Coordination chemistry of Cu(II), Co(II), Zn(II) and Ag(I) complexes of isomeric pyridine 2- and 4-carboxamides and their biological activity evaluation

Article abstract of DOI:10.1016/j.poly.2017.01.063

Eight complexes of Cu(II), Co(II), Zn(II) and Ag(I) with rarely studied ligands, N,N-diisopropyl/butylpicolinamide (L1/L2), and N,N-diisopropyl/butylisonicotinamide (L3/L4), have been synthesized and characterized spectroscopically, and their molecular and crystal structures have been reported. Diverse coordination modes of these positional isomers have been discovered, discussed and compared with the limited available literature in light of their respective convergent and divergent nature. All the complexes show stable and extended 1D, 2D or 3D coordination/H-bonded networks owing to a large number of weak C–H?O/X and other intermolecular interactions. Further, the ligands as well as the metal complexes have been evaluated for their antimicrobial activity and have turned out to be potent antimicrobial agents.

Full text of DOI:10.1016/j.poly.2017.01.063

Accepted Manuscript
Coordination chemistry of Cu(II), Co(II), Zn(II) and Ag(I) complexes of iso-
meric pyridine 2- and 4-carboxamides and their biological activity evaluation
Isha Lumb, Balkaran Singh Sran, Henna, Daljeet Singh Arora, Geeta Hundal
PII:
S0277-5387(17)30112-2
DOI:
http://dx.doi.org/10.1016/j.poly.2017.01.063
POLY 12469
Reference:
To appear in:
Polyhedron
Received Date:
Revised Date:
Accepted Date:
11 November 2016
29 January 2017
30 January 2017
Please cite this article as: I. Lumb, B.S. Sran, Henna, D.S. Arora, G. Hundal, Coordination chemistry of Cu(II),
Co(II), Zn(II) and Ag(I) complexes of isomeric pyridine 2- and 4-carboxamides and their biological activity
evaluation, Polyhedron (2017), doi: http://dx.doi.org/10.1016/j.poly.2017.01.063
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Coordination chemistry of Cu(II), Co(II), Zn(II) and Ag(I) complexes of
isomeric pyridine 2- and 4-carboxamides and their biological activity
evaluation.
a
Isha Lumb , Balkaran Singh Sran , Henna , Daljeet Singh Arora , Geeta Hundal *
a
b
b
a
a
Department of Chemistry, Centre for Advanced Studies in Chemistry, Guru Nanak Dev
University, Amritsar-143005, India
b
Department of Microbiology, Guru Nanak Dev University, Amritsar-143005, India
Email- geetahundal@yahoo.com
Keywords: convergent/divergent ligands, coordination/H-bonding complexes,
antimicrobial activity.
Abstract
Eight complexes of Cu(II), Co(II), Zn(II) and Ag(I) with rarely studied
ligands,
N,N-diisopropyl/butylpicolinamide
(L1/L2),
and
N,N-
diisopropyl/butylisonicotinamide (L3/L4), have been synthesized and characterized
spectroscopically, and their molecular and crystal structures have been reported.
Diverse
coordination
modes
of
these
positional
isomers
1

For table of content only
Coordination chemistry of Cu(II), Co(II), Zn(II) and Ag(I) complexes
isomeric pyridine 2- and 4-carboxamides and their biological activ
evaluation.
a
Isha Lumb , Balkaran Singh Sran , Heena , Daljeet Singh Arora , Geeta Hundal*
a
b
b
a,*
Department of Chemistry, Centre for Advanced Studies in Chemistry, Guru Nanak Dev University,
Amritsar-143005, India
b
Department of Microbiology, Guru Nanak Dev University, Amritsar-143005, India
Email- geetahundal@yahoo.com
Table of content
Eight complexes of Cu(II), Co(II), Zn(II) and Ag(I) ions with the rarely studied ligands, N,N-
diisopropyl/butylpicolinamide (L1/L2) and N,N-diisopropyl/butylisonicotinamide (L3/L4),
have been synthesized and characterized spectroscopically, and their molecular and crystal
structures have been reported. Diverse coordination modes of these positional isomers have
been discovered, discussed and compared with the limited available literature in the light of
their respective convergent and divergent nature.
2

have been discovered, discussed and compared with the limited available literature in
light of their respective convergent and divergent nature. All the complexes show
stable and extended 1D, 2D or 3D coordination/H-bonded networks owing to a large
number of weak C-H…O/X and other intermolecular interactions. Further, the ligands
as well as the metal complexes have been evaluated for their antimicrobial activity
and have turned out to be potent antimicrobial agents.
1
. Introduction
The distinguished role of the carboxamide group in synthetic coordination chemistry
is due to its biological significance [1]. The favourable properties of carboxamides,
such as high polarity, stability and conformational diversity, make them one of the
most popular and reliable functional groups in all branches of organic chemistry [2].
Coordination compounds of pyridine carboxamides are not only useful in
hydroxylation, epoxidation and asymmetric catalysis [3-7], NO delivery [8],
molecular recognition [9-11], dendrimer synthesis [12], spin ground state control [13-
16] and magnetic interactions [17-20], but they are amongst the most commonly
occurring heterocycles in anticancer drugs [21] and are involved in the regulation of
sirtuin activity [22]. The ligands provide an O, N donor set, which when coordinated
to transition metals may possess potentially useful biological activity, for e.g. Cu(II)
complexes with 2-picolinamide and its derivatives act as inhibitors of human
immunodeficiency virus type I protease [23] and a zinc(II) picolinamide complex
exhibits insulin mimetic activity [23]. Amide based complexes also demonstrate
antimicrobial activity against bacteria, yeast and fungi [24]. For example,
isonicotinamides possess strong antitubercular, anti-pyretic, fibrinolytic, anti-bacterial
and pharmacological properties, and also act simply as a divergent ligand for the
formation of MOFs [25]. Since the resistance to commercially available
antimicrobials is increasing at an alarming rate, there is a dire need of new
antimicrobial materials to deal with pathogenic microbes. Antimicrobial behaviour is
enhanced by incorporating transition metals in the pyridine carboxamide ligand [26].
The dimensionality and topology of coordination compounds/polymers
resulting from coordination and/or hydrogen bonding is dependent on many factors,
like the type of ligand, metal ion, anion, solvent, temperature etc. The use of divergent
polydentate ligands in metal coordination instead of the traditional convergent
chelating ligands can be very helpful in extending the coordination network through
3

space in 1, 2 or 3D architectures [21]. Additionally, the potential H-bonding of such
ligands proves to be an innovation for crystal engineering of H-bonded networks. We
have been interested in coordination complexes of N,N-disubstituted complexes of
pyridine mono/dicarboxamides for quite some time now with the aim of generating
coordination/H-bonded networks. It has been noted that among the pyridine
monocarboxamides, barring a few complexes of unsubstituted 2-picolinamide [27-32]
and a single Pd(II) complex of L1 [33], these ligands and their complexes were
largely unreported (based upon CSD and scifinder) until the reports of L3 and L4 and
some of their complexes from our lab [34-38]. To date there is no report for L2 with
any metal ion (Scheme 1). However, such fully (alkyl/phenyl) substituted pyridine
carboxamides also provide the necessary rigidity, binding sites and intermolecular
interactions to allow the creation of coordination/H-bonded polymers. A variety of
charge neutralizing anions, the electron deficient pyridine ring, the amide O atom and
abundant peripheral alkyl groups, offer just the right combination for generating a
large number of C-H···X/O/π, π···π and lone pair···π interactions [39]. The
accumulative “Gulliver effect” [40] of these weak interactions is envisaged to support,
direct or even dictate the dimensionality and topology of these compounds, leading to
diverse supramolecular architectures. The work here significantly entails the
complexes of L1-L4 with transition metal ions and the Zn(II) ion, which have been
synthesized and fully characterized by spectroscopy and X-ray diffraction methods,
and their antimicrobial activity is also reported.
R = diisopropyl, L1 and L3; R = diisopropyl, L2 and L4.
Scheme 1- Showing the ligands and their modes of coordination
2
.
Experimental details
Material and methods
2.1
4

