Synthesis and characterization of copper(II) dithiocarbamate complexes involving pyrrole and ferrocenyl moieties and their utility for sensing anions and preparation of copper sulfide and copper–iron sulfide nanoparticles

Article abstract of DOI:10.1002/aoc.4363

Bis(N-(pyrrol-2-ylmethyl)-N-butyldithiocarbamato-S,S′)copper(II) (1), bis(N-(pyrrol-2-ylmethyl)-N-(2-phenylethyl)dithiocarbamato-S,S′)copper(II) (2), bis(N-methylferrocenyl-N-(2-phenylethyl)dithiocarbamato-S,S′)copper(II) (3) and bis(N-furfuryl-N-methylfe

Full text of DOI:10.1002/aoc.4363

Received: 3 November 2017
DOI: 10.1002/aoc.4363
Revised: 10 February 2018
Accepted: 22 February 2018
F U L L P A P E R
Synthesis and characterization of copper(II)
dithiocarbamate complexes involving pyrrole and
ferrocenyl moieties and their utility for sensing anions and
preparation of copper sulfide and copper–iron sulfide
nanoparticles
Govindasamy Gurumoorthy1 | Subbiah Thirumaran1
| Samuele Ciattini2
1 Department of Chemistry, Annamalai
Bis(N‐(pyrrol‐2‐ylmethyl)‐N‐butyldithiocarbamato‐S,S′)copper(II) (1), bis(N‐
(pyrrol‐2‐ylmethyl)‐N‐(2‐phenylethyl)dithiocarbamato‐S,S′)copper(II) (2),
University, Annamalainagar, –608 002,
India
2 Centro di Cristallografia Strutturale, Polo
Scientifio di Sesto Fiorentino, Via della
Lastruccia 3, 50019 Sesto Fiorentino,
Florence, Italy
bis(N‐methylferrocenyl‐N‐(2‐phenylethyl)dithiocarbamato‐S,S′)copper(II) (3)
and bis(N‐furfuryl‐N‐methylferrocenyldithiocarbamato‐S,S′)copper(II) (4) were
prepared and characterized using elemental analysis and infrared and UV–
visible spectroscopies. X‐ray diffraction (XRD) studies on 3 show that each
copper centre adopts the square planar geometry by the coordination of four
sulfur atoms of the metalloligand N‐methylferrocenyl‐N‐(2‐phenylethyl)
dithiocarbamate. The Cu―S distances are symmetrical and are in the range
2.293–2.305 Å. The supramolecular architecture in complex 3 is sustained
in the solid state by C―H⋅⋅⋅π, C―H⋅⋅⋅S, Fe⋅⋅⋅Fe and H⋅⋅⋅H interactions. Den-
sity functional theory calculations were carried out for 3. Anion (F−, Cl−, Br−
and I−) binding studies with complex 1 were performed using cyclic volt-
ammetry. Copper sulfide, copper–iron sulfide‐1 and copper–iron sulfide‐2
nanoparticles were prepared from complexes 2, 3 and 4, respectively, and
they were characterized using powder XRD, transmission electron micros-
copy (TEM) and energy‐dispersive X‐ray, UV–visible, photoluminescence
and infrared spectroscopies. TEM images of copper–iron sulfide‐1 and cop-
per–iron sulfide‐2 reveal that the particles are spherical and oval shaped,
respectively. Photocatalytic activities of as‐prepared nanoparticles were stud-
ied by decolourization of methylene blue and rhodamine‐B under UV light.
It was found that copper–iron sulfide degrades methylene blue and rhoda-
mine‐B much better than does copper sulfide.
Correspondence
Subbiah Thirumaran, Department of
Chemistry, Annamalai University,
Annamalainagar – 608 002, India.
Email: sthirumaran@yahoo.com
Funding information
University Grants Commission (UGC),
India, Grant/Award Number: F. No. 42‐
341/2013 (SR)
KEYWORDS
anion sensing, copper sulfide, copper(II) dithiocarbamates, copper–iron sulfide, photodegradation
Appl Organometal Chem. 2018;e4363.
wileyonlinelibrary.com/journal/aoc
Copyright © 2018 John Wiley & Sons, Ltd.
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https://doi.org/10.1002/aoc.4363

