Group 10 metal complexes of dithiocarbamates derived from primary anilines: Synthesis, characterization, computational and antimicrobial studies

Article abstract of DOI:10.1016/j.poly.2018.10.073

Dithiocarbamate ligands obtained from primary amines, namely N-phenylaniline, 4-methylaniline and 4-ethylaniline, and represented as L¹, L² and L³ respectively, have been used to prepare some Ni(II), Pd(II) and Pt(II) complexes. The complexes were characterized by NMR, FTIR spectroscopy and elemental analysis. The spectroscopic data showed that the ligands were chelated to the metal ions in a bidentate mode. The complexes [Pt(L¹)₂], [Pt(L³)₂], and [Pd(L¹)₂] were further characterized by single crystal X-ray analysis and distorted square planar geometries were confirmed in all cases. The magnetic susceptibility measurements of the nickel complexes suggest a square planar geometry for the diamagnetic compounds. Thermal decomposition studies of the complexes gave their respective metal sulfides as residues. Geometry optimization and harmonic frequency calculations of the ligands and the complexes were carried out using density functional theory (DFT). Molecular electrostatic potential energy calculations were used to support the coordination through the dithio-sulfur groups of the three dithiocarbamate ligands. Non-covalent interaction (NCI) theory analysis was used to reveal the different intra and intermolecular interactions in the hydrogen bonded structures of the platinum and palladium complexes. Antimicrobial screening, conducted using selected microbes, showed that the complexes gave moderate to very active antimicrobial activities against Gram negative (Escherichia coli, Klebsiella pneumonia and Pseudomonas aeruginosa), Gram positive (Bacillus cereus and Staphylococcus aureus) and fungi (Candida albicans and Aspergillus flavus) organisms at a concentration of at 50 μg/mL. However, the [Pt(L³)₂] complex gave the best antimicrobial properties with an inhibitory zone range of 8–26 mm.

Full text of DOI:10.1016/j.poly.2018.10.073

Accepted Manuscript
Group 10 metal complexes of dithiocarbamates derived from primary anilines:
synthesis, characterization, computational and antimicrobial studies
Felicia F. Bobinihi, Damian C. Onwudiwe, Anthony C. Ekennia, Obinna C.
Okpareke, Charmaine Arderne, Joseph R. Lane
PII:
S0277-5387(18)30722-8
https://doi.org/10.1016/j.poly.2018.10.073
POLY 13551
DOI:
Reference:
Polyhedron
To appear in:
Received Date:
Accepted Date:
14 September 2018
27 October 2018
Please cite this article as: F.F. Bobinihi, D.C. Onwudiwe, A.C. Ekennia, O.C. Okpareke, C. Arderne, J.R. Lane,
Group 10 metal complexes of dithiocarbamates derived from primary anilines: synthesis, characterization,
computational and antimicrobial studies, Polyhedron (2018), doi: https://doi.org/10.1016/j.poly.2018.10.073
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Group 10 metal complexes of dithiocarbamates derived from
primary anilines: synthesis, characterization, computational and
antimicrobial studies
4
7
1,2,3 Felicia F. Bobinihi, 1,2Damian C. Onwudiwe*, Anthony C. Ekennia, 5,6Obinna C. Okpareke, Charmaine
Arderne, 6Joseph R. Lane.
1Material Science Innovation and Modelling (MaSIM) Research Focus Area, Faculty of Natural and
Agricultural Sciences, North-West University (Mafikeng Campus), Private Bag X2046, Mmabatho, South-
Africa.
2Department of Chemistry, School of Physical and Chemical Sciences, Faculty of Natural and Agricultural
Sciences, North-West University (Mafikeng Campus), Private Bag X2046, Mmabatho 2735, South Africa.
3Federal college of Education, P.M.B. 1026, Okene, Kogi-State, Nigeria
4Department of Chemistry, Alex Ekwueme Federal University Ndufu-Alike (AE-FUNAI), PMB 1010
Abakaliki, Ebonyi State, Nigeria.
5Department of Pure and Industrial Chemistry, University of Nigeria, Nsukka 410001, Enugu State, Nigeria.
6 School of Science, University of Waikato, Private Bag 3105, Hamilton, New Zealand.
7Department of Chemistry, University of Johannesburg, PO Box 524, Auckland Park, Johannesburg 2006, South
Africa.
* Corresponding author
E-mail: Damian.Onwudiwe@nwu.ac.za Tel: (+27) 18 389 2545; Fax: +27 18 389-2420
Abstract
Dithiocarbamate ligands obtained from primary amines, namely N-phenylaniline, 4-
methylaniline and 4-ethylaniline, and represented as L1, L2 and L3 respectively, have been
used to prepare some Ni(II), Pd(II) and Pt(II) complexes. The complexes were characterized
by NMR, FTIR spectroscopy and elemental analysis. The spectroscopic data showed that the
ligands were chelated to the metal ions in a bidentate mode. The complexes [Pt(L1)2],
[Pt(L3)2], and [Pd(L1)2] were further characterized by single crystal X-ray analysis and
distorted square planar geometries were confirmed in all cases. The magnetic susceptibility
measurements of the nickel complexes suggest a square planar geometry for the diamagnetic
compounds. Thermal decomposition studies of the complexes gave their respective metal
sulfides as residues. Geometry optimization and harmonic frequency calculations of the
ligands and the complexes were carried out using density functional theory (DFT). Molecular
electrostatic potential energy calculations were used to support the coordination through the
dithio-sulfur groups of the three dithiocarbamate ligands. Non-covalent interaction (NCI)
1

theory analysis was used to reveal the different intra and intermolecular interactions in the
hydrogen bonded structures of the platinum and palladium complexes. Antimicrobial
screening, conducted using selected microbes, showed that the complexes gave moderate to
very active antimicrobial activities against Gram negative (Escherichia coli, Klebsiella
Pneumonia and Pseudomonas Aeruginosa), Gram positive (Bacillus cereus and
Staphylococcus aureus) and fungi (Candida albicans and Aspergillus flavus) organisms at a
concentration of at 50 g/mL. However, the [Pt(L3)2] complex gave the best antimicrobial
properties with an inhibitory zone range of 8-26 mm.
Keywords: Dithiocarbamates; Group 10 metals; Antimicrobial; DFT
Introduction
Dithiocarbamates derived from primary and secondary amines have been reported to have
similar ligating properties to different metal ions. However, metal complexes of secondary
amine-derived dithiocarbamates [(R2CNS2)2M] have been extensively studied owing to their
greater stability than those of the primary amine analogues [(RHCNS2)2M]. The latter can
easily be decomposed to give the corresponding isothiocyanates due to the presence of an
acidic proton on the nitrogen atom of the thioureide [1, 2]. However, the stabilities of metal
complexes of dithiocarbamates derived from primary amines have been reported to improve
on coordination to metal ions and on storage under low temperatures [3]. The NH site on
metal dithiocarbamate complexes of the form [(RHCNS2)2M] stimulates the transfer of a
proton between ligands through N-H⋯S hydrogen bonds [4]. They are also capable of
providing hydrogen bonding for basic anion N-H deprotonation, which could serve as the
signal-output of certain anion-receptor interactions [5]. The presence of the acidic proton
(NH) has been reported to promote the lipophilicity of the metal complexes, thus improving
permeability through cell walls of microbial strains [3], and enhancing their antimicrobial
potentials compared to their N, N analogues of the form [(R2CNS2)2M].
Group 10 metal complexes of aryl or alkyl dithiocarbamates synthesized from secondary
amines have been extensively reported [6]. The success of cisplatin inspired an extensive
search for metal complexes of this triad with similar or better anticancer potentials. A
nickel(II) bis(dithiocarbamate) complex was first reported by Delepine, while palladium(II)
and platinum(II) bis(dithiocarbamate) complexes were first prepared by Malatesta [7]. Since
2

