Spectroscopic, solvent influence and thermal studies of ternary copper(II) complexes of diester and dinitrogen base ligands

Article abstract of DOI:10.1016/j.saa.2010.06.012

New mixed-ligand copper(II) complexes containing the bidentate dinitrogen ligands [N,N,N',N'-tetramethylethylenediamine (tmen), 2,2'-bipyridine (bipy) and 1,10-phenanthroline (phen)] and the bidentate dioxygen ligands [diethylmalonate (DEM), ethylacetoacetate (EAA) and ethylbenzoylacetate (EBA)] were prepared. The complexes were characterized by elemental analysis, infrared, mass and ESR spectral data, magnetic and molar conductance measurements and thermal gravimetric analysis. From the investigation, the geometries of the complexes are square planar for perchlorate complexes and a square pyramid or octahedral for the nitrate complexes. Solvatochromic behavior of the Cu(II) complexes indicates strong solvatochromism of their solutions in polar and non-polar solvents. The observed solvatochromism is due to the solute-solvent interaction between the chelate cation and the solvent molecules.

Full text of DOI:10.1016/j.saa.2010.06.012

Spectrochimica Acta Part A 77 (2010) 594–604
Contents lists available at ScienceDirect
Spectrochimica Acta Part A: Molecular and
Biomolecular Spectroscopy
journal homepage: www.elsevier.com/locate/saa
Spectroscopic, solvent influence and thermal studies of ternary copper(II)
complexes of diester and dinitrogen base ligands
Adel A.A. Emaraa,∗, Azza A.A. Abu-Husseinb, Ahmed A. Tahab, Nelly H. Mahmoudb
a Department of Chemistry, Faculty of Education, Ain Shams University, Roxy, Cairo 11341, Egypt
b Department of Chemistry, Faculty of Girls, Ain Shams University, Heliopolice, Cairo 11221, Egypt
a r t i c l e i n f o
a b s t r a c t
Article history:
New mixed-ligand copper(II) complexes containing the bidentate dinitrogen ligands [N,N,N’,N’-
tetramethylethylenediamine (tmen), 2,2’-bipyridine (bipy) and 1,10-phenanthroline (phen)] and the
bidentate dioxygen ligands [diethylmalonate (DEM), ethylacetoacetate (EAA) and ethylbenzoylacetate
(EBA)] were prepared. The complexes were characterized by elemental analysis, infrared, mass and ESR
spectral data, magnetic and molar conductance measurements and thermal gravimetric analysis. From
the investigation, the geometries of the complexes are square planar for perchlorate complexes and a
square pyramid or octahedral for the nitrate complexes. Solvatochromic behavior of the Cu(II) com-
plexes indicates strong solvatochromism of their solutions in polar and non-polar solvents. The observed
solvatochromism is due to the solute–solvent interaction between the chelate cation and the solvent
molecules.
Received 2 April 2010
Received in revised form 22 May 2010
Accepted 4 June 2010
Keywords:
Solvatochromism
Mixed-ligand Cu(II) complexes
Diester ligand
Dinitrogen ligands
© 2010 Elsevier B.V. All rights reserved.
1. Introduction
Several works have been published concerning the solva-
tochromism by Fukuda, Sone and others where the dioxygen
Solvatochromism in the mixed-ligand metal complexes has
received great attention because they may be used as a Lewis
acid–base color indicator [1–6] and also utilizes to develop opti-
cal sensor materials to monitor pollutant levels in the environment
[7–11].
The effect of the solvent on the spectral properties of molecules,
generally referred as solvatochromism, has been investigated for
many years, generating a copious literature [12–16].
Solvatochromism of metal complexes can be divided into three
types [2,3]; the first one comprises the case where the color changes
are by the direct attachment of solvent molecules onto metal cen-
ter, and the solvatochromic bands could be of metal-to-ligand or
ligand-to metal character and the second type is due to the attach-
ment of solvent molecules onto ligands, while the third type of the
solvatochromic bands could be of ligand-to-ligand or intra-ligand
character. Among the former whose color changes are due to those
of d–d transitions, copper(II) complexes with a strong Jahn–Teller
effect can be anticipated to show simple and regular changes in
their electronic spectra according to the strength of interactions
with solvent molecules at the axial sites [7,8,18,19].
ligands were taken as ␤-diketones [8,12,17,18].
Recently, several studies were made on the solvatochromic
behavior of Cu(II) complexes using the dioxygen ligands as о-
hydroxy benzoyl derivatives with dinitrogen ligands [19]. Also, the
solvatochromic mixed-ligand copper(II) complexes with the gen-
eral formula [Cu(Rmal)(diam)]X where Rmal = malonate deriva-
tives and diam = diamine derivatives and X− = ClO4−or Cl− were
synthesized and characterized by spectroscopic, magnetic, molar
conductance and electrochemical measurements [20]. The mass
spectra along with the analytical data suggest a bridged binuclear
structure for the chloride complexes, while perchlorate complexes
showed either mononuclear or bridged binuclear structure.
The present work reports the syntheses, spectral, ther-
mal and solvatochromic behavior of some mixed-ligand cop-
per(II) complexes of the general formula, Cu(OO)(NN)X where
(OO) = diethylmalonate, ethylacetoacetate or ethylbenzoylacetate
and (NN) = N,N,N’,N’-tetramethylethylenediamine (tmen), 2,2’-
bipyridine (bipy) or 1,10-phenanthroline (phen) and X− = NO3
,
−
ClO4 or Br−.
−
The structure of the 18 new complexes is proved using elemen-
tal analysis, infrared and electronic, mass spectra, electron spin
resonance, thermal gravimetric analysis, magnetic and molar con-
ductivity measurements.
Although all perchlorate complexes exhibit square planar geom-
etry, the nitrate and bromide complexes exhibit either square
pyramid or octahedral geometry. Fig. 1 is a representative structure
∗
Corresponding author. Tel.: +20 224551530; fax: +20 222581243.
E-mail addresses: adelaaemara@yahoo.com, adelaemara@gmail.com
(A.A.A. Emara).
1386-1425/$ – see front matter © 2010 Elsevier B.V. All rights reserved.
doi:10.1016/j.saa.2010.06.012

A.A.A. Emara et al. / Spectrochimica Acta Part A 77 (2010) 594–604
595
In case of the remaining complexes, precipitates were formed
during the addition of the dinitrogen base. The precipitates were
filtered off and air-dried. All these complexes are insoluble in most
common solvents.
