Copper(II) complexes of N,N-diisopropylpicolinamide - Solvatochromic and thermochromic phase change of a monomeric complex to a ferromagnetically coupled dimeric complex

Article abstract of DOI:10.1002/ejic.201300548

The synthesis and characterization of mononuclear [Cu(L) ₂(Cl)₂] (1), dinuclear [{CuCl(L)}₂(μ-Cl) ₂] (2), and another mononuclear complex [Cu(L)₂(ClO ₄)₂] (3) with a bidentate N,N-diisopropylpicolinamide ligand (abbreviated L) are described. Complexes 1 and 2 were obtained from the same reaction mixture and show thermo- and solvatochromic behavior. Complex 1 shows irreversible thermochromism by changing into 2 upon heating at 160 °C. These two complexes also show solvatochromism in various solvents. The coordination geometry of the metal ions in 1 and 3 is distorted octahedral, whereas 2 has Cu^II ions in a distorted square-pyramidal geometry. Variable-temperature magnetic susceptibility measurements for the dinuclear copper complex show a ferromagnetic interaction between the Cu^II ions. The chromotropism has been attributed to a change in geometry from octahedral with terminal chloride ligands to a distorted square-pyramidal coordination with bridging chloride ions. A monomer, [Cu(L)₂(Cl) ₂] (1, L = N,N-diisopropylpicolinamide), and a ferromagnetically coupled dimer, [Cu(L)(Cl)₂]₂ (2), are obtained from the same reaction mixture. They show thermo- and solvatochromic behavior owing to the change in coordination geometry from a terminally bonded octahedral copper center to a distorted square-pyramidal copper center with bridging chloride ions. Copyright

Full text of DOI:10.1002/ejic.201300548

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DOI:10.1002/ejic.201300548
Copper(II) Complexes of N,N-Diisopropylpicolinamide –
Solvatochromic and Thermochromic Phase Change of a
Monomeric Complex to a Ferromagnetically Coupled
Dimeric Complex
Isha Lumb,[a] Maninder Singh Hundal,[a] Montserrat Corbella,*[b]
Verónica Gómez,*[b] and Geeta Hundal*[a]
Keywords: N,O ligands / Copper / Thermochromism / Solvatochromism / Magnetic properties
The synthesis and characterization of mononuclear [Cu(L)2-
(Cl)2] (1), dinuclear [{CuCl(L)}2(μ-Cl)2] (2), and another mo-
nonuclear complex [Cu(L)2(ClO4)2] (3) with a bidentate N,N-
coordination geometry of the metal ions in 1 and 3 is dis-
torted octahedral, whereas 2 has CuII ions in a distorted
square-pyramidal geometry. Variable-temperature magnetic
diisopropylpicolinamide ligand (abbreviated L) are de- susceptibility measurements for the dinuclear copper com-
scribed. Complexes 1 and 2 were obtained from the same
reaction mixture and show thermo- and solvatochromic be-
havior. Complex 1 shows irreversible thermochromism by
changing into 2 upon heating at 160 °C. These two com-
plexes also show solvatochromism in various solvents. The
plex show a ferromagnetic interaction between the CuII ions.
The chromotropism has been attributed to a change in geom-
etry from octahedral with terminal chloride ligands to a dis-
torted square-pyramidal coordination with bridging chloride
ions.
state thermochromic complexes are very important owing
to their applicability as temperature sensors. Reversible
thermochromic complexes are increasingly used for textiles
and smart coatings.[2d] The irreversible ones can be useful
for the fabrication of end-of-life indicators (EOLs).[2e]
EOLs allow monitoring of the status of various systems and
devices, ranging from electronics to protective equipment,
in a noninvasive manner. Many mixed-ligand copper(II)
chloride complexes show fascinating thermochromic behav-
ior in the solid state[2f,2g] when small differences in their
electronic properties are reflected in the structures that they
prefer upon heating. Similarly, solvatochromism of metal
complexes is an attractive field because of their potential
ability to act as a Lewis acid–base color indicators, molecu-
lar switches, pollutant sensors, and imaging devices.[3] The
solvatochromism may be attributed to the binding of the
solvent to the metal center or to its effect on the ligand.
CuII is an ideal contender for the first kind as the Jahn–
Teller effect shifts the d–d transitions. However, the latter
gives very low intensities owing to the partially forbidden
nature of the d–d transitions. Thus, the preparation of sol-
vatochromic complexes with high molar absorptivity is
highly important.
Introduction
The wide range of colors exhibited by transition-metal
complexes is undoubtedly the most interesting aspect of
their coordination chemistry. However, it is more enchant-
ing to observe and explain the phenomena that lead to
color changes following some chemical or physical stimulus.
Thermochromism and solvatochromism are two such phe-
nomena and fall under the broader term chromotropism
and may or may not be overlapping in nature.[1] Thermo-
chromism has recently become a subject of interest owing
to its possible applications in imaging and sensing materi-
als.[2a,2b] A variety of thermochromic systems have been re-
ported, most of which show thermochromic properties only
in the solution state; few systems show thermochromic
functions in the solid state.[2b,2c] Solid-state thermochro-
mism can be discontinuous with an abrupt change of color
or continuous with a gradual shift of visible absorption
bands.[2b,2c] It may be reversible or irreversible in nature and
is governed by the thermodynamics of the system.[2d] Solid-
[a] Department of Chemistry, Centre for advance studies in
chemistry, Guru Nanak Dev University,
Amritsar 143005, India
E-mail: geetahundal@yahoo.com
On the other hand, dinuclear copper complexes with
hydroxo, alkoxo, and chloro bridges are interesting for the
study of magnetic exchange interactions.[4] Of these, com-
plexes with the Cu(μ-Cl)2Cu core[5] are the most fascinating
because they show a variety of geometries ranging from
square pyramidal to trigonal bipyramidal with diverse bond
Homepage: department.gndu.ac.in/department/chemistry.asp
[b] Departament de Química Inorgànica, Universitat de Barcelona
(UB),
Av. Diagonal 645, 08028 Barcelona, Spain
Homepage: http://www.ub.edu/inorgani/en/index.html
Supporting information for this article is available on the
WWW under http://dx.doi.org/10.1002/ejic.201300548.
