Competitive 7Li NMR Study on the Mn2+, Zn2+ and Cd2+ Complexes of Two New Branched Hexadentate (N6) Amines Containing the Pyridine Moiety in Nitromethane and Acetonitrile Solutions

Article abstract of DOI:10.1007/s10953-014-0196-x

Lithium-7 NMR spectroscopy was used to investigate the stoichiometry and stability of a Li⁺ complex with two new branched amines, 4,7-bis(2-pyridylmethyl)-4,7-diazadecane-1,10-diamine (L¹) and 4,8-bis(2-pyridylmethyl)-4,8-diazaundecane-1,11-diamine (L²), in acetonitrile and nitromethane. A competitive ⁷Li NMR method was also employed to probe the complexation of Mn²⁺, Zn²⁺ and Cd²⁺ ions with L¹ and L² in the same solvent systems. The formation constants of the resulting complexes were evaluated from computer fitting of the mole ratio data with an equation that relates the observed chemical shifts to the formation constant. In both solvents, the stability of the resulting 1:1 complexes was found to vary in the order Cd²⁺ > Zn²⁺ > Mn²⁺ > Li⁺.

Full text of DOI:10.1007/s10953-014-0196-x

J Solution Chem (2014) 43:1218–1231
DOI 10.1007/s10953-014-0196-x
Competitive 7Li NMR Study on the Mn2+, Zn2+ and Cd2+
Complexes of Two New Branched Hexadentate (N6)
Amines Containing the Pyridine Moiety in Nitromethane
and Acetonitrile Solutions
•
Hassan Keypour Mohammad Hasan Zebarjadian Majid Rezaeivala
•
•
•
Mojtaba Shamsipur Hamid Reza Bijanzadeh
Received: 8 July 2013 / Accepted: 5 April 2014 / Published online: 22 July 2014
Ó Springer Science+Business Media New York 2014
Abstract Lithium-7 NMR spectroscopy was used to investigate the stoichiometry and
stability of a Li? complex with two new branched amines, 4,7-bis(2-pyridylmethyl)-4,7-
diazadecane-1,10-diamine (L1) and 4,8-bis(2-pyridylmethyl)-4,8-diazaundecane-1,11-dia-
7
mine (L2), in acetonitrile and nitromethane. A competitive Li NMR method was also
employed to probe the complexation of Mn2?, Zn2? and Cd2? ions with L1 and L2 in the
same solvent systems. The formation constants of the resulting complexes were evaluated
from computer fitting of the mole ratio data with an equation that relates the observed
chemical shifts to the formation constant. In both solvents, the stability of the resulting 1:1
complexes was found to vary in the order Cd2? [ Zn2? [ Mn2? [ Li?.
Keywords Competitive 7Li NMR ꢀ Complexation ꢀ Nonaqueous solvent ꢀ Nitromethane ꢀ
Acetonitrile ꢀ Pyridine moiety ꢀ Hexadentate (N6) amines
1 Introduction
Polyamines constitute one of the categories of receptors most widely employed in coor-
dination and supramolecular chemistry [1]. This broad interest stems from their versatility.
Polyamine receptors can coordinate metal ions if the number of available amino groups is
H. Keypour (&) ꢀ M. H. Zebarjadian
Faculty of Chemistry, Bu-Ali Sina University, 65174 Hamedan, Iran
e-mail: haskey1@yahoo.com
M. Rezaeivala
Department of Chemical Engineering, Hamedan University of Technology, 65155 Hamedan, Iran
M. Shamsipur
Department of Chemistry, Razi University, Kermanshah, Iran
H. R. Bijanzadeh
Department of Chemistry, Tarbiat Modarres University, PO Box 14115-175, Tehran, Iran
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sufficient. On the other hand, they can coordinate anions or dipolar species at pH values
low enough to allow extensive protonation of the amino groups. Different books and
reviews cover this topic [2, 3]. It was anticipated that these ligands, which generate slightly
rigid chelate rings with increased torsional strain, would support a wide range of coordi-
nation geometries. Studies using nickel have thus far revealed the expected structural
diversity [4–9].
