Inorganica Chimica Acta 356 (2003) 133�/141
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Tetraazamacrocycle bearing quinoline pendant arms and itscomplexation properties
Xiuling Cui a, M. Fatima Cabral b, Judite Costa a,b, Rita Delgado a,c,*a Instituto de Tecnologia Quımica e Biologica, UNL, Apartado 127, 2781-901 Oeiras, Portugal
b Faculdade de Farmacia de Lisboa, Av. Prof. Gama Pinto, 1649-003 Lisbon, Portugalc Instituto Superior Tecnico, Av. Rovisco Pais, 1049-001 Lisbon, Portugal
Received 5 March 2003; accepted 16 June 2003
In honor of Prof. J.J.R. Frausto da Silva
Abstract
A new tetraazamacrocycle containing two 5-chloro-8-hydroxyquinoline side arms, L1, was synthesized. The protonation
constants of L1 and the stability constants of its complexes with Cu2�, Zn2�, Cd2� and Pb2� were determined by potentiometric or
spectrophotometric methods, at 20 8C in methanol:water (5:1, v/v) and ionic strength 0.10 mol dm�3 in tetrabutylammonium nitrate
(TBAN). Compound L1 has three high protonation constants, the most basic centres being both phenolates of the arms. Mono- and
dinuclear complexes were found in solution for all the metal ions, the complexes of copper are thermodynamically very stable, and
the dinuclear complexes of lead(II) are the predominant species even in 1:1 mixtures. UV�/Vis spectra of some complexes and EPR
of the copper ones confirmed the presence of dinuclear complexes.
# 2003 Elsevier B.V. All rights reserved.
Keywords: Tetraazamacrocycle; Stability constants; Dinuclear complexes; EPR copper complexes; 8-Hydroxyquinolinyl groups
1. Introduction
Macrocycles containing quinoline side arms have
found many applications as analytical reagents in
absorption spectrophotometry, fluorometry, solvent ex-
traction, and partition chromatography [1,2]. These
compounds are also used as pesticides [2].
The 8-hydroxyquinoline (HQ) forms a stable five-
membered chelate ring with metal centres and very
stable complexes with divalent transition and post-
transition metal ions of the type M(HQ)�, M(HQ)2 or
M(HQ)3� [3]. The attachment of two rigid 5-chloro-8-
hydroxyquinoline (CHQ) groups to diazamacrocycles
* Corresponding author. Tel.: �/351-214-46-97-37/38; fax: �/351-
214-41-12-77.
E-mail address: [email protected] (R. Delgado).
0020-1693/03/$ - see front matter # 2003 Elsevier B.V. All rights reserved.
doi:10.1016/S0020-1693(03)00465-1
results in an appropriate pre-organization of the mole-
cule increasing the selectivity and, in general, modifying
the co-ordination properties of the macrocycle to metal
ions [2,4�/7]. It is known that the phenol ring of CHQ is
electron-rich and the pyridine moiety electron-deficient
[4]. Therefore, if two CHQ rings approach to each other,
p�/p interactions will occur by overlapping of the phenol
moiety of one CHQ with the pyridine group of the
other, as already observed by Izatt and co-workers [4].
The p�/p interactions can be induced by co-ordination to
a metal ion so that a pseudo-second ring is formed,
resulting in a cryptate-like structure. In such a case, the
match between the cation size and the dimensions of the
pseudo-cavity formed is crucial [4]. In addition, the
CHQ molecule allows different positions to be appended
as N-substituents, by 2- or 7-positions, offering the
phenol or the pyridine as closer donor atom to the metal
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141134
centre with consequent different affinity of the resulting
molecule to metal ions. In most of the studies having
CHQ as substituents of macrocycles, this group is
appended by the 2-position [2,4�/7], and very few by
the 7-position [5,6]. In most cases, the CHQ groups were
attached to macrocycles containing oxygen donor
atoms, and consequently an special attention has been
given to the complexation of alkali and earth alkaline
metal ions, with few exceptions [2,7].
The extensive applications of quinoline derivatives
make them a very attractive field of research. We are
especially interested to obtain compounds, which could
form selective complexes with transition and post-
transition metal ions to be used as spectrophotometric
sensors for these metal ions. In the present work, we
explore CHQ groups appended by the 7-position to a
14-membered tetraazamacrocycle containing pyridine,
L1 (see Scheme 1). The 7-position CHQ derivative was
chosen because the resulting ligand has more affinity to
transition metal ions [5]. To understand the role of the
pyridine moiety of the CHQ groups, 5-bromophenol
groups were also appended to the same macrocycle, L3
[8]. The CHQ has a larger p electron area when
compared with the phenols and has the additional
nitrogen donor atom of the pyridine group then CHQ,
attached to macrocycles as side arms, can induce the
formation of di- or polynuclear complexes.