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lengths (Cu–Cl) and angles (Cu–Cl–Cu). Thus, each case Complex
3
was prepared by
a
similar reaction of
requires a specific study to conclude meaningful magneto- Cu(ClO4)2·6H2O with L in methanol. Only one blue solid
structural correlations.[6] We previously reported a chloro- was obtained in this case.
bridged dinuclear complex of CuCl2 with N-substituted
pyridine-2,6-dicarboxamide, which forms a solvent-assisted
Infrared Spectra
2D sheet in the crystal structure and offers antiferromag-
netic coupling between the metal centers.[7] Presently, the
reaction of CuCl2 with N,N-diisopropylpicolinamide (L)
has given a mononuclear and a μ-Cl dinuclear complex
from one reaction mixture. The dinuclear complex shows
ferromagnetic coupling between the two centers. At the
same time, these two complexes exhibit thermo- and solva-
tochromism related to their coordination geometry. Here,
we wish to report these results and try to correlate them
with the molecular structures. Although the crystal struc-
tures of many complexes of N,N-dialkyl-substituted pyr-
idine-2,6-dicarboxamide and a few complexes of unsubsti-
tuted 2-picolinamide[8] are known, to the best of our knowl-
edge (based upon a CSD search), there is only one crystal
structure of a PdII ion with L.[9] As the ligand is expected
to behave as a bidentate ligand and chloride is known to
give the Cu(μ-Cl)2Cu core, we expected to get both the mo-
nomeric and dimeric complexes from the reaction of L with
copper(II) chloride.
The ligand shows characteristic IR bands at 2979 and
1627 cm–1, which are assigned to νC–H stretch and amide I
νC=O bands,[10a] respectively. The IR spectra of 1, 2, and 3
display bands at 1596, 1576, and 1584 cm–1, respectively,
attributed to νC=O stretching. On complexation, a shift in
the amide stretching frequencies is observed, owing to the
coordination of the ligand through the amide carbonyl oxy-
gen atom. The band in the range 730–550 cm–1 corresponds
–
to the Cu–Cl bond. The ClO4 anion in 3 has been charac-
terized by a large band at 1059 cm–1.[10b] The sharp doublet
–
observed indicates that the ClO4 anion is in the coordina-
tion sphere.
X-Ray Crystal Structures
Molecular and Crystal Structure of [CuL2Cl2] (1)
Complex 1 (Figure 1) contains two crystallographically
independent molecules in the unit cell. In both of them,
the CuII ions are on the center of symmetry and, thus, the
asymmetric unit contains half of both molecules. The coor-
dination environment around the metal ion is the same in
Results and Discussion
The ligand L was prepared by a mixed anhydride cou- both units; two L ligands act as bidentate chelating ligands
pling reaction (Scheme 1; see Exp. Section. and ref.[26]) in- and form five-membered chelate rings with the metal ions.
stead of the reported method with thionyl chloride.[9] Com- The two trans positions around each CuII ion are occupied
plexes 1 and 2 were obtained from the same reaction mix- by chloride ions. The Cu–O bonds in the two molecules are
ture (see Exp. Section). The initial product was a blue-green significantly longer [2.661(1), 2.543(1) Å] than the Cu–N
crystalline solid mixture of the two complexes, which were [2.002(1), 1.971(1) Å] and Cu–Cl bonds [2.279(5),
separated by repeated washings with ethanol and aceto- 2.313(5) Å; Table S1]. Thus, the stereochemistry is a highly
nitrile (ACN) owing to their different solubilities. Upon tetragonally distorted octahedron with the oxygen atoms of
slow evaporation of the solvent, the blue compound yielded the ligand farthest away. The two molecules are rotated
blue crystals of [Cu(L)2Cl2] (1) and the green compound through an angle of ca. 45° with respect to each other, and
gave yellow-green crystals of [{CuCl(L)}2(μ-Cl)2] (2). there are slight differences in the bonding parameters of the
Scheme 1.
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two molecules. The two L molecules in the two crystallo- N1; C21···centroid 3.510 Å, H21C···centroid 2.790 Å,
graphically independent, centrosymmetric molecules have Є132°), which extend it to a 2D arrangement.
different Cu–N and Cu–O bond lengths, and the ligand is
Molecular and Crystal Structure of [{CuCl(L)}2(μ-Cl)2]
(2)
strongly bonded in one of the molecules.
The molecular structure of 2 is a centrosymmetric dimer
with two edge-sharing square pyramids around the CuII
center (Figure 2, a). Cl1 unsymmetrically bridges the two
metal ions with Cu1–Cl1 and Cu1–Cl1i bond lengths of
2.2705(8) and 2.4188(8) Å (i = 1 – x, –y, –z + 1). The Cu1–
O1, Cu1–N1, and Cu1–Cl2 (terminal) bond lengths are
2.249(2), 2.003(2), and 2.229(2) Å, respectively (Table S1).
The longest interaction in this case is with the bridging
chloride ion (Cl1, Figure 2, b). The copper ion shows a
quite distorted square-pyramidal geometry (τ = 0.38) with
two bridging chloride ions. The nonbonding Cu–Cu dis-
tance is 3.421 Å. There are weak intramolecular C1···Cl2
interactions (Table S2), which form a five-membered chelate
ring. Owing to the steric hindrance from the isopropyl
groups and the three chloride ions around the metal ion,
the sixth coordination position remains unoccupied.
The crystal packing shows the formation of linear tetra-
Figure 1. ORTEP diagram of 1 with 30% thermal probability ellip-
meric chains parallel to the a axis owing to C10···Cl2 H-
soids. The second crystallographically independent molecule has
bonding interactions. Each loop in this chain consists of an
eighteen-membered ring (Figure S2A) formed by the inter-
actions between dinuclear entities, which generates a super-
amolecular entity containing four Cu2+ ions. These chains
are further extended along the b axis owing to C3···O1 H-
bonding interactions and π···π interactions between the cen-
trosymmetric pyridine rings (centroid to centroid 3.92 Å) to
complete a 2D structure. The C2···Cl1 and methyl C8···Cl2
interactions extend this into 3D H-bonded network (Figure
S2B).
not been shown. Hydrogen atoms have been omitted for clarity.