All reactions were carried out in anhydrous solvents. The solvents were
distilled prior to use (THF with Na wire) following the appropriate procedures [41].
Anhydrous cobalt(II) chloride was prepared by treating a powdered sample of
hydrated CoCl
2
with SOCl
2
and refluxing the contents for a few hours until the
evolution of SO
2
ceased. The anhydrous product was washed with petroleum ether
and dried in a vacuum. Elemental analyses (C, H, N) were performed on a Perkin-
Elmer model 2400 CHN analyzer. Thermal analyses were carried out on a Perkin
Elmer Diamond TG/DTA and SII 6300 EXSTAR TG/DTA thermo analyzer in flow
−1
2
of N , in the temperature range from 20 to 800 °C, with a heating rate of 10 °C min .
IR spectra were recorded as KBr pellets on a Varian 660-IR FT-IR
1
spectrophotometer. The H NMR spectra of the ligands and complexes were recorded
on a 300 MHz JEOL FT NMR spectrometer with TMS as the reference compound.
UV-vis spectra were recorded on
a
Shimadzu Pharmaspec UV-1700
spectrophotometer. The powder XRD data were recorded on a Rigaku miniflex 2 in
the range θ = 2 to 60 °C. Cyclic voltammetric measurements were carried out on a
BAS CV 50W electrochemical analyzing system (accuracy 1.0 mV). The cyclic
voltammograms were recorded in acetonitrile with 0.1 M TBAP as the supporting
2
electrolyte using a three electrode system under a N atmosphere. A three electrode
configuration, composed of a Pt working electrode, a Pt wire counter electrode and a
Ag/AgCl reference electrode, was used for the measurements, standardized for the
redox couple ferricinium/ferrocene (E1/2 = 0.440 V, ∆E
P
= 0.052 V vs. Pt/PtCl
2
electrode). High resolution mass spectra were recorded on a Bruker’s microTOF-QII
spectrophotometer. The magnetic studies of the sample were performed using a
vibrating sample magnetometer (Microsense EZ9). The powdered sample was first
pressed into a pellet and then magnetic studies were carried out at room temperature.
Caution! Perchlorate salts are explosive in nature. The complexes synthesised here
could explode upon heating and should be handled with caution and in small amounts.
2.2
X-ray data collection
The data for all the complexes were collected on a Bruker APEX-II CCD
diffractometer using Mo Kα radiation (λ = 0.71069 Å). Complexes 1-8 were
measured at room temperature and complex 6 was measured at 153 K. The data
collected were processed by SAINT. The data were corrected for Lorentz and
polarization effects as well as for absorption. The structures of all the compounds
5

2
were solved by direct methods using SIR 97 [42] and refined on F using SHELXL 97
43]. The hydrogen atoms of the water molecules were located from the difference
[
Fourier synthesis and were refined isotropically with a distance of 0.82 Å, with Uiso
values 1.2 times that of their carrier oxygen atoms. All non-hydrogen atoms were
refined using anisotropic thermal parameters. The hydrogen atoms were included in
ideal positions with fixed isotropic Uiso values of 1.2 times that of their phenylene,
methine and methylene carbon atoms and 1.5 times that of their terminal methyl
carbon atoms,s and were treated as riding on their respective carbon atoms. There is
disorder in oxygen atoms O4-O6 of the perchlorate anion and the C9 atom of the
terminal methyl group in complex 4, while the C8 atom of the -CH group and the C14
atom of the terminal -CH group in 5 are also disordered. These disorders were
3
resolved by splitting each of these atoms in two parts (using PART) with their site
occupancy factors and Uiso values being refined as free variables and with their
distances being fixed using DFIX. In complex 6, two water molecules, O3W and
O4W, had high thermal parameters and showed a significantly short (~1.62 Å) Cu−O
distance. These were refined with their site occupancies reduced to 0.5 to give more
meaningful Ueq values, which also improved the R value. The distances still
remained anomalous, therefore these were fixed at 1.815(5) Å and refined
3
anisotropically. A residual electron density of about 1.0 e Å remained near these
atoms as a consequence and the hydrogen atoms of this disordered water molecule
were not found. In complex 8, the protons of the O2W and O4W of coordinated water
molecules could not be located from the difference Fourier map. All calculations were
performed using the WINGX package [44]. A summary of the crystal data,
experimental and refinement details are given in Table 1.
2.3 Antimicrobial activity
(
i) Test organisms
The reference strains of bacteria and one yeast, used for testing their sensitivity to
compounds, were obtained from the Microbial Type Culture Collection (MTCC),
Institute of Microbial Technology (IMTECH), Chandigarh, India. Reference strains
included the Gram positive bacteria Staphylococcus aureus (MTCC 740), the Gram
negative bacteria Klebsiella pneumoniae (MTCC 109), Pseudomonas aeruginosa
(
MTCC 741) and Escherichia coli (MTCC 119), and one yeast strain, viz Candida
6

albicans (MTCC 227). The bacterial cultures were maintained on nutrient agar slants
and Candida albicans was maintained on yeast malt agar.
(
ii) Inoculum preparation
A loopful of isolated bacterial and yeast colonies were inoculated into 5 mL of the
respective broth medium and incubated at 37 and 25 ºC respectively for 4 hrs. The
turbidity of the actively growing bacterial suspension was adjusted to match the
turbidity standard of 0.5 Mc Farland units, prepared by mixing 0.5 mL of 1.75% (w/v)
barium chloride dihydrate to 99.5 mL of 0.18 M (v/v) sulfuric acid with constant
stirring. The microbial suspension thus prepared was used for testing their sensitivity
to all the compounds used.
(
iii) Screening of compounds for antimicrobial activity by agar well diffusion
assays.
The sensitivity of different bacterial strains to different compounds was measured in
terms of the zone of inhibition using agar well diffusion assays [45]. The plates
containing Muller Hinton agar medium and yeast malt agar were spread with 100µL
of the bacterial and yeast inoculums, respectively. Wells (6 mm diameter) were cut
out from agar plates using a sterilized stainless steel cork borer and filled with 50µL
of the compound. The plates were incubated at 37 and 25 ºC for 24 hrs and the
diameter of the resultant zone of inhibition was measured. Experiments were run in
duplicate for each combination of extract and microbial strains. DMSO was used as a
control and the results were compared with that of gentamycin and amphotericin as
standard drugs.
(
iv) Minimum inhibitory concentration
The minimum inhibitory concentration was worked out by the agar dilution method.
A stock solution of all the compounds pf 2.5% (25 mg/mL) concentration were
prepared and incorporated into Muller Hinton agar medium for bacteria, and yeast
malt extract medium for Candida albicans. The final concentrations of the
compounds in the medium containing plates ranged from 0.005-2.5 mg/mL. These
plates were then inoculated with 50 µL of the activated bacterial and yeast strains by
streaking with a sterile tooth pick. The plates were incubated at 37 ºC for bacteria and
25 ºC for yeast for 24 hrs and the lowest concentration of the extract causing complete
inhibition of the bacterial growth was taken as the MIC. The results were compared
with that of a standard drug and DMSO was used as a control. The experiment was
performed in duplicate and repeated twice.
7