then, group 10 metal complexes have been explored in a number of biological areas due to
their antimicrobial and anticancer properties [8-11]. They have also been reported to be good
precursors for nickel sulfide, palladium sulfide and platinum sulfide nanoparticles through the
thermal decomposition of their respective metal dithiocarbamate complexes [12]. However,
nickel(II), palladium(II) and platinum(II) complexes of primary amine-derived aryl or alkyl
dithiocarbamates are rarely studied.
Due to the peculiar characteristics of the acidic proton in the overall antimicrobial potentials
of metal complexes of primary amine derived dithiocarbamates and the chemical diversity of
dithiocarbamate complexes, we report herein group 10 metal complexes of dithiocarbamates
derived from different N-substituted anilines. The effects of the different substituents on the
nitrogen atom of the thioureide group on the spectral properties of the complexes are
discussed and the antimicrobial potentials of the complexes were also investigated.
2.0 Experimental
2.1 Materials
Analar grade nickel(II) chloride, palladium(II) chloride, platinum(II) chloride, N-
phenylaniline, p-methylaniline, p-ethylaniline, ammonium hydroxide and tetrahydrofuran
(THF) were obtained from Sigma–Aldrich. Carbon disulfide and sodium hydroxide were
purchased from Merck. All materials were used as supplied without further purification.
2.2. Physical measurements
An Elementar Vario EL Cube was used for the elemental (CHNS) analysis. Infrared spectra
were recorded on a Bruker Alpha-P FTIR spectrometer in the 400–4000 cm-1 range. A 600
MHz Bruker Avance III NMR spectrophotometer was used for the NMR spectral analysis.
TMS was used as an internal standard. Room temperature magnetic moments and molar
conductance measurements were recorded on Sherwood susceptibility balance MSB Mark 1
and Hanna conductivity model H19991300 meters respectively.
3

2.3 Preparation of the ligands
Preparation of ammonium N-phenyldithiocarbamate (L1)
To an equimolar mixture of ammonium hydroxide (15 mL, 0.05 mol) and N-phenyl aniline
(4.6 mL, 0.05 mol) in an ice bath, carbon disulfide (3.0 mL, 0.05 mol) was added in small
portions. After stirring for 6 h, the solidified mass was filtered under suction, rinsed with
diethyl ether to give white solids, which were stored in the refrigerator.
Yield: 6.97 g, 75%; Selected FTIR, 휈(cm-1): 1452 (C=N), 1226 (C2─N), 1018 (C=S), 3035
(=C-H), 2818 (-C-H), 3283 (N-H). 1H NMR (CDCl3) δ, ppm: 6.89-7.88 (m, 10H, C6H5),
10.12 (s, 2H, NH); 13C NMR (CDCl3) δ, ppm: 121.96-128.23 (C6H5), 196.1 (–CS2).
Preparation of sodium p-methylphenyldithiocarbamate (L2)
Sodium hydroxide (1.00 g, 0.025 mol) was dissolved in 20 mL THF and added to a 20 mL
THF solution of CS2 (1.5 mL, 0.025 mol). To this solution, p-methylaniline (2.70 g, 0.025
mol) was added and the solution was stirred for 6 h under an N2 atmosphere. The resulting
precipitate was filtered and re- precipitated in acetone to give a white precipitate which was
dried under vacuum.
Yield: 4.0 g, 74%; Selected FTIR, 휈(cm-1): 1502 (C=N), 1215(C2─N), 988 (C=S), 3031 (=C-
H), 2920 (-C-H), 3326 (N-H); 1H NMR (CDCl3) δ, ppm: 6.96-7.81 (m, 10H, C6H5), 10.10 (s,
2H, NH), 2.34 (s, 6H, CH3). 13C NMR (CDCl3) δ, ppm: 123.5-128.9 (C6H5), 21.3 (CH3),
196.5 (–CS2).
Preparation of sodium p-ethylphenyldithiocarbamate (L3)
Sodium hydroxide (2.0 g, 0.05 mol) dissolved in 5 mL of water was added to carbon disulfide
(3.0 mL, 0.05 mol) and the resulting solution was stirred in an ice bath. After 30 min, 4-ethyl
aniline (6.2 mL, 0.05 mol) was added to the solution and stirred for 3 h. The resulting
precipitate was filtered and re-precipitated in acetone to give a white precipitate, which was
dried under vacuum.
Yield: 7.10 g, 65%: Selected FTIR, 휈(cm-1): 1507 (C=N), 1225 (C2─N), 1014 (C=S), 3019
(=C-H), 2921 (-C-H), 3201 (N-H):
1H NMR (CDCl3) δ, ppm: 7.09-7.46 (m, 10H, C6H5), 10.10 (s, 2H, NH), 1.19 (t, 6H, CH3),
4

