20. The pH-metric titrations in 75 % (v/v) ethanol–water of the free ligand
and its mixtures with the metal ions were carried out as described previously6-9. Using
equations of Irving and Rossoti21 the curves were used to calculate the formation constants of
the following equilibria.
M2+ + HL
and
ML+ + HL
ML+ + H+
ML2 + H+
The experimental reading were used to calculate the values of n and pL . From which the
stability constant log K1 and log K 2 were calculated with the help of the following equations
given by Irving and Rossotti(21).
n
 log K 1  pL
1 n
n 1
log
 log K 2  pL
2n
log
The metal-ligand constant was obtained using linear plot method by plotting
log
n 1
2n
log
n
1 n
or
against pL. Where n is defined by Irving and Rossitti21 as the average number of the
reagent molecules attached per metal ions. It can be obtained from the following relations
n
(V3  V2) ( N 0  E 0 )
(V 0  V2 )n ATC M0
and
n 1


1

 0
 anti log pH  . V  V3
TC L0  n TC M0
V0
 n
pL  log 10
n 0
H
3
where V2 and V3 are the volumes of alkali required to reach the same pH value, Vo is the total
volume of titrating mixture, No and Eo are the concentration of the alkali and the initial
concentration of free acid, respectively, TC0L and TC0M are the total concentration of ligand
and metal ion, respectively and pL, is the free ligand exponent. The experimental reading
were used to calculate the values of n and pL . From which the stability constant log K1 and
log K 2 were calculated with the help of the following equations given by Irving and
Rossotti(21).
The variations of pK H , log K1 , and log K2 vs. 1/ T gave straight lines which enable us to
calculate the thermodynamic parameters G o , H o , and S o of the complex formation by
the Vant Hoff equation and other relationships22. Conductemetric titrations were carried out at
room temperature by titrating 30 ml metal ion solution (1x10-3 mol dm-3) with the ligand
solution (1x10-2 mol dm-3) as titrant using WTWD-812 Weilneium-conductivity meter, model
LBR, fitted with a cell model LTA 100.
2.2.
Preparation and characterization of the organic ligand
The structure of the ligand was elucidated by IR, Mass, 1H-NMR and electronic (UV-VIS)
spectroscopy, as well as micro analytical analyses17. The mass spectra showed molecular ion
peaks at m/z: 297 corresponding to the molecular weight of the HL ligand. Also, the base
peaks were observed at m/z 178 and no meta-stable ions were observed.
The electronic absorption of HL in 75 % (v/v) ethanol-water mixture within the region 200600 nm displayed absorption bands at 208 and 224 nm assigned for  -* transitions within
the aromatic and quinoline rings, respectively. The band observed at 276 nm would be due to
the n-* transition of the C=N group. The absorption bands at 382 and 404 assigned for CT
transitions. The band at 404 nm encroaches on the visible region21 and impact the ligand its
colour.
The IR spectrum of the ligand shows a strong band at 3530 cm-1 assigned to OH of the
phenolic group. The stretching vibration of NH group appears as weak bands at 3300 cm-1.
The spectrum shows also vibrational bands at 1605, 1600 and 1140, 1130 cm-1 which are
assigned for s and as of the C=N and N-N groups, respectively. The deformation vibration,
4
, of the phenolic OH group appears at 1278 cm-1. The vibrational band observed at 2570 cm-1
is assigned to the intramolecular hydrogen bonding17.
Finally, the chemical shift (δppm) data extracted from the 1H-NMR spectrum of the ligand in
DMSO-d6 are summarized in a previous study17. These data together with the data derived
from the elemental analyses, IR, mass and electronic spectra confirmed the structure given for
these ligand17.
3. RESULTS AND DISCUSSION
3.1.
Dissociation and Formation Constants
The potentiometric titration curves of HL in presence or absence of the studied metal ions are
shown in Fig. 1. The values of the dissociation constant of the free ligand, pK H , was
determined by the Albert–Serjeant method23 while the formation constants of its metal
complexes were determined by the Irving-Rossotti method21 Only one proton dissociates
between a = 0.0 and a = 1.0 (a = number of moles of base / mole of ligand), suggesting that it
behaves as mono-protic species (HL). The values of the pK H , formation constants, as well as
the thermodynamic parameters in 75 % (v/v) ethanol-water mixture are given in Table 1.