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GURUMOORTHY ET AL.
diffraction (XRD) was performed using an EQUINOX
1000. Transmission electron microscopy (TEM) images
were recorded using a TECNAI T2 G2 (FEI). Energy‐
dispersive X‐ray spectroscopy (EDS) was performed with
a SUPRA 55VP CARL. Photoluminescence spectra were
recorded using a PerkinElmer 1555 fluorescence spectro-
photometer at room temperature.
1 | INTRODUCTION
Dithiocarbamate ligands have found ample use in coordi-
nation chemistry.[1,2] Their wide range of applications
such as in industry, agriculture and medicine have gener-
ated a large collection of crystallographic data for their
metal complexes.[3–6] Dithiocarbamates are versatile
ligands capable of stabilizing transition metals in both
high and low oxidation states,[7] and complexes of
Cu(I), Cu(II) and Cu(III) are all known, being intercon-
vertible via reversible one‐electron redox process.[8] Elec-
trochemical studies of these complexes reveal that they
can sense various anions and cations.[9] Modifications in
the R groups attached to nitrogen atom of the dithiocar-
bamate fragment can have an effect on both the structure
and chemical properties, mainly due to a change in the
2.2 | X‐ray Crystallography
Diffraction data were recorded with a SuperNova, Dual,
Cu at zero, Pilatus 200/300 K diffractometer using graph-
ite‐monochromated Cu Kα radiation (λ = 1.54184 Å) at
ambient temperature. The structure was solved and
refined by direct method using SHELXL‐2014/7.[24] All
non‐hydrogen atoms were refined anistropically and the
hydrogen atoms were refined isotropically. Details of the
crystal data and structure refinement parameters for 3
are summarized in Table 1. Selected bond lengths and
angles for 3 are given in Table 2.
acid–base nature of the dithiocarbamate fragment.[10]
A
number of copper dithiocarbamate complexes have been
shown to be efficient catalysts for the atom‐transfer radi-
cal polymerization and reverse atom‐transfer radical
polymerization of methyl methacrylate.[11,12] Recently,
copper(II) complexes containing ferrocenyl‐based
dithiocarbamates have been exploited as sensitizers in
dye‐sensitized TiO2 solar cells for converting sunlight
into electrical energy.[13] In recent years, copper dithio-
carbamate complexes have been used as single‐source
precursors for the preparation of copper sulfides[14–18]
and in ferroelectric materials,[19] electrochemical sen-
sors,[20] photoconductivity,[21] a wealth of metal‐centred
electrochemistry[2,22] and interesting photochemistry.[7,23]
Metal sulfides can also serve as important semiconductor
photocatalysts which offer the potential for complete
elimination of toxic chemicals.
Herein we report the synthesis, characterization and
anion‐sensing properties of copper(II) dithiocarbamate
complexes containing pyrrole and ferrocene moieties. In
addition, the synthesis, characterization and photocata-
lytic activities of copper sulfide and copper–iron sulfide
nanoparticles, which were synthesized from the as‐
prepared complexes, are also presented.
2.3 | Theoretical Calculations
The highest occupied molecular orbital (HOMO)–lowest
unoccupied molecular orbital (LUMO) energy was calcu-
lated with the Gaussian 03 software package using gradi-
ent‐corrected density functional theory (DFT) with the
B3LYP method.[25] The LANL2DZ basis set was used by
including effective core potential functions.
2.4 | Cyclic Voltammetry
Cyclic voltammetry was performed using a conventional
three‐electrode system. A glassy carbon electrode was
used as a working electrode. The counter electrode was
a platinum wire and the reference electrode was
Ag/AgCl. HPLC‐grade acetonitrile was used as the solvent
and tetrabutylammonium fluoroborate (0.01 M) as the
supporting electrolyte. The scan rate was 100 mV s−1
.
Complex 3 was investigated at 25 °C in an oxygen‐free
atmosphere, provided by bubbling purified nitrogen
through the solution. The concentration of the compounds
was 1 × 10−4 M in electrochemical solutions of (n‐Bu4N)I,
(n‐Bu4N)Br, (n‐Bu4N)Cl and (n‐Bu4N)F (5 × 10−4 M) in
acetonitrile. Cyclic voltammograms were recorded with a
CHI604C Electrochemical Analyser.
2 | EXPERIMENTAL
2.1 | Materials and Techniques
Reagent‐grade chemicals were purchased from commer-
cial sources and used as such. Infrared (IR) spectra were
recorded with a Thermo Nicolet Avatar 330 FT‐IR spec-
trophotometer (range: 400–4000 cm−1) as KBr pellets. A
Shimadzu UV‐1650 PC double‐beam UV–visible spectro-
photometer was used for recording the electronic spectra.
The spectra of complexes were recorded in CHCl3 and the
pure solvent was used as the reference. Powder X‐ray
2.5 | Photocatalytic Experiments
The photocatalytic activity of copper sulfide and copper–
iron sulfide was evaluated by degradation of aqueous
solutions of methylene blue and rhodamine‐B. All the
solutions were prepared using double‐distilled water.

GURUMOORTHY ET AL.
3 of 12
TABLE 1 Crystal data, data collection and refinement parameters for complex 3
Empirical formula
C40H40CuFe2N2S4
Formula weight
Crystal dimensions (mm3)
852.22
0.30 × 0.20 × x 0.20
Monoclinic
P21/n
Crystal system
Space group
a (Å)
12.5771(3)
23.3141(7)
12.9515(3)
90
b (Å)
c (Å)
α (°)
β (°)
108.725(2)
90
γ (°)
V (Å3)
3596.68(17)
4
Z
Dc (g cm−3
)
1.574
μ (cm−1
F(000)
λ (Å)
)
9.489
1756
Cu Kα (1.54184)
3.792 to 75.061
θ range (°)
Index ranges
−15 ≤ h ≤ 15, −28 ≤ k ≤ 26, −16 ≤ l ≤ 16
Reflections collected
Observed reflections [I > 2σ(I)]
Weighting scheme
R[F2 > 2σ(F2)], wR(F2)
Number of parameters refined
GOF
70 965
6452
Calc. W = 1/[σ2(Fo2) + (0.0420P)2 + 1.9063P], where p = (Fo + 2Fc )/3
2
2
0.0319, 0.0845
442
1.09
TABLE 2 Selected bond lengths (Å) and bond angles (°) of 3
Bond length
XRD
DFT/LanL2DZ
Bond angle
XRD
DFT/LanL2DZ
Cu–S4
Cu–S2
Cu–S3
Cu–S1
S1–C1
S3–C21
S2–C1
S4–C21
N1–C1
N2–C21
2.2937(6)
2.2986(6)
2.3023(6)
2.3052(6)
1.727(2)
1.723(2)
1.719(2)
1.729(2)
1.315(3)
1.316(3)
2.431
2.492
2.430
2.432
1.790
1.792
1.792
1.705
1.344
1.344
S4–Cu–S2
S4–Cu–S3
S2–Cu–S3
S4–Cu–S1
S2–Cu–S1
S3–Cu–S1
N1–C1–S2
N1–C1–S1
N2–C21–S3
N2–C21–S4
102.03(2)
77.75(2)
103.49
76.50
170.06(3)
176.90(3)
77.56(2)
179.93
179.93
76.49
103.19(2)
123.57(16)
122.86(16)
124.49 (16)
122.14 (16)
103.50
123.18
122.53
123.21
122.49
For a typical photocatalytic experiment, 0.1 g of catalyst
was added to 50 ml of an aqueous solution of rhoda-
mine‐B at a concentration of 1.0 × 10−4 M. The solution
was maintained in the dark for 30 min to reach dye
solution adsorption–desorption equilibrium. The solution
with the suspended nano‐photocatalyst was irradiated by
UV light from a mercury vapour lamp. At given time
intervals, 3 ml aliquots were withdrawn and centrifuged