Scheme
2. Experimental
2.1. Reagents
The parent macrocycle 7-methyl-3,7,11,17-tetraazabi-
cyclo-[11.3.1]heptadeca-1(17),13,15-triene (L2) was
synthesized by previously reported procedures [9,10].
Compound 7,16-bis((5-chloro-8-hydroxy-7-quinoli-nyl)methyl)-1,4,10,13-tetraoxa-7,16-diazacyclooctade-
cane (L4) was synthesized from the parent macrocycle
(Aldrich) following the described procedure [6]. All the
chemicals were of reagent grade and used as supplied
without further purification. Microanalyses were carried
out by the ITQB Microanalytical Service. IR spectra
were recorded from KBr pellets on a UNICAM
Mattson 7000 spectrometer.
2.1.1. Synthesis of 7-methyl-3,11-bis((5-chloro-8-
hydroxy-7-quinolinyl)methyl)-3,7,11,17-
tetraazabicyclo[11.3.1]heptadeca-1(17),13,15-triene
(L1)
An anhydrous toluene solution (60 cm3) of L2 (0.5 g, 2
mmol), p -formaldehyde (0.12 g, 4 mmol), and CHQ
(0.72 g, 4 mmol) were refluxed for 2 days. After
evaporation of the solvent under reduced pressure, the
residue obtained was purified by silica column chroma-
tography using CHCl3/CH3OH (20:1, v/v) as eluent. The
1.
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141 135
first fraction was CHQ and the second fraction eluted
gave a yellow powder after evaporation of the solvent,
which was verified to be the desired macrocycle (L1).
Yield: 0.3 g, 23.6%. m.p. 85�/86 8C. IR (KBr pellet):3413, 2945, 2810, 2360, 1622, 1591, 1496, 1456, 1385,
1263, 1153, 1111, 1045, 939, 789, 752 cm�1. 1H NMR
(CDCl3): d 1.41 (4H, q, NCH2CH2CH2N), 1.69 (3H, s,
NCH3), 2.16 (4H, bs, NCH2CH2), 2.77 (4H, t, J�/7.3
Hz, NCH2CH2), 3.48 (2H, s, OH ), 3.87 (4H, s, NCH2�/),
4.04 (4H, s, NCH2�/), 7.25 (2H, d, py), 7.41 (2H, Q), 7.49
(2H, m, Q), 7.50 (2H, t, py), 8.47 (2H, d, Q), 8.95 (2H, d,
Q) ppm. 13C NMR (CDCl3) and dept: d 23.7(NCH2CH2CH2N), 42.7 (CH3), 50.7, 54.1, 57.5, 59.2,
119.3 (C �/Cl), 119.9 (Q), 121.9 (py), 123.1 (Q), 126.1
(Q), 127.5 (Q), 132.7 (Q), 137.5 (py), 139.9 (C �/OH, Q),
149.3 (Q), 152.4 (Q), 156.9 (py) ppm. Found: C, 64.4; H,
5.0; N, 13.1. Calc. for C34H36Cl2N6O2: C, 64.66; H, 5.14;
N, 13.31%.
2.2. Potentiometric measurements
2.2.1. Reagents and solutions
Metal ion solutions were prepared at about 0.05 mol
dm�3 from the nitrate salts of the metals, and were
standardized as described [11]. Carbonate-free solution
of the titrant, KOH, was freshly prepared in methanol:-
water (5:1, v/v) solution, maintained in a closed bottle,
and was discarded when the percentage of carbonate
was about 0.5% of the total amount of base (tested byGran’s method) [12]. Perchloric acid solution (0.1 mol
dm�3) was prepared from a Fixanal ampoule in
methanol. The support electrolyte, tetrabutylammonium
nitrate (TBAN), was added to the solutions as a solid.
The ligand L1 was dissolved in methanol:water (5:1), in
concentrations of about 1.5�/10�3 mol dm�3. All the
solutions were prepared in a mixture of methanol:water
(5:1). This solvent was suggested from the solubility ofthe ligand. Methanol was purified by standard methods
[13] and water used was demineralized water obtained
by a Millipore/Milli-Q system.
2.2.2. Equipment and work conditions
The equipment used was described before [10,14]. The
glass electrode was pre-treated by soaking in solutions
of methanol:water of increasing rate of methanol over aperiod of 10 days, in order to prevent erratic responses.