2.4
Experimental section
L1 was prepared by a mixed anhydride coupling reaction [46] and is reported in our
previous paper [34], L2 has been prepared in exactly the same way and it is reported
here for the first time. L3 and L4 were reported earlier by our lab [35,37].
2.4.1 N,N-diisobutylpyridine-2-carboxamide, L2: Yield: 70%. B.p.: 70-72 ºC. Anal.
Calc. for C14
H
22
N
2
O: C, 71.76; H, 9.46; N, 11.95; Found: C, 71.12; H, 9.20; N, 12.04
−
1
1
%
. IR (KBr/cm ): 1631 (CO). H NMR [300 MHz; CDCl
3
, Me
), 1.70-1.83 (m, 1H, -CH),
= 7.50 Hz, J = 1.5 Hz, -CH ), 3.30,
2
), 7.26-7.30 (m, 1H, PyH), 7.52 (d, 1H,
4
Si] δ (ppm): 0.72 (d,
6
2
3
H, J = 6.60 Hz, -CH
.00-2.22 (m, 1H, -CH), 3.38, 3.39 (dd, 1H, J
.32 (dd, 1H, J = 7.50 Hz, J = 1.5 Hz, -CH
3
), 1.00 (6H, d, J = 6.60 Hz, -CH
3
1
2
2
1
2
1
3
PyH, J = 7.80 Hz), 7.75 (t, 1H, PyH, J = 7.50 Hz), 8.57 (d, 1H, PyH, J = 4.80 Hz). C
NMR [75 MHz; CDCl ] δ (ppm): 19.8 (CH ), 20.2 (CH ), 26.5 (CH), 27.3 (CH), 52.5
CH ), 56.0 (CH ), 123.6 (Py), 123.8, 136.8, 148.1, 155.6, 169.7 (C=O).
3
3
3
(
2
2
2
.4.2 Synthesis of [Co(L1)
To a suspension of ligand L1 (0.103 g, 0.5 mmol) in 5 mL isopropyl alcohol,
anhydrous CoCl (specially dried with SOCl ) (0.064 g, 0.5 mmol) dissolved in 5 mL
2
Cl
2
], 1:
2
2
isopropyl alcohol was added at once with constant stirring. Thereafter the blue
colored solution so obtained was transferred to a 50 mL conical and allowed to stand
at room temperature. Brown colored crystals were obtained after long time and these
were then filtered, air dried and analyzed for their composition. Yield: 0.17 g,
(
63.05%). M. p.: 150-153 ºC. Anal. Found: C, 55.27; H, 6.31; N, 7.54, Calc. for
−1
C
25 2 3 2
H37Cl CoN O : C, 55.46; H, 6.89; N, 7.76%. IR (υmax/cm ): 1588, 1372. µeff
(
B.M.): 3.84.
2
.4.3 Synthesis of [Zn(L1)
This complex was prepared from a reaction mixture containing
.6H O (0.062 g, 0.16 mmol) and L1 (0.068 g, 0.33 mmol) dissolved in a
2 4 2
(ClO ) ], 2:
Zn(ClO
4
)
2
2
minimum amount of methanol. The contents were stirred for 5-6 hrs and then kept at
room temperature for 2-3 days to yield colorless crystals. Yield: 0.17 g, (75.8%).
M.p.: 176-178 ºC. Anal. Found: C, 44.01; H, 5.13; N, 6.43. Calc. for
−
1
1
C
25
H
37Cl
2
ZnN
NMR (500 MHz; CDCl
.52 (s, 2H, -CH), 3.81 (s, 2H, -CH), 7.30-7.38 (m, 4H, PyH), 7.77 (d, 2H, PyH, J =
3
O
10: C, 44.43; H, 5.52; N, 6.22%. IR (υmax/cm ): 1590, 1371. H
3
; Me Si) δ (ppm): 1.13 (s, 12H, -CH ), 1.46 (s, 12H, -CH ),
4
3
3
3
8

1
3
8
.00 Hz), 8.53 (d, 2H, PyH, J = 8.00 Hz). C NMR (125 MHz, CDCl
CH ), 20.6 (CH ), 46.1 (CH), 50.9 (CH), 121.9, 124.3, 137.4, 148.8, 167.3 (C=O).
.4.4 Synthesis of [Ag(L1) NO ], 3:
3
) δ (ppm): 20.4
(
3
3
2
2
3
To a methanolic solution of L1, (0.082 g, 0.4 mmol) in 5 mL methanol,
AgNO
3
(0.039 g, 0.23 mmol) predissolved in 10.0 mL methanol was added dropwise
at room temperature. The colorless solution was filtered and kept in a conical flask for
slow evaporation. The colorless needles obtained were air dried. The compound was
analyzed to obtain its composition. Yield: 0.19 g, (72.13%). M.p.: 110-113 ºC. Anal.
Found: C, 49.39; H, 6.01; N, 12.30. Calc. for C24
H
36AgN
5
O
5
: C, 49.49; H, 6.23; N,
2.02%. IR (υmax/cm ): 1627, 1448. H NMR (500 MHz; CDCl ; Me Si) δ (ppm):
.11 (s, 12H, -CH ), 1.46 (s, 12H, -CH ), 3.51 (s, 2H, -CH), 3.67 (s, 2H, -CH), 7.32
−1
1
1
1
3
4
3
3
1
3
(
d, 4H, PyH, J = 10.00 Hz), 7.76 (s, 2H, PyH), 8.51 (s, 2H, PyH). C NMR (125
MHz, CDCl ) δ (ppm): 20.6 (CH ), 20.9 (CH ), 46.4 (CH), 50.7 (CH) 122.9, 125.3,
35.4, 148.6, 166.2 (C=O).
.4.6 Synthesis of [Ag (L2)
To a suspension of the ligand L2 (0.093 g, 0.39 mmol) in 5 mL methanol,
(0.041 g, 0.19 mmol) dissolved in 5 mL of dry methanol was added dropwise
3
3
3
1
2
2
2
2 4 2
(µ−L2) ].2ClO .2H O, 4:
AgClO
4
with constant stirring. The resulting mixture was stirred at room temperature for an
hour. The colorless solution thus obtained was transferred to a 50 mL conical flask
and allowed to stand at room temperature overnight. The colorless crystals obtained
were collected, air dried and analyzed to obtain their composition.
Yield: 0.45 g, (72.52%). M.p.: 165-167 ºC. Anal. Found: C, 48.07; H, 6.18; N, 8.41.
−
1
Calc. for C56
H
92Cl
2
Ag
2
N
8
O
14: C, 48.46; H, 6.68; N, 8.07%. IR (υmax/cm ): 1634,
; Me Si) δ (ppm): 0.81 [d, 24H, J = 4.0 Hz, -
], 0.99 [d, 24H, J = 4.0 Hz, -CH(CH ], 1.88-1.91 (m, 4H, -CH), 2.13-2.16
m, 4H, -CH), 3.25 (d, 8H, J = 8.0 Hz, -CH ), 3.42 (d, 8H, J = 8.0 Hz, -CH ), 7.55 (d,
H, PyH, J = 8.0 Hz), 7.64 (t, 4H, PyH, J = 4.0 Hz), 7.98 (t, 4H, PyH, J = 8.00 Hz),
1
1
421. H NMR (500 MHz; CDCl
3
4
CH(CH
3
)
2
3 2
)
(
2
2
4
9
2
1
2
1
3
.00 (d, 4H, PyH, J = 4.0 Hz). C NMR (125 MHz, CDCl
3
) δ (ppm): 19.5 (CH
3 2
) ,
0.3 (CH , 26.4 (CH), 27.2(CH), 52.5 (CH ), 57.2 (CH
3
)
2
2
2
), 123.5, 126.2, 138.5,
52.2, 153.2, 170.2 (C=O).
2 4 2
.4.7 Synthesis of [Cu(L2) (ClO ) ], 5:
To a methanolic solution of L2 (0.093 g, 0.4 mmol) in 5 mL methanol,
Cu(ClO .6H O (0.074 g, 0.2 mmol) pre-dissolved in 5.0 mL methanol was added
4
)
2
2
dropwise at room temperature. The blue coloured solution obtained was filtered, air
dried and left for crystallisation. The crystals thus obtained were analysed for their
9