Polyamines, both cyclic and noncyclic, are the ligands studied most with respect to their
basicity and coordination capability toward metal ions [10–13]. The structural variety of
polyamines and the stability of their complexes are reflected in their remarkable coordi-
nation chemistry. In aqueous solution, most amines are basic and there will always be
competition between the lone pairs on oxygen (OH- or H2O) and the lone pair on the
nitrogen for the metal center. This competition can be overcome to a certain extent by
using non-aqueous systems or the free amine as a solvent [14, 15]. Monodentate, and
especially polydentate, amines have played an outstanding role in the development of
coordination chemistry [16]. The study of polyamine complexes has contributed in large
measure to the rationalization of coordination compound formation. Among the variety of
spectroscopic and electrochemical methods used for the study of metal ion–amine com-
plexes [17, 18], it has been found that nuclear magnetic resonance spectrometry offers a
very sensitive technique for studying changes in the immediate chemical environment of
metal ions in solution [19–23].
In previous works we have used proton and alkali metal NMR techniques to study
the thermodynamics [24–26] and kinetics [27, 28] of metal ion complexation with some
macrocyclic ligands in nonaqueous and mixed solvents. In this study, we are interested
in investigating the complexation of Mn2?, Zn2? and Cd2? ions with 4,7-bis(2-pyri-
dylmethyl)-4,7-diazadecane-1,10-diamine (1) (L1) and 4,8-bis(2-pyridylmethyl)-4,8-di-
azaundecane-1,11-diamine (L2) in nonaqueous solvents. However, these cations have
unsuitable NMR properties, such as low receptivity, high quadropole moment, insen-
sitive chemical shifts, etc. [22, 23, 29]. Thus, we employed 7Li NMR as a very sen-
sitive probe [20, 21, 25–28] to study the metal ion–amine complexes of interest in
nitromethane (NM) and acetonitrile (AN) solutions by a competitive technique. Our
interest in the physicochemical properties of the branched hexadentate amines
encouraged us to determine the formation constants of transition metal ion complexes
with them and to gain some useful information about the properties of complex for-
mation in the solvents.
2 Experimental Section
2.1 Chemicals Used
The ethane-1,2-diamine (Fluka, 99.5 %), propane-1,3-diamine (Fluka, 99 %), pyridine-2-
carbaldehyde (Fluka, 98 %) and N-(3-bromopropyl)phthalimide (Merk, 98 %) used were
of the highest purity available. Lithium perchlorate (Merck, 99 %) was purified and dried
according to a literature method [30]. The nitrate salts of all other cations used were of the
highest purity available from Fluka and used without any further purification except for
vacuum drying over P2O5. Spectroscopic grade AN (Merk, 99.5 %) and NM (Merk, 99 %)
were purchased from Merck and used as received. All NMR measurements were made on a
Jeol FX90Q FT-NMR spectrometer with a field strength of 21.13 kG. At this field, lithium-
7 resonates at 33.74 MHz. An aqueous 4.0 molꢀL-1 LiCl solution was used as the external
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reference and the reported lithium-7 chemical shifts are referenced to this solution. The
paramagnetic (downfield) shift from the reference is designated as being positive. The
concentration of all cations used was 0.01 molꢀL-1. All chemical shift measurements were
carried out at a temperature of 25.0 0.1 °C.
2.2 General Synthesis of the Branched Hexadentate Amines (L1, L2)
The general method for the synthesis of the hexadentate amines is as follows. A solution
of diamines (10 mmol) in dry EtOH (30 mL) was added dropwise to a warm solution of
pyridine-2-carbaldehyde (2.14 g, 20 mmol) in dry EtOH (30 mL) over a period of 2 h.