The temperature was kept at 20.09/0.1 8C; atmospheric
CO2 was excluded from the cell during the titration by
passing purified argon across the top of the experi-
mental solution in the reaction cell. The ionic strength of
the solutions was kept at 0.10 mol dm�3 with TBAN.
2.2.3. Measurements
The [H�] of the solutions was determined by the
measurement of the electromotive force of the cell, E�/
E ?o�/Q log[H�]�/Ej . The term pH is defined as �/
log[H�]. E ?o, Q , Ej and K ?w were obtained as described
previously [10,14]. The value of K ?w; ionic product of
water in the methanol�/water mixture, used was found
equal to 10�14.40 mol2 dm�6 in our experimentalconditions, in agreement with that determined by
Rochester [15].
The potentiometric equilibrium measurements were
made on 20.00 cm3 of $/1.50�/10�3 mol dm�3 ligand
solutions diluted to a final volume of 30.00 cm3, in the
absence of metal ions and in the presence of each metal
ion for which the CM:CL ratios were 1:1 and 3:1. A
minimum of two replicates was performed. A correctionwas made for the small decrease in volume that occurs
on mixing methanol and water. Care has been taken to
maintain unaltered the methanol/water ratio in the
measured solution and the temperature was kept at
20 8C to prevent the evaporation of the solvent.
2.2.4. Calculation of equilibrium constants
Overall protonation constants bHi were calculated by
fitting the potentiometric data obtained for the free
ligand to the HYPERQUAD program [16].
bHi �
[HiL]
[L] � [H]i
Differences between two consecutive log bHi values
provide the stepwise protonation constants KHi :/
Stability constants of various species formed in
solution were obtained from the experimental data
corresponding to the titration of solutions of differentmetal:ligand ratios, also using the HYPERQUAD
program. The initial results were obtained in the form
of overall stability constants, bMm Hh Llvalues:
bMmHhLl�
[MmHhLl ]
[M]m � [H]h � [L]l
Mononuclear species, ML, MHiL (i�/1�/4) and also
dinuclear ones, M2L, M2HjL (j�/1�/2), M2H�1L and
M2(H�1)2L, were found for the metal ions studied,
being bM2(H�1)L�/bM2LOH�/KW and bM2(H�1)2L�/
bM2L(OH)2�/(KW)2. The species considered in a particu-
lar model were those that can be justified by the
principles of co-ordination chemistry. Differences, in
log units, between the values bMm (Hi L) [or bM2(H�1)j L]
and bMm (Hi�1L) [or bM2(H�1)j�1L] provide the stepwise
protonation reaction constants, as shown in Table 1.
The errors quoted are the standard deviations of the
overall stability constants given directly by the program
for the input data, which include all the experimental
points of all titration curves.
The protonation constants were obtained from 120
experimental points (two titration curves) and thestability constants for each metal ion were determined
from 70 to 120 experimental points (two or three
titration curves). First, each titration curve of a given
Table 1
Protonation constants of L1 and stability constants of its metal complexes
Metal ion Species m h l log bMm
Hh
Ll
L1 Equilibrium quotient log KMm
Hh
Ll
L1 log KMm
Hh
Ll
L4 a
H� 0 1 1 12.42(1) [HL]/[H]�/[L] 12.42 12.1
0 2 1 22.77(3) [H2L]/[HL]�/[H] 10.35 8.15
0 3 1 31.71(4) [H3L]/[H2L]�/[H] 8.94
0 4 1 37.80(6) [H4L]/[H3L]�/[H] 6.09
0 5 1 41.89(6) [H5L]/[H4L]�/[H] 4.09
0 6 1 44.36(9) [H6L]/[H5L]�/[H] 2.47
Cu2� 1 0 1 23.51(9) [ML]/[M]�/[L] 23.51 10.1
1 1 1 32.07(4) [MHL]/[ML]�/[H] 8.56
1 2 1 37.71(1) [MH2L]/[MHL]�/[H] 5.64
1 3 1 40.48(2) [MH3L]/[MH2L]�/[H] 2.77
2 0 1 32.79(4) [M2L]/[ML]�/[M] 9.28
2 1 1 36.86(3) [M2HL]/[MHL]�/[M] 4.07
2 �/1 1 26.14(5) [M2L]/[M2L(OH)]�/[H] �/6.65
2 �/2 1 19.10(4) [M2L(OH)]/[M2L(OH)2]�/[H] �/7.04
