Synthesis, Characterization and Applications of Selected

Transition Metal Sulfides Nanoparticles

A thesis submitted to the Department of Chemistry, Quaid-i-Azam

University Islamabad, in partial fulfillment of the requirement for the degree

of

Doctor of Philosophy

In

Inorganic/Analytical Chemistry

By

AZAM KHAN

Department of Chemistry, Quaid-i-Azam University, Islamabad 45320,

Pakistan

(2018)

I Dedicate This Work to my

Loving Parents

Table of Contents

Acknowledgements i

Abstract iii

List of Figures v

List of Tables x

List of Schemes xi

List of Abbreviations xii

Chapter 1 Introduction 1-48

1.1 Nanoscale, nanoscience, nanotechnology and nanomaterials 1

1.2 History and origin of nanotechnology 3

1.3 Transition metal sulfide nanomaterials 5

1.4 Liquid-phase synthesis of transition metal sulfide nanostructures 6

1.4.1 Hot injection method 6

1.4.2 Hydrothermal method 6

1.4.3 Solvothermal method 7

1.4.4 Microwave irradiation-assisted aqueous synthesis 8

1.4.5 Sonochemical method 9

1.4.6 Template-directed method 9

1.4.7 Colloidal solution (precipitation) method 10

1.2.8 Photochemical method 10

1.4.9 Single-source precursor method 11

1.5 Application of transition metal sulfides nanomaterials 12

1.5.1 Biological applications 13

1.5.2 Energy based applications 14

1.5.3 Catalysis 22

References 30

Chapter 2 Experimental 49-65

2.1 Materials 49

2.2 Instrumentation 49

2.3 Fabrication of rGO-PEI-MoSx electrode 50

2.4 Cyclic voltametric measurements and electrolysis 52

2.5 Calculation of the weight difference of bare and rGO-PEI-MoSx modified

glassy carbon plate electrode 53

2.6 General synthetic procedure of ligand salts and their complexes 53

2.7 Synthesis of Mercury(II) Cadmium(II) and Zinc(II) Complexes 54

2.8 Synthesis of metal sulfide (Metal = Hg, Cd, Zn) nanoparticles 60

2.9 Computational studies 62

2.10 Photocatalytic test of CdS nanospheres and nanocapsules for the

degradation of Congo red 62

2.11 Photocatalytic reduction of p-nitrophenol to p-aminophenol by CdS nano-

sheets and nanorods 63

References 65

Chapter 3 Results and Discussion 66-132

3.1 Characterization and electrocatalytic CO2 reduction activities of rGO-PEI-

MoSx 66

3.1.1 Characterization of GO-PEI by Fourier Transform Infrared

spectroscopy (FT-IR) 66

3.1.2 Fabrication and characterization of rGO-PEI-MoSx 67

3.1.3 CO2 reduction activities of rGO-PEI-MoSx 71

3.1.4 Possible origins of the activity of rGO-PEI-MoSx for CO2 reduction 77

3.2. Spectroscopic characterization of metal complexes 80

3.3 X-ray single crystal analysis of ZnL6, ZnL7, HgL1 and HgL2 82

complexes

3.3.1 Structural descriptions of ZnL6 and ZnL7 complexes 82

3.3.2 Molecular packing of ZnL6 and ZnL7 complexes 87

3.3.3 Structural descriptions of HgL1 and HgL2 complexes 91

3.3.4 Molecular packing of HgL1 and HgL2 complexes 95

3.4 Formation mechanism and characterization of metal 97

sulfide nanostructures

3.5 Solar-light driven photocatalytic degradation of Congo red dye by CdS

nanosphers and nanocapsules 111

3.6 Solar-light driven photocatalytic conversion of p-nitrophenol to p-

aminophenol on CdS nanosheets and nanorods 116

References 121

Conclusions 131

List of Publications 133

i

Acknowledgements

All praises be to ALLAH (SWT), the Creator of all the creatures of universe, Who

created us in the structure of human beings as the best creature. I owe my insightful

thanks and cordial sense of gratitude to Almighty ALLAH (SWT) who blessed me with

endurance, prospective and ability to complete my Ph.D. work. All respects to Holy

Prophet Hazrat Muhammad (PBUH) who exhorted his followers to seek for knowledge

from cradle to grave.

I feel proud to be a student of a learned supervisor Dr. Zia ur Rehman, whose personal

interest, thought-provoking guidance, valuable suggestions and discussions enabled me to

complete this challenging task. His patience and willingness as a teacher to develop my

understanding and comprehension of the topics of my Ph.D. research work is greatly

appreciated.

I am grateful to Prof. Dr. Muhammad Siddiq, Chairman, Department of Chemistry, for

providing necessary research facilities. I am obliged to Prof. Dr. Amin Badshah, Prof. Dr.

Saqib Ali and Dr. Afzal Shah for their moral support.

I am also thankful to Dr. Amir Waseem for help in TEM analysis and Dr. Azhar Iqbal for

granting access to PL analysis.

Acknowledgement would be incomplete if I do not present gratitude to all teachers of the

Department of Chemistry Quaid-i-Azam University Islamabad.

I am also really very grateful to Dr. Jie Zhang, School of Chemistry Monash University

Australia for accommodating me in his research group for six months and for his valuable

pieces of advices regarding my research work. I extend gratitude to my lab fellows in

School of Chemistry Monash University Australia, especially Fenwang Li, Sixuan Zhao

and Shu-feng Zhao for their help in undertaking the electrochemical measurements and

other analysis.

I am especially thankful to the Higher Education Commission (HEC) of Pakistan for

grant of indigenous and IRSIP scholarships.

I am also grateful to Dr. Muneeb ur Rehman, Department of Physics, Islamia College

University Peshawar, Dr. Rajwali Khan and Dr. Zulfiqar, Department of Physics Abdul

ii

Wali Khan University Mardan and Dr. Yaqoob Khan NCP, QAU Campus Islamabad for

their fruitful suggestions during my Ph.D. research.

Thanks are offered to my lab fellows M. Kashif Amir, Jamal Abul Nasir, Haseeb Ullah,

M. Imran, Faisal Hayat, Dr. Shahan Zeb, Abrar Ahmad, Noor ul Ain, Asma Amir,

Wagma Ayub, Hina Ambreen, Mehwish Arshad, Bushra and all juniors for their

company and respect.

I would like to acknowledge the nice cooperation of all employees of this department.

I have no words to render my feelings of admiration about my learned and zealous

Father, Greater Mother, my siblings, my wife and all of my cousins especially

Muhammad Islam for their unlimited patience, cordial prayers and support which gave

me confidence and courage to accomplish this target.

May ALLAH (SWT) shower his blessings on all of us.

AZAM KHAN

iii

Abstract

Currently, mankind is facing the most challenging and solution demanding energy and

environment related issues. Transition metal sulfide nanoparticles (TMS NPs) have

attracted significant attention for the mitigation of these problems since the previous

decade, due to their diverse structural types and tremendous physicochemical properties

governed by size and morphology. In this work, four different types of TMS NPs i.e,

MoSx, ZnS, CdS and HgS have been synthesized by adopting the electrochemical

deposition and single source precursor strategy. Among these four, MoSx was

electrochemically deposited on polyethylenimine (PEI) modified reduced graphene oxide

(rGO-PEI-MoSx) substrate. After characterization by FT-IR, SEM, TEM, EDS, XRD and

Raman spectroscopy it was tested as a heterogeneous electrocatalyst for the reduction of

CO2 to CO in CO2 saturated aqueous NaHCO3 solution with high efficiency and

selectivity. The catalyst is capable of producing CO at overpotential as low as 140 mV

and reaches a maximum faradaic efficiency of 85.1% at 540 mV. However, at 290 mV

syngas (CO + 2H2) was formed instead of CO formation. Detailed investigations reveal

that PEI works as a co-catalyst by synergetic effect.

A single source precursor strategy was adopted for the synthesis of HgS, CdS and ZnS

NPs by using their respective metal (II) dithiocarbamates as precursors. Prior to their use

as precursors all the complexes were characterized by CHNS, FT-IR, 1HNMR and

13CNMR spectroscopy. Four of the complexes i.e. HgL1, HgL2, ZnL6 and ZnL7 were

also characterized by single crystal XRD analysis. The conversion of parent complexes to

their off-spring HgS, CdS and ZnS NPs was achieved by thermolysis of the

corresponding precursors in ethylenediamine (en) at ambient pressure and temperature,

devoid of any externally added toxic surfactant. The morphology, structure, phase and

elemental composition of as-obtained MS (M = Hg, Cd and Zn) products were

characterized by SEM, TEM, XRD and EDS analysis. The SEM and TEM results

revealed precursor based significant morphological variation for HgS and CdS NPs

assignable to precursor’s stability/solubility in en and presence of capping agent on the

particle surface (as revealed by DFT studies and FT-IR analysis). The XRD pattern

showed that HgS NPs has grown in both typical cystalline forms i.e. black colour cubic

iv

and red colour hexagonal, however, CdS NPs into hexagonal and ZnS formed in mixed

crystalline phases. The optical properties assessment of MS (M = Hg, Cd, Zn) by UV-

Visible spectroscopy have confirmed their good absorption ability in visible (HgS and

CdS NPs) and UV (ZnS NPs) regions. Based upon the band gap suitability, CdS NPs in

different morphological forms were used as solar light driven photocatalysts for the

degradation of Congo red dye and photoconversion of an environmentally detrimental p-

nitrophenol to pharmaceutically valuable p-aminophenol. Among them, the

anisotropically grown CdS NPs showed better photocatalytic performance probably due

to their good optical absorbance and longer electron hole recombination time.

v

List of Figures

Figure No. Title Page No.

Figure 1.1: Comparison of macro, micro and nanoscale. 1

Figure 1.2: Classification of nanomaterials on the basis of dimensionality. 3

Figure 1.3: Different structures and morphologies of metal sulfide nanocrystals synthesi-

-zed by solvothermal method (a) Wire like FeS2 (b) rod-like MnS (c) belt-

like Bi2S3 (d) flowerlike CoS1.097 (e) Snowflake like Cu2S dendrite (f)

platelets like CuS (g) sphere-like In2S3 (h) Bi2S3 and (i) FeS2 nanowebs. 8

Figure 1.4: TEM images of Fe7S8 and Fe3S4 nanocrystals synthesized using a cubanetype

cluster, bis(tetra-n-butylammonium) tetrakis[benezenethiolato-μ3-

sulfidoiron] as a single source precursor. 12

Figure 1.5: (a) Cycle performance of the Ni3S2 nanowire array/Li cell. The inset shows

the SEM image of the Ni3S2 nanowire arrays grown on a nickel substrate (b)

Profiles of the normalized capacity changes with respect to the C-rate

ranged from C/10 to 10 C at the 30th cycle. (c) Cycle performance of the

Cu2S nanowire array/Li cell circulated at a high rate of 2 C. The inset shows

the SEM image of Cu2S nanowire arrays grown on the copper substrate. (d)

Profiles of the normalized capacity changes with respect to the C-rate

ranged from C/20 to 20 C of the Cu2S nanowire array/Li and Cu2S thin

film/Li cells. 16

Figure 1.6: (a,b) TEM images of NiS hierarchical hollow spheres. (c) Galvanostatic cha-

-rge discharge curves and (d) Cycling performance of NiS hollow spheres at

a current rate of 4.2 A g-1

. 18

Figure 1.7: SEM images of (a) f-NiS and (b) r-NiS. (c) Discharge/charge capacities

versus cycle number for NiS and Super P based cathodes at 75 mA g-1

(d)

Discharge-charge profiles of LABs with NiS and Super P based cathodes at

75 mA g-1

. 19

Figure 1.8: (a) Cyclic voltammograms of the MoS2, WS2, and Pt electrodes for iodide

species, scanning rate: 10 mV s-1

. (b) Photocurrent density–voltage (J–V)

curves of the DSCs using MoS2 (Blue), WS2 (Red), and Pt CEs (Black). 21

vi

Figure 1.9: Band structure alignments of the CdS–ZnS core–shell structure and

schematic of the photoexcited charge carrier distribution and related

photocatalytic reactions [158]. 23

Figure 1.10: (a) Scanning Electron Microscopy of WSe2 (scale bar, 2 μm in the main

image and 200 nm in the inset). (b) STEM image of WSe2 (scale bar, 2 nm).

(c) Catalytic performance of synthesized TMDCs in 50 vol% EMIM–BF4

and 50 vol% water electrolyte. (d) Cyclic voltammetry (CV) curves for

WSe2 NFs, bulk MoS2, Ag nanoparticles (Ag NPs), and bulk Ag in CO2

environment. Inset shows the current densities in low overpotentials. 29

Figure 2.1: Schematic illustration of the fabrication of rGO-PEI-MoSx composite on the

GCE. 51

Figure 3.1: (a) FT-IR spectra of GO (black line) and GO-PEI (red line). (b) Electrodepo-

-sition of rGO-PEI-MoSx on GCE by repeating CVs for 10 cycles. 67

Figure 3.2: (a) EDS and (b) Raman spectrum of GO-PEI and rGO-PEI-MoSx. 68

Figure 3.3: (a) CVs of bare GCE (—), rGO-PEI (—) and rGO-PEI-MoSx (—) modified

GCE in 0.1 M phosphate buffer (pH 6.81) at a scan rate of 10 mV s-1

. (b)

TEM image of rGO-PEI-MoSx (Scale bar is 400 nm). 69

Figure 3.4: (a) TEM images of rGO-PEI-MoSx, the insert is corresponding FFT pattern.

(b) XRD pattern of bare FTO (black) and rGO-PEI-MoSx on FTO (red). (c)

TEM image of electrodeposited MoSx in the absence of GO-PEI. (d) SEM

images of rGO-PEI-MoSx. 70

Figure 3.5: (a) CVs of rGO-PEI-MoSx modified GCE in N2-saturated (black curve) and

CO2-saturated (red curve) in 0.5 M NaHCO3 aqueous solution. Scan rate

was 50 mV s-1

. Insert: structure of PEI. (b) Faradaic efficiency for CO (red

bars) and H2 (blue bars) as a function of potential. 72

Figure 3.6: (a) Oxidation peak of rGO-PEI and rGO-PEI-MoSx in 0.5 M H2SO4 under

cyclic voltammetric conditions in the potential region from 0 V to 1.4 V (vs.

RHE) at a scan rate of 50 mV s-1

. (b) Linear sweep voltammogram obtained

using a rGO-PEI-MoSx modified carbon cloth (—) as working electrode in

CO2-saturated 0.5 M aqueous NaHCO3 solution. In the absence of rGO-PEI-

vii

MoSx, only a small background current was observed (—). Scan rate was 5

mV s-1

. 74

Figure 3.7: (a) Amount and faradaic efficiency of H2 (circles) and CO (squares).

Potentiostatic electrolysis at -0.4 V in CO2 saturated 0.5 M NaHCO3

aqueous solution. Measurement interval: 0.5 h. (b) Chronoamperometric

response of a rGO-PEI-MoSx modified GCE at a constant potential of -0.65

V. At the 10,000th and 50,000th second, the applied potential was switched

to open circuit potential for 500 seconds before recovering to -0.65 V. 75

Figure 3.8: (a) SEM (Scale bar is 200 nm) and (b) EDS of rGO-PEI-MoSx taken after

60,000 seconds of electrolysis. 76

Figure 3.9: (a) CVs of rGO-PEI, rGO-MoSx, rGO-PEI-MoSx modified GCE in N2

saturated and CO2-saturated 0.5 M NaHCO3 aqueous solution. Scan rate was

50 mV s-1

. (b) Faradaic efficiency for CO (red bar) and H2 (blue bars) of

rGO-PEI, rGO-MoSx, rGO-PEI-MoSx, potentiostatic electrolysis at -0.65 V

in CO2-saturated 0.5 M NaHCO3 aqueous solution. (c) CVs of rGO-PDDA-

MoSx and (d) rGO-PEG-MoSx modified GCE in N2-saturated and CO2-

saturated 0.5 M NaHCO3 aqueous solution. Scan rate was 50mV s-1

. Insert

of (c) and (d) shows the structure of PDDA and PEG, respectively. 78

Figure 3.10: HER polarization curves of rGO-MoSx (MoSx loading 6.65 nmol cm-2

,

estimated from voltammetric measurements) and rGO-PEI-MoSx (MoSx

loading 8.75 nmol cm-2

) modified GCE in (a) 0.5 M H2SO4 (b) 0.5 M

NaHCO3 and (c) 0.5 M phosphate buffer (pH 7.2) solutions respectively.

Scan rate 5 mV s-1

, with iR compensation. (d) CO production partial current

density vs. overpotential on rGO-PEI-MoSx. The partial current density of

CO is obtained by multiplying the total current density at each potential by

the faradaic efficiency of CO. The Tafel slope for rGO-PEI-MoSx is 74 mV

dec-1

. 80

Figure 3.11: Ball and stick diagram of (a) ZnL6 and (b) ZnL7 comlexes. 85

Figure 3.12: Side view of (a) ladder-like sandwiched supramolecular chains driven by C-

H···π contacts, a schematic diagram (left), a wireframe model (right) b) zig-

viii

zag arrangement of dimeric H-shaped ZnL6 in the crystal structure, a

schematic diagram (top), a wireframe model (bottom) effect. 88

Figure 3.13: A spacefill representation of two successive layers of zig-zag layered

structure of ZnL6, top view (top), side view (bottom). 89

Figure 3.14: Supramolecular parallel extension of H-shaped molecules of ZnL7 to make

a layer, top view (left), side view (right). 90

Figure 3.15: Supramolecular parallel and sideways extension of H-shaped molecules of

ZnL7 to make a layer, top view (left), side view (right). 90

Figure 3.16: Multilayered structure of compound ZnL7. 91

Figure 3.17: Ball and stick diagram of (a) HgL1 and (b) HgL2 complexes. 92

Figure 3.18: (a) Dimeric structure of complex HgL1 (b) 2D supramolecular structure and

(c) 3D supramolecular structure of complex HgL1. 95

Figure 3.19: (a) 2D supramolecular structure and (b) 3D supramolecular structure of

complex HgL2. 96

Figure 3.20: SEM images of (a) HgS-1 and (b) HgS-2 NPs. 98

Figure 3.21: FT-IR spectra of (a) HgS-1 and HgS-2 (b) HgS-4 NPs. 99

Figure 3.22: (a) SEM images of HgS-3 NPs (b) FT-IR spectrum of HgS-3 NPs. 100

Figure 3.23: SEM images of (a) HgS-4 NPs (b)(c) HgS-5 NPs (d) FT-IR spectrum of

HgS-5 NPs. 101

Figure 3.24: TEM images of (a) CdS-1(nanospheres) (b) CdS-2 (nanocapsules) (c) CdS-

3 (nanograins) (d) CdS-4 (nanosheets) and (e,f) CdS-5 (nanorods). 102

Figure 3.25: HRTEM images of (a) CdS-4 (nanosheets) and (b) CdS-5 (nanorods). 103

Figure 3.26: SEM images of the as prepared (a) ZnS-1 (b) ZnS-2 (c) ZnS-3 (d) ZnS-6

and (e,f) ZnS -7 NPs. 104

Figure 3.27: SEM images of (a,b) HgS-1 (c) HgS-2 (d) HgS-3 NPs synthesized in the

presence of OA as thermolysing solvent. 105

Figure 3.28: XRD pattern of (a) HgS-1 and HgS-4 (b) HgS-2, HgS-4 and HgS-5 NPs

synthesized in the in the presence of en (c) XRD pattern of HgS-1O, HgS-

2O and HgS-3O synthesized in the presence of OA (d) XRD pattern of

CdS-1, CdS-2 and CdS-3 NPs synthesized in the presence of en. 107

ix

Figure 3.29: (a) XRD pattern of as prepared CdS-4 and CdS-5 NPs (b) XRD pattern

CdS-4 and CdS-5 NPs after three sequential photocatalytic cycles. (c) XRD

pattern of as prepared ZnS-1 and ZnS-6 NPs (b) XRD pattern ZnS-1 and

ZnS-6 NPs after calcination at 400 oC in the atmosphere of argon. 108

Figure 3.30: (a) EDS Spectrum of HgS-1 and (b) ZnS-2 NPs. 109

Figure 3.31: UV/Visible absorption spectra of (a) HgS-1 and HgS-2 (b) CdS-1-CdS-5

and (c) ZnS-1-ZnS-3, ZnS-6-ZnS-7 NPs. 110

Figure 3.32: UV/Visible spectral changes of Congo red solution with irradiation time in

the presences of CdS-1NPs at (a) 30 oC (b) at 40

oC (c) at 50

oC and (d)

CdS-2 NPs at 30 oC. 113

Figure 3.33: (a) UV/Visible spectral changes of Congo red solution with reaction time in

the presences of CdS-1 NPs in dark (b) Photocatalytic degradation rate of

Congo red with irradiation time in the presence of CdS-1 and CdS-2 NPs (c)

The degradation kinetics of Congo red with irradiation time in the presence

of CdS-1 and CdS-2 NPs (d) Time-resolved PL of CdS-1 and CdS-2 NPs,

following excitation at 305 nm. 114

Figure 3.34: UV-Visible spectral changes associated with photoreduction of p-

nitrophenol to p-aminophenol with irradiation time in the presence of CdS-4

(a) 5 mg catalyst dose (b) 10 mg catalyst dose (c) 15 mg catalyst dose (d) 20

mg catalyst dose. 117

Figure 3.35: (a) The photocatalytic conversion of p-nitrophenol to p-aminophenol with

irradiation time at different catalyst (CdS-4) dose (b) UV-Visible spectral

changes associated with photoreduction of p-nitrophenol to p-aminophenol

with irradiation time in the presence of CdS-5 (15 mg catalyst dose) (c) The

photoconversion kinetics of p-nitrophenol to p-aminophenol with irradiation

time in the presence of CdS-1 (nanosheets) at different catalyst dose and

CdS-2 (nanorods) (d) Repetitive photocatalytic conversion of p-nitrophenol

to p-aminophenol during three sequential cycles using 15 mg of catalyst

dose of both CdS-4 and CdS-5 NPs. 119

x

List of Tables

Table No. Title Page No.

Table 2.1: Weight difference of the rGO-PEI-MoSx modified glassy carbon plate

electrode before and after electrolysis 53

Table 3.1: FT-IR stretching frequencies of dithiocarbamate complexes 81

Table 3.2: 1HNMR data of the synthesized dithiocarbamate complexes 83

Table 3.3: 13

CNMR data of the synthesized dithiocarbamate complexes 84

Table 3.4: Refinement parameter of complex ZnL6 and ZnL7 86

Table 3.5: Selected Bond lengths (Å) and Bond Angles (o) of complex ZnL6 and

ZnL7 87

Table 3.6: Refinement parameters of HgL1 and HgL2 complexes 93

Table 3.7: Selected bond lengths and bond angles of HgL1 and HgL2 complexes 94

Table 3.8: Molecular properties of HgL1-HgL4 complexes 97

Table 3.9: Reported literature data of Congo red dye degradation by CdS and hybrid

CdS NPs 115

Table 3.10: Comparison of our results with reported literature data of photocatalytic p-

nitrophenol reduction to p-aminophenol by pure CdS and CdS based

materials 120

xi

List of Schemes

Scheme No. Title Page No.

