Ability of two novel α-amino acid-based tridentate ligands to complex Fe3+ in solution

Article abstract of DOI:10.1007/s11164-014-1689-3

Two novel α-amino acid-based tridentate ligands, L¹ and L², were synthesized, and characterized by use of spectroscopic methods. Potentiometric and spectrophotometric methods were used to study the pH-dependent protonation behavior o

Full text of DOI:10.1007/s11164-014-1689-3

Res Chem Intermed
DOI 10.1007/s11164-014-1689-3
Ability of two novel a-amino acid-based tridentate
ligands to complex Fe in solution
3
+
Kirandeep Kaur Minati Baral
•
Received: 11 February 2014 / Accepted: 25 April 2014
Ó Springer Science+Business Media Dordrecht 2014
1
2
Abstract Two novel a-amino acid-based tridentate ligands, L and L , were
synthesized, and characterized by use of spectroscopic methods. Potentiometric and
spectrophotometric methods were used to study the pH-dependent protonation
behavior of the ligands and their complexation towards Fe(III) in aqueous medium
of ionic strength, l = 0.1 M KCl at 25 ± 1 °C. The stability constants of complex
1
2
species of the types FeLH , FeLH , FeLH, FeL, and FeL (L = L or L ) were
3
2
2
determined, and probable structures of the ferric complexes were proposed on the
basis of molecular modeling. The complexing ability of the ligands was compared
with that of deferiprone (a drug used for medical treatment of iron intoxication) and
transferrin (the main Fe-binding protein in plasma).
Keywords Tridentate ligands ꢀ Amino acids ꢀ Potentiometry ꢀ Spectrophotometry ꢀ
Protonation and stability constants
Introduction
Chelation therapy involves dispensing chelating organic molecules to selectively
and specifically remove excess metal ions from the body [1]. There is a need for
novel, safe, cost-effective, selectively efficient, and orally efficacious chelating
agents for Fe(III) [2–5]. Chemists have synthesized a variety of chelating agents and
investigated their complexation behavior with metal ions in a physiological
environment [6–11]. Because of the importance of selectivity for a specific metal
ion and oral efficacy, a large number of different ligands have been synthesized,
including hexadentate, bidentate, and tridentate ligands. Each of these ligand
K. Kaur ꢀ M. Baral (&)
Department of Chemistry, National Institute of Technology,
Kurukshetra 136119, Haryana, India
e-mail: minatib@gmail.com
123

K. Kaur, M. Baral
categories has its inherent advantages and disadvantages. Although siderophores,
3
?
well known hexadentate Fe -selective chelating agents secreted by microorgan-
isms, are highly selective for Fe(III), they are orally ineffective [12–15]. In contrast,
bidentate ligands are orally efficient but less selective [16, 17]. Tridentate ligands
have the advantages of both hexadentate and bidentate ligands, and might, therefore,
be ideal chelators of Fe(III). Tridentate ligands with different donor atoms, mostly
oxygen and nitrogen donors, for such biologically active metal ions as Fe(III), have
recently been synthesized [18–20].
On the basis of the results described above, we report herein the synthesis of
1
2
tridentate ligands (L and L ) containing a-amino acids by a modified literature
route (Fig. 1). The protonation behavior of the ligands and their metal complexation
with Fe(III) have been investigated for the first time. Complexation of the ligands
?
with H and Fe(III) was studied by use of potentiometric and spectrophotometric
methods. Theoretical calculation of the properties of the metal complexes was also
performed to support the experimental results.
Results and discussion
Synthesis and protonation constants of the ligands
1
2
The ligands L and L were synthesized by use of a literature procedure [21, 22]
Scheme 1) and characterized on the basis of spectroscopic data. A distinct color
(
1
2
change immediately after addition of glycine (L ) or L-alanine (L ) to salicylal-
dehyde in ethanol indicated formation of the Schiff bases. Although isolation of the
products was not possible as such, white crystalline solids could be isolated after in-
situ reduction of the Schiff base then acidification of the solution with conc. HCl.
Experimental evaluation of protonation constants is important because formation
of metal complexes is perceived as competition between protons and metal ions for
the ligand moiety. These constants are, hence, required as input data for calculation
of formation constants of the metal complexes Potentiometric and spectrophoto-
metric experiments were performed to determine the protonation constants of the
ligands. Potentiometric titrations of the ligands in aqueous solution of ionic strength
(
l) = 0.1 M KCl were performed in the pH range 2–11, at 25 ± 1 °C, after
addition of excess 0.1 M HCl; the titration curves for both ligands are shown in
Fig. 2. This particular pH range is preferred because of the types of proton to be
deprotonated. Protons in the ligands are similar to those in the corresponding Schiff
bases except for the protons of the secondary amine group; for comparison, reported
log K values of Schiff bases (pH 2–11) are listed in the Table 1 [22]. In this
investigation, no appreciable change in the curves was observed on increasing the
pH beyond 10.5, as is clearly apparent from Figs. 3a and 4a.
These curves were analyzed by use of Hyperquad 2006 software, which gave
three protonation constants for the ligands, as given in Table 1. The general form of
the equilibrium between proton and ligand, and the formula used to calculate
formation constants, are given by Eq. 1:
1
23