Zn2� 1 0 1 19.43(5) [ML]/[M]�/[L] 19.43 5.12
1 1 1 28.01(5) [MHL]/[ML]�/[H] 8.58
1 2 1 33.51(4) [MH2L]/[MHL]�/[H] 5.50
1 3 1 38.03(3) [MH3L]/[MH2L]�/[H] 4.52
1 4 1 41.98(4) [MH4L]/[MH3L]�/[H] 3.95
2 0 1 26.63(5) [M2L]/[ML]�/[M] 7.20
2 1 1 32.87(3) [M2HL]/[M2L]�/[H] 6.24
2 2 1 37.57(4) [M2H2L]/[M2HL]�/[H] 4.70
Cd2� 1 0 1 18.57(5) [ML]/[M]�/[L] 18.57
1 1 1 27.74(5) [MHL]/[ML]�/[H] 9.17
1 2 1 34.64(4) [MH2L]/[MHL]�/[H] 6.90
1 3 1 39.38(1) [MH3L]/[MH2L]�/[H] 4.74
2 0 1 26.25(4) [M2L]/[ML]�/[M] 7.68
2 1 1 33.09(4) [M2HL]/[MHL]�/[M] 6.84
Pb2� 1 0 1 16.23(3) [ML]/[M]�/[L] 15.23
1 1 1 25.75(4) [MHL]/[ML]�/[H] 9.52
1 2 1 33.49(3) [MH2L]/[MHL]�/[H] 7.74
1 3 1 38.62(2) [MH3L]/[MH2L]�/[H] 5.13
1 4 1 42.77(2) [MH4L]/[MH3L]�/[H] 4.15
2 0 1 27.38(4) [M2L]/[ML]�/[M] 11.15
2 1 1 33.22(3) [M2HL]/[M2L]�/[H] 5.84
2 2 1 38.00(3) [M2H2L]/[M2HL]�/[M] 4.78
Values in parentheses are standard deviations in the last significant figures. T�/20.0 8C; I�/0.10 M in TBAN.a Methanol at 25 8C [5].
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141136
system was treated separately and finally the data of all
the titration curves were merged and treated simulta-
neously to give the values of stability constants.
2.2.5. Spectroscopic studies1H NMR spectra were recorded with a Bruker AMX-
300 spectrometer at probe temperature. The electronic
spectra of the complexes prepared by the addition of
increasing amounts of the metal ion (in the form of
nitrate or chloride salts) to the ligands at the appropriate
pH value were recorded with a UNICAM UV�/Visspectrophotometer model UV-4 or a Shimadzu model
UV-3100 spectrophotometer for UV�/Vis�/near IR, in
methanol:water (5:1, v/v) or in DMSO. One titration of
a solution of L1 with KOH was also followed by UV�/
Vis spectroscopy from pH 7 to 14, to determine the two
higher protonation constant values of this ligand. The
protonation constants were determined using pHab [17]
or HYPERQUAD2000 programs [16].
EPR spectroscopy measurements were recorded with
a Bruker ESP 380 spectrometer equipped with contin-
uous-flow cryostats for liquid helium or liquid nitrogen,
operating at X-band. The spectra of the copper(II)
complexes, 1.0�/10�3 and 2.2�/10�2 mol dm�3, were
recorded in the ranges 102�/130 and 7.5�/32 K, in
DMSO/water mixtures and in 1.0 mol dm�3 TBAN.
Computer simulations of the spectra were carried out
with a program for a microcomputer [18].
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141 137
3. Results and discussion
3.1. Synthesis of macrocycles
Compound L1 was prepared by the Mannich reaction,
reacting the parent macrocyclic amine L2 with p -
formaldehyde and the corresponding substituted CHQ
in a particular solvent, as described to append the samegroup to other macrocycles [6,19,20]. The reaction is
quite sensitive to the solvent used [19]. Although the
reaction ortho - to the OH of phenol is preferred to that
of the para -position, the latter was blocked with Cl, the
commercial available starting material, to avoid parallel
reactions. From the two possible ways of synthesis, the
previous preparation of 3,11-bis(methoxymethyl) deri-
vative of L2 and the one-pot reaction [6,8], the latter oneled to better yields, although quite low and impossible to
improve. Izatt and co-workers [7] attempted also to
prepare CHQ derivatives of tetraazamacrocycles using
the same technique but could not obtain the desired
product.
3.2. Acid�/base behaviour of L1
The acid�/base reactions of L1 have been studied by
potentiometric and/or spectrophotometric methods at
20.0 8C and in a mixture 5:1 methanol:water (v/v). Theresults are collected in Table 1. The temperature was
chosen to minimize the evaporation of the solvent.