2.58 (q, 4H, CH2); 13C NMR (CDCl3) δ, ppm: 120.50-144.0 (C6H5), 15.73 (CH3), 27.72
(CH2), 179.30 (–CS2).
Preparation of the metal dithiocarbamate complexes (ML1-ML3, M = Ni, Pd, Pt)
An aqueous solution of 0.625 mmol of the respective metal salts (NiCl2∙6H2O, 0.149 g;
Na2PdCl4, 0.184 g; K2PtCl4, 0.259 g) was added to a stirring 25 mL aqueous solution of
1.250 mmol of the different ligands, L1 (0.233 g), L2 (0.256 g) and L3 (0.274 g). The reaction
mixture was stirred for 1 h and the different resulting colored precipitates were filtered off,
washed with cold ethanol and dried in a vacuum.
(1) [Ni(L1)2]. Yield: 1.10 g, 81%; M.Pt.: 205-207 °C; Selected FTIR, 휈(cm-1): 1468 (C=N),
1
1292 (C2─N), 958 (C=S), 3166 (=CH−H), 2953 (H2C−H), 3356 (NH), 411 (Ni-S); H NMR
13
(CDCl3) δ, ppm: 7.2-7.4 (m, 10H, C6H5), 8.56 (s, 2H, NH); C NMR (CDCl3) δ, ppm: 129-
137 (C6H5), 196.7 (–CS2); Anal. calc for C14H12N2S4Ni (395.21): C, 42.55; H, 3.06; N, 7.08;
S, 32.22. Found: C, 42.05 H, 3.46; N, 7.50; S, 32.62%; eff (BM): 0.12.
(2) [Pd(L1)2]. Yield: 0.75g, 55%; M.Pt.: 210-212 °C; Selected FTIR, 휈(cm-1): 1475 (C=N),
1298 (C2─N), 965 (C=S), 3226 (=CH−H), 2999 (H2C−H), 3365 (NH), 408 (Pd–S); 1H NMR
(CDCl3) δ, ppm: 7.3-7.6 (m, 10H, C6H5), 9.79 (s, 2H, NH); 13C NMR (CDCl3) δ, ppm: 134-
145 (C6H5), 196.7 (CS2 ); Anal. calc for C14H12N2S4Pd (442.94): C, 37.96; H, 2.73; N, 6.33;
S, 28.95. Found: C, 37.45; H, 2.34; N, 6.86; S, 28.40%; eff (BM): 0.25.
(3) [Pt(L1)2]. Yield: 0.65 g, 70%; M.Pt.: 240-243 °C; Selected FTIR, 휈(cm-1): 1494 (C=N),
1
1312 (C2─N), 977 (C=S) 3154 (=CH−H), 3001, 3375(NH); H NMR (CDCl3) δ, ppm: 7.3-
13
7.9 (m, 10H, C6H5), 9.25 (s, 2H, NH). C NMR (CDCl3) δ, ppm: 135-145 (C6H5), 196.9 (–
CS2), 424 (Pt–S); Anal. calc. for C14H12N2S4Pt (531.60): C, 31.63; H, 2.28; N, 5.27; S, 24.13.
Found: C, 32.22; H, 2.52; N, 5.20; S, 24.67%; eff (BM): 0.35.
(4) [Ni(L2)2]. Yield: 0.60g, 44%; M.Pt.: 215-217 °C; Selected FTIR, 휈(cm-1): 1508 (C=N),
1
1208 (C2─N), 996 (C=S), 3001 (=C-H), 2917 (-C−H), 3185 (N-H), 411 (Ni–S); H NMR
(CDCl3) δ, ppm: 7.0-7.4 (m, 10H, C6H5), 7.25(s, 2H, NH), 2.34 (s, 6H, CH3); 13C NMR
(CDCl3) δ, ppm: 123.5-128.9 (C6H5), 21.3 (CH3), 196.5 (–CS2); Anal calc for C16H16N2S4Ni
(423.26): C, 45.18; H, 4, 26; N, 6.59; S, 30.16; Found: C, 45.10; H, 4.23; N, 6.60; S, 30.20%;
eff (BM): 0.09.
5

(5) [Pd(L2)2]. Yield: 0.71g, 51%; M.Pt.: 230-233 °C; Selected FTIR, 휈(cm-1): 1511 (C=N),
1209 (C2─N), 979 (C=S), 3032 (=CH−H), 2920 (H2C−H), 3183 (N-H), 405 (Pd–S);1H NMR
(CDCl3) δ, ppm: 7.0-7.5 (m, 10H, C6H5), 9.68 (s, 2H, NH), 2.34 (s, 6H, CH3). 13C NMR
(CDCl3) δ, ppm 123-137 (C6H5), 21.2 (CH3), 207.2 (–CS2); Anal calc for C16H16N2S4Pd
(470.99): C, 40.63; H, 3.84; N, 5.92; S, 27.11. Found: C, 40.54; H, 3.57; N, 5.90; S, 26.93%;
eff (BM): 0.18.
(6) [Pt(L2)2]. Yield: 0.64g, 55%; M.Pt.: 248-250 °C; Selected FTIR, 휈(cm-1): 1512 (C=N),
1210 (C2─N), 974 (C=S), 3025 (=CH−H), 2980 (H2C−H), 3180 (N-H), 420 (Pt–S); 1H NMR
13
(CDCl3) δ, ppm: 7.09-7.50 (m, 10H, C6H5), 9.64 (s, 2H, NH), 2.30 (s, 6H, CH3); C NMR
(CDCl3) δ, ppm: 123-137 (C6H5), 21.04 (CH3), 210.0 (–CS2). Anal calc for C16H16N2S4Pt
(559.66): C, 34.21; H, 3,23; N,4.99; S, 2.83. Found: C, 34.03; H, 3.43; N, 5.03; S, 2.75%;
eff (BM): 0.22.
(7) [Ni(L3)2]. Yield: 0.50 g, 83 %; M.Pt.: 209-212 °C; Selected FTIR, 휈(cm-1): 1533 (C=N),
1
1298 (C2─N), 1019 (C=S), 3205 (=CH−H), 2961 (H2C−H), 3205 (N-H), 442 (Ni–S); H
NMR (CDCl3) δ, ppm: 6.50-7.18 (m, 10H, C6H5), 8.31 (s, 2H, NH), 1.25 (t, 6H, CH3), 2.6 (q,
13
4H, CH2); C NMR (CDCl3) δ, ppm: 115.0-129.0 (C6H5), 14.5 (CH3), 28.2 (CH2), 220.0 (–
CS2). Anal. calc for C18H22N2S4Ni (453.33): C, 47.70; H, 4.89; N, 6.19; S, 28.29. Found: C,
47.20; H, 4.45; N, 6.60; S, 28.65%; eff (BM): 0.10.
(8) [Pd(L3)2]. Yield: 0.56 g, 84 %; M.Pt.: 240-243 °C; Selected FTIR, 휈(cm-1): 1535 (C=N),
1
1298 (C2─N), 1018 (C=S), 3206 (=CH−H), 2961 (H2C−H), 3208 (N-H), 419(Pd– S); H
NMR (CDCl3) δ, ppm: 6.80-7.30 (m, 10H, C6H5), 8.40 (s, 2H, NH), 5.90 (t, 6H, CH3), 7.30
(q, 4H, CH2); 13C NMR (CDCl3) δ, ppm: 120.0-145.0 (C6H5), 20.8 (CH3), 45.2 (CH2), 185.0
(–CS2). Anal. calc. for C18H22N2S4Pd (501.06): C, 43.15; H, 4.43; N, 5.59; S, 25.59. Found:
C, 43.50; H, 4.80; N, 5.30; S, 25.31%; eff (BM): 0.25.
(9) [Pt(L3)2]. Yield: 0.45g, 61 %; M.Pt.: 248-250 °C; Selected FTIR, 휈(cm-1): 1534 (C=N),
1298 (C2─N), 964 (C=S), 3205 (=CH−H), 2961 (H2C−H), 2961, 3205 (N-H), 419 (Pt-S); 1H
NMR (CDCl3) δ, ppm: 7.16-7.78 (m, 10H, C6H5), 8.44 (s, 2H, NH), 1.25 (t, 6H, CH3), 2.65
(q, CH2). 13C NMR (CDCl3) δ, ppm: 120.50-144.07 (C6H5), 15.73 (CH3), 27.74 (CH2),
179.30 (–CS2). Anal. calc. for C18H22N2S4Pt (589.72): C, 36.66; H, 3.16; N, 4.75; S, 21.74;
Found: C, 36.20; H, 3.35; N, 4.20.; S, 21.20%; eff (BM): 0.29.
6