The crystal packing (Figure S1) shows stacking down the
b axis. Molecule 1 (containing Cu1 and Cl1) makes linear
chains along the a axis, owing to C3–H3···O1, C12–
H12A···Cl1, and C–H···π interactions between centrosym-
metrically related pyridine rings [C2–H2···centroid (C1–C5,
N1), 3.817, 3.62 Å, 95°; Table S2). Similarly, molecule 2
(containing Cu2 and Cl2) makes linear, centrosymmetric
chains parallel to the a axis held together by C15–H15···O2,
C16-H16···Cl2, C15-H15···Cl2, and π···π interactions be-
tween the pyridine rings (centroid to centroid of C13–C17,
N3 3.801 Å). These two kinds of chains are further con-
Molecular and Crystal Structure of [CuL2(ClO4)2] (3)
Complex 3 is a centrosymmetric molecule with the CuII
nected to each other through C20–H20B···O1, C21– ion at the center (Figure 3) and coordinated to two biden-
H21B···Cl1, and C21–H21C···π interactions between the tate chelating ligands that form a square-planar arrange-
methyl group and the centroid of the pyridine ring (C1–C5, ment about the metal ion. The Cu1–N1 [1.971(1) Å] and
Figure 2. (a) Edge-sharing coordination polyhedra of 2. (b) ORTEP diagram with 30% thermal probability ellipsoids. Hydrogen atoms
have been removed for clarity.
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Cu1–O1 [1.948(1) Å] bond lengths are significantly shorter the differential scanning calorimetry (DSC) curve (ΔH =
than those found in 1 and 2 (Table S1), especially the latter. 66.228 Jg–1). The DSC curve for 1 shows a sharp endother-
The perchlorate ions occupy the remaining two apical coor- mic peak at 160 °C (Figure S5). As the solid-state structures
dination positions around the metal ion and form rather of both 1 (blue) and 2 (green) are described above, this
long bonds [Cu–O4 2.458(1) Å] that complete a “z-out” te- thermochromism is attributed to the change of the geome-
tragonally distorted octahedron. Owing to the presence of try of the complex from monomeric, distorted octahedral
the perchlorate ions, complex 3 displays C–H···O hydrogen to dimeric square pyramidal with two bridging Cl– ions.
bonding in the crystal structure (Figure S3). C5···O5, No thermal studies were done for complex 3 owing to the
C4···O5, and C4···O2 H-bonding interactions (see Table S2) explosive nature of perchlorate ions.
create a double-helical-type crystal structure as viewed
down the a axis. This is further strengthened by phenylene
and methyl C–H···O and slightly longer C12–H12C···π
(centroid of py, C12···centroid 4.1 Å) interactions to form a
3D network.
Figure 4. TG-DTA plots of (a) 1 and (b) 2 between 20 and 800 °C.
Figure 3. ORTEP diagram of the 3 with 30% thermal probability
ellipsoids. Hydrogen atoms omitted for clarity.
Thermal Properties of 1 and 2 and Irreversible
Thermochromism of 1
The simultaneous thermogravimetric analysis and
differential thermal analysis (TG-DTA) curves of 1 and 2
are shown in Figure S4, and Figure 4 shows the thermo-
gravimetric analysis (TGA) curves of 1 and 2. Complex 1
Figure 5. Photographs of 1 revealing its irreversible thermochro-
mism. A color change from blue 1 to green 2 occurs at 160 °C.
begins to lose weight above 100 °C. In the first step, weight
loss corresponding to the loss of one L ligand is observed
at 165.19 °C (calcd. weight loss 37.62%; observed weight
loss 39.7%). The loss of another organic molecule occurs
at 229.47 °C (calcd 40%; observed 40.2%). Two chloride
ions are lost at 270.3 °C (calcd 12.90%; observed 15.38%)
to leave CuO only. The TGA and DTA curves for 2 re-
Samples 1a and 1b are intermediate stages before conversion of 1
into 2 in the solid state.
UV/Vis Spectra and Solvatochromism
The solid-state molecular structures are also well repre-
semble those of 1 beyond 200 °C, and there is no weight sented in the solution state, subject to the choice of the
loss below 200 °C. It is inferred from the comparison of solvent. The electronic spectra of the complexes were mea-
these curves that by losing one L molecule, complex 1 forms sured in methanol and acetonitrile for 1 and 2 and dimethyl
[CuLCl2], which is half of dimer 2. However, the dimeric sulfoxide/methanol (1:9) for 3. In methanol (DMSO/meth-
complex 2, which may be written as 2[CuLCl2], is stable anol for 3), all three complexes exhibit a single symmetrical
and does not undergo any weight loss up to 200 °C. There- broad d–d band in the region 700–850 nm (Figure 6, a),
after, it begins to lose weight similarly to 1 at comparable characteristic of a tetragonal geometry owing to ligand-
temperatures, again in two steps. This implies that 1, after field and Jahn–Teller effects around the Cu2+ ions. Bands
the loss of one ligand, dimerizes through one of the chloride at 794 nm (12594 cm–1, ε = 47 m–1 cm–1) for 1, 813 nm
ions and changes into 2 between 150–200 °C. Figure 5 (12300 cm–1,
ε
=
69 m–1 cm–1) for 2, and 774 nm
shows the color changes observed on heating 1, which be- (12919 cm–1, ε = 26 m–1 cm–1) for 3 are observed in the vis-
comes green after 160 °C and transforms into 2. The en- ible region (Figure 6, a). In general, these bands are as-
thalpy change for this transformation was calculated from signed[11] to the overlapping transitions dxz/dyz Ǟdx2–y2
,
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d
xy Ǟdx2–y2, and dz2 Ǟdx2–y2. The solution of 2 in methanol not be the same but are quite similar. Similarly, solutions
turns blue, as is indicated from the similarity of the spec- of both 1 and 2 in acetonitrile are yellow-green, and their
trum to that of 1, which implies that in methanol 2 is absorption spectra are also similar (Figure 6, b). This indi-
changed into a six-coordinate complex by breaking of the cates that in this solvent 1 behaves like 2, that is, it has a
labile Cu–Cl(bridging) bond and coordination of a solvent bridged structure through a chloride ion. Both of them
molecule (Scheme 2). This may result in a mononuclear show a very broad and weak band centered at ca. 800 nm
complex [CuLCl2(MeOH)2], which is very similar to 1. The (d–d band) and strong bands with shoulders at 383 and
small difference (428 cm–1) in the 10 Dq values (Table 1) of 454 nm for 1 and 380 and 454 nm for 2. The latter bands
1 and 2 in methanol indicates that the two complexes may along with their shoulders have been assigned as intense,
Figure 6. (a) Visible spectra of 1, 2, and 3 in methanol. (b) UV/Vis spectra of 1 and 2 in acetonitrile showing CT bands. (c) UV/Vis
spectra of 1 and 2 in acetonitrile showing d–d bands.
Scheme 2.