5
Table 1. Thermodynamics functions and stepwise formation constants of 1:1 and 1:2 metal complexes at 10, 20, 30 and 40 oC in 75 %
ethanol-water mixture
t oC
Complex
(HL)
10
K. cal. mol-1
20
30
-DGo30
40
cal. mol-1 K-1
-DHo30
-DSo30
logk1
logk2
logk1
logk2
logk1
logk2
logk1
logk2
-DGo1
-DGo2
-DHo1
-DHo2


10.14
10.58
10.09
10.56
9.79
10.4
9.55
10.2
58.75
61.49
50.76
33.76
192.24
202.29
Cd(II)
9.92
10.71
9.88
10.69
9.69
10.6
9.52
10.1
57.53
64.00
33.82
83.45
188.77
210.53
Ni(II)
10.23
10.59
10.18
10.57
9.8
10.32
9.55
10.11
59.28
61.55
59.28
43.23
193.69
202.48
Co(II)
Zn(II)
9.42
9.66
9.21
9.69
8.93
9.31
8.61
9
53.63
56.42
56.29
64.8
175.15
185.62
Fe(III)
7.52
8.46
7.41
8.4
6.9
8.1
6.5
7.9
43.15
48.91
85.55
47.04
139.59
160.91
Mn(II)
10.10
10.65
9.91
10.61
7.96
10.47
6.7
10.3
57.71
61.78
30.21
29.07
180.49
203.24
pka
10.78
-
10.75
--
10.54
--
10.37
--
62.60
-
35.71
--
205.45
--
6
The pK H value suggests that the proton dissociates from the OH group rather than
from the NH proton of the HL which may dissociate at higher pH values. However,
similar conclusion for similar ligands was obtained24,25. The pK H value obtained for
HL, in 75 % (v/v) ethanol-water at 30 oC is close to that reported in the literature13 for
similar ligand under similar experimental conditions. The order of basicity, pK sH ,
(Table 2), in 75 % (v/v) different mixed solvents (ethanol- water, acetone–water and
isopropanol–water) is as follows; acetone-water > isopropanol-water > ethanol-water
in contrast with the order of the dielectric constants of the organic solvents which is,
ethanol > acetone > isopropanol. Gregely26 indicated that some solvent molecules
(dioxan) progressively breakdown the hydrogen bonded structure of water molecules
whereas other can form hydrogen bonded (methanol) association with water. Thus, it
is expected that the extent of hydrogen bonding in alcohol-water is greater than
acetone-water. This may affect the measured pK sH values. The pH-metric titration
curves (Fig. 1) in the case of Mn(II) showed a sharp inflection at m = 1.0 (m =
number of moles of KOH added / mole of metal ion) corresponding to the formation
of the 1:1 MnL+ complex. This could be represented by the following equilibrium:
2+
Mn
+
HL
+
-
+
OH
MnL
+
H2O (at m = 1.0)
Such feature does not observed with other metal cations, where there was only one
overlapping buffer region between m = 0.0 to 2.0 corresponding to the formation of
bis-chelated 1M : 2L. This could be represented by the following equilibrium.
M
n+
Table 2.
+
2HL
+n-2
-
+
OH
ML2
+
2H2O (at m = 2.0)
pK sH values of the HL ligand in 75% (v/v) solvent-water and the
Dielectric constants (D) of the solvents at 30 oC
Solvent-water at 30 oC
ethanol acetone
Isopropanol
D
pK
H
24.3
20.7
18.3
10.53
10.99
10.79
7
It should be mentioned that calculation of formation constant log K1 for Cu(II)complex was not succeeded due to precipitation of the metal hydroxide at the
beginning of the titration, and since the maximum n < 1.0.
The lower stability of the Fe(III)-complex, (Table 1), could be explained by the
interaction of two ligand molecule to form [Fe(L2)]+ hindered the introduction of a
third one which may lead to lower the stability. However log K 3 was not measured.
The variation of log K1 vs. pK H is shown in Fig. 2, where a linear relation was
obtained. This linearity reflects the similarity of the ionic nature of the metal-ligand
interaction27. The data in Table 1 shows that log K 2 for the complexes is somewhat
larger than log K1 . This indicates a strong trans-effect for the second coordination28,29.