2.4 X-ray crystallography
Data collection, structure solution and refinement for [Pt(L1)2], [Pt(L3)2] and [ Pd(L1)2]
Pale yellow crystals with approximate dimensions 0.069  0.231  0.287, 0.086  0.134 
0.282 and 0.256 x 0.374 x 0.473 mm3 were used for [Pd(L1)2], [Pt(L1)2] and [Pt(L3)2]
respectively. They were selected under oil in ambient conditions and attached to the tip of a
MiTeGen MicroMount©. The crystals were each mounted in a stream of cold nitrogen at
100(2) K and centred in the X-ray beam using a video camera. The crystal evaluation and
data collection were performed on a Bruker APEXII CCD DUO diffractometer with Mo K
(λ = 0.71073 Å) radiation and the diffractometer to crystal distance was 6.00 cm [13]. The
initial cell constants were obtained from three series of  scans at different starting angles.
Each series consisted of 12 frames collected at intervals of 0.5º in a 6º range about  with an
exposure time of 10 seconds per frame. The reflections were successfully indexed by an
automated indexing routine built in the APEX2 program suite. The final cell constants were
calculated from a set of 6500, 6351 and 9210 strong reflections for [Pd(L1)2], [Pt(L1)2] and
[Pt(L3)2] respectively from the actual data collection. The data were collected using a
calculated set of parameters for the collection routine, consisting of and  scans to survey
the reciprocal space to the extent of a resolution of 0.75 Å. A total of 26240, 25200 and
15653 data for [Pd(L1)2], [Pt(L1)2] and [Pt(L3)2] respectively were harvested with exposure
times of 10 sec per frame. These highly redundant datasets were corrected for Lorentz and
polarization effects. The absorption correction was based on fitting a function to the
empirical transmission surface as sampled by multiple equivalent measurements [14]. The
systematic absences in the diffraction data were uniquely consistent for the space group Pbca,
that yielded chemically reasonable and computationally stable results of refinement [15-20].
A successful solution by Intrinsic phasing methods (SHELXT [5]) provided all non-hydrogen
atoms from the E-map. The remaining hydrogen atoms were located in an alternating series
of least-squares cycles and difference Fourier maps. All non-hydrogen atoms were refined
with anisotropic displacement coefficients. All hydrogen atoms connected to C atoms were
included in the structure factor calculation at idealized positions and were allowed to ride on
the neighbouring atoms with relative isotropic displacement coefficients. There is one
symmetry-independent molecule in the asymmetric unit, consisting of one half of the Pd(II)
complex for [Pd(L1)2], and a Pt(II) complex for [Pt(L1)2] and [Pt(L3)2]. The palladium and
platinum metals are in the +2 oxidation state and charge balance comes from the bound
7

dithiocarbamate molecule connected by only two S atoms (S1 and S2) where the negative
charge is localized on the S1 and S2 atoms. The other half of the molecule is generated by
symmetry with the Pd and Pt atoms (Figure 1) residing on a special position (0.5, 0, 0.5). The
final difference Fourier map was basically featureless for both structures.
2.5 Computational details
DFT calculation studies
Density functional theory (DFT) calculations were carried out using the Gaussian 2009
program suite [21] on the high performance computing facility at the University of Waikato.
Unless otherwise specified, all DFT calculations were completed using the 6-311G+(d,p)
basis set for all main group elements and the LANL2DZ basis set and effective core potential
(ECP) for Pt, Pd and Ni atoms. Geometry optimizations and harmonic frequency calculations
in the gas phase were run with the B3LYP functional using the default optimization criteria.
The absence of imaginary frequencies in the calculated vibrational frequencies of the ligands
and complexes showed that the optimized geometries were minima. The geometries of the
ligands were also re-optimized using three different density functionals, B3LYP, ωB97XD
and M06-2X, in the presence of an implicit solvent [22]. The frontier molecular orbital
energies and reactivity indicators of the ligands and complexes were calculated. Natural Bond
Orbital (NBO) calculations were implemented on the complexes using the CAM-B3LYP
functional. Non-covalent interaction (NCI) theory calculations were implemented on the
hydrogen bonded structures of palladium and platinum. The initial input geometries were
adapted from the crystal structures of the compounds and carried out using the ωB97XD
functional. The NCI iso-surfaces were generated using Bonder, a locally written software
program, with visualisation using VMD.
2.6 Antimicrobial studies
The microbes were clinical isolates collected from the Department of Microbiology, Federal
Teaching Hospital, Abakaliki, Nigeria. They included Gram negative (Escherichia coli,
Klebsiella Pneumonia and Pseudomonas Aeruginosa), Gram positive (Bacillus cereus and
Staphylococcus aureus) and fungi (Candida albicans and Aspergillus flavus) organisms. The
choices of the microbial strains were based on their clinical and pharmacological relevance
[23]. The agar disc diffusion method [24] was utilized for the antimicrobial screening.
8

Sterile Muller–Hinton agar (MHA) was prepared and poured on petri plates. A standard
method was adopted for the microbial culture using 80% dimethylsulfoxide as the solvent
and 1.25 mg/mL 3-(4,5-dimethylthiazol-2-yl)-2, 5-diphenyltetrazolium bromide (MTT) to
monitor the occurrence of a purple colouration [25]. A commercially available antibacterial
drug (Sulfamethoxazole) and an antifungal drug (Ketoconazole) were used as the positive
control drugs, while 80% dimethylsulfoxide was used as a negative control under similar
conditions. Measurements were conducted four times and the experimental results were given
as the mean ± S.D. of all the measurements.
3. Results and discussion
3.1 Synthesis
Scheme 1 presents the preparation of the ligand, which proceeded in the presence of a strong
alkaline solution (ammonium hydroxide or sodium hydroxide) at a very low temperature (in
ice). All the complexes were prepared using the direct reaction of solutions of the metal salts
and ligands in a 1:2 stoichiometric ratio at room temperature. All the complexes formed are
intensely coloured, as expected for d8 transition metals (due to d-d transitions), with different
shades of green for the nickel, orange for the palladium and yellow for the platinum
complexes. They are all insoluble in water, but soluble in different low boiling point organic
solvents such as chloroform, dichloromethane and methanol.
Scheme 1: Preparation of the dithiocarbamate ligands.
3.2 Infra-red spectral studies
The FT-IR spectra of the dithiocarbamate compounds show characteristic bands within the
1450-1580 cm-1 region, due to the stretching vibration of the thioureide band v (N=CS2), and
the 950-1050 cm-1 region, assigned to a v(C-S) vibration [26]. The spectral data were
compared and assigned based on similar reported compounds in the literature [27, 28]. The
FTIR spectra of the ligands gave intense stretching vibrational bands for their thioureide
9