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Table 1. The donor numbers of solvents and electronic spectra data of 1 and 2 in various solvents.[a]
Solvent
DN
λ(d–d) [nm] (cm–1)
ε(d–d) [m1 cm1]
λ(CT) [nm] (cm–1)
ε(CT) [m–1 cm1]
1
2
3
1
2
3
1
2
1
2
DCM
Chloroform 0.00
Acetonitrile 14.1
EtOH
MeOH
DMF
0.00
709 (14104)
685 (14598)
785 (12739)
824 (12136)
795 (12579)
940 (10638)
930 (10753)
709 (14104)
702 (14245)
815 (12270) 698(14327) 52
845 (11834) 68
823 (12151) 801(12484) 46
951 (10515) 785(12739)
941 (10627) 901(11099) 19
#
#
55
97
149
42
107
156
67
9
43
#
#
94
*
30
23
50
404 (24752) 403 (24814) 708
414 (24155) 410 (24390) 634
458 (21834) 458 (21834) 112
2098
423
408
–
–
159
–
22.9
23.3
26.6
29.8
*
–
–
–
–
–
–
9
440 (22727) 438 (22831) 158
DMSO
–
–
–
[a] # insoluble, * partially soluble, – no CT band observed.
low energy, Cl ligand-to-metal charge-transfer (LMCT) which broadly is the order of their increasing donor
transitions.[12a] This may occur if the most labile Cu–O number. The proximity of polar solvent molecules with high
bonds in 1 are broken by solvation and the resulting four- donor number to the axial position causes a destabilization
coordinate complex dimerizes to give a Cu2–μ-Cl2 chromo- of the electrons in the dz2 orbital of the CuII ion and results
phore (Scheme 2). These charge-transfer (CT) bands along in a redshift for the dz2 to dx2–y2 transition region. Highly
with the d–d bands impart yellow-green color to the com- coordinating solvents tend to break the most labile Cu–Cl
plex and its solution. From these observations, it may be bridging bond in 2 and form monomeric distorted-octahe-
inferred that 1 and 2 show solvatochromism that depends dral complexes by coordinating to the metal ion
on the solvent used. The halogen-to-CuII LMCT transitions (Scheme 2). On the contrary, solvents with low DNs solvate
occur at low energy and they appear relatively independent the organic ligand more and encourage bridging through
of stereochemistry.[12b–12d] In general, six-coordinate tetrag- the Cl– ions to form dimers. The most intense green color
onal CuII complexes with long axial Cu–Cl bonds do not is obtained in DCM, which has the lowest DN. As crystals
show CT absorption from such axially ligated ligands. Thus, of 1 grown from DCM have the same dinuclear structure
strong CT absorption in a tetragonal CuII complex is an as 2 and also because DCM is the least coordinating in
indicator of the equatorial binding of the halogen li- nature, the spectra of 1 and 2 in DCM (or chloroform) may
gand.[12a] The blue, monomeric 1 does not show any CT be considered as representing dimeric structures similar to
bands, but the yellow-green complex 2 with bridging chlor- that of 2 in the solid state and have the highest 10 Dq values
ide atoms at the equatorial positions does show these bands, and strong CT bands. It may be inferred from the spectro-
albeit as shoulders.
scopic data that in the polar protic solvents such as meth-
To investigate the matter further, the spectra of both anol and ethanol, the weakest bridging Cl bond is broken
complexes were recorded in different solvents, and the col- and an octahedral environment is created around the metal
ors of their solutions were noted. As shown in Figure 7, ion with a possible coordination from the solvent. A con-
both 1 and 2 give blue solutions in polar protic solvents siderable redshift of the d–d bands in DMSO and DMF
(methanol and ethanol), yellow-green solutions in acetoni- makes the solutions almost colorless; a slight yellow color
trile, green solutions in nonpolar solvents such as chloro- in DMF may be due to its own (intraligand) π–π* transition
form and dichloromethane (DCM), and almost colorless (a very weak shoulder at 438 nm). With acetonitrile, a sol-
solutions in aprotic polar solvents such as N,N-dimethyl- vent of moderate donor strength, the two structures seem to
formamide (DMF) and dimethyl sulfoxide. The blue solu- be in equilibrium, and the solution has a green-blue color.
tions have no CT bands, the yellow-green solutions have However, this equilibrium shifts totally towards green in the
shoulders as CT bands, and the green solutions have well- solid state on crystallization. To further investigate this
formed CT bands (Figure S6). The energies of the d–d point, the variable-temperature spectra of 1 in acetonitrile
bands observed (Table 1) for the two complexes in solvents were recorded.
of different donor number (DN)[13] decreases in the order
With an increase in temperature, the CT band at 458 nm
chloroform ≈ dichloromethane Ͼ acetonitrile Ͼ methanol becomes more intense (Figure 8, a). This change is linear
Ͼ ethanol Ͼ dimethyl sulfoxide Ͼ N,N-dimethylformamide, with temperature (Figure 8, b), as is clear from the inset of
Figure 7. Colors of 0.001 m solutions of (a) 1 and (b) 2 in various solvents.
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Figure 8. (a) Change in the electronic spectrum of 1 in acetonitrile on heating. (b) Change in the absorption of the CT band of 1 in
acetonitrile on heating; inset: a plot of the change in absorption vs. temperature.