Aability of two novel a-amino acid-based tridentate ligands
Fig. 1 Structure of the ligands
L and L
1
2
HO
R
R
H
N
O
O
O
1
. KOH, stirring for 2 hrs, rt
OH
H2N
HO
+
OH
2. NaBH4 3. HCl
.HCl
(ethanolic medium)
1
2
R= -H (L ), -CH3 (L )
1
Scheme 1 Synthesis of the ligands L and L
2
1
Fig. 2 Potentiometric titration curves for L (i) and L (ii) in the absence and presence of Fe(III), where
2
‘‘a’’ is moles of base added per mole of ligand present. (Circles and triangles represent theoretical values
for the ligands only and the metal complexes, respectively; solid lines represent experimental values and
dotted lines indicate where the solution became turbid)
LHnꢁ1 þ H ¼ LHn; Kn ¼ ½LHnꢂ=½ LHnꢁ1ꢂ ½Hꢂ
ð1Þ
1
where L = L or L and n = 1, 2, or 3.
2
The protonated form of the ligand under highly acidic conditions can be
?
represented as LH3 (in rest of the paper charges are omitted for clarity). As shown
in Fig. 2(i), (ii), the first inflection in the potentiometric curves for both ligands was
123

K. Kaur, M. Baral
1
23

Aability of two novel a-amino acid-based tridentate ligands
1
Fig. 3 Experimental electronic spectra of L in the absence (a) and presence (b) of Fe(III) at different pH
2
Fig. 4 Experimental electronic spectra of L in the absence (a) and presence (b) of Fe(III) at different pH
observed at a = 0, where ‘‘a’’ represents the number moles of base added per mole
of the ligand. Up to this point (a = 0) excess acid was being neutralized. Another
inflection observed at a = 3 clearly indicates that three moles of the protons had
been released by the ligands. This deprotonation can be attributed to release of
protons from the hydroxyl, secondary amine, and carboxylic acid groups of the
ligands.
1
2
The electronic spectra of L and L were recorded in the pH ranges 4.27–10.52
and 4.32–9.10.64, respectively (Figs. 3a, 4a). At low pH, ligand peaks were
observed at 200 nm, with a shoulder at 218, and at 275 nm; these can be attributed
to the primary and secondary bands of the phenyl ring. As pH was increased above
4
.27–5.94, a hyperchromic shift of the peak at 200 nm was observed without any
change in the peak at 275 nm. Above pH 5.94, a new peak started to appear at
35 nm and a bathochromic shift of the peak at 275 nm to 295 nm was observed,
2
123