composition. Yield: 0.24 g, (74.24%). M.p.: 154-156 ºC. Anal. Found: C, 46.08; H,
.13; N, 7.43. Calc. for C28 CuN 10: C, 46.00; H, 6.07; N, 7.66%. IR
max/cm ): 1584, 1454. µeff (B.M.): 1.97.
.4.8 Synthesis of [Cu(L3) (NO (H O)1.5], 6:
·3H O (0.12 g, 0.5 mmol) and L3 (0.103 g, 0.5 mmol) were dissolved in
0.0 mL of methanol, with continuous stirring. The blue coloured solution thus
6
H44Cl
2
4
O
−1
(
υ
2
2
)
3 2
2
Cu(NO
3
)
2
2
2
obtained was filtered and kept for slow evaporation. Blue coloured crystals appeared
in the filtrate after a few days.
Yield: 0.48 g, (77.79%). M.p.: 257-259 ºC. Anal. Found: C, 46.24; H, 6.10; N, 13.21.
-
1
Calc. for C24
H
38CuN
eff (B.M.): 1.98.
.4.9 Synthesis of [Co(L4)
6
O : C, 46.63; H, 6.20; N, 13.60%. IR (υmax/cm ): 1632, 1301.
9
µ
2
2 2 2 n 3
(H O) ] .2NO , 7:
3 2 2
This complex was prepared from a reaction mixture containing Co(NO ) .6H O (0.097
g, 0.33 mmol) and L4 (0.078 g, 0.33 mmol) dissolved in a minimum amount of
isopropyl alcohol. The contents were stirred for 5-6 hrs and then kept at room
temperature for 5-6 days. This yielded pink crystals, which were analysed for their
composition after air drying. Yield: 0.18 g, (69.6%). M. p.: 200-202 ºC. Anal. Found:
6
C, 48.21; H, 7.13; N, 12.51. Calc. for C28H48CoN O10: C, 48.91; H, 7.04; N, 12.22%.
-
1
IR (υmax/cm ): 1608, 1291.µeff (B.M.): 4.96
2
.4.10 Synthesis of [Zn
This complex was prepared from a reaction mixture containing equimolar amounts of
Zn(NO .3H O (0.072 g, 0.33 mmol) and the ligand L4 (0.078 g, 0.33 mmol)
2
(L4)
2
(µ−L4)
2 3 4
(NO ) ], 8:
3
)
2
2
dissolved in a minimum amount of methanol. The contents were stirred for 2-3 hrs
and then on keeping the solution at room temperature for 2-3 days, colourless crystals
were obtained, which were filtered and dried in vacuo. The compound was analysed
to obtain its composition. Yield: 0.36 g, (72.08%). M.p.: 122-125 °C. Anal. Found: C,
5
1.26; H, 6.25; N, 13.01. Calc. for C56
H
88Zn
2
N
12
O
16: C, 51.10; H, 6.74; N, 12.77%.
IR (υmax/cm ): 1617, 1308. H NMR (500 MHz; CDCl ; Me Si) δ (ppm): 0.80 [d,
], 1.01 [d,24H, J = 4.0 Hz, CH(CH ], 1.89-1.92 (m,
H,-CH), 2.13-2.15 (m, 4H, -CH), 3.02 (d, 8H, J = 8.0 Hz, -CH ), 3.39 (d,8H, J =
), 7.46 (s, 8H, PyH), 8.79 (s, 8H, PyH ), C NMR (125 MHz, CDCl ) δ
, 20.1 (CH , 26.1 (CH), 26.8 (CH), 51.0 (NCH ), 56.2 (NCH ),
44.9, 149.5, 168.4 (C=O).
−1
1
3
4
2
4
8
4H, J = 4.0 Hz, -CH(CH
3
)
2
3 2
)
2
1
3
.0 Hz, -CH
2
3
(
ppm): 19.7 (CH
3
)
2
3
)
2
2
2
1
1
0

2.4
Results and discussion
2.4.1 NMR spectral studies
1
The H NMR spectra of the ligands and the complexes were assigned via peak
1
integration, as displayed in Table S1a and shown in Fig. S1a. The H NMR spectrum
of complex 4 shows slight shifts in the aliphatic region as compared to ligand L2, as
shown in Table S1b. However, the complexation causes significant higher frequency
shifts for all the aromatic protons. Proton a shows the maximum shift (∆δ 0.43) and c
shows the minimum shift (∆δ 0.12). Similarly, complex 8 shows significant higher
frequency shifts for the aromatic (∆δ 0.29, 0.20) and aliphatic region, as compared to
the ligand L4. The significant shifts in the aromatic region indicate complexation
through the ligands. Complexes 2 and 3 (Fig. S1b) show slight shifts in the aromatic
and aliphatic regiosn as compared to the ligand L1. In the case of complexes 2 and 3
the aromatic pyridine ring protons (b and d) show slightly higher frequency shifts
(
max. ∆δ = -0.05, Table S1a) while all other protons show lower frequency shifts. As
a result of this, the b and c protons merge together to give a multiplet instead of a dd
signal. Maximum shifts (∆δ) are seen for the signals of the methine and methyl
protons, which may be attributed to the steric hindrance of these groups in the
complexes, leading to a lesser involvement in C-H…X type H-bonding relative to the
free ligand. It is clear that significant shifts are seen in 4 and 8, and incidentally both
these complexes are dimeric complexes. Detailed X-ray crystal structures of these
complexes (vide infra) shows that these dimers are π…π bonded centrosymmetric
dimers, while 2 and 3 are 0D complexes. Therefore, the anisotropy of the pyridine
rings might be a reason for the significant chemical shifts in the dimers, apart from the
coordination.
2.4.4 Infra-red spectral studies
The pyridine mono carboxamide ligands possess two potential donor sites in the
form of the pyridine nitrogen atom and the carbonyl oxygen atom (Fig. S2). The bonding
may involve either monodentate or bidentate coordination by the ligands. The IR
spectroscopy corroborates the data by showing the νC=O band of the amide ligand
−1
assigned between 1600 and1640 cm . The νC−N and νC−H stretching bands are assigned to
1
1

−1
1
450-1500 and 2900-3000 cm , respectively. The remaining peaks are given in Table
S2. All the complexes from 1-8 show a decrease in the νC=O stretching wavelength,
indicating coordination through the amide carbonyl oxygen atom (Fig. S2). Binding
through pyridine nitrogen atom increases the νC−N stretching frequency in 1-5, while it
−1
decreases in 6-8. Broad νOH bands in the region 3350-3450 cm in the complexes
indicate the presence of lattice/coordinated water molecules. Characteristic stretching
frequencies for the coordinated anions are also observed and are listed in Table S2. The
−1
band in the range 730-550 cm corresponds to the Cu–Cl bond. A strong band at 1384-
−1
1
390 cm in the nitrate containing complexes appears due to the -NO
stretch in the ONO group and is characteristics of ionic nitrates [47,48]. Such a band has
been found in complex 7 at 1385 cm . Addison et al. [49] reviewed that the NO
2
asymmetric
2
-1
¯
3
ion
coordinates to metal centres as a unidentate, symmetric and asymmetric chelating
bidentate and bridging bidentate ligand. As the symmetry of the nitrate ion differs very
S
little among them (C2ν or C ), it is difficult to distinguish these structures through
vibrational spectroscopy. In general Curtis and Curtis [50] concluded that in a similar
type of complex, the separation ꢀν of the two highest frequency bands is larger for
bidentate coordination than for unidentate coordination. Gatehouse et al. [48] observed
that both unidentate and bidentate NO
bands [ν (NO ), ν (NO ) and ν(NO)] called ν5 ,
bands differ in the two, ν > ν > ν in the former and ν
shows bands at 1483, 1347 and 1038 cm , ꢀν = 136 cm . The complex 6 shows bands at
3
groups with C2ν symmetry show three stretching
and ν2 , respectively . The order of these
> ν > ν in the latter. Complex 3
a
2
s
2
ν
1
5
1
2
1
5
2
-1
-1
-1
-1
1
423, 1349 and 1032 cm , ꢀν = 74 cm . Though the bands in both 3 and 6 are close to
each other, ꢀν shows that the mode of coordination is bidentate and unidentate,
-1
respectively. Complex 8 shows bands at 1388, 1293 and 1027 cm , obtained due to
mixing of the abovementioned bands [50,51], indicating both unidentate and bidentate
coordination. The presence of perchlorate groups in the complexes is indicated by a
−1
strong and broad band in the range 1070-1100 cm and a medium intensity band
−1
[
34,52-53] between 610 and 630 cm .
.4.5 X-ray crystal structures description
i) Molecular and crystal structure of [Co(L1)
2
(
2
(Cl) ], 1
2
1
2