The bromide complex [Cu(OO)(tmen)Br] was prepared in solu-
tion by mixing t-Bu4NBr with [Cu(OO)(tmen)]ClO4 in CH2Cl2
solvent (molar ratio of 5:1).
[N.B. Caution: the melting point, mass spectra and thermal anal-
ysis were not possible to be used in the perchlorate complexes due
to its detonation effect.]
These complexes were not isolated in the solid state, and were
used to study the effect of the bromide anion on the complexes in
solution, in different solvents.
The following detailed preparations are given as examples and
the other complexes were obtained similarly.
2.2.1. Synthesis of complex (2)
Add a mixture of (0.801 g, 5.0 mmol) of diethylmalonate (DEM)
in 20 mL absolute ethanol and (1.060 g, 10.0 mmol) of anhy-
drous Na2CO3 to an ethanolic solution of (1.853 g, 5.0 mmol)
Cu(ClO4)2•6H2O. The mixture was stirred for about 30 min result-
ing in a green solution, which was then filtered. Then a solution
of (0.581 g, 5.0 mmol) of N,N,N’,N’-tetramethylethylenediamine
(tmen) in 10 mL absolute ethanol was added dropwise to the filtrate
with a continuous stirring for an additional half an hour. A solu-
tion was formed which was filtered off and left to stand overnight
and then was recrystallized from CH2Cl2 solvent. Dark blue crystals
were formed. The yield was 1.217 g (59.3%).
Fig. 1. Schematic representation of the Cu(II) complexes.
of the complexes. Of these, four Cu(II) complexes show solva-
tochromic behavior as will be shown.
2.2.2. Synthesis of complex (7)
2. 2-Experimental
Add a mixture of (0.651 g, 5.0 mmol) of ethylacetoacetate (EAA)
in 20 mL absolute ethanol and (1.060 g, 10.0 mmol) of anhy-
drous Na2CO3 to an ethanolic solution of (1.208 g, 5.0 mmol)
Cu(NO3)2•3H2O. The mixture was stirred for about 30 min, result-
ing in a green solution, which was then filtered. Then, a solution
of (0.581 g, 5.0 mmol) of N,N,N’,N’-tetramethylethylenediamine
(tmen) in 10 mL absolute ethanol was added dropwise to the filtrate
with a continuous stirring for an additional half an hour. A solution
was formed which was filtered off and left to stand overnight and
then was recrystallized from CH2Cl2 solvent. Violet crystals were
formed. The yield was 1.660 g (79.1%) and its melting point was
200 ◦C.
2.1. Materials
Copper nitrate tri hydrate Cu(NO3)2•3H2O, copper perchlo-
rate hexa hydrate Cu(ClO4)2•6H2O, diethylmalonate (DEM),
ethylacetoacetate (EAA), ethylbenzoylacetate (EBA), N,N,N’,N’-
tetramethylethylenediamine (tmen), 2,2’-bipyridine (bipy), 1,10-
phenanthroline (phen), anhydrous sodium carbonate (Na2CO3)
and tertiary tetrabutylammonium bromide t-Bu4NBr were ana-
lytical reagent grade and obtained from either Merck or Aldrich
and were used without further purification. Organic solvents
used for spectral studies were nitrobenzene (PhNO2), acetoni-
trile (MeCN), methanol (MeOH), ethanol (EtOH), acetone (Me2CO),
methylenechloride (CH2Cl2), N,N-dimethylformamide (DMF) and
dimethylsulfoxide (DMSO) were spectro-grade and used without
further purification.
2.2.3. Synthesis of complex (14)
Add a mixture of (0.864 g, 5.0 mmol) of ethylbenzoylacetate
(EBA) in 20 mL absolute ethanol and (1.060 g, 10.0 mmol) of anhy-
drous Na2CO3 to an ethanolic solution of (1.853 g, 5.0 mmol)
Cu(ClO4)2•6H2O. The mixture was stirred for about 30 min result-
ing in a green solution, which was then filtered. Then a solution
of (0.581 g, 5.0 mmol) of N,N,N’,N’-tetramethylethylenediamine
(tmen) in 10 mL absolute ethanol was added dropwise to the filtrate
with a continuous stirring for an additional half an hour. A solution
was formed which was filtered off and left to stand overnight and
then was recrystallized from CH2Cl2 solvent. Violet crystals were
formed. The yield was 2.332 g (77.3%).
2.2. Syntheses of Cu(OO)(NN)X complexes
These complexes were prepared by adding
a mixture of
5.0 mmol of dioxygen ligand in 20 mL absolute ethanol and
10 mmol of anhydrous Na2CO3 to an ethanolic solution of 5.0 mmol
of CuX2•nH2O (where X− = NO3−or ClO4−). The mixture was stirred
for about 30 min resulting in a green solution, which was then
filtered. Then a solution of 5.0 mmol diam (where diam = tmen,
bipy or phen) in 10 mL absolute ethanol was added dropwise to the
filtrate with a continuous stirring for an additional half an hour. The
nine complexes [Cu(DEM)(tmen)]NO3 and [Cu(DEM)(tmen)]ClO4,
[Cu(EAA)(tmen)(NO3)]•2H2O, [Cu(EAA)(tmen)]ClO4, [Cu(EAA)-
2.3. Physical measurements
The analysis of carbon, hydrogen and nitrogen were carried out
at the Microanalytical Center, Cairo University, Giza, Egypt. Cop-
per ions were determined by EDTA solution using muroxide, as an
indicator. FT-IR spectra (4000–400 cm−1) of the compounds were
recorded as KBr discs using FT-IR (shemedzo) spectrophotometer
model 4000. The UV–vis spectra of the compounds were obtained
on a JASCO model V-550 UV–vis spectrophotometer. Mass spectra
(bipy)(NO3)(H2O)]•H2O,
[Cu2(EAA)2(phen)2(NO3)2(H2O)]•H2O,
[Cu(EBA)(tmen)(NO3)(H2O)], [Cu(EBA)(tmen)]ClO4 and [Cu(EBA)
(bipy)(NO3)(H2O)]•H2O were freely soluble in ethanol. These
complexes were filtered off and left to stand overnight and then
were recrystallized from CH2Cl2 solvent, where dark blue crystals
were obtained.