This trans-effect was also supported by the given thermodynamic where it was found
that H1>H2 which is related to the bond strengths in these complexes. The order of
stability constants, log K1 , at 30 oC for HL- complexes is as follow
Ni(II) > Co(II) > Cd(II) > Zn(II) >Mn(II) > Fe(III)
3.2.
Effect of Temperature
The dissociation constants of the ligand pK H , as well as the stability constants of its
metal complexes with Co(II), Ni(II), Mn(II), Cd(II), Zn(II) and Fe(III) ions in 75 %
(v/v) ethanol-water were evaluated at 10, 20, 30, and 40 oC. The slope of the plot
( pK H or log K vs. 1/ T ) was utilized to calculate the enthalpy change (H) for the
dissociation or complexation process, respectively, and given in Table 1. From the
free energy change (G) and the enthalpy change (H), the entropy change (S) can
be calculated using the well-known relationships:
d log K / dt  H o / 4.57T
 G  2.303RT log K
G o  H o  TS o
8
Tables 1 reveals that: (i) The pK H value of the free ligand decrease with increasing
temperature indicating that the acidity of the ligand increases with increasing
temperature, i.e., a higher temperature favours the dissociation or ionization process.
(ii) On the basis of (i), a positive value of H would be expected indicating that the
dissociation process is accompanied by the absorption of heat and it is endothermic
process. (iii) The large positive value of G indicates that the dissociation process is
non spontaneous. (iv) The large negative values of S of the dissociation process
indicate that the ionization of the ligand is entropically unfavourable. (v) The large
positive values of S of the complexation process indicate that the complex formation
is entropically favourable and the mechanism of complexation is based on hydrogen
ion (H+) liberation and the release of water28 (vi) The negative values of both G and
H of the complexation process indicate that the complexation process proceeds
spontaneously and exothermically, i.e., the complex formation is enthalpically
favourable. A Similar observation have been reported by Abdel-Moez et al30.
In general, the positive values of S for all the complex systems are consistent with
the hypothesis that a large number of water molecules are released upon complexation
with the probability of a change of the coordination number28. This was supported by
the values of H, where it was found that -H1>-H2 for the complexes, except for
Zn(II) and Cd(II)-complexes. The lower negative values of -H2 were taken as an
evidence for a change in the dentate characters from tridentate, (ONN-donors in the
1:1 complexes to bidentate (ON-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 al31. For Zn(II) and Cd(II)-complexes, where -H1 < -H2, This can be
considered as an evidence for the trans-effect28,29 of the second coordination.
It is convenient to analyze any thermodynamic function X o (X = Go; Ho, and So)
into two parts: a) X c0 (critic, non-electrostatic part), a temperature-independent
component intrinsic to the molecules or ions and arising out of short range or covalent
forces insensitive to the environment, b) X el0 , a temperature-dependent part, owing
to the interaction of the dipoles or ions with long-range electrostatic forces of the
solvent medium; then
9
X 0  X el0  X co
and calculated according to the following equations36.
Gnon  nRT ln M  Rca ; Gel  Rc exp (T /  )
H non  Rca ; H el  Rc (1  T /  ) exp (T /  )
S non  nRT ln M ; S el  ( Rc /  ) exp (T /  )
G  nRT ln M  Rc (a  exp (T /  ) )
H  Rc[a  (1  T /  ) exp (T /  ) ]
S  nRT ln M  ( Rc /  ) exp (T /  )
Where  is a temperature characteristic of solvent (=219 K); M is the number of
moles contained in 1000 g of water (=55.5); C and a constants(36). The separation of
thermodynamic parameters into temperature-dependent and temperature-independent
was first suggested by Gurney32 for proton ionization reactions and has been extended
to metal complexes formation by Anderegg33 and Nancollas34. The nature of
temperature dependence X el0 can be approximated from the Born model35. Table 3
summarizes the temperature-dependent and temperature-independent data. which
were calculated as described above36.