K. Kaur, M. Baral
presumably because of deprotonation of the phenolic –OH to form phenolate ion,
which stabilizes the p* excited state, because of charge delocalization, and brings
the lowest excited state closer to the highest ground state, thus enabling a lower-
energy (longer-wavelength) transition. Above pH 10.52, no appreciable changes in
the spectra were observed. The spectral titrations revealed three isosbestic points at
2
18, 257, and 279 nm (Fig. 3a). The presence of isosbestic points and the shifts in
the peaks might be because of the presence of equilibria between the different
protonated and deprotonated species of the ligand, as suggested in Table 1. The
2
spectral behavior of ligand L was similar to that of ligand L . At low pH ligand L
1
2
had absorption bands at 203 nm, with a small shoulder at 220 nm, and at 274 nm;
these shifted as the pH was increased (Fig. 4a). The log K values calculated from
the spectrophotometric data by use of the pHAb software are given Table 1. The
values obtained are in good agreement with those obtained by use of the
potentiometric method.
The species-distribution diagrams, which represent percentage formation of the
1
different deprotonated and protonated species of L and L at different pH, are
2
1
2
shown in Fig. 5. Ligands L and L can both be regarded as triprotic acids. The
diagrams indicate that in acidified solution (up to pH 3), the ligands are fully
1
protonated, i.e. L H and L H . The carboxylic acid group, being highly acidic, is
2
3
3
deprotonated first to form the species LH with 100 % formation between pH 3.5
2
and 7.5. Increasing the pH further causes deprotonation of the phenolic proton of the
1
2
ligands, forming the species LH, with 80 and 84 % formation for L and L ,
respectively. Complete deprotonation of the ligands, forming L, occurs above
pH 10. The low protonation constant, log K = *2.43–2.46, for carboxylic acid
protonation indicates its high tendency to deprotonate. The high log K values of
*
the ammonium ion of zwitterions (LH ). The possible form [A] is in dynamic
11.71–11.78 for formation of the free amine at higher pH suggest dissociation of
2
equilibrium with form [B]; the percentage of each varies with pH. Because of the
possibility of formation of many micro species by formation of zwitterions, the
?
stepwise dissociation of the protonated species LH3 is depicted in Scheme 2 for
the convenience of readers; this emphasizes the coordination priority of the donor
atoms during chelation of the metal.
For the Schiff base analogs, the imine group is expected to deprotonate before the
phenolic group, and its log K value is assigned as 7.57. Deprotonation of the amine
group to the imine (–C=N–) group resulted a decrease in the log K value,
2
presumably because of the lower density of negative charge on the sp -hybridized
3
nitrogen in –C=N– than on the sp -hybridized amine nitrogen, and also because of
intramolecular hydrogen bonding between the phenolic –OH and the imine nitrogen
1
atom [22, 23]. In the reduced forms (L and L ), the secondary amine exists in the
2
?
form –NH2 at low pH, is highly basic in nature, and is expected to deprotonate at a
higher pH, after deprotonation of the phenolic –OH.
Complexation studies with Fe(III)
1
Potentiometric titrations of the ligands (L and L ) in the presence of Fe(III) were
2
conducted in aqueous solution at l = 0.1 M KCl and 25 ± 1 °C for 1:1 and 1:2
1
23

Aability of two novel a-amino acid-based tridentate ligands
1
Fig. 5 Species-distribution curves for ligands L (a) and L (b)
2
?
3
Scheme 2 Dissociation scheme for protonated species of ligands L (LH )
molar ratio of metal to ligand. Potentiometric curves of the ligands in the presence
of Fe(III) are shown in Fig. 2. The deviation of the metal–ligand curves from the
free ligand curves indicates the potential of the ligands to form complexes with
1
2
Fe(III) in solution. For both L and L , the first inflection point is obtained at a = 1,
which indicates release of one mole of protons from the ligands on complexation
with Fe(III). Turbidity appeared at a = 6, indicating release of six protons. These
potentiometric data were analyzed by use of Hyperquad 2006. Several different
possible models were tested by the software; the best fit was obtained for the models
123