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GURUMOORTHY ET AL.
to remove catalyst. The concentration of dye solution was
determined using a UV–visible spectrophotometer.
was added to the dithiocarbamic acid solution resulting
in the formation of a brown precipitate. The precipitate
obtained was filtered, washed several times with cold
water and then air‐dried (Scheme 1). Yield: 78%, m.p.
149–151°C. IR (KBr, ν, cm−1): 3381 (νN―H), 1494
(νC―N), 1029 (νC―S). UV–visible (CHCl3, λ, nm): 634,
439, 275, 244: Anal. Calcd for C20H23CuN4S4 (%): C,
46.35; H, 5.83; N, 10.81; found (%): C, 46.18; H, 5.77; N,
10.75.
3 | PREPARATION OF COMPLEXES
3.1 | Preparation of Amines
N‐(Pyrrol‐2‐ylmethyl)‐N‐butylamine,
ylmethyl)‐N‐(2‐phenylethyl)amine, N‐methylferrocenyl‐
N‐(2‐phenylethyl)amine and N‐furfuryl‐N‐
N‐(pyrrol‐2‐
methylferrocenylamine were prepared by general
methods reported earlier.[26]
3.3 | Preparation of Bis(N‐(pyrrol‐2‐
ylmethyl)‐N‐(2‐phenylethyl)
dithiocarbamato‐S,S′)copper(II) (2)
3.2 | Preparation of Bis(N‐(pyrrol‐2‐
ylmethyl)‐N‐butyldithiocarbamato‐S,S′)
copper(II) (1)
A method similar to that described for the synthesis of 1
was adopted; however, N‐(pyrrol‐2‐ylmethyl)‐N‐(2‐
phenylethyl)amine was used instead of N‐(pyrrol‐2‐
ylmethyl)‐N‐butylamine (Scheme 1). Colour: brown;
yield: 80%; m.p. 153–155°C. IR (KBr, ν, cm−1): 3388
(νN―H), 1490 (νC―N), 1025 (νC―S). UV–visible (CHCl3, λ,
nm): 639, 438, 275, 237: Anal. Calcd for C28H30CuN4S4
(%): C, 54.74; H, 4.92; N, 9.12; found (%): C, 54.60; H,
4.89; N, 9.06.
N‐(Pyrrol‐2‐ylmethyl)‐N‐butylamine (4.0 mmol, 0.9 ml) in
ethanol was mixed with carbon disulfide (4.0 mmol,
0.3 ml) under ice‐cold conditions (5°C). The solution
was stirred for 30 min. This produced the N‐(pyrrol‐2‐
ylmethyl)‐N‐butyldithiocarbamic acid solution. An aque-
ous solution of Cu(CH3COO)2⋅H2O (2.0 mmol, 0.3993 g)
SCHEME 1 Preparation of complexes
1, 2, 3 and 4