From the eight basic centres of the molecule, four
amines of the macrocycle, two phenolate and two
pyridine centres of the arms, it was possible to determine
six protonation constants, the other two were very low
to be determined by the techniques used. The value of
the first protonation constant was determined byspectrophotometric titration and the second one con-
firmed by the same technique (see Fig. 1). The bands at
Fig. 1. Spectrophotometric titration curves of L1 in methanol:water
(5:1, v/v), at 20.0 8C in 0.1 mol dm�3 in TBAN, between pH 7.11 and
13.55: 1*/pH 7.11; 2*/pH 7.86; 3*/pH 8.67; 4*/pH 9.70; 5*/pH
10.26; 6*/pH 10.72; 7*/pH 11.45; 8*/pH 12.01; 9*/pH 12.77; 10*/
pH 13.10; 11*/pH 13.55.
250 and 335 nm at pH 8 disappear at pH about 10.5 to
give rise to the bands at 260, 345 and 380 nm which
increase with the pH. The same happens with the CHQ
itself, which exhibits, in the same experimental condi-tions, a band at 335 nm that is converted to the bands at
345 and 380 nm with the increase of the pH.
The most basic centres of the molecule are the
phenolate groups, for which the values of 12.42 and
10.35 can be ascribed. The third and fourth values, 8.94
and 6.09, correspond to the protonation of amine
centres of the macrocycle. The third one would corre-
spond to the protonation of the amine opposed to thepyridine of the macrocycle, and the protonation of the
second amine probably implies a rearrangement of
charges into the macrocyclic cavity with both proto-
nated amines trans to each other, as found by 1H NMR
titrations for compounds containing the same macro-
cycle but different side arms [21]. The next one, 4.09, can
be attributed to the protonation of the pyridine of one
quinoline group or the third amine of the macrocycle.The sixth and the others correspond to the protonation
of the remaining centres, which are very acidic due to
the presence of positive charges at short distance of
already protonated amines.
The protonation constants of 8-hydroxyquinoline are
9.63 and 4.95 in water (in log units) [3b] and, 11.30 and
4.33 in 75% ethanol [3b,22]. For CHQ, they are slightly
lower: 9.29 and 3.79 (in water) [3a]. On the other hand,the parent macrocycle, L2, presents the following values:
9.74, 8.67, 4.67 and B/1 in water [10], and slightly lower
in a mixture methanol:water (1:1, v/v): 9.60, 8.12, 3.95
and B/2 [23]. Then, the presence of the CHQ groups
decreases the basicity of the amine centres of the
macrocycle and, on the other hand, the presence of the
amines of the macrocycle at short distance increases the
basicity of the phenolate groups of CHQ. These effectsare well established for linear amines having N-o -
hydroxybenzyl substituents [24,25], and the same effects
were found for cyclic amines with the same type of
substituents [5,26]. It is assumed to be due to the
formation of intramolecular hydrogen bonding interac-
tions between the phenolic hydrogen and the amine
nitrogen. These intramolecular hydrogen bonding inter-
actions were found in the crystal X-ray diffractionstructures of L4 or L5, involving the hydroxy groups
of CHQ of the first compound or the phenol groups of
the second one and the nitrogen atoms of the macro-
cycle, resulting in the elongation of the macrocycle along
the N� � �N direction. On the other hand, the same
macrocycle with appended CHQ groups bound by the
2-position, where the pyridine moiety is closer to the
nitrogen of the ring, does not exhibit intramolecularhydrogen bonding, but intermolecular ones between the
hydroxyl groups that point away from the macrocycle
and a similar group of other molecule [5]. The intramo-
lecular hydrogen bonds stabilize the deprotonated form
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141138
of the amine contributing to decrease its basic character,
and, at the same time, they stabilize the phenol form
conferring a more basic character to the phenolate. It
can be inferred that the same effects should occur withL1 taking into account the relative position of the
phenol group in CHQ substituents to the nitrogen
atom of the macrocycle (see Table 1).
Compound L3 was synthesized as a model for the
study of L1, but it was impossible to find a common
medium for the study of both. However, we believe that
they should present rather similar behaviour in solution,
although the basicity of the phenolate groups of L3
would be slightly higher and its metal complexes would
be thermodynamically less stable [5].
Fig. 3. Species distribution curves calculated for the Cu(II):L1
complexes in methanol:water (5:1, v/v): (a) 1:1 ratio, CL�/CM�/
8.333�/10�4 mol dm�3; (b) 2:1 ratio (M:L), CL�/8.333�/10�4 mol
dm�3 and C �/1.666�/10�3 mol dm�3.