Figure 8 (b), which shows a plot of absorbance vs. tempera- +1.7 to –0.7 V, and a AgCl/Ag reference electrode. All three
ture at this wavelength. As the CT band is due to chloride complexes display very similar quasireversible to irrevers-
bridging, at ambient temperature in acetonitrile both 1 and ible redox waves. Complex 3 (Figure 9, c) with E1/2
2 are present, and the amount of 2 increases as the tempera- +0.661 V and ΔEp = 0.716 V should represent a redox cou-
ture increases. ple in a tetragonally distorted monomeric copper complex
=
The absorption spectra of 3 in ACN, methanol, DMF, with a CuN2O2(ACN)2 core in acetonitrile. Redox peaks in
and DMSO represent a tetragonally distorted octahedral the range +1.5 to 0.25 V are attributed to a CuIII/CuII cou-
structure, and the solutions are blue only (Figure S7). The ple, and those in the range –0.5 to –1.2 V are attributed to
most labile Cu–O(perchlorate) bond (see crystal structure) CuII/CuI and CuI/Cu0 couples.[14–17] Therefore, the values
on the axial position is replaced by a molecule of the coor- observed for 3 represent a CuIII/CuII redox couple. Interest-
dinating solvent. As explained above, the d–d band shifts ingly, the CV traces of 1 and 2 are similar to each other
towards lower energy as the donor strength of the solvent and partially similar to that of 3. As discussed above, in
increases (Table 1). Although the complex is insoluble in acetonitrile, both 1 and 2 are supposed to have similar
DCM and chloroform, the absence of any CT band in ace- structures in which both the monomer and the chloro-
tonitrile indicates that 3 does not show solvatochromism of bridged dimer are in equilibrium in solution. Therefore,
the kind shown by 1 and 2. This further ratifies that a Cu2– they show similar redox behavior with two irreversible re-
μ-Cl2 chromophore is responsible for this effect in 1 and 2. dox waves, which indicates that the CuII ions are in two
To rule out the possibility of CT bands from the solvent environments. The E1/2 and ΔEp values for the quasirevers-
itself, solid-state reflectance spectra of 1 and 2 were re- ible waves are E1/2 = 0.780 V, ΔEp = 1.10 V and E1/2
=
corded (Figure S8). Although the range of the instrument 0.850 V, ΔEp = 0.983 V for 1 and 2, respectively (Figure 9,
did not allow measurements of the d–d bands, the spectra a and b). Therefore, in all three complexes this may corre-
show CT bands near 450 nm for 2 and also for 2b, obtained spond to a CuIII/CuII couple in a tetragonal monomeric
after heating 1 at 160 °C, which are conspicuously absent in complex. In all three, ΔEp is much greater than 0.0580 and,
1. These spectra also confirm that the solid-state thermo- therefore, they are irreversible processes. The small differ-
chromism involves a transformation from 1 to 2 at high ences in their parameters may be due to the different coor-
temperature, as indicated by the thermal studies and the dination cores: {CuN2Cl2(ACN)2}, {CuNOCl2(ACN)2},
colors of the compounds obtained. The thermochromically and {CuN2O2(ACN)2} in 1–3 in solution, respectively.
obtained yellow-green species retains its identity in the solid Apart from this, both 1 and 2 show another irreversible
state and reverts only on recrystallization from methanol/ oxidative response at higher positive anodic potentials (Epa)
ethanol. The spectra of 2 formed after heating 1 in meth- of 0.866 and 1.009 V, respectively. This has been assigned
anol and ACN is in accordance with the above-described to the CuIII II
/
redox system for the CuII ion in a square-
solvatochromism effects (Figure S9).
pyramidal chloro-bridged environment and E1/2 = 0.701 V,
ΔEp = 0.330 V for 1 and E1/2 = 0.801 V, ΔEp = 0.403 V for
2. Thus, it may be inferred from these data that in aceto-
nitrile there is a mixture of monomeric and dimeric struc-
Cyclic Voltammetric Studies
Cyclic voltammetry (CV) studies were performed in dry tures.[9] To further investigate, cyclic voltammograms of
CH3CN with a Pt working electrode (Figure 9), Fc+/Fc as both 1 and 2 were recorded in DCM (Figure 9, e and f),
reference potential, tetrabutylammonium perchlorate and both showed only one redox wave. The ΔEp values of
(TBAP) as supporting electrolyte in the potential range 0.457 and 0.446 V for 1 and 2, respectively, are quite close
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Figure 9. Cyclic voltammograms of (a–c) 1–3, (d) a mixture of 1 and 2 in CH3CN at 50 mV/s, and (e–f) 1–2 in DCM.
to those found for the redox couples in acetonitrile assigned the homogeneity of the compounds (Figure S10). Pattern 4
to the chloro-bridged CuII species. Finally, a cyclic voltam- in Figure S10 represents a powder pattern taken for a solid
mogram of a mixture of 1 and 2 in acetonitrile (Figure 9, obtained immediately after evaporation of the solvent (at
d) shows two peaks, similar to those observed for them room temperature) from a solution of 1 in acetonitrile. The
separately in acetonitrile (Figure 9, a, b). The respective val- pattern is close to that of 2 but also shows a few peaks
ues of E1/2 = 0.800 V, ΔEp = 0.400 V and E1/2 = 0.850 V, matching those of 1. For example, (0,0,2)* at θ = 6.8° repre-
ΔEp = 1.01 V for 1 and 2, respectively, fully endorse the sents the first strong peak in the powder pattern of 1. No
observation that in acetonitrile these are present in equilib- such peak is present in the XRD pattern of 2 in the pure
rium. Easy oxidation has been observed previously in a CuII state. Therefore, a solution in acetonitrile has a mixture of 1
species bonded to an oxygen atom from an amide[18] formed and 2. The XRD patterns of samples of 1 heated at various
from a primary amine with no deprotonation of the amide temperatures are presented in Figure 10. Firstly, the PXRD
nitrogen atom. The situation is quite similar to the present pattern of 1 was recorded at room temperature (the lowest
case, for which no deprotonation is possible. As the dinu- black colored line). Then, complex 1 was heated at different
clear complex is a centrosymmetric complex with both CuII fixed temperatures (70, 100, 120, 140, and 160 °C) in a pro-
ions in the same environment, it is logical to infer that both grammable oven, and PXRD patterns were recorded for the
of them show only one redox potential.
heated samples. These PXRD patterns were then compared
with the generated PXRD patterns for 1 and 2. The PXRD
pattern of a sample heated at 160 °C matched well with that
generated for 2 and demonstrated the complete conversion
of 1 into 2. The (0,0,2) reflection at θ = 6.64° represents the
Powder XRD Studies
The measured powder XRD (PXRD) patterns for 1–3
match their generated powder patterns exactly, which shows
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first strong peak in the powder pattern for 1. This peak is state. The experimental data were fitted by the Bleaney–
also observed in the XRD pattern taken at 70 °C. The Bowers equation[20] with the spin Hamiltonian H = –JS1S2
(1,0,–1) peak at θ = 9.44° and the (0,2,0) peak at θ = 16.4° and by considering the presence of intermolecular interac-
(characteristic of 2) are also present, which shows that the tions (zJЈ). The temperature-independent paramagnetism
samples contains a mixture of both 1 and 2 at 70 °C. At was fixed to 45ϫ10–6 cm–3 mol–1. The best fit (R =
100 °C, the (0,0,2) peak at θ = 6.64° disappears. The 7.2ϫ10–6) was obtained with J = 43.1 cm–1, zJЈ = –0.21,
(1,0,–1) peak at θ = 9.44°, the (0,0,2) peak at θ = 13.3°, and and g = 2.14.
the (2,0,0) peak at θ = 18.7° became more prominent at 120
and 140 °C. Finally, at 160 °C, 1 was converted into 2, and
the powder pattern matched the generated powder pattern
of 2. Therefore, it has been proved that the monomer con-
verts into the dimer upon gradual heating in the solid state.