GURUMOORTHY ET AL.
5 of 12
this region is characteristic of bidentate chelation of
dithiocarbamate ligands. The IR spectra of complexes 1–
4 exhibit a distinct vibrational band at around 1020 cm
3.4 | Preparation of Bis(N‐
methylferrocenyl‐N‐(2‐phenylethyl)
dithiocarbamato‐S,S′)copper(II) (3)
−1
which is associated with the bidentate νC―S vibration
A method similar to that described for the synthesis of 1
of the dithiocarbamate ligands. The spectral band
observed in the region 1450–1550 cm−1 is associated with
νC―N vibrational mode. The position of this band in IR
spectra in general is found between values for single
was
adopted;
however,
N‐methylferrocenyl‐N‐(2‐
phenylethyl)amine used instead of N‐(pyrrol‐2‐ylmethyl)‐
N‐butylamine (Scheme 1). Colour: brown; yield: 76%; m.
p. 165–167°C. IR (KBr, ν, cm−1): 1490 (νC―N), 1025
(νC―S). UV–visible (CHCl3, λ, nm): 630, 428, 271, 237:
Anal. Calcd for C40H40CuFe2N2S4 (%): C, 56.38; H, 4.74;
N, 3.28; found (%): C, 56.16; H, 4.73; N, 3.26.
(1250–1350 cm−1
)
and double (1640–1690 cm−1
)
bonds.[29,30] For complexes 1–4, the νC―N (thioureide)
band appeared in the range 1487–1494 cm−1, indicating
the partial double bond character of C―N bond.[31] While
the calculated values of νC‐S and νC―N vibrations are
observed at 1488 and 1025 cm−1 for 3. Additionally, the
band observed at around 3380 cm−1 for complexes 1 and
2 is indicative of the free NH function in the pyrrole moiety.
3.5 | Preparation of Bis(N‐furfuryl‐N‐
methylferrocenyldithiocarbamato‐S,S′)
copper(II) (4)
A method similar to that described for the synthesis of 1
4.2 | Electronic Spectral Studies
was
adopted;
however,
N‐furfuryl‐N‐
methylferrocenylamine was used instead of N‐(pyrrol‐2‐
ylmethyl)‐N‐butylamine (Scheme 1). Synthesis and
characterization of this complex have been reported
earlier.[13] To study its utility for the preparation of
bimetallic sulfide nanoparticles, complex 4 was prepared
using a slightly modified procedure. Colour: brown; yield:
78%; m.p. 170–172°C. IR (KBr, ν, cm−1): 1487 (νC―N),
1017 (νC―S). UV–visible (CHCl3, λ, nm): 632, 437, 275,
228: Anal. Calcd for C34H36CuFe2N2O2S4 (%): C, 50.79;
H, 4.01; N, 3.48; found (%): C, 50.62; H, 4.01; N, 3.45.
Electronic spectra of 1–4 are shown in Figures S6–S9. The
electronic absorption spectra of the complexes are charac-
terized by three absorption bands associated with the
dithiocarbamate fragment. Two peaks observed in the
range 225–275 nm are assigned to the π → π* intraligand
transitions of dithiocarbamate.[32] The broad absorption
bands of medium intensity with λmax at around 435 nm
in the spectra are associated with charge transfer transi-
tions of the type ligand → metal and metal → ligand
between Cu(II) ion and dithiocarbamate ligand.[33,34] The
weak and broad band with appearance of a shoulder in
the region around 630–640 nm belongs to d–d transitions
of Cu(II) (2B1g → 2Eg) in a square planar geometry.[35]
3.6 | Preparation of Copper and Copper–
Iron Sulfide
An amount of 0.5 g of 2 was mixed with 15 ml of
triethylenetetraamine in a round‐bottom flask and then
the contents of the flask were refluxed for 15 min. The
black precipitate obtained was filtered off and washed
with methanol.
A similar procedure was adopted for the preparation
of copper–iron sulfide‐1 and copper–iron sulfide‐2 from
complexes 3 and 4.
4.3 | Single‐Crystal X‐ray Structural
Analysis
Generally Cu(II) dithiocarbamate complexes display
three types of structural motifs: they are (i) square planar
monomeric,[36,37] (ii) five‐coordinate dimeric[36,38] and
(iii) infinite polymeric chain.[39] Complex 3 was obtained
as a monomeric species. XRD studies of 3 show that each
copper atom is surrounded by four sulfur atoms from two
chelating N‐methylferrocenyl‐N‐(2‐phenylethyl)dithiocar-
bamate ligands with isobidentate coordination mode
(Figure 1). The Cu―S1, Cu―S2, Cu―S3 and Cu―S4 dis-
tances are approximately symmetric with values of 2.305,
2.298, 2.302 and 2.293 Å, respectively. The copper centre
of 3 has a distorted square planar four‐coordinate
environment.[14,40–42] The distortion is due to small bite
4 | RESULTS AND DISCUSSION
4.1 | IR Spectral Studies
IR spectra of 1–4 are shown in Figures S1–S4 and the the-
oretical IR spectrum of 3 is displayed in Figure S5. The
νC―S band in the range 950–1050 cm−1 is used to deter-
mine the mode of coordination of the dithiocarbamate
ligand to the metal centre.[27,28] The splitting of the
angles
of
the
chelating
dithiocarbamates
(S1―Cu―S2 = 77.56(2)° and S3―Cu―S4 = 77.75(2)°)
which results in secondary deviations of the other angles
from ideal square planar angles. The delocalization of π‐
ν
C―S band is characteristic of a monodentate dithiocarba-
mate ligand whereas the appearance of a single band in

6 of 12
GURUMOORTHY ET AL.
formation
of
two
intermolecular
C―H⋅⋅⋅S
(C3―H3A⋅⋅⋅S4 = 2.795 Å and C23―H23A⋅⋅⋅S1 = 2.835 Å)
bonds giving rise to a one‐dimensional self‐arrangement
(Figure S10). A C39―H39⋅⋅⋅π interaction occurs between
H39 and a centroid of a cyclopentadienyl ring (C16–20).
The parameters associated with this interaction are a
H39–ring centroid distance of 3.025 Å and angle of
159.89° at H (Figure S11). The H7 atom forms an inter-
molecular C7―H7⋅⋅⋅S2 interaction with S2 which results
in a polymeric chain structure (Figure S12). The ferro-
cene groups are linked as layers. The distances between
two adjacent iron atoms are in the range 6.297–6.655 Å.
In this layer, each ferrocene group is surrounded by four
other ferrocene groups which are perpendicular to the
central ferrocene group. The cyclopentadienyl planes in
each ferrocene moiety are nearly parallel and the confor-
mations nearly eclipsed. In this layer, the distance between
H15 and H39 is 2.478 Å, a distance which shows the forma-
tion of rare closed‐shell hydrogen–hydrogen interactions
in 3 shorter than the sum of van der Waals radii for the
hydrogen atoms[43] (Figure 3). In complex 3, four intramo-
lecular C―H⋅⋅⋅S interactions occur between the hydrogen
atoms of CH2 and sulfur atoms (Figure S13).
FIGURE 1 Molecular structure of complex 3
electron density over the S2CN moiety is obvious from the
shortening of the C―S (mean: 1.724 Å) and C―N (1.315
and 1.316 Å) bond lengths considering the typical C―N
and C―S single bonds.[2,14,40–42]
In complex 3, molecules pack through C14―H14⋅⋅⋅π
(2.766 Å) interaction between the monomers making up
the polymeric chain (Figure 2; Table 3). There is also
5 | DFT STUDIES
5.1 | Molecular Geometry and Structural
Properties
Figure S14 exhibits the optimized structure of complex 3
obtained from the B3LYP/LanL2DZ method. The selected
optimized geometrical parameters of 3 from DFT method
calculations and the XRD values are listed in Table 2. For
3, each copper atom is surrounded by four sulfur atoms
from
two
chelating
N‐methylferrocenyl‐N‐(2‐
phenylethyl)dithiocarbamate ligands with bidentate coor-
FIGURE 2 Intermolecular C―H⋅⋅⋅π interaction in 3
dination mode. The Cu―S1, Cu―S2, Cu―S3 and Cu―S4
TABLE 3 Geometric details of hydrogen bonding (Å, °) in 3
Interactions
D―H (Å)
H···A (Å)
D···A (Å)
D―H···A (°)
C2―H2B···S1a
0.97
0.97
0.97
0.97
0.97
0.93
0.93
0.93
2.630
2.703
2.560
2.665
2.795
2.921
2.766
3.025
3.021 (2)
3.053 (2)
3.083 (2)
3.019 (2)
3.753 (2)
3.776 (3)
3.671
104.41
101.81
113.85
102.01
169.06
153.52
164.79
159.89
C10―H10A···S2a
C30―H30B···S3a
C22―H22B···S4a
C3―H3A···S4b
C7―H7···S2b
C14―H14···Cg(C24‐C29)c
C39―H39···Cg(C16‐C20)c
3.912
aIntramolecular C―H···S hydrogen bonding.
bIntermolecular C―H···S interactions.
cIntermolecular C―H···π interactions.