Scheme 2.1: Synthesis of ligands (L1-L5) and their complexes {Cd (CdL1-CdL5) and

Hg (HgL1-Hg5)}. Z = N (L1-L3) and C (L4 and L5). 54

Scheme 2.2: Synthesis of Zn(II) dithiocarbamate complexes. 56

Scheme 2.3: Proposed thermolysis mechanism of (a) monomeric precursor complexes

and (b) dimeric precursor complexes by en. 61

xii

List of Abbreviations

NTMS Nanostructured Transition 0-D Zero Dimensional

- Metal Sulfides 1-D One Dimensional

2-D Two Dimensional 3-D Three Dimensional

STEM Scanning Tunneling Electron AFM Atomic Force

- Microscopy - Microscopy

CNTs Carbon Nanotubes TMS Transition Metal

MS Metal Sulfides - Sulfides

TOP Tri-n-Ortho-Phosphine TOPO Tri-n-Ortho-

PV Photovoltaic - Phosphine Oxide

CTAB Cetyltrimethylammonium EG Ethylene Glycol

- Bromide LED Light Emitting Diode

UV Ultra Violet SSP Single Source

TW Terawatt - Precursor

MC Metal Chalcogenide QDs Quantum Dots

QYs Quantum Yields NIR Near Infrared

PL Photoluminescence PTA Photo thermal

NCs Nanocrystals - Ablation

ECS Energy Conversion and Storage LIBs Lithium-ion Batteries

TMDs Transition-Metal Dichalcogenides SEM Scanning Electron

OA Octyl Amine - Microscopy

TEM Transmission Electron Microscopy HCs Hydrocarbons

ORR Oxygen Reduction Reaction DOE Department of Energy

RHE Reversible Hydrogen Electrode DSSCs Dye-sensitized solar

en Ethylene Diamine - cells

SHE Standard Hydrogen Electrode NPs Nanoparticles

HER Hydrogen Evolution Reaction FE Faradaic Efficiency

DFT Density Functional Theory NFs Nanoflakes

rGO Reduced Graphene Oxide PEG Polyethylene Glycol

PDDA Polydiallyldimethyl-ammonium GO Graphene Oxide

xiii

- Chloride PEI Polyethylenimine

PXRD Powder X-Ray Diffraction CV Cyclic Voltammetry

EDS Energy Dispersive Spectroscopy IR Infra-Red

HOMO High Occupied Molecular Orbital CR Congo Red

LUMO Low Unoccupied Molecular Orbital FDH Formate

GCE Glassy Carbon Electrode - Dehydrogenases

TOF Turnover Frequency FT-IR Fourier Transform

IL Ionic Liquid - Infrared

HR-TEM High Resolution Transmission Electron Microscopy

1

Chapter 1

INTRODUCTION

1.1 Nanoscale, nanoscience, nanotechnology and nanomaterials

The word nano, Greek in origin, means something tiny. One nanometer means billionth

of a meter (10-9

m), a length equivalent to 5 silicon or 10 hydrogen atoms placed in line

[1]. To rationalize the concept of 1 nm, one should know that the growth rate of human

fingernails is about 1 nm/second, the diameter of a pen dome is almost 105

nm and the

width of DNA molecule is approximately 1-2 nm (Figure 1.1). The nanometer scale is

conventionally defined as 1-100 nm. 1 nm (1 nm = 10 Å) was considered large by

chemists, while working with the world of molecules, however, engineering physics was

taking 1 μm (1000 nm) as a smaller unit till the fourth quarter of the previous century [2].

Matter exist in the fuzzy interface in the middle of these big and small limits of length

scale instigated as a science of nanomaterials, which has been developed tremendously

into a most exciting and active field of fortitude, imparting a revolutionary impact on

materials.

Figure 1.1: Comparison of macro, micro and nanoscale.

Nanoscience doesn’t simply means the knowledge of tiny particles, but here materials

with reduced dimensions unveil novel physical phenomena, communally called quantum

effects, that are size dependent and radically dissimilar from macroscale material. Thus

the more prcise definition of nanoscience is ‘the study of phenomena and management of

materials at atomic, molecular and macromolecular scales, where properties change

2

considerably from those at a larger scale [3]. By switching to nanometer length scale, the

quantum effects appear and surface-area effects turn out to be exceedingly important that

causes remarkable variations in the properties of materials and devices. Decreasing the

size of the materials to nanoscale, result in high percentage of surface atoms of the

objects. Wolfgang Pauli expressed a quote long ago, saying: “God made the bulk; the

surface was invented by the Devil”.

In nanoscienc, the rationality of Devil's domain has been stretched massively as the

surface physics and chemistry flinch governing bulk properties. This must be properly

taken into account. The reactivity of nanoscale materials due to high percentage of

surface atoms is considered as a principal factor that distinguishes the properties of

nanostructures and bulk material [4].

Applying nanoscience in device fabrication is called nanotechnologies, which include

production, characterization, design and applications of structures, devices and systems

by adjusting size and morphology at the nanometer length scale. The first corner stone in

the field of nanotechnology is the development of facile and cost effective methods to

fabricate and process nanomaterials in order to search novel physical properties and

phenomena for the purpose of their potential applications [1].

Nanomaterial is defined as an object/particle having at least one dimension in the range

of nanoscale (approximately 1-100 nm), including nanoparticles, nanorods, thin films,

nanofibers, nanowires, nanotubes and bulk materials composed of nanoscale particles or

containing nanoscale structures [1]. The minimum size range of 1 nm is set to circumvent

single atom or their small groups to be categorized as nano-objects, however, the upper

limit (100 nm) is fluid and normally objects with larger dimensions (even 200 nm) are

considered as nanomaterial [5]. In the previous decade’s large variety of nanomaterials

have been obtained. As the major discriminatory feature of nanostructured material is

dimensionality, therefore, for further development and understanding nanomaterials have

classified into four groups on the basis of dimensionality i.e. 0-D, 1-D, 2-D and 3-D

(Figure 1.2).

Zero-dimensional nanomaterials (0-D): Those materials whose all dimensions are in

nanoscale. The most common representative example of 0-D nanomaterials are

nanoparticles (fullerenes), quantum dots, nanoshells and nanorings etc.

3

One-dimensional nanomaterials (1-D): Those materials whose one dimension is

outside the nanoscale i.e. one dimension is larger than 100 nm. This lead to needle shaped

nanomaterials. The representative examples of 1-D nanomaterials include nanotubes,

nanorods and nanowires etc.

Two-dimensional nanomaterials (2-D): Those particles whose two dimensions are

outer of the nanoscale i.e. two dimensions are larger than 100 nm. 2-D nanomaterials

exhibit plate like shape and the common examples are nanofilms, nanocoatings and

nanolayers etc.

Three-dimensional nanomaterials (3-D): Those nanomaterials whose, all the three

dimensions are outside the nanoscale. 3-D nanomaterials are composed of

nanocrystalline structures or contain nanoscale features. Hence such materials can be

either composed of various arrangements of nanosized crystals or grains or contain

bundles of nanowires and nanotubes, dispersions of nanoparticles along with

multinanolayers. Common examples of this kind of nanomaterials include fullerites,

clathrates and powder skeletons [6].

Figure 1.2: Classification of nanomaterials on the basis of dimensionality.

1.2 History and origin of nanotechnology

Modern sciences and technology give birth as a result of human dreams and imagination.

Nanotechnology is the outcome of such imagination and dreams which is continuously

developing, as commercial and academic interests continues to rise and as new research

is presented to the scientific community. The study of nanoparticles is not new, for

instance the fact that small particles of a substance have different properties than the large

particles was known for a long time, however, with no clear justification. Thus people

4

subconsciously remained involved with the nano-world phenomena since the early

history of human civilizations. The secrets of premature nano-production simply

delivered from one to other generation without explanations of their unique properties.

The famous Lycurgus cup in the British museum is one of the good examples, which is

an exceptional creation of glass manufacturers of ancient Rome. The cup possesses

unusual optical properties i.e. green in natural light but turns red if illuminated from

inside. When the research technique became advanced the researchers found that the

presences of gold and silver nanoparticles are accountable for the unfamiliar color of the

cup. The multi-colored church glass windows in Europe achieved high perfection in the

Middle ages, which contained gold and other metal nanoparticles as showed by recent

research. In 2006 P. Paufler studied the sabre fragments made from Damask steel using

electron microscope in order to investigate the reasons of ultra-strength it owned. The

results showed that the steel is composed of nano-fiberous structure, which is supposed to

induce in the structure by special thermodynamical treatment [7]. The term ‘‘nanometer’’

was first coined by Richard Zsigmondy, the 1925 Nobel Prize Laureate in chemistry. He

used the term while measuring the particle size of gold colloids using a microscope [8].

Although nanotechnology has its background in the distant past, yet did not really get

going until the 2nd

half of previous century. The idea of modern nanotechnology was

presented by Richard Feynmen in 1959 for the first time, while delivering a talk in the

conference of American Physical society with a title “There is Plenty of Room at the

Bottom”. He presented the concept of controlling matter at the atomic level. He

anticipated writing of entire Encyclopedia Britannica on the head of a pin and fitting all

worlds’ books in a pamphlet. He justified his arguments by saying that the laws of

physics are not in contradiction of the opportunity of maneuvering things atom by atom.

The term ‘‘nanotechnology’’ was used for the first time by a Japanese scientist, Norio

Taniguchi, at the international session on industrial production in Tokyo in 1974. He

described the super thin dealing out of materials with nanometer precision and the

construction of nano-sized mechanisms. He endorsed that nanotechnology consist of the

processing, separation, consolidation and deformation of materials by one atom or one

molecule.

5

Inspired by Feynmen ideas, Eric Drexler conveyed the conception of molecular

manufacturing to the general public in the form of a book “Engines of Creation: The

Coming Era of Nanotechnology” (1986). In his book he pointed out that molecular

manufacturing will lead to revolutions in medicine, artificial intelligence, and the

conquest of space. Drexler technically explained the possibility of molecular

manufacturing [9].

The Drexler's work motivated other researchers towards the field of nanotechnology. Dr.

Richard Smalley specifically said Engines of Creation encouraged him to peruse

nanotechnology and that he was a "fan of Eric", as a result Smalley and coworkers

discovered fullerenes which practically started the golden era of nanotechnology [8,10].

The invention of microscopy technology (STM & AFM) in the early to mid-eighties

produced a critical effect on the development of nanotechnology, as it became possible

for the first time to clearly identify individual atoms. The discovery of carbon nanotubes

by Japanese scientist Iijima in 1992 further advanced the science of nanotechnology [11].

Since then a substantial strengthening of nano-technological exploration and projects are

ensuing and the figure of publications on nanotechnolgy subjects rises abruptly. The

word of the possibilities of nanotechnology is diffusing promptly, and the atmosphere is

dense with news of innovations in this science. Billions of dollars have been put in by

governments and businesses for research and developments in nanoscience, and

governmental alliances and encounter lines are beginning to form. Social communities

are becoming aware and getting introduction to nanotechnology due to the references to it

are becoming more common in movies, books, video games and television [9].

1.3 Transition metal sulfide nanomaterials

Nanostructured transition metal sulfides (NTMS) represent an important and interesting

class of inorganic materials owed to their fascinating optical, electrical, photovoltaic,

photocatalytic and electro-catalytic properties [12]. NTMS are generally cheap and

abundantly occurs in nature in the form of minerals such as molybdenite (MoS2),

tungstenite (WS2), heazlewoodite (Ni3S2), chalcocite (Cu2S), pyrite (FeS2), greenockite

(CdS), cinnabar (HgS) and so on, making them attractive for large scale technological

applications [13]. Controlling the morphology, size, composition, and structure of NTMS

6

by developing scalable synthetic routes is the core issues to be resolved before their

practical applications. The traditional solid-state or gaseous-state synthetic approaches for

achieving these goals are often inadequate. In contrast the liquid-based systems have

been proved as dominant approach for producing anticipated NTMS. In this method,

nucleation and growth process, which ultimately affect particles size and shape, can be

tuned by the judicial choice of solvent and/or surfactant and by altering the

thermodynamic and kinetic parameters of reaction. Furthermore, the synthesis can be

accomplished under normal conditions. Thus, many liquid-based synthetic procedures

have been developed for producing high-quality NTMS, salient of them are discussed as

follow [14].

1.4 Liquid-phase synthesis of transition metal sulfide nanostructures

1.4.1 Hot injection method

Thermal decomposition of single molecular precursors (comprising together the metal

and chalcogen atoms) in coordinating solvent is an attractive direction towards the

synthesis of metal chalcogenides nanoparticles. In this technique, the precursor is

dispersed in an organometallic reagent tri-n-ortho-phosphine (TOP) tracked by

decomposition at higher temperature (typically > 250 oC) [15]. The nanoparticles (NPs)

formation (nucleation) is initiated by the precursor’s decomposition whereas the growth

is repressed when the precursor amount is exhausted. The NPs formation is comprised of

three steps:

Preliminary addition of precursor

Fast nucleation

Organized evolution of the nuclei by Ostwald ripening.

NPs synthesized by this method show good crystallinity and good morphological

variation i.e. nanospheres, nanorods and nanoplates. Size and morphological tuning can

be easily done by monitoring the reaction conditions i.e. temperature, time, nature and

ratios of the initial reagents and aging period [14].

1.4.2 Hydrothermal method

7

In the hydrothermal process, generally applied for the production of metal sulfides

nanostructures (MSNs), a precursor is heated in airtight steel pressure vessel in the

presence of water to a desired reaction temperature. The sealed vessel is usually heated

above 100 oC to achieve the pressure of vapor saturation, to facilitate autogenous pressure

to be established in a sealed system. The reaction temperature, volume of liquid and

dissolved salts collectively contributes to the in situ generated pressure. Hydrothermal

method has been proved an effective method in synthesizing various nanostructures of

metal sulfides up to now, such as hexagonal plates of CuS were prepared from copper

chloride and sodium thiosulphate precursors in the presence of CTAB and HNO3 as

assisting agents [16]. In another study molybdenum sulfide, iron sulfide and copper

sulfide NPs have been created by this technique, using sodium thiosulphate pentahydrate

(Na2S2O3.5H2O) and hydroxylamine sulphate (NH2OH)2.H2SO4 as initial materials and

reacted with the respective metal salts [17]. The various morphologies of other metal

sulfides obtained with hydrothermal method include MoS2 hollow cubic cages [18] beta-

In2S3 nanoflowers [19] CuS microtubules [20] SnS2 nanocrystals [21] Sb2S3 nanoribbons

[22] and so on.

1.4.3 Solvothermal method

The solvothermal method differs from hydrothermal one in the sense of using non-

aqueous solvent as a reaction medium. A well-known fact that polarity, viscosity and

softness of solvent affect the solubility and transport behavior of the predecessors in

liquid-based synthesis, which governs the reactivity, size, morphology, and phase of the

final product. Thus, diverse carbon-based solvents with special physico-chemical

properties are chosen as reaction medium. The literature contains many reports about the

synthesis of metal sulfides (MS) nanostructures by solvothermal method with

sophisticated control of the size, morphological distribution and crystallinity, including

wire-like FeS2 [23] rod-like MnS [24] belt-like Bi2S3 [25] flowerlike CoS1.097 [26]

snowflake like Cu2S dendrite [27] platelets like CuS [28] sphere-like In2S3 [29] Bi2S3 [30]

and FeS2 nanowebs [31] as shown in figure 1.3a-i, respectively. The temperature of the

process is generally kept in the range of 100–250 oC and the sulphur sources are usually

elemental sulfur, CS2 and thiourea etc.

8

Figure 1.3: Different structures and morphologies of metal sulfide nanocrystals

synthesized by solvothermal method. (a) Wire like FeS2 (b) rod-like MnS

(c) belt-like Bi2S3 (d) flowerlike CoS1.097 (e) Snowflake like Cu2S dendrite

(f) platelets like CuS (g) sphere-like In2S3 (h) Bi2S3 (i) and FeS2 nanowebs

[23-31].

1.4.4 Microwave irradiation-assisted aqueous synthesis

In this technique microwave irradiations are used for initiating chemical reactions. The

radiations create electromagnetic field which imposes force over charged particles in

reaction medium causing charge movement or molecular rotation. This could induce

more polarization in case of polar molecules. The total friction and collisions between

molecules have been reckoned in relations of thermal and non-thermal effects [32]. Using

this technique, different morphologies of copper sulfide having sub micrometers

dimensions were synthesized by using diverse copper and sulphur ion sources [33]. Other

MS (M = Co, Pb, Cd, Cu, Zn), M2S (M = Ag), M2S3 (M = Sb, Bi) were also effectively

9

produced under the microwave irradiation by the reaction of respective metal salts with

thiourea in ethylene glycol (EG) [34]. Other morphologies of CuS including spherical,

tubular and leaf like have been created through this technique [35,36]. Molar ratio of

precursors was found to be a critical factor tuning morphology of the NPs [35]. The small

reaction time, reduced particle size and purity of the products are the claimed benefits of

this method [37].

1.4.5 Sonochemical method

Sonochemical method has progressed as dynamic technique for the fabrication of

nanomaterials. Herein the interaction between ultrasonic radiation and solution induces

molecular vibrations and the energy is conveyed to the solution, which triggers the

reaction [38]. The ultrasound radiation can rise the reactivity of metal powders upto 105-

times in a liquid–solid heterogeneous system [39]. It is suggested that the chemical

effects of ultrasound emanate from acoustic cavitation instead of direct interaction with

precursors. By exposing liquid to ultrasound stimulations, bubbles passes through the

course of development, evolution and implosive breakdown.

The breakdown of effervesce delivers powerful local heating (~5000 K), huge pressures

(over 1800 atm), and immense cooling rates (~1010 K s-1

), which facilitate chemical

reactions to take place [40]. In this method, Usually the solution is irradiated with

ultrasonic radiation of 20- 40 kHz, which hold both metal and chalcogenide ions source.

The fast reaction rates, controllable reaction parameters, uniform shape and size

distribution are the major associated advantages of the technique. The literature contains

many reports about synthesis of the transition metal sulfide nanomaterials such as HgS,

PbS, CuS, NiS [41,42] by this route. NPs of various morphologies such as nanorods,

nanowires or nanotubes have been produced in the presence of growth assisting agents.

[43,44].

1.4.6 Template-directed method

This method is among the popular approaches for the synthesis of functional materials of

numerous nanostructures. Commonly, two types of templates are employed in this

method, hard and soft template. Some of the hard templates functions only as a physical

framework for the deposition of preferred material coating while others in addition to

10

shape defining also acts as reactant i.e. react with other compounds/elements to generate

required nanostructures. Soft template contains ligands, surfactants, polymers, and

organogelators. Various TMSs nanostructures with different morphologies have been

fabricated using the pre-formed nanostructures as hard templates, such as porous

aluminum oxide template was used for manufacturing MoS2 nanofibers and nanotubules

[45]. Cu2S nanoribbons prepared from Bi2S3 nanoribbons [46] mesoporous WS2 and

MoS2 prepared from mesoporous silica [47] and Ni chains were used as template for

making Ni/Ni3S2 peapods and NiS hollow chains [48].

1.4.7 Colloidal solution (precipitation) method

The precise precipitation reaction resulting in dilute suspension of semiconductors

nanoparticles is a common synthetic route. In this method the metal ions solution is

slowly injected to aqueous or non-aqueous solution having chalcogen ion. Arresting the

growth step after nucleation by adjusting the reaction equilibrium between the solvated

ions and nanocrystals of anticipated chalcogenides affords controlled size nanocrystals.

The parameters such as solvent role, molar ratio and concentration of precursor ions and

temperature have been identified to be very significant for controlling such equilibrium.

In addition, the capping agents (surfactants, polymers etc) are helpful in achieving

monodispersed and stabilized nanoparticles, specifically, the thiol-capped chalcogenides

have been found very stable. In this method commonly used S2-

source includes H2S,

Na2S, thiourea and thioacetamide [49-51]. The poor crystallinity, surface defects and lack

of uniform size distribution are the major disadvantages associated with this method

[52,53].

1.4.8 Photochemical method

In photochemical synthesis the metal sulfide nanocrystals are obtained by using a

powerful laser to initiate the breakdown of starting compound and then commencement

of nucleation and evolution of the NPs. The selection of the light source like UV, g-ray

and LED is made on the basis of the photochemical stability of the precursors. O’Brien

et al. carried out the photolysis of Na[Eu(S2CEt2)4]∙3.5H2O by using white LED light

and obtained EuS NCs of 9 nm in diameter [54]. The synthesis of other metal sulfides

11

nanostructures such as ZnS [55] NiS, PdS [56] WS2 [57] etc. have also been reported in

literature by photochemical method. Beside the long reaction time and often low yield,

the advantages of low cost, safety, low reaction temperature and environmental

friendliness brand this procedure striking for the synthesis of MS nanomaterials [54].

1.4.9 Single-source precursor method

Single source precursor (SSPs) method is among the most effective and standard route

for synthesis of high quality MS nanocrystals. SSPs can be defined as a molecule which

encloses both metal and chalcogen elements, and upon decomposition gives desired MC

NPs. By modifying the reaction temperature and time, shape and size of MC

nanoparticles can be effectively switched [14]. Furthermore, SSPs carry the advantage of

metal-chalcogen bond template, which can be useful for controlling the stoichiometry and

shape of the nanocrystals [58,59]. The preformed M-S bond existence in the SSP

molecule facilitates and tunes the growth kinetics of the nanocrystals, and ultimately size

and morphology of the required end product. Additionally, the different binding strength

of ligands may be responsible for the dissimilarity in the stability of precursors or

decomposition kinetics; hence this anomalous decomposition pattern can be exploited for

tunning the NPs size and morphology [60]. So far, many size and shape controlled MS

have been generated from SSPs such as nanoparticles of Ag2S [61] alpha-MnS [62]

nanoplates of Fe3S4 [63] nanosheets of MoS2 and WS2 [64] and so on.

O’Brien and Vanitha [65] decomposed cubane-type cluster (Figure 1.4)

[NnBu4]2[Fe4S4(SPh)4], bis(tetra-n-butylammonium) tetrakis[benezenethiolato-μ3-

sulfido-iron] in the temperature range of 180-250 oC in various solvents

(hexadecylamine, dodecylamine, oleylamine and octylamine,) and achieved iron sulfide

NPs with good shape and size distribution (Figure 1.4). By increasing the temperature

above 180 oC in alkylamine, a fascinating phase transfer from pyrrhotite type Fe7S8 to

Fe3S4 with a greigite structure was observed. In another study, Korgel et al. thermally

degraded nickel thiolate predecessors with octanoate and achieved rhombohedral NiS

(millerite) nanorods and triangular nanoprisms [66]. The decomposition of [Cu

(mdpa)2][CuCl2] in ethylenediamine and ethylene glycol afforded hexagonal chalcocite

Cu2S nanostructures [67]. In another study spherical and rod-like CdS NPs were obtained

12

by thermolysis of [Cd(mdpa)2Cl2] and [Cd(bdpa)2Cl2] respectively, without using any

external surfactants, In situ generated thiols in thermolysis of the SP governed the growth

of nanocrystals [58]. The above achievements indicate the effectiveness of this technique

for directing the size and shape of MS nanomaterials.

Figure 1.4: TEM images of Fe7S8 and Fe3S4 nanocrystals synthesized using a cubane

type cluster, bis(tetra-n-butylammonium) tetrakis[benezenethiolato-μ3-

sulfido-iron] as a single source precursor [65].

1.5 Application of transition metal sulfides nanomaterials

As discussed in the forgoing section, the synthesis of NTMS of different sizes and shapes

can be successfully achieved by different liquid-phase synthetic methods. Being an

important and typical class of semiconductors, these materials have received significant

attention owing to their diverse structural types and tremendous physicochemical

properties, which are influenced by their size and morphology [68]. Reducing the particle

size of MS to nanoscale, the emergence of well known quantum size effect and high

surface to volume ratio (huge number of surface atoms) [69] makes them highly potent

for variety of applications as discussed below.

13

1.5.1 Biological applications

The distinctive optical properties of MS QDs i.e. large extinction coefficients, good

quantum yields (QYs), notable photo-stability, and narrow emission spectra brand them

suitable candidates for biological imaging, targeting and therapy etc. [70,71].

a. Bioimaging

The manipulation of fluorescence emission property of MS QDs for bioimaging has

developed as prevailing area of research. CdSe@ZnS core-shell QDs are presently the

most common material used for in vitro imaging, which offers modifiable optical

properties in range of 400 - 700 nm (visible spectrum) [72,73]. However, the CdSe@ZnS

QDs visible emission is interrupted by the autofluorescence of bodily tissues, creating

uncertainty in imaging. To solve this problem, some near infrared (NIR) emitting MS

QDs, in the first near-infrared window (NIR-I, 650-950 nm), like CdHgTe/CdS QDs,

CdTexSe1-x@CdS and CuInS2@ZnS QDs were used for bioimaging [74-76].

Furthermore, exhausting QDs that can release radiations in the second near-infrared

window (NIR-II, 1000 nm-1400 nm) improve the signal-to-noise ratio up to 100-folds

than NIR-I emitters, a fact well explained by simulations and modeling studies [77].