groups at 1452, 1507 and 1508 cm-1 for L1, L2 and L3 respectively, which shifted to higher
wavenumbers in the spectra of the metal complexes. The increase in wavenumber of the C=N
bands compared to those obtained in the spectra of the free ligands could be attributed to the
chelation of the sulfur atoms to the metal ions, which reinforces the double bond character of
the C=N group by strengthening the force constant of the bond, thus giving a higher
wavenumber [29]. In addition, the vibration bands for v(C= N) increase from Ni to Pt due to
the increase in the Lewis acid properties of the metals. The π-donor property also gets weaker
down the group 10 metals, thus raising the C=N vibrations [30]. The effect of the strength of
the electron donating substituents (H, CH3 and CH2CH3) in the para position of the phenyl
ring was observed as the values of the thioureide bands appeared in the order L1 < L2 < L3.
The v(C–S) vibrational band was observed as a single band around 1018-1040 cm-1 in the
spectra of the ligands, and it appeared at relatively lower wavenumbers in the complexes. The
single band indicated a symmetrical bidentate coordination of the ligand to the metal center
[30]. The magnitude of the C-S bands were also in the order L1 < L2 < L3.
3.3 NMR spectroscopic studies
The NMR spectra of the ligands showed the protons of the phenyl ring as multiplets in the
range δ 6.96-7.88 ppm. The peaks were slightly higher in the spectra of the metal complexes.
The increase in chemical shift could be attributed to a weak deshielding brought about by the
shift of electron density towards the nitrogen atom of the NHR′ group, thereby enhancing the
electron density on the sulfur atom via the thioureide π-system [31]. The NH proton peaks
appeared at δ 10.12-10.10 ppm for the uncoordinated ligands but were deshielded to lower
13
chemical shifts in the spectra of the complexes. The C NMR spectra exhibited peak values
within the expected range, but in different environments. In the spectra of the uncoordinated
ligands the -NCS2 peak positions were observed at δ 196.1, 196.5 and 179.3, ppm for L1, L2
and L3 respectively. These peaks resonated around δ 196.7-220.0, 196.7-206.9, 196.9-210
ppm in the spectra of the Ni(II), Pd(II) and Pt(II) complexes. An increase in the -NCS2 peak
position is generally observed down the group of each series, confirming that the heavier
atom complexes have higher δ values than the lighter atoms in the same group [31].
3.4 X-ray crystallography studies
10

The X-ray structures of the bis-(N-phenyldithiocarbamate)palladium(II), bis-(N-
phenyldithiocarbamate)platinum(II), and bis-(4-ethylphenyldithiocarbamate)platinum(II)
complexes are presented in Figure 1, and their crystal and structure refinement data are
summarized in Tables 1 and 2. The coordination environment of the Pd complex shows
symmetric bond lengths with the Pd-S distances Pd(1)-S(1) = 2.3232(5) Å and Pd(1)-S
(2) = 2.3399(5) Å.
In
bis-(N-phenyldithiocarbamate)platinum(II)
and
bis-(4-
ethylphenyldithiocarbamate)platinum(II) complexes, symmetric bond lengths were also
exhibited with Pt-S distances Pt1-S1=2.3201(7) Å and Pt1-S 2=2.3313(7) Å in the former;
and Pt(1)-S(1) = 2.3163 (6) and Pt(1)-S(2) = 2.3216 (6) Å in the later. All the complexes
have a distorted square planar geometry due to the small S-Pd/Pt-S2 bite angles of
75.134(17), 74.75(3) and 74.93 (2)° for [Pd(L1)2], [Pt(L1)2] and [Pt(L3)2], respectively. The
phenyl groups make angles of 86.17(7) and 86.03(7) with the dithiocarbamate planes in the
palladium complex, [Pd(L1)2]. However, in the Pt complex [Pt(L1)2], analogous angles of
87.20(10) and 86.95(10) were made; while in [Pt(L3)2] they were 87.40(8) and 87.47(9) .
The bond lengths for both Pd-S and Pt-S are relatively close, 2.32-2.34 and 2.32-2.33 Å.
Comparing the Pt-S bond lengths in [Pt(L1)2] (2.3201-2.3202 Å), with an unsubstituted
phenyl group in the ligand L1, and the Pt-S bond lengths in [Pt(L3)2] (2.3163-2.3216 Å),
where the ligand L3 bears a p-ethylphenyl group, there is a slight variation. The bond lengths
were slightly reduced by the presence of the ethyl group, which may influence the
intramolecular interactions between the atoms. However, their intermolecular hydrogen
bonding interactions are similar (details of which appear in Table 3 for [Pd(L1)2], Table 4 for
[Pt(L1)2], and Table 5 for [Pt(L3)2]; selected interactions are also shown in Figure 1), which
supports their similar coordination modes and chemical properties and also explains why they
are usually studied as a group [32].
3.5 Thermal studies of the compounds
Figures 2-4 show the overlapped TGA/DTG graphs of the complexes. Relevant data from the
decomposition are presented in Table 6. The metal complexes show similar single step
decomposition patterns in the range 163-294 C. The Ni and Pd complexes of the ligand L1
resulted in residues with a mass of 2.78/2.60 and 3.84/4.10 mg (calculated/found,)
respectively, indicating a 1:1 ratio of the metal to sulfide ((NiS and PdS). The [PtL1 ]
2
complex yielded a 1:2 metal to sulfur ratio with a residual mass of 5.66/5.57 mg, which
confirmed PtS2. The metal complexes of the ligands L2 and L3 followed a similar pattern of
11

single step decomposition, with the onset around 114 and 174 C, which proceeded up to 295
and 423 C, respectively, with various percentage weight losses. All the obtained residues
conformed to a 1:1 metal-sulfur ratio, indicating NiS, PdS, and PtS. The presence of the
methyl and ethyl substituents on the phenyl ring increased the thermal stability of the
complexes and the stability increased in the order L3 > L2 > L1. To show the variation in the
temperature of the maximum rate of decomposition between complexes of the same ligand
moiety, their respective DTG graphs were overlapped. Figures 2b, 3b and 4b present the
overlapped DTG graphs which conspicuously show the changes in their decomposition peaks
as a function of the type of metal ions and also the ligand type. All the complexes are
potential single source precursors for the synthesis of their respective metal sulfide
nanoparticles [33].
3.6 Antimicrobial studies
The compounds were screened at the three different concentrations of 10, 25 and 50 µg/mL.
Significant zones of inhibition were obtained at 50 µg/mL and these were taken as the
minimum inhibitory concentration of the metal complexes against the microbes: Gram
positive (S. Aureus and B. cerues), Gram negative (K. pneumonia, P. aeruginosa and E. coli)
and two fungi organisms (C. albican and A. flavus). The results are presented in Table 7.
The metal complexes gave varied antimicrobial activities which ranged from moderate to
very active. E. Coli was susceptible to all the complexes, while P. Aeruginosa and K.
Pneumonia were susceptible to all the complexes except for [Pd(L1)2] and [Ni(L2)2]
respectively. In general, the Gram negative bacteria organisms were more vulnerable to the
complexes as compared to the Gram positive bacteria organisms. This is probably due to the
presence of an outer protective lipopolysaccharide membrane in Gram positive bacteria
strains which does not permit lipophobic materials into the cell, making penetration of the
complexes intricate [34]. However, the metal complexes of L3 were mostly very active
against both Gram positive and Gram negative bacterial strains. They were the most active of
the metal complexes screened against the microbes.
The compounds displayed better antibacterial activity compared to their antifungal activity.
[Pt(L3)2] had the best antifungal activity against C. albican and A. flavus compared to the
other complexes, but lower activity than Ketoconazole, used as the standard drug. Similarly,
Sulfamethoxazole exhibited the best antibacterial activity against all the bacteria strains
12