GURUMOORTHY ET AL.
7 of 12
FIGURE 3 Layer formed via Fe⋅⋅⋅Fe interactions and H⋅⋅⋅H
interactions in 3
distances are approximately symmetric with values of
2.431, 2.492, 2.430 and 2.432 Å, respectively. A compari-
son of bond parameters (Table 2) of experimental and cal-
culated values reveals that the differences between
experimental and calculated bond lengths are in the
range 0.03–0.07 Å except Cu―S bond distances (0.13–
0.19 Å). Similarly the bond angles are almost equal for
both methods (0.35–1.5°) except S2―Cu―S3 (9.87°) and
S4―Cu―S1 (3.03°). These differences between bond
parameters arise due to the calculated geometric parame-
ters using the DFT method considering the gas phase
only and that the molecule is free of interactions.
FIGURE 4 Graphical images of (a) HOMO and (b) LUMO of
complex 3
5.2 | Frontier Molecular Orbital Analysis
The HOMO and LUMO are the most important orbitals
in a molecule. HOMO is mainly located over four sulfur
atoms and the LUMO of π nature is delocalized over both
the MS2CN rings (Figure 4). The hardness (η), chemical
potential (μ) and electronegativity (χ) are calculated using
HOMO and LUMO energies. For complex 3, η =
2.2611 eV, μ = −3.4583 eV and χ = 3.4583 eV.[44] It is
seen that the chemical potential of 3 is negative. This
indicates that the complex is stable. The hardness sig-
nifies the resistance towards the deformation of electron
clouds of chemical systems under small perturbations
encountered during chemical processes.
< green < blue (positive potential). The blue colour indi-
cates the strongest attraction of electrons and red indi-
cates the strongest repulsion of electrons.[45,46] The
molecular electrostatic potential map of 3 (Figure S15)
shows that regions having negative potential are over
the four S atoms and the regions having the positive
potential are over the two N atoms. This supports the
presence of canonical structure R1R2N+―CS2
.
2−
6 | ANION BINDING STUDIES
The recognition ability of 1 towards halide ions in the
form of their corresponding tetrabutylammonium salts
was investigated using cyclic voltammetry in CH3CN
5.3 | Molecular Electrostatic Potential
The various values of the electrostatic potential at the
molecular electrostatic potential surfaces are represented
by different colours. Potential increases in the order red
(negative potential) < orange < yellow (zero potential)
solution containing 0.01
M
tetrabutylammonium
fluoroborate as a supporting electrolyte. Complex 1 shows
a one‐electron quasi‐reversible Cu(II) to Cu(I) reductive
response (Figure 5). Similar observations for copper(II)