Through the bio-conjugation of silver sulfide quantum dots (PL emission in the NIR-II

region) [78, 61] with particular ligands, the targeted labeling and imaging of various cell

lines were accomplished. The cytotoxicity of silver sulfide quantum dots were

systematically investigated in terms of apoptosis and necrosis, cell proliferation, reactive

oxygen species and DNA damage assessments, ended with observation that the

cytotoxicity is insignificant for Ag2S QDs with bright PL and high biocompatibility [79].

b. Therapy

Photothermal ablation (PTA) therapy is considered potentially an effective treatment for

tumors as it is minimally offensive to the surrounding healthy tissues than the common

tactics i.e. chemotherapy, hormone therapy surgery, radiation therapy [80,81]. PTA

therapy is composed of three key components i.e. light, photothermal transformation

agent, and heat exchange efficacy. The absorbance capability of NIR laser (λ=700-1100

nm) by biological tissues is small and can reach up to several centimeters depth in tissues,

14

however, the improvement of NIR laser-induced PTA require access to biocompatible

and proficient photothermal coupling agents. The little cytotoxicity, cost and intrinsic

NIR region absorption of copper sulfide NPs motivated Chen et al. to use them for the

first time in PTA therapy [82]. The low photothermal transformation efficiency of CuS

NCs prompted researchers to devise a strategy for developing novel structured CuS and

use new NIR light to improve photothermal conversion efficiency.

Hu and coworkers achieved high NIR photothermal exchange efficiency by irradiating

hydrophilic flower-like CuS ( energy gap of 1.85 eV) with a 980 nm NIR laser (power

density = 0.51 W cm-2

), which destroyed the cancer cells in 5-10 min [83]. The same

group further improved the photothermal conversion efficiency (25.7 %, induced by a

980 nm laser) by using hydrophilic Cu9S5 NPs (size ~70×13 nm) with a band gap of 1.57

eV, and observed the in vivo killing of cancer cells by photothermal effects of the Cu9S5

NPs using the irradiation of 980 nm laser for a short period (∼10 minutes) [84]. The work

can be extended to other metal sulfide NPs of diverse particle sizes and morphologies

with the aim to get more efficient photothermal conversion reagents.

1.5.2 Energy based applications

Energy and environmental problems are among the most challenging issues facing

mankind in this century. The global energy demand is projected to increase from 15–17

TW in 2010 [85] to 25–27 TW by 2050 [86]. The contribution of fossil fuels to this total

world energy demand is estimated 81 % while nonfossil fuels collectively accounted for

13 % in 2011 [87]. If the reliance for energy demand is continued on increasingly

diminishing fossil fuel reserves, worldwide competition is predicted to increases in the

next 50 years, which will results in higher costs, both commercially and politically [88].

In spite of the heavy reliance on fossil fuels for energy production, they are not

considered ideal because of the emission of CO2 upon combustion, which is the most

notorious greenhouse gas causing catastrophic climate change, [89]. To alleviate these

energy and environmental problems, there is urgent need to develop efficient,

economical, and pollution free energy conversion and storage (ECS) devices that can be

used to run power demanding areas such as cell phones, camcorders, laptops, electric

vehicles and hybrid electric vehicles etc. [90]. Nanostructured transition metal sulfides

15

have an important status in electrochemical devices owing to their outstanding

electrochemical activities, mechanical and thermal stabilities and recyclability. The

applications of transition metal sulfide nanostructures in typical electrochemical devices

such as supercapacitors, Li ion batteries, solar and fuels cells are discussed below.

a. Li ion batteries

Lithium-ion batteries (LIBs) are trustworthy power storing devices for electric vehicles,

smart grids and have also become the main power source in most of portable electronic

devices [91]. The energy density of graphitic anode in commercial LIBs is still low with

hypothetical capacity of only 372 mA h g-1

, thus different kind of materials have been

designed and investigated as electrode in LIBs for the purpose to accomplish the

challenge of low energy density. Transition metal sulfides are potentially strong

candidates to be used as electrode material in LIBs due to low cost, great lithium loading

capability and extended cycling life [92]. MoS2 was evaluated in 1980 for the first time as

Li-storage electrode material [93], since then it has been studied as electrode material for

LIBs in different shapes and structures [94]. A variety of other TMSs NPs have been

investigated as cathode materials, which include CuS [95], Cu2S [96], NiS [97], TiS2

[98], VS2 [99] and so on. Chen and co-workers achieved remarkable rate capability and

capacity by using Ni2S3 [100] and Cu2S [101] nanowire arrays as cathode materials,

which was 2.5 times (Ni2S3) and 1.6 times (Cu2S) higher than commercially used LiCoO2

(<145 mA h g-1

). The recycling efficacy and the disparity of the normalized capacity with

respect to the C rate of Ni3S2 and Cu2S nanowire arrays/Li cells cycled are displayed in

figure 1.5a&b and b&c, respectively. This exceptional electrochemical activity of

nanowire arrays was accredited to the increased reaction sites and efficient charge

transport. A number of other TMSs have also been considered as anode material in LIBs,

such as In2S3 [102] ZnS [103] and two dimensional (layered) TMDs (MS2, M = Mo, Co,

Zr, W) [104-107]. The layered transition metal sulfides have analogous structure to

graphene i.e.

16

Figure 1.5: (a) Cycling performance of Ni3S2 nanowire array/Li cell. The inset shows the

SEM image of the Ni3S2 nanowire arrays grown on a nickel substrate (b)

Profiles of the normalized capacity changes with respect to the C-rate

ranged from C/10 to 10C at the 30th

cycle. (c) Cycle performance of the

Cu2S nanowire array/Li cell circulated at a high rate of 2C. The inset shows

the SEM image of Cu2S nanowire arrays grown on the copper substrate. (d)

Profiles of the normalized capacity changes with respect to the C-rate

ranged from C/20 to 20C of the Cu2S nanowire array/Li and Cu2S thin

film/Li cells [100,101].

having three stacked atomic layers (S–M–S) which can function as superb intercalation

hosts for lithium, thus making them specifically important as anode material in LIBs. In

the course of lithium intercalation, thorough charge transfer takes place, which involve

the simultaneous reduction of M4+

to M3+

and transmission of Lithium ions to the van der

Waals gaps [108].

17

b. Supercapacitors

The amazing characteristics of electrochemical capacitors or supercapacitors such as fast

charging and discharging, elevated power density, amazing durability (>100K cycles),

variable temperature operational range and safety have appealed for intensive research

interest in these energy storage/conversion devices. Currently, the effective electrode

materials for supercapacitors as reported in the literature are conducting polymers, metal

oxides and carbon-based compounds [109].

Nonetheless, research efforts for the development of novel materials to be used as

electrode with better energy-power proficiency and reduced price are still continued to

recognize their widespread practice. Research interests in TMS NPs as electrode material

has increased recently owing to their stimulating fundamental properties in conjunction

with decent performance [110,111]. Lou et al. [112] synthesized hierarchical NiS hollow

spheres (Figure 1.6a&b) by template-engaged conversion method which revealed

prominent specific capacitances of 583–927 F g-1

at different current densities of 4.08–

10.2 A g-1

(Figure 1.6c). They also found 70 % retention of the original capacitance next

to 1000–3000 cycles (Figure 1.6d).

The same group [113] prepared CoS2 ellipsoids with and achieved high specific

capacitance of 1040 F g-1

at a current density of 0.5 A g-1

, the better performance of

material was ascribed to the new anisotropic voids, which assist the infusion of

electrolyte and offer high active surface.

Wu et al. achieved an excellent cycling stability and high capacitance with ultra-small

Cu1.96S nanoparticles (~10 nm) consistently rooted in porous carbon octahedra [114]. In a

recent study, Tang et al. obtained a great capacitance of 704.5 F g-1

and good retention

(62.6 %) next to 5000 cycles by using MnS nanocrystals as supercapacitor materials

[115]. MoS2 has also been inspected for supercapacitance by testing it as electrode

material [116]. Chhowalla’s et al. studied monolayer nanosheets of MoS2 [117] as

electrode material and observed the effective intercalation of ions in MoS2 nanosheets.

The as synthesized material demonstrated the capacitance rate reaching from ~400 to

~700 F cm-3

in different kinds of electrolytes in water.

18

Figure 1.6: (a,b) TEM images of NiS hierarchical hollow spheres. (c) Galvanostatic

charge discharge curves and (d) Cycling performance of NiS hollow spheres

at a current rate of 4.2 A g-1

[112].

c. Fuel cells

Fuel cells produce electricity for automotive impulsion by oxidizing fuels at the anode

and reducing O2 at the cathode. Fuel cells have attained great advancement in numerous

technical aspects [118], but the intrinsic sluggish kinetics and low reversibility of oxygen

reduction deter their hands-on use [119]. To overcome these downsides, the creation of

more efficient and inexpensive oxygen reduction reaction (ORR) catalyst is needed. The

materials recognized as strong ORR catalysts include noble Pt-alloys [120] perovskites

oxide [121] nanocrystalline spinels [122] and carbonaceous materials [123]. Interestingly,

transition metal sulfides nanomaterials are also strong candidates as ORR catalysts

attributable to their reduced price and analogous catalytic activity. For ORR cobalt

19

sulfides (CoxSy) are commonly established prospective catalysts, but the issues of its low

conductivity and structural instability are to be solved for practical application. To handle

these issues the various morphologies and composites of cobalt sulfide has been

synthesized and investigated for ORR, such as CoS2 NPs [124], Co9S8 microspheres

[125], Co3S4 nanotubes-graphene composite [126], Co1-x/rGO hybrid [127] etc. and

obtained a fruitful improvement for ORR.

Figure 1.7: SEM image of (a) f-NiS and (b) r-NiS. (c) Discharge/charge capacities

versus cycle number for NiS and Super P based cathodes at 75 mA g-1

(d)

Discharge-charge profiles of LABs with NiS and Super P based cathodes at

75 mA g-1

[129].

Beside the advancement in cobalt sulfides based ORR catalysis; the other MS NPs have

also been studied. Surendranath et al. tested thin films of Ni3S2 for ORR which was

electrodeposited on gold electrodes. The electrode exhibited better catalytic activity than

polycrystalline platinum by showing onset potential at 0.8 V (vs. RHE) and faradaic

20

efficiency of approximately 90% for four-electron reduction of O2 in neutral phosphate

buffer [128]. Zhang and co-workers prepared NiS having flower and rod like

morphologies (Figure 1.7a&b) by one-step hydrothermal approach and tested the as

obtained materials as cathode in Li-air batteries (non-aqueous). Interestingly, they

observed that flower-like NiS showed better activity and stability than pure Super P and

rod-like NiS by means of capacity of 6733 mA h g-1

, the charge voltage of 4.24 V at the

current density of 75 mA g-1

(Figure 1.7c&d) [129].

Inspired by the remarkable results of MoS2, as was shown by it in the field of energy

storage devices, its potential as ORR catalyst was also investigated. Li et al. prepared

MoS2 nanoparticles of various sizes and tested them as ORR catalysts. They discerned

the effect of particles size on the performance and durability of the synthesized material.

The MoS2 nanoparticles of 2 nm size demonstrated superior ORR performance and

durability than that of other particles, which was credited to the sufficiently uncovered

Mo edges, which are considered advantageous for reduction of O2 to H2O by four-

electron reduction mechanism [130]. Beside the above mentioned MS, Iron sulfide [131]

and copper sulfide [132] have also been recognized as possible ORR catalysts.

d. Solar cells

The direct transformation of solar energy into electricity by photovoltaic (PV) solar cells

is clean and technologically advanced energy source in this century. Silicon solar cells

have been widely scrutinized and applied for hands-on purposes, but its pronounced

manufacturing budget and long energy payback period have restricted their commercial

use for PV power production. Thus, it makes it essential to acquire elevated conversion

efficacy from cheap constituents having reduced manufacturing budget and small energy

ingestion. The transition metal sulfide nanomaterials with efficient optical adsorption,

appropriate band gaps, nonlinear optical properties and little noxiousness have greatly

appealed the consideration of researchers for solar cells applications.

Wu et al. obtained 1.6 % photovoltaic conversion efficiency from Cu2S/CdS

heterojunction solar cells. The material was placed on glass substrates and flexible plastic

substrates, which showed amazing durability (4 months testing period) on both substrates

[133]. CuInS2 NCs is one of the most capable material for PV solar cells due to its direct

21

band gap matching well with the solar spectrum, incredible radiation constancy, great

absorption coefficients (5×105 cm

-1) and little harmfulness [134]. More than 10 %

conversion efficiency has been achieved with this material [135], however the

theoretically predicted conversion proficiency of CuInS2 single junction cells is around

28% [136]. Cu2ZnSnS4 being comprised of earth-abundant and nontoxic constituents with

direct band gap energy of ∼1.5 eV and elevated absorption coefficient (>104 cm-1

) [137]

have been used in PV devices to obtain conversion efficiency of up to 6.7 % [138].

The high optical absorption coefficient and low cost of transition metal sulfide

nanostructures also make them suitable to be used in dye-sensitized solar cells (DSSCs)

[139,140]. The literature also contains reports about investigating FeS, FeS2, NiS and

CoS NPs as photocathodes in tandem solar cells with dye-sensitized TiO2 NPs used as the

parallel photoanode [141-143]. The layered metal sulfide i.e. MoS2 and WS2 have also

Figure 1.8: (a) Cyclic voltammograms of the MoS2, WS2, and Pt electrodes for iodide

species, scanning rate: 10 mV s-1

. (b) Photocurrent density-voltage (J–V)

curves of the DSCs using MoS2 (Blue), WS2 (Red), and Pt CEs (Black)

[146].

been studied for use in photovoltaic cells because of their distinctive electronic features

which are attributable to the wide band edge excitation of the metal centered d-d

transition [144,145]. Wu et al. achieved around 7.73 % and 7.59 % [146] efficiency by

using WS2 and MoS2 as counter electrode materials in dye sensitized solar cells (DSCs)

22

as indicated by their electrochemical results displayed in figure 1.8a &b. Further studies

revealed that the electrocatalytic performance of the counter electrode can be increased

by adding carbon materials to MoS2. The forgoing details suggest that the objective of

reduced price and eminent efficiency solar cell composed of earth rich, nonhazardous

elements can be achieved by using transition metal sulfides nanomaterials and their

binary, alloyed and doped compositions.

3.5.3 Catalysis

a. Photocatalysis

Semiconductor NPs are promising photocatalysis for a number of essential and beneficial

reactions such as organic pollutants degradation, decontamination of air and water,

generation of green and renewable fuel (H2) and organic transformation. Transition metal

sulfides are the most promising and extensively studied materials for photocatalysis due

to their visible or UV-light absorbing capability, arrangement of electronic structure,

charge transport characteristics and electron-hole recombination time. Cadmium sulfide

is an important semiconductor and has been the focus of considerable interest due to

direct and narrow band gap (2.42 eV) and nonlinear optical properties [147]. Owing to

these characteristics it can be checked as a visible-light-driven photocatalyst for

instigating photochemical reactions [148,149]. In the previous decade various types of

CdS based nanocomposite materials have been synthesized and studied as photocatalyst

such as CdS/Pt [150] CdS/TiO2 [151] CdS/MoS2 [152] CdS/CdSe hetero nanorods [153]

and so on. Recently, Gong et al. obtained a high value (1.12 mmol h-1

) of H2 production

under visible light by using CdS clusters decorated graphene nanosheets as a

photocatalyst, which is nearly 4.87 times greater as compared to pure cadmium sulfide

NPs. This remarkable photoactivity was credited to the presence of graphene which

prolonged the recombination time of the photogenerated charge carriers from CdS by

their efficient collection and transportation [154]. The suitable band gap in the visible

region, little toxicity and decent chemical stability of group II, III and VI ternary metal

sulfides i.e. ZnIn2S4 NCs (Eg = 2.3 eV), enabled them to be used as solar light assisted

photocatalyts for water reduction [155]. Consequently, Liu et al. obtained even better

efficiency and photo activity for hydrogen production by using Pt co-catalyst with

23

ZnIn2S4 floriated microspheres [156]. To examine the influence of other metal sulfides

on the activity of Pt loaded ZnIn2S4 for hydrogen production Shen et al. synthesized and

investigated a chain of Pt-loaded MS/ZnIn2S4 (MS = CuS, , CoS, Ag2S, NiS, SnS and

MnS) photocatalysts. They observed that the photocatalytic activity of ZnIn2S4 was

lowered by loading CoS, NiS, or MnS whereas CuS, Ag2S, SnS loading boosted the

photocatalytic performance. The highest activity for hydrogen production was

accomplished by charging 1.0 wt.% of copper sulfide together with 1.0 wt.% platinum

on ZnIn2S4, which was estimated up to 201.7 μmolh-1

, 1.6 times greater than that of

ZnIn2S4 loaded with 1.0 wt.% platinum only. The better performance was attributed to

the dual co-catalytic role of CuS and Pt [157].

Figure 1.9: Band structure alignments of the CdS–ZnS core-shell structure and schematic

of the photoexcited charge carrier distribution and related photocatalytic

reactions [158].

Xie et al. proposed a unique charge transfer mechanism for hydrogen evolution (729

μmol/h/g) from water by CdS-ZnS core-shell particles. In terms of band positions, the

photoexcited electrons and holes in the CdS core should not transfer to the ZnS shell

owing to its higher CBE (conduction band energy) and lower VB (valence band) position.

Nevertheless, the spatial transfer of holes from photoexcited CdS to the ZnS is allowed

due to the existence of acceptor states within the band-gap (Figure 1.9), analogous to the

24

mechanism of charge transfer in DSSCs, thus facilitating charge separation and higher

photocatalytic activity [158]. Hierarchically porous ZnIn2S4 microspheres also have been

shown to be an effective visible light driven photocatalyst for the degradation of toxic

organic dyes such as methylene blue, methyl orange, Congo red, and rhodamine B

[159,160]. In light of the above discussion the suitable band gaps, low toxicities and high

stabilities of transition metal sulfides and their alloys are strong candidates to be used as

photocatalysts to initiate photochemical reactions.

b. Electrocatalysis (Electrochemical CO2 reduction)

The combustion of fossil fuels as energy source is causing a sharp rise in atmospheric

CO2 level, which increased from δ = 270 ppm in pre-industrial age [161] to δ = 401.3

ppm by July 2015, which is high enough the safety limit of δ = 350 ppm for atmospheric

CO2 [162]. The current atmospheric CO2 level are at the uppermost level since records

began and is perhaps the maximum in approximately 15 million years based on Boron

isotopes ratio in planktonic shells [163]. Further increase in greenhouse CO2 level in

atmosphere may cause catastrophic climate change, which outspread further than climate

to ocean acidification for the reason that the latter is a main sink for atmospheric CO2

[164].

To cope with energy demands of ever-increasing global population and environmental

threats owed to CO2 emission, an ultimate way out is to transform atmospheric CO2 into

fuels i.e., hydrocarbons (HCs) and Methanol or fuel precursors such as CO/H2 (syn gas)

[165]. Implementing this approach will not only reduce our dependency on fossil fuels

but will also prevent the external addition of CO2 from fossil fuels, as the carbon cycle is

closed by recycling CO2 into fuels and commodity chemicals. The various CO2 reduction

tactics which have been anticipated and extensively explored in the past decade include

photochemical, biochemical, thermochemical and electrochemical methods [166-169].

Among these methods electrochemical and photo-electrochemical reduction of carbon

dioxide is considered as appropriate approach to store the solar energy in chemical bonds

of portable molecular fuels [170]. Additionally the present energy production and

transportation infrastructure could be used with slight modification [171].

25

Currently, the practical application of CO2 capture, conversion and utilization is hindered

by certain barriers which include,

(1) CO2 capture, separation, purification, and transportation to consumer spots are costly

processes.

(2) Chemical/electrochemical conversion of CO2 requires a high energy input.

(3) Lack of investment incentives and limitations in market size.

(4) Insufficiency of industrial obligation to develop CO2-based chemicals.

(5) Poor socio-economic driving forces [172,173].

Regardless of such challenges, the scientific community still recognizes CO2 seizure,

conversion and consumption as a reasonable and auspicious front-line area for the

purpose of energy exploration and environmental stability.

The conversion of CO2 using electrochemical catalysis, in this context has received great

consideration in latest years due to numerous advantages:

(1) The electrode potentials and reaction temperature can be used to control the process.

(2) The overall chemical intake can be decreased to only water or wastewater as the

supporting electrolytes can be completely reused.

(3) For driving the process the energy can be attained from wind, hydroelectric, sun, tidal,

geothermal, and thermoelectric processes devoid of producing any new CO2.

(4) The electrochemical reaction systems are modular, compact, on-demand and easy to

scale up [174].

Nonetheless certain obstacles still need to be overcome such as sluggish kinetics of CO2

electroreduction even at higher electrode reduction potential in the presence of

electrocatalyst, small energy efficiency owed to parasitic or decomposition reaction of the

solvent at eminent reduction potential and great energy intake.

According to 2008 Bell/DOE report the biggest challenge in CO2 electro reduction into

energy bearing products is low performance electrocatalyst (i.e. small catalytic activity

and poor durability) [175]. This background exemplifies the challenges that must be

addressed.

Electroreduction of CO2 is classified into homogeneous and heterogeneous reactions. The

former is usually comprised of organics or metal–organic molecules as electrocatalyst,

which has distinctive active centers for CO2 interaction and its solution is prepared in

26

electrolytes [165, 176]. In this case the electrocatalyst often demonstrate wonderful

selectivity to CO2 reduction due to their unique molecular structure, thus have attracted

substantial consideration and have widely studied since the 1970s. But the high budget,

noxiousness, and complex separation separation process after use hindered their practice

at industrial scale.

In heterogeneous electrocatalytic reduction of CO2, the reaction usually takes place at the

electrode-electrolyte interface, whereas the electrocatalyst is a solid electrode whereas the

electrolyte is normally CO2 saturated aqueous solution. Typically, this process is

composed of three key steps,

i) CO2 adsorption (chemical) on electrocatalyst.

ii) Electron conduction and/or proton movement to split C–O links and/or form C–H

bonds.

iii) Configuration rearrangement of products to desorb them from the electrocatalyst

surface and transfer into the electrolyte [177].

Researchers have developed more interest in inorganic heterogeneous electrocatalyst

recently, on account of their simplistic fabrication, environmental safety, excellent

productivity, and great prospective for commercial applications.

Electrochemical reduction of carbon dioxide is multi-step process commonly proceed

through two-, four-, six-, or eight-electron reduction machanism [178]. The kinetics of

CO2 electroreduction involves complex mechanisms and sluggish reaction rates, even in

the presence of electrocatalyst. The first step involves the formation of key intermediate

CO2•-

through single electron flow to CO2 molecule. Owing to the huge energy

requirements for linear CO2 molecule to structurally rearrange into bent radical anion, a

single electron reduction of CO2 occur at high negative potential (-1.90V versus SHE,

Reaction R1) Nonetheless the proton coupled multielectron reactions take place at

relatively moderate potentials (Reaction R2-R7) [88]. In general, the reduction products

are combination of carbon compounds, including carbon, methane (CH4), carbon

monoxide (CO), formate (HCOO−) or formic acid (HCOOH), ethane (C2H6),

formaldehyde (CH2O) ethene (C2H4 ), ethyl alcohol (C2H5OH) and methyl alcohol

(CH3OH), whose formation depends on type and selectivity of electrocatalyst and applied

electrode potential [171, 179].

27

CO2 + e- → CO2

•- E

o = -1.9 V (R1)

CO2 + 2H+ + 2e

- → CO + H2O

E

o = -0.52 V (R2)

CO2 + 2H+ + 2e

- → HCOOH

E

o = -0.61 V (R3)

CO2 + 4H+ + 4e

- → HCHO + H2O

E

o = -0.51V (R4)

CO2 + 6H+ + 6e

- → CH3OH + H2O

E

o = -0.38 V (R5)

CO2 + 8H+ + 8e

- → CH4 + 2H2O

E

o = -0.24V (R6)

CO2 + 12H+ + 12e

- → C2H4 + 4H2O

E

o = -0.34 V (R7)

2H+ + 2e

- → H2

E

o = -0.42 V (R8)

As displayed in reactions (R2-R8) the equilibrium potentials of CO2 reduction are

thermodynamically close to that of HER. This means that HER is a major side reaction

go together with electroreduction of CO2 in aqueous electrolytes. Furthermore, the minor

thermodynamic potential variances for reduction yields (Reactions R2-R7) designate the

hurdle of reducing CO2 to the required compound with better selectivity [180]. The above

discussion make it indispensable to find promising electrocatalysts to realize cost

effective CO2 electroreduction at lesser overpotential with pronounced current density,

decent durability and selectivity. The CO2 electroreduction studies in the past several

decades have been devoted on research and development of electrocatalysts to conquer

the mentioned challenges [181,182].