K. Kaur, M. Baral
FeLH , FeLH , FeLH, FeL, and Fe(L) . Log b values for formation of the different
3
2
2
complex species in solution are tabulated in Table 2. The general equations which
represent the equilibrium reactions for formation of the complexes are:
M þ L ¼ M Lb110 ¼ ½MLꢂ=½Mꢂ½Lꢂ
ð2Þ
ð3Þ
M þ L þ H ¼ MLH; b111 ¼ ½MLHꢂ=½Mꢂ½Lꢂ½Hꢂ
2
M þ L þ 2H ¼ MLH2; b112 ¼ ½MLH2ꢂ=½Mꢂ½Lꢂ½Hꢂ
ð4Þ
ð5Þ
ð6Þ
3
M þ L þ 3H ¼ MLH ; b ¼ ½MLH ꢂ=½Mꢂ½Lꢂ½Hꢂ
3
113
3
2
M þ 2L ¼ MðLÞ ; b ¼ ½ML2ꢂ=½Mꢂ½Lꢂ
120
2
Spectrophotometric titrations were also performed to calculate formation constants
1
2
for the metal complexes formed in solution. The electronic spectra of L and L in the
presence of Fe(III) were recorded at different pH (Figs. 3b, 4b). Shifting of the ligand
peaks were observed, with formation of isosbestic points as the pH was continuously
1
increased. Peaks at 200, 218, and 275 nm for ligand L were shifted bathochromically
to 205, 235, and 297 nm, respectively. The absorbance intensity of the peak at 200 nm
also increased slightly. An appreciable hyperchromic shift was observed for a new
peak obtained at 235 nm. An appreciable bathochromic shift with increasing pH was
also observed for the phenyl ring band at 275 nm. These shifts and formation of
isosbestic points are indicative of interaction of Fe(III) with the ligand. The
2
1
spectrophotometric results for ligand L were similar to those for L , which is
2
indicative of similar modes of complexation in solution (Fig. 4b). The peaks for L at
218, 220, and 274 nm were shifted to longer wavelengths, 237 and 295 nm, with
formation of four isosbestic points at 209, 219, 245, and 279 nm.
1
2
The pH-dependent species-distribution diagrams for ligands L and L in the
presence of Fe(III) are shown in Fig. 6, and indicate that complexation starts below
pH 2. The very first species formed was FeLH . As pH was increased, successive
3
formation of species FeLH and FeLH, in very small amounts, occurred in the pH
2
range 3–4. At pH 4, the species FeL is predominant. Later, formation of FeL2
occurs, and it predominates above pH 5. It is well established that the salt FeCl
3
3
exists as [Fe(H O) ] in aqueous solution. Complexation of the metal starts with
?
2
6
formation of FeLH , the other sites being occupied by water molecules. As the pH is
3
increased the ligand releases the protons successively to form complexes of the type
FeLH , FeLH, FeL, and FeL ; possible modes of complexation are shown in Fig. 7.
The ligands coordinate to Fe(III) through amine-N, phenolate-O, and carboxylate-O
2
2
atoms. On examination of log b values for formation of the different species it was
observed that L formed stronger complex with Fe(III) than L . The presence of the
2
1
2
methyl group in the alanine moiety of L increases the basicity of the ligand, hence
the ability to form complexes with Fe(III) was enhanced. The total energy (E )
T
1
calculated by use of the semi-empirical PM6 method for the complex Fe(L )
2
2
ET = -5126.2456 kcal/mol) is higher than that for Fe(L )
(
(E =
T
2
2
5669.9114 kcal/mol), which also suggests better complexing by L than by L .
1
-
The optimized structures of the complexes are shown in Fig. 8.
1
23

Aability of two novel a-amino acid-based tridentate ligands
1
2
Table 2 Stability constants (log b) for the Fe(III) complexes formed in solution with ligands L and L
1
2
Species
log b (Fe-L )
log b (Fe-L )
Potentiometry
Spectrophotometry
Potentiometry
Spectrophotometry
FeLH
FeLH
FeLH
FeL
3
24.89
21.78
18.98
16.34
28.67
24.83
21.69
18.97
16.32
28.66
25.79
22.62
19.71
16.98
29.63
25.76
22.63
19.74
16.94
29.65
2
FeL
2
The complexation ability of the ligands was also examined by calculating pFe
under specific pH, [L], and [Fe] conditions. In this study pFe was calculated at
-
L] = 10 M, [Fe] = 10 M, and pH 7.4 for both ligands. The pFe values for L
5
-6
1
[
and L are 24.51 and 23.86, respectively. One unit difference in pFe is the evidence
2
2
1
of the better complexation ability of L than of L for Fe(III). To check the
complexation ability of the ligands under acidic and basic conditions, pFe values of
1
2
L and L were plotted against pH in the range 2–11. From Fig. 9 it is clear that this
order of complexation ability is followed over the whole pH range and that the
ligands had the ability to complex Fe(III) in a wide range of pH. Next, the pFe
values of the ligands were compared with those of three N-carboxyalkyl 3-hydroxy-
4
-pyridines [24], deferiprone [25] (Fig. 10), and transferrin [26]. The pFe values at
3
4
5
pH 7.4 for ligands L , L , and L are 21.3, 19.9, and 19.8, respectively. For
deferiprone and transferrin, the pFe values are 19.3 and 20.3, respectively. Under
1
2
similar conditions, the pFe values for the ligands L and L are almost three to four
times greater, which clearly implies they are good chelating agents for Fe(III). So,
1
2
ligands L and L can be investigated for their potential application for iron
overload chelation therapy.
Conclusions
1
2
Two amino acid-based ligands, glysal (L ) and alasal (L ), coordinate the metal ion
Fe(III) through N, O, and O donor atoms of sec-amine, phenol, and carboxylic acid
groups, respectively. The species formed by complexation in solution are MLH3,
MLH , MLH, ML, and ML . It is apparent from their formation constants and pM
2
2
2
values that both ligands are efficient chelators of Fe(III), with L being slightly
better, as confirmed by theoretical studies. The pM values of the ligands at pH 7.4
show that their ability to complex Fe(III) is better than that of transferrin (an Fe-
binding protein in plasma), deferiprone (a standard drug), and N-carboxyalkyl
3
-hydroxy-4-pyridines (derivatives of deferiprone). Substitution of a hydrogen atom
by a methyl group (higher ?I effect) leads to the formation of a more stable
complex, implying that use of groups with sufficient electron-donating nature may
result in ligands which are more efficient chelators of Fe(III), assuming interference
by the bulkiness of the group is taken into consideration.
123