8 of 12
GURUMOORTHY ET AL.
(222) planes (JCPDS card no, 73–1667). In the case of cop-
per–iron sulfide‐2, the diffraction peaks at 2θ = 29.45°,
48.67° and 58.0° are assigned to the (111), (220) and
(311) planes of cubic structure (JCPDS card no.
81–1378). The XRD patterns of copper sulfide and cop-
per–iron sulfide‐1 are of poor quality being indicative of
low crystallinity and the presence of other phases.
The morphologies of the as‐prepared metal sulfide
nanoparticles were investigated by TEM analysis. TEM
images of copper sulfide and copper–iron sulfides are
shown in Figure 6. The TEM images of copper sulfide
particles obtained from complex 2 show that they are
spherical with diameters in the range 10–18 nm
(Figure 6a,b). The TEM image of copper–iron sulfide‐1
reveals that the shapes of the particles are spherical
(diameters of 11–20 nm; Figure 6c). The shape of the cop-
per–iron sulfide‐2 nanoparticles is oval morphology
(Figure 6d,e).
The elemental composition of copper sulfide and cop-
per–iron sulfide nanoparticles was investigated. EDS
spectra of products obtained from complexes 2, 3 and 4
are shown in Figure S17. The EDS trace of the product
obtained from solvothermal decomposition of complex 2
reveals the formation of copper sulfide. The Cu:S atomic
ratio (9:5.6) shows that there are some vacancies of
Cu2+ ions or some sulfur dangling bonds are present in
the sample. EDS analysis of copper–iron sulfide‐1 and ‐2
indicates the presence of two metals (Cu and Fe) and sul-
fur. The Cu:Fe:S ratio of copper–iron sulfide‐1 and ‐2 are
3.5:1.0:3.7 and 1.0:2.0:2.7, respectively.
FIGURE 5 Changes in cyclic voltammogram of complex 1 on the
addition of anions (TBA, tetrabutylammonium)
dithiocarbamates were reported by Gupta et al.[35]
Complex 1 exhibits a marked electrochemical change to
F − (ΔE1/2 = 0.2689 V) followed by I− (ΔE1/2 = 0.2011 V)
(Table 4). In contrast, addition of the other anions to a
CH3CN solution of complex 1 does not induce any
significant change in the signals of cyclic voltammogram
of complex 1.
7 | CHARACTERIZATION OF
COPPER SULFIDE AND COPPER–
IRON SULFIDE NANOPARTICLES
Copper sulfide and copper–iron sulfide nanoparticles
obtained from complexes 2, 3 and 4 are represented as
copper sulfide, copper–iron sulfide‐1 and copper–iron sul-
fide‐2, respectively. The XRD patterns of copper sulfide
and copper–iron sulfides are shown in Figure S16. The
diffraction pattern of the dominant phase in copper sul-
fide is cubic Cu9S5 with major diffraction peaks of
(0015), (1010), (0114), (110), (0027), (1112) and (2017)
planes (JCPDS card no. 47–1748). The diffraction patterns
of copper–iron sulfide‐1 are indexed to the cubic phase of
Cu5FeS4 with characteristic (111), (200), (220), (331) and
The expected Cu:S and Cu:Fe:S ratios for copper
sulfide (Cu9S5) and copper–iron sulfide −1 (Cu5FeS4)
from powder XRD studies are 9:5 and 5:1:4, respectively.
But the ratios observed from EDS are different. This sup-
ports the presence of various phases. In the case of cop-
per–iron sulfide‐2 (CuFe2S3) the ratios of Cu:Fe:S
determined from powder XRD (1:2:3) and EDS (1:2:2.7)
are almost the same.
Figure S18 illustrates the UV–visible absorption
spectra of copper sulfide and copper–iron sulfide
nanoparticles dispersed in distilled ethanol at room
temperature. A broad band appears at 272 nm (4.56 eV)
for copper sulfide whereas a band at 270 nm (4.59 eV)
and a weak band at 260 nm (4.77 eV) are observed for
copper–iron sulfide‐1 and ‐2, respectively. Compared with
bulk copper sulfide (1033 nm, 1.2 eV),[47] the absorption
maxima of copper sulfide and copper–iron sulfide
nanoparticles exhibit a large blue shift, which is attributed
to the quantum confinement of charge carriers in the
nanoparticles.
TABLE 4 Electrochemical anion recognition data of complex 1
Compounda
Epc (V)
ΔE1/2 (V)b
Complex
−1.5024
−1.2331
−1.4485
−1.4635
−1.7035
—
TBA fluoride
TBA chloride
TBA bromide
TBA iodide
0.2693
0.0539
0.0389
0.2011
Figures S19 and S20 show the photoluminescence
spectra recorded at room temperature for copper sulfide
and copper–iron sulfide‐2 nanoparticles. The emission
aTBA, tetrabutylammonium.
bShift of Cu(II)/Cu(I) reduction potential produced by anions.

GURUMOORTHY ET AL.
9 of 12
FIGURE 6 TEM images of copper sulfide, scale bars (a) 100 nm and (b) 20 nm; copper–iron sulfide‐1 (c); and copper–iron sulfide‐2, scale
bars (d) 100 nm and (e) 20 nm
spectra of both types of nanoparticles exhibit a peak at
around 400 nm corresponding to the band‐edge emission.
Along with this emission another red‐shifted intense
peak is observed at 432 nm in the spectrum of
copper–iron sulfide‐2. The red‐shifted emission results
from trap‐related electron–hole recombination.
IR spectra of copper sulfide and copper–iron sulfide
nanoparticles (Figures S21–S23) exhibit three bands in
the region 2850–2964 cm−1 due to aliphatic νC―H. The
bands in the region 3415–3485 cm−1 are assigned to
N―H stretching vibrations. These data indicate the
presence of triethylenetetraamine capping agent in

10 of 12
GURUMOORTHY ET AL.
copper sulfide and copper–iron sulfide‐1 and ‐2. The lack
of aromatic C―H and N―CS2 stretching vibrations
indicates the absence of dithiocarbamate ligands in the
as‐synthesized copper–iron sulfide nanoparticles.
7.1 | Photocatalytic Activity of Copper
Sulfide and Copper–Iron Sulfide
The photocatalytic activity performances of as‐prepared
copper sulfide and copper–iron sulfide‐1 were evaluated
by photocatalytic degradation of methylene blue and rho-
damine‐B aqueous solutions. The degradation of methy-
lene blue and rhodamine‐B was carried out using UV
irradiation as followed by spectrophotometric monitoring.
Figure 7 shows the temporal evolution of the absorption
spectra during the photocatalytic degradation of methy-
lene blue and rhodamine‐B in the presence of copper sul-
fide and copper–iron sulfide‐1. As the irradiation time
increased the absorption peaks decreased. Figure 8 shows
the photodegradation efficiency of methylene blue and
rhodamine‐B as a function of irradiation time. C is the
absorption of methylene blue and rhodamine‐B at 662
and 554 nm, respectively, at time t and C0 is the absorption
of methylene blue and rhodamine‐B before irradiation.
The experiments showed the good photocatalytic
activity of copper sulfide and copper–iron sulfide‐1 for
the degradation of methylene blue and rhodamine‐B
under UV irradiation. It was observed that copper sulfide
degraded 89% of methylene blue in 180 min while cop-
per–iron sulfide‐1 degraded 93% of methylene blue in
180 min under UV light. It was also observed that copper
sulfide and copper–iron sulfide degraded 82 and 86% of
FIGURE 7 Absorption spectral changes of methylene blue and
rhodamine‐B using (a, b) copper sulfide and (c, d) copper–iron
sulfide‐1 under UV light
FIGURE 8 Photocatalytic degradation of methylene blue (M.B)
and rhodamine‐B (R.B) using copper sulfide and copper–iron
sulfide‐1