Complex 1 (Fig. 1) is single molecular complex which crystallizes in the
monoclinic centrosymmetric space group P2 /n. In this complex, the Co(II) ion lies on
1
the centre of symmetry and thus the asymmetric unit contains only one half of the
molecule. The Co(II) ion is in a tetragonally distorted octahedral environment, where
two L1 ligands are coordinating in the typical chelating mode, with the remaining two
sites being occupied by chloride anions (Table S3). The dihedral angle between the
pyridine ring of L1 and its amide group is 35.87(9)°, which shows that these two
groups are rotated with respect to each other. In the crystal structure of complex 1,
terminal methyl groups, methine and phenylene carbon atoms as well as the
coordinated chloride anion show intermolecular H-bonding interactions between the
i
single molecules. C4−H4···Cl(1) H-bonds form 1D chains along the a axis (Fig. S3a)
and H-bonding between the methane C7 and methyl C12 atoms with a Cl atom forms
a 2D network in the ab plane (Fig. S3b). Important hydrogen bond parameters are
listed in Table S4.
Fig. 1. ORTEP diagram of complex 1 with 40% thermal probability ellipsoids.
(
ii)
Molecular and crystal structure of [Zn(L1)
2 4 2
.(ClO ) ], 2
In complex 2 (Fig. 2), the Zn(II) ion lies on the centre of symmetry and thus
the asymmetric unit contains only one molecule of the L1 ligand and one perchlorate
anion. The Zn(II) ion is in a tetragonally distorted octahedral geometry, with two
chelating L1 ligands (Table S5) and two perchlorate ions present in trans positions
relative to each other. The dihedral angle between the pyridine ring of L1 and its
amide group is 30.4(2)°. The H-bonding interactions in this complex are due to the
coordinated perchlorate anion, where the C4−H4···O4 interaction gives rise to 1D
chains along the a axis (Fig. S4a) and other weak H-bonding interactions of the
1
3

pyridine and terminal methyl group with the perchlorate anion (Table S6) give rise to
a 2D network in the ac plane (Fig. S4b). This is further strengthened by the CH···π
interaction C11−H11B···π (C11···centroid 3.462 Å, H11B···centroid 2.545 Å).
Fig. 2. ORTEP diagram of complex 2 with 30% thermal probability ellipsoids.
(
2 3
iii) Molecular and crystal structure of [Ag(L1) NO ], 3
The single molecular complex 3 (Fig. 3) has one Ag(I) ion, two L1 ligands and one
−
NO
3
ion present in the asymmetric unit. The Ag(I) ion has a square pyramidal geometry
(τ ≈ 0.2), connected to the N1 atom of the pyridine ring and O1 atom of one L1 ligand and
the N2 atom of the pyridine ring of another L1 ligand. The nitrate ion is coordinating in a
chelating mode (Table S7). The Ag(1)···O1(amide) bond 2.666(2) Å is significantly longer
than the Ag−O(nitrate) bonds. The dihedral angles between the pyridine rings of two L1
ligands are 51.9(2) and 81.8(2)°. The dihedral angle between the pyridine rings of two L1
−
ligands and the coordinated NO
3
ion are 83.4(2) and 28.2(2)°, which shows that one
pyridine ring is almost perpendicular to the nitrate anion and other is parallel to it. The
intermolecular H-bonding interactions between the coordinated nitrate oxygen O3 atom
i
and the pyridine ring hydrogen H2 atom, C2−H2···O3 , and C−H···π (Py) interactions,
C13−H13···π (C13···centroid 3.826 Å, H4···centroid 3.012 Å) and C23−H23C···π
(C23···centroid 3.996 Å, H23C···centroid 3.240 Å) help in forming a linear 1D chain
along the c axis (Fig. S5a). The other intermolecular H-bonding interaction act via the O4
atom with the hydrogen H22 atom of a terminal methyl group (C22−H22···O4 = 3.496(5)
Å (#1 = −x+1,−y,−z+2) (Fig. S5b, Table S8), further helping to extend it into a 2D
hydrogen bonded network in the ac plane.
1
4

Fig. 3. ORTEP diagram of complex 3 with 30% thermal probability ellipsoids.
Hydrogen atoms have been omitted for clarity.
(
iv)
Molecular and crystal structure of [Ag
2 2 2 4 2
(L2) (µ−L2) ].2ClO .2H O, 4
The centrosymmetric, edge sharing dimeric complex 4 (Fig. 4) has one
crystallographically independent Ag(I) ion, two L2 ligands, one perchlorate and
one water molecule present in the asymmetric unit. The Ag(I) ion has a trigonal
coordination geometry (Fig. 4 inset) through the N1 and N3 atoms of pyridine
rings from two ligands, while the third position is occupied by the oxygen atom
O1 of the amide group of a third centrosymmetric ligand (Table S9). Thus, one L2
ligand is behaving as a bridging bidentate ligand, while second is monodentate
through the pyridine nitrogen atom. The amide group of second ligand is non-
coordinating, but it does show H-bonding interactions with the lattice water
molecules. The dihedral angle between the pyridine rings and amide groups of the
L2 ligands are 58.87(13) and 50.65(38)°, which show that the pyridine rings are
gauche to the amide groups. Ag···Ag non-bonding distances generally lie in the
range 2.78-3.44 Å, as found in the literature [54-56]. The sum of the van der Waals
radii of two silver atoms is 3.44 Å and the Ag···Ag distance in the dimeric unit 4
is found to be 3.393 Å, which may be considered as argentophilic interactions.
There are strong π···π interactions between two unique pyridine rings with a
1
5

centroid···centroid distance of 3.729 Å (Fig. S6a). Consequently, CH···π (Py)
interactions, C3−H3···π (C3···centroid 3.578 Å, H3···centroid 3.750 Å) and
C4−H4···π (C4···centroid 3.559 Å, H4···centroid 3.762 Å), are observed (Fig.
S6a). The perchlorate anion H-bonds with the methyl groups of the ligand to give
rise to a 1D chain along the b axis and a 2D network in the ab plane (Fig. S6b).
Important hydrogen bond parameters are listed in Table S10.
Fig. 4. ORTEP diagram of complex 4 with 30% thermal probability ellipsoids.
Hydrogen atoms and anions (perchlorate) are omitted for clarity.
(
v)
Molecular and crystal structure of [Cu(L2)
2
.(ClO
4
)
2
], 5
The single molecular complex 5 (Fig. 5) has one crystallographically
independent Cu(II) ion, one L2 ligand and one coordinated perchlorate present in the
asymmetric unit. The Cu(II) ion is tetragonally coordinated through the pyridine and
amide group (Table S11) of the L2 ligand, with the O2 atom of a perchlorate anion
with significantly long bond distances on the z-axis. The dihedral angle between the
pyridine ring and amide group is 18.7(1)°, which shows that both are almost parallel
to each other. The H-bonding interactions between the perchlorate ion and the
methylene and phenylene protons form a 2D network in the bc plane (Fig. S7, Table
S12).
1
6

Fig. 5. ORTEP diagram of complex 5 with 40% thermal probability ellipsoids.
Hydrogen atoms are omitted for clarity.
(
vi)
Molecular and crystal structure of [Cu(L3)
The complex 6 (Fig. 6) is a single molecular complex having two
crystallographically independent [Cu(L3) .(NO .(H O)1.5] units in the unit cell.
2 3 2 2
(NO ) (H O)1.5], 6
2
3
)
2
2
Since both units have the same structure, only one of them is discussed here. In each
molecular unit, the Cu(II) ion adopts an octahedral geometry in which two L3
ligands are coordinated through the pyridine nitrogen atoms N1 and N3 in trans
positions. The remaining four positions are occupied by two unidentate nitrate anions
(
Table S13) and one water molecule O1W(or O2W) and half of a disordered water
molecule O3W (or O4W) (Cu(1)···O1W = 2.234(5) Å and Cu(1)···O3W = 1.764(5)
Å) (Table S14). The dihedral angles between the pyridine rings and amide groups in
both molecules are 86.64(3), 89.61(3) and 89.75(4), 84.41(5)° which shows both are
almost perpendicular to each other. The coordinated nitrate group shows H-bonding
interactions with the phenylene protons (C37−H37···O4, C16−H16···O12,
C28−H28···O6, C4−H4···O14), while the coordinated water molecules are H-
bonded to each other (O1W−H11W···O4W, O2W−H22W···O3W), forming H-
bonded 1D zigzag polymeric chains along the a axis (Fig. S8). The water O2W
molecule shows H-bonding interactions with the carbonyl oxygen atom
1
7