3.3. Metal complexes studies
The stability constants of L1 with Cu2�, Zn2�, Cd2�
and Pb2� have also been studied by potentiometric and/
or spectrophotometric methods in the conditions used
for the determination of the protonation constants. The
results are also listed in Table 1.Compound L1 forms in solution mono- and dinuclear
complexes, each of them presenting several protonated
species, with all the metal ions studied. For the Cu2�
complexes, two hydroxocomplexes, M2L(OH)1 or 2 were
also found at pH�/5.5 in the 2:1 ratio (metal:ligand).
Fig. 2. Species distribution curves calculated for the Pb(II):L1 com-
plexes in methanol:water (5:1, v/v): (a) 1:1 ratio, CL�/CM�/8.333�/
10�4 mol dm�3; (b) 2:1 ratio (M:L), CL�/8.333�/10�4 mol dm�3
and CM�/1.666�/10�3 mol dm�3.
M
With the other metal ions studied, the hydroxocom-
plexes could not be found due to precipitation of the
ML complex, which is a neutral species, and conse-
quently less soluble in the polar solvent mixture used.
The ML1 complexes (1:1 ratio) follow the usual trend of
thermodynamic stability: CuL1�/ZnL1�/CdL1�/PbL1,
but an unexpected order of stability constants was found
for the dinuclear complexes: [Pb2L1]2��/[Cu2L1]2��/
[Cd2L1]2��/[Zn2L1], revealing a clear tendency of Pb2�
to form dinuclear species. The [Pb2L1]2� species not
only appear in the 2:1 ratio, but already in the 1:1 ratio
starting to form at pH 5 and having a maximum
percentage at pH:/7.5 (see the speciation diagram of
Fig. 2). For the complexes of the other metal ions with
the same ligand, L1, the dinuclear species only appear in
the 2:1 ratio, as can be seen in Fig. 3 for the copper(II)
complexes, which exhibits a behaviour similar to that
found for the other metal ions.
It is interesting to emphasize that in the 1:1 ratio, the
completely deprotonated complexes only form at high
pH values (pH�/7 for the copper and �/8 for the lead
complexes), implying a very basic centre on the ligand,
the protonation constant being 8.56 for Cu2�, 8.58 for
Zn2�, 9.17 for Cd2� and 9.52 for Pb2� complexes (in
log units) (see Table 1). However, an excess of metal ion
(2:1 ratio) leads to the complete deprotonation of the
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141 139
complex at much lower pH due to the formation of
dinuclear complexes (see Figs. 2 and 3).
It is reasonable to think that the remaining proton on
the MHL species is bound to the oxygen of one
phenolate, which is the most basic centre of the ligand,
and supported by the crystal X-ray structures of the
monoprotonated nickel and zinc complexes of the model
compound L3, [M(HL3)]� with M�/Cu2� and Zn2�
[8]. More difficult to rationalize is the relative position
of the two metal ions in the dinuclear complexes,
although the formation of these species at low pH
values implies the involvement of the CHQ pendent
arms.
The formation of the dinuclear species was also
confirmed by UV�/Vis spectroscopy of the copper and
lead complexes and by EPR spectroscopy of the copper
ones.
The ligand L1 exhibits an UV�/Vis spectrum with
bands at 380, 345 and 260 nm (Fig. 1). Upon addition of
metal ions, the peaks shift or change their absorptivity,
the largest shift was verified for the copper complex.
Upon addition of increasing amounts of a Cu2�
solution, the bands of the ligand decrease and those of
the CuL1 complex at 628 (170), 415 (1515), 345 (1890)
and 270 (11870) nm appear. Further addition of the
Cu2� solution leading to [Cu2L1]2� shift two of the
previously formed bands to 410 (2214) and 280 (11 720)
nm [the values given correspond to lmax (nm) (omolar
(dm3 mol�1 cm�1))]. The electronic spectra of the Zn2�
and Pb2� complexes of L1, and the Cu2� and Zn2�
complexes of L4 were also carried out in methanol and
in all cases the maximum of absorption bands of the ML
complexes displayed shifts similar to those described for
the Cu2� complexes of L1 (5�/35 nm) when the dinuclear
complexes were formed.