The various materials studied as heterogeneous electrocatalyst for CO2 reduction include

metals, metal alloys, metal chalcogenides, metal oxides and carbon based materials,

whose performance has been recently reviewed by us and many others [178, 183,184].

Among these materials transition metal sulfides are recently attracting the attention of

researchers to investigate these for electrochemical reduction of CO2. Hara et al. in 1997,

noted a marked change in product selectivity by treating metallic electrodes (Zn, Ni, Fe,

Cu, Ag, Pd, and Pt) with aqueous solution of Na2S during the electrochemical reduction

of CO2 [185]. The authors made a suggestion on the bases of current–potential curve that

the adsorbed S2−

on metal electrodes surface prevent the CO adsorption. However, the

assumption was made solely on the basis of electrochemical investigation lacking further

characterization of the tested metallic electrodes. The Probable reason for product

selectivity may be the partial oxidation of metal surface into sulfides due to the Na2S

treatment. The change in color of the electrodes particularly Ag and Cu electrodes whose

28

surfaces transformed to black and gray respectively give the direct evidence for such

behavior. Thus it made it essential to devote more efforts to investigate TMC for

electrocatalysis. Motivated by the statistic that MoS2 show comparable water splitting

performance to Pt catalysts [186], Asadi et al. investigated and obtained outstanding

CO2-coversion activity in terms of the current density (65 mA cm−2

at −0.764 V vs RHE),

FE (ca. 98% for CO at −0.764 V vs RHE), overpotential (ca. 54 mV), selectivity and

stability on the Mo-terminated boundaries of molybdenum sulfide in EMIM-BF4 [187].

These results were better than silver NPs and the bulk silver, and also certified the

hypothetical expectation by Chan et al. [188]. Additionally, density functional theory

(DFT) scheming exposed that higher catalytic activity can be endorsed to the small work

function (3.9 eV), metallic character and greater d-electron density on the Mo-terminated

edges of MoS2. In order to ratify their understanding, they then fabricated perpendicularly

aligned MoS2 nanosheets with additional Mo atoms on the boundaries, and observed

further enhancement in the current density of CO2 reduction (i.e. 130 mA cm−2

at −0.764

V vs RHE). On the basis of these theoretical and experimental results the same group

synthesized a series of nanostructured transition-metal dichalcogenides (TMDs) i.e.

MoS2, MoSe2, WS2, and WSe2 nanoflakes (NFs) (representative WSe2 shown in Figure

1.10a) by chemical vapor transport technique. The as synthesized nanoflakes were

assessed for electrochemical CO2 reduction in aqueous electrolyte composed of 50 vol%

EMIM-BF4 [189]. Among the TMD-based NFs, WSe2 NFs terminated with W atoms

(Figure 1.10b) exhibited superior activity in CO2 reduction to CO (Figure 1.10c), with

24% FE and current density of 18.95 mA cm−2

. The turnover frequency of 0.28 s−1

was

calculated for CO formation at −0.164 V versus reversible hydrogen electrode (RHE)

(overpotential of 54 mV). These amazing results shown by WSe2 are better than formerly

testified bulk MoS2 and other electrocatalysts (Figure 1.10d), which are attributed to large

number of exposed edge sites in the 2D nanoarchitecture, lower charge-transfer resistance

(Rct) and lower work function (3.52 eV). The catalyst also showed outstanding stability

(over 27 h of chronoamperometry), which was accredited to the exceptionally stable W

edge atoms in WSe2 nanoflakes.

Beside electrochemical CO2 reduction transition metals sulfides have also investigated as

OER and HER electrocatalyst to replace precious noble metals. The OER electrocatlytic

29

activity of TMS is discussed in the in the previous section (fuel cells). For HER

electrocatalysis nanostructured MoSx (x ≥ 2) is the most extensively studied and has been

recognized as efficient and low-cost substitutes to platinum for the electrochemical

generation of H2 from water, thus making a promise to meet the urgent demand of

renewable energy [190-192]. A series of other TMS, such as WS2 [193] and MSx (M =

Fe, Co, Ni; x ≥ 2) [194, 195] have also shown promising results for hydrogen evolution

reaction.

Figure 1.10: (a) Scanning Electron Microscopy of WSe2 (scale bar, 2 μm in the main

image and 200 nm in the inset). (b) STEM image of WSe2 (scale bar, 2

nm). (c) Catalytic performance of synthesized TMDCs in 50 vol% EMIM–

BF4 and 50 vol% water electrolyte. (d) Cyclic voltammetry (CV) curves for

WSe2 NFs, bulk MoS2, Ag nanoparticles (Ag NPs), and bulk Ag in CO2

environment. Inset shows the current densities in low overpotentials [189].

30

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49

Chapter 2

EXPERIMENTAL

2.1 Materials

Chemicals used in the synthetic and other reactions were of high purity and used without

further purification. NaHCO3 (ACS grade), NaH2PO4 (AR grade), Cd(NO3)2 ∙ 4H2O,

Hg(NO3)2 ∙ H2O, Zn(NO3)2 ∙ 6H2O, Congo red (M.W= 696.67, C32H24N6O6S2Na2) were

obtained from Merck and natural graphite (crystalline, 300 mesh) from Alfa Aesar. 1-

(Diphenylmethyl)piprazine, 1-benzylpiperazine, CS2 and NaOH were purchased from

Fluka. Polyethylenimine (PEI, Mn = 10,000, branched), polyethylene glycol (PEG, Mn =

400), polydiallyldimethylammonium chloride (PDDA, 20 wt% in water, Mw = 100,000 ~

200,000), (NH4)2MoS4 (99.97% trace metals basis), 4-benzylpiperadine, 1-(2-

hydroxyethyl)piperazine, 4,4′-trimethylendipiperadine, methanol, p-nitrophenol, sodium

borohydride, and ethylenediamine were acquired from Sigma-Aldrich. Graphene oxide

(GO) was prepared from natural graphite applying the procedure described by Hummers

et al [1]. Phosphate buffer was prepared by adding 5 M NaOH solution to 0.1 M

NaH2PO4 solution to adjust the pH to 6.81. All the aqueous solutions used in

electrochemical measurements were prepared in deionized water.

2.2 Instrumentation

Melting points of the synthesized complexes were determined by Gallenkamp (UK)

electrothermal melting point apparatus in capillary tube while elemental analysis were

undertaken on CHNS analyzer 932 Lesco, USA. Raman spectroscopy was undertaken by

Renishaw InVia Microscope with a 532 nm laser source and IR spectra were obtained

with Thermoscientific-6700 FT-IR spectrophotometer and Digilab (7000) Stingray

Imaging FT-IR Spectrometer. 1H and

13C NMR were documented on a Bruker 300 MHz

instrument in which TMS was used as internal reference. The coupling constant (J)

values and chemical shifts in the 1H NMR are presented in Hz and ppm respectively.

Furthermore the multiplicities of signals are specified with chemical shifts (s = singlet, d

= doublet, t = triplet, m = multiplet). The X-ray single crystal analysis was done on a

Bruker Kappa APEX-II CCD diffractometer equipped with a CCD detector set at 40.0

50

mm as of the crystal. The various intensities were determined from a sealed ceramic

diffraction tube (SIEMENS) having graphite monochromatic Mo-Kα radiation and the

structure was solved by applying Patterson and DIRDIF procedure. Further refinement on

F2 was conceded by full matrix least squares by means of SHELXL-97.

The PXRD data was collected either by a Bruker D8 discover Germany or by Philip X-

ray powder diffractometer (Cu Kα radiation). Scanning electron microscopy (SEM) and

energy dispersive spectroscopy (EDS) of the NPs were undertaken either using SEM

JEOL Model JSM6510 or by FEI Nova NanoSEM 450 FEG SEM instrument equipped

with Bruker Quantax 400 X-ray analysis system. The morphology and size of the NPs

were investigated either on a Philips EM208 transmission electron microscopy or on FEI

Tecnai G2 T20 TWIN TEM Instrument. Shimadzu 1800 spectrophotometer was

employed for UV/Visible absorption spectra of NPs and to investigate the progress of CR

dye degradation with time. Steady-state and time-resolved photoluminescence (PL) of

some of the NPs were commenced by employing standard instrument, Flau Time 300

(FT-300) steady state and life-time spectrometer, PicoQuant GmbH, Germany. The PL

was revealed following pulsed LED laser excitation source, PLS-300, whose pulses are

fixed at 305 nm with full width half maximum (FWHM) of ~416 ps and pulse energy

0.077 pJ. The measured tails of the PL kinetics were fitted by using Easy Tau and FluoFit

software. Gas chromatography (GC) was accomplished with an Agilent 7820. A gas

chromatography system and an Agilent HP-5 column with dimensions of 30 m x 0.32 µm

and film thickness of 0.25 µm or a SGE 25QC3/BP20 WAX column with dimensions of

25 m x 0.32 µm and film thickness of 1.0 µm. The retaining period was equated with

already identified compounds. CHI 760E electrochemical workstation (CH Instruments,

Austin, Texas, USA) was used for performing the electrochemical experiments at

ambient temperature (22 ± 2 °C).

2.3 Fabrication of rGO-PEI-MoSx electrode

The whole procedure for fabricating the rGO-PEI-MoSx electrode is schematically

illustrated in figure 2.1. Firstly 10 mg graphene oxide (GO) powder was dispersed

in 10 mL water and pH was set near 9.0 by dropwise addintion of right volume of

aqueous NaOH solution (5 M). The as prepared solution was sonicated for at least

51

1 hour and then 500 µL of a 5 wt% polyethylenimine (PEI) aqueous solution was

added quickly and sonicated for another 30 minutes. Afterwards, it was stirred at

60 °C for 12 hours to form the GO-PEI solution. In order to obtain a well-

dispersed solution, centrifugation at 3000 rpm was applied for 20 minutes to

remove any large aggregates. Then 4 μL of GO-PEI solution was casted onto a

glassy carbon electrode (GCE) to form the GO-PEI modified GCE, which was then

put under infrared lamp for drying. The as fabricated electrode was then dipped

into a solution composed of 2 mM (NH4)2MoS4 and 0.1 M phosphate buffer for

electrodeposition of MoSx by cycling the potential from 0.3 V to -1.3 V (vs.

Ag/AgCl (3 M KCl)) at a scan rate of 50 mV s-1

for 10 times. During this process,

GO is electroreduced to

Figure 2.1: Schematic illustration of the fabrication of rGO-PEI-MoSx composite

on the GCE.

52

reduced graphene oxide (rGO), while MoSx is electrodeposited onto the PEI-rGO.

RGO-PEI, rGO-MoSx was prepared by a similar procedure. In some experiments

the PEI was replaced by polyethylene glycols (PEG) and

polydiallyldimethylammonium chloride (PDDA) and similar method was adopted

for the fabrication of rGO-PEG-MoSx and rGO-PDDA-MoSx electrodes.

2.4 Cyclic voltametric measurements and electrolysis

Cyclic voltammteric (CV) measurements of all electrodes were carried out in N2 or

CO2 saturated NaHCO3 (0.5 M) electrolyte. The saturation of the solution was

carried out by aerating N2 or CO2 for minimum of 30 minuutes in a conventional

three-electrode cell containing an Ag/AgCl (3 M KCl) reference electrode, a Pt

wire counter electrode and a modified or unmodified glassy carbon working

electrode (GCE, 3 mm in diameter, CH Instruments, USA). All data were recorded

after 4 CV cycles. In some experiments for determining the capability of the

catalyst on a high surface area electrode, a slice of carbon cloth (1 x 0.5 cm)

modified with rGO-PEI-MoSx was used as working electrode. A gas-tight two-

compartment electrochemical cell was engaged for performing electrolysis whose

compartments were separated with a glass frit. Each compartment was filled with

12.5 mL electrolyte leaving nearly 10 mL headspace. In this experiment a carbon

cloth of 1 cm x 1.5 cm dimensions was employed as counter electrode, a modified

glassy carbon plate (0.4 cm x 1.5 cm) as working electrode and Ag/AgCl (3 M

KCl) as reference electrode. The cell was degassed afore electrolysis by frothing

CO2 gas for about 30 minutes. Stirring of the solutions was continued throughout

electrolysis. Once the electrolysis was completed, a slight fraction of the cell's

head space (200 µL) was taken out by gas-tight syringe and analysed by Gas

chromatography.The durability experiment was conducted by Chronoamperometry

at a constant potential of -0.65 V with CO2 bubbled into both compartments

continuously and the cell semi-covered. All potentials were changed to the

reversible hydrogen electrode (RHE) reference scale using the equation (2.1).

E (vs RHE) = E (vs Ag/AgCl) + 0.210 V + 0.0591 V x pH (2.1)

53

All the working electrodes were polished using 0.3 µm aqueous Al2O3 slurry on a

polishing cloth, sonicated in water and acetone, rinsed with acetone, and then dried

under a flow of nitrogen before use.

2.5 Calculation of the weight difference of bare, before and after electrolysis of

rGO-PEI-MoSx modified glassy carbon plate electrode

The weight difference of the rGO-PEI-MoSx modified glassy carbon plate electrode

before and after electrolysis was measured by a high-precision micro balance (Sartorius,

max weight 31 g, d = 0.001 mg) after careful rinsing and drying. The data are shown in

table 2.1. The electrolysis was conducted in 0.5 M aqueous NaHCO3 solution at -0.65 V

(vs. RHE) for 60,000 s.

Table 2.1: Weight difference of the rGO-PEI-MoSx modified glassy carbon plate

electrode before and after electrolysis.

Bare electrode

(mg)

Before electrolysis (mg) After electrolysis (mg) passed Q (C)

818.763 818.910 818.896 672.2

2.6 General synthetic procedure of ligand salts and their complexes

In this work, seven sodium dithiocarbamates (L1-L7) were synthesized. Among them

L1-L5 were allowed to react with Cd(NO3)2 and Hg(NO3)2 (Scheme 2.1) while L1-L3,

L6 and L7 with Zn(NO3)2 (Scheme 2.2) to obtain fifteen new metal(II) dithiocarbamates.

Typically, the synthesis of the ligands (L1-L4, L6, L7) was conceded by reacting

substituted piperazine/piperadine dissolved in methanol with equimolar aqueous solution

of NaOH, which was added dropwise to substituted piperazine/piperadine solution and

stirred for 1 hour. To the reaction mixture was then added equimolar methanolic solution

of CS2 at 0 oC in a dropwise manner. The reaction mixture was further stirred for four

hours. The precipitate of the ligand salt was obtained after rotary evaporation of solvent,

washed with diethyl ether and dried under IR lamp. The corresponding ligand salt in

distilled water was treated with aqueous metal nitrate (Zn, Cd, Hg) in 2 : 1 molar ratios

54

except L5 (1 : 1 ratio). Mixing of reactants abruptly resulted in the formation of the

precipitate in the medium, whose stirring was continued for 4 hours at ambient

temperature (scheme 2.1). After filtering and washing, the precipitate of the complexes

was placed under IR lamp for drying. These complexes were recrystallized from DMSO

(HgL1, HgL2, ZnL6 and ZnL7) and CHCl3:CH3OH (remaining complexes) mixture.

Scheme 2.1: Synthesis of ligands (L1-L5) and their complexes {Cd (CdL1-CdL5) and

Hg (HgL1-Hg5)}. Z = N (L1-L3) and C (L4 and L5).

2.7 Synthesis of Mercury(II) Cadmium(II) and Zinc(II) Complexes

Bis(4-benzhydrylpiperazine-1-carbodithioate-κ2 S, Sʹ)Mercury(II) (HgL1)

Amount of reagents used: (1 g, 2.84 mmol) Sodium 4-benzhydrylpiperazine-1-

carbodithioate and (0.488 g, 1.42 mmol) Hg(NO3)2 ∙ H2O. Yield: 0.85 g, 74 %. m.p. 232 -

233 °C. Anal. Calcd (Found) for C36H38HgN4S4 (Mol. Wt.: 855.56 g mol-1

): C, 50.54

(50.50); H, 4.48 (4.51); N, 6.55 (6.53); S, 14.99 (14.98) %.

55

Bis (4-benzylpiperazine-1-carbodithioate-κ2 S, Sʹ)Mercury(II) (HgL2)

Amount of reagents used: (1 g, 3.6 mmol) Sodium 1-benzylpiperazine-1-carbodithozate

and (0.62 g, 1.8 mmol) Hg(NO3)2 ∙ 2H2O. Yield: 1g, 77%. m.p. 202-203 °C. Anal. Calcd

(Found) for C48H60Hg2N8S8 (Mol. Wt.: 1406.74): C, 40.98 (40.95); H, 4.30 (4.34); N,

7.97 (7.96); S, 18.24 (18.25) %.

Bis (4-hydroyethylpiperazine-1-carbodithioate-κ2 S, Sʹ)Mercury(II) (HgL3)

Amount of reagents used: (1 g, 4.4 mmol) Sodium 1-hydoxyethylpiperazine-1-

carbodithioate and (0.75 g, 2.2 mmol) Hg(NO3)2 ∙ 2H2O. Yield: 1g, 75%. m.p. 272-273

°C. Anal. Calcd (Found) for C28H52Hg2N8O4S8 (Mol. Wt.: 1222.46): C, 27.51 (27. 50); H,

4.29 (4.28); N, 9.17 (9.15); S, 20.98 (20.96) %.

Bis(4-benzylpiperadine-1-carbodithioate-κ2 S, Sʹ)Mercury(II) (HgL4)

Amount of reagents used: (0.54 g, 2 mmol) Sodium 1-benzylpiperadine-1-

carbodithozate and (0.342 g, 1 mmol) Hg(NO3)2 ∙ 2H2O. Yield: 0.55 g, 78%. m.p. melts

and decompose at 188 oC. Anal. Calcd (Found) for C52H64Hg2N4S8 (Mol. Wt.: 1402.79):

C, 44.52 (44.49); H, 4.60 (4.62); N, 3.99 (4.00); S, 18.29 (18.28) %.

56

Scheme 2.2: Synthesis of Zn(II) dithiocarbamate complexes.

Propane-1,3-diyl bis (piperidincarbamodithioate)Mercury(II) (HgL5)

Amount of reagents used: (1 g, 2.4 mmol) of sodium propane-1,3-diyl bis

(piperidincarbamodithioate) and (0.82 g, 2.4 mmol) Hg(NO3)2 ∙ 2H2O. In this case the

ligand to metal ion ratio was kept in 1:1 due to the tetra dentate coordinating mode of

sodium propane-1,3-diyl bis (piperidincarbamodithioate). Yield: 0.56 g, 80%. m.p.

decomposes from 295-298 oC. Anal. Calcd (Found) for (C15H24HgN2S4)n {Mol. Wt.:

(561.21)n}: C, 32.10 (32.08); H, 4.31 (4.33); N, 4.99 (4.98); S, 22.85 (22.83) %.

57

Bis(4-benzhydrylpiperazine-1-carbodithioate-κ2 S, Sʹ)Cadmium(II) (CdL1)

Amount of reagents used: (0.7 g, 2 mmol) Sodium 4-benzhydrylpiperazine-1-

carbodithioate and (0.308 g, 1 mmol) Cd(NO3)2 ∙ 4H2O. Yield: 0.60 g, 78 %. m.p. 290-

292 °C. Anal. Calcd (Found) for C36H38CdN4S4 (Mol. Wt.: 767.38): C, 56.35 (56.32); H,

4.99 (5.05); N, 7.30 (7.28); S, 16.71(16.72) %.

Bis (4-benzylpiperazine-1-carbodithioate-κ2 S, Sʹ)Cadmium(II) (CdL2)

Amount of reagents used: (0.55 g, 2 mmol) Sodium 4-benzylpiperazine-1-

carbodithioate, (0.308 g, 1 mmol) Cd(NO3)2 ∙ 4H2O. Yield: 0.46 g, 76 %. m.p. 320-321

°C. Anal. Calcd (Found) for C48H60Cd2N8S8 (Mol. Wt.: 1230.39): C, 46.86 (46.83); H,

4.92 (4.94); N, 9.11(9.11); S, 20.85 (20.86) %.

Bis (4-hydroyethylpiperazine-1-carbodithioate-κ2 S, Sʹ)Cadmium(II) (CdL3)

Amount of reagents used: (0.92 g, 4 mmol) Sodium 1-hydoxyethylpiperazine-1-

cabodithioate and (0.616 g, 2 mmol) Cd(NO3)2 ∙ 4H2O. Yield: 0.85 g, 80%. m.p. 287 °C.

Anal. Calcd (Found) for C28H52Cd2N8O4S8 (Mol. Wt.: 1046.11): C, 32.15 (32.14); H,

5.01 (5.05); N, 10.71 (10.71); S, 24.52 (24.51) %.

58

Bis(4-benzylpiperadine-1-carbodithioate-κ2 S, Sʹ)Cadmium(II) (CdL4)

Amount of reagents used: (1.63 g, 6 mmol) Sodium 1-benzylpiperadine-1-

carbodithozate and (0.92 g, 3 mmol) Cd(NO3)2 ∙ 4H2O. Yield: 1.45 g, 79%. m.p. 194-195

oC. Anal. Calcd (Found) for C52H64Cd2N4S8 (Mol. Wt.: 1226.43): C, 50.92 (50.90); H,

5.26 (5.29); N, 4.57 (4.55); S, 20.92 (20.89) %.

Propane-1,3-diyl bis (piperidincarbamodithioate)Cadmium(II) (CdL5)

Amount of reagents used: (1.2 g, 3 mmole) sodium propane-1,3-diyl bis

(piperidincarbamodithioate) and (0.92 g, 3 mmol) Cd(NO3)2 ∙ 4H2O. Yield: 1.1 g, 78%.

m.p. above 300 oC. Anal. Calcd (Found) for (C15H24CdN2S4)n {Mol. Wt.: (473.04)n}: C,

38.09 (38.07); H, 5.11 (5.14); N, 5.92 (5.90); S, 27.11 (27.12) %.

Bis(4-benzhydrylpiperazine-1-carbodithioate-κ2 S, Sʹ)Zinc(II) (ZnL1)

Amount of reagents used: (1.05 g, 3 mmol) Sodium 4-benzhydrylpiperazine-1-

carbodithioate and (0.45 g, 1.5 mmol) Zn(NO3)2 ∙ 6H2O. Yield: 0.84 g, 79 %. m.p. 201-

202 °C. Anal. Cal. (Found) for C36H38N4S4Zn (Mol. Wt.: 720.36): C, 60.02 (59.98), H,

5.32 (5.35), N, 7.78 (7.77), S, 17.80 (17.77) %.

59

Bis (4-benzylpiperazine-1-carbodithioate-κ2 S, Sʹ) Zinc(II) (ZnL2)

Amount of reagents used: (1.1 g, 4 mmol) Sodium 4-benzylpiperazine-1-carbodithioate

and (0.616 g, 2 mmol) Zn(NO3)2 ∙ 6H2O. Yield: 0.8 g, 70 %. m.p. 188-189 °C. Anal. Cal.

(Found) for C48H60N8S8Zn2 (Mol. Wt.: 1136.34): C, 50.73 (50.72); H, 5.32 (5.36); N,

9.86 (9.87); S, 22.57 (22.55) %.

Bis (4-hydroyethylpiperazine-1-carbodithioate-κ2 S, Sʹ)Zinc(II) (ZnL3)

Amount of reagents used: (0.91 g, 4 mmol) Sodium 1-hydoxyethylpiperazine-1-

cabodithioate and (0.616 g, 2 mmol) Zn(NO3)2 ∙ 6H2O. Yield: 0.76 g, 80%. m.p. 237-238

°C. Anal. Cal. (Found) for C28H52N8O4S8Zn2 (Mol. Wt.: 952.06): C, 35.32 (35.29); H,

5.51(5.54); N, 11.77 (11.75); S, 26.94 (26.95) %.

Bis 4-(3-methoxyphenyl)piperazine-1-carbodithioate-κ2 S, Sʹ)Zinc(II) (ZnL6)

Amount of reagents used: (0.87 g, 3 mmol) Sodium 4-(3-methoxyphenyl) piperazine-1-

carbodithioate and (0.45 g, 1.5 mmol) Zn(NO3)2 ∙ 6H2O. Yield: 0.67 g, 74%. m.p. 250-

251 °C. Anal. Cal. (Found), for C48H60N8O4 S8Zn2 (Mol. Wt.: 1200.34): C, 48.03

(47.99), H, 5.04 (5.06), N, 9.33 (9.31), S, 21.37 (21.34) %.