GURUMOORTHY ET AL.
11 of 12
[6] R. Singh, N. K. Kaushik, Spectrochim. Acta Mol. Biomol.
Spectrosc. 2008, 71, 669.
rhodamine‐B, respectively, under UV light irradiation in
180 min. The degradation efficiency of copper–iron
sulfide is greater than that of copper sulfide for the inves-
tigated dyes. The presence of iron in copper–iron sulfide‐
1 can also enhance the photocatalytic degradation activity
due to smaller crystal size, higher efficiency for electron–
hole regeneration and charge trapping.
[7] A. M. Bond, R. L. Martin, Coord. Chem. Rev. 1984, 54, 23.
[8] A. R. Hendrickson, R. L. Martin, N. M. Rohde, Inorg. Chem.
1976, 15, 2115.
[9] J. Cookson, P. D. Beer, Dalton Trans. 2007, 1459.
[10] G. Gurumoorthy, S. Thirumaran, Phosphorus, Sulfur Silicon
Relat. Elem. 2017, 192, 330.
8 | CONCLUSIONS
[11] P. Li, K.‐Y. Qui, J. Polym Sci, Polym. Chem. 2002, 40, 2093.
[12] P. Li, K.‐Y. Qui, Macromolecules 2002, 35, 8906.
Complexes 1–4 containing pyrrole and ferrocene moieties
were prepared and characterized. The copper centre in 3
has distorted square planar four‐coordinate environ-
ments. The chemical potential of 3 calculated using
DFT is negative (−3.4583 eV). This indicates that the
complex is stable. Anion binding studies show that com-
plex 1 prefers to bind with F−. This study demonstrates
that spherical copper sulfide and copper–iron sulfide
nanoparticles and oval‐shaped copper–iron sulfide nano-
particles can be prepared from the copper dithiocarbamate
complexes. Copper–iron sulfide (bimetallic) nanoparticles
revealed better photocatalytic activity for the
photodegradation of methylene blue and rhodamine‐B
under UV light compared to copper sulfide (monometal-
lic). This study indicates that the new copper(II) dithiocar-
bamate complexes containing various N‐bound organic
moieties are useful for sensing anions and for preparing
effective photocatalysts (copper sulfide and copper–iron
sulfide nanoparticles with various shapes).
[13] V. Singh, R. Chauhan, A. N. Gupta, V. Kumar, M. G. B. Drew,
L. Bahadur, N. Singh, Dalton Trans. 2014, 43, 4752.
[14] P. D. Beer, N. Berry, M. G. B. Drew, O. D. Fox, M. E. Padilla‐
Tosta, S. Patell, Chem. Commun. 2001, 199.
[15] R. Nomura, K. Miyakawa, T. Toyosaki, H. Matsuda, Chem.
Vap. Deposition 1996, 2, 174.
[16] M. Kemmler, M. Lazell, P. O'Brien, D. J. Otway, J.‐H. Park, J.
R. Walsh, J. Mater. Sci. Mater. Electron 2002, 13, 531.
[17] P. O'Brien, J. R. Walsh, I. M. Watson, L. Hart, S. R. P. Silva,
J. Cryst. Growth 1996, 167, 133.
[18] M. R. Lazell, P. O'Brien, D. J. Otway, J.‐H. Park, Chem. Mater.
1999, 11, 3430.
[19] A. L. Abdelhady, K. Ramasamy, M. A. Malik, P. O'Brien, S. J.
Haigh, J. Raftery, J. Mater. Chem 2011, 21, 17888.
[20] T. Okubo, R. Kawajiri, T. Mitani, T. Shimoda, J. Am. Chem.
Soc. 2005, 127, 17598.
[21] R. H. Bube, W. H. McCarroll, J. Phys. Chem. Solids 1959, 10,
333.
[22] D. V. Konarev, A. Y. Kovalevsky, D. V. Lopatin, A. V.
Umrikhin, E. I. Yudanova, P. Coppens, R. N. Lyubovskaya, G.
Saito, Dalton Trans. 2005, 1821.
ACKNOWLEDGEMENTS
S.T. is grateful to the University Grants Commission
(UGC), India (no. 42‐341/2013 (SR)) for providing funding
for this research study. We are grateful to SAIF, Panjab
University, Chandigarh, India, for recording TEM images.
[23] V. F. Plyusnin, A. V. Kolomeets, V. P. Grivin, S. V. Larionov, H.
Lemmetyinen, J. Phys. Chem. A 2011, 115, 1763.
[24] G. M. Sheldrick, Acta Crystallogr. C 2015, 71, 3.
[25] M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria, M. A.
Robb, J. R. Cheeseman, J. A. Montgomery, Jr., T. Vreven, K. N.
Kudin, J. C. Burant, J. M. Millam, S. S. Iyengar, J. Tomasi, V.
Barone, B. Mennucci, M. Cossi, G. Scalmani, N. Rega, G. A.
Petersson, H. Nakatsuji, M. Hada, M. Ehara, K. Toyota, R.
Fukuda, J. Hasegawa, M. Ishida, T. Nakajima, Y. Honda, O.
Kitao, H. Nakai, M. Klene, X. Li, J. E. Knox, H. P. Hratchian,
J. B. Cross, C. Adamo, J. Jaramillo, R. Gomperts, R. E. Stratmann,
O. Yazyev, A. J. Austin, R. Cammi, C. Pomelli, J. W. Ochterski,
P. Y. Ayala, K. Morokuma, G. A. Voth, P. Salvador, J. J.
Dannenberg, V. G. Zakrzewski, S. Dapprich, A. D. Daniels,
M. C. Strain, O. Farkas, D. K. Malick, A. D. Rabuck, K.
Raghavachari, J. B. Foresman, J. V. Ortiz, Q. Cui, A. G. Baboul,
S. Clifford, J. Cioslowski, B. B. Stefanov, G. Liu, A. Liashenko,
P. Piskorz, I. Komaromi, R. L. Martin, D. J. Fox, T. Keith,
M. A. Al‐Laham, C. Y. Peng, A. Nanayakkara, M. Challacombe,
P. M. W. Gill, B. Johnson, W. Chen, M. W. Wong, C. Gonzalez,
J. A. Pople, Gaussian 03, Revision B.03, Gaussian Inc., Pittsburgh,
PA 2003.
ORCID
Subbiah Thirumaran
http://orcid.org/0000-0002-5771-
5328
REFERENCES
[1] P. J, Prog. Inorg. Chem. 2005, 53, 1.
[2] G. Hogarth, Prog. Inorg. Chem. 2005, 53, 71.
[3] F. H. Allen, Acta Crystallogr B 2002, 58, 71.
[4] F. Shaheen, A. Badshah, M. Gielen, C. Gieck, M. Jamil, D. de
Vos, J. Organometal. Chem. 2008, 693, 1117.
[5] V. Alverdi, L. Giovagnini, C. Marzano, R. Seraglia, F. Bettio, S.
Sitran, R. Graziani, D. Fregona, J. Inorg. Biochem. 2004, 98, 1117.