The mixture was refluxed under stirring for 12 h and then allowed to cool to the room
temperature. Solid sodium borohydride (3.78 g, 100 mmol) was then added slowly and
the reaction mixture was heated at reflux for a further 2 h. The solution was filtered and
the filtrate volume reduced to 20 mL by rotary evaporation. Excess water was added and
the product was extracted with chloroform (3 9 25 mL); the combined chloroform
solutions were separated and dried over magnesium sulfate. The chloroform was removed
by rotary evaporation to leave a yellow oil (90–95 %). The resulting product (10 mmol)
was dissolved in AN (20 mL), and K2CO3 (2.76 g, 20 mmol) was added. The mixture
was refluxed and then a solution of N-(3-bromopropyl)phthalimide (3.28 g, 12 mmol) in
AN (40 mL) was added. The mixture was refluxed for 24 h and then filtered while hot.
The filtrate was reduced to dryness by rotary evaporation. The resulting product was
boiled under reflux for 20 h in aqueous HCl (25 %, 100 mL). After this time the solution
was evaporated to a small volume (ca. 25 mL) under vacuum and cooled in a refrigerator
for several hours. The solid present was filtered off and the filtrate was evaporated to
dryness under vacuum. Water (50 mL) was added to the mixture and the pH was adjusted
to 12 with sodium hydroxide. The product was extracted with chloroform (3 9 25 mL),
and then the combined chloroform solutions were separated and dried over magnesium
sulfate. The chloroform was removed by rotary evaporation to leave a brown oil
(Scheme 1).
2.2.1 4,7-Bis(2-pyridylmethyl)-4,7-diazadecane-1,10-diamine (L1)
Yield: 2.35 g (63 %). Analysis, found: C, 66.9; H, 8.73; N, 23.08 %. Calculated for
C20H32N6: C, 67.38; H, 9.05; N, 23.57. IR (Nujol, cm-1): 3,350, 1,642 t(C=N), 1,591,
1,444 [t(C=C) and t(C=N)py], 2,908, 2,851 [t(CH2)], 3,350 [t(NH2)]. EI-MS: m/z (%),
1
[L1]?, 356 (100 %). H NMR (CDCl3 ppm): dH = 1.49 (m, 4H, Hi, i’), 2.33 (m, 4H,
Hj,j’), 2.45 (t, 4H, Hh,h’), 2.76 (s, 4H, Ha,a’), 3.52 (s, 4H, Hb,b’), 3.84 (m, 4H, -NH2),
7.20 (d, 2H, Hd,d’), 7.44 (d, 2H, He,e’), 7.19 (m, 2H, Hf,f’), and 8.36(d, 2H, Hg,g’)
(Scheme 1).
2.2.2 4,8-Bis(2-pyridylmethyl)-4,8-diazaundecane-1,11-diamine (L2)
Yield: 2.01 g (54 %). Analysis, found: C, 68.0; H, 9.1; N, 22.5 %. Calculated for
C22H36N6: C, 68.1; H, 9.3; N, 22.7. IR (Nujol, cm-1): 3,351, 3, 288 [t(NH2)], 1,590, 1,434
1
[t(C=C) and t(C=N)py]. EI-MS: m/z (%), [L1]?, 386 (100 %). H NMR (CDCl3 ppm):
dH = 1.56 (m, 2H, Ha), 1.59 (m, 4H, Hj,j’), 2.40 (m, 4H, Hb,b’), 2.49 (m, 4H, Hi,i’), 2.69
(m, 4H,Hk,k’), 3.59 (s, 4H, Hc.c’ Hd,d’), 4.50, 4.64 (m, 4H, Hl,l’),7.03 (m, 2H, He,e’),
7.23 (m, 2H, Hg,g’), and 7.53 (m, 2H, Hf,f’) (Scheme 1).