60

Bis 4-(4-methoxyphenyl)piperazine-1-carbodithioate-κ2 S, Sʹ)Zinc(II) (ZnL7)

Amount of reagents used: (0.87 g, 3 mmol) Sodium 4-(4 methoxyphenyl) piperazine-1

carbodithioate and (0.45 g, 1.5 mmol) Zn(NO3)2 ∙ 6H2O. Yield: 0.7 g, 77%. m.p. 290-

291°C. Anal. Cal. (Found), for C48H60N8O4S8Zn2 (Mol. Wt.: 1200.34): C, 48.03 (47.97),

H, 5.04 (5.05), N, 9.33 (9.32), S, 21.37 (21.32) %.

2.8 Synthesis of metal sulfide (Metal = Hg, Cd, Zn) nanoparticles

All the prepared compounds were employed as single source precursors (SSPs) for the

production of their respective metal sulfide nanoparticles (NPs). Typically, 0.4-1 g of the

metal complexes were mixed with 8-15 mL of ethylenediamine (en) taken in a 50 mL

two necked flask and heated slowly till the boiling point of en (118 oC) with constant

stirring, and the reflux was continued for 2 minutes. In each experiment, as soon as the

reflux temperature was achieved the precipitates of metal sulfides of different colors,

depending on the nature and crystalline phase of MS, start to appear (Scheme 2.3 a&b).

Additionally, reaction progress was noted by the formation of PbS suspension when the

evolved gas from the flask was allowed to bubble into aqueous lead nitrate solution. The

metal sulfide thus obtained in the form of precipitate were collected by centrifugation,

repeatedly rinsed with methanol and dried in a furnace at 50 oC for two hours.

61

Scheme 2.3 Proposed thermolysis mechanism of (a) monomeric precursor complexes and

(b) dimeric precursor complexes by en.

a

b

62

2.9 Computational Studies

The chemical reactivity and stability of the complexes HgL1-HgL4 and CdL1-CdL2 were

theoretically calculated using Density Functional Theory (DFT) calculations. DFT is a

computational quantum mechanical modeling which is usually exploited in chemistry

material science and physics, and has developed theoretical study of chemical reactivity

and stability of various species [2]. In the present study, the aforesaid complexes were

optimized at the start with the support of DFT using Gaussian 09 set of program using

B3LYP/LanL2DZ basis set of theory and Frontier molecular orbital (HOMO-LUMO)

energies of the complexes under study were estimated with the identical procedures and

basis set [3]. The molecular properties of the compounds were calculated by the

following formulae, which have been previously reported in literature [5].

Electron affinity (E.A) = -ELUMO

Chemical Potential (μ) = ½ (EHOMO+ELUMO)

Ionization potential (I.P) = -EHOMO

HOMO-LUMO gap (ΔE) = ELUMO-EHOMO

Global Hardness (η) = ½ (ELUMO-EHOMO)

2.10 Photocatalytic test of CdS nanospheres and nanocapsules for the degradation

of Congo red

The photocatalytic activities of the CdS NPs were investigated for degradation of Congo

red (CR) dye at natural pH (pH=6.5) under solar light in unclouded days in October 2015,

in the middle of 10 am and 3 pm (GPS coordinates N=33o43’ 5”, E= 73

o3’38”). All

experiments were performed in the Pyrex glass vessel of 50 mL capacity containing 30

mL of 50 mg/L of CR solution and 0.5 mg/mL of CdS NPs as a catalyst. Before exposure

to sunlight, in each experiment, the suspension was stirred in the absence of light for 0.5

hours to build adsorption equilibrium between the catalyst surface and dye molecules and

change in concentration of the CR due to adsorption on the catalyst surface was measured

by Shimadzu 1800 spectrophotometer scanning from 200-800 nm. Magnetic stirring of

the suspension was continued during exposure in sunlight to ensure the contact between

the catalyst active sites and dye molecules in solution. The sample CdS-1 was evaluated

at three different temperatures, 30 oC (±1), 40

oC (±1) and 50

oC (±1), by carefully

63

optimizing the hotplate temperature under direct sunlight, however, the other sample

(CdS-2) was studied at normal day temperature. The progress of degradation and

decolorization was continuously monitored with the interval of 10 minutes by

withdrawing 4 mL of suspension, followed by centrifugation and measuring the

shrinkage in absorption. The degradation rate of the dye was calculated by the following

equation (2.2).

Degradation rate (%) = 𝟏 −𝑪𝒊

𝑪𝒐× 𝟏𝟎𝟎 (2.2)

Where 𝐶𝑜 and 𝐶𝑖 is the initial concentration of CR dye and sample at time T,

respectively. The kinetics of degradation reaction was articulated by Langmuir-

Hinshelwood model. The expression for reaction rate is given in equation (2.3).

ln(𝑪𝒊

𝑪𝒐) = - kappt (2.3)

Where 𝑡 is the reaction time, kapp is the pseudo first order kinetic constant. Plot of ln(𝐶𝑖

𝐶𝑜)

vs 𝑡 give -kapp [6].

2.11 Photocatalytic reduction of p-nitrophenol to p-aminophenol by CdS nanosheets

and nanorods

The CdS nanosheets and nanorods mediated photoreduction of p-nitrophenol to p-

aminophenol was studied under direct sunlight in May 2016 from 9 am to 3 pm (GPS

coordinates N = 33o43’ 5”, E = 73

o3’38”). Pyrex glass vessel of 50 mL capacity was used

for conducting the experiments. To build the adsorption equilibrium between the p-

nitrophenol and catalyst surface the mixture was stirred in the dark for 30 minutes before

each test. Typically, 0.1g (10 ppm) of p-nitrophenol was dissolved in 100 mL distilled

water. 1.5 mL of this solution was further diluted to 25mL and scanned by Shimadzu

1800 spectrophotometer, then to it 0.003 g of NaBH4 was added, as a result the light

yellow color intensified due to the formation of phenolate ion. To measure the

concentration of the phenolate ion, the solution was scanned again, followed by the

addition of appropriate amount of CdS NPs and then exposed to direct sunlight with

magnetic stirring of the suspension to guarantee the contact between the catalyst active

64

sites and p-nitrophenol. The conversion was spectrophotometrically checked with the

interval of two minutes by withdrawing 2 mL of suspension, followed by centrifugation

and measuring the contraction in the intensity of distinctive peak of p-nitrophenol at 400

nm with parallel growth of new peak at 298 nm corresponding to p- aminophenol [7].

The kinetics of the photocatalytic reduction was calculated using equation (2.2).

65

REFERENCES

1. Hummers, W. S.; Offeman, R. E., Preparation of Graphitic Oxide, J. Am. Chem. Soc.

1958, 80, 1339-1339.

2. Torrent‐Sucarrat, M.; Salvador, P.; Geerlings, P.; and Sola, M., On the quality of the

hardness kernel and the Fukui function to evaluate the global hardness, J. Comput.

Chem. 2007, 28, 574-583.

3. Jian, F. F.; Li, C. L.; Zhao, P. S.; and Zhang, L., Experimental Studies on

Isonicotinato Cadmium (II) Complex [Cd (C6H4NO2) 2 (H2O)4] and Density

Functional Calculations, Struct. Chem. 2005, 16, 529-533.

4. Zhan, C. G.; Nichols, J. A.; and Dixon, D. A., Ionization potential, electron affinity,

electronegativity, hardness, and electron excitation energy: molecular properties from

density functional theory orbital energies, J. Phys. Chem. A. 2003, 107, 4184-4195.

5. Sangeetha, C. C.; Madivanane, R.; Pouchaname, V., The vibrational spectroscopic

(FT-IR & FTR) study and HOMO & LUMO analysis of 6-Methyl Quinoline using

DFT Studies, Arch. Phys. Res. 2013, 4, 67-77.

6. Kaur, S.; Singh, V., TiO2 mediated photocatalytic degradation studies of reactive red

198 by UV irradiation, J. Hazard. Mater. 2007, 141, 230-236.

7. Kong, X. K.; Sun, Z. Y.; Chen, M.; Chen, Q. W., Metal-free catalytic reduction of 4-

nitrophenol to 4-aminophenol by N-doped graphene, Energy. Environ. Sci. 2013, 6,

3260-3266.

66

Chapter 3

RESULTS AND DISCUSSION

This chapter comprises of two sections: (a) electrochemically produced MoSx for CO2

reduction (3.1) and (b) metal sulfide nanoparticles (HgS, CdS and ZnS) from single

source precursors, their formation mechanism and photocatalytic application of CdS in

environmental remediation and organic transformation reaction (3.2-3.6).

3.1 Characterization and electrocatalytic CO2 reduction activities of rGO-PEI-MoSx

In nature, formate dehydrogenases (FDH), which contain molybdenum or tungsten

centers coordinated by a selenocysteine residue and amino acid residues, can catalyze the

oxidation of formate to CO2 [1,2]. Hirst and co-workers [3,4] showed that they could

efficiently and specifically catalyze the reversible transformation between CO2 and

formate under mild conditions. So, we considered it worthwhile to explore new catalysts

resembling the FDH active site for electroreduction of CO2. Molybdenum disulphide

(MoS2), a two-dimensional transition metal dichalcogenide, consists of a planar triangular

lattice of Mo atoms packed in between two planes of S atoms and each Mo (IV) is

coordinated to six S ligands, while each S center is connected to three Mo atoms. From a

biomimic point of view, MoS2 shares a similar structure with the active center of FDH,

making it potentially an electrocatalyst for CO2 reduction. In the present work amorphous

molybdenum sulfide was electrochemically deposited on polyethylenimine modified

reduced graphene oxide (rGO-PEI-MoSx) substrate, characterized by various

spectroscopic techniques and evaluated for electrochemical CO2 reduction.

3.1.1 Characterization of GO-PEI by Fourier Transform Infrared

spectroscopy (FT-IR)

Fourier Transform Infrared (FT-IR) spectroscopic characterization of graphene oxide

(GO) and GO-PEI was undertaken to reveal the nature of the interaction between PEI and

GO. As shown in figure 3.1a, these two materials exhibit distinctly different FT-IR

spectra. Two additional peaks at 2841 cm-1

and 2932 cm-1

were found with GO-PEI,

67

which were assigned to symmetric and asymmetric stretching modes of the CH2 of the

PEI chains [5]. Furthermore, a significant decrease of the peak at 1724 cm-1

(accredited to

C=O from the carboxyl group in pristine GO) and the appearance of a new band at 1632

cm-1

(endorsed to the C=O stretching of the primary amide) propose that the carbonyl

groups are converted to amides on reaction with the PEI [6]. This is further confirmed by

the new bands observed at 1450 cm-1

(C-N stretching vibration) and 1572 cm-1

(N-H

bending vibration) with GO-PEI [7]. This covalent interaction between PEI and GO is

beneficial in maintaining the high stability of the modified electrode under long term

electrolysis conditions. It is important to point out that the band at 1592 cm-1

ascribed to

the C=C vibration of the aromatic rings was observed in both GO and GO-PEI, implying

that the sp2 character in GO-PEI is preserved after the chemical conjugation [8] which is

crucial to preserve the conductivity of the fabricated electrode.

3.1.2 Fabrication and characterization of rGO-PEI-MoSx

The rGO-PEI-MoSx electrode was fabricated by electrodeposition of MoSx onto a rGO-

PEI modified electrode using the detailed procedure described in the experimental

section. The voltammogram (Figure 3.1b) clearly demonstrates that a new reduction

Figure 3.1: (a) FT-IR spectra of GO (black line) and GO-PEI (red line) and (b)

Electrodeposition of rGO-PEI-MoSx on GCE by repeating CVs for 10

cycles.

68

process at 0.1 V (vs. RHE, all potentials henceforward are with respect to this reference)

and a new oxidation process at 0.4 V emerge and their current magnitude increases with

cycling. The oxidation peak is attributed to the oxidation of [MoS4]2-

to form MoS3 and

S8 according to the equation (3.1) [9,10]:

[𝐌𝐨𝐒𝟒]𝟐− → 𝐌𝐨𝐒𝟑 +𝟏

𝟖𝐒𝟖 + 𝟐𝐞− (𝟑. 𝟏)

The reduction peak is assigned to the reduction of the electrogenerated S8. At further

negative potential than -0.25 V, reductive deposition of amorphous MoIV

S2 (i.e. MoSx)

takes place according to the previous electrochemical quartz crystal microbalance studies

(Eq. 3.2) [9].

[𝐌𝐨𝐒𝟒]𝟐− + 𝟐𝐇𝟐𝐎 + 𝟐𝐞− → 𝐌𝐨𝐒𝟐 + 𝟐𝐇𝐒− + 𝟐𝐎𝐇− (𝟑. 𝟐)

In the potential range of -0.4 V to -0.69 V, GO is electroreduced to rGO [11]. Thus, the

deposition of MoSx and reduction of GO occurred simultaneously in this potential region,

which was confirmed by energy dispersive spectroscopy (EDS) and Raman spectroscopy.

Figure 3.2: (a) EDS and (b) Raman spectrum of GO-PEI and rGO-PEI-MoSx.

69

In the EDS spectrum the existence of peaks of Mo and S indicates the successful

deposition of MoSx on GO-PEI (Figure 3.2a). Moreover, the N peak remains after the

electrodeposition, which implies the stability of PEI anchored on GO. The Raman

spectrum show the presence of D and G peaks at 1348 and 1596 cm-1

afterward the

electrodeposition of MoSx (Figure 3.2b), which correspond to the E2g mode witnessed for

sp2 carbon domains and edges and/or defects of the sp

2 domains, respectively [12]. The

intensity ratio (ID/IG) of D and G bands of GO-PEI is around 0.83, whereas the ID/IG of

rGO-PEI-MoSx is 1.21 owed to the occurrence of unrepaired defects that remain

afterwards the subtraction of enormous total of oxygen-containing functional groups.

This value of ID/IG ratio is in agreement with maximum chemical reduction reports

[12,13], suggesting the successful electroreduction of GO to rGO accompanied by the

electrodeposition of MoSx. The presence of hydrophilic PEI on GO could increase the

space between each GO layer and make it more easily accessible for [MoS4]2-

to achieve

efficient deposition of MoSx [14].

Figure 3.3: (a) CVs of bare GCE (—), rGO-PEI (—) and rGO-PEI-MoSx (—) modified

GCE in 0.1 M phosphate buffer (pH 6.81), scan rate of 10 mV s-1

and (b)

TEM image of rGO-PEI-MoSx (Scale bar is 400 nm).

Moreover, both rGO-PEI and rGO-PEI-MoSx modified glassy carbon electrode (GCE)

show a large increase in capacitance compared to bare GCE (Figure 3.3a) due to a

significantly enlarged conductive surface area introduced by rGO [15]. The further

70

enhanced capacitance of rGO-PEI-MoSx arises from pseudo-capacitance processes

associated with MoSx [16]. To confirm the interaction between rGO and MoSx,

transmission electron microscopic (TEM) and scanning electron microscopic (SEM)

images were used to characterize the morphology. As revealed in figure 3.3b, rGO-PEI-

MoSx shows a shining and crumpled texture related with flexible graphene sheets. Figure

3.4a shows the magnified TEM image, and MoSx nanoparticles of diameter 17 ±3 nm can

be clearly seen. The fast Fourier transform (FFT) pattern indicates the amorphous nature

of MoSx nanoparticles (Figure 3.4a inset), which is further confirmed by X-ray

Diffraction (XRD) with no characteristic peaks of any crystalline structures is observed

(Figure 3.4b). Moreover, in the absence of GO-PEI, the MoSx tended to aggregate into

Figure 3.4: (a) TEM images of rGO-PEI-MoSx, the insert is corresponding FFT

pattern, (b) XRD pattern of bare FTO (black) and rGO-PEI-MoSx on

FTO (red), (c) TEM image of electrodeposited MoSx in the absence

of GO-PEI and (d) SEM images of rGO-PEI-MoSx.

71

much larger particles as can be evidenced in figure 3.4c. By contrast, as can be seen from

figure 3.4d the MoSx nanoparticles grew uniformly on the surface of rGO-PEI allowing

the exposure of more active sites of MoSx to CO2.

3.1.3 CO2 reduction activities of rGO-PEI-MoSx

To evaluate the electrocatalytic ability of rGO-PEI-MoSx for CO2 electroreduction, cyclic

voltammetric (CV) measurement were primarily undertaken in 0.5 M aqueous NaHCO3

solutions saturated with N2 and CO2 (pH values 8.5 and 7.2, respectively) at a scan rate of

50 mV s-1

(Figure 3.5a). Under both atmospheres, rapid increases of current were

observed at around -0.22 V. However in the case of CO2 saturated solution, a wide

process also appears, which is centered at -0.65 V. This kind of process was attributed to

CO2 reduction in the previous literature reports, where nitrogen doped carbon nanofibers

[17] or Ag [18,19] were used as catalysts. Since CV cannot provide convincing proof for

the authentication of this reduction process, potentiostatic electrolysis was conducted by

applying a potential in the range of -0.25 V to -0.95 V (corresponding it curve is depicted

in figure 3.5c) and the faradaic efficiency (F.E.) is plotted as a function of the applied

potential in figure 3.5b. The results reveal that hydrogen and carbon monoxide are the

dominant products with a shared F.E. of about 100% over the whole potential range.

Therefore, the liquid part, if any, was not analysed.

The selectivity headed for H2 and CO is intensely reliant on the applied potential. At

applied potential of -0.25 V (corresponding to an overpotential of 140 mV for the

formation of CO, since the equilibrium potential is -0.11 V vs. RHE [20], all the

overpotentials hereafter are reported with respect to this equilibrium potential.), CO with

a F.E. of ~ 5% was identified however H2 is the principal product. This overpotential for

CO formation is significantly lower than that of polycrystalline Au or Ag [20] and as low

as that achieved in oxide-derived Au nanoparticles under the same conditions [21]. F.E.

of CO formation surpasses H2 production when the potential turn into more negative and

it touches the maximum of 85.1% at an applied potential of -0.65 V (overpotential of 540

mV), which exceeds the highest F.E. of the polycrystalline Ag electrode reported in the

72

Figure 3.5: (a) CVs of rGO-PEI-MoSx modified GCE in N2-saturated (black curve) and

CO2-saturated (red curve) 0.5 M NaHCO3 aqueous solution. Scan rate was

50 mVs-1

. Insert: structure of PEI, (b) Faradaic efficiency for CO (red bars)

and H2 (blue bars) as a function of potential and (c) Total current

density vs. time curve during electrolysis at different applied potentials.

same electrolyte medium [20]. H2 reclaims domination in the further negative potential

region. This phenomenon was also witnessed in the case of Ag and Au catalysts in

previous reports [19, 20, 22, 23] and has been described as follows: In the potential

region of low CO2/CO overpotentials, hydrogen evolution reaction (HER) is leading

since it is thermodynamically and kinetically more feasible [20, 24]. In the potential

region of CO2/CO overpotentials (i.e. -0.35 – -0.75 V vs. RHE), CO2 reduction take place

to form CO, and HER is slowed down on account of the slow desorption of CO from the

active sites [25]. In the potential region of high CO2/CO overpotentials, kinetics of CO2

73

reduction as well as HER rises. Nevertheless, HER regains dominance since H2 evolution

is the kinetically preferred process and H+ is much more freely available in the aqueous

HCO3

- medium. Therefore, CO2 reduction reaction progresses relatively slowly due to

either kinetic or mass transport limitations [19, 23, 26, 27].

The source of CO is an important issue that needed to be confirmed. In our study, no CO

was detected when the electrolysis was undertaken in N2 atmosphere, which confirms

that the direct source of CO is not HCO3-. Another possible source could be the catalyst

materials, which include support electrode, rGO and PEI. 13

C labelled CO2 and NaHCO3

could be used to provide direct evidence. However, the formal potential is 0.52 V vs.

RHE for the reaction CO(g) + 2H+ + 2e

- ↔ C(s) + H2O, which is much more positive than

the applied potential (-0.25 – -0.95 V vs. RHE) used. Therefore, direct oxidation of C to

generate CO is not a thermodynamically favorable process. Under the mildly reducing

electrochemical conditions, extensive oxidation of rGO and PEI to produce the amount of

CO detected is also very unlikely. Moreover, the weight difference of the rGO-PEI-MoSx

modified glassy carbon plate electrode was measured before and after electrolysis which

was found 0.014 mg. Assuming that the total lost material is due to the formation of CO

from carbon, then the amount of generated CO should be 1.17 µmol. However, the actual

amount of CO produced is 2.96 mmol (672.2*85.1%/2/96500 based on the Faraday’s

law), which is far more than 1.17 µmol; the rGO-PEI used was 0.157 mg. Even carbon

with the same mass as expected for all of the modified catalyst materials cannot produce

sufficient CO even at 100% conversion yield. Therefore, neither rGO nor PEI can be the

major source of CO.

The intrinsic catalytic activity toward CO2 reduction for each MoSx active site is

indicated by turnover frequency (TOF), calculated using the equation (3.3):

𝐓𝐎𝐅 =𝑰𝐭𝐨𝐭𝐚𝐥𝐅. 𝐄.

𝟐𝐅𝒏 (𝟑. 𝟑)

where Itotal is the current (in A) in CV measurement, F.E. is faradaic efficiency for

CO formation, F is Faraday’s constant (in C mol-1

). The factor 2 arises from the

fact that two electrons are prerequisite to form CO molecule from one CO2

74

molecule. n is active sites number (in mol) and was estimated from voltammetric

measurements. To quantify the total amount of MoSx, freshly prepared rGO-PEI-

MoSx electrode was electrochemically oxidised in 0.5 M H2SO4 under cyclic

voltammetric conditions in the potential region from 0 V to 1.4 V (vs. RHE) at a

scan rate of 50 mV s-1

(Figure 3.6a). The amount of electron (Q) consumed to

oxidise MoIV

and S2-

associated with MoSx to MoVI

and SO42-

, respectively, is

calculated by integrating the area of the oxidation peak, which corresponds to the

overall 18-electron oxidation process as described in equation (3.4) [12, 28].

𝐌𝐨𝐒𝟐 + 𝟏𝟏𝐇𝟐𝐎 → 𝐌𝐨𝐎𝟑 + 𝟐𝐒𝐎𝟒𝟐− + 𝟐𝟐𝐇+ + 𝟏𝟖𝐞− (3.4)

Figure 3.6: (a) Oxidation peak of rGO-PEI and rGO-PEI-MoSx in 0.5 M H2SO4 under

cyclic voltammetric conditions in the potential region from 0 V to 1.4 V (vs.

RHE) at a scan rate of 50 mV s-1

and (b) Linear sweep voltammogram

obtained using a rGO-PEI-MoSx modified carbon cloth (—) as working

electrode in CO2-saturated 0.5 M aqueous NaHCO3 solution. In the absence

of rGO-PEI-MoSx, only a small background current was observed (—).

Scan rate was 5 mV s-1

.

The electric quantity Q is integrated to be 1.08 x 10-3

C. Thus, the amount of MoSx is

6.18 x 10-10

mol based on the Faraday’s law. Taking Itotal = 3.4 x 10-4

A, F.E. = 85.1%

and n = 6.18 x 10-10

mol the TOF at the applied potential of -0.65 V is calculated to be

75

2.4 s-1

. Such a large TOF value implies that this catalyst could potentially be used for

practical applications when it is immobilized on a porous support with a large surface

area. Indeed, a current density as high as 55 mA cm-2

at -0.65V could be reached when

the rGO-PEI-MoSx composite catalyst was immobilized on a carbon cloth electrode

(Figure 3.6b).

It is also worth mentioning that the molar ratio of H2 and CO is about 2:1, which is the

composition of syngas, when the applied potential is -0.4 V (equivalent to 290 mV

overpotential for CO formation; Figure 3.7a). Thus, this produced gas mixture could

readily serve as precursor gas for the Fischer-Tropsch process, an important gas to liquids

technology for generating methanol, synthetic petroleum and other liquid fuels [29].

More importantly, the F.E. of H2 and CO remains constant at this potential over at least 3

hours of electrolysis (Figure 3.7a) showing great promise as a practical catalyst for this

process.

Figure 3.7: (a) Amount and faradaic efficiency of H2 (circles) and CO (squares).