12 of 12
GURUMOORTHY ET AL.
[26] E. Sathiyaraj, G. Gurumoorthy, S. Thirumaran, New J. Chem.
2015, 39, 5336.
[41] A. Meleros, A. Cachapa, R. Hernantoz‐Molina, M. T. Armas, P.
Gili, M. Sokolov, J. Gonazalez‐Plates, F. Brito, Inorg. Chem.
Commun. 2003, 6, 498.
[27] F. Bonati, R. Ugo, J. Organometal. Chem. 1967, 10, 257.
[42] W. E. Hatfield, P. Singh, F. Nepveu, Inorg. Chem. 1990, 29,
[28] L. Ronconi, L. Giovagnini, C. Marzano, F. Bettio, R. Graziani,
4214.
G. Pilloni, D. Fregona, Inorg. Chem. 2005, 44, 1867.
[43] C. F. Matta, J. Hernandez‐Trujillo, T.‐H. Tang, R. F. W. Bader,
Chem. – Eur. J. 2003, 9, 1940.
[29] J. O. Hill, R. J. Magee, Rev. Inorg. Chem. 1981, 3, 141.
[30] A. W. Herlinger, S. L. Wenhold, T. V. Long, J. Am. Chem. Soc.
1970, 92, 6474.
[44] A. Bendjeddou, T. Abbaz, A. Gouasmia, D. Villemin, Am. J.
Appl. Chem. 2016, 4, 104.
[45] J. Luque, J. M. Lopez, M. Orozco, Theor. Chem. Acc. 2000, 103,
[31] J. Criado, I. Fernandez, B. Macias, J. M. Salas, M. Medarde,
343.
Inorg. Chim. Acta 1990, 174, 67.
[46] E. Scrocco, J. Tomas, Adv. Quantum Chem. 1978, 11, 115.
[32] N. Manav, A. K. Mishra, N. K. Kaushik, Spectrochim. Acta A
2006, 65, 32.
[47] Q. W. Tian, M. H. Tang, Y. G. Sun, R. J. Zou, Z. G. Chen, M. F.
Zhu, S. P. Yang, J. L. Wang, J. H. Wang, J. Q. Hu, Adv. Mater.
2011, 23, 3542.
[33] C. K. Jorgensen, J. Inorg. Nucl. Chem. 1962, 24, 1571.
[34] D. Coucouvanis, Prog. Inorg.Chem 1980, 26, 302.
[35] A. N. Gupta, V. Singh, V. Kumar, A. Rajput, L. Singh, M. G. B.
SUPPORTING INFORMATION
Drew, N. Singh, Inorg. Chim. Acta 2013, 408, 145.
Additional Supporting Information may be found online
in the supporting information tab for this article.
[36] G. M. de Lima, D. C. Menezes, C. A. Cavacanti, J. A. F. dos
Santos, I. P. Ferreira, E. B. Paniago, J. L. Wardell, S. M. S. V.
Wardell, K. Krambrock, I. C. Mendes, J. Mol. Struct. 2011,
988, 1.
How to cite this article: Gurumoorthy G,
Thirumaran S, Ciattini S. Synthesis and
[37] I. P. Ferreira, G. M. de Lima, E. B. Paniago, J. A. Takahashi, K.
Krambrock, C. B. Pinheiro, J. L. Wardell, L. C. Visentin, J. Mol.
Struct. 2013, 1048, 357.
characterization of copper(II) dithiocarbamate
complexes involving pyrrole and ferrocenyl
moieties and their utility for sensing anions and
preparation of copper sulfide and copper–iron
sulfide nanoparticles. Appl Organometal Chem.
2018;e4363. https://doi.org/10.1002/aoc.4363
[38] F. F. Jain, Z. X. Wang, Z. P. Bai, X. Z. You, H. K. Fun, K.
Chinnakali, I. A. Razak, Polyhedron 1999, 18, 3401.
[39] F. W. B. Einstein, J. S. Field, Acta crystalllogr. B, 2929, 1974, 30.
[40] S. C. Ngo, K. K. Banger, M. J. DelaRosa, P. J. Toscano, J. T.
Welch, Polyhedron 2003, 22, 1575.