Figure 11. χMT vs. T and M vs. H (inset) plots for 2. The solid line
is the best fit to the experimental data.
Several studies have been performed to achieve magneto-
structural correlations for compounds with a [Cu2(μ-Cl)2]2+
core; the value of the magnetic coupling constant has been
Figure 10. Comparison of the powder XRD patterns of 1 [experi-
related to different structural parameters such as the Cu–
mental at room temperature (black), generated (red) and at dif-
ferent temperatures] and 2 with the pattern generated from its crys- Cl–Cu angle and the Cu–Cl distance.[21] Hatfield and co-
tal structure.
workers proposed an empirical correlation between the
magnetic coupling constant and the α/R ratio (α is the Cu–
Cl–Cu angle and R is the longest Cu–Cl bridge dis-
tance).[21b] They found that for a value of α/R lower than
Magnetic Properties
32.6 or higher than 35.8 °Å–1, the magnetic interaction is
The magnetic susceptibilities of the CuII complexes 1 and
antiferromagnetic, and for values between these limits, the
3 were determined in the solid state at room temperature.
The μeff values of 1.83 and 1.98 μB for 1 and 3, respectively,
lie within the range normally found for CuII complexes and
show that they have S = 1/2 states.
interaction is ferromagnetic. According to this correlation,
compound 2 with an α/R value of 38.7 °Å–1 should show a
J value of –65 cm–1, which is completely opposite to the
experimental one (43.1 cm–1).
For this kind of compound, it is quite difficult to estab-
lish simple magneto-structural correlations because their
Variable-Temperature Magnetic Susceptibility of 2
The magnetic susceptibility of 2 was recorded from 300 magnetic properties vary drastically depending on the rela-
to 2 K, and the χMT vs. T plot is shown in Figure 11. At tive geometry of the copper ions, the distortion from the
300 K, the χMT value is 0.92 cm3 mol–1 K, slightly higher ideal geometry, and the orientation of the two CuII polyhe-
than expected for two uncoupled CuII ions (S = 1/2, χMT dra.[22] As has been indicated, the geometry of compounds
= 0.75 cm3 mol–1 K, assuming g = 2). As the temperature with five-coordinate metal ions can be square pyramidal
decreases, the χMT values increases to a maximum of χMT (SP), trigonal bipyramidal (TBP), or distorted with dif-
= 1.12 cm3 mol–1 K at 15 K. This behavior is characteristic ferent disposition of coordination polyhedra. Three dif-
of ferromagnetic coupling, and this is confirmed by the field ferent SP orientations for dinuclear complexes with [Cu2(μ-
dependence of the magnetization at 2 K, which shows M/ Cl)2]2+ cores are reported in the literature: (1) square pyra-
Nμβ values indicative of two unpaired electrons (Figure 11, mids that share one base-to-apex edge with the two bases
inset). The M/Nμβ vs. H plot was fitted with the Brillouin nearly perpendicular to one another (SP-I), (2) square pyra-
equation for a ground spin state S = 1, and a g value of 2.06 mids that share one base-to-apex edge with parallel basal
was obtained.[19] The decrease observed at low temperature planes (SP-II), and (3) square pyramids that share a basal
could be caused by intermolecular antiferromagnetic inter- edge with coplanar basal planes (SP-III).[22] For CuII ions
actions or the zero-field splitting (ZFS) of the S = 1 ground with TBP geometry, two situations are reported: (1) the
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Table 2. Selected structural parameters and magnetic data for compounds with a [Cu2(μ-Cl)2]2+core and coplanar magnetic orbitals.
Compound[a]
α[b] [°]
σ[c] [°]
τ[d]
J[e] [cm–1]
Ref.
[24a]
[24b]
[24c]
[21a]
[Cu2(μ-Cl)2L1 Cl4] (A)
97.1
95.8
95.3
94.5
93.6
5.72
0
13.84
7.81
34.9
0.10
0.005
0.36
0.15
0.38
–58.4
–30.2
–3.7
33.7
43.1
2
[Cu2(μ-Cl)2(morphH)2Cl4] (B)
[Cu2(μ-Cl)2L2 Cl2] (C)
2
[Cu2(μ-Cl)2{HB(l-pz)3}2] (D)
[Cu2(μ-Cl)2(L)2Cl2] (2)
this work
[a] Ligand abbreviations: L1 = o-bis(amido)-appended dimethylthiotetrathiafulvalene; morph = morpholine; L2 = N,N,NЈNЈ-tetraisoprop-
ylpyridine-2,6-dicarboxamide; HB(l-pz)3 = hydrotris(1-pyrazolyl)borate. [b] Cu–Cl–Cu angle. [c] Bifold angle. [d] Parameter of distortion
from ideal geometry in pentacoordinate ions (τ = 0 for SP and τ = 1 for TBP). [e] Magnetic exchange coupling parameter based on the
Hamiltonian H = –JS1S2.
chloride bridges are in axial and equatorial positions (TBP- small and sharp peak is observed at g ≈ 2, probably owing
I) and (2) both chloride bridges are in equatorial positions to an impurity (Figure S11). A shoulder at ca. 1700 G is
(TBP-II).[21b] In 2, the copper ions show a quite distorted also observed and could be assigned to the half-field transi-
square-pyramidal geometry (τ = 0.38; τ = 0 for SP and τ = tion ΔMs = Ϯ2.
1 for TBP geometries)[23] with a coplanar arrangement of
the pyramids (SP-III). The magneto-structural correlation
mentioned above was performed for CuII ions with SP-II
and TBP-I geometries and, consequently, it is not appropri-
ate for 2.
Conclusions
Two compounds of copper(II)chloride with N,N-di-
isopropylpicolinamide (L) were synthesized from the same
mixture and have different structures; 1 is mononuclear, and
2 is dinuclear. The irreversible thermochromism shown by
1 is attributed to a phase change to 2. This phase change is
triggered by the loss of one ligand L from 1 on heating,
followed by dimerization and bridging by two chloride ions.
Both 1 and 2 show solvatochromism based on the geometry
around the CuII center, the presence of a Cu2–μ-Cl2 moiety,
and the polarity of the solvent. In nonpolar solvents, the
low energy LMCT transitions of the bridging chloride ions
in the equatorial positions are responsible for the solvato-
chromism. As the color change is based on strong CT tran-
sitions, the solvatochromism is highly intense. The dinuclear
complex shows ferromagnetic coupling between the two
copper centers, which is favored by the distortion from a
SP geometry to a TBP geometry. The corresponding per-
chlorate complex 3 does not show any chromotropism.