compared to the metal complexes except for [Pt(L3)2] against S. aureus. In addition [Pt(L3)2]
exhibited 89, 82 and 77% of the antibacterial activity of Sulfamethoxazole against B. cereus,
P. aeruginosa and S. aureus, respectively. The Ni(II) and Pd(II) complexes of L2 and L3
showed better antimicrobial activities compared to literature reports involving similar
compounds [35]. Similarly, [Pt(L3)2] exhibited better antibacterial properties compared to
related literature studies [36]. Interestingly, the antimicrobial results of the metal complexes
were better than those reported in literature [37, 38]. The compounds could be probable lead
agents in antimicrobial research.
3.7 Computational studies
3.7.1 Molecular electrostatic potential (MEP)
The molecular electrostatic potential (MEP) is a three dimensional cloud around a molecule
created by its nuclei and electrons. It is very helpful in predicting the reactivity of a range of
chemical systems for both electrophilic and nucleophilic sites. A number of chemical
interactions, including drug-receptor interactions, enzyme-substrate interactions and
hydrogen bonding in both chemical and biological systems, are predicted with the use of
molecular electrostatic potentials. The molecular electrostatic potential V(r) is represented as:
푀
푍퐴
|푟 − 푅퐴|
푑ŕ휌(푟)
|ŕ − 푟|
∫
−
푉(ŕ) = ∑
퐴−1
where ZA is the charge on the nucleus A located at RA and ρ(r) is the electronic density
function of the molecule obtained from its computed wave function. V(ŕ) is a real physical
property that can be determined experimentally as well as computationally [21]. In order to
investigate the reactive sites in the compounds, the molecular electrostatic potentials were
evaluated at the B3LYP/6-311++G(d,p) level of theory. The values of the electrostatic
potential at the surface are defined by different colours, with the potential increasing in the
order from red < orange < yellow < green < light blue < blue. Figure 5 shows the
electrostatic potential maps and contour diagrams for the dithiocarbamate ligands. The
negative yellow regions of the MEP map are related to the negative electrostatic potential and
sites for electrophilic attack. These regions are localized around the dithio functionality with
potential values of -0.0585, -0.05836 and -0.06083 au for the three ligands L1, L2 and L3,
13

respectively. The negative blue regions of the electrostatic potential map, on the other hand,
are related to the nucleophilic reactivity and they are localized around the imine hydrogen
atom of the dithiocarbamate with positive potential values of 0.0485, 0.0483 and 0.0508 au
for L1, L2 and L3, respectively. The electrochemical potential maps for the three ligands show
that the dithio-group has the most positive electrostatic potential in the three ligands and, thus,
is the most likely site for metal attack. There was also a slight increase in the electrostatic
potential value with the increase in the alky substituent attached to the phenyl ring.
3.7.2. Frontier molecular orbital energies and chemical reactivity
The frontier molecular orbitals, which includes the highest occupied molecular orbital
(HOMO) and the lowest unoccupied molecular orbital (LUMO), are very important quantum
chemical parameters because they indicate the possibility of electronic transitions in
molecules. They are crucial for the understanding of electron transfer processes in a molecule
during chemical reactions [22]. The HOMO orbital primarily acts as the electron donor and
the LUMO largely acts as the electron acceptor, and the gap between the HOMO and LUMO
determines the chemical stability of the molecule [39]. The HOMO and LUMO orbital
diagrams for the three dithiocarbamate ligands are presented in Figure 6. The orbital
diagrams show that the HOMO orbitals are located on one of the sulfur atoms of the
dithiocarbamate ligand, while the LUMO orbitals are located on the dithiocarbamate carbon
atom for all three ligands. This is an indication of electron donation from the occupied sigma
orbitals of the sulfur atom to the unoccupied π orbitals of the carbon atom. The HOMO-
LUMO energy gap can be used as a stability index, providing an estimation of the reactivity
of the molecule [22]. A large HOMO-LUMO energy gap indicates high stability and low
chemical reactivity. A smaller energy gap indicates less stability and higher reactivity of the
molecule and, thus, an easier electronic transition [21, 40]. The values of the HOMO-LUMO
energy gap presented in Table 8 show a low energy gap at the B3LYP level of theory and this
increases as the density functional changes from B3LYP to M062X, and then to ωB97XD.
The nature of the alkyl substituent on the phenyl ring of the dithiocarbamate ligands did not
have any significant effect on the reactivity of the molecules. Apart from the HUMO and
LUMO energies, there are other reactivity descriptors that can be obtained from the HOMO-
LUMO relationship. The global reactivity descriptors based on Koopmans theorem [41]
14

include the ionisation potential (IP), electron affinity (EA), chemical potential (μ) and the
chemical hardness (η). The global electrophilicity index (ω), proposed by Parr et al [42], was
also evaluated. Lower values of the chemical hardness, electrophilicity index and chemical
potential are related to better electron donating properties. The chemical reactivity parameters
for the dithiocarbamate ligands using three different functionals were investigated and the
results are presented in Table 8. The results indicate no significant effect of using different
functionals on the reactivity parameters. Small values of the chemical hardness η are an
indication of the dominance of the electron donating properties of the ligands. There is,
however, no significant difference in the chemical hardness as the substituent at the para
position of the phenyl ring of the dithiocarbamate ligand increased from H, to CH3 and then
to C2H5 across the three levels of theory. The implication of this is that the three
dithiocarbamate ligands are of similar reactivity and the size or nature of the alkyl substituent
on the phenyl group has little or no significant effect on the kinetic stability of the compounds.
The frontier molecular orbital molecular orbital diagrams for the complexes are also
presented in Figure 7, and the energy gap and reactivity parameters are presented in Table 9.
The molecular orbital diagrams show that the HOMO orbitals are located around the metal
coordination environment for all the complexes while the LUMO orbitals are spread around
most of the compound for the platinum complexes. The LUMO orbitals for the palladium and
nickel complexes are also concentrated around the metal coordination area. The orbital
diagrams and the reactivity parameters in Table 9 show a lot of similarities in the metal
complexes formed from the three dithiocarbamate ligands.
3.7.3. Non-covalent interactions in the hydrogen bonded platinum and palladium complexes
The supramolecular architecture of a compound in the solid state is a function of the nature
and strength of the different non-covalent interactions in the molecule [43]. These non-
covalent interactions, which may include hydrogen bonding, halogen bonding, Van der Waals
or dipole-dipole interactions, are dependent on the electron density around the compound. A
number of topological methods can be applied in order to identify the relationship between
the electron density and molecular reactivity [44-46]. Non-covalent interaction (NCI) theory
is based on analysis of the reduced density gradient, which is represented as:
1
|∇휌|
푠 =
2(3휋2)1/3 휌(퐫)4/3
15