(
O2W−H21W···O2), the nitrate group (C14−H14···O8, C26−H26···O11) with the
phenylene protons and the carbonyl oxygen atom (C25−H25···O2) with the
phenylene protons, further extending the complex into a 2D H-bonded chain.
Fig. 6. ORTEP diagram of the complex 6 with 40% thermal probability ellipsoids.
Hydrogen atoms are omitted for clarity.
(
2 2 2 n 3
vii) Molecular and crystal structure of [Co(L4) (H O) ] .2NO , 7
The 1D polymeric complex 7 has an asymmetric unit consisting of one Co(II)
−
ion, one L4 ligand, one coordinated water molecule and one lattice NO
3
ion. The
Co(II) ion has an octahedral environment (Fig. 7), connected to two pyridine nitrogen
and two amide oxygen atoms from four different L4 ligands. The remaining two
positions are covered by two O1W water molecules (Table S15). Thus each L4 ligand
is bridging between two Co(II) ions and the arrangement gives a centrosymmetric, 1D
polymeric tape (Fig. 8) The dihedral angle between the pyridine ring and amide group
of the L4 ligand is 74.39(23)°. In the crystal structure of complex 7, the terminal
methyl groups and nitrate groups show intermolecular H-bonding interactions and
π···π interactions between two pyridine rings (Fig. S9b), with a centroid···centroid
distance of 3.468 Å. The CH···π (Py) interactions, C1−H1···π (C1···centroid 3.240
Å, H1···centroid 3.405 Å) and C2−H2···π (C1···centroid 3.310 Å, H2···centroid
3.516 Å) (Fig. S9b), further strengthen the 1D tape. The nitrate groups show H-
bonding interactions with coordinated water protons and terminal methylene protons
of the ligand (Fig. S9 c,d and Table S16) and form corrugated chains in the ac plane.
Topologically, the structure consists of [0 1 0] chains with a 2-connected uninodal net
with the 2C1 topological type (Fig. S10).
1
8

Fig. 7. ORTEP diagram of the asymmetric unit with the atom numbering scheme used
for complex 7.
Fig. 8. The π···π interactions between the antiparallel pyridine rings and the
polyhedra around the metal centres in 7.
(
xi)
2 3 2 2
Molecular and crystal structure of [Zn(L4) .(NO ) ] , 8
The dimeric complex 8 (Fig. 9) has one crystallographically independent
−
Zn(II) ion, two L4 ligands and two coordinated NO
3
ions present in the asymmetric
unit. Each Zn(II) ion shows a distorted octahedral geometry (Fig. S11a), connected to
two pyridine nitrogen atoms from two L4 ligands and one position occupied by the
amide O2 atom from one of the centrosymmetrically related L4 ligands. The
remaining three positions are occupied by oxygen atoms of nitrate anions; one nitrate
group acts as a bidentate chelating ligand and the other one is unidentate (Table S17).
Thus, one L4 ligand behaves as a 2-connected linker, whereas the other one is a
1
9

monodentate ligand. The dihedral angle between the pyridine rings and amide groups
of the L4 ligands are 54.1(2) and 85.5(3)°, which shows that for the former, the
pyridine ring is gauche to the amide group, while in the latter case both are
perpendicular to each other. There are strong π···π and CH···π interactions between
the pyridine rings and alkyl groups (Fig. S11b). C-H...O interactions with the nitrate
oxygen atoms O6 and O8 and amide oxygen atom O1 form a 1D tape along the a axis
(
Fig. S11c). A unidentate nitrate ion further extends these parallel tapes into a 2D H-
bonded network in the ac plane (Fig. S11c). The hydrogen bond parameters are listed
in Table S18.
Fig. 9. ORTEP diagram of complex 8 with 30% thermal ellipsoids probability.
Hydrogen atoms are omitted for clarity.
2
0

Table 1. Crystallographic data for complexes 1-8.
Identification code
Empirical formula
Formula weight
T (K)
1
C24H36Cl2CoN4O2
542.40
2
C24H36Cl2N4O10Zn
676.84
3
C24H36AgN5O5
582.45
4
C56H92Ag2Cl2N8O14
1388.02
5
C28H44Cl2CuN4O10
731.12
6
C48H76Cu2N12O19
1252.29
153(2)
7
C28H48CoN6O10
687.65
8
C56H88N12O16Zn2
1316.16
296(2)
298(2)
295(2)
296(2)
296(2)
296(2)
296(2)
Crystal system
Space group
Monoclinic
P21/n
Monoclinic
P21/n
Triclinic
P−1
Monoclinic
P21/n
Triclinic
P−1
Triclinic
P−1
Monoclinic
P21/n
Monoclinic
P21/n
Absorption
mm-1
coefficient
0.835
1.042
0.772
0.687
0.858
0.773
0.557
0.795
(
)
Crystal Size (mm3)
0.21 x 0.18 x 0.12
8.2914(5)
16.1445(11)
10.6129(8)
90
0.20 x 0.16 x 0.10
8.3307(3)
14.1874(6)
12.8931(6)
90
0.22 x 0.18 x 0.13
7.683(4)
0.22 x 0.19 x 0.12
12.485(3)
24.953(4)
12.740(5)
90
0.16 x 0.14 x 0.14
8.4573(2)
9.7343(4)
10.4636(3)
85.503(2)
84.5250(10)
83.350(3)
849.75(5)
1
0.18 x 0.18 x 0.16
13.5412(4)
15.0791(5)
16.5619(6)
105.868(2)
90.082(3)
109.282(2)
3054.71(18)
2
0.22 x 0.16 x 0.12
15.8030(11)
7.6020(4)
15.2690(11)
90
0.11 x 0.08 x 0.06
9.3617(4)
12.0419(7)
29.4879(19)
90
a (Å)
b (Å)
c (Å)
α (°)
13.531(5)
14.275(3)
103.842(3)
105.110(4)°
94.021(5)°
1377.1(9)
2
β (°)
γ (°)
101.711(4)
90
94.448(4)
90
116.562(5)
90
109.770(4)
90
93.964(2)
90
V (Å3)
1391.08(17)
2
1519.26(11)
2
3550.1(17)
2
1726.2(2)
2
3316.3(3)
2
Z
F(0 0 0)
570
704
604
1448
383
1316
730
1392
Reflections collected
Independent Reflections
Final R indices
13062
11294
24353
37479
7807
10668
15709
26951
3462
2983
6594
9272
2991
6290
3957
6913
R1 = 0.0324,
wR2 = 0.0811
R1 = 0.0494,
wR2 = 0.0922
0.71073
R1 = 0.0667,
wR2 = 0.1788
R1 = 0.0733,
wR2 = 0.1816
0.71073
R1 = 0.0412,
wR2 = 0.0828
R1 = 0.0831,
wR2 = 0.0955
0.71073
R1 = 0.0549,
wR2 = 0.1523
R1 = 0.0866,
wR2 = 0.1762
0.71073
R1 = 0.0377,
wR2 = 0.1220
R1 = 0.0430,
wR2 = 0.1268
0.71073
R1 = 0.0744,
wR2 = 0.2216
R1 = 0.1262,
wR2 = 0.2430
0.71073
R1 = 0.0654,
wR2 = 0.1490
R1 = 0.1489,
wR2 = 0.1901
0.71073
R1 = 0.0576,
wR2 = 0.1348
R1 = 0.1157,
wR2 = 0.1605
0.71073
R indices (all data)
λ(Å)
Dc(Mg/m3)
1.295
1.480
1.405
1.298
1.429
1.361
1.323
1.318
Data
/
restraints
/
3462 / 0 / 151
2983 / 0 / 188
6594 / 0 / 316
9272 / 11 / 390
2991 / 0 / 209
6290 / 6 / 751
3957 / 8 / 216
6913 / 8 / 398
parameters
Goodness of fit on F2
CCDC No.
1.026
1.194
1.000
1.021
1.260
1.098
0.969
1.024
1474746
1474748
1474747
1002905
1474768
1031564
1474769
1474770
2
1