According to some authors,35,36 electrostatic (el) or environment components
represent longe-rang electrostatic forces depending on the environment and
temperature; while non-electrostatic (non) or critic components represent short-ramge
or quantum mechanical forces; insensitive to the environment and independent of
o
temperature. Also, H non
arises from changes in the ligand field stabilization (LFS)
accompanying complex formation34.
o
As seen from Table 3, H non
has high negative values while H elo has positive values
o
and, since H non
reflects the covalent nature of the formed complexes, the expected
o
degree of covalency exists in such systems. Also, Gelo >> Gnon
for all complex
10
indicating that the complexation process is highly influenced by the temperature and
environment (medium).
Table 3. Electrostatic (el) and critic (c) thermodynamic parameters for the
reaction of metal ions with HL ligand in 75 % (v/v) ethanol – water
at 30 oC
o
G30
( K.J .mol1)
complex
(HL)
Co(II)
Cd(II)
Ni(II)
Zn(II)
Fe(III)
Mn(II)
 Gelo
log K1
47.86
47.12
48.31
44.24
36.70
45.40
log K 2
49.70
51.30
49.87
46.13
41.01
50.15
o
H30
( K.J .mol1)
o
 Gnon
log K1
10.90
10.46
11.14
9.40
6.50
12.35
log K 2
11.60
12.70
11.70
10.30
7.91
11.65
H elo
o
 H non
log K1
21.05
20.57
21.27
19.55
16.72
22.50
log K 2
21.76
22.86
21.84
20.87
18.07
21.81
o
S30
(.J .mol1K 1)
log K1
20.55
20.23
20.68
18.99
15.75
19.40
Selo
log K 2
log K1
log K 2
21.42
22.02
21.41
19.80
17.60
21.52
225.77
222.44
227.22
208.66
173.11
214.0
235.33
241.96
235.24
217.61
193.41
236.48
The complex formation reactions of most metal ions with HL are found to be highly
exothermic due to a higher stability constant values of these complexes. It is known
that the X c0 values reflect the covalency of the bonding and the structural changes
on complexation36. Regarding the covalency, X c0 increases according to the order
Ni(II), Co(II), Cd(II), Mn(II), Zn(II), Fe(III). This order is in a good agreement with
the softness of the metal ions37. The observed entropy changes, S 0 , are, however,
positive due to the large positive contribution of the temperature-dependent part
S el0 , which reflects the release of bounded water molecules from the hydrated ions
on complex formation. The increase of the value of S 0 results in an increased
endothermicity of H el0 and hence a lower observed exothermic change in H 0 .
The plot of m (electronegativity of metal ion) against the logarithm of the overall
formation constant ( log K1 K 2 ) (Fig. 3) shows a linear relationship with the exception
of the Fe(II) complexes. It may be inferred that the stability of the metal complexes
increases with increasing the m values and consequently the metal-ligand bond
would be more covalent. This interpretation is supported by the linear correlation
between H c vs. m (Fig. 4). On the other hand, the covalency character of these
11
complexes can be concluded from the linear correlation of H c vs. AN M of the
cations (Fig 5). This agrees well with similar conclusions cited by Linert38, Van
Uitert39 and Selbin40.
It is known that a softer metal ion has a greater affinity for a softer donor41. In Fig. 6,
o
the values of H non
are plotted against the quantity E n# , introduced by Klopman42 as a
measure of the softness of a metal ion in solution. A soft metal ion is characterized by
a large negative value of E n# and vice versa. As can seen from Fig. 6, a linear
correlation with a number of scattered points (for the Ni(II), Mn(II) and Cd(II)o
complexes) was obtained. In general, H non
increases with the softness of the metal
ion.
The metal-ligand interactions can be divided into two groups; (i) those of nickel and
cobalt which have high affinities for N-donor atoms and (ii) those of manganese, zinc
and cadmium with intermediate affinities for both N- and O-donor atoms43. This is
consistent with the above linear correlation with some scattered points (Fig. 3).
3.3.