The EPR spectra of the copper complexes of L1 and
L3 were performed in water and in a mixture of
water:DMSO (1:1) in the 1:1 and 2:1 M:L ratios. For
L4 only the 2:1 spectrum was recorded, in the same
conditions. The spectra in both solvents are identical,
but better defined in the mixture of solvents. In the case
of 1:1 metal to ligand ratio, the spectra of the L1 and L3
complexes are similar, and already described for L3 [8]
Table 2
Spectroscopic EPR data for the Cu2� complexes of L1, and other similar co
Complex lmax (nm) Species EPR Ai (�/104 cm
gx gy
CuL1 628 A 2.049 2.110
B 2.037 2.057
CuL3 614 A 2.049 2.068
B 2.050 2.060
[CuL2]2� 560 2.034 2.060
[CuL6] 630 2.027 2.084
(see Table 2). The spectra indicate the presence of two
species (A and B), as easily observed by the splitting of
the bands at low field. Taking into account the pH of
the solutions, the speciation diagram suggests that the
signals can be attributed to [Cu(HL1)]� and [CuL1]
species. Each species exhibits three well-resolved lines at
low field without superhyperfine splitting, and the
strong and not resolved band of the high field part of
the spectra overlaps the fourth line. The hyperfine
coupling constants (A) and g values of glassy solutions
of these complexes, obtained by simulation of the
spectra [18], are collected in Table 2. The parameters
of both species are typical of copper(II) complexes in
rhombic symmetry axially elongated and a dx2�y2
ground state, consistent with elongated rhombic-octa-
hedral or distorted square-based pyramidal stereoche-
mistries [27,28]. These parameters, together with the
position of the d�/d absorption band, of both species of
CuL1 indicate penta- or hexaco-ordination environ-
ments for the complexes. Complex B exhibits para-
meters and also the visible band maximum value similar
to those of [CuL6] (see Table 2), which adopts an
octahedral structure where the four nitrogen atoms of
the macrocycle determine the equatorial plane and the
six-co-ordination is completed via the two oxygen atoms
of the carboxylate arms [30]. Therefore, a similar
structure should be expected for B, suggesting that this
is the species [CuL1]. However, the gz value for the A
species is higher than suggested for a square-pyramidal
geometry and the Az value higher than expected for an
octahedral arrangement, then important distortions of
these geometries can be ascribed to this copper(II)
complex, which should be ascribed to the monoproto-
nated complex [Cu(HL1)]� [8].
The increase of the amount of copper(II) to 2:1 ratio
(metal to ligand) in solutions of CuL1 leads to an EPR
spectrum which can be interpreted as that of a triplet
state of a coupled dinuclear copper centre with an S�/1
state and zero-field splitting, having two types of signals
in the DMs�/1 and DMs�/2 regions (see Fig. 4a). The
hyperfine structure corresponding to the first transition
is only observed at temperatures of �/100 K and diluted
solutions (of about 1.1�/10�3 mol dm�3), while the
mplexes
�1) Reference
gz Ax Ay Az
2.236 13.5 20.3 173.5
2.210 19.8 31.8 166.9 This work
2.229 13.6 13.5 178.1
2.194 3.0 11.3 163.8 [8]
2.188 0.5 3.4 192.9 [29]
2.221 14.9 21.3 165.4 [30]
Fig. 4. EPR X-band spectra of the Cu2� complexes of L1 (a), of L3 (b)
and of L4 (c), in the 2:1 M:L ratio in H2O:DMSO (1:1). The high-field
part of the spectra were recorded at 114 K, frequency (n ) 9.408 GHz
for (a), at 102 K, frequency (n ) 9.643 GHz for (b) and at 102 K,
frequency (n ) 9.404 GHz for (c). The low-field part of the spectra were
recorded at 9 K, frequency (n ) 9.643 GHz for (a) and (b), and at 8 K,
frequency (n ) 9.644 GHz for (c). Microwave power of 2.4 mW and
modulation amplitude of 0.9 mT for all complexes. Concentration of
the complexes: 2.00�/10�3 mol dm�3.
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141140
signal at DMs�/2 transition is only observed at tem-
peratures lower than 30 K and concentrations of about
2�/10�2 mol dm�3. Similar signals, but with hyperfine
structures better defined, were observed for the 2:1
solutions of copper complexes of L4 (see Fig. 4c).
However, the spectrum of corresponding solutions of
Cu2�:L3 (2:1 ratio) recorded in the same conditions
only exhibits a signal of very low intensity in the DMs�/
2 transition at 9 K which disappear at 19 K (see Fig. 4b).