Non-covalent interactions are revealed by regions where the reduced density gradient is
small and the density is low but non-negligible. [47]. The reduced density gradient is plotted
as a function of the electron density ρ, oriented by the sign of λ2 to identify the interaction
types. Attractive interactions appear at λ2 < 0 while repulsive interactions appear at λ2 > 0. A
pictorial 3D representation of these interactions with an RGB (red-green-blue) colouring
scheme is used to rank these interactions. Red represents strong repulsive/destabilizing
interactions, which include ring closure interactions and steric interactions, while blue
regions are associated with highly attractive/stabilizing interactions, including hydrogen
bonding interactions. In between the two extremes are the green regions, indicating weaker
van der Waals interactions. The 3D isosurface plot for the palladium dithiocarbamate
complex is shown in Figure 8. The plot shows only the relevant intermolecular interactions
that are responsible for the hydrogen-bonded chain found in the molecule. The strongest
interaction found in the dimer is the strong stabilizing intermolecular hydrogen bonding
interaction, shown as a blue pill-like isosurface between the dithiosulfur group of one
complex and the N-H group of another complex. Similar intermolecular hydrogen bonding
interactions were reported for
a
hydrogen bonded 2-{[2-(phenylsulfonyl)
hydrazinylidene]methyl}benzoic acid dimer [46]. Another interaction, depicted in the
palladium complex dimer in Figure 7, is the weak intermolecular n-π interaction, shown as a
light green flat isosurface between the dithiosulfur group and the π bond of the adjacent
phenyl ring. This interaction can be classified as van der Waals forces in terms of strength.
Finally, the last interaction, with a large flat boat-shaped isosurface, has a mixture of a
slightly attractive green coloured Van der Waals interaction from the intermolecular Pd-S
interaction and a slightly repulsive yellow coloured interaction resulting from intermolecular
S–S interaction. Similar interactions were found for the platinum dithiocarbamate complex
(picture not shown).
3.7.4 Natural Bond Orbital (NBO) analysis
The natural bond orbital theory is an explicit method of studying intermolecular and
intramolecular interactions in a compound. It provides the basis for investigating the charge
transfer and hyperconjugative interactions in a molecular system [48]. NBO analysis was
carried out for the Pt, Pd and Ni dithiocarbamate complexes by examining the possible
interactions between the filled orbitals (donor) Lewis type NBOs and the empty orbital
16

(acceptor) non-Lewis type NBOs. The hyperconjugative energy of this interaction can be
estimated by second order perturbation theory [49]. The strength of the bond in a molecule
can be estimated by the value of the stabilization energy E2 which is a pointer to the level of
interaction between the electron donor and the electron acceptor [50]. NBO analysis for the
six dithiocarbamate complexes was performed with the CAM-B3LYP functional. Selected
donor acceptor interactions and their stabilization energies are presented in the supplementary
Table S1. The table shows a number of donor – acceptor interactions from the σ lone pair
orbitals of the sulfur atom to the π* orbitals of the conjugated dithiocarbamate carbon atom,
with the highest stabilization energies ranging from 153 to 201 kcal/mol in all the complexes.
This is an indication of electron delocalisation in the complexes. The other donor-acceptor
interactions with high stabilization energies of between 162 and 180 kcal/mol are the n–π*
interactions. The n–σ* (LpS - LP*M, M = metal) interactions, with stabilizations energies of
5.5 − 7.7 kcal/mol, are also prominent in all the metal complexes. Apart from that, there are
other π – π* (BDC – BD*C) interactions with high stabilization energies of between 18.0 and
25 kcal/mol, indicative of relative bond strength. Some strong metal to ligand charge transfer
(MLCT) interactions were also recorded for the platinum dithiocarbamate complexes with
high stabilization energies of between 10 and 20 kcal/mol.
Conclusion
The biological relevance of the acidic proton (NH) in dithiocarbamate derived from primary
amines has motivated the synthesis of substituted and unsubstituted phenyldithiocarbamates.
The ligands were obtained from N-phenylaniline, 4-methylaniline and 4-ethylaniline, which
yielded dithiocarbamate ligands denoted as L1, L2 and L3 respectively. Metal complexes of
Ni(II), Pd(II) and Pt(II) were subsequently synthesized and characterised; quantum
mechanical calculations were used to explore their bonding properties and interactions. NCI
analysis of the complexes indicated the presence of strong stabilizing inter-molecular N-H---
S hydrogen bonding and weaker intermolecular n---π and Pd---S van der Waals interactions.
Frontier molecular orbital analysis indicated the delocalisation of electrons around the
coordination sphere of the metal complexes and an increase in the length of the alkyl
substituent on the phenyl groups showed no significant effect on the HOMO-LUMO energy
gaps. The high stabilisation energies of the metal complexes from second order perturbation
analysis of the fork matrix in the NBO is an indication of a very strong metal-ligand
17

interaction for the complexes. The antimicrobial results showed that antimicrobial activities
followed the order L1 < L2 < L3 (for their metal complexes). The ethyl substituted
dithiocarbamate complexes had the best antimicrobial activity compared to the unsubstituted
and methyl substituted dithiocarbamate complexes. These compounds could be probable lead
agents in antimicrobial research.
Supplementary data
Crystallographic data of the complexes have been deposited with the Cambridge
Crystallographic Data Center allocated with the deposit numbers CCDC 1849649, 1848912
and 1875081. Copy of the data can be obtained free of charge on application to CCDC, 12
Union
Road,
Cambridge
CB2
1EZ,
UK,
fax:
+44
1223
336033,
email:deposit@ccdc.cam.ac.uk.
REFERENCES
[1] A. C. Ekennia, A. J. Odola, Inter. J. Pharm. Bio. Chem. Sci. 2 (2013) 21.
[2] D.C. Onwudiwe, P.A. Ajibade, Mater. Lett. 65 (2011) 3258.
[3] K. Ghosh, V. Mohan, P. Kumar, S.W. Ng, E.R.T. Tiekink, Inorg. Chim. Acta. 416 (2014)
76.
[4] D.E. Lynch, I. McClenaghan, Cryst. Eng. 6 (2003) 1.
[5] K.C. Chang, S.S. Sun, M.O. Odago, A.J. Lees, Coord. Chem. Rev. 284 (2015) 111.
[6] A. C. Ekennia, J. Applied Chem. (IOSR-JAC), 5 (2013) 36.
[7] G. Hogarth. Transition Metal Dithiocarbamates: 1978–2003. (2005) 53. Ed. Kenneth D.
Karlin.in. Progress in Inorganic Chemistry. Wiley
[8] F. Fu, H. Zeng, Q. Cai, R. Qiu, J. Yu, Y. Xiong, Chemosphere 69 (2007) 1783.
[9] S. Kanchi, P. Singh, K. Bisetty, Arab. J. Chem. 7 (2014) 11.
[10] G .Hogarth, Mini Rev. Med. Chem. 12 (2012) 1202.
[11] V. Bala, G. Gupta, V. L. Sharma. Mini Rev. Med. Chem. 14 (2014) 1021.
[12] D. C. Onwudiwe, A. C. Ekennia. Res Chem. Intermed. 43 (2017) 1465.
[13] Bruker-AXS (2014). APEX2. Version 2014.11-0. Madison, Wisconsin, USA.
[14] L. Krause, R. Herbst-Irmer, G. M. Sheldrick, D. Stalke, J. Appl. Cryst. 48 (2015) 3.
18