Conductemetric Titration
To obtain meaningful information about the composition of the prepared complexes,
conductometric titrations were performed. The technique has in common the fact that
the measured quantity is directly proportional to the concentration of one or more of
the ions of interest. If the reaction between the metal ions and the complexing titrant
(ligand) is essentially complete, a titration curve is obtained that consists of two or
more straight, or nearly straight, line segments intersecting at the equivalence point
(the stoichiometric ratio) of the complex formed44. The conductogram of HL is shown
in Fig. 7. It is clear that the conductance value of the metal ion solution increases
steadily as the ligand solution is added. The increasing continues until the equivalence
point of the titration. The behavior obtained is attributed to the replacement of some
ions by complex molecules with different mobility values. With the next drop of
added titrant, the ligand remains unreacted and a slight increase is sometimes
observed due to the accumulation of excess ligand solution. The conductogram
12
exhibit two obvious slopes, suggesting that the probable stoichiometric ratios of the
complexes are M:L and M:2L.
3.4.
Spectroscopic Measurements
Stoichiometry
of
the
complexes
under
investigation
was
determined
spectrophotometrically by applying the mole ratio method45. This method was
measured at pH 8.0 and = 420 nm for Co(II) and Ni(II)-HL complexes, in 75 (v/v) %
ethanol-water mixture, µ = 0.05 mol dm-3 (KNO3) using reference ligand a) [Ni(II)] =
6 × 10-5 mol dm-3 and [HL] = 2 × 10-5 - 2 × 10-4 mol dm-3. Analysis of the data at
various wavelengths proved that Co(II), and Ni(II) ions form two main complex
species with stoichiometric ratios of 1 : 1 and 1:2 n(M) : n(L) (Fig. 8).
Under the optimum conditions, using a constant concentration of HL(1×10-3 mol dm3
) and varying Co(II), and Ni(II) concentrations in ethanol–water mixture, it was
found that the system obeys Beer's law within the metal concentration range 1 1×10-52×10-4 mol dm-3 for Co(II) and Ni(II) complexes. The molar absorptivity () values at
500 nm for Co(II), and Ni(II) complexes were 8.0×103 dm3 mol-1 cm.-1, and 1.4×104
dm3 mol-1 cm-1, respectively, indicating that HL can be readily utilized as a sensitive
reagent for micro analytical determination of cobalt(II), and nickel(II) (Fig. 9).
CONCLUSION
(i)
The ligand behaves as monobasic (monoprotic) species (HL) towards the
metal ions as evidenced from the titration curves.
(ii)
The maximum n values were found to be  2, revealing that both ML and
ML2 species are formed in solution.
(iii)
For all complexes, log K2 > log K1 indicating that the vacant sites of the metal
ions are more freely available for the binding of two ligand molecules
(iv)
More stable complexes, will be formed with:
– hard-hard or soft-soft interactions of the metal ions and the ligand.
– higher charges and small sizes of the metal ions.
– high basicity of the ligand.
13
– lower temperatures. This is consistent with exothermic complexation.
(v)
The dissociation process is non-spontaneous, endothermic and entropically
unfavourable while the complexation (chelation) process is spontaneous,
exothermic and entropically favourable.
REFERENCES
1.
P. W. Heilman R. D. Geilman A. J. Scozzie R. J. Wayner M. J. Gullo and S.
Z. Aryan, J. Pharm. Sci. 69 (1980) 282
2.
A. A. Schilt and W. E. Dunbar, Talanta 16 (1969) 519
3.
A. K. Sen Gupta N. Srivastava and A. A. Gupta, Indian J. Chem., Sect. B 21
(1982) 793
4.
S. Kirschner Y. K. Wei D. Francis and J. G. Bergman, J. Med. Chem. 9
(1966) 369
5.
R. Fusco and R. Trave, Rend. Lomnardo Sci. Pt. I. 91 (1957) 202
6.
A. A. T. Ramadan M. A. El-Behairy A. I. Ismail and M. Mahmoud, Monatsh.
Chem. 125 (1994) 1171
7.
A. A. T. Ramadan R. M. Abdel-Rahman and M. H. Seda, Asian J. Chem.
4(1992) 569
8.
A. A. T. Ramadan R. M. Abdel-Rahman M. A. El-Behairy A. I. Ismail and M.
Mahmoud, Thermochim. Acta 222 (1992) 291
9.
A. Taha B.El-Shetary and W. Linert, Monatsh. Chem. 124 (1993) 135
10.
J. R. Dilworth, Coord. Chem. Rev., 21 (1976) 29
11.
J. R. Merchant and D. S. Clothia, J. Med. Chem., 13 (1970) 335
12.