The DMs�/2 transition signal reveals a spin exchange
interaction between the two copper atoms. If assumed
that two unpaired electrons, each located on a Cu2�, are
exchanged with an energy larger than the hyperfine
coupling energy, each unpaired electron (/S� 12) couples
with two copper nuclei (/ICu�32) in an equivalent
manner, the hyperfine structure consists of seven lines.
Indeed, seven hyperfine lines can be observed in the
DMs�/2 region of the L4 complexes, with an average
hyperfine spacing of 95.1 cm�1. The hyperfine constant
is approximately half the value found in mononuclear
copper(II) complexes having the same donor atoms and
geometry.
The triplet state (S�/1) involving the coupling of two
copper(II) ions can be represented by the spin Hamilto-
nian, assuming that the system is axially symmetric:
H�gIbHzSz�g�b(HxSx�HySy)�D[S2z �
13S(S�1)]
where D describes the zero-field splitting constant
within the triplet state and can be considered to be built
up from a dipolar Ddip and an exchange, Dex contribu-
tions:
D�Dex�Ddip
where
Dex��1
8J
�1
4(gI�2)2�(g��2)2
�and
Ddip��(g2
I � 12g2�)b2
r3
where �/J is the energy separation between the singlet
ground state and the first excited triplet state, and r is
the Cu�/Cu distance, which appears in the dipole�/dipolecontribution [31�/33].
The J values were determined by the signal intensity
of the EPR spectra at different temperatures by the
expression: I�/c{T [3�/exp(�/J /kT )]}, where I is the
spectral intensity of the DMs�/2 signal, which is
proportional to the paramagnetic susceptibility of the
dinuclear system, c being the proportionality constant
and k the Boltzmann constant. As the linewidth did notdepend on the temperature, the relative intensity of the
triplet spectra is equal to the relative peak height of a
certain line in the first derivative spectra of the DMs�/2
transition [34]. The determined values were 5.03 cm�1
for the L1 complex and �/7.76 cm�1 for that of L4.
The 2D value was calculated by the difference
between the Hz2and Hz1
transitions, respectively
[32,35], yielding to 0.0185 and 0.0262 cm�1 for L1 andL4 complexes, respectively. Taking gI�2:240 and g��/
2.082 for both complexes, obtained by simulation of the
DMs�/2 transitions signal of the L4 complex, and
considering the same values for the L1 complex, due to
the complete superimposition of the signals of both
complexes (see Fig. 4), the values of 6.04 and 6.12 A
were determined for the copper�/copper distance from
the above expression for the L1 and L4 complexes,respectively.
On the other hand, as the intensity of the DMs�/2
signal of the EPR spectrum of the L3 copper(II) complex
is very low, even at 9 K, then a different kind of
dinuclear species is probably formed.
4. Conclusions
Compound L1, which has a very high overall basicity,
forms in solution stable mono- and dinuclear metal
complexes. Unfortunately, it was not possible to obtain
X. Cui et al. / Inorganica Chimica Acta 356 (2003) 133�/141 141
crystals with adequate size for X-ray diffraction deter-
mination of metal complexes of L1. However, the crystal
structures of metal complexes of the model compound
L3, [Ni(HL3)]� and [Zn(HL3]� [8], indicate that in bothcases the nitrogen donors of the macrocycle determine a
square-planar arrangement around the metal centre. But
while the nickel complex can be described as a distorted
octahedron with the axial positions occupied by the
oxygen atoms of the phenolic pendant arms, in
[Zn(HL3]� the oxygen atom of the protonated pheno-
late is also directed towards to zinc but at a long
distance from the metal suggesting a square-pyramidalgeometry rather than an octahedral one. Furthermore,
DFT theoretical calculations suggested that the cop-
per(II) complex adopts a geometry which can be
considered as between distorted octahedral and
square-pyramidal [8]. Our EPR results indicate that
the monoprotonated copper(II) complex of L1 also
adopts a comparable distorted arrangement, although
the deprotonated complex should adopt an octahedralgeometry.
Our potentiometric measurements for the complexes
of all the metal ions studied and the EPR spectroscopy
of the copper(II) complexes clearly indicate the forma-
tion of dinuclear complexes. However, on the basis of
these results it is impossible to predict the co-ordination
sphere and the relative position of the second metal ion,
but they suggest that the quinoline arms are necessaryinvolved.
Acknowledgements
The authors acknowledge the financial support fromFundacao Para a Ciencia e Tecnologia (FCT) and
POCTI, with coparticipation of the European Commu-
nity fund FEDER (Project No. POCTI/1999/QUIM/
35396). X.C. also acknowledges FCT for the grant
(SFRH/BPD/1502/2000).
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