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O
(CH2)n
n=2,3
N
i: EtOH
NH2-(CH2)n-NH2
n=2,3
+
2
ii: NaBH4
N
N
N
HN
H
d'
d
e'
e
a'
a
c
c' b'
b
f'
f
i: CH3CN
ii: K2CO3
N
N
N
N
g
g'
h'
h
O
h'
i
i'
j
iii:
Br
N
NH2
H2N
O
L1
iv:HCl 25%
v:NaOH
e'
e
f'
f
b'
b
d
d' c'
c
g'
g
a
N
N
N
N
h
h'
i'
i
j'
j
k'
k
NH2
H2N
L2
Scheme 1 Synthesis of the branched hexadentate amines showing the lettering scheme for NMR
assignments
2.3 Preparation of Solutions
Stock solutions of the amines were prepared by weighting appropriate amounts at con-
centrations of 0.12 molꢀL-1 in AN and NM, respectively. Appropriate quantities of metal
salts at concentrations of 0.01 molꢀL-1 (dissolved in the same solvents, respectively) were
added to the [ligand]/[metal] mole ratio of 4, using a microsyringe. Then, the mixture was
7
under ultrasonic irradiation for 5 min and after that the Li NMR spectra of the resulting
solutions were recorded at 25.0 ( .1) °C.
3 Results and Discussion
The 7Li chemical shifts were monitored as a function of the mole ratio of complexant, 4,7-
bis(2-pyridylmethyl)-4,7-diazadecane-1,10-diamine (L1) and 4,8-bis(2-pyridylmethyl)-4,8-
diazaundecane-1,11-diamine (L2), to lithium ion in AN and NM solutions, in the absence
and presence of equimolar concentration of the different M2? ions used. In all cases, only a
7
single Li resonance was observed, indicating that the exchange rate of Li? ion between
the bulk solution and the complexed sites is fast on the NMR time scale.
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J Solution Chem (2014) 43:1218–1231
All the resulting chemical shift–mole ratio plots are illustrated in Figs. 1, 2 and 3. As
can be seen, in the absence of other metal ions (Fig. 1), addition of L1 and L2 amines to the
lithium ion solution causes an almost linear paramagnetic shift in the observed 7Li
chemical shift, where dobs, a population-averaged combination of the chemical shifts of
Li? ion in the bulk solvent, dLi, and that of the Li? ion complexed with ligand, dLiL, which
are quite different from each other due to differences in the electron densities surrounding
the probe nucleus in the two sites, is given as dobs = PLi dLi ? PLiL dLiL (where,
PLi ? PLiL = 1 in the course of titration of metal ion with the ligand). However, as is
obvious from Fig. 1, the dobs value begins to level off at mole ratios greater than unity,
because after formation of a 1:1 complex almost all of the Li? ions are present as the LiL
complex (PLi = 0 and PLiL = 1 so that dobs = dLiL). Figure 1 shows that the slope of the
corresponding mole ratio plots changes significantly at the point where the ligand-to-cation
mole ratio is equal to one, indicating the formation of a relatively stable 1:1 complex in
solution. It is interesting to note that, in NM solutions (Figs. 2b, 3b), the inflection points of
the mole ratio plots are clearly sharper than those obtained in AN solutions (Figs. 2a, 3a),
indicating the higher stability of the resulting complexes in the NM solvent.
The formation constants of 1:1 (LiL)? complexes were calculated from the variation of
the 7Li chemical shift with the L/Li? mole ratio. It has been shown previously [31] that the
observed chemical shift of Li? at a fixed concentration, CLi, upon titration with varying
ligand concentrations, CL, is given by:
ÈÂ
dobs
¼
ðKLiCLi ꢁ KLi ꢁ KLiCL ꢁ 1Þ þ ðKL2iCL2 þ KL2iCL2i ꢁ 2KL2iCLCLi
þ 2KLiCL þ 2KLiCLi þ 1Þ ꢂ ꢃ ðdLi ꢁ dLiꢁLÞ=2KLiCLig þ dLiꢁL
ð1Þ
1=2
where dLi and dLi–L are the respective chemical shifts of the free and complexed Li? ion.
The nonlinear least-squares curve fitting program KINFIT [32] was used to evaluate the
7
KLi and dLi–L values. A sample computer fit of the Li chemical shift–mole ratio data is
shown in Fig. 4 and the calculated log10 Kf values are given in Table 1. The fair agreement
between the observed and calculated chemical shifts shown in Fig. 4a further supports the
formation of a complex between Li? and amines with 1:1 stoichiometry.
It is well known that, most transition metals nuclei cannot be accurately used for an
NMR study of their macrocyclic complexes. This is mainly because of their high quadruple
moments and low receptivities [33]. Thus, in this work, we decided to employ 7Li NMR as
a very sensitive probe [34–38] to investigate the complex formation of Mg2?, Zn2? and
Cd2? ions with L1 and L2 in the solvents used. In this case, known concentrations of each
of the M2? ions and Li? (i.e., 0.01 molꢀL-1), present in solution, are simultaneously
titrated with each of ligands L1 and L2, while monitoring the changes in the chemical shift
7
7
of the Li nucleus in solution. The resulting plots of variation of the Li chemical shift in
the presence of equimolar concentrations of M2? ions and Li? (0.01 molꢀL-1) as a
function of the L/Li? mole ratio in AN and NM solutions are shown in Figs. 2 and 3. As is
obvious from Figs. 2 and 3, unlike the cases involving in the formation of a 1:1 Li?–L
complex in which the change in chemical shift with the L/Li? mole ratio is quite linear at
mole ratio \1, here the mole ratio plots in the presence of other M2? ions used show a
small change in chemical shift of Li? ion at L/Li? mole ratios between 0 and 1, followed
by a sharp linear increase in dobs, which begins to level off at mole ratios[2. More or less
similar dobs–L/Li? mole ratio behavior observed for all three Mg2?, Zn2? and Cd2? ions in
the presence of Li? ion clearly indicate that all 1:1 M2?– L complexes formed in both AN
and NM are more stable than their corresponding Li?–L complex. Here the resulting mole
ratio plots show that, due to higher stability of M2?–L complexes over Li?–L, the majority
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Fig. 2 Lithium-7 chemical shifts as a function of the L/Li? mole ratio in acetonitrile and nitromethane
solutions
of ligand molecules interact with M2? ions at L/Li? mole ratios\1 so that most of the Li?
ions present remain free in solution and, consequently, the 7Li chemical shift changes only
slightly. However, at mole ratios [1, where the number of free M2? ions in solution has
largely diminished, the Li? ions present begin to form their 1:1 complexes with added
ligand molecules and, thus, the dobs–mole ratio behavior at mole ratios [1 shows a trend
more or less similar to that observed in the case of titration of Li? ion alone with the ligand
(see Fig. 1), and levels off at mole ratios [2. This indicates quantitative formation of 1:1
complexes of the ligands with both Li? and all three M2? ions studied.
Another point which worth noting in Figs. 2 and 3 is that, in all cases, the addition of nitrate
salts of Mg2?, Zn2? and Cd2? ions to a solution containing equimolar amounts of LiClO4
causes some changes of about 0.1 to about 0.9 ppm in the 7Li NMR chemical shift of free Li?
ions at the L/Li? mole ratio = 0, as compared with the case of the lithium salt alone. This is
most probably due to increasing the ionic strength of the solution by 100 % and, conse-
quently, in increasing the extent of ion pairing of the Li? ion with the counter anions present.
The competitive complexation equilibria for the case of 1:1 complexation between
amines and Li? and proposed ions can be written as:
Liþ þ L ꢀ LiþL KLi ¼ ½LiþL]/[Liþꢂ½Lꢂ
ð2Þ
ð3Þ
M
2þ þ L ꢀ M2þ
The free amine concentration is then obtained from Eq. 4 [39]:
L
KM ¼ ½M2þL]/[M2þꢂ½Lꢂ
3
2
KLiKM½Lꢂ ꢁ fKLiKMðCL ꢁ CLi ꢁ CMÞ ꢁ KLi ꢁ KMg½Lꢂ
ð4Þ
ꢁ fKLiðCL ꢁ CLiÞ þ KMðCL ꢁ CMÞ ꢁ 1g½L ꢁ 1ꢂ½Lꢂ ꢁ CL ¼ 0
where CL, CLi, and CM are the analytical concentrations of the amine, Li? and M2? ion,
respectively. In this case, the observed [7Li] NMR chemical shift is obtained from Eq. 5 [35]:
dobs ¼ PLidLi þ PLiLdLiL
ð5Þ
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1225
Fig. 3 Lithium-7 chemical shifts as a function of the L1/Li? mole ratio in acetonitrile (a) and nitromethane
(b) solutions in the absence and presence of different M2? ions
where PLi and PLiL are the mole fractions of the free and complexed lithium ion,
respectively. By substitution from Eq. 2 and the mass balance equation CLi
[Li?] ? [Li?L], Eq. 5 can be written as:
=
dobs ¼ fdLi þ dLiL½LꢂKLig=f1 þ KLi½Lꢂg:
ð6Þ
The KM values were evaluated by obtaining the free amine concentration from Eq. 4
and fitting the chemical shift–mole ratio data to Eq. 6 by using the KINFIT program.
A sample computer fit of the data is shown in Fig. 4b and the resulting log10 Kf values
are included in Table 1. Our assumption of 1:1 stoichiometry for both Li?–L and M2?–L
complexes seems reasonable in the light of a fair agreement between the observed and
calculated chemical shifts.
The apparent stability constants of the complexes formed between the two ligands and
all metal ions used in AN and NM, as solvents of much lower dielectric constants than
water, are expected to be significantly affected by both the nature of the counter anions and
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Table 1 Formation constants for metal ion complexes with L1 and L2 in AN and NM solutions at 25 °C
Cationa
log10 Kf
Acetonitrile
A1
Nitromethane
A1
A2
A2
Li? (0.76)
Mn2? (0.83)
Zn2? (0.74)
Cd2? (0.95)
a
2.65 0.05
3.97 0.05
4.32 0.09
4.93 0.32
2.49 0.04
3.85 0.06
3.99 0.05
4.28 0.10
3.87 0.04
4.45 0.07
4.87 0.06
5.58 0.20
3.36 0.04
4.35 0.09
4.48 0.02
5.29 0.17
˚
Values in parentheses are the ionic sizes of cations in A
the concentrations of the salts used, due to variations in extent of ion pairing between the
metal ions and their counter anions. Thus, in this work, we used the same salt concentration
of 0.01 molꢀL-1 with the same counter anion, in order to have a constant effect of salt
concentration and counter anion on all apparent stability constants of the complexes of
interest.
The results given in Table 1 clearly demonstrated that, in both NM and AN solutions
and for all four metal ions studied, the resulting complexes with the Schiff bases L1 and L2
vary in the order of ML1 [ ML2. This is due to the fact that the ligands L1 and L2 differ in
the size of their central chelate rings so that they can form five- and six- membered rings,
respectively (see Fig. 5). As it is well documented in the literature [40, 41], the complexes
forming a six-membered chelate ring are expected to be less stable than those forming a
five-membered chelate ring.
In fact, N1,N10-(ethane-1,2-diyl)bis(N1-(pyridin-2-ylmethyl)propane-1,3-diamine) (L1)
and N1,N10-(propane-1,3-diyl)bis(N1-(pyridin-2-ylmethyl)propane-1,3-diamine) (L2) are
both hexadendate (N6) amine ligands containing two pyridine moiety groups at terminal
positions. It has been shown that the complexation ability of the amines to metal ions can
be considerably increased by attaching a more rigid carbon chain between the donor
nitrogen groups [43–45]. In the cases of ligands L1 and L2, the stability of the resulting
metal ion complexes arises from the two types of nucleation centers: (1) the amine
nitrogens of the carbon chain and (2) the pyridine moiety residues, which control the
strength of complexation by p-electron interactions, steric influences and by supplying two
terminal nitrogen atoms [45–47]. Subsequently, the metal ion is surrounded by the flexible
carbonic chain, which can easily adapt to different cation sizes, and is better shielded from
the solvent and the counter ions by the two terminal pyridine groups [46–49]. Therefore,
as is obvious from Table 1, in all cases the Mn?–L1 complexes are more stable than the
M
n?–L2 complexes in the solvents.
The data given in Table 1 also illustrate the influence of solvent properties on the
stability of the resulting complexes between ligands L1 and L2 and Li? and all three M2?
metal ions used. It is obvious that the nature of solvent plays a fundamental role in complex
formation; in all cases, the complexes are more stable in NM than in AN. It should be
noted that, while NM and AN have about the same dielectric constants, their solvating
abilities as expressed by the Gutmann donor number [42], are quite different (i.e., e = 35.9
and DN = 2.7 for NM and e = 36.1 and DN = 14.1 AN). AN as a solvent of higher
solvating ability that can compete more strongly with amines for the metal ions than NM as
a poor solvating solvent and, thus, the complexes formed in NM are expectedly to be more
stable than those in AN.
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Fig. 5 Computer fit of lithium-7 chemical shift versus the L1:Li? mole ratio for Li?–L2 (a) and Li?–L2–
Mn2? (b) systems in acetonitrile solution: (x) experimental point; (o) calculated point; (=) experimental and
calculated points are the same within the resolution of the plot
The data given in Table 1 reveal that, in both solvents used, the stability of the resulting
complexes varies in the order Cd2? [ Zn2? [ Mn2? [ Li?. The complex stability order
results from the superposition of several factors including the extent of interaction of donor
groups of the ligand with the cation, the consonance between the size of metal ion and the
steric size of the ligand cavity formed, the extent of ligand conformational changes as a
consequence of complex formation, the hard-soft acid–base characters of the metal ion and
the donating groups of the ligand, desolvation of the ligand and cation and solvation of the
resulting complex, the two later factors being strongly dependent on the nature of solvents
used.
The cationic charge and size (and the consequent charge density) can play a funda-
mental role in the complexation process from the viewpoints of ionic solvation and binding
energy with the ligand’s donating atoms. By increasing the charge density of a cation, both
its solvation energy and binding energy to a ligand are expected to increase, so that the
increased binding energy of the cation to the ligand acts against its increased desolvation
energy. On the other hand, since the change in topology of the open chain ligand from a
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J Solution Chem (2014) 43:1218–1231
1229
flexible free state in solution to a helical rigid conformation in the complexed form leads to
a large loss of entropy, the decrease in cation size and increase in charge density may
induce steric deformations of the ligand and reduce the entropy of complexation.
As is obvious from Table 1, among the different cations used, the Li? ion forms the
least stable complexes with ligands L1 and L2 in both solvent systems. In fact, the lithium
ion is a single charged hard acid that is expected to have weak interactions with the
intermediate–soft amine and pyridine nitrogen’s of the ligand molecules. On the other
hand, the bivalent transition metal ions form stronger complexes with both ligands than
does the Li? ion. This is due to the soft character of the transition metal ions, which result
in stronger interactions with the donating nitrogen’s of the ligands [50].
Meanwhile, the sequence of stability of the Zn2? and Mn2? complexes with both
ligands (i.e., Zn2? [ Mn2?) follows the Irving–Williams order [51], which generally
governs the equilibrium constants of first row transition metal ions. Finally, the results
show the highest stability for Cd2? complexes with ligands L1 and L2, among all three M2?
cations used. This is most probably due to the proper size of the large Cd2? ion with soft
nature which could favor a suitable spatial fit with flexible donating atoms of the ligands
(i.e., the four amine- and two pyridine-nitrogens of L1 and L2) so that, in this case, the
central metal ion induces the lest amount of steric deformation to the ligands’ structures
[52, 53].
Acknowledgments We are grateful to the Faculty of Chemistry of Bu-Ali Sina University, Ministry of
Science, Research and Technology of Iran for financial support.
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