Potentiostatic electrolysis at -0.4 V in CO2-saturated 0.5 M NaHCO3

aqueous solution. Measurement interval: 0.5 h and (b) Chronoamperomet-

ric response of a rGO-PEI-MoSx modified GCE at a constant potential of -

0.65 V. At the 10,000th and 50,000th second, the applied potential was

switched to open circuit potential for 500 seconds before recovering to -

0.65 V.

Moreover, the stability of rGO-PEI-MoSx was also inspected by chronoamperometry at -

0.65 V, which is the optimum potential with respect to maximum CO selectivity. As

76

shown in figure 3.7b, the current residues constant for a period of 16 hours (60,000 s).

Initial decline of current, which cannot be completely endorsed to the double layer

charging process, is mainly due to the adsorption of gaseous products. This current was

largely restored after the electrode was kept under open circuit condition for 500 Seconds

afore applying the same potential for electrolysis, proposing long-term stability of the

rGO-PEI-MoSx. SEM images and EDS of the electrode surface after 60,000 s of

electrolysis are shown in figure 3.8a & b, respectively. Compared with SEM image taken

before the electrolysis (Figure 3.4d), not much difference in morphology was observed

after the long-term electrolysis experiment. The elemental analysis reveals the

preservation of C, N, Mo and S elements, which also indicates the stability of the

material.

Figure 3.8: (a) SEM (Scale bar is 200 nm) and (b) EDS of rGO-PEI-MoSx taken after

60,000 seconds of electrolysis.

In HER catalysis, amorphous MoSx often suffers from activity decay due to mechanical

loss of catalyst materials or poisoning by surface adsorbates from the electrolyte or

reference electrode rather than inherent property changes [30,31]. Chorkendorff et al.

showed that catalyst support, such as carbon fibers, could play a central role in improving

the mechanical stability of the amorphous MoSx catalyst [32]. Therefore, in present case,

77

the presence of the rGO support enhances the mechanical stability of the electrodes under

catalytic turnover conditions, as suggested in regard to other types of support [16, 32-34].

3.1.4 Possible origins of the activity of rGO-PEI-MoSx for CO2 reduction

To reveal the possible origins of the activity of rGO-PEI-MoSx for CO2 reduction, the

CV measurements were undertaken using rGO-PEI, rGO-MoSx and rGO-PEI-MoSx

modified GCEs, respectively. As shown in figure 3.9a, in the absence of MoSx, rGO-PEI

modified GCE demonstrates a reduction onset potential that is about 500 mV more

negative than the other two electrodes in both nitrogen and CO2-saturated 0.5 M aqueous

NaHCO3 solutions which indicates that MoSx is the effective catalyst towards the

reduction reactions observed here. Electrolysis results (Figure 3.9b) show only H2 was

detected with a F.E. of ~93% at -0.85 V. The rGO-MoSx modified GCE presents similar

CV curves in N2 and CO2 atmospheres and potentiostatic electrolysis at -0.65 V reveals

that H2 is the only product with a F.E. of ~95%. Interestingly, although the incorporation

of PEI into rGO-MoSx system results in a decrease in current density under either N2 or

CO2 atmospheres, which will be discussed later, there is a significant difference in

voltammetric characteristics and a considerably higher F.E. for CO at -0.65 V. Hence, a

synergetic effect of PEI and MoSx is responsible for the catalysis of CO2 reduction.

Crystalline MoS2 has demonstrated its capability for effectively electroreducing CO2 in

an imidazolium based ionic liquids (IL) mixed with water as proton source [27]. Similar

observations have also been made previously by the same group in their study of

electrocatalytic reduction of CO2 at an Ag electrode in this medium [35]. They found that

imidazolium cation could act as a co-catalyst via the formation of IL-CO2 complex to

reduce the activation energy barrier. Norskov et al. recently anticipated theoretically by

DFT calculations that the edge sites of crystalline MoS2 could be the catalyzing sites for

CO2 reduction [36]. They claimed that the bridging S atom at the edges of MoS2 could

selectively bind the intermediate HCOO over the product CO, which resulted in a

deviation from the transition-metal scaling relationship of intermediates in CO2 reduction

process and hence significantly improved CO2 reduction activities over transition-

78

Figure 3.9: (a) CVs of rGO-PEI, rGO-MoSx, rGO-PEI-MoSx modified GCE in N2

saturated and CO2-saturated 0.5 M NaHCO3 aqueous solution. Scan rate

was 50 mV s-1

, (b) Faradaic efficiency for CO (red bar) and H2 (blue bars)

of rGO-PEI, rGO-MoSx, rGO-PEI-MoSx, potentiostatic electrolysis at -0.65

V in CO2-saturated 0.5 M NaHCO3 aqueous solution, (c) CVs of rGO-

PDDA-MoSx and (d) rGO-PEG-MoSx modified GCE in N2-saturated and

CO2-saturated 0.5 M NaHCO3 aqueous solution. Scan rate was 50mV s-1

.

Inset of (c) and (d) shows the structure of PDDA and PEG, respectively.

metal catalysts. Although there are no well-defined edges in amorphous MoSx, there are

many structurally and coordinatively unsaturated Mo and S atoms [31,37,38] which are

equivalent to the edge sites in crystalline MoS2 from a catalysis point of view. On this

basis, in principle, the amorphous MoSx used here may be expected to be a better

electrocatalyst for CO2 reduction due to abundance of such active sites. However, in the

absence of PEI, rGO-MoSx modified GCE only showed activity towards HER, which

infers the crucial role of PEI for CO2 reduction. To investigate its role, PEI was replaced

79

with other polymers, polydiallyldimethylammonium chloride (PDDA) or polyethylene

glycol (PEG) for the fabrication of rGO-polymer-MoSx modified electrodes. Figure 3.9c

&d present the CVs of rGO-PDDA-MoSx and rGO-PEG-MoSx modified GCE in nitrogen

and carbon dioxide saturated 0.5 M aqueous NaHCO3 solution at the scan rate of 50 mV

s-1

respectively. Again, both electrodes show similar CV curves in N2 and CO2

atmospheres with the same onset potential and only a slight decrease of current density in

the presence of CO2, and further potentiostatic electrolysis demonstrated H2 as the only

product. Thus, it is clear that the PEI possesses unique properties that produce synergetic

effects with MoSx for CO2 reduction. Two factors may account for this effect: (a) PEI can

suppress HER and thus enhance the competing CO2 reduction: To probe this further, the

HER catalysis properties of rGO-MoSx and rGO- PEI-MoSx were investigated in

aqueous media, containing 0.5 M H2SO4, 0.5 M bicarbonate (pH 8.5) or 0.5 M phosphate

buffer (pH 7.2, at which catalytic CO2 reduction was observed). Results obtained are

shown in figure 3.10a-c. The kinetics of HER reaction were retarded in the presence of

PEI in all media, as suggested by a larger overpotential. This effect of PEI on the HER is

most likely due its amine groups. Rezaei et al. have shown that amine groups in IL

electrolytes can inhibit hydrogen formation in aqueous H2SO4 solution [39]. Rosen et al.

also found this effect when they used an ionic liquid as a co-catalyst for CO2 reduction in

mixed solvents containing water and [EMIM][BF4] [40]. (b) PEI may stabilize the

transitional species in the CO2 reduction process: Numerous researchers have reported a

Tafel slopes of ~120 mV dec-1

in aqueous media at metallic electrodes, which indicated

the one-electron reduction of CO2 to form CO2•–

as the rate determining step [20].

Nonetheless a smaller value of 74 mV dec-1

was achieved with rGO-PEI-MoSx modified

electrode in the present work (Figure 3.10d). In view of the existence of a competing

HER, the Tafel slope obtained in this study is very near to the theoretical value of 59 mV

dec-1

estimated for the case in which a reversible CO2/CO2•–

process take place before a

rate determining chemical step [21, 23, 41]. This enhanced kinetics of the CO2/CO2•–

process was attributed in studies of nitrogen-doped carbon nanotubes, to the stabilization

of CO2•–

by amines on PEI, through hydrogen bonding and electrostatic interactions [42].

80

Figure 3.10: HER polarization curves of rGO-MoSx (MoSx loading 6.65 nmol cm-2

,

estimated from voltammetric measurements) and rGO-PEI-MoSx (MoSx

loading 8.75 nmol cm-2

) modified GCE in (a) 0.5 M H2SO4 (b) 0.5 M

NaHCO3 and (c) 0.5 M phosphate buffer (pH 7.2) solutions, respectively.

Scan rate 5 mV s-1

, with iR compensation. (d) CO production partial

current density vs. overpotential on rGO-PEI-MoSx. The partial current

density of CO is obtained by multiplying the total current density at each

potential by the faradaic efficiency of CO. The Tafel slope for rGO-PEI-

MoSx is 74 mV dec-1

.

3.2 Spectroscopic characterization of metal complexes

Metal(II) dithiocarbamates {ML1-ML5 (M = Cd, Hg) and ZnL1-ZnL3, ZnL6, ZnL7}

were easily synthesized under ambient conditions by mixing the corresponding metal salt

and sodium dithiocarbamate. Interestingly, the Zn (II) and Hg (II) complexes (except

HgL5 which was insoluble at all) showed good solubility in common organic solvents

81

such as CHCl3, CHCl2, acetone, dimethyl sulfoxide (DMSO) and dimethyl formamide

(DMF), however, the Cd (II) complexes were found either partially soluble or insoluble

(CdL5) in the forgoing solvents.

Fourier transform infrared (FT-IR) spectroscopy is a suitable method for obtaining useful

information (Table 3.1) about the bonding nature and coordination mode of the

dithiocarbamates. The presence of bands in between 300-385 cm-1

are assignable to M-S

stretch (M = Zn, Cd, Hg) [43] and single peak for SCS present at 967-1030 cm-1

indicated the bidentate dithiocarbamate coordination to M (M = Zn, Cd, Hg).

Furthermore, the C-N stretch in the range of 1472-1525 cm-1

designate the induction of

double bond features by reason of delocalization of electronic pair of piperazinyl nitrogen

to SCS moiety upon complexation [44].

Table 3.1: FT-IR stretching frequencies of dithiocarbamate complexes

Compound ν (cm-1

) C-N ν (cm-1

) CSS ν (cm-1

) M-S

HgL1 1507 969 317

HgL2 1482 995 300

HgL3 1525 1030 320

HgL4 1473 985 311

HgL5 1522 971 322

CdL1 1434 1014 358

CdL2 1485 996 361

CdL3 1503 1024 370

CdL4 1478 991 347

CdL5 1513 967 334

ZnL1 1489 1003 375

ZnL2 1472 1007 381

ZnL3 1493 1014 385

ZnL6 1496 1010 385

ZnL7 1497 1015 369

82

The allotment of the proton resonances was done by their peak multiplicity, intensity

pattern and comparison of their integration values with the expected composition (Table

3.2). In 1H-NMR spectrum of complexes {ML1, ML2, ML4 (M = Cd, Hg) and ZnL1,

ZnL2, ZnL6 and ZnL7} the phenyl protons gave multiplets in the range of 6.39 ppm-7.44

ppm. However, ML1-ML3 (M = Cd, Hg, Zn) and ZnL6, ZnL7 showed two triplets for

piperazinyl protons. The methine proton (4.30 ppm) of ML1, methylene protons (~3.54)

of ML2 (M = Zn, Cd, Hg) and hydroxyl proton ( 4.52 ppm) of ML3 (M = Zn, Cd, Hg)

gave a singlet. Furthermore, in ML3 hydroxyl connected methylene protons gave two

independent triplets at 3.92 ppm and 2.28 ppm. For ML4 (M = Cd, Hg) a triplet (3.27

and 3.24 ppm), a quartet (1.74 and 1.75 ppm) and a multiplet (1.40-1.51. and 1.38-1.47

ppm) was observed for piperadinyl protons. A doublet appeared for methylene protons of

benzyl moiety at 4.63 and 4.66 ppm. In ZnL6 and ZnL7, methoxy protons observed as a

singlet at 3.73 and 3.80 ppm. The 13

C NMR data (Table 3.3) is in consonance with the

proposed structure for each complex. An upfield shift of 7-11 ppm for the SCS carbon

has confirmed metal-ligand coordination [45].

3.3 X-ray single crystal analysis of ZnL6, ZnL7, HgL1 and HgL2 complexes

3.3.1 Structural descriptions of ZnL6 and ZnL7 Complexes

For complexes ZnL6 and ZnL7, crystal data and refinement parameters are listed in table

3.4 whereas bond lengths and bond angles are shown in table 3.5. The centrosymmetric

and dimeric molecular structure [ZnL2]2 was observed for both ZnL6 and ZnL7, featuring

four dithiocarbamate chelates with two sulfur atoms bridging two Zn atoms resulting in

pentacoordinated geometry as was describe previously [46]. Due to this dimeric nature of

ZnL6 and ZnL7, an additional four membered ring (Zn-S-Zn-S) is formed along with Zn-

S-C-S which appears due to the anisobidentate coordination mode of 1,1-dithiolate

ligands. The geometry around Zn atom in ZnL6 and ZnL7 is distorted square-pyramidal

with a τ value of 0.28 and 0.29, respectively.

The distance between Zn and all its five coordinating sulfur atoms is different from each

other in both ZnL6 [Zn(1)-S(2) = 2.35 Å)], [Zn(1)-S(1) = 2.44 Å], [Zn(1)-S(4) = 2.36

Å)], [Zn(1)-S(3) = 2.81 Å] and [Zn(1)-S(3)(bridge) = 2.38 Å] and ZnL7 [Zn(1)-S(2) =

83

Table 3.2: 1HNMR data (ppm) of the synthesized metal(II) dithiocarbamate complexes

3J (

1H,

1H) in Hz are listed in parenthesis.

Compound

Chemical shift δ (ppm)

CH2/CH

N

N

-OH N

-OCH3

HgL1 7.20-7.44,

m 4.30, s 4.2, t, (5.2), 2.54, t (5.1) - - -

HgL2 7.27-7.38,

m 3.56, s 4.06, t (5.1), 2.57, t (5.1) - - -

HgL3 - 2.18, t (5.1), 3.94, t

(5.1) 3.27, t (5.1), 3.99, t (5.1)

4.52,

s - -

HgL4 7.17-7.29,

m 4.63, d (5) - -

3.27, t (5.1), 1.74, q (5.1), 1.3,

m -

CdL1 7.20-7.44,

m 4.30, s 4.21, t (5.2), 2.54, t (5.1) - - -

CdL2 7.28-7.38,

m 3.57, s

4.16, t (5.1), 2.56, t

(5.1). - - -

CdL3

2.29, t (5.), 3.86, t (5.1) 3.26, t (5.1), 3.95, t (5.1) 4.43,

s - -

CdL4 7.12-7.31,

m 4.66, d (5) -

3.24, t (5), 1.75, q (5), 1.4, m -

ZnL1 7.19-7.44,

m 4.29, s 4.11, t (5.2), 2.53, t (5.1) - - -

ZnL2 7.22-7.4, m 3.51, s 4.19, t (5.2), 2.55, t (5.2) - - -

ZnL3

2.39, t (5.1), 3.98, t

(5.1) 3.92, t (5.1), 3.29, t (5.1)

4.33,

s - -

ZnL6 6.39-7.16,

m - 4.16, t (5.2), 3.25, t (5.2) - - 3.73, s

ZnL7 6.61-7.20,

m - 4.19, t (5.2), 3.55, t (5.2) - - 3.80, s

84

Table 3.3: 13

CNMR data of the synthesized metal(II) dithiocarbamate complexes

Compound

Chemical shift δ(ppm)

SCS

CH2/

CH N

N

CH2-CH2-OH N

-OCH3

HgL1 203.2 127.3, 128.1,

129.5, 142.1 76.7 53.0, 51.4 - - -

HgL2 202.4 127.5, 128.7,

129.3, 141.3 62.4 53.1, 52.2 - - -

HgL3 200.8 - - 53.2,52.8 58.8, 59.7 - -

HgL4 204.7 124.9, 127.4,

130.9, 136.7 52.1 - -

57.4, 42.2,

45.3 -

CdL1 203.2 127.3, 127.8,

128.7, 142.3 76.7 53.0, 51.4 - - -

CdL2 201.2 127.8, 128.4,

129.2, 141.7 62.7 52.3, 51.2 - - -

CdL3 201.5 - - 52.5, 51.6 58.1, 59.2 - -

CdL4 203.9 124.7, 127.9,

131.4, 136.3 50.1 - -

58.3, 45.6,

42.2 -

ZnL1 202.2 128.7, 127.3,

127.8, 141.7 75 51.3, 51.2 - - -

ZnL2 201.8 127.2, 128.4,

129.5, 142.1 61.9 52.6, 51.3 - - -

ZnL3 203.5 - - 52.3, 51.7 57.8, 58.9 - -

ZnL6 203.9

160.6,151.9,

130.2, 108.6,

105.0,102.2

- 50.9, 48.2

- 55.3

ZnL7 202.7 161.6,145.7,

115.6, 115.3 - 52.7, 49.1 - - 55.7

85

2.46 Å)], [Zn(1)-S(1) = 2.33 Å], [Zn(1)-S(4) = 2.32 Å)], [Zn(1)-S(3) = 2.90 Å] and

[Zn(1)-S(3)(bridge) = 2.35 Å] (Figure 3.11). However, the Zn-Sshort and Zn-Slong are

comparable with similar structure [47].

Evidently, a change in the position of methoxy groups from meta (ZnL6) to para (ZnL7)

position on peripheral phenyl rings changes the electronic nature as well as it affects the

stereochemical conformations of the molecules. This can be visualized clearly by two

Figure 3.11: Ball and stick diagram of (a) ZnL6 and (b) ZnL7 complexes.

different modes of intramolecular C-H···π interactions in both ZnL6 [C(21)-H(21)···Cg

3.34 Å) and ZnL7 [C(22)-H(22)···Cg 2.81 Å). Compound ZnL6 containing m-methoxy,

p-H and compound ZnL7 having p-methoxy, m-H on peripheral phenyl rings adopts such

a conformation that can be stabilized by intramolecular C-H···π interactions. It is further

86

demonstrated by the change of perfect chair conformation of piperazinyl moiety in ZnL6

to twist boat conformation in ZnL7 (Figure 3.11).

Table 3.4: Refinement parameters of complex ZnL6 and ZnL7

Crystal Data ZnL6 ZnL7

Chemical formula C48H60N8O4S8Zn2 C48H60N8O4S8Zn2

Crystal size (mm3 )

0.35 × 0.2 × 0.12

0.35 × 0.2 × 0.12

Formula weight (g/mole) 600.13 1200.26

Crystal system Monoclinic Monoclinic

Space group P21/c P21/c

T(K) 296(2) 296(2)

Wavelength (A)

0.71073

0.71073

Cell parameters

a/Å 11.4918(4) 14.892(2)

b/Å 8.4640(3) 14.3674(16)

c/Å 27.9084(9) 12.3605(16)

α/° 90.00

90.00

β/° 98.325(2) 94.859(5)

γ/° 90.00 90.00

V(Å3 ) 2685.95(16) 2635.2(6)

Z 4 2

Absorption coefficient (mm-1

) 1.255 1.279

F(000) 1248.0 1248

Data collection

2Θ range for data collection 4.3 to 56.66° 2.74 to 52°

Number of collected reflections 22112 20480

Number of independent

reflections 6580

5174

Goodness-of-fit on F2 0.994 1.020

Final R indexes [I>=2σ (I)] R1 = 0.0458, wR2 = 0.0773 R1 = 0.0408, wR2

= 0.0893

Final R indexes [all data] R1 = 0.1056, wR2 = 0.0918

R1 = 0.0667, wR2

= 0.1020

87

Table 3.5: Selected Bond lengths (Å) and Bond Angles (o) of complex ZnL6 and

ZnL7

Type of

Bond

Bond Length (Å) Type of Bond Bond Angle (o)

Zn1 -S1

ZnL6 ZnL7

S1 -Zn1 -S31

ZnL6 ZnL7

2.4442(7)

2.3335(10)

153.94(3)

-

Zn1 -S2

2.3524(7) 2.4628(9) S2 -Zn1 -S1

75.38(2)

75.77(3)

Zn1 -S3

2.3802(8) 2.3537(9) S2 -Zn1 -S31

91.63(2)

-

Zn1 -S31

2.8134(7)

- S2 -Zn1 -S3

117.87(3) 107.81(3)

Zn1 -S4

2.3616(8)

2.3172(9) S2 -Zn1 -S4

136.75(3)

103.26(3)

S1 -C1

1.719(2) 1.726(3) S3 -Zn1 -S1

114.64(3)

116.19(3)

S2 -C1

1.727(2) 1.717(3) S3 -Zn1 -S31

91.37(2)

-

N1-C1 1.329(3)

1.316(4) S4 -Zn1 -S1

104.78(3) 133.65(4)

S3-C13 1.744(3) 1.735(3) S4 -Zn1 -S3

101.60(3) 108.23(3)

S4-C13 - 1.718(3) S4 -Zn1 -S31

69.32(2) -

S4 C131

1.713(2)

- - - -

N3-C13 1.320(3)

1.320(4) - - -

3.3.2 Molecular packing of ZnL6 and ZnL7 complexes

A significant effect of electronic and conformational changes on molecular packing was

also observed. There are some general considerations, which are worth discussing before

going into the individual molecular packing of ZnL6 and ZnL7. In molecular structure

ZnL6, where methoxy groups are at meta position, electron donating mesomeric effect of

both amino and methoxy group is additive on ortho/para positions. However, more

electronegative oxygen atom of methoxy group partially withdraws electron density from

nearby ortho positions making them slightly less electron rich as compared to its para

88

Figure 3.12: Side view of (a) ladder-like sandwiched supramolecular chains driven by C-

H···π contacts, a schematic diagram (left), a wireframe model (right) (b)

zig-zag arrangement of dimeric H-shaped ZnL6 in the crystal structure, a

schematic diagram (top), a wireframe model (bottom) effect.

positions. So, both ortho and para positions of these substituents are potential hydrogen

bond acceptors for C-H···π interactions. In ZnL7, where methoxy groups are at para

position of peripheral aromatic rings, electron density is almost symmetrically distributed

favoring C-H···π interactions. Nevertheless, dominating mesomeric effect of amino

group, not only make its ortho positions electron rich, it also prevent methoxy group to

show its mesomeric effect. Therefore, methoxy group only expresses its inductive effect

a

b

89

making its ortho positions slightly electron deficient. The four carbons of piperazinyl

moiety in both ZnL6 and ZnL7 are electron deficient, partially due to the

electronegativity and partially due to positive character of nitrogen achieved through

delocalization of its lone pairs to the neighboring unsaturated systems. This is clearly

demonstrated by the shorter N-C bond distances in both ZnL6 [N(1)–C(1)S(1)S(2)

1.33Å, N(3)–C(13)S(3)S(4) 1.32Å, N(4)–C(18) 1.41Å & N(2)–C(6) 1.42Å] and ZnL7

[N(1)–C(1)S(1)S(2) 1.32Å, N(3)–C(13)S(3)S(4) 1.32Å, N(4)–C(18) 1.41Å & N(2)–C(6)

1.40Å]. Therefore, the hydrogens connected to these carbon atoms are slightly acidic and

potential hydrogen bond donors.

Figure 3.13: A spacefill representation of two successive layers of zig-zag layered

structure of ZnL6, top view (top), side view (bottom).

A zig-zag layered structure of ZnL6 in the solid state stabilized by four different types of

hydrogen bonds namely, C-H···π, CH···S, CH···O, CH···C hydrogen bonds is

conceived. Every dimeric H-shaped molecule of ZnL6 is sandwiched between another

two molecules providing infinite 1D ladder-like supramolecular chain driven by C- H···π

[C(5)-H(5B)···C(23) 2.73 Å], [C(2)-H(2A)···C(19) 2.73 Å] and [C(2)-H(2A)···C(23)

2.83 Å] interactions (Figure 3.12a). Interestingly, these H-shaped infinite 1D

supramolecular chains of ZnL6 move in zig-zag fashion cemented by multiple weak

hydrogen bond contacts [C(4)-H(4A)···O(2) 2.66 Å], [C(5)-H(5A)···C(24) 2.83 Å] and

[C(12)-H(12B)···S(3) 2.93 Å] (Figure 3.12b). Further, the sideways extension of these

zig-zag H-shaped infinite 1D supramolecular chains of ZnL6 by means of another weak

90

intermolecular CH···S hydrogen bonds [C(3)-H(3A)···S(1) 2.97 Å] provide zig-zag

layered structure (Figure 3.13).

Fig. 3.14: Supramolecular parallel extension of H-shaped molecules of ZnL7 to make a

layer, top view (left), side view (right).