[15] G. M. Sheldrick, (2013). XPREP. Version 2013/1. Georg-August-Universität Göttingen,
Göttingen, Germany.
[16] G. M. Sheldrick, Acta Cryst. D66 (2010) 479.
[17] G. M .Sheldrick, Acta Cryst. A71 (2015) 3.
[18] G. M. Sheldrick, Acta Cryst. C71 (2015) 3.
[19] O. V. Dolomanov, L. J. Bourhis, R. J. Gildea, J. A. K. Howard, H. Puschmann, J. Appl.
Cryst. 42 (2009) 339.
[20] I. A. Guzei, (2007-2013). Programs Gn. University of Wisconsin-Madison, Madison,
Wisconsin, USA.
[21] A. Zülfikaroğlu, H. Batı, N. Dege, J. Mol. Struct. 1162 (2018) 125.
[22] A.O. Zacharias, A. Varghese, K. Akshaya, M. Savitha, L. George, J. Mol. Struct. 1158
(2018) 1.
[23] W.P. Ross, P.W. Malbrook, Clinical and oral microbiology, Hemispher Publishing
Corporation, New York, 1982.
[24] M. Balouiri, M. Sadiki, S.K. Ibnsouda, J. Pharm. Anal. 6 (2016) 71.
[25] H.J. Simon, E.J. Yin, Appl. Envir. Microbiol. 19 (1970) 573.
[26] A. Manohar, K. Karpagavel, A. Murugan, Int. J. Chem. Tech. Res. 6 (2014) 474.
[27] J. P. Fuentes-Martínez, I. Toledo-Martínez, P. Román-Bravo, P.G.y. García, C. Godoy-
Alcántar, M. López-Cardoso, H. Morales-Rojas, Polyhedron 28 (2009) 3953.
[28] N. Singh, S. Bhattacharya, J. Organomet. Chem.700 (2012) 69.
[29] J. O. Adeyemi,, D. C. Onwudiwe, A. C. Ekennia, R. C. Uwaoma, E. C. Hosten.. Inorg.
Chim. Acta 477 (2018) 148.
[30] E. Sathiyaraj, G. Gurumoorthy, S. Thirumaran, New J. Chem. 39 (2015) 5336.
[31] B.A. Prakasam, K. Ramalingam, G. B. Andrea, Polyhedron 26 (2007) 4489.
[32] A. L. Spek, Acta Cryst. D65 (2009) 148.
[33] D. C. Onwudiwe, J. N. Mugo, H. Madalina, H. Eric, J. Sulfur Chem. 36 (2015) 36.
[34] F. Javed, M. Sirajuddin, S. Ali, N. Khalid, M.N. Tahir, N.A. Shah, Z. Rasheed, M.R.
Khan, Polyhedron 104 (2016) 80.
[35] S. I. Islam, S. B. Das, S. Chakrabarty, S. Hazra, A. Pandey, A. Patra. Adv. Chem. 2016
(2016) 1.
[36] M. Mijajlovic, S. Vasic, I. Radojevic, J. Maksimovic, L. Comic, M. Nikolic, G. Radic.
2nd international electronic conference on medicinal chemistry (2016).
[37] P. A. Ajibade, O. G. Idemudia, A. I. Okoh, Bull. Chem. Soc. Ethiop. (2013) 27.
[38] M. Gaber, H. A. El-Ghamry, M. A Mansour. J. Photochem. Photobio. 354A (2018) 163.
19

[39] A. Eşme, S.G. Sağdınç, Acta A188 (2018) 443.
[40] V. Balachandran, M. Murugan, V. Karpagam, M. Karnan, G. Ilango, Spectrochim. Acta,
130A (2014) 367.
[41] T. Koopmans, Physica 1 (1934) 104.
[42] R.G. Parr, L.v. Szentpaly, S. Liu, J. Am. Chem. Soc. 121 (1999) 1922.
[43] E.E. Oyeka, J.N. Asegbeloyin, I. Babahan, B. Eboma, O. Okpareke, J. Lane, A. Ibezim,
H.H. Bıyık, B. Törün, D.C. Izuogu, J. Mol. Struct. 1168 (2018) 153.
[44] J. Cioslowski, S.T. Mixon, J. Am. Chem. Soc. 114 (1992) 4382.
[45] B. Silvi, A. Savin, Nature 371 (1994) 683.
[46] J.N. Asegbeloyin, D.C. Izuogu, E.E. Oyeka, O.C. Okpareke, A. Ibezim, J. Mol. Struct.
1175 (2019) 219.
[47] J.N. Asegbeloyin, E.E. Oyeka, O. Okpareke, A. Ibezim, J. Mol. Struct. 1153 (2018) 69.
[48] F. Weinhold, C.R. Landis, Chem. Educ. Res. Prac.2 (2001) 91.
[49] J. Chocholoušová, V. Špirko, P. Hobza, Phys. Chem. Chem. Phys. 6 (2004) 37.
[50] M.K. Awad, M.F. Abdel-Aal, F.M. Atlam, H.A. Hekal, Spectrochim. Acta 206A (2019)
78.
20

Figures
(a)
(b)
21

(c)
(d)
22

(e)
(f)
Fig 1: (a) Molecular structure of [Pd(L1)2] with atomic displacement ellipsoids drawn at the
50% probability level; (b) intermolecular hydrogen bonding contacts of [Pd(L1)2]; (c)
molecular structure of [Pt(L1)2] with atomic displacement ellipsoids drawn at the 50%
probability level; (d) intermolecular hydrogen bonding contacts of [Pt(L1)2]; (e) molecular
structure of [Pt(L3)2] with atomic displacement ellipsoids drawn at the 50% probability level;
(f) intermolecular hydrogen bonding contacts of [Pt(L3)2]. In all three diagrams, the NH∙∙∙S
interactions are also shown as red dashed lines and symmetry positions for all three structures
are identical at: (i) 3/2x, ½+y, z, (ii) 1x, y, 1z, (iii) ½+x, ½y, 1z
23

120
100
80
25
20
15
10
5
(a)
(b)
60
40
0
-5
20
0
200
400
600
800
0
200
400
600
800
Universal V4.5A TA Instruments
Universal V4.5A TA Instruments
Temperature (°C)
Temperature (°C)
Fig.2: (a) TG and (b) DTG of the Ni (green), Pd (purple) and Pt (red) complexes of L1
40
100
(a)
(b)
30
20
10
0
80
60
40
20
-10
0
200
400
600
800
0
200
400
600
800
Universal V4.5A TA Instruments
Universal V4.5A TA Instruments
Temperature (°C)
Temperature (°C)
Fig.3. (a) TG and (b) DTG of the Ni (green), Pd (purple)and Pt (red) complexes of L2
40
30
20
10
0
120
100
80
60
40
20
0
(b)
(a)
-10
0
200
400
600
800
0
200
400
600
800
Universal V4.5A TA
Temperature (°C)
Universal V4.5A TA
Temperature (°C)
Fig.4: (a) TG and (b) DTG of the Ni (green), Pd (purple), and Pt (red) complexes of L3
24

Fig. 5: Electrostatic potential maps and contour diagrams for the dithiocarbamate ligands.
Contour value = 0.004 au.
25

CODE
HL1
HOMO ORBITALS
LUMO ORBITALS
HL2
HL3
Fig. 6: Frontier molecular orbital diagrams for the 3 dithiocarbamate ligands. M.O. contour
value = 0.02 au
26

HOMO
LUMO
1
[Pt(L )2]
2
[Pt(L )2]
3
[Pt(L )2]
1
[Pd(L )2]
2
[Pd(L )2]
3
[Pd(L )2]
27

1
[Ni(L )2]
2
[Ni(L )2]
3
[Ni(L )2]
Fig. 7: Frontier molecular orbital diagrams for the platinum, palladium and nickel complexes
of the dithiocarbamate ligands. M.O. contour value = 0.02 au
28

Fig. 8: 3-D isosurface plots of relevant non-covalent interactions in the palladium
dithiocarbamate complex. Isovalue = 0.03 au
29