2.4.6 Thermogravimetric analysis
The simultaneous thermogravimetric and differential thermal analyses (TG-
DTA) curves of complexes 1, 3, 6, 7 and 8 are shown in Fig. S12. Complexes 2 and 5
were not studied as they contain the perchlorate anion. Complex 1 (Fig. 10) begins to
lose weight above 100 °C. The two lattice water molecules are lost up to 200 °C
endothermically, as shown in the DTA curve (Calc. 6.2%; obs. 6.8%). In the next
step, a loss corresponding to 1/2L1 ligand and two chloride ions is observed at 300 °C
(
Calc. 48.0%; obs. 48.9%). Loss of another L1 ligand starts at 373 °C (Calc. 32.8%;
obs., 32.2%), leaving 13.5% CoO behind.
The TGA curves for complex 3 show the weight loss of a nitrate anion and
N,N-diisopropyl chain at 200 °C (Calc. 25.0%; obs. 24.5%). The remaining organic
part is lost up to 300°C (Calc. 55.7%; obs. 56.7%) , leaving AgO as the residue.
Complex 6 loses two coordinated water molecules between 90 and 200 °C
endothermically (calc. 6.0%; obs. 6.7%). Beyond this temperature, all the organic part
and a nitrate moiety are lost at 270 °C exothermically (calc. 72.5%; obs. 73.0%). Finally,
the complex loses –NO
residue at 600 °C.
Complex 7 loses one lattice and two coordinated water molecules between 90 and
2
(Calc. 7.6%; obs. 8.2 %), leading to the formation of a black CuO
200 °C endothermically (Calc. 7.6%; obs. 8.2%). Beyond this temperature there is a weight
loss in two steps between 220 and 400 °C, corresponding to the loss of two ligand
2
molecules and a nitrate moiety in first step (Calc. 81.8%; obs. 81.4%) and a -NO group in
the second step (Calc., 9.1%; obs., 8.3%), leaving behind a black CoO residue at 600 °C
Complex 8 is devoid of any coordinated water. Weight begins to decrease at
‾
1
00 °C with the loss of two NO
3
anions (Calc., 9.5%; obs., 9.9%). A subsequent
weight loss is observed up to 450 °C, corresponding to the organic moieties, leaving
behind two equivalents of zinc(II) nitrate (calc. 71.1%; obs. 69.8%).
2
2

100
1
3
6
7
8
80
60
40
20
0
100
200
300
400
500
600
700
o
Temp C
Fig. 10. TGA plots of 1, 3, 6, 7 and 8 as a function of temperature between 20 and
800 °C.
2.4.7 UV-vis spectral studies
The electronic spectra of complexes 1, 5, 6 and 7 were taken in acetonitrile
(
Fig. 11a,b,). Complexes 5 and 6 present a relatively broad band in the visible region,
−1
−1
−1
with λmax. centred at about 666 nm (15015 cm , ε = 74 M cm ) and 693 nm (14430
−1
−1
−1
cm , ε = 77 M cm ), respectively, indicating a distorted octahedral geometry in
each case and the broadness may also be attributed to Jahn−Teller distortions. In
2
2
general, these bands are assigned [57] to the overlapping transitions, dxz, dyz → dx −y
,
2
2
2
2
2
d
xy → dx −y and d → dx −y . In Cu(II) octahedral or pseudo-octahedral environments,
z
this band is usually composed of at least three components that in all the present cases
2
remain unresolved. These bands originate from transitions of the dxy, d
z
and dxz, dyz
2
2
pair to the σ antibonding and half-filled dx − y level and the relative order of these
transitions depends upon the extent of the axial metal-ligand interaction and on the
overall geometry around the metal centre. The λmax values corresponding to single
transitions in these complexes give their 10 Dq values. Thus, crystal field splitting is
more for complexes 5 and 6 due to a stronger N
2
O
2
(solvent)
2
coordination
environment.
2
3

Complex 1 shows a multiple band in the visible region. The absorption spectra
of the six coordinated octahedral and pseudo-octahedral cobalt species exhibit three
4
4
4
4
4
4
1 2 3 2
transitions: T2g← T1g (F) (ν ), A2g← T1g (ν ) [58]. The ν
) and T1g(P) ← T1g (ν
transition occurs when two states cross [1], is very weak (formally two electrons) and
is often not observed. Therefore, generally two principal regions near 8000-10,000
−
1
−1
cm (ν
1
) and the other lying in the visible region near 20,000 cm (ν
3
) are observed
[
58]. The blue coloured solution of complex 1 gave two bands in the visible region at
−
1
−1
−1
λ
max. 687 and 591 nm (14556 and 16920 cm , ε = 251 and 190 M cm ),
transition and indicating a distorted octahedral geometry. The
presence of a doublet due to the ν transition in 1 indicates a considerable lowering of
the O symmetry towards a tetragonally distorted D h symmetry. This is supported by
corresponding to the ν
3
3
h
4
the ‘z out’ tetragonal solid state of the complex (Fig. 11b) as well.
Complex 7, on other hand, is almost perfectly octahedral and shows much less
tetragonal distortion, hence its absorption spectrum conforms to that of an octahedral
3
case with an unsplit ν band. A pink coloured solution of the six coordinated
octahedral polymeric complex 7 in acetonitrile gave a broad band in the visible region
−1
−1
−1
with a maximum near 534 nm (18,726 cm , ε = 74 M cm ) corresponding to the
4
4
T
1g(P) ← T1g (ν
3
) transition, indicating a Co(II) distorted octahedral geometry.
0
0
0
0
.15
.10
.05
.00
0.3
1
7
5
0.2
0.1
0.0
6
4
00
600
800
1000
400
600
800
1000
Wavelength (nm)
Wavelength (nm)
a
b
Fig. 11. (a) UV-vis spectra of copper(II) complexes 5 and 6. (b) UV-vis spectra of
cobalt(II) complexes 1 and 7 in acetonitrile.
2.4.8 Cyclic voltammetric studies
CV studies were performed in dry CH CN with a Pt working electrode (Fig.
3
+
1
2), using the Fc /Fc reference potential and TBAP as a supporting electrolyte in the
potential range +3.0 to −1.0, with an AgCl/Ag reference electrode. The peak
2
4

separation for the signals accounts for electrochemical reversibility. The redox peaks
in the range +1.5 to 0.25 V are reported to be due to the Cu(III)/Cu(II) couple and in
the range −0.5 to −1.2 V, due to the Cu(II)/Cu(I) and Cu(I)/Cu(0) couples [59-62].
Starting from the mononuclear copper(II) complex 5, two-stepped cyclic
voltammograms were observed. Both the redox peaks are irreversible with ∆Ep =
1
.34 V, E1/2 = 0.75 and ∆E = 2.53 V, E1/2 = 0.72, which have been assigned as the
p
metal oxidation from Cu(III)/(II) couple and ligand centred oxidation respectively. 6
also shows two irreversible oxidative responses, the first oxidative response at ∆Ep =
1.42 V, E1/2 = 0.94 corresponds to the Cu(III)/(II)couple and second oxidative
response at ∆Ep = 2.58 V, E1/2 = 0.82 V corresponds to a ligand centred oxidation.
The CV scan of 1 shows an irreversible redox peak at E1/2 = 0.41 V (∆Ep =
2.82 V) corresponding to the Co(III)/Co(II) redox response and indicating easy
accessibility of the Co(III) complex [63,64]. The polymeric complex 7 shows an
irreversible redox couple at E1/2 = −0.28 V (∆Ep = 0.89 V) attributed to the
Co(II)/Co(I) process [65].
For 3, an irreversible redox process appears at E1/2 = 0.14 V (∆Ep = 0.38 V),
corroborating the Ag(0)/Ag(I) process [66]. An additional irreversible anodic signal
appears at Ep
a
= 1.99 V which may be due to the oxidation of the coordinated ligand
[
68]. The voltammogram 4 shows one irreversible redox couple at E1/2 = 0.23 V (∆Ep
=
0.28 V) which is attributed to the reduction process Ag(0)/Ag(I).
2
5