596
A.A.A. Emara et al. / Spectrochimica Acta Part A 77 (2010) 594–604
were carried out by a shimadzu-GC–MS–QL mass spectrome-
ter model 1000EX using a direct inlet system. TGA curves were
obtained using NETZSCH-gerateban Bestell-Nr 348472c. Thermal
analyzer equipped with a thermobalance. Samples (∼50 mg) were
heated at a programmed rate of 5 ◦C min−1 in a dynamic N2 atmo-
sphere. The sample was contained in a boat-shaped platinum pan
suspended in the center of a furnace. Meting points reported in
this work were not corrected. Magnetic susceptibilities were mea-
sured at room temperature using a magnetic susceptibility balance
(Johnson Matthey Alfa product, Model No. MKI). ESR spectra were
recorded on a Bruker Model EMX, X-band spectrometer. Diamag-
netic corrections calculated from Pascal’s constants [21]. Molar
conductance of 10−3 M solutions of the complexes in CH2Cl2 and
DMF solutions were measured on a Corning conductivity meter NY
14831 model 441 (USA)
There are five conceptual features in the infrared spectra of the
complexes. In case of the free perchlorate, ClO4−, there are four
vibrational bands observed at 935, 460, 1050–1170 and 630 cm−1
.
The third and fourth bands appear strongly in the infrared spec-
tra. The strong bands observed in the perchlorate Cu(II) complex−e1s
2, 4, 6, 8, 10, 12, 14, 16 and 18 at the range 1084–1117 cm
(antisymmetric stretch) and the sharp band at 621–625 cm−1
(antisymmetric bend) suggest uncoordinated perchlorate anions
[23]. The free nitrate, NO3−, ions exhibited two bands which are
observed at 1384 and 856 cm−1. Three strong vibrational bands
for the nitrate groups were observed at 1357–1383, 1272–1277
and 952–972 cm−1 which indicate that it is coordinated in com-
plexes 1, 3, 5, 7, 9, 11, 13, 15 and 17. Unidentate nitrate groups
show C2V symmetry, with three non-degenerate modes of vibration
(ꢀs, ꢀs and ꢀas), where ꢀs(NO3−) is in the range 1357–1383 cm−1
,
ꢀ
ꢀ
ꢀs (NO3−) 1272–1296 cm−1 and ꢀas−(NO3−) 905–1026 cm−1 [24].
The ꢀs(NO3−) of the unidentate NO3 is markedly shifted to lower
frequency compared to that of the free nitrate (1384 cm−1) [24]
providing a measure of the covalent bond strength due to transfer
3. Results and discussion
The ternary metal complexes were prepared in the molar
ratio 1:1:1 from diethylmalonate, dinitrogen bases (tmen, bipy
−
of electron density from NO3 to the metal ion.
The second feature is related to the ethyl groups which appeared
at the range 2801–3118 cm−1. This indicates that it still persist in
the complexes compared with their precursors [25].
−
−
and phen) and copper ions associated with NO3 or ClO4 as
counter-anions in absolute ethanol. The diester possesses monoba-
sic bidentate dioxygen ligand, while all dinitrogen bases are neutral
bidentate ligands. All Cu(II) complexes carried one positive charge
due to the presence of ˛-hydrogen in the diester where the struc-
ture of the diester exists in several tautomeric forms due to the
presence of ˛-hydrogen of the methylene group, where the depro-
tonated (anionic form IV) is the active form, which binds to the
Cu(II) center compensating one positive charge of Cu(II). Scheme 1
represents the tautomeric forms of the diethylmalonate.
The third feature is the dioxygen (OO) ligand in the ternary com-
plexes exhibited two partially resolved bands at 1602–1665 and
1514–1586 cm−1, which assigned to ꢀas(COO) and a single band
at 1457–1484 cm−1 for ꢀs(COO). Generally, the free dioxygen lig-
and showed a strong broad band at 1739 cm−1 [26]. This band
was shifted to lower frequencies which appeared at 1602–1628
and 1565–1586 cm−1. This lowering shift may be due to the chela-
tion of the oxygen atoms of the carbonyl group in the dioxygen
ligand.
The fourth feature is about the characteristic bands of the dini-
trogen base ligands. Changes are observed in the infrared spectrum
of bipy, when it is coordinated with the metal centers. The char-
acteristic band of the C N group in complexes 3, 4, 9, 10, 15
and 16 are shifted to lower frequencies compared with the free
ligand at 1570 cm−1 [27]. This suggests that both nitrogens of 2,2ꢀ-
bipy are coordinated to the metal ion [28]. Similar behavior has
been observed for the C N vibration of phen, which has a band at
1560 cm−1. This band shifts on coordination to lower frequencies in
complexes 5, 6, 11, 12, 17 and 18; respectively, presumably because
of coordination to the metal [29].
Eighteen new Cu(II) complexes were identified by infrared,
electronic, mass, electron spin resonance spectra, and magnetic
and molar conductivity measurements. Solvatochromic properties
were studied to investigate the effect of ligands and anions on the
structure of the complex and its chromotropic behavior. Based on
the effect of the anion and solvents, the Cu(II) complexes exhibit
different geometrical arrangements such as a square planar, square
pyramidal and octahedral arrangements. Table 1 lists the physical
and analytical data of the complexes.
3.1. Infrared spectra
The last feature is the weak to medium bands in the two
ranges 411–518 and 512–593 cm−1, which could be assigned to
the stretching frequencies of the ꢀ(Cu–O) and ꢀ(Cu–N) bands [30],
respectively, supporting that the bonding of the ligands to the metal
ions is achieved by the oxygen of the carbonyl group of the dioxygen
ligands and nitrogen atoms of the base ligands.
The vibrational assignments of the Cu(II) complexes were aided
by comparison with the vibrational frequencies of the starting
materials [22–27]. Table 2 lists the characteristic infrared bands for
1–18 complexes. A comprehensive vibrational survey of the deriva-
tive ligands are published by other authors and the references are
cited in the text.
3.2. Electronic spectra, magnetic and molar conductivity
measurements
The electronic spectra for the complexes [Cu(DEM)
(tmen)(NO3)]•2H2O (1), [Cu(DEM)(tmen)]ClO4 (2), [Cu(EAA)
(tmen)NO3]•2H2O (7), [Cu(EAA)(tmen)]ClO4 (8), [Cu(EAA)(bipy)
(NO3)(H2O)]•H2O (9), [Cu2(EAA)2(phen)2(NO3)2(H2O)]•H2O (11),
[Cu(EBA)(tmen)(NO3)(H2O)] (13), [Cu(EBA)(tmen)]ClO4 (14) and
[Cu(EBA)(bipy)(NO3)(H2O)]•H2O (15) were obtained in CH2Cl2
solutions, while the other complexes were insoluble in most
common organic solvents and were not possible to get electronic
spectra in solution. The electronic spectra were obtained by
reflectance within the range 450–800 nm. It is possible to draw up
the electronic transitions and predict the geometry with the aid
of the magnetic moment and molar conductivity measurements.
Table 3 lists the characteristic electronic spectral bands and the
Scheme 1. Schematic representation of the tautomeric forms (I–III) of the diester
ligands (R = OC2H5 for DEM, R = CH3 for EAA and R = C6H5 for EBA). Form IV is a
general form.

Table 1
The physical and analytical data of the Cu(DEM)(diam)X complexes.
Complex
Weight; g (mmol)
Empirical formula M.wt. Wt. (g)
(yield %)
Color
m.p.
( C)
Elemental analysis Found (Calc.) %
◦
OO ligand
0.801 (5.0)
NN ligand
Metal Salt
1.208 (5.0)
C
H
N
Cl
–
Cu
•
1
DEM + tmen + Cu(NO
)
3H
O
0.755 (5.0)
0.755 (5.0)
0.781 (5.0)
0.781 (5.0)
0.991 (5.0)
0.991 (5.0)
0.755 (5.0)
0.755 (5.0)
0.781 (5.0)
0.781 (5.0)
0.991 (5.0)
0.991 (5.0)
0.755 (5.0)
0.755 (5.0)
0.781 (5.0)
0.781 (5.0)
0.991 (5.0)
0.991 (5.0)
CuC
CuC
CuC
CuC
H
N
N
N
N
O
9
436.5 1.624 (39.6)
410.0 1.217 (59.3)
476.5 1.065 (37.4)
478.0 1.752 (73.2)
947.0 3.181 (67.1)
520.0 1.172 (45.0)
406.5 1.660 (79.1)
408.0 1.790 (87.0)
446.5 1.412 (63.2)
448.0 1.270 (56.5)
905.0 3.221 (71.0)
472.0 1.940 (81.9)
450.5 2.682 (82.4)
486.0 2.332 (77.3)
508.5 2.313 (67.9)
446.0 1.703 (57.0)
514.5 2.230 (59.4)
534.0 2.829 (79.1)
Dark blue
Dark blue
245
–
35.83 (35.74)
6.60 (6.84)
9.63 (9.62)
14.65 (14.54)
3
2
2
13 35
3
2
3
2
•
→ [Cu(DEM)(tmen)(NO )] 2H
O
3
2
•
2
DEM + tmen + Cu(ClO
)
6H
O
0.801 (5.0)
0.801 (5.0)
0.801 (5.0)
0.801 (5.0)
0.801 (5.0)
0.651 (5.0)
0.651 (5.0)
0.651 (5.0)
0.651 (5.0)
0.651 (5.0)
0.651 (5.0)
0.864 (5.0)
0.864 (5.0)
0.864 (5.0)
0.864 (5.0)
0.864 (5.0)
0.864 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
1.208 (5.0)
1.853 (5.0)
H
ClO
32.10 (32.20)
37.86 (37.78)
42.48 (42.67)
47.86 (48.15)
45.98 (45.40)
39.57 (41.33)
35.10 (35.29)
43.41 (43.00)
42.48 (42.86)
46.68 (46.80)
45.98 (45.76)
45.63 (45.28)
41.86 (41.97)
49.64 (49.56)
56.37 (56.50)
54.26 (53.65)
51.69 (51.08)
5.66 (5.61)
8.76 (8.81)
4.07 (3.97)
4.55 (4.22)
3.93 (3.80)
7.18 (7.63)
6.66 (6.13
4.26 (4.70)
4.07 (3.79)
4.01 (4.33)
3.93 (3.60)
6.25 (6.45)
5.61 (5.56)
4.47 (4.52)
4.33 (4.26)
4.15 (4.08)
3.41 (3.56)
6.63 (6.83)
7.65 (7.39)
5.31 (5.90)
8.79 (8.87)
5.52 (5.58)
10.41 (10.33)
6.63 (6.86)
9.83 (9.41)
6.31 (6.25)
9.35 (9.10)
5.52 (5.93)
9.64 (9.32)
5.56 (5.36)
8.26 (8.17)
6.41 (6.28)
10.01 (10.11)
5.12 (5.24)
8.90 (8.70)
15.41 (15.48)
13.46 (13.33)
13.12 (13.30)
13.55 (13.41)
12.00 (12.20)
16.03 (15.62)
15.04 (15.56)
13.97 (14.22)
14.00 (14.18)
13.97 (13.76)
13.22 (13.45)
14.01 (14.09)
12.99 (13.07)
12.49 (12.36)
14.24 (14.43)
12.47 (12.34)
11.95 (11.89)
4
2
2
11 23
8
→ [Cu(DEM)(tmen)]ClO
4
3H
•
3
DEM + bipy + Cu(NO
)
O
H
O
8
Bluish-violet 330
Bluish-violet
Bluish-violet 320
–
3
2
2
15 21
•
→ [Cu(DEM)(bipy)(NO )(H O)] H
O
3
2
2
•
4
DEM + bipy + Cu(ClO
)
6H
O
H
ClO
–
7.61 (7.43)
4
2
2
17 19
8
→ [Cu(DEM)(bipy)]ClO
4
•
5
DEM + phen + Cu(NO
)
3H
O
Cu
C
H
N
O
–
3
2
2
2
38 40
6
15
→ [Cu (DEM) (phen) (NO
)
(H O)]
2
2
2
)
3
6H
2 2
•
6
DEM + phen + Cu(ClO
O
CuC
CuC
CuC
CuC
CuC
H
N
ClO
Bluish-violet
Violet
–
200
–
7.10 (7.07)
4
2
2
19 22
2
3
2
3
2
9
→ [Cu(DEM)(phen)]ClO
4
•
7
EAA + tmen + Cu(NO
)
3H
O
H
N
N
N
N
O
8
–
3
2
2
12 29
•
→ [Cu(EAA)(tmen)(NO )] 2H
O
3
2
O
•
8
EAA + tmen + Cu(ClO
)
6H
H
12 25
ClO
Violet
8.90 (8.70)
4
2
2
→ [Cu(EAA)(tmen)]ClO
4
3H
•
9
EAA + bipy + Cu(NO
)
O
H
O
8
Green
250
–
–
3
2
2
16 21
•
→ [Cu(EAA)(bipy)(NO ) (H O)] H
O
3
2
2
•
10 EAA + bipy + Cu(ClO
)
6H
O
H
ClO
Bluish-violet
Green
7.61 (7.93)
4
2
2
16 17
7
→ [Cu(EAA)(bipy)]ClO
4
•
11 EAA + phen + Cu(NO
)
3H
O
Cu
C
H
N
O
340
–
–
3
2
2
2
36 38
6
14
•
(H O)] H O
2 2 2
→ [Cu (EAA) (phen) (NO
)
2
2
2
3
•
12 EAA + phen + Cu(ClO
)
6H
O
CuC
CuC
CuC
CuC
CuC
CuC
CuC
H
N
ClO
Bluish-violet
Violet
7.10 (7.52)
4
2
2
18 17
2
3
2
3
2
3
2
→ [Cu(EAA)(phen)]ClO
4
•
13 EBA + tmen + Cu(NO
)
3H
O
H
N
N
N
N
N
N
O
7
155
–
–
3
2
2
17 29
→ [Cu(EBA)(tmen)(NO )(H O)]
3
2
•
14 EBA + tmen + Cu(ClO
)
6H
O
H
ClO
Violet
7.51 (7.31)
4
2
2
17 27
8
8
9
→ [Cu(EBA)(tmen)]ClO
4
3H
•
15 EBA + bipy + Cu(NO
)
O
H
O
8
Green
200
–
–
3
2
2
21 23
•
→ [Cu(EBA)(bipy)(NO )(H O)] H O
3
2
2
•
16 EBA + bipy + Cu(ClO
)
6H
O
H
17 19
ClO
Blue
7.73 (7.96)
–
4
2
2
→ [Cu(EBA)(bipy)]ClO
4
•
17 EBA + phen + Cu(NO
)
3H
O
O
O
H
O
8
Bluish-green 180
Blue
3
2
2
23 21
•
→ [Cu(EBA)(phen)(NO )] H
3
2
•
18 EBA + phen + Cu(ClO
)
6H
H
32 19
ClO
–
6.95 (6.65)
4
2
2
→ [Cu(EBA)(phen)]ClO
4

598
A.A.A. Emara et al. / Spectrochimica Acta Part A 77 (2010) 594–604
Table 2
Infrared frequencies of the characteristic bands of Cu(II) complexes (cm−1).
Complexes
␯(CH3CH2), ␯(CH3)
and ␯(CH)aromatic
␯(CO−
)
␯(NO3)
␯(ClO4)
␯(Cu-N)
␯(Cu-O)
1
2
3
4
5
6
7
8
[Cu(DEM)(tmen)(NO3)]•2H2O
[Cu(DEM)(tmen)]ClO4
2929–3016
2929–2989
2921–3109
2924–3118
2927–3113
3067
2806–3008
2807–2993
2933–3109
2934–3116
2937–3055
2914–3086
2801–2984
2850–2990
2983–3113
2934–3015
2900–3046
3086–3118
1612 1567 1465
1626 1586 1469
1602 1565 1475
1603 1571 1473
1628 1586 1484
1625 1584 1457
1610 1564 1465
1613 1517 1472
1607 1521 1474
1608 1567 1473
1607 1514 1477
1606 1566 1475
1613 1576 1479
1609 1574 1470
1665 1577 1446
1608 1566 1497
1648 1586 1475
1610 1568 1475
1357 1272 952
–
1380 1296 1026
–
1383 1277 972
–
1383 1273 989
–
1381 1286 987
–
1383 1282 972
–
1357 1272 952
–
1364 1293 1023
–
–
593
512
523
515
588
544
516
513
585
486
488
494
554
556
513
518
597
544
518
445
444
411
430
424
477
454
419
420
429
433
484
486
478
482
480
477
1201 1090 1021 624
–
1160 1084 1011 621
–
1146 1089 1020 622
–
1179 1089 1016 623
–
1174 1095 1061 622
–
1150 1117 1064 622
–
1200 1092 1021 625
–
1161 1093 1011 621
–
1146 1089 1020 622
[Cu(DEM)(bipy)(NO3)(H2O)]•H2O
[Cu(DEM)(bipy)]ClO4
[Cu2(DEM)2(phen)2(NO3)2(H2O)]
[Cu(DEM)(phen)]ClO4
[Cu(EAA)(tmen)(NO3)]•2H2O
[Cu(EAA)(tmen)]ClO4
9
[Cu(EAA)(bipy)(NO3)(H2O)]•H2O
[Cu(EAA)(bipy)]ClO4
10
11
12
13
14
15
16
17
18
[Cu2(EAA)2(phen)2(NO3)2(H2O)]•H2O
[Cu(EAA)(phen)]ClO4
[Cu(EBA)(tmen)(NO3)(H2O)]
[Cu(EBA)(tmen)]ClO4
[Cu(EBA)(bipy)(NO3)(H2O)]•H2O
[[Cu(EBA)(bipy)]ClO4
[Cu(EBA)(phen)(NO3)]•H2O
[Cu(EBA)(phen)]ClO4
1363 1277 1014
–
magnetic moments of the Cu(II) ions in their complexes in the
solid state, and the molar conductance measurements.
ethylbenzoylacetate. The last two bands correspond to the n → ꢁ*
transitions of the oxygen and nitrogen atoms which are overlapped
with the intermolecular charge transfer from the aromatic ring.
The one unsymmetrical band for 2, 4, 6, 8, 10, 12, 14, 16, and 18
complexes at 564, 584, 521, 609, 597, 587, 567, 578 and 589 nm due
The electronic transitions due to the organic ligands are not
mentioned in the tables, where in all electronic transitions of the
Cu(II) complexes, the absorption bands due to the organic ligands
were appeared and compared with the electronic transition of the
free ligands. Generally, the intensities of the electronic transition
bands due to the d–d transitions in the metal complexes are lower
than the electronic transitions due to the organic ligands.
2
2
2
to 2A1g ← B1g
,
2B2g ← B1g
,
2Eg ← B1g transitions in a square pla-
nar geometry [32]. The one band for 1 and 7 complexes at 656 and
2
2
573 nm due to T2g ← Eg transition in square pyramid geometry.
The observed one band for 3, 5, 9, 11, 13, 15 and 17 complexes at
708, 693, 678, 702, 645, 678 and 614 nm due to 2T2g ← Eg transition
2
Electronic spectral data of the ligands in the Cu(II) complexes
were predicted in CH2Cl2 solution. In the electronic spectrum of
the ligands, five absorption bands at 272–290, 307–319, 363–373,
390–410 and 428–430 nm were characterized. The former two
in octahedral geometry.
The obtained complexes are violet, dark blue, bluish-violet
and green crystals. The analytical and molar conductance data
suggest the mononuclear structure of the complexes except
for [Cu2(DEM)2(phen)2(NO3)2(H2O)] (5) and [Cu2(EAA)2(phen)2
(NO3)2(H2O)]•H2O (11) complexes. The molar conductance values
obtained for solutions of these complexes in CH2Cl2 were in the
range 10, 83, 43, 85, 15, 19, 11, 83 and 14 Ohm−1 cm2 mol−1 for
complexes 1, 2, 7, 8, 9, 11, 13, 14 and 15; respectively, while in
1
1
bands correspond to 1La → A1 and 1Lb → A1 transitions of the
aromatic ring of the 1,10-phenanthroline, 2,2ꢀ-bipyridine and
ethylbenzoylacetate ligands [31]. The third band corresponds to
the ꢁ → ꢁ* transition of the carbonyl group of the dioxygen ligands
(diethylmalonate, ethylacetoacetate and ethylbenzoylacetate) and
also, the phenyl rings of 1, 10-phenanthroline, 2,2ꢀ-bipyridine and
Table 3
Electronic absorption bands (nm) in CH2Cl2 solutions, molecular ion peak, magnetic moment and molar conductance of the copper(II) complexes and their assignments.
Complex
ꢂmax (nm)
Molecular ion peakb
Magnetic moment, ꢃ eff (B.M.)
Molar conductance (ꢄ)
m/z (intensity%)
CH2Cl2
DMF
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
[Cu(DEM)(tmen) (NO3)]•2H2O
[Cu(DEM)(tmen)]ClO4
656 (102)
564 (71)
708a
438 (0.5)
–
477 (4.0)
–
480 (7.0)
–
411 (14.7)
–
447 (11.9)
–
471 (1.9)
–
453 (15)
–
509 (18.9)
–
1.98
2.17
1.92
2.09
2.43c
2.14
1.99
2.10
1.91
2.04
3.07c
2.11
1.99
2.20
1.94
2.11
1.93
2.13
10
83
–
–
–
97
102
91
105
140
95
98
110
91
106
135
96
97
103
91
[Cu(DEM)(bipy)(NO3)(H2O)]•H2O
[Cu(DEM)(bipy)]ClO4
584a
[Cu2(DEM)2(phen)2(NO3)2(H2O)]
[Cu(DEM)(phen)]ClO4
693a
521a
–
[Cu(EAA)(tmen)(NO3)]•2H2O
[Cu(EAA)(tmen)]ClO4
573 (84)
609 (74)
678 (91)
597a
43
85
15
–
19
–
11
83
14
–
[Cu(EAA)(bipy)(NO3)(H2O)]•H2O
[Cu(EAA)(bipy)]ClO4
[Cu2(EAA)2(phen)2(NO3)2(H2O)]•H2O
[Cu(EAA)(phen)]ClO4
702 (88)
587a
[Cu(EBA)(tmen)(NO3)(H2O)]
[Cu(EBA)(tmen)]ClO4
645 (69)
567 (42)
678 (77)
578a
[Cu(EBA)(bipy)(NO3)(H2O)]•H2O
[[Cu(EBA)(bipy)]ClO4
104
96
95
[Cu(EBA)(phen)(NO3)]•H2O
[Cu(EBA)(phen)]ClO4
614a
563 (7.2)
–
–
–
589a
Electronic spectra in the visible region were recorded in CH2Cl2 solvent and the values (εmax) are in parentheses (mol−1 cm−1). The electronic transitions for the ligands were
observed in all copper(II) complexes within the range 200–450 nm, and is not cited in this table.
Molar conductivities (Ohm−1 cm2 mol−1) were measured in 10−3 mol−1 cm−1 in CH2Cl2 and DMF solvents.
aValues were measured in Nujol mull in the solid state.
bMass spectra for the perchlorate complexes were not performed due to its detonation effect.
cꢃeff is the magnetic moment of one Cu(II) ion in the complex.

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599
DMF the values were in the range 91–130 Ohm−1 cm2 mol−1 for
the same complexes; respectively.
of magnetically dilute complexes are in the range 1.9–2.2 B.M.,
with compounds whose geometries approach octahedral having
moments at the lower end, and those with geometries approaching
square planar or tetrahedral having moments at the higher end [34].
This indicates that all the perchlorate complexes are square planar
while the nitrate complexes have octahedral geometry except 1, 7
and 17 complexes which are square pyramid.
In CH2Cl2 solvent, the values of the perchlorate complexes are
in the range of 1:1 electrolytes, while the values of the nitrate com-
plexes are in the non-electrolytic range. In DMF solution, the molar
conductance values are in the range 1:1 or 1:2 electrolytes of both
the perchlorate and nitrate counter-anions; respectively. This may
be due to the fact that the DMF solvent replaced the NO3− anion in
the complexes, which results in the 1:1 or 1:2 electrolytes due to
the uncoordinated nitrate ions [33].
3.3. Mass and ESR spectra
The molar conductance for 3, 4, 5, 6, 10, 12, 16, 17 and 18 com-
plexes were measured only in DMF and gave values in the range
91–140 Ohm−1 cm2 mol−1; respectively. In these complexes DMF
replaced the NO3− anion and all the complexes are in the range 1:1
electrolyte except complexes 5, 11 which are 1:2 electrolyte.
The magnetic moments data of the investigated Cu(II) com-
plexes are found in the range 1.91–2.20 B.M. In general, moments
The mass spectra of the Cu(II) complexes of the nitrate
were performed and the molecular ion peak were recorded in
Table 2. The molecular ion peaks are coincident with the for-
mula weight of the complexes (See Table 1). This supports the
identity of the structure. Schemes 2 and 3 illustrate the frag-
mentation patterns of the complexes, where complexes 5 and 9
showed the fragmentation of the mononuclear pattern only. The
Scheme 2. Fragmentation patterns of the [Cu2(DEM)2(phen)2(NO3)2(H2O)] complex (5).

600
A.A.A. Emara et al. / Spectrochimica Acta Part A 77 (2010) 594–604
Scheme 3. Fragmentation patterns of the [Cu(EAA)(bipy)(NO3)(H2O)]•H2O (9).
mass spectra [Cu(DEM)(bipy)(NO3)(H2O)]•H2O complex (3) and
[Cu(EAA)(tmen)(NO3)]•2H2O complex (7) were depicted in Fig. 2.
The mass spectra of the perchlorates were not performed due to
the detonation of the perchlorate compounds.
of the later complex (5) exhibited two bands, one of them is sharp
with g = 2.1948 and the other is very weak at g = 2.0514. The shape
of the spectrum is consistent with the octahedral geometry around
each Cu(II) center in the complex.
X-band ESR spectrum of [Cu(DEM)(bipy)]ClO4 complex (4) and
[Cu2(DEM)2(phen)2(NO3)2(H2O)] complex (5) were recorded in the
solid state at 25 ◦C. The spectrum of the later complex (4) exhibited
one band with g value 2.0811, which is in consistent with the square
planar geometry around the Cu(II) ion in the complex. The spectrum
3.4. Thermal gravimetric analysis
Thermal gravimetric analysis (TGA) was used as a probe to
proof the associated water or solvent molecules to be in the

A.A.A. Emara et al. / Spectrochimica Acta Part A 77 (2010) 594–604
601
Fig. 2. Mass spectra of (A) [Cu(DEM)(bipy)(NO3)(H2O)]•H2O complex (3) and (B) [Cu(EAA)(tmen)(NO3)]•2H2O complex (7).
Fig. 3. TGA-DrTGA of (A) [Cu(DEM)(tmen)(NO3)]•2H2O (1) and (B) [Cu(EAA)(bipy)(NO3)(H2O)]•H2O (9).
Table 4
Thermal behavior of the mixed-ligand Cu(II) nitrate complexes.
Compound
DrTGA peak (◦C)
Temperature range (◦C)
Decomposition product lost (formula Wt.)
Wt. loss% Found (Calcd.)
1
3
5
[Cu(DEM)(tmen)(NO3)]•2H2O
70.9
184.3
214.1
80.0
50–89
2H2O (36)
DEM (159)
tmen + NO2 (162)
H2O (18)
tmen + H2O + NO2 (223)
bipy (156)
bipy (156)
H2O (18)
2NO2 + 4C2H4 (204)
2CO2 (82)
EAA + 2H2O (165)
tmen (116)
7.874 (8.248)
118–205
210–520
32–142
37.110 (36.426)
35.964 (37.113)
3.500 (3.778)
47.110 (46.799)
32.423 (32.739)
32.423 (32.739)
2.359 (1.901)
20.992 (21.540)
8.281 (9.290)
42.922 (40.291)
28.135 (28.536)
11.233 (11.316)
5.225 (4.032)
[Cu(DEM)(bipy)(NO3)(H2O)]•H2O
[Cu2(DEM)2(phen)2(NO3)2(H2O)]
242.4
405.4
146.3
268.2
292.7
461.5
193.0
334.1
442.2
75.4
148–249
255–413
140–165
214–277
280–314
328–490
145–267
314–410
437–524
41–106
7
9
[Cu(EAA)(tmen)(NO3)]•2H2O
NO2 (46)
H2O (18)
[Cu(EAA)(bipy)(NO3)(H2O)]•H2O
132.2
212.1
258.3
474.2
51.6
117–176
179–221
226–323
339–534
33–114
H2O + C2H6 (46)
bipy (156)
NO2 (46)
CH3COCHCOOH (101)
H2O (18)
9.904 (10.302)
35.221 (34.938)
11.092 (10.301)
21.277 (22.620)
2.525 (1.989)
11
[Cu2(EAA)2(phen)2(NO3)2(H2O)]•H2O
137.5
216.9
308.4
372.6
211.6
442.5
75.6
121–157
168–234
246–313
321–426
175–347
425–504
65–106
C2H4 (28)
3.255 (3.094)
2NO2 + C2H6 (122)
2CO2 + C2H4 + H2O (134)
2phen + 2CH2CHO (446)
EBA + H2O + tmen (325)
NO2 (46)
12.963 (13.481)
14.703 (14.807)
48.317 (49.282)
71.588 (72.142)
8.882 (10.211)
3.509 (3.355)
13
15
[Cu(EBA)(tmen)(NO3)(H2O)]
[Cu(EBA)(bipy)(NO3)(H2O)]•H2O
H2O (18)
231.1
417.1
87.8
235.2
468.9
145–302
349–484
106–179
209–319
395–563
EBA + H2O + NO2 (255)
bipy (156)
H2O (18)
phen (180)
NO2 + EBA (237)
50.551 (50.148)
31.558 (30.678)
3.470 (3.498)
31.994 (34.985)
50.00 (46.069)
17
[Cu(EBA)(phen)(NO3)]•H2O

602
A.A.A. Emara et al. / Spectrochimica Acta Part A 77 (2010) 594–604
Fig. 4. Electronic absorption spectra of 4.0 × 10−3 mol L−1 of (A) [Cu(DEM)(tmen)]ClO4 complex (2) and (B) [Cu(EAA)(tmen)(NO3)]•2H2O complex (7) solutions at 25 ◦C.
Fig. 5. Electronic absorption spectra of 4.0 × 10−3 mol L−1 of (A) [Cu(EAA)(tmen)]ClO4 complex (8) and (B) [Cu(EBA)(tmen)]ClO4 complex (14) solutions at 25 ◦C.
coordination sphere or in the crystalline form [35]. Thermo-
gravimetric (TGA) was carried out for the nitrate complexes
only (complexes 1 and 9 chosen as examples), since the per-
chlorate ones expected to be explosive if dried. TGA–DrTGA
curves of the [Cu(DEM)(tmen)(NO3)]•2H2O (1) and [Cu(EAA)-
(bipy)(NO3)(H2O)]•H2O (9) were depicted in Fig. 3.
steps (Fig. 3(B)). Decomposition of the complex started at 41 ◦C
and finished at 524 ◦C with five stages. The first stage of decom-
position involves the loss of uncoordinated H2O molecule in the
41–106 ◦C temperature range, and is accompanied by a weight
loss of 5.225%. The second stage of decomposition occurs in the
The stages of decomposition, temperature ranges, decomposi-
tion product loss as well as the found and calculated weight loss
percentages of the complexes are given in Table 4.
The thermal decomposition process of the [Cu(DEM)(tmen)
(NO3)]•2H2O complex
1 involves three decomposition steps
(Fig. 3(A)). Decomposition of the complex started at 50 ◦C and fin-
ished at 520 ◦C with three stages. The first stage of decomposition
involves the removal of two H2O molecules in the 50–89 ◦C tem-
perature range, and is accompanied by a weight loss of 7.874%. The
second stage of decomposition occurs in the 118–205 ◦C temper-
ature range, corresponding to the loss of the DEM ligand, and is
accompanied by a weight loss of 37.110%. While the third stage
involves the removal of the tmen and the nitrate group a form of
NO2 in the 210–520 ◦C temperature range, and is accompanied by
a weight loss of 35.964%.
The thermal decomposition process of [Cu(EAA)(bipy)
(NO3)(H2O)]•H2O complex (9) involves five decomposition
Fig. 6. Solvation free energy relationship (DN vs ␯max/1000).

A.A.A. Emara et al. / Spectrochimica Acta Part A 77 (2010) 594–604
603
Fig. 7. Electronic absorption spectra of 4.0 × 10−3 mol L−1 of (A) [Cu2(DEM)(tmen)Br] complex (19), (B) [Cu(EAA)(tmen)Br] complex (20) and (C) [Cu(EBA)(tmen)Br] complex
(21) solutions at 25 ◦C.
117–221 ◦C temperature range, corresponding to the loss of C2H4
and coordinated H2O molecule, and is accompanied by a weight
loss of 9.904%. While the third stage involves the removal of bipy
ligand in the 179–221 ◦C temperature range, and is accompanied by
a weight loss of 35.221%. The fourth stage occurs in the 226–323 ◦C
temperature range, corresponding to the loss of the nitrate group
in the form of NO2. The last stage of decomposition occurs in the
339–534 ◦C temperature range, corresponding to the removal of
CH3COCHCOOH molecule, and is accompanied by a weight loss of
21.277%.
ꢀmax/103 = ꢀ◦ + a (DN). Where ꢀmax is the measured d–d absorption
frequency, ꢀ◦ is the extrapolated frequency and (a) is the slope,
represents the sensitivity of the complex toward solvent. Linear-
ity of the ꢀmax vs DN, confirms the solvatochromic behavior in
perchlorate complexes. Solvation free energy relationship (DN vs
ꢀmax/1000) was depicted in Fig. 6.
The slope values (a) are in the order ([Cu(EAA)(tmen)]ClO4 <
[Cu(EBA)–(tmen)]ClO4 < [Cu(EAA)(tmen)(NO3)]•2H2O < [Cu(DEM)
(tmen)(NO3)]•H2O The effect of the anion was studied using
t-Bu4NBr. The bromide complexes were prepared in solution in the
molar ratio 5:1 (t-Bu4NBr: perchlorate complexes) by replacing
the perchlorate anion by the bromide anion.
3.5. Electronic spectra of the complexes in various solvents
The different anions (NO3−, ClO4 or Br−) used in the present
−
study chosen in such away, since the large size of the perchlorate
anion makes it so difficult to coordinate to the Cu(II) ion, which
lead to the formation of mainly square planar complexes, while
the bromide anion with the small size and less electronegativity
than the perchlorate anion make it so easy for the bromide to
coordinate to the Cu(II) ion changing the geometry form square
planar to mainly octahedral Fig. 7. Electronic absorption spectra
of [Cu2(DEM)(tmen)Br] complex (19), [Cu(EAA)(tmen)Br] complex
(20) and [Cu(EBA)(tmen)Br] complex (21) solutions at 25 ◦C were
depicted in Fig. 7. The nitrate anion may be inside or outside the
coordination sphere leading to the formation of square planar,
square pyramid or octahedral geometry.
The absorption spectra of complexes 1, 2, 7, 8, 9, 11, 13, 14
and 15 were measured in different organic solvents with differ-
ent donor numbers (DN), while the complexes 3, 4, 5, 6, 10, 12
and 16 are sparingly soluble in most organic solvents thus the d–d
bands of these complexes were recorded in Nujol mulls. These elec-
tronic spectra show only one broad band in the visible region due
to the promotions of the electron in the lower energy orbital to
the hole in dx2–y2 orbital of the copper(II) ion (d9). The position of
this band is shifted to longer wavelength (red shift) as the donor
number of solvent increases. Assuming that the approach of the
solvent occurs along with z-axis, the solvent molecules are repelled
by the two electrons in d2z orbital of copper(II) and only the more
species can force their way to form strong bonds. Upon coordina-
tion, they exert a z-component field proportional to their position
in the spectrochemical series.
4. Conclusions
The donor number proposed by Gutmann [26–38] is a measure
of the donor ability or Lewis basicity of a solvent molecule mea-
sured on the basis of the amount of heat evolved when the same
solvent is added to a solution of SbCl5 in DCE.
Thus, while a solvent of low donor number (basicity) has less
effect in the band maxima, others with higher values induce larger
red shift.
The positions of the d–d absorption bands of the perchlorate
complexes were in the range 521–597 nm, while for the nitrate
complexes these absorption band ranges from 614 to 708 nm. These
absorption bands suggest a square planar geometry for all per-
chlorate complexes and square pyramid (complexes 1 and 7) or
octahedral geometry for the nitrate complexes.
Solvatochromism appeared in [Cu(DEM)(tmen)]ClO4 (2),
[Cu(EAA)(tmen)(NO3)]•2H2O (7), [Cu(EAA)(tmen)]ClO4 (8) and
[Cu(EBA)(tmen)]ClO4 (14) complexes. These complexes exhibit
a color change from violet or bluish-violet through blue to
green according to the donor number (DN) of the solvent used.
Figs. 4 and 5 confirm the solvatochromic behavior in complexes 2,
7, 8 and 14.
Mixed-ligand Cu(II) complexes derived from three dioxygen
ester ligands (diethylmalonate, ethylacetoacetate and ethylben-
zoylacetate), and dinitrogen ligands (tmen, bipy and phen) with
nitrate, perchlorate and bromide anions, were prepared and char-
acterized by different spectroscopic techniques and TGA and molar
conductance and magnetic measurements. The perchlorate com-
plexes are square planar while the nitrate ones are either square
pyramidal or octahedral structure. In solution, four mixed-ligand
copper complexes showed solvatochromism. This behavior was
confirmed by applying the linear salvation free energy relation-
ship. The effect of anion was studied using t-Bu4N+Br− as a source
of bromide anion that replaced the perchlorate anions converting
the geometry of the complexes from mainly square planar to mainly
octahedral.
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