N. S. Biradar and B. R. Havinale, Inorg. Chim. Acta, 17 (1976) 157
13.
H. N. Fox, Science, 116 (1952) 129
14.
R. K. Agarwal and R. K. Sarin, Polyhedron, 12 (1993) 241
15.
G. Durgaprasad and C. C. Patel, Indian J. Chem., 11A (1973) 1300
16.
K. Sone and Y. Fukuda, “Inorganic Thermochromism (Inorganic Chemistry
Concept)”. Vol. 10. Springer, Heidelberg, 1987
17.
M. El-Behery and H. El-Twigry, Spectrochemical Acta A, 64 (2006) in press
14
18.
A. I. Vogel, “Quantitative Inorganic Analysis”, Longman, London, 1978.
19.
Yoe, J. H. and Jones, A. L., Ind. Eng. Chem., Anal. Ed. 16 (1944) 111
20.
P. Job, Ann. Chim. 9 (1928) 113; 11 (1936) 97
21.
H. M. Irving and H. S. Rossotti, J. Chem. Soc. (1954) 2904
22.
A. A. T. Ramadan, Thermochim. Acta 186 (1991) 235
23.
A. Albert, E. P. Serjeant “Ionization Constants of Acids and Basses”,
Chapman, Hall, Edinburgh (1971)
24.
A. M. Hassan, F. Taha, A. El-Roudi, M. T. Quenawy, J. Indian Chem. Soc. 73
(1996) 325
25.
253
A. M. Hassan, A. El-Roudi, M. T. Quenawy, Egypt. J. Pharm. Sci. 34 (1993)
26.
A. Gergely and T. Kiss, J. Hnorgan. Nucl.Chem., 39 (1976) 109
27.
A. T. Ramadan, B. A. El-Shetary, M. S. Abdel-Moez, H. S. Seleem, Acta
Chim. Hung. 13 (1993) 25
28.
W. Linert, A. Taha, J. Coord. Chem. 29 (1993) 265
29.
A. McAuley, G. H. Nancollas, K. Torrance, Inorg. Chem. 6 (1967) 136
30.
M. S. Abdel-Moez S. Abo El-Wafa and H. Sleem, Thermochim. Acta 149317
(1989) 317
31.
A. G. Evans, J. C. Evans, B. A. El-Shetary, C. C. Rowlands, P. H. Morgan, J.
Coord. Chem. 9 (1979) 19
32.
R. W. Gurney “Ionic Processes in Solution”, McGraw-Hill, New York, 1953
33.
G. Anderegg, Helv. Chim. Acta 51 (1968) 1856
33.
G. H. Nancollas “Interactions in Electrolytic Solutions”, Elsevier, Amsterdam,
1966
34.
C. Degisch and G. H. Nancollas, J. Chem. Soc. A (1970) 1125
35.
M. Z. Born, Phys. 1 (1920) 45; P. Paoletti and A. Viacca, J. Chem. Soc.,
(1964) 5051
36.
U. B. Rao and H. B. Makhur, J. Inorg. Nucl. Chem. 33 (1971) 2919; C. K.
Jorgenson, Acta Chem. Scnd. 10 (1956) 887
15
37.
R. G. Pearson, J. Am. Chem. Soc. 85, 3533 (1963).
38.
W. Linert R. F. Jameson.G. BauerA. and Taha, J. Coord. Chem. 42 (1997) 211
39.
L. G. Van Uitert W. C. Fernelius and B. E. Douglas, J. Am. Chem. Soc. 75
(1953) 2736
40.
M. C. Dayand J. Selbin “Theoretical Inorganic Chemistry”, P. 114. Reinhold,
New York, 1967.
41.
R. G. Pearson, J. Am. Chem. Soc. 85 (1963) 3533
42.
G. Klopman, J. Am. Chem. Soc. 90 (1968) 233
43.
A. T. Ramadan, M. H. Seada, E. N. Rizkalla, Talanta 30 (1983) 245
44.
F. W. Fifield, D. Kealey “Prinicples and Practice of Analytical Chemistry”,
Blackie Academic and Professional, London, 1995, pp. 261, 371
45.
J. H. Yoe and A. L. Jones, Ind. Eng. Chem., Anal. Ed. 16 (1944) 111
16