Fig. 12. Cyclic voltammograms of the copper complexes 5 and 6, the cobalt
complexes 1 and 7, and the silver complexes 3 and 4.
2.4.9 Discussion
A summary of the coordination parameters of 1-8 and a few other known complexes
of these four ligands is presented in Table S19. Roman numerals represent complexes
taken from the literature with same ligands.
Coordination mode of ligands: It is seen that the picolinamides L1 and L2 prefer a
bidentate chelating mode as it results in the formation of stable five membered
chelate ring on coordination. The geometry found is generally a tetragonally distorted
octahedron, with the assistance from counter anions, forming ion paired (IP)
complexes. Interestingly the M−O (amide) bond length seems to be dependent on the
dihedral angle between the pyridine and amide planes. A larger rotation of the amide
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6

plane with respect to the pyridine ring takes the amide O atom away from the plane of
the chelate ring formed and thus gives longer M−O bonds. For e.g. in the iso-
structural complexes I, III and 5, the M-O distance decreases significantly with a
decreasing dihedral angle. In I the effect is strong enough to elongate the M−O bonds
so much that they move to the elongated axis in the ‘Z- out’ tetragonal structure,
while the M−N and M−Cl bonds lie in an approximate square plane. In III, 1 and 5
however, it is the anion which forms longer bonds. This fact further contributes
significantly in the tetragonal distortion of an already Jahn-Teller distorted system of
the Cu(II) ion.
The coordination modes of L1 and L2 in 3 and 4 are unique in the sense that
one ligand behaves as monodentate through the pyridine N atom in both complexes,
while the other behaves as bidentate, being chelating in the former and bridging in the
latter. With the metal being the same (Ag (I)), they serve as good examples to see the
influence of steric factors and anions on the coordination modes adopted. A poorly
coordinating perchlorate ion and a sterically hindered L2 ligand forms the dimeric
complex 4, while L1 with a coordinating nitrate ion prefers to give the single
molecular complex 3. The flexibility in coordination requirements around the Ag(I)
ion facilitates these modes.
The isonicotinamide derivatives L3 and L4, in the bidentate bridging
coordination mode, generally end up giving 1 or 2D coordination polymers (except in
complex 8, which is a dimer), with or without the help of counter anions. Unidentate
coordination, when found, is invariably through the pyridine N atom and yields 0D
coordination complexes with the assistance from the counter ions or solvent. The
coordination geometries found are either tetrahedral or octahedral. The dihedral
angles between the pyridine and amide groups do not seem to have any obvious
correlation to the M-L distances in the complexes of L3 and L4. However, these
dihedral angles are much larger than those found in the complexes of L1 and L2. This
may be to facilitate the favourite bidentate chelating mode of L1 and L2. For this
reason, the dihedral angles are relatively more when they act in the monodentate
(
complex 3) or bidentate bridging (complex 4) modes.
Non-covalent interactions: The choice of ligands and counteranions in the present
study is based upon maximizing the weak H-bonding interactions, which do not
contribute significantly towards any crystal structure on an individual basis but gain
importance on the ground of sheer numbers. A plethora of weak interactions many a
times offsets the effect of fewer strong interactions.
2
7

C−H···O(amide) interactions: The hydrogen bonding interactions between amide and
alkyl groups play a significant role in organic crystal engineering and are also
important in metal-organic systems because of the presence of these fuctional groups
in most the organic compounds [40]. Presently, the fully substituted amide nitrogen
atoms forbid the possibility of strong amide-amide H-bonding [68], nonetheless the
presence of alkyl groups ensures a gain in terms of a large number of weak H-bonding
interactions in the periphery domain. Further, the pyridyl nitrogen atom is invariably
involved in connecting metal centre,s whereas the amide groups may or may not be
coordinated and thus are an important H-bonding asset. This is more so because of the
changed basicity and hence of the H-bond acceptance of a metal coordinated amide
group in the ligand domain [68]. In the present work, all complexes show C−H···O
H-bonding interactions between the amide O atom and alkyl groups.
δ+
δ-
C−H···X(Halide) interactions: The polarity of the metal-halogen bond (M −X ) in
halide containing complexes makes the halogen atom a better hydrogen bond acceptor
than their carbon-bound counterparts. This hydrogen bond acceptor tendency, along
with the directional properties of metal halides, can be used as molecular building
blocks in the field of crystal engineering [40]. As put forth by Aakeroy, “C−H···X
interactions can tilt the balance between several options of stronger bonded networks
thus acting as an important steering force in solid state assembly”. In all the
complexes having halide anions (I, II, IV, V, VII, IX, X and 1), C−H···X type
interactions are present and play a significant role in increasing the dimensionality of
the crystal structures. For an evaluation of the strength of this type of force, the mean
normalized hydrogen bond distances (RHX = (dH···X)/(r
for all these complexes on basis of van der Waals radii of hydrogen (1.09Å) , chlorine
1.75Å) and bromine (1.85 Å) atoms.The results are summarized in Table S20, which
H
+ r )) [69] were calculated
X
(
shows that the RHX distances for one interaction (for complexes 1, I, VII and IX) and
two interactions (for complex IV) are less than the reported values for strong
interactions [70]. For complexes II, V and X these values are slightly larger than the
reported values. For complex 1, the C−H···Cl−M interaction is even less than the RHX
value for C−H··· F−M (0.943 Å), which indicates that a relatively strong interaction is
present.
C−H(alkyl)···π and C−H(aromatic)···π interactions: C−H···π interactions lack
directionality because of poor polarization of the donor acceptor system, but in the
presence of strong donor and acceptor groups they participate in supramolecular
assemblies [40]. In the present study, C−H···π interactions between pyridine rings and
2
8

alkyl groups help to increase the dimensionality of the structures. C−H···π
interactions between pyridine rings and hydrogen atoms of other rings are also present
in many complexes. These interactions are summarized in Table S21. For complex 3
and X, such interactions between a “T shaped motif” (edge to face) are also found.
Many times C−H···π interactions also supplement these non-covalent interactions to
form 2D/3D H-bonded networks.
π···π interactions: Aromatic moieties can contribute to self-assemblies through π ···π
interactions. The T-shaped motif is the most stable state because of π – σ attraction.
The ‘face to face’ conformation is unstable because of electrostatic repulsions
between the negative charge of ring cloud systems. Due to these repulsions, a slipping
packing arrangement takes place, which is also a stable conformation [40]. The
presence of substituents and heteroatoms disturbs the uniformity of the charge
distribution and develops a partial atomic charge and a permanent dipole which give
rise to dipole-dipole and induce dipole interactions. In addition, an aromatic nitrogen
atom coordinated to the metal atom enhances the electron withdrawing effect due to
its positive charge, resulting in an increase in the tendency to stack because of low π
electron density. For the present study, π ···π interactions are found in complexes 4, 7,
8, I, II, V, VII, VIII and XI. A summary of the displacement angles and d offset
values is shown in Table S22. Centroid to centroid distances lying in the range 3.3-3.8
Å show strong to moderate π ···π interactions. Similarly, displacement angles and d
o
offset values in the range 17-28 and 1.0-1.92 Å show that slipping is taking place. In
complexes I, II, V and XI high values of the displacement angle and d offset are
present as compared to other complexes. Incidentally all these complexes contain the
copper(II) ion which is highly electronegative (χ =2.00) compared to other metals in
the first transition series [71]. More shifting of the π cloud also becomes responsible
for the high d offset value. This decrease in the overlap region between the rings is
driven by the π–σ attraction, thus these interactions are more C−H···π (Table S21b)
rather than π ···π interactions [72].
2.5
Antimicrobial activity
Various infectious diseases and increasing microbial resistance are major
st
concerns in the 21 century, revealing the steadily decreasing potencies of prevalent
antibiotics. To combat this problem, there is a need to explore various resources
which can expand the antibiotic spectrum [73]. Keeping this in mind, all the twelve
2
9