K. Kaur, M. Baral
1
Fig. 6 pH-dependent species-distribution diagrams for FeL (a) and FeL (b)
2
1
Fig. 7 Formation of complex species and proposed coordination modes of L and L
2
1
Fig. 8 Optimized structures of complexes Fe(III)–L (a) and Fe(III)–L (b), calculated by use of the
2
semi-empirical PM6 method
1
23

Aability of two novel a-amino acid-based tridentate ligands
1
Fig. 9 pH vs pFe plot for L
and L . pFe was calculated for
2
-
5
[
1
L] = 1 9 10 and [Fe] = 1 9
0
-6
by use of protonation
constants of the ligands and
complexation constants, b11n
O
R
L3 R=
N
N
O
O
O
L4
R=
O
O
O
L5
R=
O
O
O
Deferiprone
Some derivatives of Deferiprone
3
Fig. 10 Structures of deferiprone and its derivatives L , L , and L
4
5
Experimental
All chemicals and solvents were of analytical grade and were available commer-
cially. Glycine, alanine, KOH, and HCl were obtained from Fisher Scientific.
Salicylaldehyde and sodium borohydride were obtained from Sigma Aldrich.
Anhydrous ferric chloride was procured from Qualigens Fine Chemicals (Thermo
Fisher Scientific), and solvents were from Merck. Chemicals were used as received,
except that all solvents were dried by use of standard methods [27].
1
UV–visible spectra were recorded on an Hitachi U0080D spectrophotometer. H
NMR spectra were recorded on a Bruker Avance II 400 NMR spectrometer at SAIF
Panjab University, Chandigarh, using tetramethylsilane (TMS) as internal reference.
Chemical shifts are expressed in parts per million (ppm). Potentiometric studies
were performed with a Sension 02 pH meter. C, H, N analysis was performed with a
Euro EA 3000 elemental analyzer.
1
Synthesis of L
1
-[(2-Hydroxyphenyl)methylamino]acetic acid (L ) was synthesized by use of a
2
literature procedure [28]. Glycine (6.0 g, 80 mmol) and potassium hydroxide (4.5 g,
0 mmol) were dissolved in 40 mL water, forming a clear solution, and the mixture
8
123

K. Kaur, M. Baral
was then added to 20 mL ethanolic solution of salicylaldehyde (9.6 g, 80 mmol).
The reaction mixture was stirred for 2 h. A dark yellow solution was obtained, to
which one equivalent of sodium borohyride (3.0 g, 80 mmol) and 3 or 4 drops of
saturated aqueous sodium hydroxide solution were added. On completion of the
addition, the solution became white and turbid. The solution was stirred overnight.
The solution was filtered, and the pH was adjusted to 4 with concentrated
hydrochloric acid. Under refrigeration a white solid was obtained, which was
collected by filtration, washed with ethanol, and vacuum dried. White solid
(
1
82.01 %): H NMR: (D O/DCl) 6.93, (d, J = 8.12, 1H, Ar–H ortho to alkyl chain),
2
7
–
4
.01, (t, J = 11.32, 1H, Ar–H para to –OH), 7.22, (d, J = 7.48, 1H, Ar–H ortho to
OH),7.37, (t, J = 14.4, 1H, Ar–H meta to –OH), 4.47, (t, J = 4.4, 2H, –CH–),
?
.07, (t, J = 10.32, 2H, –CH ), 7.701, (s, 2H, –NH2 ), Anal. Calcd. C H NO ꢀHCl
2
9
11
3
(
found) C-49.70 (49.17), H-5.61(5.07), N-6.40 (6.89).
2
Synthesis of L
2
2
-[(2-Hydroxyphenyl)methylamino]propanoic acid (L ) was synthesized by use of a
similar method. Alanine (3.0 g, 33.6 mmol) and potassium hydroxide (1.9 g,
3.6 mmol) were dissolved in 25 mL water, forming a clear solution, and the
mixture was then added to 10 mL ethanolic solution of salicylaldehyde (4.1 g,
3.6 mmol). The reaction mixture was stirred for 2 h. One equivalent of sodium
3
3
borohyride (1.3 g, 33.6 mmol) and 3 or 4 drops of saturated aqueous sodium
hydroxide solution were added dropwise to the dark yellow solution. The solution
was stirred overnight. The solution was filtered and the pH was adjusted to 4 with
concentrated hydrochloric acid. Under refrigeration a white solid was obtained,
which was collected by filtration, washed with ethanol, and vacuum dried. White
1
solid (81.67 %): H NMR: (D O/DCl) d = 1.75 (d, J = 7.28, 3H, –CH ), 4.13, (q,
2
3
?
H, –CH–), 4.39, (m, 2H, –CH ), 7.91, (s, 2H, –NH ), 6.89, (d, J = 10.24, 1H, Ar–
1
H ortho to alkyl chain), 6.98, (m, 1H, Ar–H para to –OH), 7.16, (q, 1H,Ar–H meta
2 2
to –OH), 7.35, (m, 1H, Ar–H ortho to –OH), Anal. Calcd. C H NO ꢀHCl (found)
1
0
13
3
C-51.81 (51.32), H-6.10 (5.89), N-6.01 (5.73).
Titration procedure
Potentiometric studies of the ligands and ferric complexes were performed at
5 ± 1 °C, maintained by use of a double-wall glass-jacketed titration cell
2
connected to a constant temperature circulatory bath. A glass electrode coupled
with a Sension 02 pH meter was used to perform the titration. The pH meter was
calibrated by the standard method of using buffer solutions [29]. Titrations were
performed between standard solutions of HCl and KOH; the pH meter reading was
converted into hydrogen ion concentration by use of the calculated hydrogen ion
concentration (pK = 13.77 ± .05).
w
Triply distilled and deoxygenated water was used for preparation of all solutions.
Standardization of the 0.1 M solution of KOH was against potassium hydrogen
phthalate. HCl solution (0.1 M) was standardized against standardized KOH
solution; 0.1 M ionic strength was maintained by adding a calculated amount of
1
23

Aability of two novel a-amino acid-based tridentate ligands
1
M KCl solution. Because of poor solubility of the ligands in water, 0.01 M ligand
solution was prepared in 0.1 M HCl solution; 0.01 M ferric ion solution was
prepared by dissolving anhydrous ferric chloride in water. Final concentrations of
the ligands and metal were 0.001 M. Initial total volume of the titration mixture was
5
0 mL. pH meter readings were taken after each addition of base, with sufficient
time being allowed to achieve equilibrium. Titrations were performed with metal to
ligand ratios, L:M, 1:1 and 2:1. The computer software Hyperquad 2006 [30, 31]
was used to calculate the protonation constants of the ligands and the formation
constants of Fe(III) complexes. HYSS software [32, 33] was used to obtain pH-
dependent species distribution diagrams from measured equilibrium constants.
Spectrophotometric titrations were conducted under the same conditions as
potentiometric titrations. In these studies, dilute solutions of the ligands
1
-4
2
-4
-4
(
L = 1 9 10 M, L = 5 9 10
M) and metal ion (1 9 10
M,
-
4
5
0
0
9 10 M) were made acidic by addition of an appropriate amount of
.1 M HCl at an ionic strength of 0.1 M KCl, and titrated against standard
.1 M KOH solution. After each addition of base, sufficient time was allowed to
achieve equilibrium. When constant pH was achieved, a sample of the solution was
removed and its spectrum was recorded.
Molecular modeling
All calculations were performed by use of the Windows-based software Hyperchem,
version 8.0 [34] The initial minimum strain geometry of the proposed metal
complex was obtained by use of molecular mechanics methods, using the MM?
force field. The minimum strain energy structure was used as input file for re-
optimization of the structure by use of the PM6 Hamiltonian in a semi-empirical
method coupled with Polak-Rebiere algorithms in the gas phase.
Acknowledgments The authors would like to thank the Council of Scientific and Industrial Research
(CSIR), Human Resource Development Group, PUSA, New Delhi, India, for financial Support and the
National Institute of Technology, Kurukshetra, India, for providing infrastructure and research facilities.
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