Figure 3.15 Supramolecular parallel and sideways extension of H-shaped molecules of

ZnL7 to make a layer, top view (left), side view (right).

The molecular packing of ZnL7 is altogether different from ZnL6. The layered structure

of ZnL7 is again obtained, stabilized by three different types of multiple hydrogen bonds

91

namely, CH···π, CH···S, CH···O hydrogen bonds. However, H-shaped molecules of

ZnL7 move in a parallel zig zag fashion in a specific layer to provide an uneven surface

of the layer as compared to head to tail zig zag movement of H-shaped molecules in

forming smooth zig zag layer of ZnL6 (Figure 3.14). A bifurcated CH···S [C(5)-

H(5A)···S(1) 2.92 Å], [C(5)-H(5A)···S(3) 2.95 Å], another CH···S [C(14)-H(14B)···S(4)

3.00 Å], and CH···π [C(17)-H(17B)···C(19) 2.86 Å] interactions stabilize this parallel

zig zag arrangement of molecules in packing of ZnL7. Sideways extension of this parallel

zig zag molecular arrangement by means of weak CH···O contacts [C(10)-H(10)···O(2)

3.34 Å] completes a layer (Figure 3.15). Various such layers are connected with each

other by means of CH···π [C(11)-H(11)···C(20) 2.86 Å, C(17)-H(17B)···C(20) 2.86 Å],

A bifurcated CH···S [C(5)-H(5A)···S(1) 2.92 Å], [C(5)-H(5A)···S(3) 2.95 Å], another

CH···S [C(14)-H(14B)···S(4) 3.00 Å] and CH···O [C(3)-H(3B)···O(2) 2.65 Å, C(10)-

H(10)···O(2) 2.71 Å] interactions to provide a multi-layered structure (Figure 3.16).

Figure 3.16: Multilayered structure of ZnL7 complex.

3.3.3 Structural descriptions of HgL1 and HgL2 complexes

The complexes HgL1 and HgL2 were also characterized by single crystal XRD, the

crystallographic parameters and, selected bond angles and bond length are presented in

tables 3.6 & 3.7. The crystal system of HgL1 and HgL2 is monoclinic P21/C and triclinic

92

Figure 3.17: Ball and stick diagram of (a) HgL1 and (b) HgL2 complexes.

P-1, respectively. HgL1 has a distorted tetrahedral geometry mediated by two

dithiocarbamate bidentate chelates (Figure 3.17a.) as evident from the angles; S(1) Hg(1)

S(3) 111.19(5)o and S(2) Hg(1) S(4) 139.15(6)

o [48]. Two sulfur atoms of one

dithiocarbamate are coordinated by isobidentate manner with Hg-S bond lengths [Hg(1)-

S(1) 2.534Å and Hg(1)-S(2) 2.572Å] and remaining two sulfur atoms of other

dithiocarbamate are coordinated by anisobidentate manner with short Hg-S [Hg(1)-S(3)

2.708Å] and long Hg-S [Hg(1)-S(4) 2.410Å] bond lengths. In the both dithiocarbamate

93

Table 3.6: Refinement parameters of HgL1 and HgL2 complexes

Crystallographic Complex HgL1 Complex HgL2

Parameter

Empirical formula C24H30HgN4S4 C38H44HgN4OS5

Formula weight 703.35 933.66

Temperature/K 296(2) 296(2)

Crystal system monoclinic triclinic

Space group P21/c P-1

a/Å 20.417(3) 9.7914(9)

b/Å 9.1364(10) 14.8953(19)

c/Å 14.6595(15) 15.3991(19)

α/° 90 67.256(4)

β/° 94.422(5) 73.727(3)

γ/° 90 85.211(4)

Volume/Å3 2726.5(6) 1987.5(4)

Z 4 2

ρcalcg/cm3 1.713 1.560

μ/mm-1

5.971 4.173

F(000) 1384.0 936.0

Crystal size/mm3 0.360 × 0.280 × 0.200 0.42 × 0.16 × 0.14

Radiation MoKα (λ = 0.71073) MoKα (λ = 0.71073)

2Ө 4.002 to 51.994 4.938 to 54.272

Index ranges -25 ≤ h ≤ 25, -11 ≤ k ≤ 11, -12 ≤ h ≤ 11, -19 ≤ k

≤ 19,

Reflections collected -11 ≤ l ≤ 18 -19 ≤ l ≤ 19

19465 29916

Independent reflections 5300 [Rint = 0.0592, Rsigm= 8630 [Rint = 0.0602,

Rsigma

0.0774] = 0.0749]

Data/restraints/parameters 5300/0/231 8630/7/362

Goodness-of-fit on F2 1.178 0.999

Final R indexes [I >=2σ (I)] R1 = 0.0719, wR2 = 0.1440 R1 = 0.0489, wR2 =

0.1000

Final R indexes [all data] R1 = 0.1410, wR2 = 0.1843 R1 = 0.1050, wR2 =

0.1181

Largest diff. peak/hole/e Å-3

3.29/-1.66 0.95/-0.97

ligands the C-N bond lengths [C(1)-N(1) 1.324Å and C(19)-N(3) 1.33Å] are similar and

are shorter than single C-N bond (1.47 A° ) but larger than C-N double bond (1.28 A° )

[44]. The dimeric molecular structure was observed for compound HgL2 featuring four

dithiocarbamate chelates with two sulfur atoms bridging two Hg atoms resulting in

distorted square-pyramidal geometry, having different Hg-S bond lengths [Hg(1)-S(1) =

94

2.408Å)], [Hg(1)-S(2) = 3.164Å], [Hg(1)-S(3) = 2.772Å)], [Hg(1)-S(4) = 2.454Å] and

[Hg(1)-S(2)(bridge) = 2.747Å] (Figure 3.17b) [46]. Due to this dimeric nature of HgL2,

an additional four membered ring (Hg-S-Hg-S) is formed along with Hg-S-C-S which

appears due to the anisobidentate coordination mode of 1,1-dithiolate ligands. In the both

dithiocarbamate ligands the C-N bond lengths [C(1)-N(1) 1.319Å and C(13)-N(3)

1.328Å] are similar and are shorter than single C-N bond (1.47 A° ) however larger than

C-N double bond (1.28 A° ) as observed earlier [44].

Table 3.7: Selected bond lengths and bond angles of HgL1 and HgL2 complexes

Type of Bond Bond Length (Å) Type of Bond Bond Angle (o)

HgL1 HgL2 HgL1 HgL2

Hg1-S1

2.4442(7) 2.408(4) S1-Hg1-S21 - 102.93(12)

Hg1-S2 2.572(16) 3.164(4) S1-Hg1-S2 71.13(5) 63.62

Hg1-S21 - 2.747(4) S3-Hg1-S2

1 - 99.68(11)

Hg1-S3 2.708(15) 2.772(3) S3-Hg1-S2 110.70(6) 155.15

Hg1-S4 2.410(15) 2.454(3) S4-Hg1-S21 - 101.47(11)

C1-S1 1.726(6) 1.756(14) S4-Hg1-S2 139.15(6) 95.44

C1-S2 1.725(6) 1.695(13) S2-Hg1-S21 - 102.36

C1-N1 1.323(7) 1.319(16) S1-Hg1-S3 111.19(5) 121.96(11)

S3-C19 1.701(6) - S4-Hg1-S3 70.65(5) 68.85(10)

S3-C13 - 1.695(13) S1-Hg1-S4 148.63(6) 150.69(13)

S4-C19 1.741(5) - - - -

S4-C13 - 1.743(12) - - -

C19-N3 1.33(6) - - - -

C13-N3 - 1.328(15) - - -

95

3.3.4 Molecular packing of HgL1 and HgL2 complexes

The crystal structure of compound HgL1 disclosed that two homoleptic molecules are

joined by non-covalent S···S [S(1)···S(4) 3.598 Å] interactions to form a dimer (Figure

3.18a). Each dimer is linked to a dimethylsulfoxide molecule through CH···S [C(38A)-

H(38B)···S(3) 2.877Å], CH···CH [C(23)-H(23B)···C(37A)-H(37C) 2.315Å] and CH···π

[C(38A)-H(38C)···C(33A) 2.817Å] interactions and two dimethylsulfoxide molecules

are linked together by CH···O [C(37A)-H(37A) ···O(1A) 2.212Å] and C···O[C(37A)

···O(1A) 3.161Å] attractions to create 2D-supramolecular chain (Figure 3.18b). The 2D-

supramolecular chains are linked through CH···CH [C(9A)-H(9A)···C(37A)-H(37B)

2.061Å, C(5)-H(5A)···C(37A)-H(37C) 2.003Å], CH···π [C(24)-H(24)···C(10A) 2.874Å,

C(28)-H(28)···C(35A) 2.890Å, C(28)-H(28)···C(34A) 2.818Å, C(28)-H(28)···C(33A)

2.875Å], CH···C [C(5)H(5A)···C(37A) 2.650Å], CH···S [C(23)-H(23A)···S(2) 2.952Å]

and S···C [S(2)···C(23) 3.488Å] interactions to form overall 3D-supramolecular structure

(Figure 3.18c).

Figure 3.18: (a) Dimeric structure of complex HgL1 (b) 2D supramolecular structure and

(c) 3D supramolecular structure of complex HgL1.

96

The crystal structure of compound HgL2 disclosed that molecules are linked through

CH···C [C(6A)H(6B)···C(13) 2.817Å], CH···S [C(2)-H(2B)···S(1) 2.927Å, C(17)-

H(17A)···S(1) 2.848Å, C(5)-H(5A)···S(4) 2.892Å], CH···CH [C(22A)-

H(22A)···C(10A)-H(10A) 2.204Å] and CH···π [C(22A)-H(22A)···C(10A) 2.838Å]

interactions to form 2D-supramolecular chain (Figure 3.19a). The 2D-supramolecular

chains are linked through CH···S [C(2)-H(2B)···S(1) 2.927Å, C(17)-H(17A)···S(1)

2.848Å, C(5)-H(5A)···S(4) 2.892Å] interactions to form overall 3D-supramolecular

structure (Figure 3.19b).

Figure 3.19: (a) 2D supramolecular structure and (b) 3D supramolecular structure of

complex HgL2.

97

3.4 Formation mechanism and characterization of metal sulfide nanostructures

The physico-chemical behavior of group II-VI semiconductor nanoparticles (NPs) is

known to be influenced by their size and shape [49,50]. High surface to volume ratio

(enormous amount of superficial atoms) and or the three dimensional (3D) confinements

of electrons are reasons for their novel and enhanced properties. Altering size of NPs

modify the degree of confinement of electrons and hence the electronic structure

(particularly the band edge) [51]. Therefore, controlling size and morphology of NPs is a

hot area of research in the field of nanoscience and technology [52]. Solvothermal

method has been proved an effective and extensively studied technique for controlling the

shape and size of the semiconductor NPs, in recent years. It has been proposed that the

NPs morphology is directed by a delicate balance between the kinetic and

thermodynamic growth regimes [53]. Additionally, formation of nuclei during the

decomposition of precursor is a key step in directing shape of the nanocrystals [54].

To better realize impact of the precursor nature on size and morphology of the metal

sulfide (MS) NPs, the as-synthesized M(II) dithiocarbamates were used as SP to get MS

(M = Zn, Cd, Hg) NPs.

Table 3.8: Molecular properties of HgL1-HgL4 complexes.

HgL1 HgL2 HgL3 HgL4

ЄHOMO(eV) -5.7462 -5.7462 -5.8460 -5.8354

ЄLUMO(eV) -2.2877 -1.7666 -1.9730 -1.7136

ΔЄ=(ЄLUMO-ЄHOMO)(eV) 3.4585 3.9795 3.873 4.1218

IE= -ЄHUMO(eV) 5.7462 5.7462 5.8460 5.8354

EA= -ЄLUMO(eV) 2.2877 1.7666 1.9730 1.7136

Global Hardness

(η)=1/2(ЄLOMO-ЄHOMO),

1.7292 1.9898 1.9365 2.0609

Chemical Potential

μ=1/2(ЄHOMO + ЄLUMO)

-4.0169 -3.7564 -3.9095 -3.7745

98

The kinetics of decomposition of a complex, a controlling factor to tune size and

morphology, depends on the ligand binding strength which in turns relay on ligand

structure [55]. So size and morphology can be tuned by altering the ligand structure. To

investigate this fact, the Hg (II) precursors were chosen, which differ in stability and

other molecular properties as determined by the density functional theory calculations

(Table 3.8). The HOMO-LUMO energy gap, an index of stability and chemical

reactivity, of the complexes was the main focus of this study [56].

Generally, greater is this energy gap, greater is the stability or low chemical reactivity

and vice versa. On the basis of HOMO-LUMO gap the stability of the title complexes

was found in the following order, HgL4 > HgL2 > HgL3 > HgL1.

The HOMO-LUMO gap (Table 3.8) of HgL1 precursor (predicted by DFT) suggests its

lower stability, so it can be assumed that HgL1 rapidly decomposes to release the HgS

nuclei. Thus, fast decomposition of HgL1 results in fast nucleation and growth of the

NPs, however, further growth is hindered due to diminishing of the surface energy by

adsorption of the in situ generated organic molecules at a suitable stage [57]. Hence, the

thermolysis of HgL1 has resulted in the formation of spherical particles having 15-20 nm

diameters with uniform size distribution as evident in the SEM image (Figure 3.20a). The

fast nucleation and growth commonly lead to the formation of zinc blend phase of the

crystals [58], as observed in the powder diffraction pattern of HgS-1 (Figure 3.28a).

Figure 3.20: SEM image of (a) HgS-1 and (b) HgS-2 NPs.

99

To confirm the existence of any organic molecule on the surface of NPs, the FT-IR

analysis were carried out, which suggests presence of the in situ generated Imidazolidine-

2-thione on the surface of HgS-1, HgS-2 and HgS-4 NPs. Both HgS-1 and HgS-2 (Figure

3.21a) show a single broad band ranging from about 3300 cm-1

to 3600 cm-1

indicating

presence of moisture that may overlap the N-H stretch (both primary and secondary

amine). However, for HgS-4, N-H band was observed at 3393 cm-1

(Figure 3.21b). The

other bands at 2992-2997 cm-1

and 2885-2886 cm-1

(symmetric and asymmetric stretches

of methylene), 1642-1647cm-1

(N-H bending), 1292-1295 cm−1

(C-N stretch) and three

bands in the range of 1532-1570 cm-1

, 1355-1415 cm-1

and 1080-1154 cm-1

due to the

mixed stretching of –N-C=S were also observed [59]. All these bands indicate presence

of Imidazolidine-2-thione on the surface. Identical capping agent on the surface of HgS1,

HgS2 and HgS4 negates any role of capping agent in the formation of different sizes and

morphologies.

The precursors HgL2 has afforded anisotropically grown rectangular slabs (Figure 3.20b)

of about 20 nm thickness and length of about 30 nm. As HgL1 and HgL2 have different

Figure 3.21: FT-IR spectra of (a) HgS-1 and HgS-2 (b) HgS-4 NPs

thermodynamic stabilities (as evident from DFT data) and different crystalline geometry.

Slow decomposition of HgL2 may result in low concentration of nuclei with different

shape [57], the slow growth of which might facilitate the oriented attachment of the

crystallites terminated in rectangular slabs.

100

Large cubic shaped crystals of 120-130 nm in dimensions and sharp edges can be

observed in the SEM image (Figure 3.22a) of HgS-3 NPs. The SEM image shows

uniform size distribution of the particles whereas sharp and narrowness of XRD peaks

confirm its better crystallinity than the forgoing NPs (Figure 3.28b). A sharp band at

1085 cm-1

, in FT-IR spectrum of the HgS-3 NPs (Figure 3.22b), can be granted to C-O

stretch of the primary alcohol, signifying presence of the hydroxy-ethylpiperazine, an

expected byproduct of decomposition , on the surface of NPs. So, the high affinity of

hydroxy-ethylpiperazine for {100} and {110} planes, due to their high energy, than

{111} plane may be the reason for growth along later plane to form cubes [55].

The SEM image HgS-4 NPs (Figure 3.23a) confirms the quasi spherical and diverse

particles size (40-150 nm). Likewise HgS1 and HgS2, the presence of Imidazolidine-2-

thione on the surface of HgS4 (Figure 3.21b) and relatively stable nature of its parent

precursor HgL4 (Table 3.8), allow us to assume that slow decomposition may be the

reason for different size and shape. The slow decomposition of HgL4 leads to low

concentration of HgS monomers in the course of the reaction, which may quickly grow

into seeds. These seeds further grow into large particles either through deposition of the

newly generated Hg-S nuclei on its surface or undergo Ostwald’s ripening ended in large

size quasi spherical particles [60,61].

Figure 3.22: (a) SEM images of HgS-3 NPs (b) FT-IR spectrum of HgS-3 NPs.

101

The SEM images of the HgS-5 NPs (Figure 3.23 b&c) show nanorods of around 100 nm

in diameter and 700 nm in length. Usually oriented attachment, structure directing agent

or a strongly coordinating solvent is believed to stimulate anisotropic growth [53].

However, herein, unlike HgL1-HgL4, which dissolved in en before 85 oC, the HgL5

remain undissolved up to the en boiling point (117 oC). So no nucleation was observed up

till 117 oC. However, at boiling temperature of en b.p, the precursor instantly decomposed

inducing a nucleation burst, thus provided a kinetic drive for anisotropic particles growth

or HgS nanorods [62].

Figure 3.23: SEM images of (a) HgS-4 NPs (b)(c) HgS-5 NPs (d) FT-IR spectrum of

HgS-5 NPs.

Additionally, in FT-IR spectrum of HgS-5 (Figure 3.23d) NPs, absence of any significant

band in the functional group region indicates that the particles surface is devoid of any in

102

situ generated organic capping agent, thus supporting the foresaid proposed mechanism

of nanorods formation.

The thermolysis of Cd(II) dithiocarbamates (CdL1-CdL5) composed of similar

dithiocarbamate ligands in en has given five different morphologies i.e. nanospheres (size

Figure 3.24: TEM images of (a) CdS-1(nanospheres) (b) CdS-2 (nanocapsules) (c) CdS-

3 (nanograins) (d) CdS-4 (nanosheets) and (e,f) CdS-5 (nanorods).

103

20-70 nm), nanocapsules (20 nm diameter and 30 nm length), nanograins (10-15 nm),

nanosheets with thickness of less than 5 nm and nanorods (80 nm length and 4-6 nm in

diameter) as shown in figure 3.24a-f. The mechanism of the morphological evolution of

CdS nanostructures is also governed by the similar factors as discussed for HgS

nanostructures.

However, CdL4 precursor has afforded nanosheets (Figure 3.24d) which seem to be

continuous and extend to several hundred micrometers in length and width. The

transparent image suggests the transmission of electrons through sheets of thickness less

than 5 nm. No evident boundary among the individual sheets suggests stacking over one

another. The formation of nanosheets from CdL4 can be attributed to its high solubility in

en, completely dissolved at 60 °C. At 70 °C the solution turned slightly turbid due to

nucleation of the CdS monomers resulting in individual CdS dot formation. A gradual

increase in the temperature resulted in more CdS dots formation that probably adopted

spherical assembly through common crystallographic facet {111} and then at the en

boiling point (b.p) grew into sheet-like structure in {111} direction through oriented

attachment [63]. Moreover, a lattice spacing of 0.33 nm was measured for both CdS

nanosheets and nanorods from their HR-TEM images (Figure 3.25 a,b), which is in

agreement to the (002) plane of hexagonal CdS confirming their preferential growth

along the c-axis.

Figure 3.25: HRTEM images of (a) CdS-4 (nanosheets) and (b) CdS-5.

104

Figure 3.26: SEM images of the as prepared (a) ZnS-1 (b) ZnS-2 (c) ZnS-3 (d) ZnS-6

and (e,f) ZnS -7 NPs.

Surprisingly, the thermolysis of Zn(II) complexes of the forgoing dithiocarbamate ligands

did not produce ZnS NPs of significant morphological and size variation. As shown in

figure 3.26a&b, the products obtained from the thermolysis of precursor complexes ZnL1

105

and ZnL2 exhibit sheets like morphology, consisting of small and large sheets of irregular

thickness. However, the thermolysis of ZnL3, ZnL6 and ZnL7 has resulted in coalesced

particles and bulk structures (3.26c-f), so suggesting that this method is inappropriate for

the synthesis of ZnS NPs under the employed experimental conditions. The thermolysis

of precursor complexes (HgL1-HgL5) was also examined in octlyamine (OA) for the

purpose to explore the effect of thermolysing solvent on the morphology and size of the

NPs. It was found that all five Hg(II) dithiocarbamate precursors afforded nearly

monodispersed spherical particles of small size (20 ± 5 nm) as can be seen in the SEM

images (Figure 3.27a-d).

Figure 3.27: SEM images of (a,b) HgS-1 (c) HgS-2 (d) HgS-3 NPs synthesized in the

presence of OA as thermolysing solvent.

The evolution of almost similar morphology and size from different Hg(II)

dithiocarbamates can be attributed to the dominant role of thermolysing solvent, which

106

being a long chain organic amine may act as a surfactant leading to the formation of

small NPs of the spherical shape [52]. Thus, it can be concluded that long chain amines

like OA can be used to get particles of the smaller size, but with no obvious

morphological variation under the employed experimental conditions.

The reflections from the (111), (200), (220) and (311) planes in the diffraction pattern of

HgS-1 and HgS-4 (Figure 3.28a) are characteristics of cubic mercury sulfide, which show

decent match with the reported pattern (JCPDS NO.00-006-0261). However, minor peaks

at 31.29o, 45.89

o, 51.87

o (2θ) indicate the presence of hexagonal phase of HgS. The

diffraction pattern of mercury sulfide nanoparticles synthesized from the precursors

HgL2, HgL3 and HgL5 are shown in (Figure 3.28b). The reflections from the different

planes agree well with hexagonal phase (JCPDS 01-080-2192). The relatively broad

XRD lines in the diffraction pattern of nanocrystals formed from precursor HgL2 indicate

reduced particle size as evident from the SEM. By changing the thermolysis solvent from

en to OA, all the precursors exclusively afforded cubic phase of HgS (Figure 3.28c), as

all the reflections can be completely indexed to the cubic phase of HgS using the standard

JCPDS card No. 00-006-0261. Moreover, the observed sharp and broad peaks in

diffraction pattern of HgS NPs prepared with OA suggest superior crystallinity and

reduced particle size as compared to HgS NPs synthesized with en.

X-ray diffraction analysis of CdS NPs revealed that the thermolysis of all the precursors

has exclusively come up with wurtzite (hexagonal) phase. The reflections in the XRD

pattern of CdS-1, CdS-2 and CdS-3 NPs (Figure 3.28d) are consistent with the standard

JCPDS card No. 00-042-1049 while the reflections in the XRD pattern of CdS-4 and

CdS-5 can be indexed with standard JCPDS card No. 01-077-2306, in consonance with

the typical hexagonal phases of CdS (Figure 3.29a). All the diffraction pattern of both

HgS and CdS NPs are devoid of any detectable impurities such as HgO, CdO, Hg, Cd or

free sulphur etc.

The broad peak in XRD pattern of the CdS nanostructures reveals small size and less

crystallinity as compared to the sharp and narrow peaks in the diffraction pattern of HgS,

which points toward good crystallinity and large size, as was also observed in their SEM

and TEM results.

107

Figure 3.28: XRD pattern of (a) HgS-1 and HgS-4 (b) HgS-2, HgS-4 and HgS-5 NPs

synthesized in the in the presence of en (c) XRD pattern of HgS-1OA,

HgS-2OA and HgS-3OA synthesized in the presence of OA (d) XRD

pattern of CdS-1, CdS-2 and CdS-3 NPs synthesized in the presence of en.

The comparatively less crystallinity of CdS than HgS NPs can be accredited to the lesser

atomic mass and size of the Cd, a fact further supported by the formation of

semicrystalline ZnS by the thermolysis of the Zn-precursor composed of ligands similar

to those of Cd- and Hg-complexes. The multiple and weak reflections in XRD pattern of

ZnS (Figure 3.29c) is an evidence of the mixed crystal structure formation with poor

crystallinity. Some of the reflections can be assigned to wurtzite while others corresponds

to zinc blende structure of ZnS along with some unknown peaks, however, annealing of

the samples at 400 oC in the atmosphere of argon showed few and sharp reflections

108

(Figure 3.29d), which is in close agreement to hexagonal phase (wurtzite) of ZnS (JCPDS

card No. 01-075-1534).

Figure 3.29: (a) XRD pattern of as prepared CdS-4 and CdS-5 NPs (b) XRD pattern

CdS-4 and CdS-5 NPs after three sequential photocatalytic cycles. (c)

XRD pattern of as prepared ZnS-1 and ZnS-6 NPs (b) XRD pattern ZnS-1

and ZnS-6 NPs after calcination at 400 oC in the atmosphere of argon.

So, in addition to other parameters, size and mass of the respective metal ion may also

contribute a substantial effect in controlling the crystallinity, phase purity and

morphology of the resultant NPs.

109

Figure 3.30: (a) EDS Spectrum of HgS-1 and (b) ZnS-2 NPs.

The elemental construction of NPs was analysized by energy dispersive X-ray

spectroscopy (EDS), which confirmed the presence of metal (metal = Hg, Cd, Zn) to

sulphur ratio of 1:1. The representative EDS spectra of HgS and ZnS are displayed in

figure 3.30, beside the presence of respective metal and sulphur; traces of oxygen and

carbon have also been detected. This may be due to the existence of capping molecule on

the surface of NPs.

UV-Visible absorption spectroscopy is an important method to assess the optical

properties of the semiconductor NPs. The UV/Visible absorption spectra of HgS, CdS

and ZnS nanostructures are shown in figure 3.31 a-c. From the UV/Visible spectra it can

be observed that both HgS and CdS NPs show good photo absorption in the visible

region with well-defined excitonic absorption maxima at 372 (HgS-1), 390 nm (HgS-2)

and 425-470 nm (CdS NPs). However, ZnS NPs displayed closed excitonic absorption

maxima in the UV region centered at 315-325 nm.

110

Figure 3.31: UV/Visible absorption spectra of (a) HgS-1 and HgS-2 (b) CdS-1 - CdS-5

and (c) ZnS-1 - ZnS-3, ZnS-6 - ZnS-7 NPs.

111

The blue shift in the excitonic absorption maxima of NPs as equated to respective bulk

materials (ca. HgS = 620 nm, CdS = 512, ZnS = 340) can be indorsed to well-known

quantum size effect.

3.5 Solar-light driven photocatalytic degradation of Congo red dye by CdS

nanospheres and nanocapsules

The toxic organic dyes released into wastewater from textiles, food and cosmetics

industries are causing water pollution, a noteworthy environmental problem [64]. These

colored compounds upset aquatic biological processes either by blocking light

penetration or direct destruction of aquatic communities due to associated noxiousness

[65]. To find an effective and long lasting solution for the exclusion of these aquatically

harmful compounds a range of technologies have been tested which include

biodegradation, adsorption and photocatalytic oxidation [66-68]. The photocatalytic

oxidation of these toxic pollutants using the nano-sized semiconductors as photocatalyst

seems to be more appropriate strategy [69]. Among the semiconductors NPs, TiO2 has

been certainly regarded as the most efficient catalyst for the degradation of toxic organic

compounds under UV radiation. However, its poor response to visible light due to wide

band gap hinders its application as photocatalyst [70]. In recent years, tremendous efforts

have been witnessed for maximum utilization of clean, safe and abundant solar energy

through the development of visible light driven photocatalyst [71-73]. Cadmium sulfide,

an important chalcogenide having direct band gap (2.42 eV and nonlinear optical

properties) material, has been extensively used in recent years as visible-light driven

photocatalyst for the treatment of aqueous dye contamination [74]. To compare the

visible-light-driven photocatalytic performance, cadmium sulfide nanospheres and

nanocapsules were used for the degradation of Congo red dye in direct sunlight.

Mechanistically, the photocatalytic reaction of CdS is initiated when an electron jump to

conduction band from valence band, when it is irradiated with light of energy (hν)

equivalent to or larger than the band gap energy of CdS leaving hole behind (Eq. 3.5).

The photogenerated hole reacts with water to generate hydroxyl radical (•OH) (Eq. 3.6),

which is a strong oxidizing agent that attack on adsorb or nearby dye molecules to the

catalyst surface, thereby degrading it. The electrons in the conduction band are hunted by

112

molecular O2 and get converted into superoxide radical anion (Eq. 3.7), which is

protonated and yield hydrogen peroxide (Eq. 3.8). The hydrogen peroxide dissociates

into hydroxyl radical and further accelerates the degradation process (Eq. 3.9 and Eq.

3.10).

𝐶𝑑𝑆 + ℎν → 𝐶𝑑𝑆 {𝑒−(𝐶𝐵) + ℎ+(𝑉𝐵)} (3.5)

𝐻₂𝑂 + ℎ+ → 𝐻�� + 𝐻⁺ (3.6)

𝑂₂ + 𝑒⁻ → 𝑂₂⁻ (3.7)

𝑂₂⁻ + 2𝐻⁺ → 𝐻₂𝑂₂ (3.8)

𝐻₂𝑂₂ → 𝐻�� + 𝐻�� (3.9)

𝑂�� + 𝐷𝑦𝑒 → 𝑑𝑒𝑔𝑟𝑎𝑑𝑎𝑡𝑖𝑜𝑛 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 (3.10)

In this work, the photocatalytic performance of cadmium sulfide CdS-1 and CdS-2 was

evaluated for the decontmination of CR in direct sunlight and in dark. The degradation

progress of the CR was monitored by observing depletion in the intensity of absorption

bands at 497 nm (azo link) and at 338 nm (naphthalene group) as a function of irradiation

time. The UV/ Visible spectral changes in the concentration of CR dye with reaction time

in the presence of CdS-1 and CdS-2 NPs is displayed in figure 3.32a-d. In the presence of

CdS-1 NPs 99.3% decolorization of the dye is achieved in 60 minutes due to the

fragmentation of azo links (497nm), additionally the decay at 338 nm provide an

evidence that 93% of the of aromatic fragments in the dye molecule have been degraded

along with azo links. However, in the presence of CdS-2 NPs 99.2% decolorization is

reached in 30 minutes along with 95% degradation of the aromatic fragments. To further

evaluate the influence of temperature on the catalytic performance of CdS NPs, CdS-1

was selected and its activity was tested at 40 oC and 50

oC keeping the other experimental

conditions the same. A significant catalytic activity was observed at higher temperatures

i.e. 84.7% and 96.7% decolourization was observed in 20 minutes at 40 oC and at 50

oC,

respectively.

At these two elevated temperatures almost complete decolorization was witnessed

sequentially in 50 and 40 minutes (Figure 3.33a). An increase in photocatalytic activity of

113

Figure 3.32: UV/Visible spectral changes of Congo red solution with irradiation time in

the presences of CdS-1NPs at (a) 30 oC (b) at 40

oC (c) at 50

oC and (d)

CdS-2 NPs at 30 oC.

cadmium sulfide NPs at elevated temperature is due to acceleration in electron-hole

transport, a well-established behavior of semiconductor materials. The photocatalytic

behavior of the CdS-1 and CdS-2 nanoparticles was confirmed by the fact that doing the

same reaction in the dark with CdS-1 NPs, only 26% and 65% of the dye molecules

adsorbed on the particles surface in 30 and 70 minutes, but afterward no appreciable

increase in adsorption/degradation was noted till 100 minutes as shown in figure 3.33b.

The reaction followed pseudo first order kinetics as a linear relation subsists amid the dye

concentration and irradiation time (Figure 3.33c). The calculated k values for the

degradation reaction by CdS-1(nanospheres) at 30 oC, 40

oC and 50

oC are -0.0645

114

Figure 3.33: (a) Photocatalytic degradation rate of Congo red with irradiation time in the

presence of CdS-1 and CdS-2 NPs (b) UV/Visible spectral changes of

Congo red solution with reaction time in the presences of CdS-1 NPs in

dark (c) The degradation kinetics of Congo red with irradiation time in the

presence of CdS-1 and CdS-2 NPs (d) Time-resolved PL of CdS-1 and

CdS-2 NPs, following excitation at 305 nm.

min-1

, -0.0683 min-1

and -0.178 min-1

indicating a rapid increase in the photodegradation

rate with increasing temperature. However, CdS-2 (nanocapsule) showed high

degradation kinetics even at low temperature, with k value of -0.1224 min-1

(Figure

3.33c). At the same temperature (30 oC), the photocatalytic activity of CdS-2

(nanocapsules) is 2 times more than CdS-1 (nanospheres). In order to get an

understanding about the superior photocatalytic activity of CdS-2 than CdS-1, the time-

resolved photoluminescence (PL) measurements were undertaken. The PL kinetics are

115

displayed in figure 3.33d, the CdS-1 shows shorter PL life-times as compared to CdS-2,

and it was best fitted with a bi-exponential decay model. This model suggested that the

PL originates from two distinct states with times constants of t1 = 3.908 ns ± 0.025 ns and

t2 = 11.771 ns ± 0.069 ns. It also suggests almost 94 % of PL decay within 3.908 ns. This

is an indication of an efficient electron-hole pair trapping in CdS-1. The PL kinetics of

CdS-2 was best fitted with a tri-exponential model, suggesting time-constants of t1=1.155

ns ± 0.067 ns, t2= 4.078 ns ± 0.076 ns and t3= 12.799 ns ± 0.143 ns. The PL life-times of

CdS-2 is ~ 1 ns longer than the PL of CdS-1. Thus, signify the higher photocatalytic

activity of CdS-2 than CdS-1, because the photo-injected electron hole pairs of CdS-2

have longer recombination time and hence the efficiency of electron transfer from the

CdS-2 is higher than CdS-1. Additionally, the CdS-2 (nanocapsule) also showed good

adsorption of dye on surface as was observed during the establishment of adsorption

equilibrium between dye and catalyst (Figure 3.32d). Furthermore, CdS-2 (nanocapsules)

was found better photocatalyst than analogous reported ones (Table 3.9) [65, 75-79].

Table 3.9: Reported literature data of Congo red dye degradation by CdS and

hybrid CdS NPs.

S.No

.

CdS/CdS Based

material

Catalyst

loading

(mg/mL)

Dye

Concentration

mg/L

Time

(minutes)

%

Degradation

Reference

1 CdS/Dendrite 0.125l 10 60 91 [75]

2 CdS/nanospheres 0.5 10 120 92 [76]

3 CdS/nanocapsules 0.5 50 30 99 This work

4 CdS-MMA/

agglomerates of

ultrafine particles

0.5 30 90 94.34 [77]

5 CdS- PEG/

agglomerates

of ultrafine

particles

0.5 30 90 97 [78]

6 CdS- Chitosan/

Spherical

1.5 20 180 85.9 [65]

7 CdS-Ti4O9 1 50 45 100 [79]

116

3.6 Solar-light driven photocatalytic conversion of p-nitrophenol to p-aminophenol

on CdS nanosheets and nanorods

The organic contaminants present in the agricultural and industrial waste-water is

adversely affecting the environment. The nitro aromatic compounds are probably the

more frequent and stable enough against the natural degradation process due to their

chemical and biological stability. Among the nitroaromatic compounds p-nitrophenol is a

major organic pollutant, so the conversion of this disagreeable chemical to p-

aminophenol is of great industrial significance as the latter is used for the production of

aniline, paracetamol, phenacetin and acetanilide [80].

This reaction of environmental and pharmaceutical value was first reported by Pal in

2002 in the presence of Ag nanoparticles and BH-4 [81]. After then many other noble

metals such as Pt, Pd, Ru, Au and certain metal alloys have been used to catalyze this

reaction [82]. Nitrogen doped graphene, metal organic frameworks (MOFs) and

nanoparticles embedded metal organic frameworks have also been studied as catalyst for

the conversion of p-nitrophenol to p-aminophenol [83-85]. However, the industrial

applicability of these catalysts is hampered by their high cost, toxicity and recyclability.

Therefore, the development of an economical and efficient reaction route is yet an

unresolved dilemma. The use of metal chalcogenides as photocatalyst, which harvest the

abundant and inexhaustible solar energy, is a green and promising strategy for such

reactions.

There are only few reports regarding the visible light driven photocatalytic conversion of

p-nitrophenol to p-aminophenol. Agileo et al used CdS nanorods and nanofibers for this

reaction and the conversion was achieved in 270 minutes and 30 minutes, respectively

[86, 87]. Pahari et al tested CdS nanoflowers for this reaction under the household CFL

lamp, and achieved the conversion in 90 minutes [88]. Few research groups also tested

CdS-based composite materials under visible light, including CdS-TiO2, CdS/graphene

hybrid and noble metal deposited CdS [89-91]. The photocatalytic conversion of p-

nitrophenol to p-aminophenol is initiated with the amination of nitro group, which takes

place by taking the conduction band electrons from CdS and H+ from NaBH4. The photo-

generated hole in the CdS valence band is filled by electrons generated by the

decomposition of NaBH4.

117

Among the various morphologies of as synthesized CdS in this work, we tested CdS

nanosheets (CdS-4) and nanorods (CdS-5) for solar light-driven photocatalytic

conversion of the p-nitrophenol to p-aminophenol using water as a solvent and NaBH4 as

a reducing agent. The p-nitrophenol shows absorption maxima at 317 nm in water, which

shifts to higher wavelength (400 nm) upon NaBH4 addition due to the formation of

phenolate ion of the light yellow color.

Figure 3.34: UV-Visible spectral changes associated with photoreduction of p-

nitrophenol to p-aminophenol with irradiation time in the presence of

CdS-4 (a) 5 mg catalyst dose (b) 10 mg catalyst dose (c) 15 mg catalyst

dose (d) 20 mg catalyst dose.

The conversion progress was spectrophotometrically monitored; the intensity of the

characteristic peak of p-nitrophenol at 400 nm depleted gradually with irradiation time

118

accompanied by a parallel growth of a new peak at 298 nm (Figure 3.34a-d) for p-

aminophenol [92, 93]. The absence of peaks at 388 or 302 nm due to 4-benzoquninoe

monoxime or 4-nitrosophenol demonstrate the clean conversion process without

generating any by product [94].

To study the effect of catalyst amount, an important parameter in catalysis, the amount of

CdS nanosheets was varied from 5 mg to 20 mg keeping the other experimental

conditions the same. A sharp rise in photoreduction rate of the p-nitrophenol was

observed up to 15 mg, a dose at which approximately 100% reaction occurred in 10

minutes (Figure 3.35a). However, beyond this amount (20 mg) the photoreduction rate

was adversely affected (12 min). Low catalytic activity at high catalyst load can be

explained in terms of particles aggregation (reduction in the number of active catalyst

sites), light scattering and screening effect of the dense particles, reducing the light flux

received by the inner particles [95]. Furthermore, under similar conditions, 15 mg CdS

nanorods were observed to do the conversion in 8 min (Figure 3.35b), performance better

catalytic than the nanosheets.

The linear correlation between ln(c/co) and t for both nanostructures (Figure 3.35c)

suggest that the reduction of nitrophenol follows pseudo first order kinetics. The

measured k values for nanosheets and nanorods are -0.3638 min-1

and -0.4035 min-1

,

respectively confirming the relatively fast kinetics for the latter. It was further noted that

kapp (-0.2028 min-1

to -0.3638 min-1

) increased by the gradual increase in the catalyst

amount up to 15 mg, but diminished beyond this (-0.2727 min-1

at 20 mg).

The reusability or stability of a catalyst is a seminal factor for its practical application. In

this work, the reusability of both catalysts was measured for 15 mg load (optimum

amount) for three consecutive cycles. The results show slight decrease in photocatalytic

activities in the initial cycle, though remain above 92% even in the 3rd

cycle (Figure

3.35d). X-ray diffraction (XRD) patterns indicate that the crystal structure of both

morphologies is undamaged after the reaction (Figure 3.29b). The better recyclability

results and intact crystal structures before and after performing the activity designate

stability of both NPs to photo-corrosion. Moreover, our results are comparable with the

119

Figure 3.35: (a) The photocatalytic conversion of p-nitrophenol to p-aminophenol with

irradiation time at different catalyst (CdS-4) dose (b) UV-Visible spectral

changes associated with photoreduction of p-nitrophenol to p-aminophenol

with irradiation time in the presence of CdS-5 (15 mg catalyst dose) (c) The

photoconversion kinetics of p-nitrophenol to p-aminophenol with

irradiation time in the presence of CdS-1 (nanosheets) at different catalyst

dose and CdS-2 (nanorods) (d) Repetitive photocatalytic conversion of p-

nitrophenol to p-aminophenol during three sequential cycles using 15 mg

of catalyst dose of both CdS-4 and CdS-5 NPs.

precious noble metals and better than the previously reported CdS (Table. 3.10) [86-91,

96]. The relatively high photocatalytic activity of both anisotrophically grown CdS

120

morphologies can be most probably ascribed to higher quantum confinement effect that

promotes the interfacial charge transfer, large surface area and good optical absorbance.

Table 3.10: Comparison of our results with reported literature data of

photocatalytic p-nitrophenol reduction to p-aminophenol by pure

CdS and CdS based materials.

S.No CdS/CdS based materials Light source Time taken for 100%

reduction

Ref

1 CdS-nanosheets Natural solar light 10 min This work

2 CdS-nanorods (D=5,L=80 Natural solar light 8 min This work

3 CdS-nanorods(D=16,L=75) 3 W LED lamp

(blue light)

270 min [86]

4 CdS-nanofibres 3 W LED lamp

(blue light)

30 min [87]

5 CdS-flower 40 W CFL lamp 90 min [88]

6 CdS quantum dots-TiO2

nanotubes array

300 W Xe lamp 60 min [89]

7 CdS nanospheres /graphene

hybrid

300 W Xe arc lamp 16 min (N2 atm,

NH4HCO2 as sacrificial

agent)

[90]

8 Au deposited CdS

nanostructures

150 W halogen lamp 240 min [91]

9 CdS nanowires-reduced

graphene oxide

nanocomposite

300 W Xe arc lamp 20 min (N2 atm,

NH4HCO2 as

sacrificial agent

[96]

121

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131

Conclusions

In summary, electrochemical (MoSx) and single source precursor (HgS, CdS and ZnS)

methods were used to generate metal sulfide NPs of the metals Mo, Hg, Cd and Zn. MoSx

was electrochemically deposited on polyethylenimine (PEI) modified reduced graphene

oxide (rGO-PEI-MoSx) substrate and tested as electrocatalyst for the reduction of CO2 to

CO in CO2 saturated aqueous NaHCO3 solution with high efficiency and selectivity. The

FT-IR analysis of the substrate (rGO-PEI) revealed covalent interaction between rGO and

PEI, while SEM and TEM results showed the uniform growth of MoSx nanoparticles on

rGO-PEI surface. The Fast Fourier Transform (FFT) pattern indicated the amorphous

nature of MoSx NPs, which was further confirmed by XRD pattern of the material as no

characteristic peaks of any crystalline structures were observed. The catalyst was found

to reduce CO2 to CO at overpotential as low as 140 mV and reached a maximum faradaic

efficiency of 85.1% at overpotential of 540 mV resulted in high TOF value (2.4 s-1

) for

the formation of CO. At overpotential of only 290 mV with respect to CO formation, this

catalyst is also capable of producing syngas (2H2:1CO). The catalyst is highly stable in

the long-term electrolysis (16 h). The efficiency and selectivity towards CO2 reduction

rather than hydrogen evolution at the optimal applied potential are attributed to the

synergetic combination of the properties of MoSx and PEI.

The dithiocarbamate complexes of Hg(II), Cd(II) and Zn(II) were easily synthesized at

ambient conditions by simply mixing respective metal salt and corresponding sodium

dithiocarbamate. The single crystal XRD analysis of HgL2, ZnL6 and ZnL7 showed their

growth into dimeric molecular structure with distorted square pyramidal geometry while

HgL1 showed monomeric molecular structure with distorted tetrahedral geometry. The as

synthesized complexes were used as molecular precursors for the synthesis of ZnS, CdS

and HgS NPs by simply refluxing the respective precursors with en under ambient

pressure and temperature without the use of any external toxic surfactant; a safe, green

and one pot synthetic approach. The SEM and TEM images showed precursor based

significant variation in the size and morphology of CdS and HgS NPs. The formation of

different sizes and morphologies of the resultant NPs was governed by stability

(thermolysis rate) of the precursors (determined by DFT studies), solubility in en and in

132

situ generated capping agents (FT-IR analysis of NPs). In contrast, the similar

dithiocarbamate complexes of Zn(II) did not produce substantial morphological disparity

and size control upon thermolysis in en. The possible reason may be the low atomic mass

and size of Zn as compared to Cd and Hg. The crystallinity of the resultant NPs was

found to increase down the group from ZnS to HgS and may be correlated with

increasing size and atomic weight of the metal. The absence of any detectable impurities

in the XRD pattern of all the NPs and the presence of M (M = Zn, Cd, Hg) to S in 1:1

ratio as shown by EDS affirm the effectiveness of the synthetic method for the formation

of pure metal sulfide NPs. The band gap of CdS NPs (as calculated from their UV/

Visible absorption spectroscopy) corresponds well with the spectrum of sunlight,

therefore, exhibited remarkable performance (especially, the anisotrophically grown NPs)

as solar light driven photocatalyst for the degradation of Congo red dye and reduction of

p-nitrophenol to p-aminophenol. In nutshell, these catalysts in current form or after

appropriate modification can be used for organic transformation reactions of

pharmaceutical and industrial significance, clean energy production from renewable

sources and environmental remediation.

133

List of Publications

1. Azam Khan, Zia-ur-rehman, Abdullah Khan, Hina Ambreen, Haseeb Ullah,

Sayed Mustansar Abbas, Yaqoob Khan and Rajwali Khan. Solar-light driven

photocatalytic conversion of p-nitrophenol to p-aminophenol on CdS nanosheets

and nanorods. Inorg. Chem. Commun. 79 (2017) 99-103.

2. Azam Khan, Haseeb Ullah, Jamal Abdul Nasir, Suzanne Shuda, Chen Wei, M.

Abdullah and Zia-ur-rehman. Metal- and Carbon-based materials as heterogeneous

electro-catalysts for CO2 reduction. J. Nanosci. Nanotecnol. 18 (2018) 3031-3048.

3. Azam Khan, Zia-ur-rehman, Muneeb-ur-Rehman, Rajwali Khan, Zulfiqar, Amir

Waseem, Azhar Iqbal, and Zawar Hussain Shah. CdS nanocapsules and

nanospheres as efficient solar light-driven photocatalysts for degradation of

Congo red dye. Inorg. Chem. Commun. 72 (2016) 33-41.

4. Fengwang Li, Shu-Feng Zhao, Lu Chen, Azam Khan, Douglas R. MacFarlane,

and Jie Zhang. Polyethylenimine promoted electrocatalytic reduction of CO2 to

CO in aqueous medium by graphene-supported amorphous molybdenum

sulphide. Energy Environ. Sci. 9 (2016) 216-223.

5. Zia-ur-Rehman, Shaista Ibrahim, Azam Khan, Muhammad Imran, Muhammad

Moazzam Naseer, Imran Khan, Afzal Shah, Muhammad Nawaz Tahir, Muneeb-

ur-Rahman, and Iqra Zubair Awan. Homobimetallic zinc (II) dithiocarbamates:

synthesis, characterization and in vivo antihyperglycemic activity. J. Coord.

Chem. 69 (2016) 551-561.

6. Zafar Ali Khan Khattak, Zia‐ur Rehman, Azam Khan, Afzal Shah, Waqar Zaid,

Amin Badshah, M. Fettouhi, and Atif Fazal. "Self‐Assembled Heteroleptic Zn (II)

Dithiocarbamate‐Based 2D‐Interwoven Supramolecular Giant Macrocycles and

Their Redox Properties. Heteroat. Chem. 25 (2014) 238-244.

7. Jamal Abdul Nasir, Shaheen Gul, Azam Khan, Abrar Ahmad, Zulfiqar, Zia-ur-

Rehman. Efficient Solar Light Driven Photocatalytic Degradation of Congo Red

Dye on CdS Nanostructures Derived from Single Source Precursor. J. Nanosci.

Nanotecnol. 18 (2018) 7405-7413.

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8. Jamal Abdul Nasir, Hina Ambareen, Azam Khan, M. Abdullah Khan, Wei Chen,

M. S. Akhter, and Zia-ur-Rehman. Photoreduction of 4-Nitrophenol to 4-

Aminophenol Using CdS Nanorods . J. Nanosci. Nanotechnol. 18 (2018) 7516–

7522