The most relevant structural parameters and the mag-
netic coupling constants for 2 and analogous compounds
with the same basal-plane arrangement are collated in
Table 2. For these compounds, the single electron of each
copper ion is mostly in the dx2–y2 orbital (basal), and the
interaction between the magnetic orbitals of the two CuII
ions will occur through the central chloride ions through a
σ* interaction. As can be seen, there is a wide range of
J values from –58 to 43 cm–1. With this magnetic orbital
disposition, the magnetic interaction should be very sensi-
tive to the Cu–Cl–Cu angle (α), and the antiferromagnetic
coupling decreases as this angle approaches 90° (accidental
orthogonality). As can be seen in Table 2, the crossover
point would be ca. 95° for these compounds. However, com-
pounds B and C show very similar α angles and different
magnetic behavior; thus, other factors should also be con-
sidered. In this case, the distortion from a SP geometry to
a TBP geometry in compound C favors the ferromagnetic
contribution, as has been reported previously.[22b] This ef-
fect is also observed for 2. Moreover, another important
Experimental Section
factor is the bifold angle (σ); this is the angle formed by Materials and Methods: All reactions were performed in anhydrous
solvents. The solvents were distilled prior to use (THF with Na
wire) by following appropriate procedures.[27] Anhydrous copper(II)
chloride was prepared by treating a powdered sample of hydrated
CuCl2 with SOCl2 and heating the mixture to reflux for a few hours
until the evolution of SO2 ceased. The anhydrous product was
washed with petroleum ether and dried in vacuo. Elemental analy-
ses (C, H, N) were performed with a Perkin–Elmer model 2400
CHN analyzer. IR spectra were recorded with samples as KBr pel-
lets with a Varian 660-IR FTIR spectrophotometer. Thermal
analyses were performed with a Perkin–Elmer Diamond TG/DTA
thermo analyzer in N2 flow from 20–800 °C with a heating rate
of 10 °Cmin–1. DSC was performed with a Perkin–Elmer DSC 7
two planes: one containing the copper centers and the basal
ligands of the core and the lateral plane defined by the Cl1,
Cu1, N1, and Cl2 atoms in this case. A ferromagnetic inter-
action has been observed[25] at bifold angles greater than
25°. Therefore, the existence of a ferromagnetic coupling in
2 is consistent with these results.
EPR Spectra
The electron paramagnetic resonance (EPR) spectra of
powdered samples of 2 were recorded at different tempera-
tures. At room temperature, the spectrum shows a band
centered at g ≈ 2.14 with a peak-to-peak line width of ΔHpp
= 400 G. Below 100 K, the most intense band appears at g
≈ 2.26 with ΔHpp ≈ 400 G; at these temperatures, a new
1
instrument. The H NMR spectrum of the ligand L was recorded
with a 300 MHz JEOL FT NMR spectrometer with tetramethylsil-
ane (TMS) as the reference compound. UV/Vis spectra were re-
corded with a Shimadzu Pharmaspec UV-1700 UV/Vis spectropho-
tometer. Variable-temperature UV/Vis studies were performed with
band becomes visible at lower field (ca. 1000 G), and a the same spectrophotometer by circulating hot water at a fixed tem-
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perature. Solid-state UV spectra was recorded with a Shimadzu
UV-2450 UV/Vis spectrophotometer. Molar conductance values of
millimolar solutions were measured with a conductivity bridge digi-
tal conductivity meter (Systronics 306). The XRD was performed
with a Rigaku Miniflex 2 diffractometer in the range θ = 2–60°.
The variable-temperature XRD were recorded after samples of 1
had been heated at different fixed temperatures in a programmable
oven. Magnetic susceptibility measurements (2–300 K) and magne-
tization measurements at 2 K (0–50000 G) were performed with a
Quantum Design MPMP superconducting quantum interference
device (SQUID) magnetometer at the Unitat de Mesures
Magnètiques (Universitat de Barcelona). Two different magnetic
fields were used for the susceptibility measurements, 500 (2–5 K)
and 10000 G (2–300 K), with superimposable graphs. Pascal’s con-
stants were used to estimate the diamagnetic corrections for the
compound. The fit was performed by minimising the function R =
Σ[(χMT)exp – (χMT)calc]2/Σ[(χMT)exp]2. Cyclic voltammetric measure-
ments were performed with a BAS CV 50W electrochemical analyz-
ing system (accuracy 1.0 mV). The cyclic voltammograms were re-
corded in acetonitrile with 0.1 m TBAP as the supporting electro-
lyte and a three-electrode system under a N2 atmosphere. A three-
electrode configuration composed of a Pt working electrode, a Pt
wire counterelectrode, and a Ag/AgCl reference electrode was used
for the measurements, which were standardized to the ferricinium/
ferrocene redox couple (E1/2 = 0.440 V, ΔEP = 0.052 V vs. Pt/PtCl2
electrode). Solid-state diffuse-reflectance UV/Vis spectra were by
using 10% compound in a BaSO4 pellet. Solid-state EPR spectra
were recorded from room temperature to 4 K at X-band (9.4 GHz)
frequency by using a Bruker ESP-300E spectrometer. High-resolu-
tion mass spectra were recorded with a Bruker microTOF-QII
spectrophotometer.
bath to solidify. The creamish solid obtained was recrystallized
with heptane to afford transparent colorless crystals (yield 0.23 g,
70%), m.p. 70 °C. C12H18N2O (206.29): calcd. C 69.90, H 8.73, N
13.59; found C 69.83, H 8.82, N 13.65. IR: ν = 1627(CO) cm–1. 1H
˜
NMR (300 MHz, CDCl3, Me4Si): δ = 8.58 (d, 1 H, Py), 7.75 (t, 1
H, Py), 7.44 (d, 1 H, Py),7.29 (dd, 1 H, Py), 3.83 (m, 1 H, NCH),
3.55 (m, 1 H, NCH), 1.579 [d, 6 H, CH(CH3)2], 1.215 [d, 6 H,
CH(CH3)2] ppm. 13C NMR (75 MHz, CDCl3): δ = 168.6 (C=O),
156.5, 148.6, 136.8, 123.4, 121.7 (Py), 50.6 (NCH), 45.9 (NCH),
20.6 [CH(CH3)2], 20.4 [CH(CH3)2] ppm.
[Cu(L)2Cl2] (1) and [{CuCl(L)}2(μ-Cl)2] (2): These complexes were
prepared from a reaction mixture containing anhydrous CuCl2
(dried with SOCl2, 0.03 g, 0.22 mmol) and L (0.164 g, 0.4 mmol)
dissolved in a minimum amount of dry methanol. The contents
were stirred for 5–6 h and then kept at room temperature for 2–
3 d. The product obtained was a blue-green crystalline solid mix-
ture (yield 0.096 g, 79.1%), which was separated by washings with
ethanol as the green compound is insoluble in ethanol at ambient
temperature. Thus, the blue compound was extracted (twice) in eth-
anol, and the green residue was dissolved in acetonitrile to yield
the green compound. The blue complex yielded blue crystals of
[Cu(L)2Cl2] (1) upon slow evaporation from ethanol, and the green
complex gave yellow-green crystals of [{CuCl(L)}2(μ-Cl)2] (2) from
acetonitrile.
1: M.p. 174 °C, yield 0.032 g (33.3%). C24H36Cl2CuN4O2 (547.03):
calcd. C 52.6, H 6.58, N 10.24; found C 52.86, H 6.13, N 10.14.
IR: ν = 1596, 1349 cm–1. Molar conductance (Ω–1 cm2 mol–1): 39.10
˜
(CH3OH). μeff = 1.83 μB.
2: M.p. 220 °C, yield 0.046 g (47.9%). C24H36Cl4Cu2N4O2 (681.48):
calcd. C 47.21, H 5.90, N 9.18; found C 47.39, H 6.01, N 9.93. IR:
N,N-Diisopropylpicolinamide: The ligand L was prepared by a
mixed anhydride coupling reaction,[26] instead of the reported
method with thionyl chloride.[9] To a stirred solution of picolinic
acid (0.2 g, 1.6 mmol) in dry tetrahydrofuran (THF) at –15 °C was
added N-methyl morpholine (0.232 mL, 2.1 mmol) and ethyl
chloroformate (0.232 mL, 2.4 mmol). The mixture was stirred for
15 min, N,N-diisopropylamine (0.250 mL, 1.78 mmol) was added,
and the mixture was stirred for 5 h. The THF was evaporated, and
the product was extracted with ethyl acetate and water. The product
was further extracted with a saturated sodium carbonate solution,
a 5% citric acid solution, and brine. The solution was dried with
sodium sulfate, concentrated under vacuum, dried, and put in ice
ν = 1576, 1445, 1360 cm–1. Molar conductance (Ω–1 cm2 mol–1): 41
˜
(CH3OH). μeff = 2.71 μB.
[Cu(L)2(ClO4)2] (3): Cu(ClO4)2·6H2O (0.370 g, 1 mmol) and L
(0.412 g, 2 mmol) were dissolved in methanol (20 mL) with con-
tinuous stirring. The blue solid that precipitated from the reaction
mixture after mixing was collected by filtration and dried in vacuo.
Blue crystals formed in the filtrate as well as by recrystallization
from acetonitrile after a few days. C24H36Cl2CuN4O2 (547.03):
calcd. C 42.69, H 5.3, N 8.3; found C 42.78, H 5.8, N 7.96. IR: ν
˜
=
1584, 1447, 1369, 1059, 626 cm–1. Molar conductance
(Ω–1 cm2 mol–1): 49 (CH3OH). μeff = 1.98 μB.
Table 3. Crystallographic data for 1–3.
1
2
3
Empirical formula
Formula weight
Crystal size [mm]
Crystal system
Space group
a [Å]
C24H36Cl2CuN4O2
547.01
C24H36Cl4Cu2N4O2
681.45
0.18ϫ0.18ϫ0.16
monoclinic
P21/n
10.2909(7)
10.7794(8)
C24H36sCl2CuN4O10
675.01
0.20ϫ0.16ϫ0.14
monoclinic
P21/n
8.237(5)
14.275(5)
0.16ϫ0.14x0.10
triclinic
¯
P1
7.8700(2)
9.8014(3)
b [Å]
c [Å]
α [°]
18.3066(4)
92.780(2)
13.942(1)
90.00
2.989(3)
90.00
β [°]
98.162(2)
105.906(3)
1338.61(5)
106.926(5)
90.00
1479.58(18)
94.053(5)
90.00
1523.5(11)
γ [°]
V [Å3]
Z
F(000)
2
574
2
700
2
702
Reflections
Final R indices
R indices (all data)
5266/4429
R1 = 0.0385, wR2 = 0.0780
R1 = 0.0298, wR2 = 0.0757
2484/2015
R1 = 0.0410, wR2 = 0.0766
R1 = 0.0290, wR2 = 0.0705
3741/3350
R1 = 0.0328, wR2 = 0.0820
R1 = 0.0286, wR2 = 0.0790
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Caution! Perchlorate salts are explosive in nature. They can explode
upon heating and should be handled with caution and in small
amounts.
X-ray Data Collection: The data for 1, 2, and 3 were collected with
a Bruker Apex-II CCD diffractometer. The data were corrected
for Lorentz and polarization effects as well as for absorption. The
structures of all three compounds were solved by direct methods
by using SIR97[28a] and refined on F2 by using SHELXL-97.[28b]
All non-hydrogen atoms were refined with anisotropic thermal pa-
rameters. The hydrogen atoms were included in ideal positions with
a fixed isotropic U value and were riding on their respective non-
hydrogen atoms. The data measurement and other refinement pa-
rameters for these three compounds are given in Table 3.
[5]
CCDC 906810 (for 1), -906811 (for 2), and -906812 (for 3) contain
the supplementary crystallographic data for this paper. These data
can be obtained free of charge from The Cambridge Crystallo-
graphic Data Centre via www.ccdc.cam.ac.uk/data_request/cif.
[6]
Supporting Information (see footnote on the first page of this arti-
1
cle): H and 13C NMR spectra of L, HRMS spectra of 1 and 2 in
various solvents, TG-DTA plots for 1 and 2, DSC curves for 1,
UV/Vis absorption spectra for 1–3 in various solvents, powder
XRD patterns, and powder EPR spectrum of 2.
Acknowledgments
G. H. is thankful to the Council of Scientific and Industrial Re-
search (CSIR), New Delhi, Scheme no. 01(2406)/10/EMR-II for
financial support. I. L. thanks the University Grant Commision
(UGC-SAP), New Delhi for a research fellowship. The Spanish au-
thors thank the Ministerio de Educación y Ciencia (MEC)
(CTQ2009-07264/BQU) and the Comissió Interdepartamental de
Recerca i Innovació Tecnològica de la Generalitat de Catalunya
(CIRIT) (2009SGR1454) for financial support.
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Received: April 25, 2013
Published Online: July 19, 2013
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© 2013 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim