Redox behavior, spectroscopic characterization and ...
Transcript of Redox behavior, spectroscopic characterization and ...
Redox behavior, spectroscopic characterization and
biological activities of some new robustly electro-active
compounds
A Thesis Submitted to the Department of Chemistry, Quaid-i-Azam University, Islamabad, in partial fulfillment
of the requirements for the degree of
Doctor of Philosophy
in
Physical Chemistry
By
ABDUR RAUF
Department of Chemistry Quaid-i-Azam University
Islamabad, Pakistan 2016
II
Dedicated to
My Beloved
Parents & Teachers
Acknowledgements
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All adoration and glory to the Almighty Allah who is gracious and clement upon us with every
breath in our life. Myriad of Darood and Salam upon the torch bearer of mankind, Holy Prophet
Muhammad صلى الله عليه وسلم who is the only complete and absolute personality in the mankind. I feel proud
of being apostle of the greatest personality of the world.
I am highly grateful to my affectionate supervisor, Dr. Afzal Shah, Assistant Professor,
Department of Chemistry, Quaid-i-Azam University, Islamabad. His comprehensive knowledge
and logical discussion was of great value and it was a wonderful period of learning not only in
perspective of scientific aspect but also in management of my work. His perpetual inspiration
and friendly behavior assisted me to complete this study. I am highly grateful to Prof. Dr.
Muhammad Siddiq, Chairman, Department of Chemistry, Quaid-i-Azam University,
Islamabad, for providing all the necessary facilities during my research work. I am highly
indebted to Prof. Dr. Saqib Ali for guiding me in the synthetic work, Dr. Hidayat Hussain for
providing quinones and Dr. Rumana Qureshi and Dr. Zahid Rashid for their help in
computational studies. Their cheering attitude and fruitful discussion will always be
acknowledged. Many thanks to Prof. Dr. Irfan Zia Qureshi, Qamar Abbas and Abdul Aziz,
Department of Animal Sciences, Quaid-i-Azam University, Islamabad, for the antidiabetic
studies. I am also thankful to Dr. Rashda Abbasi, KRL hospital, Islamabad for her timely and
enthusiastic cooperation for cytotoxicity and antioxidant activities. Countless thanks to Dr. Zia-
ur-Rehman and crystallographer, Dr. Francine Belanger-Gariepy, Department de Chimie,
Universite de Montreal, Montreal, Canada for single crystal analysis (data collection, structure
solution and structure analysis) and fruitful collaboration. Many thanks and acknowledges to the
Higher Education Commission, Islamabad, Pakistan for providing me the indigenous Ph.D.
scholarship.
During the completion of this work I collaborated with many colleagues for whom I have
great regard and I wish to extend my warmest thanks to Mr. Saghir Abbas and Mr. Khurram
Shahzad Munawar and my lab fellows Dr. Imdad Ullah, Mr. Latif-ur-Rehman, Mr. Zaheer
Ahmad, Amir Hassan Shah, Mr. Khurshid Ahmad and Mr. Asif at the Department of
Chemistry, Quaid-i-Azam University, Islamabad. I shall always reminisce their laudable
assistance and encouragement during my studies.
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I am greatly honored to mention the nice cooperation of all employees and para-teaching
staff of the department, especially Mr. M. Sharif Chohan, Mr. Rana Tahir and Mr. Rana
Matloob for providing chemicals and computer facilities and Boys hostel administration for
providing me wonderful place to stay during my studies.
I am sincerely thankful to my beloved father, mother and all other family members, without
their cooperation I was not able to complete this work. Heartiest prayers for my parents, who
always encouraged, supported and prayed for my success.
Abdur Rauf
Abstract
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The correlation of biological activities of compounds with their redox properties is the
subject of extensive investigations of bioelectrochemists. Schiff bases, quinones and
naphthalenes contain electroactive moieties and their broad range biological activities
are closely related with the ability of these compounds to donate and/or accept
electrons. We synthesized Schiff bases 1-((4-bromophenylimino) methyl) naphthalen-
2-ol (BPIMN) and 1-((2,4-dimethylphenylimino)methyl)naphthalen-2-ol (HL) and
used HL as a ligand for the preparation of its metallic complexes. All the synthesized
compounds were confirmed by 1H NMR, 13C NMR, FTIR, TGA and UV–Vis
spectroscopy. Structures of Schiff bases were also characterized by X-ray analysis and
the experimental findings were supported by quantum mechanical calculations. The
results of BPIMN were compared with a structurally related Schiff base, 1-((4-
chlorophenylimino) methyl) naphthalen-2-ol (CPIMN). The photometric and
electrochemical fate of all these Schiff bases were investigated in a wide pH range
and the obtained results helped in proposing the redox mechanistic pathways. The
synthesized compounds were subjected to numerous biological applications and the
results revealed that Schiff base HL and its metal complexes other than oxovanadium
complex remarkably decrease the blood glucose, triglyceride and cholesterol levels.
The metal complexes were found to exhibit significant inhibition against alkaline
phosphatase enzyme as compared to Schiff bases. The zinc complex was found as the
most potent inhibitor of bacteria/fungi while the vanadyl product displayed the least
activity among all the metal coordinated products.
Quinones, another biologically important class were also investigated due to
their robust electrochemical properties and wide range of biological activities. The
laxative and therapeutic activities of quinones are related to their redox
characteristics. Electrochemically unexplored hydroxy substituted quinones including
4-hydroxy-5-methoxynaphthalene-1-ylacetate (HMNA), 1,4-dihydroxy-2-(3-hydroxy-
3-(trichloromethyl)pentyl)-8-methoxyanthracene-9,10-dione (HCAQ), 1-hydroxy-
2(hydroxymethyl)anthracene-9,10-dione (HAC), 1,8-dihydroxy-4,5-dinitro
anthracenedione (DHDN) 4,8-dihydroxy-9,10-dioxo-9,10-dihydroanthracen-1-yl
acetate (HACAD), 1,4,5-trihydroxyanthracene-9,10-dione (HAD) and 1,4,5-
trihydroxy-2-methyl-3-(3-oxobutyl)anthracene-9,10-dione (HOAD) were selected and
their redox behavior was studied in a wide pH range using modern
electrochemical techniques. Kinetic parameters such as diffusion coefficient and
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heterogeneous electron transfer rate constant and thermodynamic parameters of the
electron transfer processes such as ∆G#, ∆H# and ∆S# were electrochemically
evaluated. Their redox mechanisms were proposed on the basis of experimental
findings supported by computational calculations. Moreover, a detailed UV–vis
spectroscopy was carried out in a wide pH range for photometric characterization
and acid-base dissociation constant, pKa determination.
Though naphthalene by itself is toxic, however, some of its derivatives are bestowed
with medicinal properties. Two biologically important naphthalene derivatives,
naphthalene-2,3-dicarboxylic acid (NDA) and 1,8-dimethoxynaphthalene (DMN)
were characterized by electrochemical techniques and screened for their antioxidant
and anti-diabetic activities. NDA was found less toxic to HeLa cells and biological
antioxidant studies revealed it as a more effective antioxidant as compared to DMN
and standard antioxidant, ascorbic acid. Both NDA and DMN significantly increased
the cholesterol level in blood but showed varied biological activities as regards to
glucose and triglyceride concentrations. The cytotoxicity results evidenced DMN to
significantly inhibit the cell proliferation in a dose dependent manner. Like the
biological antioxidant studies, the electrochemical results also witnessed NDA as
stronger antioxidant than DMN. pH dependent oxidation of NDA revealed its
antioxidant role to be exerted both by the donation of electrons and protons. Although
the oxidation potential of NDA is greater than the widely used natural antioxidant,
ascorbic acid, yet it is capable of donating two electrons as compared to one electron
donating ability of ascorbic acid. The redox mechanistic pathways proposed in this
work are expected to provide useful insights about the unexplored mechanisms by
which Schiff bases, quinones and naphthalenes exert their biochemical actions.
Contents
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Chapter 1-Introduction…………………………………………. 1
1.1 Schiff Bases……………………………………………………………. 2
1.2 Quinones……………………………………………………………….. 5
1.3 Naphthalene Derivatives………………………………………………. 9
References……………………………………………………………... 11
Chapter 2- Theoretical Background of the
Electrochemical Techniques........................................................ 20
2.1 Cyclic Voltammetry…………………………………………………… 20
2.1.1 Basic Principle of Cyclic Voltammetry……………………………….. 20
2.1.2 Nature of Electron Transfer Process…………………………………... 21
2.1.3 Reversible Process Diagnostic Criteria of Reversibility…………….. 21
2.1.4 Quasi-Reversible Process……………………………………………… 22
2.1.5 Irreversible Process……………………………………………………. 23
2.1.6 Multi-Electron Transfer Process………………………………………. 24
2.1.7 Types of Two-Electron Transfer Reactions…………………………… 25
2.2 Differential Pulse Voltammetry (DPV)………………………………... 26
2.3 Square wave voltammetry (SWV)…………………………………….. 27
References……………………………………………………………... 29
Chapter 3-Experimental………………………………………... 32
3.1 Chemicals and reagents……………………………………………….. 32
3.2 Synthesis of Schiff Bases and complexes……………………………... 32
3.2.1 1-((4-bromophenylimino) methyl) naphthalen-2-ol (BPIMN)………... 32
3.2.2 1-((4-chlorophenylimino) methyl) naphthalen-2-ol (CPIMN)..……….. 33
3.2.3 1-((2, 4-dimethylphenylimino)methyl)naphthalen-2-ol (HL)…………. 33
3.2.4 Oxovanadium(IV) complex of 1-((2, 4-dimethylphenylimino)methyl)
naphthalen-2-ol (L2VO)……………………………………………….. 34
3.2.5 Dibutyltin complex of 1-((2, 4-dimethylphenylimino)methyl)
naphthalen-2-ol (L2Sn)…..…………………………………………….. 35
3.2.6 Zinc complex of 1-((2, 4-dimethylphenylimino)methyl)naphthalen-2-ol
(L2Zn)…..…………………………………………………………………. 35
3.2.7 Cobalt complex of 1-((2, 4-dimethylphenylimino)methyl)naphthalen-2-ol
(L2Co)….……………………………………………………………………... 36
3.3 Hydroxy substituted quinones…………………………………………. 37
Contents
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3.4 Naphthalene derivatives……………………………………………….. 39
3.5 Voltammetric studies…………………………………………………... 39
3.6 Spectroscopic measurements…………………………………………... 40
3.7 Computational Studies………………………………………………… 40
3.8 Anti-diabetic activity…………………………………………………... 40
3.8.1 Animals and maintenance……………………………………………... 41
3.8.2 Induction of diabetes…………………………………………………... 41
3.8.3 Experimental design…………………………………………………… 41
3.8.4 Control and experimental groups……………………………………… 41
3.8.5 Glucose determination…………………………………………………. 41
3.8.6 Total serum cholesterol and triglycerides determination……………… 42
3.8.7 Statistical analysis……………………………………………………... 42
3.9 Cytotoxicity and antioxidant activity………………………………….. 42
3.9.1 Cell lines and cell culture……………………………………………… 42
3.9.2 Cell survival and proliferation assay…………………………………... 42
3.9.3 Biological antioxidant activity: lipid peroxidation…………………….. 43
3.10 Assay of alkaline phosphatase activity………………………………… 44
3.11 Antimicrobial
ac4tivities………………………………………………... 44
3.11.1 Growth medium, culture and inoculum preparation…………………... 44
3.11.2 Antimicrobial assay by disc diffusion method………………………… 45
References……………………………………………………………... 46
Chapter 4- Result & Discussion…...…………………………... 50
4.1 Synthesis, photometric and electrochemical fate of Schiff bases……... 50
4.1.1 Structural characterization……………………………………………... 50
4.1.2 Crystal structures description of BPIMN and CPIMN………………… 50
4.1.3 UV-vis spectroscopy of BPIMN and CPIMN…………………………. 54
4.1.4 Cyclic voltammetry of BPIMN and CPIMN…………………………... 56
4.1.5 Differential pulse voltammetry of BPIMN and CPIMN………………. 58
4.1.6 Square wave voltammetry of BPIMN and CPIMN…………………… 59
4.1.7 Redox mechanism of BPIMN and CPIMN……………………………. 59
4.1.8 Computational study of BPIMN and CPIMN…………………………. 61
4.1.9 Inhibition of ALP……………………………………………………… 62
Contents
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4.2 Synthesis, redox behavior and biological applications of a novel Schiff
base 1-((2,4-dimethylphenylimino)methyl)naphthalene-2-ol (HL) and
its metallic derivatives…………………………………………………. 63
4.2.1 Structural Confirmation of HL and its metallic complexes…………… 63
4.2.2 Crystal structures description of HL…………………………………... 63
4.2.3 Electronic absorption spectroscopy of HL and complexes……………. 65
4.2.4 Cyclic voltammetry of HL and complexes……………………………. 66
4.2.5 Square wave voltammetry of HL and complexes……………………... 69
4.2.6 Differential Pulse Voltammetry of HL and complexes………………... 70
4.2.7 Redox mechanism of HL and complexes….….……………………….. 75
4.2.8 Anti-diabetic Activity of HL and complexes…….……………………. 78
4.2.9 In Vitro studies: Cytotoxicity analysis on HeLa………………………. 80
4.2.10 Inhibition of ALP of HL and complexes………......………………....... 83
4.2.11 Antimicrobial activities of HL and complexes...……………………… 84
4.2.12 Computational study of HL and complexes…………………………… 85
4.3 Redox behavior of a novel menadiol derivative………………………. 88
4.3.1 Cyclic voltammetry of HMNA………………………………………... 88
4.3.2 Differential pulse voltammetry of HMNA…………………………….. 92
4.3.3 Square wave voltammetry of HMNA…………………………………. 93
4.3.4 Redox mechanism of HMNA………………………………………….. 94
4.3.5 Electronic absorption spectroscopy of HMNA………………………... 97
4.3.6 Computational studies……………………...……………………...…... 99
4.4 Redox behavior of two anthracenedione derivatives………………….. 102
4.4.1 Cyclic voltammetry of HAC and DHDN……………………...………. 102
4.4.2 Differential pulse voltammetry of HAC and DHDN…………………. 105
4.4.3 Square wave voltammetry HAC and DHDN…………………….......... 108
4.4.4 Proposed redox mechanism HAC and DHDN……………………........ 110
4.5 Redox mechanism, kinetic and thermodynamic parameters of a novel
anthraquinone……………………...……………………...…………… 113
4.5.1 Cyclic voltammetry of HCAQ……………………...…………………. 113
4.5.2 Differential pulse voltammetry of HCAQ……………………...……… 119
4.5.3 Square wave voltammetry of HCAQ……………………...…………... 121
4.5.4 Electronic absorption spectroscopy of HCAQ……………………........ 122
Contents
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4.5.5 Redox mechanism of HCAQ……………………...…………………… 125
4.6 Temperature responsive redox behavior of hydroxyanthracenediones... 127
4.6.1 Cyclic voltammetry……………………...……………………...……... 127
4.6.2 Evaluation of kinetic parameters from Cyclic voltammetry…………... 128
4.6.3 Thermodynamic parameters from cyclic voltammetry………………... 130
4.6.4 Differential pulse voltammetry……………………...………………… 132
4.6.5 Square wave voltammetry……………………...…………………….... 135
4.6.6 Electronic absorption spectroscopy……………………...…………….. 137
4.6.7 Redox Mechanism……………………...……………………................ 140
4.7 Biological activity, pH dependent redox behavior and UV–Vis
spectroscopic studies of naphthalene derivatives. …………………….. 145
4.7.1 Electrochemical studies……………………...……………………........ 145
4.7.2 Electronic absorption spectroscopy……………………......................... 150
4.7.3 Anti-hyperglycemic activity……………………...……………………. 151
4.7.4 Cytotoxicity analysis on HeLa……………………...………………… 153
4.7.5 Biological antioxidant activity……………………...…………………. 155
Conclusions…..……………...……………………...……………… 157
References…………………...……………………...………………. 160
List of Figures
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Fig. 2.1 (a) Potential sweep in cyclic voltammery (b) A typical cyclic
voltammogram for an electroactive species…………………………. 21
Fig. 2.2 Cyclic voltammogram of Quasi-reversible redox processes………. 23
Fig. 2.3 Cyclic voltammogram for a reversible two-step system (a) ∆E = -180
mV (b) ∆E = -90 mV (c) ∆E = 0 mV. (d) ∆E = 180 mV……… 24
Fig. 2.4 (a) Scheme of application of potential (b) General differential pulse
voltammogram………………………….……………………………. 26
Fig. 2.5 SWV showing total, forward and backward current……………….. 27
Fig. 4.1 Ball and sticks diagrams of BPIMN (A) and CPIMN (B)…………... 50
Fig. 4.2 (a) 2D microporous structure of the compound BPIMN mediated by
C7…Br1 = 3.311 Å and H13…O1= 2.482 Å (b) 2D microporous
structure of the compound CPIMN mediated by Cl1…C7 = 3.323 Å
and H13…O1 = 2.481 Å………………………….…………………. 53
Fig. 4.3 UV-Vis spectra 6of 20 μM BPIMN and CPIMN
………………………………………….…………………………….. 54
Fig. 4.4 Electronic absorption spectra of 20 μM BPIMN (A) and 20 μM
CPIMN (B) recorded in different pH media……………………….. 55
Fig. 4.5 (A)The CVs of CPIMN recorded at 100mVs-1 starting from 0 V (a) -
0.6 V (b) and -1.4 V (c) with 1mM concentration of analyte in 50%
ethanol and 50% BR buffer of pH-7, (B) Voltammogram of BPIMN
and CPIMN obtained under same conditions……………………… 57
Fig. 4.6 Differential pulse voltammograms for oxidation of CPIMN at 10
mVs-1 in 50% ethanol and 50% BR buffers of wide pH range used as
supporting electrolyte………………………………………. 58
Fig. 4.7 Multiple scan SWVs of 1mM CPIMN recorded at 100 mV s-1 in 50%
ethanol and 50% BR buffer of pH-7, used as supporting
electrolyte………………………….………………………………… 59
Fig. 4.8 Concentration dependent inhibition of alkaline phosphatase (ALPs)
by BPIMN and CPIMN………………………….…………………... 62
Fig. 4.9 Ellipsoid diagram of Schiff base HL……………………………….. 63
Fig. 4.10 UV-Vis spectra of 20 μM HL, (L2VO), (L2Sn), (L2Zn) and (L2Co).. 65
Fig. 4.11 UV-Vis spectra of 20 μM HL (A) and 30 μM (L2Sn) (B) in different
media of pH range 2-12………………………….…………………
66
List of Figures
x
Fig. 4.12 Cyclic Voltammograms of (A) HL in a buffer of pH 4.0, 7.0 and 10.0
(B) HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) in a buffer of pH-7 at
100 mVs-1………………………………………………………….. 67
Fig. 4.13 Cyclic Voltammograms of (L2Zn) at 100mVs-1 in a buffer of pH 4.0,
7.0 and 10.0………………………….…...……………………..……. 68
Fig. 4.14 Plot of Ip vs concentration for (L2Sn) in a buffer of pH 7.0………….. 68
Fig. 4.15 Square wave voltammograms of different concentrations of (L2Sn)
obtained at a scan rate 100 mV/s in pH 7.0, inset shows Ip as a
function of concentration………………………….………………… 69
Fig. 4.16 Differential pulse voltammograms recorded in different pH media at
5 mVs-1 showing oxidation of HL (A) and (L2Co) (B) …………….. 71
Fig. 4.17 Plots of oxidation potentials versus pH for HL (A) and (L2Co) (B)…. 71
Fig. 4.18 Differential pulse voltammograms recorded in different pH media at
5 mVs-1 showing reduction of HL (A) and (L2Sn) (B) ……………… 72
Fig. 4.19 Differential pulse voltammograms recorded in different pH media at
5 mVs-1 showing reduction of (1) (A) and (L2Zn) (B) ……………... 73
Fig. 4.20 Differential pulse voltammograms recorded at 5 mVs-1 showing
forward and reverse scans of reduction process of (L2Zn) at pH-6.0... 74
Fig. 4.21 Plots of Epc versus pH for HL (A), (L2Zn) (B) and (L2Sn) (C)…...… 74
Fig. 4.22 Plasma glucose in mice treated with HL, (L2VO), (L2Sn), (L2Zn) and
(L2Co) compound groups……………...……………………............... 79
Fig. 4.23 (A) Triglyceride and (B) Cholesterol level in mice treated with HL,
(L2VO), (L2Sn), (L2Zn) and (L2Co) compound groups……………… 80
Fig. 4.24 Effect of concentration of HL and its complexes (L2VO), (L2Sn),
(L2Zn) and (L2Co) on HeLa cells…………………………………… 81
Fig. 4.25 Viability curves for HL, (L2VO), (L2Sn), (L2Zn) and (L2Co)
complexes in HeLa…………………………………………………... 82
Fig. 4.26 Inhibition profile of alkaline phosphatase (ALPs) at different
concentrations of HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) ………… 83
Fig. 4.27 Numbering scheme of HL used in computational studies……….. 86
Fig. 4.28 CVs (scan 1–5) of 0.2 mM HMNA obtained at 100 mV s-1 in a
solution of pH 4.0, using 0.1 M acetate buffer…………………….
88
List of Figures
xi
Fig. 4.29 CVs obtained in Ar saturated solution of 0.2 mM HMNA at a GCE at
100 mVs-1 in different supporting electrolytes of pH ranging from
1.22 to 7.0 (A) and 8.07-12.93 (B) ………………………………..
89
Fig. 4.30 (A) CVs obtained with the GCE in a solution of 0.2 mM HMNA
in pH 4, 0.1 M acetate buffer at different scan rates ranging from
20-400 mV s-1 (B) CVs obtained with the GCE in a solution of
0.2 mM HMNA in pH 7, 0.1 M Phosphate buffer at different scan
rates………………………….…………………………… 90
Fig. 4.31 (A) Plot of Ipa (µA) vs. square root of scan rate (V/s) at pH 4.0 (B)
Plot of Ipa (µA) vs. square root of scan rate (V/s) at pH 7.0……. 90
Fig. 4.32 (A) Plots of log Ipa (µA) vs. log ν (V/s) at pH 4.0 and (B) pH 7.0 91
Fig. 4.33 (A) CVs of different concentration of HMNA (B) Plot of Ipa (µA)
vs. concentration (µM/cm3) obtained at 100 mV s-1 scan rate in a
medium buffered at pH 4.0 using 0.1 M acetate buffer………..……. 92
Fig. 4.34 DPVs of 0.2 mM HMNA obtained at 10 mVs-1 scan rate in
supporting electrolytes of pH 1.22 to 12.93…………………………. 93
Fig. 4.35 SW voltammograms obtained in solution of 0.2 mM HMNA in pH
4.2 (A) 1st scan showing total current (It), forward current (If) and
backward current (Ib) (B) SWVs of first three scans of 0.2 mM
HMNA recorded at 100 mVs-1 scan rate in pH 7.4………………… 94
Fig. 4.36 HPLC of HMNA carried out at 300 nm in acidic and neutral
conditions………………………….……………………………….. 96
Fig. 4.37 Absorption spectra of 25 µM HMNA obtained in pH 4.0-9.2 using
20% ethanol and 80% aqueous supporting electrolytes………….. 99
Fig. 4.38 (A) Absorption spectra of 25 µM HMNA obtained in pH 9.2-10.2
using solvent system of 20% ethanol and 80% aqueous supporting
electrolytes (B) Plot of absorbance of HMNA vs. pH at 300 nm 99
Fig. 4.39 (A) Charge distribution of HMNA and (B) 5-methoxynaphthalene-
1,4-diol…………………….………………………….…………… 101
Fig. 4.40 (A) Graphical representation of EHOMO of HMNA and (B) 5-
methoxynaphthalene-1,4-diol………………………….………….. 101
Fig. 4.41 CVs of 0.2 mM solution of HAC scan-I (1) and scan-II (2) recorded
in solution of pH 7.0 at 100 mVs-1 scan rate…………………………. 102
List of Figures
xii
Fig. 4.42 (A) CVs of 0.2 mM HAC obtained at 100 mVs-1 in different pH
media: (1) pH 2.0 (2) pH 3.0 (3) pH 4.0 (4) pH 5.0 (5) pH 6.0 (6) pH
7.0 (7) pH 8.0 (8) pH 10.0 (9) pH 11.0 (10) pH 12. 8 (B) Plot of Ep vs
pH of both anodic and cathodic peaks…………………………. 103
Fig. 4.43 Cyclic voltammograms of (1) the solvent and (2) 1 mM DHDN
showing anodic peaks (a1 and a2) and cathodic peaks (c1 and c2) in a
medium of pH 7.0 at 100 mVs-1…………………………. 104
Fig. 4.44 (A) DPVs of 0.2 mM HAC showing reduction in various pH media
(B) Plot of Epc vs pH of HAC…………………………. 106
Fig. 4.45 DPVs representing oxidation of 0.2 mM HAC in different pH media 107
Fig. 4.46 DPVs of 0.5 mM DHDN obtained in solution of pH 7.0 at a scan rate
of 5 mV s-1 showing (A) reduction and (B) oxidation labelled by
prominent anodic peaks a1 and a2…………………………. 107
Fig. 4.47 (A) SWVs representing reduction of 0.2 mM HAC in different pH
media (B) SWV of 0.2 mM HAQ in pH 7.0 showing forward and
backward components of the total current…………………………… 108
Fig. 4.48 SWVs obtained at100 mVs-1 showing the nature of cathodic (A) and
anodic (B) peaks of 0.5 mM DHDN in pH 7.0………………………. 110
Fig. 4.49 1st scan CVs of 1 mM HCAQ obtained in pH 3.0 and 7.0 at 100 mVs-
1 ………………………….……………………………………... 114
Fig. 4.50 DPVs of 0.4 and 0.6 mM HCAQ recorded in pH 7.0 at 5 mVs-1 after
cathodic scan without cleaning the electrode surface………………... 114
Fig. 4.51 Effect of (A) scan rate at 323±1K and (B) temperature at 80 m Vs-1
on the cyclic voltammetric behaviour of 0.7 mM HCAQ…………… 115
Fig. 4.52 Plots of peak current of (A) 1a and (B) 2c against the square root of
scan rate using 0.7 mM solution of HCAQ…………………………... 116
Fig. 4.53 Plot of log ks as a function of 1/T…………………………. 118
Fig. 4.54 (A) DPVs of 0.7 mM HCAQ recorded at 5 mVs-1 in 50% ethanol and
50% BR buffer of pH 3-11 (B) Plot of Ep versus pH using DPV data
of peak 1a………………………….……………………………. 120
Fig. 4.55 (A) 1st and (B) 2nd scan DPVs showing reduction of 0.7 mM HCAQ
and oxidation of the reduction product in 50% ethanol and 50% BR
buffer of pH 3-11 at 5 mVs-1………………………………………….
120
List of Figures
xiii
Fig. 4.56 Ep as a function of pH using DPV data of peak 1c……..…………… 121
Fig. 4.57 (A) SWV of different concentrations of HCAQ obtained at 100 mV
s-1 in a medium of pH 10 (B) Plot of peak current versus
concentration……………………………….…………………………. 122
Fig. 4.58 (A) UV-Vis spectra of different concentration of HCAQ obtained in
pH 7 (B) Absorption spectra of 16 µM HCAQ recorded in pH 3-11 122
Fig. 4.59 Color variation of 16 µM HCAQ in pH 3-12……………………… 124
Fig. 4.60 (A) Plots of absorbance versus concentration using UV-Vis
spectroscopic data obtained in pH 7 (B) Absorbance of 16 µM HCAQ
as a function of pH………………………….……………… 124
Fig. 4.61 (A) Consecutive CVs of 0.6 mM HACAD obtained in pH 7.0 at 100
mVs-1 (B) CVs of 0.6 mM HAD and (C) HACAD at different scan
rates………………………….……………………………………… 128
Fig. 4.62 Plots of peak current corresponding to peak 1a against square root of
scan rate using 0.6 mM solution of (A) HACAD, (B) HAD and (C)
HOAD………………………….………………………… 129
Fig. 4.63 Plot of logko as a function of 1/T…………………………. 131
Fig. 4.64 (A) DPV’s of HACAD obtained in 50% ethanol and 50% BR buffer
of pH 3-11 at 5 mVs-1 (B) Plots of peak potential as function of pH.. 132
Fig. 4.65 DPV’s of (A) HAD and (B) HOAD in media of different pH at 5
mVs-1 (C) plot of peak potentials of HOAD against pH………..……. 134
Fig. 4.66 SWVs of different concentrations of (A) HACAD and (B) HAD
obtained at 100 mVs-1 in a medium of pH 10.0. Insets show plots of
peak current versus concentration…………………...……………….. 135
Fig. 4.67 Square wave voltammograms of different concentrations of HOAD at
100 mVs-1 in a medium buffered at pH 10.0. Inset shows plot of peak
current versus concentration……………………………………. 136
Fig. 4.68 SWVs of 0.6 mM HACAD showing forward, backward and total
current………………………….…………………………...…………. 137
Fig. 4.69 UV-Vis spectra of 0.02 mM solutions of the selected
anthracenediones in pH 7.0………………………….……...………… 137
Fig. 4.70 Absorption spectra of 0.02 mM solutions of (A) HACAD and (B)
HAD in media of pH 3-12………………………….………………… 139
List of Figures
xiv
Fig. 4.71 Absorbance of 20 µM (A) HAD and (B) HOAD as a function of pH 139
Fig. 4.72 (A) CVs at 100 mVs-1 (B) DPVs showing reduction and (C) oxidation
of 1 mM solution of DMN at 5 mVs-1………………………………… 146
Fig. 4.73 (A,B) CVs of 1 mM solution of NDA obtained at 100 mVs-1 (C)
DPVs at 5 mVs-1 in different pH media and (D) plots of peak
potentials vs. pH……..…………………….…………………………. 147
Fig. 4.74 Electronic absorption spectra of 10 μM NDA and 22 μM DMN…….. 150
Fig. 4.75 Electronic absorption spectra of 15 µM NDA (A) and 30 µM DMN
(B) recorded in different pH media using 50% ethanol and 50% BR
buffer………………………….…………………………….………… 151
Fig. 4.76 Glucose (A), cholesterol (B) and triglycerides (C) concentrations
following treatment with naphthalene derivates NDA and DMN;
control groups were treated with glibenclamide, an antidiabetic drug,
whereas negative control was treated with Alloxan Monohydrate…... 152
Fig. 4.77 Percent viability of HeLa cells treated with different concentrations
(10, 50, 100 and 150 µM) of NDA and DMN…………………...…… 154
Fig. 4.78 MDA (µM, mean ± SD) levels relative to the NTC sample in (A)
NDA (B) DMN treated HeLa cell line……………………………..…. 155
List of Tables
xv
Table 3.1 Thermo gravimetric data of complexes (1) and (4) ……………… 37
Table 3.2 Names and Structures of selected quinones ………………….. 38
Table 4.1 Crystal data and details of the structure refinement for BPIMN
and CPIMN ……………… ……………… …………………. 51
Table 4.2 Selected experimental and calculated geometric parameters for
BPIMN and CPIMN ……………… ……………… ………… 52
Table 4.3 Hydrogen-bonding geometry (Å, ◦) for crystal BPIMN……….. 53
Table 4.4 Selected parameters of BPIMN and CPIMN calculated by
RB3LYP/3-21G set using DFT method in GUSSIAN03W
package ……………… ……………… ……………… ……………… ……………… 61
Table 4.5 Crystal data and details of the structure refinement for HL…….. 64
Table 4.6 Selected experimental and calculated geometric parameters for
HL ……………… ……………… ……………… ……………… ……………… ……… 64
Table 4.7 Hydrogen-bonding geometry (Å, ◦) for crystal DMN ……………… 65
Table 4.8 Diffusion coefficient, rate constant, LOD and LOQ values for
HL, (1), (2), (3) and (4) ……………… ……………… ……………… ………… 70
Table 4.9 IC50 value (µM) of HL, (1), (2), (3) and (4) complexes for HeLa
cell line ……………… ……………… ……………… ……………… ……………… 82
Table 4.10 Comparison of antibacterial activities of HL and complexes…… 84
Table 4.11 Comparison of antifungal activities of HL and complexes……… 85
Table 4.12 Selected parameters of BPIMN and CPIMN calculated by
RB3LYP/3-21G set using DFT method in GUSSIAN03W
package. ……………… ……………… ……………… ……………… ……………… 86
Table 4.13 Observed and calculated hydrogen chemical shifts (δ, in ppm)…. 87
Table 4.14 EHOMO and ELUMO values of compounds ……………… …………………… 100
Table 4.14 Parameters obtained from SWV ……………… …………………………… 109
Table 4.15 Diffusion coefficient and heterogeneous rate constant values of
HCAQ evaluated on the basis of oxidation peak 1a at different
temperatures ………………………………………………………………………… 116
Table 4.16 Heterogeneous electron transfer rate constant values of HCAQ
evaluated on the basis of quasi-reversible redox peaks at different
temperatures ………………………………………………………………………………
117
List of Tables
xvi
Table 4.17 Thermodynamic parameters of HCAQ ……………… ……………… 118
Table 4.18 Diffusion coefficient and heterogeneous rate constant values of
the selected compounds at different temperatures ……………… 130
Table 4.19 Thermodynamic parameters of the redox processes of the selected
compounds ……………… ……………… ……………… ……………… ……………… 132
Table 4.20 Values of half peak width, slope of Ep – pH plots, number of
electrons and protons and LOD and LOQ ……………… ……… 135
Table 4.21 Molar extinction coefficient and pKa values of the analytes……. 140
Table 4.22 Percent viability of HeLa cells following treatment with NDA
and DMN ……………… ……………… ……………… … 154
List of Schemes
xvii
Scheme 3.1 Synthesis of 1-((4-bromophenylimino)methyl)naphthalen-2-ol
and 1-((4-chlorophenylimino)methyl)naphthalen-2-ol………. 33
Scheme 3.2 Synthesis of Schiff base HL and its numbering scheme……. 34
Scheme 3.3 Synthesis of metal complexes of HL …………………………………….. 36
Scheme 3.4 Hydrolysis of alkaline phosphatase………………………………………… 44
Scheme 4.1 Intramolecular hydrogen bonding in BPIMN and CPIMN….. 54
Scheme 4.2 Tautomeric forms of BPIMN………………………………………………….. 55
Scheme 4.3 Proposed mechanism for oxidation of CPIMN for peak 1a….. 60
Scheme 4.4 Proposed mechanism for oxidation of CPIMN for peak 2a….. 61
Scheme 4.5 Proposed oxidation mechanism of HL for peak 1a………….. 75
Scheme 4.6 Proposed oxidation mechanism of HL for peak 2a………….. 76
Scheme 4.7 Proposed reduction mechanism for peak 1c and 2c…………. 77
Scheme 4.8 Proposed redox mechanism of HMNA in slightly acidic, neutral
and alkaline conditions…………………………………………………………… 95
Scheme 4.9 Proposed redox mechanism of HMNA and its hydrolyzed
product under strongly acidic conditions………………………………. 97
Scheme 4.10 Proposed reduction mechanism of HAC……………………………… 110
Scheme 4.11 Suggested oxidation mechanism of HAC…………………………….. 111
Scheme 4.12 Proposed reduction mechanism of DHDN…………………. 112
Scheme 4.13 Suggested oxidation mechanism of DHDN……………………… 112
Scheme 4.14 Proposed oxidation mechanism of HCAQ corresponding to
peak 1a…………………………………………………………………………………… 125
Scheme 4.15 Proposed oxidation mechanism corresponding to peak 2a
showing proton-coupled electron loss in acidic conditions…. 125
Scheme 4.16 Suggested CE type oxidation mechanism of HCAQ
corresponding to peak 2a in neutral and alkaline conditions… 126
Scheme 4.17 Proposed oxidation mechanism of HCAQ corresponding to
peak 2a’ in neutral and alkaline conditions……………………… 126
Scheme 4.18 Proposed redox mechanism of HCAQ corresponding to peaks
appearing in the negative potential window of GCE……….. 126
Scheme 4.19 Chemical structure of HOAD showing hydrogen bonding….. 134
Scheme 4.20 Proposed redox mechanism of HACAD in media having pH
≤ 6.0………………………………………………… 141
List of Schemes
xviii
Scheme 4.21 Proposed redox mechanism of HACAD in media having pH ≥
7.0…………………………………………………………………………………………… 142
Scheme 4.22 Proposed redox mechanism of HAD in media having pH ≤ 4.0 142
Scheme 4.23 Suggested redox mechanism of HAD in media having pH ≥
5.0………………………………………………………… 143
Scheme 4.24 Proposed redox mechanism of HOAD in pH 3-12………………… 143
Scheme 4.25 Proposed redox mechanism corresponding to peak 1c and 2c of
all the selected compounds 144
Scheme 4.26 Proposed oxidation and reduction mechanism of
DMN………………………………………………………………………………………… 148
Scheme 4.27 Proposed reduction and oxidation mechanism of NDA…….. 149
Chapter 1 Introduction
Page 1
Electrochemistry offers useful insights about the chemistry of redox molecules
especially small organic and biomolecules which have the ability to closely approach
to the electrode surface and thus directly transfer electron to the electrode surface. On
the other hand, large biomolecules like proteins and enzymes are usually stagnant to
transfer electrons at ordinary electrodes [1, 2]. The voltammetric behavior of
molecules depends on the pH of the medium and their potential-pH diagrams can give
useful informations about the stability of different forms of molecules [3]. Moreover,
the pH dependent redox behavior have a significant role in controlling the biocidal
applications of different drugs. Based on these considerations, the redox mechanistic
pathways of robustly electroactive compounds belonging to Schiff bases, quinones
and naphthalene derivatives were studied in media of different pH.
Litreature survey reveals that many classes of organic and inorganic
compounds are used as redox probes and electron transfer mediators, however, low
efficiency and high cost limit their use in bulk processes. Therefore, scientists are
struggling to search efficient and cost effective redox probes for use in many
industrial processes. In this context Frank et al. have provided a list of quinones and
phenazines bestowed with the potential of acting as electron shuttle [4]. Chen et al.
have documented the effects on the electrochemical activities of microorganisms
treated against phenazines and quinones [5]. Riboflavin, 2- and 4-aminophenol, humic
acid and different substituted benzenes and anthraquinones have been reported to act
as efficient redox mediators [6]. The present work is an effort to search for efficient
electron transfer mediators, redox probes and to explore the electrochemical
properties and pH dependent redox behavior of biologically important compounds
capable of registering exquisite voltammetric signatures at the surface of glassy
carbon electrode. Aromatic hydroxy derivatives from different classses have been
selected for the investigation of their electron transfer reactions in a wide pH range.
Literature survey reveals that the researchers have mostly used a single
electrochemical technique for the investigation of the redox properties of compounds.
However, for getting reliable information about the redox properties, it obligatory to
use a variety of electrochemical techniques as cyclic voltammetry provides a general
portrait of the redox response, square wave voltammetry ensures the
reversibility/quasireversibility and differential pulse voltammetry helps in the
determination of number of electrons involved in oxidation or reduction processes.
Most of the electrochemists working on evaluation of redox properties investigate
Chapter 1 Introduction
Page 2
their analytes in a single medium, while, for unfolding the proton coupled electron
transfer reactions, it is important carry out electrochemical experiments im media of
different pH. Keeping these ideas in mind we utilized several sensitive
electrochemical techniques and carried out experiments in media of a wide pH range
with the objective of unfolding the mechanisms of proton coupled electron transfer
reactions. Moreover, we used computational studies for the identification of the most
oxidizable or reducible moieties of the compounds. The redox mechanisms proposed
in this work are expected to provide useful insights about the biological and medicinal
roles of the selected compounds. A brief introduction about these biologically
important electroactive compounds is given in the subsequent sections.
1.1 Schiff Bases Schiff bases are the condensation products of the reaction between primary amines
and carbonyl compounds. They are characterized by the presence of azomethine or
imine (–C=N) group in their structures. They were named after their discoverer, a
well known chemist Hugo Schiff. Schiff bases attached to alkyl systems are relatively
unstable and readily polymerize but those attached with aromatic systems are quite
stable owing to conjugation [7]. Schiff bases are bestowed with wide range of biocidal
and technological applications due to their photometric, electrochemical and
thermochromics properties [8, 9]. They can act as excellent ligands in the field of
coordination chemistry due to the available binding sites in their structures [10]. Thus
they play an important role in the formation of stable complexes with metal ions.
These types of compounds are also capable of exhibiting excellent activity in
homogenous [11] and heterogeneous catalysis [12]. They are capable of exhibiting
catalytic activity even in the presence of moisture and at temperatures higher
than100oC. Schiff bases and their metal complexes have been reported to catalyse a
large number of biological reactions. It is found that Schiff bases are involved as
catalyst in hydrogenation of olefins. They also play important role in enzyme
preparation [13]. They are employed as charge carriers in potentiometric sensors
because of their excellent stability, sensitivity and selectivity [14]. Moreover they can
effectively inhibit corrosion [15] by a mode involving the formation of a monolayer
on the surface to be protected.
Chapter 1 Introduction
Page 3
Schiff bases contains imine group which is associated with wide range of
biological applications [16]. Imine complexes are found to exhibit anti-viral [17],
anti-malarial, anti-tumour [18], anti-bacterial effects. They are widely employed in
the treatment of diabetes and AIDS. They can work as model for understanding the
activity and action of various biologically important compounds. Isatin Schiff base
ligands are known to exhibit anti-viral activity and are found quite effective for HIV
treatment [18]. Other Schiff bases derivatives like salicyaldehyde Schiff base
derivatives may find use in vaccination of various viral infections including small
pox, poliomyelitis (polio) etc [19]. They are used in anti-epiletic drugs due to their
capability of exhibiting anticonvulsant activity.
Photochromic [20] and thermochromic properties of Schiff bases have made
them useful materials in modern technological devices. They are in routine usage of
optical computers and molecular memory storage devices and imaging systems where
they are employed to measure as well as to control the intensity of radiations [21].
Their photometric properties make it possible to use them as photodetectors in
biological systems as well as photo stabilizers [22]. They are also used as dyes in
solar collectors. They are likewise applied in optical sound recording technology, in
molecular memory storage and in reversible optical memories. Other properties of
Schiff bases which make them suitable material in modern appliances are their
chelating ability, optical non-linearity, thermal stability, liquid crystal property and
possessing a structure having electrical properties to transfer protons.
The antibacterial, anticancer, antitumor, anti-inflammatory agents, anti-
tuberculosis, insecticidal, antimicrobial, anticonvulsant [23] activities of Schiff bases
have been extensively reported as they are applied against a number of organisms
including Plasmopora viticola, Candida Albicans, Bacillus polymxa, Mycobacteria,
Trychophyton gypseum, Erysiphe graminis and Escherichia coli Staphylococcus
aureus. Aromatic schiff bases find extensive applications in medicinal and inorganic
chemistry. Thus, spurred on by the broad range applications of Schiff bases, the
synthesis of new electroactive chemotherapeutic candidates of this class was one of
the prime objective of our work and that’s why we synthesized and characterized
different Schiff bases and some of their metal complexes.
They are used for enhancing the sensitivity and selectivity of electrochemical
and optical sensors [24]. Their metal complexes serve as models for anticancer and
herbicidal applications [25]. Isatin Schiff bases possess antiviral, antiprotozoal, anti-
Chapter 1 Introduction
Page 4
HIV, anticonvulsant and anthelmintic activities. Schiff bases of 4-dimethylamine
benzaldehyde show excellent antibacterial activity and N-(Salicylidene)-2-
hydroxyaniline is used for the treatment of mycobacterium tuberculosis [26]. 2,4-
dichloro-5-fluorophenyl group containing Schiff base is an effective inhibitor of
bacterial growth [27]. Schiff base obtained from p-toluidene and furylglyoxal show
excellent inhibitory action against Bacillus subtilis, Escherichia coli, Staphylococcus
aureus and Proteus vulgaris [28]. Schiff bases containing β-keto esters and o-
aminobenzoic acid are used for the treatment of S. epidermidis, E. coli, B. cinerea and
A. niger [29]. Silva et al. documented the antiviral activity of Schiff base produced
form salicaldehyde and 1-amino-3-hydroxyguanidine tosylate [26]. Ancistrocladidine,
a metabolitic product of ancistrocladaceae and dioncophyllaceae plants has been
reported to show antimalarial properties [26]. In medicines they are used as antibodies
and anti-inflammatory agents [30]. New chemotherapeutic Schiff bases are now
attracting the attention of biochemists [31]. Some Schiff bases have been documented
to protect metals from aggressive medium by adsorption at metals surfaces [32].
Schiff bases find use as catalyst for oxygenation, hydrolysis, electro-reduction and
decomposition of various compounds. Electrochemistry provides highly useful details
of catalysis as it involves the change in oxidation state and structure of complex.
Therefore, their electronic and steric effects offers basic key role in the design of new
catalysts in numerous fields [33].
Azomethine dyes comprising the largest percentage in organic dyes are well
known in cosmetic and food industries as coloring agents [34]. Beside their
applications in emerging optical and analytical fields [3] their pharmacological
applications are also reported due to their structural similarity with natural biological
species [35]. They are widely distributed in organic substances and azomethine
linkage (C=N) present in Schiff bases has been shown as responsible for most of their
biological and industrial applications [36]. Therefore, our electrochemical
investigations of the pH dependent redox behavior of azomethine are expected to
offer valuable insights about getting a control over the modulation of biological and
industrial applications of Schiff bases.
Schiff bases are reported to be good chelating agents [37], especially when
–OH or –SH functional group is present because of their ability of existing in a five or
six membered ring including metal ion. Hydroxyl group containing Schiff bases have
received the utmost attention of physicists and chemists due to their specfic
Chapter 1 Introduction
Page 5
thermochromic/photochromic characteristics and electrochemical applications [38].
These pH dependent proton exchanging compounds are also applied in the designing
of different molecular electronic devices [39]. Moreover, hydroxyl-substituted Schiff
bases possess antioxidant activities [40]. Metal complexes of these bases are also
fascinating due to their stability in different oxidation states [41]. Based on these
considerations, we synthesized Schiff bases and their metal complexes and explored
their electrochemical and photochromic properties.
Induction of desired biological action of a drug needs its structural evaluation.
Therefore, extensive research for the determination of structure and consequences of
structural modification are the basic requirements in this regard. Electrochemical data
provides valuable information and allows the prediction of redox mechanism. This
knowledge helps in finding suitable donor–acceptor combination. Although it may not
be free of uncertainty but helps to achieve the targets. Steric and inductive effects of
the substituent groups strongly influence the probability of electronic colud on metal
ion and basicity of nitrogen atom hence, change the redox properties of whole Schiff
base molecule. Based on these considerations we synthesized different Schiff bases
and characterized them by a variety of techniques for getting useful insights about
their biological action mechanism and role in cellular milieu.
1.2 Quinones
Benzoquinones and their chloro and methyl derivatives are reported to exhibit good
redox properties and are used to decrease the band gap of photovoltaic cells [42].
Naphthoquinones and their reduced forms are well known parts of biological electron
transfer chains present in the membranes of bacteria, mitochondria and chloroplast
[43]. Some naphthoquinone derivatives like menadiols are used as vitamins
depending upon their molecular structure [43]. Two important derivatives of
naphthoquinones, menadione and β lapachon have achieved clinical status as anti-
tumor drugs. These type of compounds are gaining mounting attention of chemists,
biologists and pharmacologists because many anticancerous drugs contain quinone
functionality [44]. However, it is difficult to generalize the overall biological action
mechanism of these large number of quinones with diverse structures [45]. In the
context of biological action mechanism, the ability of accepting electrons to form
radical anion or dianion makes these naphthoquinoes quite significant. Therefore,
Chapter 1 Introduction
Page 6
their redox properties are governed by the electron donating or withdrawing
substituents attached to their basic framework [46]. Vitamin K3 is a naturally
occurring naphthoquinone associated with the production of prothrombin, a blood
clotting protein. Vitamin K includes K1 and K2 which are formed from provitamin K3.
These vitamins play an integral part in bone calcification. Deficiency of these can lead
to serious health problems like excessive bleeding and hemorrhage. Plumbagin and
other structural analogs of vitamin K have been reported to have anticancerous
properties [47]. Thus, the detailed electrochemical studies of quinones seems
conceivable for for getting information about their role in different fields.
Anthracenediones (ACDs) are bestowed with broad range applications.
Depending upon the position of keto groups, ACDs have various isomers but the most
studied and common representative of this class is 9,10-anthracenedione. Natural
ACDs are diverse and their diversity can be related to the existence of different
substituents, such as -OH, -CH3, -NH2, -OCH3, -CH2OH, -CHO, -COOH, etc., as well
as the reduction of carbonyl groups to anthracenes and double bonds in benzene ring
to hydro-anthracenediones. AQs found in some plants and animals are used as
constituents of many dyes [48]. A number of their derivatives play a key role in
catalyzing reduction of biologically important molecules like vitamin K [49]. Some
AQs find use as algaecides and fungicides. Their derivatives, being electron
acceptors, play a major role in photosynthesis and are able to transport a negative
charge in biological systems [50]. Due to the adsorption ability of AQs at a glassy
carbon electrode, their derivatives are used for electrode modification [51]. The
modified electrode can be used as redox catalyst for the reduction of oxygen to
hydrogen peroxide [52]. On complexation with macrocyclic polyethers, AQs form
lumophores that can be used to selectively detect many transition metal ions in
solution, thus offering an alternative photophysical detection mechanism [53]. Some
AQs like rufigallol and aclarubicin are used as antimalarial [54] and anticancer drugs
[55]. A plethora of anthracenediones have been studied for their anticancer activities.
The results show that some members of this class have strong potency against
leukemia and tumor [56]. Glycosidic derivatives of anthracenediones are used as
laxatives and for the treatment of skin diseases caused [57]. Some anthracenediones
are used for the treatment of malaria because of their effective role against other
quinine resistant plasmodium species in human body [58]. An thracenediones such as
9-methoxydihydrodeoxybostrycin and 10-deoxybostrycin have been reported to
Chapter 1 Introduction
Page 7
exhibit strong potency against bacillus cereus [59]. Asongalem et al. have
documented some anthracenediones as analagesic and hypotensive [60]. AQs, being
redox catalysts, play a vital role in the production of paper pulp [12]. Alizarin, an AQ
derivative, is extensively used in the laboratory as pH indicator, an indicator of bone
growth and a spot test reagent [61] due to its color variation in media of different pH.
Apart from pH-dependent photophysical properties, the voltammetric behavior of
AQs also depends on pH and their potential-pH diagrams can give useful information
about the stability of different forms of AQs in acidic, basic and neutral conditions
[62]. AQs can be used for electrochemical switchable devices and pH-responsive drug
delivery systems on the basis of the strong pH dependency of their redox behavior
[63]. AQs are commonly used as electron transfer mediators and redox probes [64].
Feng and coworkers demonstrated the enhancement in power output of a microbial
fuel cell by the mediator role of an AQ derivative for efficiently transferring electrons
from bacteria to anode [65]. The results of Kumamoto et al. revealed that an
oligonucleotide modified with AQ can act as a suitable redox reporter for the
detection of a single-base mismatch in DNA [66]. Abi and Ferapontova also used an
AQ as redox probe for the investigation of the electron transfer mechanism within
loosely compact monolayers of DNA [67].
A number of anthracenediones are used as natural pigments [68]. Chaoyan Li
et al. have introduced the use of anthracenedione based dyes as photo sensitizer in
photosensitized dye solar cell with some excellent results [69]. Solvent violet 13, a
synthetic anthracenedione dye is used to dye hydrocarbon products such as
thermoplastics, polystyrenes and synthetic fiber [70]. It is also used in hair and skin
care products [71]. Literature survey reveals that some anthracenediones are used for
the production of gas in satellite balloons. Moreover, the industrial applications of this
class are further enhanced by its use in paper making industries as digestive additives
[72]. Another important industrial application of anthracenediones is their role in the
production of H2O2. Alkylated derivatives including 2-ethyl-9, 10- anthracenedione
are used for this purpose rather than simple anthracenediones [73].
ACDs are used as useful nucleotide specific ligands for the purification of
proteins [41]. The enzyme encoded by UGT1A8 gene has glucuronidase activity with
many substrates including anthracenediones [74]. Anthracenedione sulfonate is used
to catalyze cyclic photophosphorylation reactions [75]. These are employed for
labeling purposes in biological systems during analytical studies. Modification of
Chapter 1 Introduction
Page 8
deoxynucleosides or deoxynucleoside triphosphates by linking the anthracenedione
molecule at 5-position in pyrimidine or 7-position in 7-deazaadenine yields a suitable
substrate for polymerase and incorporation of a label in DNA [76]. ACDs are also
used for labeling peptides via attachment to their N terminus [77]. These compounds
being the inhibitors of topoisomerase find use as pharmacological target for cancer
prevention [78]. Anthracene-9,10-dione, a naturally occurring irritant produces
different taste during illness thus, widely used as bird repellent [79]. Due to the ability
of ACDs as carrier of electron and hydrogen, the electrochemical behavior of this
class of compounds has been the main focus of researchers. Allietta et al reported
electrochemical-polymerization of pyrrole-substituted anthracenedione at indium-tin
oxide electrode [80]. These are also potentially useful semiconductors due to their
intriguing properties for low cost organic electronics [81]. Many researchers have
reported that the biological activities of ACDs are strongly linked with their redox
behavior [9]. To get insights about their biological action, assessment of reaction
mechanistic pathways and evaluation of physicochemical parameters, several
researchers have studied the reduction of such molecules under different conditions.
The quinones and anthracenediones exibit two one-electron reduction steps
successivly producing semiquinone and dianion and generat two separate cathodic
waves in which the first step follows completely reversible redox behavior and the
second step either reversible or quasi-reversible depending upon the experimental
conditions [82]. Anthracenediones containing hydroxyl groups are interesting from
electrochemical point of view because the position of these electropores has been
reported to modify the typical redox behavior of the quinonoid group [83] due to
intramolecular hydrogen bonds formation. In addition, these moieties are related to
the biological activity of ACDs [84].
The current work is expected to assist in explaining some of the redox
properties of quinones which are still unexplored by the scientific community.
Although anthracenediones have a broad spectrum of applications and their reduction
is widely studied [82, 83], yet information about their electro-oxidation is meager and
the pH dependent oxidation of quinones at different temperatures is a largely
unexplored area. So to bridge this gap, the pH-dependent electro-oxidation of
different anthraquinone derivatives was investigated at different temperatures with the
objectives of adding new candidates to the literature of the quinone family and
providing useful insights into the pH-dependent electrochemical fate of quinones.
Chapter 1 Introduction
Page 9
Most quinones exhibit redox behavior in the negative potential domain of glassy
carbon or gold electrodes, but interestingly our quinone derivatives, can also show
redox response in the positive potential windows of the electrode. Hence, in
comparison to simple quinones, these substituted quinones has extended applications
as redox probe and electron transfer mediator. The biological activities of quinones
are related to their redox reactions [84]. The biological activity of hydroxyquinones is
considered to correspond to their hydroxyl group [85]. Therefore, these selected
compounds are also screened for several biological activities. By modifying the
position of the hydroxyl group, the redox properties of the quinonoid moiety can be
altered [86]. Thus, a greater alteration in the properties of these compounds may be
due to the possibility of changing the position of several hydroxyl groups in the
structures of our studies anthraquinones.
1.3 Naphthalene Derivatives Naphthalene derivatives are naturally occurring compounds present in droseraceae,
nepenthaceae, polygonaceae and many other carnivorous plants [87]. They are
endowed with broad range of biological activities [88]. Various studies have reported
naphthalene derivatives as potent anti-diabetic agents [89]. It is well established that
excess glucose in the plasma due to pancreatic damage or non-production of insulin or
insulin resistance leads to the production of free radicals or reactive oxygen species
(ROS). The production of ROS is highly damaging to tissues and cells. Accordingly
many natural or synthetic anti-diabetic agents possess the properties of upregulating
and inducing antioxidant enzymes that readily scavenge ROS [90]. With this in mind
we have screened two unexplored derivatives of naphthalene, naphthalene-2,3-
dicarboxylic acid (NDA) and 1,8-dimethoxynaphthalene (DMN) (Schemes 1 and 2)
for their anti-diabetic activities. These experiments were conducted in vivo by
inducing diabetes in Sprague–Dawley rats followed by treatment with the test
compounds for measurement of the end points like plasma glucose, total serum
cholesterol and triglycerides.
The high concentration of free radicals is the main cause of oxidative stress.
They accelerate the mutation rate which increases the risk of cancer and
neurodegenerative diseases [91]. They also decrease the immunity which in turn leads
to an increase in susceptibility to bacterial and viral infections [92]. In this scenario,
Chapter 1 Introduction
Page 10
antioxidants play an important role by scavenging and neutralizing free radicals by the
donation of electrons and protons, thus, maintaining sound health [93]. Antioxidants
therefore have gained enormous attention in the current research due to their ability to
combat different diseases [94]. Naphthalene, although by itself exhibits toxic effects
yet some of its derivatives are bestowed with medicinal value [95], and several of
them have been characterized as antioxidants. Considering the cytotoxic and
antioxidant properties of naphthalene derivatives, the pharmaceutical and medicinal
chemists are trying to produce less toxic drugs which can exhibit maximum
therapeutic index and could therefore be used as therapeutics. Based on these
considerations, we screened the selected naphthalenes for their antioxidant activity.
Moreover, these were also analyzed for their cytotoxic effects against cervix
adenocarcinoma cells (HeLa).
The primary metabolites of aromatic compounds like naphthalene derivatives
have been reported to act as potent antioxidants [95]. Some of their metabolites work
in the Mitchellian redox loops of chloroplast and mitochondrial electron transport
chains. Their secondary metabolites are structurally various compounds that play
opposing roles in many organisms. The broad range biological activities and
inhibitory effects of such compounds are related to their redox behavior [96].
Moreover, the physicochemical properties of naphthalenes are greatly dependent on
the nature, number and position of different substituents. The biological activities of
naphthalenes depend strongly on the electronic properties of the substituents attached
at the aromatic ring [97]. Thus, spurred on by the wide spectrum of biological
activities of naphthalenes and their relationship with redox behavior and electronic
properties of the groups attached at the aromatic rings, we investigated the pH
dependent electrochemical and electronic absorption spectroscopic behavior of the
two naphthalene derivatives.
Chapter 1 Introduction
Page 11
References [1] Itoh S, Biomimetic electron transfer reactions (1) 1. Redox functions of novel o-
quinone coenzymes, Electrochemistry 2000; 68: 807-811.
[2] Tatsumi H, Nakase H, Kano K, Ikeda T. Mechanistic study of the autoxidation of
reduced flavin and quinone compounds, Journal of Electroanalytical Chemistry 1998;
443: 236-242.
[3] Susan MABH, Begum M, Takeoka Y, Watanabe M, Effect of pH and the extent of
micellization on the redox behavior of non-ionic surfactants containing an
anthraquinone group, Journal of Electroanalytical Chemistry 2000; 481: 192-199.
[4] Van der Zee FP, Cervantes FJ. Impact and application of electron shuttles on the
redox (bio) transformation of contaminants: a review. Biotechnology Advances 2009;
27: 256-277.
[5] Chen B-Y, Hsueh C-C, Liu S-Q, Hung JY, Qiao Y, Yueh P-L, Unveiling
characteristics of dye-bearing microbial fuel cells for energy and materials recycling:
redox mediators, International Journal of Hydrogen Energy 2013; 38: 15598-15605.
[6] Sun J, Li W, Li Y, Hu Y, Zhang Y, Redox mediator enhanced simultaneous
decolorization of azo dye and bioelectricity generation in air-cathode microbial fuel
cell, Bioresource Technology 2013; 142: 407-414.
[7] Pandeya S, Sriram D, Nath G, DeClercq E, Synthesis, antibacterial, antifungal and
anti-HIV activities of Schiff and Mannich bases derived from isatin derivatives and
N-[4-(4′-chlorophenyl) thiazol-2-yl] thiosemicarbazide, European Journal of
Pharmaceutical Sciences 1999; 9: 25-31.
[8] Brodowska K, Lodyga‐Chruscinska E, Schiff Bases—interesting range of applications
in various fields of science, ChemInform 2015; 46-50.
[9] Hadjoudis E, Vittorakis M, Moustakali-Mavridis I, Photochromism and
thermochromism of schiff bases in the solid state and in rigid glasses, Tetrahedron
1987; 43: 1345-1360.
[10] Dharmaraj N, Viswanathamurthi P, Natarajan K, Ruthenium (II) complexes
containing bidentate Schiff bases and their antifungal activity, Transition Metal
Chemistry 2001; 26: 105-109.
[11] De Clercq B, Verpoort F, Assessing the scope of the introduction of Schiff bases as
co‐ligands for monometallic and homobimetallic ruthenium ring‐opening metathesis
Chapter 1 Introduction
Page 12
polymerisation and ring‐closing metathesis initiators, Advanced Synthesis &
Catalysis 2002; 344: 639-648.
[12] Cozzi PG, metal–salen schiff base complexes in catalysis: practical aspects, Chemical
Society Reviews 2004; 33: 410-421.
[13] Kumar S, Dhar DN, Saxena P, Applications of metal complexes of schiff bases—a
review, Journal of Scientific and Industrial Research 2009; 68: 181-187.
[14] Faridbod F, Ganjali MR, Dinarvand R, Norouzi P, Riahi S, Schiff's bases and crown
ethers as supramolecular sensing materials in the construction of potentiometric
membrane sensors, Sensors 2008; 8: 1645-1703.
[15] Gomma GK, Wahdan MH, Schiff bases as corrosion inhibitors for aluminium in
hydrochloric acid solution, Materials Chemistry and Physics 1995; 39: 209-213.
[16] Aptula AO, Roberts DW, Mechanistic applicability domains for nonanimal-based
prediction of toxicological end points: General principles and application to reactive
toxicity, Chemical Research in Toxicology 2006; 19: 1097-1105.
[17] Garoufis A, Hadjikakou S, Hadjiliadis N, Palladium coordination compounds as anti-
viral, anti-fungal, anti-microbial and anti-tumor agents, Coordination Chemistry
Reviews 2009; 253: 1384-1397.
[18] Tiekink ER, Gold compounds in medicine: potential anti-tumour agents, Gold
Bulletin 2003; 36: 117-124.
[19] Jagt DL, Royer RE, Gossylic iminolactones and gossylic lactones and their anti-viral
activities, US Patents No. 5026726; 1991.
[20] Bamfield P, Hutchings MG, Chromic phenomena: technological applications of
colour chemistry, Royal Society of Chemistry; 2010.
[21] Yersin H, Highly efficient OLEDs with phosphorescent materials, John Wiley &
Sons; 2008.
[22] Banks RE, Smart BE, Tatlow J, Organofluorine chemistry: principles and commercial
applications, Springer Science & Business Media; 2013.
[23] Dettrakul S, Surerum S, Rajviroongit S, Kittakoop P, Biomimetic transformation and
biological activities of globiferin, a terpenoid benzoquinone from cordia globifera,
Journal of Natural Products 2009; 72: 861-865.
[24] Dave H, Ledwani L, A review on anthraquinones isolated from cassia species and
their applications, Indian Journal of Natural Products and Resources 2012; 3:
291-319.
Chapter 1 Introduction
Page 13
[25] Kim Y-M, Lee C-H, Kim H-G, Lee H-S, Anthraquinones isolated from cassia tora
(leguminosae) seed show an antifungal property against phytopathogenic fungi,
Journal of Agricultural and Food Chemistry 2004; 52: 6096-6100.
[26] Salimi A, Eshghi H, Sharghi H, Golabi SM, Shamsipur M, Electrocatalytic reduction
of dioxygen at the surface of glassy carbon electrodes modified by some
anthraquinone substituted podands, Electroanalysis 1999; 11: 114-119.
[27] Campos‐Martin JM, Blanco‐Brieva G, Fierro JL, Hydrogen peroxide synthesis: an
outlook beyond the anthraquinone process. Angewandte Chemie International Edition
2006; 45: 6962-6984.
[28] Santacesaria E, Di Serio M, Russo A, Leone U, Velotti R, Kinetic and catalytic
aspects in the hydrogen peroxide production via anthraquinone, Chemical
Engineering Science 1999; 54: 2799-2806.
[29] Nissim R, Batchelor-McAuley C, Li Q, Compton RG, The anthraquinone mediated
one-electron reduction of oxygen in acetonitrile, Journal of Electroanalytical
Chemistry 2012; 681: 44-48.
[30] Huang Q, Lu G, Shen HM, Chung M, Ong CN, Anti‐cancer properties of
anthraquinones from rhubarb, Medicinal Research Reviews 2007; 27: 609-630.
[31] Hou Y, Cao S, Brodie PJ, Callmander MW, Ratovoson F, Rakotobe EA,
Antiproliferative and antimalarial anthraquinones of Scutia myrtina from the
Madagascar forest, Bioorganic & Medicinal Chemistry 2009; 17: 2871-2876.
[32] Yang K-L, Wei M-Y, Shao C-L, Fu X-M, Guo Z-Y, Xu R-F, Antibacterial
anthraquinone derivatives from a sea anemone-derived fungus Nigrospora sp, Journal
of Natural Products 2012; 75: 935-941.
[33] Shadi IT, Chowdhry BZ, Snowden MJ, Withnall R, Semi‐quantitative analysis of
alizarin and purpurin by surface‐enhanced resonance Raman spectroscopy (SERRS)
using silver colloids, Journal of Raman Spectroscopy 2004; 35: 800-807.
[34] Hoyte D, Alizarin as an indicator of bone growth, Journal of Anatomy 1960; 94:
432-436.
[35] Takahashi S, Anzai J, Recent progress in ferrocene-modified thin films and
nanoparticles for biosensors, Materials 2013; 6: 5742-5762.
[36] Feng C, Ma L, Li F, Mai H, Lang X, Fan S, A polypyrrole/anthraquinone-2, 6-
disulphonic disodium salt (PPy/AQDS)-modified anode to improve performance of
microbial fuel cells, Biosensors and Bioelectronics 2010; 25: 1516-1520.
Chapter 1 Introduction
Page 14
[37] Abi A, Ferapontova EE, Unmediated by DNA electron transfer in redox-labeled DNA
duplexes end-tethered to gold electrodes, Journal of the American Chemical Society
2012; 134: 14499-14507.
[38] Li C, Yang X, Chen R, Pan J, Tian H, Zhu H, Anthraquinone dyes as photosensitizers
for dye-sensitized solar cells, Solar Energy Materials and Solar Cells 2007; 91: 1863-
1871.
[39] Gonzalez-García S, Moreira MT, Artal G, Maldonado L, Feijoo G, Environmental
impact assessment of non-wood based pulp production by soda-anthraquinone
pulping process, Journal of Cleaner Production 2010; 18: 137-145.
[40] Shang H, Zhou H, Zhu Z, Zhang W, Study on the new hydrogenation catalyst and
processes for hydrogen peroxide through anthraquinone route, Journal of Industrial
and Engineering Chemistry 2012; 18: 1851-1857.
[41] Shamsipur M, Siroueinejad A, Hemmateenejad B, Abbaspour A, Sharghi H, Alizadeh
K, Cyclic voltammetric, computational, and quantitative structure–electrochemistry
relationship studies of the reduction of several 9, 10-anthraquinone derivatives,
Journal of Electroanalytical Chemistry 2007; 600: 345-358.
[42] Gupta N, Linschitz H, Hydrogen-bonding and protonation effects in electrochemistry
of quinones in aprotic solvents, Journal of the American Chemical Society 1997; 119:
6384-6391.
[43] Capon A, Parsons R, The rate of a simple electron exchange reaction as a function of
the electrode material, Journal of Electroanalytical Chemistry and Interfacial
Electrochemistry 1973; 46: 215-222.
[44] Guin PS, Das S, Mandal P, Electrochemical reduction of sodium 1, 4-dihydroxy-9,
10-anthraquinone-2-sulphonate in aqueous and aqueous dimethyl formamide mixed
solvent: a cyclic voltammetric study, International Journal of Electrochemical Science
2008; 3: 1016-1028.
[45] Piljac I, Murray RW, Nonaqueous Electrochemistry of 1‐Hydroxy‐9, 10‐
Anthraquinone and Its Conjugate Base, Journal of The Electrochemical Society 1971;
118: 1758-1764.
[46] González FJ, Aceves JM, Miranda R, Gonzalez I, The electrochemical reduction of
perezone in the presence of benzoic acid in acetonitrile, Journal of Electroanalytical
Chemistry and Interfacial Electrochemistry 1991; 310: 293-303.
Chapter 1 Introduction
Page 15
[47] Johnson J, Gandhidasan I, Murugesan R, Cytotoxicity and superoxide anion
generation by some naturally occurring quinones, Free Radical Biology and Medicine
1999; 26: 1072-1078.
[48] Abreu FCd, Ferraz PAdL, Goulart MO, Some applications of electrochemistry in
biomedical chemistry. Emphasis on the correlation of electrochemical and bioactive
properties, Journal of the Brazilian Chemical Society 2002; 13: 19-35.
[49] Bartoszek A, Metabolic activation of adriamycin by NADPH-cytochrome P450
reductase; overview of its biological and biochemical effects, Acta Biochimica
Polonica 2002; 49: 323-332.
[50] Watanabe N, Forman HJ, Autoxidation of extracellular hydroquinones is a causative
event for the cytotoxicity of menadione and DMNQ in A549-S cells, Archives of
Biochemistry and Biophysics 2003; 411: 145-157.
[51] Aguilar-Martínez M, Bautista-Martinez J, Macias-Ruvalcaba N, González I, Tovar E,
Marín del Alizal T, et al. Molecular structure of substituted phenylamine α-OMe-and
α-OH-p-benzoquinone derivatives, Synthesis and correlation of spectroscopic,
electrochemical, and theoretical parameters, The Journal of Organic Chemistry 2001;
66: 8349-8363.
[52] Venugopala K, Jayashree B, Microwave-induced synthesis of Schiff bases of
aminothiazolyl bromocoumarins as antibacterials, Indian Journal of Pharmaceutical
Sciences 2008; 70: 88-95.
[53] Wadher S, Puranik M, Karande N, Yeole P, Synthesis and biological evaluation of
Schiff base of dapsone and their derivative as antimicrobial agents, International
Journal of PharmTech Research 2009; 1: 22-33.
[54] Cates LA, Rashed MS, Phosphorus GABA analogues as potential prodrugs,
Pharmaceutical Research 1984; 1: 271-274.
[55] Ashraf MA, Mahmood K, Wajid A, Maah MJ, Yusoff I, Synthesis, characterization
and biological activity of Schiff bases, International Conference on Chemistry and
Chemical Process 2011; 10: 1-7.
[56] Cimerman Z, Miljanic S, Galic N, Schiff bases derived from aminopyridines as
spectrofluorimetric analytical reagents, Croatica Chemica Acta 2000; 73: 81-95.
[57] Kabak M, Elmali A, Elerman Y, Keto–enol tautomerism, conformations and structure
of N-(2-hydroxy-5-methylphenyl), 2-hydroxybenzaldehydeimine, Journal of
Molecular Structure 1999; 477: 151-158.
Chapter 1 Introduction
Page 16
[58] Patel PR, Thaker BT, Zele S, Preparation and characterisation of some lanthanide
complexes involving a heterocyclic~diketone, Indian Journal of Chemistry 1999; 38:
563-567.
[59] Valcarcel M, De Castro ML. Flow–Through (Bio) Chemical Sensors: Elsevier; 1994.
[60] Asongalem E, Foyet H, Ekobo S, Dimo T, Kamtchouing P, Antiinflammatory, lack of
central analgesia and antipyretic properties of Acanthus montanus (Ness) T.
Anderson, Journal of Ethnopharmacology 2004; 95: 63-68.
[61] da Silva CM, da Silva DL, Modolo LV, Alves RB, de Resende MA, Martins CV,
Schiff bases: A short review of their antimicrobial activities, Journal of Advanced
Research 2011; 2: 1-8.
[62] Yang X-B, Wang Q, Huang Y, Fu P-H, Zhang J-S, Zeng R-Q, Synthesis, DNA
interaction and antimicrobial activities of copper (II) complexes with Schiff base
ligands derived from kaempferol and polyamines, Inorganic Chemistry
Communications 2012; 25: 55-59.
[63] Kalaivani S, Priya NP, Arunachalam S, Schiff bases: facile synthesis, spectral
characterization and biocidal studies, International Journal of Applied Biology and
Pharmaceutical Technology 2012; 3: 219-223.
[64] Liu S-Y, Nocera DG, A simple and versatile method for alkene epoxidation using
aqueous hydrogen peroxide and manganese salophen catalysts, Tetrahedron Letters
2006; 47: 1923-1926.
[65] Budakoti A, Abid M, Azam A, Synthesis and antiamoebic activity of new 1-N-
substituted thiocarbamoyl-3, 5-diphenyl-2-pyrazoline derivatives and their Pd (II)
complexes, European Journal of Medicinal Chemistry 2006; 41: 63-70.
[66] Kumamoto S, Watanabe M, Kawakami N, Nakamura M, Yamana K, 2′-
Anthraquinone-Conjugated Oligonucleotide as an Electrochemical Probe for DNA
Mismatch, Bioconjugate chemistry 2007; 19: 65-69.
[67] Naderi E, Ehteshamzadeh M, Jafari A, Hosseini M, Effect of carbon steel
microstructure and molecular structure of two new Schiff base compounds on
inhibition performance in 1M HCl solution by DC, SEM and XRD studies, Materials
Chemistry and Physics 2010; 120: 134-141.
[68] Şafak S, Duran B, Yurt A, Türkoğlu G, Schiff bases as corrosion inhibitor for
aluminium in HCl solution, Corrosion Science 2012; 54: 251-259.
Chapter 1 Introduction
Page 17
[69] Behpour M, Ghoreishi S, Mohammadi N, Soltani N, Salavati-Niasari M, Investigation
of some Schiff base compounds containing disulfide bond as HCl corrosion inhibitors
for mild steel, Corrosion Science 2010; 52: 4046-4057.
[70] Kianfar AH, Sobhani V, Dostani M, Shamsipur M, Roushani M, Synthesis,
spectroscopy, electrochemistry and thermal study of vanadyl unsymmetrical Schiff
base complexes, Inorganica Chimica Acta 2011; 365: 108-112.
[71] Kianfar A, Mohebbi S, Synthesis and electrochemistry of vanadium (IV) Schiff base
complexes, Journal of the Iranian Chemical Society 2007; 4: 215-220.
[72] Zollinger H, Front Matter: Wiley Online Library; 1995.
[73] Hamon F, Djedaini-Pilard F, Barbot F, Len C, Azobenzenes—Synthesis and
carbohydrate applications, Tetrahedron 2009; 65: 10105-10123.
[74] Pathak P, Jolly V, Sharma K, Synthesis and biological activities of some new
substituted arylazo Schiff bases, Oriental Journal of Chemistry 2000; 16: 161-162.
[75] Xu H, Zeng X, Synthesis of diaryl-azo derivatives as potential antifungal agents,
Bioorganic & Medicinal Chemistry letters 2010; 20: 4193-4195.
[76] Samadhiya S, Halve A, Synthetic utility of Schiff bases as potential herbicidal agents,
Oriental Journal of Chemistry 2001; 17: 119-122.
[77] Tonelli M, Vazzana I, Tasso B, Boido V, Sparatore F, Fermeglia M, Antiviral and
cytotoxic activities of aminoarylazo compounds and aryltriazene derivatives,
Bioorganic & Medicinal Chemistry 2009; 17: 4425-4440.
[78] Rani P, Srivastava V, Kumar A, Synthesis and antiinflammatory activity of
heterocyclic indole derivatives, European Journal of Medicinal Chemistry 2004; 39:
449-452.
[79] Rao GK, Kaur R, Pai PS, Synthesis and biological evaluation of dibenzo [b, f]
azepine-5-carboxylic acid [1-(substituted-phenyl)-ethylidene]-hydrazides. Journal for
Medicinal Chemistry, Pharmaceutical Chemistry and Computational Chemistry 2011;
3: 323-329.
[80] Rai B, Synthesis, Spectral and Antimicrobial Study of Co (II), Ni (II) and Cu (II)
Complexes with Schiff Bases of 3-Pyridinyl n-Pentyl Ketone, The Indian Council of
Chemists 2008; 25: 137-141.
[81] Munoz-Hernandez M-A, McKee ML, Keizer TS, Yearwood BC, Atwood DA, Six-
coordinate aluminium cations: characterization, catalysis, and theory, Dalton
Transactions 2002; 3: 410-414.
Chapter 1 Introduction
Page 18
[82] Kumar S, Dhar DN, Saxena P, Applications of metal complexes of Schiff bases—a
review, Journal of Scientific and Industrial Research 2009; 68: 181-187.
[83] Gerdemann C, Eicken C, Krebs B, The crystal structure of catechol oxidase: new
insight into the function of type-3 copper proteins, Accounts of Chemical Research
2002; 35: 183-191.
[84] Golcu A, Dolaz M, Demirelli H, Diorak M, Serin S, Spectroscopic and analytic
properties of new copper (II) complex of antiviral drug valacyclovir, Transition Metal
Chemistry 2006; 31: 658-665.
[85] Jungreis E, Thabet S, Analytical applications of Schiff bases, Chelates in Analytical
Chemistry 1969; 2: 149-177.
[86] Abbaspour A, Esmaeilbeig A, Jarrahpour A, Khajeh B, Kia R, Aluminium (III)-
selective electrode based on a newly synthesized tetradentate Schiff base, Talanta
2002; 58: 397-403.
[87] Kishore N, Mishra BB, Tiwari VK, Tripathi V, An account of phytochemicals from
Plumbago zeylanica (Family: Plumbaginaceae): A natural gift to human being,
Chronicles of Young Scientists 2012; 3: 178-185.
[88] Svoboda J, Novotna V, Kozmík V, Glogarova M, Weissflog W, Diele S, A novel type
of banana liquid crystals based on 1-substituted naphthalene-2, 7-diol cores, Journal
of Materials Chemistry 2003; 13: 2104-2110.
[89] Somsak L, Czifrak K, Toth M, Bokor E, Chrysina E, Alexacou K-M, New inhibitors
of glycogen phosphorylase as potential antidiabetic agents, Current Medicinal
Chemistry 2008; 15: 2933-2983.
[90] Johansen JS, Harris AK, Rychly DJ, Ergul A, Oxidative stress and the use of
antioxidants in diabetes: linking basic science to clinical practice, Cardiovascular
Diabetology 2005; 4: 1-11.
[91] Lagouge M, Larsson NG, The role of mitochondrial DNA mutations and free radicals
in disease and ageing, Journal Of Internal Medicine 2013; 273: 529-543.
[92] Rahal A, Kumar A, Singh V, Yadav B, Tiwari R, Chakraborty S, Oxidative stress,
prooxidants, and antioxidants: the interplay, Biomed Research International 2014;
2014.
[93] Raiola A, Rigano MM, Calafiore R, Frusciante L, Barone A, Enhancing the health-
promoting effects of tomato fruit for biofortified food, Mediators of Inflammation
2014; 2014.
Chapter 1 Introduction
Page 19
[94] Hamdy NA, Anwar MM, Abu-Zied KM, Awad HM, Synthesis, tumor inhibitory and
antioxidant activity of new polyfunctionally 2-substituted 5,6,7,8-
tetrahydronaphthalene derivatives containing pyridine, thioxopyridine and
pyrazolopyridine moieties, Acta Poloniae Pharmaceutica 2013; 70: 987-1001.
[95] Banerjee S, Veale EB, Phelan CM, Murphy SA, Tocci GM, Gillespie LJ, Recent
advances in the development of 1, 8-naphthalimide based DNA targeting binders,
anticancer and fluorescent cellular imaging agents, Chemical Society Reviews 2013;
42: 1601-1618.
[96] Bolton JL, Trush MA, Penning TM, Dryhurst G, Monks TJ, Role of quinones in
toxicology, Chemical Research in Toxicology 2000; 13: 135-160.
[97] Macıas-Ruvalcaba N, Gonzalez I, Aguilar-Martınez M, Evolution from hydrogen
bond to proton transfer pathways in the electroreduction of α-NH-quinones in
acetonitrile, Journal of The Electrochemical Society 2004; 151: E110-E118.
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 20
The history of electrochemistry may be tracked back to Michael Faraday in the first
half of nineteen century or even to Volta and Galvani in eighteen century. However it
expanded exponentially in 1960s and 1970s [1, 2]. All such techniques involve
recording of the resulting current by the application of potential to the electrodes.
Several sensitive voltammetric techniques have been invented and applied for
electrochemical investigations of both organic and inorganic species. A brief theory of
some useful electrochemical techniques is given below.
2.1 Cyclic Voltammetry Cyclic voltammetry is one of the basic and simplest electrochemical techniques for
the study of electroactive species based on mathematical treatment of reactions
involving electron transfer at an electrode surface [3]. It provides the basic criteria for
the investigation of redox portrait of analytes and identification of oxidizable and
reducible electrophores present in the target species [4].
2.1.1 Basic Principle of Cyclic Voltammetry An elementary single step electron transfer process may be represented as:
O + ne- R (2.1)
where ‘O’ and ‘R’ represent oxidized and reduced forms of electroactive species,
respectively.
Cyclic voltammetry based on scanning the potential of working electrode in a
triangular potential wave is shown in Fig. 2.1a. The current resulting from the applied
potential is measured. The plot of potential versus current is termed a cyclic
voltammogram. For a reversible redox process, a simplest cyclic voltammogram
related to a single potential cycle is depicted in Fig. 2.1b. When the potential is
scanned in positive direction, anodic current starts to flow and an oxidation peak is
observed at a potential represented as Epa. When the potential scanning is reversed,
the oxidized product gets reduced back to the original molecule (in case of reversible
reaction) at potential Epc. Forward and reverse reactions can be represented as:
A → An+ + ne- (2.1a)
An+ + ne- → A (2.1b)
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 21
Fig. 2.1 (a) Potential sweep in cyclic voltammery (b) A typical cyclic voltammogram
for an electroactive species
Cyclic voltammetry involves the scanning of potential in a cyclic way i.e.
direction of potential scan is reversed after time t = λ (at switching potential, Eλ).
Potential E (t) at any time t is given by:
E (t) =Ei + νt t<λ (2.2a)
E (t) = Ei + 2νλ - νt t≥λ (2.2b)
Here “ν’’ is scan rate in V/s. “Ei” is initial potential.
2.1.2 Nature of Electron Transfer Process Several electrochemical parameters including peak potential (Ep), peak current (ip)
and half peak potential (Ep/2) are used to ensure the nature of electron transfer process.
Mainly electron transfer processes are categorized into three types.
a. Reversible process.
b. Quasi-reversible.
c. Irreversible process
2.1.3 Reversible Process: Diagnostic Criteria of Reversibility Reversible reaction exhibits rapid electron transfer rate constant with a magnitude
greater than 10-2 cm s-1. Here, peak potential (Ep) is related to half wave potential
(E1/2) by the following relationships:
nFRTEE cp /)002.0109.1(2/1 ±−= (2.3a)
1/ 2 (1.109 0.002) /apE E RT nF= + ± (2.3b)
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 22
where E1/2 is potential corresponding to 88% of peal current (ip). ΔEp value calculated
from above equations is given by
Δ Ep= 2.22 /c ap pE E RT nF− = − (2.4)
At 298 K, equations 2.3a and 2.3b can be rearranged as
nEE ap /0285.02/1 += (2.5a)
nEE cp /0285.02/1 −= (2.5b)
At 298K, ΔEp is given by
∆𝐸𝑝 = 𝐸𝑝𝑐 − 𝐸𝑝𝑎 = −0.057𝑛
𝑉 (2.6)
For reversible reaction, peak potential and ΔEp does not depend on scan rate. In case
of a broader peak it is more useful to use Ep/2 which is given by
nFRTEE cp /09.12/12/ += (2.7)
At 298 K
nEE cp /028.02/12/ += (2.8)
Combining equations (2.3a) and (2.7) we get,
nFRTEE cp
cp /2.22/ =− (2.9)
At 298K, it is given by
nEE cp
cp /0565.02/ =− (2.10)
The above equations can be used to ensure the reversibility. Peak potential Ep
and peak separation ∆Ep are independent of sweep rate for reversible system [5]. It
can be seen that increase in number of electrons lowers the peak separation. Its values
decrease as 57 mV, 29 mV and 19 mV for one, two and three electrons respectively.
2.1.4 Quasi-Reversible Process Quasi-reversible is an intermediate case between reversible and irreversible processes
as shown in Fig. 2.2. Heterogeneous electron transfer rate constant varies between
10-2-10-5 cm sec-1. It can be represented as follows:
O + ne- R Separation of peak potentials ensures the quasi-reversible nature of the process
according to the following relation:
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 23
∆𝐸𝑝 = 𝐸𝑝𝑐 − 𝐸𝑝𝑎 ≥ �59𝑛�𝑚𝑉 ≤ �120
𝑛�𝑚𝑉 at 298 K (2.11)
Fig. 2.2 Cyclic voltammogram of quasi-reversible redox processes
2.1.5 Irreversible Process In totally irreversible reactions, only one cathodic (or anodic) peak appears with no
corresponding anodic (or cathodic) peak in the reverse direction. Reinmuth [6], Ayabe
and Matsuda [7] and Delahay [8] have offered detailed explanation for irreversible
process. Generally it can be described as:
O + ne- R Irreversibility of an electron transfer process obeys the following equations/rules.
i. Peak width is given by
/ 2 1.857p pa
RTE En Fa
− = (2.12)
At 298K, it can be solved to the form,
�𝐸𝑝 − 𝐸𝑝/2� = 47.7𝛼𝑛𝑎
𝑚𝑉 (2.13)
α is the charge transfer coefficient and na represents number of electrons in
charge transfer process.
ii. Peak potential is dependent on scan rate.
iii. Peak current can be related to scan rate and electron transfer rate constant (ks) by the
following equations:
𝐼𝑃 = 2.99 × 105𝑛(𝛼𝑛𝑎)1/2𝐴𝐶𝑜∗𝐷𝑜1/2𝜈1/2 (2.14)
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 24
For an irreversible system process, Ip in terms of rate constant is given as [9]:
hso kC
nFAI
,∗= (2.15)
Where I represents peak current while kos,h represents standard heterogeneous rate
constant for forward reaction. For a totally irreversible process, the standard
heterogeneous electron transfer rate constant ≤ 10-5 cm.sec-1.
2.1.6 Multi-Electron Transfer Process Usually multi-electron transfer process occurs in two separate steps, each step is
associated with its own electrochemical parameters [10]. Stepwise reversible electron
transfer followed by electron transfer called “EE mechanism” is given by the
following reaction
1 1 1( )oO n e R E+ ↔ (2.16)
1 2 2R n e R+ → (E2o) (2.17)
Fig. 2.3 Cyclic voltammogram for a reversible two-step system (a) ∆E = -180 mV (b)
∆E = -90 mV (c) ∆E = 0 mV. (d) ∆E = 180 mV (Source: Polcyn, D.S.; Shain, I. Anal.
Chem. 1966, 38, 370)
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 25
The value of ks, for first reversible electron transfer limiting case can be
calculated as [11]:
𝑘𝑠 = 𝑘𝑠°𝑒𝑒𝑒 �−𝛼1𝐹∆𝐸°
2𝑅𝑅� � (2.18)
Where,
2 1( )o o oE E E∆ = − (2.19)
Magnitude of ∆Eo ensures the dependent or independent nature of peaks. On the basis
of ∆Eo values, voltammograms can be in the forms shown in Fig. 2.3.
2.1.7 Types of Two-Electron Transfer Reactions Case 1: |𝐸20 − 𝐸10| ≥ 150 𝑚𝑉 (separate Peaks)
The EE mechanism is known as “disproportionate mechanism” if |∆Eo| values are
greater than 150 mV [12]. Two distinct peaks appear in cyclic voltammograms. Each
disproportionation reaction can be described as,
1 1R O n e→ + (2.20a)
1 2 2R n e R+ → (2.20b)
22
1
[0][ ][ ]disp
Rk
R= (2.21)
0 02 1ln ( )disp
nFk E ERT = −
(2.22)
Case 2: 2 1o oE E− < 100 mV (Peaks Overlapped)
A single wave appears in cyclic voltammograms and the peak current remains
intermediate between those of single step one electron and two electron transfer
reactions.
Case 3: 2 1o oE E− = 0 mV (Single peak)
In this case, cyclic voltammograms contain only single wave with intermediate peak
current and Ep-Ep/2 = 21 mV.
Case 4: 1 2o oE E< → 2nd Reduction is Easy than 1st one
∆Eo < -180 mV shows that energy for the first electron transfer process is greater than
the second one ( 1 2o oE E< ) and ∆Ep = 28.5 mV. A single wave appears with
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 26
characteristics of a direct two electron transfer (O + 2e → R2). Cyclic voltammetry
helps in getting the general redox portrait of the electroactive compounds.
2.2 Differential Pulse Voltammetry (DPV) Differential pulse voltammetry is more accurate and sensitive pulse electroanalytical
technique in which the potential of working electrode is applied in terms of
millisecond pulse and the current produced is measured in last few milliseconds of
each pulse [13, 14]. The resulting current difference is plotted against the applied
potential to get the differential pulse voltammogram as shown in Fig. 2.4.
Fig. 2.4 (a) Scheme of application of potential (b) General differential pulse
voltammogram
In DPV, the effect of charging current is minimized to enhance sensitivity and Faradic
current is extracted to get more reliable results. Therefore, these measurements can
record the redox properties of very small amounts of the analytes. DPV can detect the
analyte up to 0.01 µM concentration. Number of electrons involved in the redox
processes can be calculated from peak width at half peak current (W1/2) according to
following equation [15].
W1/2 = 3.52RTαnaF
(2.23)
At 298K, by considering αn=1 for reversible process the above equation is reduced to
𝑊1/2 = 90.4𝑛
(2.24)
where, n represents the number of electrons transferred in the process, R the general
gas constant, F the Faraday’s constant and T the temperature in Kelvin. Equation 2.24
depicts that the half peak width of magnitude 90, 45 and 30 mV involves one, two and
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 27
three electrons at 298 K in reversible process respectively [16-18]. Due to the
sensitivity of DPV, it is used to study the pH effect on the redox behavior of analytes
in order to obtain the number of protons involved in the process using the equation
[19-22]. 𝑑𝐸𝑝
𝑑𝑒𝑑� = 0.059𝑃 𝛼𝑛� (2.25)
2.3 Squarewave voltammetry (SWV) Squarewave voltammetry is fast and sensitive electroanalytical technique [23-25].
SWV consumes least amount of the analyte and it can measure forward and backward
current components at the same time. Another advantage of SWV is that a current
can be obtained at a fairly high effective scan rate due to its fast nature, thus
minimizing the scan time and it avoids the poisoning of electrode [26].
It
If
Ib
I/A
E / V Fig. 2.5 SWV showing total, forward and backward current
The current obtained from SWV can be resolved into forward and backward
current components and by using these components, reversible, quasi-reversible
and irreversible nature of the redox processes can be determined. As the
electroactive compounds find use in pharmaceutical and clinical applications against
numerous diseases, therefore, development of a sensitive method is essential for the
detection of their least amount. SWV is more sensitive and fast technique, capable of
detecting species produced for a very short time, so it can be employed for
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 28
determination of limit of detection (LOD) and limit of quantification (LOQ) values
using the following equations [27-30].
𝐿𝐿𝐷 = 3𝑆/𝑚 (2.26)
𝐿𝐿𝐿 = 10𝑆/𝑚 (2.27)
where S is the standard deviation of intercept and m the slope of the plot of peak
current versus concentration.
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 29
References [1] Stock JT, Orna MV, Electrochemistry, past and present, ACS Publications; 1989.
[2] Wahl D, A short history of electrochemistry-part 1, Galvanotechtnik; 2005.
[3] Kissinger PT, Heineman WR, Cyclic voltammetry, Journal of Chemical Education
1983; 60: 702.
[4] Klingler R, Kochi J, Electron-transfer kinetics from cyclic voltammetry: Quantitative
description of electrochemical reversibility, The Journal of Physical Chemistry 1981;
85: 1731-1741.
[5] Compton RG, Banks CE, Understanding voltammetry, World Scientific; 2007.
[6] Reinmuth W, Theory of stationary electrode polarography, Analytical Chemistry
1961; 33: 1793-1794.
[7] Matsuda H, Ayabe Y, Theoretical analysis of polarographic waves. I. Reduction of
simple metal ions, Bulletin of the Chemical Society of Japan 1955; 28: 422-428.
[8] Delahay P, Theory of irreversible waves in oscillographic polarography. Journal of
the American Chemical Society 1953; 75: 1190-1196.
[9] Shah AH, Zaid W, Shah A, Rana UA, Hussain H, Ashiq MN, Qureshi R, Badshah A,
Zia MB, Kraatz H-B, ph dependent electrochemical characterization, computational
studies and evaluation of thermodynamic, kinetic and analytical parameters of two
phenazines, Journal of the Electrochemical Society 2015; 162: H115-H123.
[10] Marcus RA, On the theory of oxidation‐reduction reactions involving electron
transfer. I, The Journal of Chemical Physics 1956; 24: 966-978.
[11] Fawcett WR, Opallo M, Kinetic parameters for heterogeneous electron transfer to tris
(acetylacetonato) manganese (III) and tris (acetylacetonato) iron (III) in aproptic
solvents, Journal of Electroanalytical Chemistry 1992; 331: 815-830.
[12] Lah J, Bezan M, Vesnaver G, Thermodynamics of berenil binding to poly [d (AT)]
poly [d (AT)] and poly [d (A)] poly [d (T)] duplexes, Acta Chimica Slovenica 2001;
48: 289-308.
[13] Gupta VK, Jain R, Radhapyari K, Jadon N, Agarwal S, Voltammetric techniques for
the assay of pharmaceuticals—a review, Analytical biochemistry 2011; 408: 179-196.
[14] Shuman LM, Sparks DL, Page AL, Helmke PA, Loeppert RH, Soltanpour PN,
Differential Pulse Voltammetry: Methods of Soil Analysis Part 3, Chemical Methods
1996: 247-268.
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 30
[15] Janeiro P, Brett AMO, Solid state electrochemical oxidation mechanisms of morin in
aqueous media, Electroanalysis 2005; 17: 733-738.
[16] Diculescu VC, Vivan M, Brett AMO, Voltammetric behavior of antileukemia drug
glivec. Part I-Electrochemical study of glivec, Electroanalysis 2006; 18: 1800-1807.
[17] Janeiro P, Brett AMO, Redox behavior of anthocyanins present in vitis vinifera L,
Electroanalysis 2007;19: 1779-1786.
[18] Jorge S, Pontinha A, Oliveira‐Brett A, Electrochemical redox behavior of omeprazole
using a glassy carbon electrode, Electroanalysis 2010; 22: 625-631.
[19] ShangGuan X, Zhang H, Zheng J, Electrochemical behavior and differential pulse
voltammetric determination of paracetamol at a carbon ionic liquid electrode,
Analytical and Bioanalytical Chemistry 2008; 391: 1049-1055.
[20] Lau O-W, Luk S-F, Cheung Y-M, Simultaneous determination of ascorbic acid,
caffeine and paracetamol in drug formulations by differential-pulse voltammetry
using a glassy carbon electrode, Analyst 1989; 114: 1047-1051.
[21] Sanghavi BJ, Srivastava AK, Adsorptive stripping differential pulse voltammetric
determination of venlafaxine and desvenlafaxine employing Nafion–carbon nanotube
composite glassy carbon electrode, Electrochimica Acta 2011; 56: 4188-4196.
[22] Oliveira-Brett AM, Silva LAd, Brett CM, Adsorption of guanine, guanosine, and
adenine at electrodes studied by differential pulse voltammetry and electrochemical
impedance, Langmuir 2002; 18: 2326-2330.
[23] Dogan-Topal B, Ozkan SA, Uslu B, The analytical applications of square wave
voltammetry on pharmaceutical analysis, The Open Chemical and Biomedical
Methods Journal 2010; 3: 56-73.
[24] Goyal RN, Oyama M, Singh SP, Fast determination of salbutamol, abused by athletes
for doping, in pharmaceuticals and human biological fluids by square wave
voltammetry, Journal of Electroanalytical Chemistry 2007; 611: 140-148.
[25] Karimi-Maleh H, Ensafi AA, Allafchian A, Fast and sensitive determination of
captopril by voltammetric method using ferrocenedicarboxylic acid modified carbon
paste electrode, Journal of Solid State Electrochemistry 2010; 14: 9-15.
[26] Lopes IC, Santos PV, Diculescu VC, Peixoto FM, Araújo MC, Tanaka AA, Oliveira-
Brett AM, Microcystin-LR and chemically degraded microcystin-LR electrochemical
oxidation, Analyst 2012; 137: 1904-1912.
Chapter 2 Theoretical Background of the Electrochemical Techniques
Page 31
[27] Bozic RG, West AC, Levicky R, Square wave voltammetric detection of 2,4,6-
trinitrotoluene and 2,4-dinitrotoluene on a gold electrode modified with self-
assembled monolayers, Sensors and Actuators B: Chemical 2008; 133: 509-515.
[28] Uslu B, Ozkan SA, Senturk Z, Electrooxidation of the antiviral drug valacyclovir and
its square-wave and differential pulse voltammetric determination in pharmaceuticals
and human biological fluids, Analytica Chimica Acta 2006; 555: 341-347.
[29] Blanco-Lopez M, Lobo-Castanon M, Miranda-Ordieres A, Tunon-Blanco P,
Electrochemical sensors based on molecularly imprinted polymers, TrAC Trends in
Analytical Chemistry 2004; 23: 36-48.
[30] Levent A, Yardim Y, Senturk Z, Voltammetric behavior of nicotine at pencil graphite
electrode and its enhancement determination in the presence of anionic surfactant,
Electrochimica Acta 2009; 55: 190-195.
Chapter 3 Experimental
Page 32
This chapter provides information about the chemicals and reagents used for synthesis
and characterization. The synthetic schemes of Schiff bases and metal complexes are
presented here. Structures of the selected quinones and naphthalene derivatives are
given. Moreover, the procedures used for electrochemical, photometric and biological
investigations have been discussed in this chapter.
3.1 Chemicals and reagents 2-Hydoxynaphthaldehyde, 4-chloroaniline, 4-bromoaniline, dimethyl sulfoxide acetic
acid, zinc acetate, cobalt acetate, magnesium chloride, ethylenediaminetetraacetic acid
disodium salt (Na2-EDTA), FeSO4, L-glutamine, hydrochloric acid, pyruvic acid,
penicillin-G, sodium chloride, sodium dodecyl sulfate (SDS), sodium sarcosinate,
streptomycin sulfate, 1,1,3,3-tetramethoxypropane, 2-thiobarbituric acid,
sulforhodamine B (SRB), trichloroacetic acid (TCA), thiazolyl blue tetrazolium
bromide (MTT), triton X-100, trizma-base and trypsin/EDTA (5%) were purchased
from Sigma–Aldrich. Keeping solubility in concentration fresh working solutions of
all the analytes were prepared in a mixture of ethanol and Britton–Robinson (BR)
buffer (1:1). BR buffers of wide pH range (2–12) were used as supporting electrolytes
for electrochemical studies.
3.2 Synthesis of Schiff Bases and complexes 3.2.1 1-((4-bromophenylimino) methyl) naphthalen-2-ol (BPIMN) The Schiff base BPIMN was synthesized by reacting 2-hydroxynaphthaldehyde (0.03
mmol, 5.0 mg) and 4-chloroaniline (0.03 mmol, 5.16 mg) in dry ethanol as shown in
Scheme 3.1. The obtained yellow crystalline solid was cooled, filtered and
recrystallized in a mixture of chloroform and pet ether (3:1). The formation of
BPIMN was ensured from the following data:
C17H12NOBr: Yield: 75 %, M.P.:162.5-164 oC; Mol. Wt.: 326, IR υ/cm-1: 3387
(OH), 1615 (C=N), 1299 (C-O). 1H NMR (CDCl3, 300 MHz) δ: 15.35 (s, 1H, -OH),
9.36 (s, 1H, H7, HC=N), 8.12 (d, 1H, H10, J = 8.4 Hz), 7.75 (d, 1H, H11, J = 8.4 Hz),
7.84 (d, 1H, H13, J = 9.3 Hz), 7.55 (t, 1H, H14, J = 8.4 Hz), 7.13 (d, 1H, H16, J = 9.3
Hz), 7.28-7.46 (m, 5H, H2, H6, H3, H5, H15 ) 13C NMR (75 MHz) δ: 144.7 (1C,
Chapter 3 Experimental
Page 33
C1), 121.8 (2C, C2, C6), 129.8 (2C, C3, C5), 123.7 (1C, C4), 155.8 (1C, C7), 109.0
(1C, C8), 168.5 (1C, C9), 119.0 (1C, C10), 128.2 (1C, C11), 136.6 (1C, C12), 133.0
(1C, C13), 127.5 (1C, C14), 129.4 (1C, C15), 121.5 (1C, C16), 132.0 (1C, C17).
3.2.2 1-((4-chlorophenylimino) methyl) naphthalen-2-ol (CPIMN) The Schiff base abbreviated as CPIMN was synthesized and recrystallized by the
method reported in literature [1-3] (Scheme 3.1). The yield and spectroscopic data of
CPIMN are given below:
CHO
OH
stirr 1 hr
NH2
X
HO
N
X
+C2H5OH
X=Br, Cl Scheme 3.1 Synthesis of 1-((4-bromophenylimino)methyl)naphthalen-2-ol and 1-((4-
chlorophenylimino)methyl)naphthalen-2-ol
C17H12NOCl: Yield: 75 %, M.P.:157-160 oC; Mol. Wt.: 281.5, IR υ/cm-1: 3399
(OH), 1620 (C=N), 1H NMR (CDCl3, 300 MHz) δ: 15.25 (s, 1H, -OH), 9.40 (s, 1H,
H7, HC=N), 8.12 (d, 1H, H10, J = 8.4 Hz), 7.75 (d, 1H, H11, J = 8.4 Hz), 7.84 (d, 1H,
H13, J = 9.3 Hz), 7.37 (t, 1H, H14, J = 7.8 Hz), 7.13 (d, 1H, H16, J = 9.3 Hz), 7.52-
7.61 (m, 3H, H2, H6, H15), 7.22-7.28 (m, 2H, H3, H5). 13C NMR (75 MHz) δ: 145.1
(1C, C1), 122.1 (2C, C2, C6), 132.7 (2C, C3, C5), 129.5 (1C, C4), 155.8 (1C, C7),
109.0 (1C, C8), 168.6 (1C, C9), 119.9 (1C, C10), 136.7 (1C, C11), 128.2 (1C, C12),
127.5 (1C, C13), 123.7 (1C, C14), 119.9 (1C, C15), 119.0 (1C, C16), 133.0 (1C,
C17).
3.2.3 1-((2, 4-dimethylphenylimino)methyl)naphthalen-2-ol (HL) Stoichiometric amounts of 2,4-dimethylbenzenamine and 2-hydroxy-1-
naphthaldehyde (5 mmol of each) were added to freshly dried ethanol. The mixture
was refluxed for 2 hrs. The obtained yellow crystalline solid was cooled, filtered and
crystallized in mixture of chloroform and pet ether (3:1) (Scheme 3.2).
Chapter 3 Experimental
Page 34
CH NH2
OH CH3
CH3
O
C2H5OHReflux
CH3H3C
NCH
OH
1
23
4
5
6
9
78
10 11
12
13
14
1516
17
1819
Scheme 3.2 Synthesis of Schiff base HL and its numbering scheme
The formation of HL was ensured from the following data.
Yield: 85 %, M.P.: 155.5 oC; Mol. Wt.: 275.33, IR υ/cm-1: 3446 (OH), 1615 (C=N),
1299 (C-O). 1H NMR (CDCl3, 300 MHz) δ: 15.09 (s, 1H, -OH), 9.27 (s, 1H, H9,
HC=N), 7.24 (d, 1H, H3, J = 1.2 Hz) 7.09-7.12 (m, 3H, H5, H6, H12), 2.35 (s, 3H,
H7), 2.44 (s, 3H, H8), 8.07 (d, 1H, H13 J = 8.8 Hz), 7.67 (d, 1H, H15, J = 8.8 Hz),
7.30 (d, 1H, H16, J = 8.0 Hz), 7.50 (d, 1H, H17, J = 8.0 Hz), 7.69 (d, 1H, H18, J = 8.0
Hz); 13C NMR (75 MHz) δ: 140.6 (1C, C1), 130.7 (1C, C2), 133.3 (1C, C3), 136.8
(1C, C4), 127.8 (1C, C5), 118.6 (1C, C6), 18.1 (1C, C7), 21.0 (1C, C8), 171.7 (1C,
C9), 108.6 (1C, C10), 153.1 (1C, C11), 116.9 (1C, C12), 131.8 (1C, C13), 129.3 (1C,
C14), 128.0 (1C, C15), 122.7 (1C, C16), 127.0 (1C, C17), 123.4 (1C, C18), 136.4
(1C, C19).
By using HL as ligand, its four metal complexes were also synthesized according to
Scheme 3.3.
3.2.4 Oxovanadium(IV) complex of 1-((2,4-dimethylphenylimino)
methyl)naphthalen-2-ol (L2VO) Oxovanadium complex (L2VO) was synthesized by treating the vanadyl(V)
isopropoxide with HL in 1:2 using acetonitrile as a solvent at room temperature as
shown in Scheme 3.3. The green colored product (L2VO) was precipitated out
immediately, filtered, washed with acetonitrile and then air dried.
Chapter 3 Experimental
Page 35
Complex (L2VO) : Yield: 76 %, M.P.:> 300 oC; Mol. Wt.: 615.61, IR υ/cm-1: 1599
(C=N), 1301 (C-O), 978 (V=O), 572 (V-O). μeff (BM) 2.04. Anal. Calc. for
[C38H32N2O3V] (%); C, 74.14; H, 5.24; N, 4.55; O, 7.80; Found: C, 74.55; H, 5.20;
N, 4.52; O, 8.01
3.2.5 Dibutyltin complex of 1-((2,4-dimethylphenylimino)methyl)
naphthalen-2-ol (L2Sn) Dibutyltin complex (L2Sn) was synthesized by refluxing the dibutyltindichloride with
HL in 1:2 ratio using toluene as a solvent at elevated temperature for four hours. The
yellow solution was concentrated by the process of evaporation under reduced
pressure to get the precipitates of (L2Sn).
Sn CH2-CH2-CH2-CH3Oα β γ δ
Complex (L2Sn) : Yield: 70 %, M.P.: 249.3oC; Mol. Wt.: 781.61 , IR υ/cm-1: 1590
(C=N), 1303 (C-O), 540 (Sn-C), 551 (Sn-O), 452 (Sn-N). 1H NMR (CDCl3, 300
MHz) δ: 9.28 (s, 1H, H9, HC=N), 7.28 (s, 1H, H3), 7.11-7.15 (m, 3H, H5, H6, H12),
2.39 (s, 3H, H7), 2.49 (s, 3H, H8), 8.10 (d, 1H, H13, J = 8.4 Hz), 7.72 (d, 1H, H15, J
= 7.8 Hz), 7.35 (t, 1H, H16, J = 8.1 Hz), 7.54 (t, 1H, H17, J = 8.1 Hz), 7.82(d, 1H,
H18, J = 8.4 Hz); 1.40-1.48 (m, 2H, αH) 1.78-1.85 (m, 4H, βH, γH), 0.97 (t, 3H, δH, J
= 7.2 Hz). 13C NMR (75 MHz) δ: 140.2 (1C, C1), 130.7 (1C, C2), 133.4 (1C, C3),
137.2 (1C, C4), 127.8 (1C, C5), 118.6 (1C, C6), 18.1 (3C, C7), 21.0 (3C, C8), 172.6
(1C, C9), 117.0 (1C, C10), 154.0 (1C, C11), 118.6 (1C, C12), 131.9 (1C, C13), 129.4
(1C, C14), 128.2 (1C, C15), 123.1 (1C, C16), 127.0 (1C, C17), 123.5 (1C, C18),
136.6 (1C, C19), 26.8 (2C, αC), 26.9 (2C, βC), 26.3 (2C, γC), 13.5 (3C, δC), 1J =
423.75/405.0 Hz (119/117Sn, 13C).
3.2.6 Zinc complex of 1-((2,4-dimethylphenylimino)methyl)
naphthalen-2-ol (L2Zn) Zinc complex (L2Zn) was synthesized by adding zinc acetate in HL solution in 1:2
using ethanol as a solvent at room temperature. The yellow colored product (L2Zn)
was precipitated out in 5-10 minutes, filtered, washed with ethanol and then air dried.
Chapter 3 Experimental
Page 36
Complex (L2Zn) : Yield: 78 %, M.P.:173 oC; Mol. Wt.:614.06 , IR υ/cm-1: 1591
(C=N), 1310 (C-O), 474 (Zn-O), 465 (Zn-N). 1H NMR (CDCl3, 300 MHz) δ: 9.30 (s,
1H, H9, HC=N), 7.30 (s, 1H, H3), 7.13-7.16 (m, 3H, H5, H6, H12), 2.35 (s, 3H, H7),
2.44 (s, 3H, H8), 8.09 (d, 1H, H13, J = 8.4 Hz), 7.70 (d, 1H, H15, J = 7.8 Hz), 7.33 (t,
1H, H16, J = 8.1 Hz), 7.56 (t, 1H, H17, J = 8.1 Hz), 7.80(d, 1H, H18, J = 8.4 Hz). 13C
NMR (75 MHz) δ: 140.1 (1C, C1), 130.6 (1C, C2), 133.3 (1C, C3), 136.9 (1C, C4),
127.8 (1C, C5), 118.6 (1C, C6), 18.1 (3C, C7), 21.0 (3C, C8), 171.8 (1C, C9), 108.5
(1C, C10), 155.1 (1C, C11), 117.0 (1C, C12), 129.5 (1C, C13), 129.3 (1C, C14),
128.1 (1C, C15), 122.8 (1C, C16), 127.0 (1C, C17), 123.4 (1C, C18), 136.4 (1C,
C19).
HC N
OH
HC N
O
CH
N
OM
CH3
CH3
H3C
CH3
H3C
H3C
2X
X= VO(OCHMe2)3, (C4H9)2SnCl2, (CH3COO)2Zn, (CH3COO)2Co
M= V=O, (C4H9)2Sn, Zn, Co
Scheme 3.3 Synthesis of metal complexes of HL
3.2.7 Cobalt complex of 1-((2,4-dimethylphenylimino)methyl)
naphthalen-2-ol (L2Co) Cobalt complex (L2Co) was synthesized by treating cobalt acetate with HL in 1:2 ratio
using ethanol as a solvent at room temperature according to Scheme 3.3. The reddish
brown colored product (L2Co) was precipitated out in few minutes, filtered, washed
with ethanol and then air dried. The formation of oxovanadium and cobalt complex
was ensured from the following thermo gravimetric analysis (TGA) data shown in
Table 3.1.
Complex (L2Co) : Yield: 80 %, M.P.:183 oC; Mol. Wt.: 607.61 , IR υ/cm-1: 1600
(C=N), 1305 (C-O), 474 (Co-O), 507 (Co-N). μ eff (BM) 1.81. Anal. Calc. for
Chapter 3 Experimental
Page 37
[C38H32N2O2Co] (%); C, 75.12; H, 5.31; N, 4.61; O, 5.27; Found: C, 75.35; H, 5.22;
N, 4.75; O, 5.43.
Table 3.1 Thermo gravimetric data of complexes (L2VO) and (L2Co)
Complex Temp.
range oC
Evolved
components
% Weight loss % Residue
Observed Calculated Observed Calculated
(L2VO) 24-765 2C8H9N 38.61 38.71 C9H6O3V
765-951 C13H8 26.91 26.67 34.48 34.62
(L2Co) 198-658 C38H32N2 84.55 85.03 CoO2
15.45 14.97
3.3 Hydroxy substituted quinones Quinones are promising candidates for the development of a variety of drugs due to
their therapeutic [4-6], laxative [7] and other pharmacological characteristics [8-10].
The present study also includes investigations of some substituted quinones for their
electrochemical and biological activities. The reduction of anthraquinones is
extensively reported. However, meager information is available about their electro-
oxidation. To bridge this gap in literature, we carried out their detailed pH dependent
electrochemical investigation. Names and structures of selected quinone derivatives
are given in Table 3.2.
Chapter 3 Experimental
Page 38
Table 3.2 Names and Structures of selected quinone derivatives
No Code Name Structure
1 HMNA 4-hydroxy-5-
methoxynaphthalen-1-
yl acetate
O
OHOCH3
C
O
CH3
2 HAC 1-hydroxy-
2(hydroxymethyl)anthr
acene-9,10-dione
O
O
OH
CH2OH
3 DHDN 1,8-dihydroxy-4,5-
dinitroanthracene-9,10-
dione
O
O
OHOH
NO2 NO2 4 HCAQ 1,4-dihydroxy-2-(3-
hydroxy-3-
(trichloromethyl)pentyl)
-8-methoxyanthracene-
9,10-dione
O
O
OH
OH
CCl3
OHOCH3
5 HACAD 4,8-dihydroxy-9,10-
dioxo-9,10-
dihydroanthracen-1-yl
acetate
O
O
OH
OOH C
O
CH3
Chapter 3 Experimental
Page 39
6 HAD 1,4,5-
trihydroxyanthracene-
9,10-dione
O
O
OH
OHOH 7 HOAD 1,4,5-trihydroxy-2-
methyl-3-(3-
oxobutyl)anthracene-
9,10-dione
O
O
OH
OHOH
CH3
CH2CH2COCH3
3.4 Naphthalene derivatives A number of naphthalene derivatives are endowed with broad range biological
activities [11-15]. Various studies have reported naphthalene derivatives as potent
anti-diabetic agents. In order to explore their electrochemical, biological and other
pharmaceutical properties, the following two representative derivatives of
naphthalene, 1,8-dimethoxynaphthalene and naphthalene-2,3-dicarboxylic acid were
investigated in detail.
3.5 Voltammetric studies DY2100 Series Mini Potentiostat, USA was used for electrochemical measurements.
Glassy carbon with electro-sensing area of 7.1×10-6 m2 was used as working
electrode. Ag/AgCl (saturated with KCl) and platinum wire (1 mm diameter) were
used as reference and counter electrodes, respectively. The surface of working
electrode was cleaned before each run using 0.3 micron diamond powder. The
polished electrode surface was carefully washed with doubly distillated water for 30 s.
The cleaned electrode was subsequently placed in the solvent and various
voltammograms were recorded for getting steady state baseline. This practice ensured
the reproducibility of experimental results. The measurement cell was immersed in a
water circulating bath (I-2400, IRMECO GmbH, Germany) in order to control the
temperature. The pH measurements were carried out using INOLAB pH meter with
Model no pH 720. Differential pulse voltammetric (DPV) experiments were carried
out at 5 mV s-1 scan rate. In square wave voltammetry (SWV), a scan rate of
Chapter 3 Experimental
Page 40
100 mV s-1 was used by setting the experimental conditions of 20 Hz frequency and
5 mV potential increments [16-18]. Before each experiment, oxygen was flushed out
by bubbling nitrogen gas through the solution for ten minutes. Nitrogen saturated air
was maintained over the solution during all the electrochemical experiments.
3.6 Spectroscopic measurements Melting points were determined using Gallen Kamp apparatus. Infrared measurements
(4000–400 cm-1) were performed on thermoscientific NICOLET 6700 FTIR
spectrophotometer. Multinuclear (1H and 13C NMR) spectra were recorded in solution
on Bruker ARX 300 MHz using tetramethylsilane (TMS) as internal reference.
Absorption spectroscopy was carried out using Shimadzu 1601 spectrophotometer.
Crison micro pH 2001 pH-meter with an Ingold combined glass electrode was used
for the pH measurements. Fresh working solutions of the analytes were prepared in
50% ethanol and 50% BR buffers.
3.7 Computational Studies Gaussian 5.0 package was used for computational studies. Density functional theory
at Becke three parameter Lee–Yang–Parr exchange correlation functional (B3LYP)
and 6-311++G** minimal bases set were used for the geometric optimization of non-
metallic compounds while 6-311G++ minimal bases set were chosen for all metal
complexes due to their larger structures. Charge distribution, energy calculations of
HOMO and LUMO, structural parameters and NMR spectrum of Schiff base were
measured using most stable conformation. 1H NMR and 13C NMR calculated
chemical shifts of HL obtained by GIAO method and using TMS as reference were
compared with experimental shift values. DFT and 6-311G++ /6-311++G** were
selected due to its success in reliable calculation of charge and energy [19-23].
3.8 Anti-diabetic activity Anti-diabetic activity of the analytes was determined in the laboratory of Animal and
Human Physiology, Department of Animal Sciences, Quaid-i-Azam University
Islamabad. All animal handling was done according to the guidelines provided by the
‘‘Ethics Committee of the Department of Animal Sciences on humane use of animals
for scientific research’’ [24-26].
Chapter 3 Experimental
Page 41
3.8.1 Animals and maintenance Healthy adult male Sprague–Dawley rats (n = 20, average body wt = 240 ± 10 g)
obtained from the National Institute of Health, Islamabad were maintained at the
Animal House Facility of Biological sciences, Quaid-i-Azam University, Islamabad.
Five rats were kept in cage and provided rodent diet and water. Light/dark cycle was
kept for 12:12 h.
3.8.2 Induction of diabetes Diabetes was induced by a single intraperitoneal injection of alloxan monohydrate
(Sigma Aldrich, USA) at the dose of 150 mg/kg b.w. Fasting plasma glucose levels
>200 mg/dl were considered diabetic. These animals were then used for the current
study.
3.8.3 Experimental design Four groups of rats were formulated (n = 5); groups 3 and 4 included those that were
treated with acute doses of compounds; negative control group was treated with only
alloxan while positive control group was treated with glibenclamide (10 mg/kg b.w in
distilled H2O), (Euglucon, Roche Pharma). The following are the control and
experimental treatment groups.
3.8.4 Control and experimental groups Group 1: Negative control (diabetic, only alloxan pretreated)
Group 2: Positive control (diabetic, alloxan pretreated followed by glibenclamide
treatment).
Group 3: Compound treated (diabetic, alloxan pretreated)
3.8.5 Glucose determination Blood was collected through caudal venepuncture using a 26 gauge butterfly cannula.
Blood samples were drawn at 0 (Pre Alloxan)–0 (Post Alloxan) h, 1, 2, 3, 4, 5, 6 and 7
h after dosage. Plasma glucose was determined with a dextrostix using glucometer
(Accu-Check Active, Roche).
Chapter 3 Experimental
Page 42
3.8.6 Total serum cholesterol and triglycerides determination Total serum cholesterol and triglycerides were determined by using commercially
available kits obtained from Globe Diagnostics S.r.1 (Italy).
3.8.7 Statistical analysis Data were analyzed through one-way analysis of variance (ANOVA) using the
Statistical Package for Social Sciences (SPSS version 16.0 Inc. Chicago, Illinois,
USA) Post-hoc Tukey–Kramer test. P < 0.05 was considered statistically significant
difference. Data are presented as line or bar diagrams constructed using the GraphPad
Prism 5 (Version 5.01 GraphPad Software Inc. USA). Data from the biological
activity experiments are presented as mean ± standard deviation (SD). One-way
ANOVA and two tailed t-test were performed to compare differences in NTC, ethanol
(solvent) treated and compound treated samples. The results were considered
significant when P < 0.05.
3.9 Cytotoxicity and antioxidant activity 3.9.1 Cell lines and cell culture Human cervix adenocarcinoma (HeLa, ATCC CCl-2) cells were cultured in
Dulbecco’s Modified Eagle Medium (DMEM) supplemented with 10% FBS, 2 mM
L-glutamine, 1 mM Na-pyruvate, 100 U/ml penicillin, 100 lg/ml streptomycin at
37 ºC in a humidified 5% CO2 atmosphere. Cells were harvested by trypsinization
with 1 mL 0.5 mM trypsin/EDTA for 1 min at room temperature.
3.9.2 Cell survival and proliferation assay MTT assay was used to quantitatively assess cell viability by using the reported
method of Mossman with some modifications [27-29]. In brief, 1.5×105 cells/mL
were seeded in 96-well plates and allowed to grow over night. Cultures were
incubated with different dilutions of the compounds (10, 50, 100 and 150 µM) or
absolute ethanol for 24 h as controls untreated cultures (NTC) were included in the
experiment. Samples representing the non-cellular background i.e., media only and
compounds only samples containing MTT and acidified 10% SDS solution were also
included in each experiment. After the exposure, the culture medium was removed
Chapter 3 Experimental
Page 43
from each well and replaced with 100 µl of new medium containing MTT
(0.5 mg/mL) and incubated for 3 h at 37 ºC. The resulting formazan product was
dissolved by adding equal amount of acidified 10% SDS and further incubation for
18 h at 37 ºC. Absorbance was measured at 565 nm by using a microplate reader
(AMP PLATOS R-496). The non-cellular background was subtracted from the
respective samples and percent viability was measured relative to the NTC sample
using the following formula:
Percent (%)viability = �Corrected Absorbance of test samplesCorrected Absorbance of NTC
� × 100
The experiments were performed twice with triplicates for each sample. IC50 values
representing the concentrations that inhibit 50% of cell growth were calculated where
required.
3.9.3 Biological antioxidant activity: lipid peroxidation The extent of membrane lipid peroxidation (LPO) was estimated by measuring the
formation of malondialdehyde (MDA) using the thiobarbituric acid-reactive species
(TBARs) assay as described by Ohkawa et al., with some modifications [30]. MDA is
one of the products of membrane LPO. To assay the capacity of the compounds for
protecting HeLa cells from ROS-mediated oxidative injury, briefly, pre-seeded cells
(>90% viability; 1.5 × 105 cells/ml) were incubated with different dilutions of the
compounds (10, 50 and 100 µM) for 24 h. Oxidative stress was induced by the
addition of Fe++ (as FeSO4) to the cells at a final concentration of 100 µM, for 1 h.
NTC and non-cellular background (media only and compounds only samples
containing cell lysis buffer, 10% SDS and TBA solution) were included in each
experiment as controls. After the treatment, cells were harvested and lysed in cell
lysis buffer (2.5 M NaCl, 10 mM trizma-base at pH 10.0, 100 mM Na2EDTA, 1%
sodium sarcosinate, 1% Triton X-100, 10% DMSO). To the cell lysate (0.1 ml) an
equal amount of 10% SDS was added and samples were incubated at room
temperature for 5 min. TBA (5.2 mg/ml; 0.25 ml) was added to the mixture and
incubated at 95 ºC for 45 min. After cooling to room temperature, absorbance of the
mixture was taken at 532 nm by using NanoDrop 2000 (Thermo Fisher Scientific).
The non-cellular background was subtracted from the respective samples and MDA
(µM) for each sample were calculated from 1,1,3,3-tetramethoxypropane standard
curve (range of 0–100 µM) using the following formula.
Chapter 3 Experimental
Page 44
MDA (µM) = �(Corrected absorbance) − (y − intercept)
Slope�
MDA (µM) was calculated relative to the NTC sample. The experiments were
performed twice with triplicates for each sample.
3.10 Assay of alkaline phosphatase activity For alkaline phosphatase activity of the synthesized Schiff bases, the reported method
[31-33] with slight modification was used. Working substrate was made by mixing
four parts of reagent A (ethanolamine pH 9.8 mol/dm3 and Magnesium chloride
0.5 mmol/dm3) and one part of reagent B (p-nitrophenyl phosphate 50 mmol/dm3).
Substrate was incubated for 5 min at 25 ºC. A 2mL of the substrate was taken in a
cuvette and 40 µL of human serum having activity of 165 IU/L was added. After
incubation for 1 min, absorbance was measured to check the activity of enzyme. ALP
hydrolyzed p-NPP and produced a yellow colored p-nitrophenol that showed
absorbance at 405 nm (Scheme 3.4). After that various amounts of Schiff bases
(20 µL, 40 µL, 60 µL) were periodically added from 6.25 mmolar stock solution and
again incubated for 3 min. Absorbance was recorded after 1–5 min and in the end
percentage inhibition was calculated from the average.
p-Nitropheyl phosptahe + Mg2+ + H2O p-Nitrophenol + PiALP
Pi represents inorganic phosphate Scheme 3.4 Hydrolysis of alkaline phosphatase
3.11 Antimicrobial activities The fungicidal and bactericidal activities of HL and its metallic complexes were
carried out by disc diffusion method as following [29].
3.11.1 Growth medium, culture and inoculum preparation The bacterial strains (Escherichia coli, Bacillus subtilis, Staphlocuccus aureus and
Pasturella multocida) were cultured overnight at 310 K in a nutrient agar. The pure
bacterial cultures gained were maintained in the medium in slants and petri plates. For
inoculums preparation, 13 g of the nutrient broth was suspended in doubly distilled
water (1 litre), mixed homogenously and autoclaved for 15 minutes at 394 K. Then
Chapter 3 Experimental
Page 45
10 µL of the pure culture of a bacterial strain was added to a 100 mL of freshly
prepared nutrient broth medium and shaked (140 revolutions per minute) at 310 K for
24 hours. The prepared inocula were kept at 277 K. The inocula with 1×108
spores/mL were utilized for activity measurement [34-36].
The fungal strains (Alternaria alternata, Ganoderma lucidum, Aspergillus
niger and Penicillium notatum) were cultured overnight at 310 K using potato
dextrose agar. The pure cultures were maintained in a sabouraud dextrose agar (SDA)
medium in slants and petri plates which were presterilized in a hot air oven at 453 K
for 3 hours. These cultured slants were incubated at 301 K for 3 to 4 days for the
growth of fungal strains.
3.11.2 Antimicrobial assay by disc diffusion method Antimicrobial activities were measured by using the disc diffusion method [37-39].
For antibacterial action 2.8 g nutrient agar or for antifungal investigations, 3.9 g
potato dextrose agar was suspended in 0.1 liter of doubly distilled water and sterilized
by autoclaving at 394 K for 900 seconds. It was then mixed well with 100 μL
inoculums and was added into sterilized petri plates. Finally a 9 mm filter paper discs
soaked with 100 μL of a specific solution were laid flat on growth medium. The petri
plates were then incubated at 310 K for 24 hours or 301 K for 48 hours for the
manipulation of bacteria or fungi, respectively. Inhibition zones formed around the
discs were recorded in millimeters (mm) using a zone reader [40, 41]. Following
equation was used to calculate inhibition percentage.
Linear growth in test sample (mm)% Growth inhibition = 100 - × 100Linear growth in control (mm)
Chapter 3 Experimental
Page 46
References [1] Pierens GK, Venkatachalam T, Bernhardt PV, Riley MJ, Reutens DC, A solid state
study of keto-enol tautomerismin three naphthaledene Schiff bases, Australian Journal
of Chemistry 2012; 65: 552-556.
[2] Naeimi H, Safari J, Heidarnezhad A. Synthesis of Schiff base ligands derived from
condensation of salicylaldehyde derivatives and synthetic diamine, Dyes and
Pigments 2007; 73: 251-253.
[3] Sinha D, Tiwari AK, Singh S, Shukla G, Mishra P, Chandra H, Synthesis,
characterization and biological activity of Schiff base analogues of indole-3-
carboxaldehyde, European Journal of Medicinal Chemistry 2008; 43: 160-165.
[4] Monks TJ, Hanzlik RP, Cohen GM, Ross D, Graham DG, Quinone chemistry and
toxicity, Toxicology and Applied Pharmacology 1992; 112: 2-16.
[5] Kamei H, Koide T, Kojima T, Hashimoto Y, Hasegawa M, Inhibition of cell growth
in culture by quinones, Cancer Biotherapy & Radiopharmaceuticals 1998; 13:
185-188.
[6] Huang Q, Lu G, Shen HM, Chung M, Ong CN, Anti‐cancer properties of
anthraquinones from rhubarb. Medicinal Research Reviews 2007; 27: 609-630.
[7] Hubacher MH, Doernberg S, Horner A, Laxatives: chemical structure and potency of
phthaleins and hydroxyanthraquinones, Journal of the American Pharmaceutical
Association 1953; 42: 23-30.
[8] Ghosheh OA, Houdi AA, Crooks PA, High performance liquid chromatographic
analysis of the pharmacologically active quinones and related compounds in the oil of
the black seed (Nigella sativa L.), Journal of Pharmaceutical and Biomedical Analysis
1999; 19: 757-762.
[9] Loadman PM, Phillips RM, Lim LE, Bibby MC, Pharmacological properties of a new
aziridinylbenzoquinone, RH1 (2, 5-diaziridinyl-3-(hydroxymethyl)-6-methyl-1, 4-
benzoquinone), in mice. Biochemical Pharmacology 2000; 59: 831-837.
[10] Monks TJ, Jones DC, The metabolism and toxicity of quinones, quinonimines,
quinone methides, and quinone-thioethers, Current Drug Metabolism 2002; 3:
425-438.
[11] Xie X, Kozlowski MC, Synthesis of the naphthalene portion of the rubromycins,
Organic Letters 2001; 3: 2661-2663.
Chapter 3 Experimental
Page 47
[12] Dai J, Krohn K, Flörke U, Draeger S, Schulz B, Kiss‐Szikszai A, Metabolites from
the endophytic fungus nodulisporium sp. from juniperus cedre, European Journal of
Organic Chemistry 2006; 3: 498-506.
[13] Reddy RA, Baumeister U, Keith C, Hahn H, Lang H, Tschierske C, Influence of the
core structure on the development of polar order and superstructural chirality in liquid
crystalline phases formed by silylated bent-core molecules: lateral substituents, Soft
Matter 2007; 3: 558-570.
[14] Kim W, Khil L, Clark R, Bok S, Kim E, Lee S, Naphthalenemethyl ester derivative of
dihydroxyhydrocinnamic acid, a component of cinnamon, increases glucose disposal
by enhancing translocation of glucose transporter 4, Diabetologia 2006; 49:
2437-2448.
[15] Svoboda J, Novotna V, Kozmík V, Glogarova M, Weissflog W, Diele S, A novel type
of banana liquid crystals based on 1-substituted naphthalene-2, 7-diol cores, Journal
of Materials Chemistry 2003; 13: 2104-2110.
[16] Behpour M, Valipour A, Keshavarz M, Determination of buprenorphine by
differential pulse voltammetry on carbon paste electrode using SDS as an
enhancement factor, Materials Science and Engineering: C 2014; 42: 500-505.
[17] Guerreiro GV, Zaitouna AJ, Lai RY, Characterization of an electrochemical mercury
sensor using alternating current, cyclic, square wave and differential pulse
voltammetry, Analytica Chimica Acta 2014; 810: 79-85.
[18] Zeng L, Zhang A, Zhu X, Zhang C, Liang Y, Nan J. Electrochemical determination of
nonylphenol using differential pulse voltammetry based on a graphene–DNA-
modified glassy carbon electrode, Journal of Electroanalytical Chemistry 2013; 703:
153-157.
[19] Lizarraga E, Romano E, Raschi AB, Leyton P, Paipa C, Catalán CA, A structural and
vibrational study of dehydrofukinone combining FTIR, FTRaman, UV–visible and
NMR spectroscopies with DFT calculations, Journal of Molecular Structure 2013;
1048: 331-338.
[20] Saheb V, Sheikhshoaie I, Setoodeh N, Rudbari HA, Bruno G, A new tridentate Schiff
base Cu (II) complex: synthesis, experimental and theoretical studies on its crystal
structure, FT-IR and UV–Visible spectra, Spectrochimica Acta Part A: Molecular and
Biomolecular Spectroscopy 2013; 110: 124-129.
[21] Karthikeyan N, Prince JJ, Ramalingam S, Periandy S, Electronic [UV–visible] and
vibrational [FT-IR, FT-raman] investigation and NMR–mass spectroscopic analysis
Chapter 3 Experimental
Page 48
of terephthalic acid using quantum Gaussian calculations, Spectrochimica Acta Part
A: Molecular and Biomolecular Spectroscopy 2015; 139: 229-242.
[22] de Toledo T, Pizani P, da Silva L, Teixeira A, Freire P, Spectroscopy studies on
Schiff base N, N′-bis (salicylidene)-1, 2-phenylenediamine by NMR, infrared, raman
and DFT calculations, Journal of Molecular Structure 2015; 1097: 106-111.
[23] Nielsen CJ, Klaeboe P, Andersen NH, Guirgis GA, Hickman DV, Morris TB, Infrared
and raman spectra, DFT-calculations and spectral assignments of 1,1,3,3,5,5-
hexafluoro-1, 3, 5-trisilacyclohexane, Journal of Molecular Structure 2015; 1099:
399-406.
[24] Foxman B, Molecular tools and infectious disease epidemiology: Academic Press;
2010.
[25] Touitou Y, Portaluppi F, Smolensky MH, Rensing L, Ethical principles and standards
for the conduct of human and animal biological rhythm research, Chronobiology
International 2004; 21: 161-170.
[26] Kilkenny C, Browne W, Cuthill IC, Emerson M, Altman DG, Animal research:
reporting in vivo experiments: the arrive guidelines, the Journal of Gene Medicine
2010; 12: 561-563.
[27] Mosmann T, Rapid colorimetric assay for cellular growth and survival: application to
proliferation and cytotoxicity assays, Journal of Immunological Methods 1983; 65:
55-63.
[28] Kaennakam S, Siripong P, Tip-pyang S, Velucarpins A–C, three new pterocarpans
and their cytotoxicity from the roots of dalbergia velutina, Fitoterapia 2015; 105:
165-168.
[29] Sun X-Y, Ouyang J-M, Liu A-J, Ding Y-M, Gan Q-Z, Preparation, characterization,
and in vitro cytotoxicity of COM and COD crystals with various sizes, Materials
Science and Engineering: C 2015; 57: 147-156.
[30] Ohkawa H, Ohishi N, Yagi K, Assay for lipid peroxides in animal tissues by
thiobarbituric acid reaction, Analytical Biochemistry 1979; 95: 351-358.
[31] Malik MR, Vasylyeva V, Merz K, Metzler-Nolte N, Saleem M, Ali S, Synthesis,
crystal structures, antimicrobial properties and enzyme inhibition studies of zinc (II)
complexes of thiones, Inorganica Chimica Acta 2011; 376: 207-211.
[32] Shrestha S, Diogenes A, Kishen A, Temporal-controlled release of bovine serum
albumin from chitosan nanoparticles: effect on the regulation of alkaline phosphatase
Chapter 3 Experimental
Page 49
activity in stem cells from apical papilla, Journal of Endodontics 2014; 40:
1349-1354.
[33] Santos IC, Mesquita RB, Bordalo AA, Rangel AO, Use of solid phase extraction for
the sequential injection determination of alkaline phosphatase activity in dynamic
water systems, Talanta 2012; 98: 203-210.
[34] Jabeen M, Ali S, Shahzadi S, Shahid M, Khan Q, Sharma S, Homobimetallic
complexes of ligand having O-and S-donor sites with same and different di-and
trialkyl/aryltin (IV) moiety: their synthesis, spectral characterization and biological
activities. Journal of the Iranian Chemical Society 2012; 9: 307-320.
[35] Tan H, Zhou S, Deng Z, He M, Cao L. Ribosomal-sequence-directed selection for
endophytic streptomycete strains antagonistic to Ralstoniasolanacearum to control
tomato bacterial wilt, Biological Control 2011; 59: 245-254.
[36] Khan AV, Ahmed QU, Mir MR, Shukla I, Khan AA, Antibacterial efficacy of the
seed extracts of Melia azedarach against some hospital isolated human pathogenic
bacterial strains, Asian Pacific Journal of Tropical Biomedicine 2011; 1: 452-455.
[37] Rizwan K, Zubair M, Rasool N, Riaz M, Zia-Ul-Haq M, de Feo V, Phytochemical
and biological studies of Agave attenuata, International Journal of Molecular Sciences
2012; 13: 6440-6451.
[38] Kronvall G, Giske CG, Kahlmeter G, Setting interpretive breakpoints for
antimicrobial susceptibility testing using disk diffusion, International Journal of
Antimicrobial Agents 2011; 38: 281-290.
[39] Fujisaki M, Sadamoto S, Ikedo M, Totsuka K, Kaku M, Tateda K, Development of
interpretive criteria for tebipenem disk diffusion susceptibility testing with
staphylococcus spp. and haemophilus influenzae, Journal of Infection and
Chemotherapy 2011; 17: 17-23.
[40] Bhalodia NR, Shukla V, Antibacterial and antifungal activities from leaf extracts of
cassia fistula l.: an ethnomedicinal plant. Journal of Advanced Pharmaceutical
Technology & Research 2011; 2: 104-109.
[41] Huynh Q, Borgmeyer J, Smith C, Bell L, Shah D, Isolation and characterization of a
30 kDa protein with antifungal activity from leaves of engelmannia pinnatifida.
Biochemal Journal 1996; 316: 723-727.
Chapter 4 Results & Discussion
Page 50
This chapter is devoted to the interpretation of the results obtained from the detailed
characterization of the synthesized Schiff bases, selected quinones and naphthalene
derivatives. The outcome of the biological studies related to the evaluation of
antioxidant, antidiabetic, cytotoxic, enzyme inhibition and antimicrobial potential of
the compounds have also been discussed in this chapter.
4.1 Synthesis, photometric and electrochemical fate of
Schiff bases
4.1.1 Structural characterization The Schiff bases 1-((4-bromophenylimino) methyl) naphthalen-2-ol (BPIMN) and 1-
((4-chlorophenylimino) methyl) naphthalen-2-ol (CPIMN) were structurally
characterized by spectroscopic techniques. FT-IR spectrum of the newly synthesized
Schiff base, BPIMN, showed a strong absorption band around 1734 cm-1 due to
azomethine group. The appearance of proton resonance signal of -CH=N at 9.15 ppm
in the 1H NMR spectrum confirmed the formation of Schiff base. The absence of any
resonance peak corresponding to primary amine supported the purity of the product
i.e. Schiff base. The characteristic resonance signal of -CH=N at 168.63 in the 13C
NMR spectrum also verified the formation of the product.
4.1.2 Crystal structures description of BPIMN and CPIMN
A B
Fig. 4.1 Ball and sticks diagrams of BPIMN (A) and CPIMN (B)
Chapter 4 Results & Discussion
Page 51
Table 4.1 Crystal data and details of the structure refinement for BPIMN and CPIMN
Structure refinements BPIMN CPIMN Empirical formula C17H12NOBr C17H12NOCl Formula weight 326.19 281.73 Temperature/K 100 100 Crystal system Monoclinic Monoclinic Space group P21/n P21/n a/Å 4.7671(3) 4.7136(6) b/Å 20.5291(11) 20.3076(18) c/Å 13.5235(7) 13.5174(18) α/° 90 90 β/° 93.808(3) 93.830(8) γ/° 90 90 Volume/Å3 1320.55(13) 1291.0(3) Z 4 4 ρcalcmg/mm3 1.641 1.449 m/mm-1 4.184 2.557 F(000) 656.0 584.0 Crystal size/mm3 0.2 × 0.05 × 0.03 0.2 × 0.05 × 0.05 2Θ range for data collection
7.84 to 140.572° 7.868 to 139.968°
Index ranges -5 ≤ h ≤ 4, -24 ≤ k ≤ 24, -16 ≤ l ≤ 16
-4 ≤ h ≤ 5, -24 ≤ k ≤ 24, -13 ≤ l ≤ 11
Reflections collected 26005 4152 Independent reflections 2505[R(int) = 0.0696] 1721[R(int) = 0.0493] Data/restraints/parameters 2505/2/189 1721/0/184 Goodness-of-fit on F2 1.044 1.045 Final R indexes [I>=2σ (I)] R1 = 0.0331, wR2 =
0.0885 R1 = 0.0480, wR2 = 0.1319
Final R indexes [all data] R1 = 0.0359, wR2 = 0.0908
R1 = 0.0544, wR2 = 0.1382
Largest diff. peak/hole / e Å-3 0.97/-0.67 0.34/-0.44
The molecular structures of BPIMN and CPIMN along with crystallographic
numbering schemes are shown in. Summary of crystal data and details of structural
refinements can be seen in Table 4.1. Selected bond distances and bond angles of
BPIMN and CPIMN are listed in Table 4.2. It is evident from the data that
geometrical parameters of the synthesized compounds are comparable to each other
Chapter 4 Results & Discussion
Page 52
and both have monoclinic structures. Smaller N-C11 bond length as compared to N-
C12 confirms the marked double bond character in the N-C11 bond. Intramolecular
hydrogen bonding is possible in both molecules between H atom of hydroxyl group
and N atom in enolic as well zwitterionic forms (Table 4.3). The molecules have non-
planar geometry as revealed by the torsional angles. The aromatic rings are not in the
form of regular hexagons due the attached moieties.
Table 4.2 Selected experimental and calculated geometric parameters for BPIMN and
CPIMN
BPIMN
Exp.
BPIMN
Calc.
Δ CPIMN
Exp.
CPIMN
Calc.
Δ
Bond distances (Å)
O1 C2 1.308(3) 1.386 0.078 1.305(3) 1.385 0.080
N1 C11 1.304(3) 1.293 -0.011 1.308(3) 1.293 -0.015
N1 C12 1.412(3) 1.416 0.004 1.411(4) 1.416 0.005
C1 C2 1.418(3) 1.399 -0.019 1.419(4) 1.400 -0.019
C1 C10 1.447(3) 1.443 -0.004 1.446(3) 1.443 -0.003
C1 C11 1.423(3) 1.461 0.038 1.419(4) 1.460 0.041
C15 X 1.900(2) 1.931 0.031 1.750(3) 1.834 0.084
Bond angles (◦)
C11 N1 C12 124.0(2) 120.8 -3.2 124.1(3) 120.8 -3.3
C2 C1 C10 119.9(2) 118.8 -1.1 119.9(2) 118.8 -1.1
C2 C1 C11 118.6(2) 115.8 -2.8 118.6(2) 115.8 -2.8
C11 C1 C10 121.6(2) 125.4 3.8 121.5(3) 125.4 3.9
O1 C2 C1 122.5(2) 118.0 -4.5 122.2(2) 118.0 -4.2
O1 C2 C3 118.1(2) 120.7 2.6 118.5(3) 120.7 2.2
N1 C11 C1 122.1(2) 125.6 3.5 122.1(3) 125.6 3.5
C13 C12 N1 122.6(2) 124.1 1.5 123.4(2) 124.0 0.6
Chapter 4 Results & Discussion
Page 53
Table 4.3 Hydrogen-bonding geometry (Å, ◦) for crystal BPIMN
D H A d(D-H)/Å d(H-A)/Å d(D-A)/Å D-H-A/°
O1 H1 N1 0.83 1.79 2.516(3) 145
N1 H1A O1 0.87 1.79 2.516(3) 140
Two kinds of non-covalent interactions [C7…Br1 = 3.311 Aº and H13…O1 =
2.482 Aº] and [Cl1…C7 = 3.323 Aº and H13…O1 = 2.481 Aº] lead to 2D
supramolecular structures to the compounds. The molecules are arranged in a manner
to produce rectangular cavities of dimension ca. 10.249 × 6.564 Aº and 15.434 ×
6.680 Aº (Fig. 4.2 A and B). Such microporous structures are expected to find
applications in separation and purification science, host–guest chemistry, gases
storage and catalysis.
Fig. 4.2(a) 2D microporous structure of the compound BPIMN mediated by
C7…Br1 = 3.311 Å and H13…O1= 2.482 Å (b) 2D microporous structure of the
compound CPIMN mediated by Cl1…C7 = 3.323 Å and H13…O1 = 2.481 Å
Chapter 4 Results & Discussion
Page 54
4.1.3 UV-Vis spectroscopy of BPIMN and CPIMN
Electronic absorption spectra of BPIMN and CPIMN depicted in Fig. 4.3 show five
bands at 226, 318, 363, 441 and 460 nm in ethanol as a solvent. In case of CPIMN,
the band at 226 nm shifts to 223 nm. The bands at 226 and 223 nm exhibit maximum
absorbance with molar extinction coefficient, ε of 3.9 × 104 and 4.2 × 104 M-1cm-1 for
BPIMN and CPIMN respectively. The intense bands at 226 nm and 223 nm are
thought to be due to aromatic rings present in the structures involving π-π* transitions.
Less intense bands around 318 nm are probably due to π-π* transitions of –C=N
group. As intramolecular hydrogen bonding is possible between -C=N and -OH
(Scheme 4.1) so there is a decrease of π-π* energy in comparison to first band. Peaks
in the range of 363–460 nm include n-π* transitions due to –C=N and –C=O moieties.
The location of these peaks at longer wavelengths can be related to the electron
releasing nature of -OH and –N=C groups.
Fig. 4.3 UV-Vis spectra of 20 μM BPIMN and CPIMN
N
Br
OH
N
Cl
OH
Scheme 4.1 Intramolecular hydrogen bonding in BPIMN and CPIMN
Chapter 4 Results & Discussion
Page 55
The existence of BPIMN and CPIMN in zwitterionic forms (Scheme 4.2) as
evidenced by crystal studies result in the separation of peaks in the range of 360–460
nm. Ortho-hydroxy Schiff bases exist in zwitterionic forms as reported in
literature [1].
O
N
Br
HO
N
Br
O
N
Br
H
Scheme 4.2 Tautomeric forms of BPIMN
Normally, such zwitterionic forms are produced by a proton transfer from ortho-
hydroxy group to the nitrogen lone pair in the phenol-imine form through a six
membered transition state. Neutralization of the negative and positive charges in the
zwitterionic forms results in the formation of neutral keto-amine forms.
Fig. 4.4 Electronic absorption spectra of 20 μM BPIMN (A) and 20 μM CPIMN (B)
recorded in different pH media
Fig. 4.4 shows the electronic absorption spectra of 20 μM BPIMN and CPIMN in
acidic and basic conditions. The results obtained from the spectra recorded in
different pH media can help in the determination of pKa of the -OH group present in
the molecules. BPIMN shows an intense doublet at 231–247 nm in the pH range
4–12, while the intense band of CPIMN moves to much shorter wavelengths with the
Chapter 4 Results & Discussion
Page 56
appearance of less intense band at 245 nm. This spectral behavior can be explained on
the basis of electronegativity values of chlorine and bromine atoms. Chlorine being
more electronegative deactivates benzene ring so imparts hypsochromic shift to the
first band along with hypochromic effect on the second band, while bromine being
less electronegative imparts bathochromic and hyperchromic shifts to the second
band, hence, both bands of BPIMN appear as doublet.
In strongly acidic medium, hypsochromic effect occurs in the most intense
band due to possible protonation of -OH and –C=N groups. These moieties can get
positive charges after proton capture and hence, cause to deactivate the aromatic ring.
This effect can also be confirmed by the disappearance of band at 318 nm in CPIMN
spectra due to –C=N and hypochromic effect in the third band. In pH 6.0 and 8.0, the
probability of protonation reduces, therefore; both these molecules show well
resolved bands at 440 and 460 nm. Further increase in pH results in more blue shift of
the intense band toward the UV region. With the rise in pH of the medium, the peak at
245 nm (due to π-π* transitions) intensifies, the band at 318 nm decays and a new
intense peak corresponding to n-π* transitions appears at 397 nm. These peculiar
spectral characteristics can be explained by the deprotonation of -C-OH group which
gets converted to –C=O in alkaline medium and gives new peaks at 245 and 397 nm.
Moreover, BPIMN shows three isosbestic points at 239, 299 and 337 nm, indicating
its existence in different tautomeric forms. pKa of BPIMN and CPIMN with values
9.2 and 6.7 were calculated from the plots of absorbance versus pH.
4.1.4 Cyclic voltammetry of BPIMN and CPIMN The cyclic voltammetry of 1 mM solutions of BPIMN and CPIMN was performed at
a scan rate of 100 mV s-1 in medium of pH 7.0 using GCE in the potential range of
-1.5 to +1.5 V. CPIMN registered one oxidation peak at +0.73 V and three reduction
peaks at -0.92, -1.24 and -1.48 V, while BPIMN oxidized at +0.71 V.
Voltammograms of both these compounds were recorded in different potential
windows in order to ensure the dependent or independent nature of the peaks as
shown in Fig. 4.5. Cathodic and anodic peaks of CPIMN were found independent of
each other and hence, generated by separate reducible and oxidizable moieties of the
compound.
Chapter 4 Results & Discussion
Page 57
-2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0 1.5-32
-24
-16
-8
0
8
16
24
32
I p /µA
E / V vs. Ag/AgCl
cb
a
-1.4 -0.7 0.0 0.7 1.4
-20
-10
0
10
20
Curre
nt (µ
A)
Potential (V)
BPIMN CPIMN
Fig. 4.5 (A)The CVs of CPIMN recorded at 100mVs-1 starting from 0 V (a) -0.6 V
(b) and -1.4 V (c) with 1mM concentration of analyte in 50% ethanol and 50% BR
buffer of pH-7, (B) Voltammogram of BPIMN and CPIMN obtained under same
conditions
A remarkable feature of these voltammograms is that the Schiff base having
chlorine atom has three reduction signals while bromine containing Schiff base has no
prominent peak in the cathodic region. This clear voltammetric distinction can be
related to the more electronegative character of chloro group in encouraging the N of
azomethine to donate its lone pair of electrons toward the phenyl ring thus rendering
the C=N group to reduce in the negative potential domain of GCE. In case of
comparatively less electron withdrawing ability of bromo group, the azomethine
functionality may remain electron rich hence no signals of reduction appear at the
GCE.
In order to investigate the adsorption of CPIMN at the surface of working
electrode, consecutive scans were recorded without cleaning the electrode surface. A
reasonably small decrease in peak current indicated slight adsorption of the analyte.
The plot between log of scan rate and log of peak current evidenced the redox
processes to be mainly controlled by diffusion and slightly by adsorption. Moreover,
there is indication of the appearance of another oxidation peak at higher positive
potential in the second scan which gets more prominent with further scans. Multiple
SW voltammograms (see Section ‘Square wave voltammetry’) offered more clear
evidence of the appearance of this second peak due to more sensitivity.
Chapter 4 Results & Discussion
Page 58
4.1.5 Differential pulse voltammetry of BPIMN and CPIMN
DPVs of 1 mM BPIMN and CPIMN were obtained in a broad pH range (2–12) in
order to ensure the involvement of protons during electron transfer processes. The
width at half peak height (W/2) of all oxidation and reduction peaks were determined
at different pH values. In case of CPIMN, the oxidation signal was broad in the pH
range 3–8 due to two overlapping peaks, which split into two well defined separate
peaks of 88 mV half peak heights at pH 9.0 as shown in Fig. 4.6. These W1/2 values
are very close to the theoretical value (90.4 mV) of one electron transfer process
according to the criteria reported in literature [2, 3]. BPIMN showed almost the same
behavior with peak splitting occurring at pH-8.0. Thus, in the electrochemical
processes of these Schiff bases, one electron is involved in all signals as confirmed
from their W1/2 values. The peak potentials shifted with change in pH of the medium.
Epa1 as a function of pH with a slope of 55 mV per unit pH (closer to 59 mV
theoretical value) demonstrated the involvement of one electron and one proton in the
first oxidation step [4]. The participation of the same number of electrons and protons
in the second oxidation process was shown by 57 mV per unit pH slope of Epa2 versus
pH plot [5].
0.0 0.4 0.8 1.20.01.53.0
E / V vs. Ag/AgCl
pH-30.01.83.6
pH-50.00.51.0
pH-7.4I pa /
µA
0.00.81.6
pH-90.01.93.8
pH-100.02.55.0
pH-110.02.34.6
pH-12
Fig. 4.6 Differential pulse voltammograms for oxidation of CPIMN at 10mVs-1 in
50% ethanol and 50% BR buffers of wide pH range used as supporting electrolyte
Chapter 4 Results & Discussion
Page 59
0.0 0.5 1.0 1.50.02.55.0
0.01.53.00.00.91.80.00.51.00.00.20.4
-0.0450.0000.045
Curre
nt (µ
A)
Potential (V)
Scan 1
Scan 2
Scan 3
Scan 4
Scan 5
Scan 6
Fig. 4.7 Multiple scan SWVs of 1mM CPIMN recorded at 100 mV s-1in 50% ethanol
and 50% BR buffer of pH-7, used as supporting electrolyte
4.1.6 Square wave voltammetry of BPIMN and CPIMN
Square wave voltammograms of 1 mM CPIMN were obtained in the pH range of 3–
12 at a scan rate of 100 mV s-1. Like CV, SWV showed irreversible nature of
oxidation and reduction processes of CPIMN. Five successive potential scans
recorded without cleaning the working electrode surface are displayed in Fig. 4.7. An
examination of the voltammograms shows that a new peak appears at more positive
potential in the second scan, which becomes more prominent in the third and fourth
scans, while height of the original peak diminishes in later scans. On the basis of these
electrochemical findings, redox mechanisms of BPIMN and CPIMN were proposed
as discussed in subsequent section.
4.1.7 Redox mechanism of BPIMN and CPIMN
1-((4-Chlorophenylimino) methyl) naphthalen-2-ol exists in enolic and zwitterionic
forms as confirmed by FTIR studies. The oxidation of CPIMN gave a broad oxidation
peak in acidic medium, which split into two distinct peaks in basic medium. This is
because when the medium is acidic, proton combines with negatively charged oxygen
of zwitterionic form and thus the enolic form of the compound exists in acidic
Chapter 4 Results & Discussion
Page 60
medium. The half peak width values of both peaks indicated the transfer of one
electron each. The oxidation of hydroxyl part occurs by the loss of one electron and
one proton below pH-9. This peak becomes pH independent above pH-9 and hence,
abstraction of one electron occurs without involvement of proton in highly alkaline
conditions.
N
OH
Cl
-H+
N
O
Cl
N
O
Cl
N
O
Cl
pH < 9
pH > 91e-
1e-
Scheme 4.3 Proposed mechanism for oxidation of CPIMN for peak 1a
The appearance of another oxidation peak at pH ≥ 9 can be related to the
presence of zwitterionic form of CPIMN. The oxidation corresponding to this peak
occurs by the loss of one electron and one proton as witnessed by 88 mV half peak
width and 57 mV pH-1 slope of Ep versus pH plot. The appearance of an extra peak in
the SW voltammogram of CPIMN at more positive potential in 2nd and higher scans is
due to the shifting of equilibrium toward enolic form. The proposed mechanism of
CPIMN oxidation is presented in Scheme 4.3 and Scheme 4.4.
Chapter 4 Results & Discussion
Page 61
HN
O
Cl
-H+N
O
Cl
N
O
Cl
N
O
Cl
pH > 7
pH > 111e-
1e-
Scheme 4.4 Proposed mechanism for oxidation of CPIMN for peak 2a
4.1.8 Computational study of BPIMN and CPIMN
The DFT optimized structures of the investigated compounds were obtained using
3-21G basis set [6]. Computational study of BPIMN and CPIMN was done to
calculate the Mullikan charges, optimization energies, dipole moments, spin
multiplicity and point groups of all three forms (see Table 4.4). The results reveal that
keto and zwitterionic forms of both compounds are more stable than their enolic
forms. From Mullikan Charge distribution values, it is obvious that oxygen atom of -
OH group has the most negative charge which justify our attribution of electron
abstraction from the same electropores.
Table 4.4 Selected parameters of BPIMN and CPIMN calculated by RB3LYP/3-21G
set using DFT method in GUSSIAN03W package
BPIMN CPIMN enol form keto form ZI form enol form keto form ZI form Charge 0 0 0 0 0 0 Spin Singlet Singlet Singlet Singlet Singlet Singlet T. Energy (a.u.)
-3342.49 -3342.49 -3342.50 -1238.68 -1238.68 -1238.68
Dipole Moment (D)
4.736 4.205 4.077 5.653 4.434 4.422
EHOMO(a.u.) -0.20541 -0.19498 -0.20082 -0.20912 -0.20418 -0.20395 ELUMO(a.u.) -0.06855 -0.06887 -0.07414 -0.07098 -0.07751 -0.07659 Point Group C1 C1 C1 C1 C1 C1
Chapter 4 Results & Discussion
Page 62
4.1.9 Inhibition of ALP
Alkaline phosphatases (ALP) occur widely in nature and found in many organisms. In
humans, it is produced in liver, bone, and placenta and normally present in high
concentration in bile and growing bone. It is released into the blood during injury and
during normal activities such as bone growth and pregnancy. The enzyme is termed
alkaline phosphatase because it works under alkaline conditions, as opposed to acidic
phosphatase [7].With few exceptions, ALP are homodimeric enzymes and each
catalytic site contains three metal ions, i.e., two Zn and one Mg, necessary for
enzymatic activity [8]. Increased level of ALP is indicative of hepatobiliary disease,
osteoblastic activity and tumor formation [9]. Schiff bases are inhibitors of ALP and
exert their inhibition effect by blocking the active sites of the enzyme. The results of
our experiments reveal that chloro-substituted Schiff base inhibit the enzyme more
effectively as compared to bromo-substituted Schiff base. ALP was inhibited in a
concentration-dependent manner as shown in Fig. 4.8.
0.00 0.05 0.10 0.15 0.20 0.250
9
18
27
36
45
Inhi
bitio
n of
ALP
(%)
Concentrtation of Inhibitors (mM)
BPIMN CPIMN
Fig. 4.8 Concentration dependent inhibition of alkaline phosphatase (ALP) by
BPIMN and CPIMN
Chapter 4 Results & Discussion
Page 63
4.2 Synthesis, redox behavior and biological applications of
1-((2,4-dimethylphenylimino) methyl)naphthalene-2-ol
(HL) and its metallic derivatives
4.2.1 Structural Confirmation of HL and its metallic complexes FTIR spectrum of HL showed a strong absorption band at 1615 cm-1 due to
azomethine group which indicated the formation of Schiff base (HL). Moreover,
appearance of resonance signals at 9.27 ppm and 172.1 ppm in 1H and 13C NMR
spectra respectively confirmed the formation of HL. Absence of any resonance peak
related to primary amine revealed the purity of the synthesized product. In case of
metal complexes, the absorption bands related to azomethine group in FTIR spectra
appearing at lower wavelengths in the range of 1591-1600 cm-1 suggested the
coordination of azomethine nitrogen with metal atoms. Data obtained from elemental
analysis and thermo gravimetric analysis of (L2VO) and (L2Co) indicated the
formation of complexes (L2VO) and (L2Co). For tin and zinc complexes, downfield
shifting of azomethine group (HC=N) resonance signals was observed and
disappearance of proton resonance signal of hydroxyl group confirmed the bonding of
ligand anion with the metallic cationic parts. Moreover, Sn-C coupling constant for
dibutyltin complex was calculated from the satellite peaks appearing at 23.9 and 29.6
ppm in 13C NMR spectrum.
4.2.2 Crystal structures description of HL
Fig. 4.9 Ellipsoid diagram of Schiff base HL
The molecular structure of HL along with crystallographic numbering scheme is
shown in Fig. 4.9. Summary of crystal data and details of structural refinements can
be seen in Table 4.5. Selected bond distances and bond angles of HL are listed in
Chapter 4 Results & Discussion
Page 64
Table 4.6. It is evident from the data that the synthesized Schiff base has a monoclinic
structure. Double bond between C and N can be confirmed by the shorter N-C11 bond
length as compared to N-C12 bond length.
Table 4.5 Crystal data and details of the structure refinement for HL
Crystal Data Formula C19H17NO α, β, γ [Deg] 90, 96.938, 90 Formula Weight 275.34 F(000) 292 Crystal System Monoclinic Crystal Size [mm] 0.00×0.00×0.00 Space group P21 Temperature (K) 0 a, b, c [A°] 5.9821, 9.0445, 13.5252 Theta Min-Max [Deg] 0.0, 0.0
Position of H1 represents that HL exists in zwitterionic form as well as in
enolic form (Fig. 4.9). Intramolecular hydrogen bonding is present between H atom of
hydroxyl group and N atom in enolic as well zwitterionic forms (Table 4.7). Tortional
angles revealed that the molecular geometry is non-planar. The aromatic rings are not
in the form of regular hexagons due the attached moieties.
Table 4.6 Selected experimental and calculated geometric parameters for HL
Bond distances (Å) Exp. Calc. ∆ O1 C1 1.292 1.310 0.018 N1 C11 1.306 1.315 0.009 N1 C12 1.408 1.402 -0.006 N1 H1 1.020 1.090 0.070 C13 C18 1.488 1.510 0.022 C15 C19 1.509 1.514 0.005 C18 H18A 0.960 1.070 0.110 C19 H19C 0.960 1.070 0.110 Bond angles (◦) Exp. Calc. ∆ C11 N1 C12 126.0 121.8 -4.2 C11 N1 H1 112.0 109.4 -2.6 O1 C1 C2 119.6 119.4 -0.2 N1 C11 C10 122.8 124.1 1.3 N1 C12 C17 122.8 122.4 -0.4 Torsion angles (◦) Exp. Calc. ∆ C12 N1 C11 C10 179.2 178.6 -0.6 C11 N1 C12 C13 170.9 177.0 6.1 N1 C12 C13 C18 -1.2 -1.2 0.0
Chapter 4 Results & Discussion
Page 65
Table 4.7 Hydrogen-bonding geometry (Å, ◦) for crystal DMN
D H A d(D-H)/Å d(H-A)/Å d(D-A)/Å D-H-A/° N1 H1 O1 1.0200 1.6700 2.5413 141.00
4.2.3 Electronic absorption spectroscopy of HL and complexes
260 325 390 455 520 5850.0
0.2
0.4
0.6
0.8
1.0
Ab
sorb
ance
Wavelength (nm)
HL (1) (2) (3) (4)
Fig. 4.10 UV-Vis spectra of 20 μM HL, (L2VO), (L2Sn), (L2Zn) and (L2Co)
UV-visible spectra of HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) obtained in
ethanol having 8 % DMSO can be seen in Fig. 4.10. HL shows four bands at 318,
350, 442 and 461 nm. First band at 318 nm is probably due to π-π* transitions of
–C=N group. Second band at 350 nm can be related to π-π* transitions of –C=O
group. This effect is consistent with the crystal studies which show the existence of
HL in keto form. Peaks at 442 and 461 nm correspond to n- π* transitions due to
–C=N and –C=O moieties. In case of metal complexes a new band appears within the
range of 260-270 nm which may be due to π-π* transitions of aromatic rings [10].
Appearance of this band can be related to the inductive effect raised by metal ions
giving bathochromic shift in the absorption band of aromatic rings. Hyperchromic
effect in the spectrum of (L2Sn) can be explained by the inductive effect of butyl
groups attached to tin metal atom, hence activating the chromophores. Molar
extinction coefficients, ε of HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) are 2.78 × 104,
3.21 × 104, 2.08 × 104, 1.91 × 103 and 1.29 × 104 M-1cm-1 respectively.
Chapter 4 Results & Discussion
Page 66
300 350 400 450 500 5500.0
0.1
0.2
0.3
0.4
0.5
Abs
orba
nce
Wavelength (nm)
pH-2 pH-4 pH-6 pH-7 pH-8 pH-10 pH-12
250 300 350 400 450 500 5500.0
0.2
0.4
0.6
0.8
1.0
Abs
orba
nce
Wavelength (nm)
pH-2.0 pH-4.0 pH-6.0 pH-7.0 pH-8.0 pH-10.0 pH-12.0
A B
Fig. 4.11 UV-Vis spectra of 20 μM HL (A) and 30 μM (L2Sn) (B) in different
media of pH range 2-12
Fig. 4.11(A) represents the absorption spectra of HL in different pH media.
HL has shown an intense doublet band at 321 nm. Increase in pH of the medium
results in hypochromic effect and a new band appeared at 414-461 nm which further
shifted to 396 nm at pH > 9. Appearance of new band can be related to the
deprotonation of HL. An isosbestic point at 371 nm also appeared in the spectra
which confirmed the transformation of HL into different tautomeric forms. Tin
complex has shown similar photometric response with two isosbestic points at 298
and 374 nm confirming the existence of its different mesomeric forms (Fig. 4.11(B)).
No significant change was observed in photometric studies of complexes (L2VO),
(L2Zn) and (L2Co) in different pH media.
4.2.4 Cyclic voltammetry of HL and complexes
Cyclic voltammetric behavior of I mM solution of HL and its complexes was
examined at a scan rate of 100 mV s-1 in different pH media using GCE as working
electrode. In medium buffered at pH 7.0, HL registered two oxidation signals at +0.76
V and +1.01 V in the forward scan. The appearance of no peak in reverse scan
ensured the irreversibility of both the oxidation processes.
An observation of Fig. 4.12 (A) reveals that the redox nature of HL is strongly
pH dependent. In basic medium (pH-10.0) the two oxidation peaks overlap into a
single broader peak at +0.83 V. While no significant shift in reduction peak is
Chapter 4 Results & Discussion
Page 67
observable with change in pH. However, in acidic medium a new reduction peak at -
1.13V can be seen.
-1.0 -0.5 0.0 0.5 1.0 1.5-100
-50
0
50
100
150
Cu
rren
t (µA
)
Potential (V)
pH-4.0 pH-7.0 pH-10.0
-1.0 -0.5 0.0 0.5 1.0-150
-100
-50
0
50
100
150
Curr
ent (
µA)
Potential (V)
HL 1 2 3 4
(A) (B)
Fig. 4.12 Cyclic Voltammograms of (A) HL in a buffer of pH 4.0, 7.0 and 10.0 (B)
HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) in a buffer of pH-7 at 100mVs-1
Fig. 4.12(B) depicts the cyclic voltammograms of HL and its complexes in a medium
buffered at pH 7.0. In voltammograms of all metal complexes, the two oxidation
peaks overlap and get converted into one broader peak at +0.76 V, +0.79 V, +0.74 V
and +0.79 V for (L2VO), (L2Sn), (L2Zn) and (L2Co) respectively. In the negative
potential window, instead of one sharp cathodic signal, two overlapping peaks appear
for (L2VO), (L2Sn) and (L2Co). Compound (L2Zn) exhibited a sharp reduction peak at
-0.44 V and the reduced product oxidized irreversibly at -0.24 V (Fig. 4.13). The
regular shifting of peaks with the change in pH is a clear evidence of the pH
dependent nature of all these redox processes (Fig. 4.12 & Fig. 4.13).
Chapter 4 Results & Discussion
Page 68
-1.0 -0.5 0.0 0.5 1.0 1.5
-40
-20
0
20
40
60
Cur
rent
(µA
)
Potential (V)
pH-4.0 pH-7.0 pH-10.0
Fig. 4.13 Cyclic Voltammograms of (L2Zn) at 100mVs-1 in a buffer of pH 4.0,
7.0 and 10.0
The effect of scan rate on the cyclic voltammetric response was investigated
for all compounds in a medium of pH-7.0. By using Randles-Sevcik, equation the
values of diffusion coefficients were calculated from the slopes of the plots between Ip
and υ1/2. The straight lines of the plots suggested the diffusion controlled behavior of
the redox processes [5]. CVs of different concentrations of the analytes were
examined and used to calculate heterogeneous electron transfer rate constant ks using
the following equation for an irreversible process [11].
𝐼𝑝 = 𝑛𝑛𝑛𝑛𝑘𝑠 (4.1)
0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0
3.0
4.5
6.0
7.5
9.0
Peak
curr
ent (
Ip)
Concentration (M)
R2= 0.9991
Fig. 4.14 Plot of Ip vs concentration for (L2Sn) in a buffer of pH 7.0
Chapter 4 Results & Discussion
Page 69
Where n represents number of electrons calculated from DPV studies and A the
surface area of GC working electrode. The value of ks was calculated from the slope
of plotting Ip versus concentration plot (Fig. 4.14). The values of Do and Ks for HL,
(L2VO), (L2Sn), (L2Zn) and (L2Co) are listed in Table 4.8.
4.2.5 Square wave voltammetry of HL and complexes
To ensure the reversible/irreversible nature of the redox processes of HL and its
complexes, SWVs was obtained. The same direction of forward and backward current
components of all the redox processes proved the irreversible nature. Due to its
sensitive nature, SWV was used for the estimation of limit of detection (LOD) and
limit of quantification (LOQ) by observing the effect of concentration on peak current
values (Fig. 4.15) using equations 4.2 and 4.3.
LOD = 3𝑆 𝑚� (4.2)
LOQ = 10𝑆 𝑚� (4.3)
Where S represents standard deviation and m the slope of linear segments. LOD and
LOQ values are given in Table 4.8.
0.00 0.25 0.50 0.75 1.000
2
4
6
8
0 150 300 450 600
Concentration (µM)
N = 12 Mean = 2.85467E-6SD = 2.63667E-6
Mean2.85467E-6
+1 SD5.49133E-6
-1 SD2.18E-7
Cur
rent
(µA
)
Potential (V)
0.10µM 2.60µM 7.80µM 15.6µM 62.5µM 125µM 167µM 250µM 333µM 500µM 571µM 667µM
Fig. 4.15 Square wave voltammograms of different concentrations of (L2Sn) obtained
at a scan rate 100 mV/s in pH 7.0, inset shows Ip as a function of concentration
Chapter 4 Results & Discussion
Page 70
Table 4.8 Diffusion coefficient, rate constant, LOD and LOQ values for HL,
(L2VO), (L2Sn), (L2Zn) and (L2Co)
Compounds Do × 106 (cm2.sec-1) ks × 105 (cm.sec-1) LOD (µM) LOQ (µM)
HL 7.514 4.750 7.21 24.03
(L2VO) 0.040 0.021 17.6 58.60
(L2Sn) 0.144 0.260 3.48 11.60
(L2Zn) 0.507 0.390 0.10 0.33
(L2Co) 0.620 0.563 5.63 18.77
4.2.6 Differential Pulse Voltammetry of HL and complexes
As the redox processes of HL and its complexes were found pH dependent, so
differential pulse voltammetry was carried out for getting more reliable information
about the involvement of protons during electron transfer processes. The number of
electrons was calculated from half peak width values for irreversible processes using
the following equation [12].
𝑊1/2 = 3.52 𝑅𝑅𝛼𝛼𝛼
(4.4)
Fig. 4.16 shows the DPVs corresponding to the oxidation of HL and its cobalt
complex. In strongly acidic conditions, two well defined pH dependent oxidation
peaks were observed. Intensity of 1st oxidation peak started decreasing with the
increase in pH and finally this signal disappeared at pH higher than 4.0, while 2nd
oxidation peak intensified and became broader with the increase in pH. In all other
metal complexes, the same two oxidation signals were observed but peak 3a was not
present in all cases. Vanadium complex of HL showed a new oxidation pH dependent
signal at low potential value of +0.26 V in strongly acidic media. Shifting of peak 2a
with pH was significant up to pH-5.0, after that it became less pH dependent, hence
pH-5.0 was selected for the calculation of pKa values of the compounds. Shifting of
peak potential towards more negative potential explains the facile oxidation of
analytes with increase in pH. A new pH independent and broad oxidation peak in the
range of pH 4.0-9.0 suggested another one electron oxidation step of HL without
Chapter 4 Results & Discussion
Page 71
involvement of proton. Number of protons involved in each process was calculated
using the following equation [12]. 𝑑𝐸𝑝
𝑑𝑑𝑑� = 0.059𝑃 𝛼𝑛� (4.5)
(2a)
0.0 0.5 1.0 1.502402460.0
2.55.00.03.57.00.02.55.00.03.26.40.03.57.00.09.5
19.0
Cur
rent
(µA
)
Potential (V)
pH-2.0
pH-3.0
pH-4.0
pH-5.0
pH-6.0
pH-7.4
pH-9.0
pH-11.0
(1a)
0.0 0.5 1.0 1.50.02.14.20.01.73.40.00.80.01.20.01.10.01.20.00.61.20.00.51.00.00.80.00.91.80.01.42.8
Cur
rent
(µA
)
Potential (V)
pH-2.0
pH-3.0
pH-4.0
pH-5.0
pH-6.0
pH-7.0
pH-8.0
pH-9.0
pH-10.0
pH-11.0
pH-12.0 (B)
(A)
(1a) (2a) (3a)
Fig. 4.16 Differential pulse voltammograms recorded in different pH media at
5 mVs-1 showing oxidation of HL (A) and (L2Co) (B)
2 4 6 8 100.5
0.6
0.7
0.8
0.9
1.0
1.1
Peak 1a Peak 2a Peak 2a Peak 3a
Ep
pH
Slope= -0.045
Slope= -0.1205
Slope= -0.017
Slope= -0.0023(A)
3.0 4.5 6.0 7.5 9.0 10.50.4
0.5
0.6
0.7
0.8
0.9
1.0
Peak 1a Peak 2a Peak 3a
Ep
pH
Slope= -0.019
Slope= -0.0606
Slope= -0.00
(B)
Fig. 4.17 Plots of oxidation potentials versus pH for HL (A) and (L2Co) (B)
Chapter 4 Results & Discussion
Page 72
Slopes of plots of Ep vs pH were used to calculate number of protons involved in the
oxidation processes of HL and its complexes (Fig. 4.17).
(2c)
-1.2 -0.8 -0.4 0.0-1.17-0.78-0.390.00
-1.17-0.78-0.390.00
-1.17-0.78-0.390.00
-1.11-0.74-0.370.00
-1.17-0.78-0.390.00
-1.17-0.78-0.390.00
-1.41-0.94-0.470.00
-1.44-0.96-0.480.00
Curre
nt (µ
A)
Potential (V)
pH-2.0
pH-3.0
pH-4.0
pH-5.0
pH-6.0
pH-7.0
pH-9.0
pH-11.0
-1.2 -0.8 -0.4 0.0
-1.92-1.28-0.640.00
-1.74-1.16-0.580.00
-1.59-1.06-0.530.00
-1.44-0.96-0.480.00
-1.35-0.90-0.450.00
Cur
rent
(µA
)
Potential (V)
pH-2.0
pH-3.0
pH-4.0
pH-5.0
pH-6.0
(1c)
(A) (B)
(1c)(2c)
Fig. 4.18 Differential pulse voltammograms recorded in different pH media at
5 mVs-1 showing reduction of HL (A) and (L2Sn) (B)
Chapter 4 Results & Discussion
Page 73
-1.0 -0.5 0.0-3.2-1.60.0-2.10.0-4.6
-2.30.0
-2.60.0
-5.0-2.50.0
-5.8-2.90.0-6.2
-3.10.0
-6.4-3.20.0
-3.30.0
-2.90.0
-2.30.0
Cur
rent
(µA
)
Potential (V)
pH-2.0
pH-3.0
pH-4.0 pH-5.0
pH-6.0
pH-7.0
pH-8.0
pH-9.0
pH-10.0
pH-11.0
pH-12.0
-1.0 -0.5 0.0-2.46-1.64-0.820.00
-1.71-1.14-0.570.00
-1.62-1.08-0.540.00
-0.99-0.66-0.330.00
-1.41-0.94-0.470.00
-1.17-0.78-0.390.00
-1.35-0.90-0.450.00
-1.41-0.94-0.470.00
-1.71-1.14-0.570.00
-1.44-0.96-0.480.00
-1.29-0.86-0.430.00
Cur
rent
(µA
)
Potential (V)
pH-2.0
pH-3.0
pH-4.0 pH-5.0
pH-6.0
pH-7.0
pH-8.0
pH-9.0
pH-10.0
pH-11.0
pH-12.0(A) (B)
(1c)(2c) (2c) (1c)
Fig. 4.19 Differential pulse voltammograms recorded in different pH media at
5 mVs-1 showing reduction of (L2VO) (A) and (L2Zn) (B)
Fig. 4.18 and Fig. 4.19 show DP voltammograms of reduction processes of HL
and complexes. In negative potential window, HL registered two pH dependent
signals in the acidic pH range while no reduction signals in basic media. Shifting of
peak potential to more negative values can be attributed to the difficulty of reduction
process due to de-protonation of the electroactive moiety of HL [5]. Complex (L2VO)
registered two reduction peaks (1c and 2c) both of which are independent of pH. Peak
1c remained the same in the whole pH range while intensity of 2nd peak decreased
with increase in pH and disappeared in the basic region. Tin complex (L2Sn) also
exhibited two step reductions in which first process was just an electron transfer while
the second process involved the accompaniment of proton along with electron
transfer. Cobalt complex (L2Co) registered only one peak at +4.9 V in acidic media
which was pH independent and no peak in basic pH range. Zinc complex (L2Zn)
exhibited a reduction peak at a very low potential value of 0.0 V in strong acidic pH-
2.0 which showed significant shifting by increasing pH of the medium. The reduced
Chapter 4 Results & Discussion
Page 74
product of zinc complex oxidized in an irreversible manner as its oxidation signal also
appeared in the reverse scan (Fig. 4.20).
-1.2 -0.9 -0.6 -0.3 0.0 0.3
-6
-4
-2
0
2
Curr
ent (
µA)
Potential (V)
Forward Reverse
Fig. 4.20 Differential pulse voltammograms recorded at 5 mVs-1 showing forward and
reverse scans of reduction process of (L2Zn) at pH-6.0
1.0 1.2 1.4 1.6 1.8 2.010121416182022
2 4 6 8 10 12-0.7
-0.6
-0.5
-0.4
-0.3
-0.2
-0.1
0.0
peak 1c peak 1c Peak 2c
Ep
pH
Slope= -0.0711
Slope= -0.0456
Slope= -0.0080
2 3 4 5 6-1.2
-1.1
-1.0
-0.9
-0.8
-0.7
-0.6
Peak 1c Peak 2c
Ep
pH
Slope= -0.0954
Slope= -0.0727
B
A
(A) (B)
2 4 6 8 10 12-1.3
-1.2
-1.1
-1.0
-0.9
-0.8
-0.7
Peak 1c Peak 2c
Ep
pH
Slope= -0.0014
Slope= -0.0310
(C)
Fig. 4.21 Plots of Epc versus pH for HL (A), (L2Zn) (B) and (L2Sn) (C)
Chapter 4 Results & Discussion
Page 75
These finding are in agreement with the cyclic voltammetric investigations.
Another reduction signal appeared at -0.67 V in acidic range which diminished with
increase in pH. Number of electrons and protons involved in all these reduction
processes (Fig. 4.21) were also calculated and used in finding the mechanistic path
ways of these processes (mechanism is given in section 4.2.7).
4.2.7 Redox mechanism of HL and complexes
HCNH
H3C
CH3
O
-1e-
-1H+ HCN
H3C
CH3
O
HCN
H3C
CH3
OH
HCN
H3C
CH3
O-GCE
HCN
H3C
CH3
O CHN
CH3
CH3
O
(Dimer)
Scheme 4.5 Proposed oxidation mechanism of HL for peak 1a (according to Figs.
4.12, 4.16 & 4.17)
Chapter 4 Results & Discussion
Page 76
HCNH
H3C
CH3
O
-2e-
-2H+N
H3C
CH3
O
1-((2,4-dimethyl phenylimino) methylene)naphthalen-2(1H)-one
HCNH
H3C
CH3
O
Scheme 4.6 Proposed oxidation mechanism of HL for peak 2a (according to
Figs. 4.12, 4.16 & 4.17)
The number of electrons and protons involved in redox process were calculated from
half peak width and slope values of Ep versus pH plot. The HL exhibited three
oxidation peaks in positive potential window. Peak 1a was found to involve loss of
one electron and one proton. On the basis of voltammetric results it is suggested that
one electron and one proton has been lost from hydroxyl group at ortho position
which then converted to free radical and gets stabilized by dimerization or by
adsorption at electrode surface. The one possibility is, it can form five membered ring
with N atom. Secondly, it is converted to keto form with free radical resonating in the
whole benzene ring. It can also get adsorb on the GCE surface or dimerize as shown
in Scheme 4.5. Similarly, Peak 2a involves the loss of two electrons and two protons.
This peak is assigned to azomethine group of HL present in keto form as shown in
Scheme 4.6. On the basis of DPV results, the formation of 1-((2,4-dimethyl
phenylimino)methylene)naphthalen-2(1H)-one is proposed.
Reduction of HL exhibited two separate peaks. W1/2 value of peak 1c suggested loss
of three protons along with four electrons per two molecules. This peak was related to
the conversion of azomethine group into -C-N bond (Scheme 4.7). Reduction peak 2c
involves gain of one electron and one proton is assigned to the -C=O moiety (keto
form) which got converted to oxygen free radical and get stabilized by adsorption on
the electrode surface.
Chapter 4 Results & Discussion
Page 77
HCN
H3C
CH3
OH
H2CNH
H3C
CH3
O-GCE
HCNH
H3C
CH3
OH
4e- 3H+
H2CNH
H3C
CH3
OH
2
H2CNH
H3C
CH3
O
1e-
1H+
Peak 2C
Scheme 4.7 Proposed reduction mechanism for peak 1c and 2c (according to Figs.
4.12, 4.17, 4.18 & 4.21)
Oxidation of oxovanadium complex occurs in three different steps. DPV results show
that peak 1a involves loss of two electrons along with two protons. Peak 2a remains
independent in the whole pH range showing no involvement of proton while loss of
three protons and four electrons per two molecules for peak 3a was calculated from
DPV results. Organotin complex of HL exhibits oxidation in two steps in the similar
manner. pH independent peak 1a involves the loss of one electron while peak 2a
involves loss of one proton along with one electron. Cobalt and zinc complexes also
oxidize in a similar manner but in case of zinc complex pH independent peak 2a
Chapter 4 Results & Discussion
Page 78
appears only in the basic region. In case of HL, the two peaks are assigned to the
oxidizable moieties –OH and –C=N present in the molecule.
4.2.8 Anti-diabetic Activity of HL and complexes
The anti-diabetic activity of HL and its metal complexes were investigated. Mice
treated with ligand HL showed significant decrease in blood glucose level as
compared to the negative control group but showed increase as compared to positive
control group (p<0.001; q= 8.329 and p<0.001; q= 7.722; Fig. 4.22A). Fig.4.23 A
shows that serum triglycerides concentration decreased significantly as compared to
negative and positive control groups (p<0.001; q= 8.717; and p<0.001; q= 15.415;),
with no significant effects on serum cholesterol concentration (Fig. 4.23B). An
observation of Fig. 4.22B reveals that animals treated with oxovanadium complex
(L2VO) significantly increase the blood glucose concentration as compared to positive
control group (p<0.001; q= 12.907). Fig.4.23A shows that serum triglycerides
concentration gets decreased as compared to both positive and negative control
groups (p<0.001;q= 13.336 and p<0.001; q= 6.638).
Chapter 4 Results & Discussion
Page 79
-1 0 1 2 3 4 5 6 70
200
400
600
800Negative ControlPositive ControlHL
**
Time(hr)G
luco
se(C
once
ntra
tion
mg/
dl)
(A)
-1 0 1 2 3 4 5 6 70
200
400
600
800
Negative Control
(1)
Positive Control
**
Time (hr)
Glu
cose
(con
cent
rati
on m
g/dl
)
-1 0 1 2 3 4 5 6 70
200
400
600
800
C
(2)
Negative Control
Positive Control
**
Time (hr)
Glu
cose
(con
cent
rati
on m
g/dl
)
-1 0 1 2 3 4 5 6 70
200
400
600
800
D
(3)
Negative Control
Positive Control
**
Time (hr)
Glu
cose
(con
cent
rati
on m
g/dl
)
-1 0 1 2 3 4 5 6 70
200
400
600
800
E
Negative Control
Positive Control
(4)
**
Time (hr)
Glu
cose
(con
cent
rati
on m
g/dl
)
A
(B) (C)
(D) (E)
Fig. 4.22 Plasma glucose in mice treated with HL, (L2VO) (1), (L2Sn) (2), (L2Zn) (3)
and (L2Co) (4) compound groups
Except oxovanadium complex (L2VO), mice treated with compounds HL, (L2Sn),
(L2Zn), and (L2Co) treated group demonstrated a significant decrease in blood glucose
concentration as compared to the negative control group (p<0.001; q= 18.037). Data
are presented as mean ± SEM.
Compound (L2Sn) treated group demonstrated a significant decrease in blood
glucose concentration as compared to the negative control group (p<0.001; q= 18.037;
Fig. 4.22C), while a significant decrease occurred in serum triglycerides concentration
as compared to positive and negative control groups (p<0.001;q= 14.455 and
Chapter 4 Results & Discussion
Page 80
p<0.001; q= 7.758; Fig. 4.23A). Serum cholesterol concentration decreased in
comparison to the positive control group (p<0.05; q= 4.430; Fig. 4.23B).
10
100
200
300
400
500
Tri
glyc
erid
es(C
once
ntra
tin
mg/
dl)
****
**
**
10
100
200
300
400
500Negative Control
Positive Control
HL
(1)(2)(3)(4)C
hole
ster
ol(C
once
ntra
tion
mg/
dl)
* * *
(A) (B)
Fig. 4.23 (A) Triglyceride and (B) Cholesterol level in mice treated with HL,
(L2VO) (1), (L2Sn) (2), (L2Zn) (3) and (L2Co) (4) compound groups
Except (L2Co), mice treated with HL, (L2VO), (L2Sn), and (L2Zn) decreased serum
triglycerides concentration as compared to both control groups (p<0.001), (L2VO)
treated group has a significant effect on decreasing serum cholesterol concentration as
compared to both control groups (p<0.05), while groups treated with (L2Sn) and
(L2Zn) decreased cholesterol concentration as compared to positive control group
(p<0.05). Data are presented as mean ± SEM.
(L2Zn) treated group showed a significant decrease in blood glucose level as
compared to the negative control groups (p<0.001; q= 16.161; Fig. 4.22D), a
significant decrease in serum triglycerides concentration relative to positive and
negative control groups (p<0.001; q= 13.416 and p<0.05; q= 6.718, Fig. 4.23A),
while a decrease was noticeable in serum cholesterol concentration as compared to the
positive group (p<0.05; q= 4.408, Fig. 4.23B). (L2Co) treated group showed
significant decrease in blood glucose level as compared to negative control (p<0.001;
q= 19.361; Fig. 4.22E), while no effect was evident on serum triglycerides
concentration and cholesterol concentrations (Fig. 4.23A & B).
4.2.9 In Vitro studies: Cytotoxicity analysis on HeLa The compounds HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) were tested for their
cytotoxic and anticancer activity against HeLa cell line at different concentrations
(150, 100, 50, 10, and 5 µM) using SRB assay. SRB stained samples were visualized
Chapter 4 Results & Discussion
Page 81
with Olympus IMT-2 inverted microscope (magnification 60X) and photographed
(Fig. 4.24 A-Q). The viability curves (Fig. 4.25) were generated and IC50 values for
the complexes are given in Table 4.9.
Fig. 4.24 Effect of concentration of HL and its compexes (L2VO), (L2Sn), (L2Zn)
& (L2Co) on HeLa cells*
* Cultures treated with different concentrations of the compounds for 24 h were stained with SRB,
visualized with Olympus IMT-2 inverted microscope (magnification 60X) and photographed. A.
Untreated control, B. Solvent (5% DMSO + 95% Ethanol) treated sample, C, D, E, treated with 100,
50, and 10 μM of HL, respectively, F, G, H, treated with 100, 50, and 10 μM of (L2VO), respectively, I,
Chapter 4 Results & Discussion
Page 82
J, K, treated with 100, 50, and 10 μM of (L2Sn), respectively, L, M, N, treated with 100, 50, and 10
μM of (L2Zn), respectively, O, P, Q, treated with 100, 50, and 10 μM of (L2Co), respectively.
Fig. 4.25 Viability curves for HL, (L2VO) (1), (L2Sn) (2), (L2Zn) (3) and (L2Co)
(4) complexes in HeLa.
Cultures were exposed to serial dilutions of the compounds (150, 100, 50, 10, and 5
µM) for 24 h and relative (%) viabilities (mean±SD) were measured using SRB
assay
Table 4.9 IC50 value (µM) of HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) complexes
for HeLa cell line
Compound IC50 value (µM)
HL 106.70
(L2VO) 40.66
(L2Sn) 5.92
(L2Zn) 42.82
(L2Co) 107.68
In comparison to the untreated culture, HL was found toxic to the cells as
observable from Fig. 4.24 C, D, E. A significant decrease in relative viability
(p˂ 0.05) was observed with an IC50 value of 106.7 µM. (L2VO) strongly influenced
the cellular morphology and cells became more fiber like as shown in Fig. 4.24 F, G,
H. A pronounce decrease in relative viability (p˂ 0.001) of the treated cultures was
0
50
100
4.5 45 450Rel
ativ
e (%
) via
bilit
y
Concentration (µM)
HL(2)(3)(1)(4)
Chapter 4 Results & Discussion
Page 83
observed and this complex was twice toxic to the cells compared to HL, with an IC50
value of 40.66 µM.
Cells were very sensitive to (L2Sn) where few live cells can be seen in Fig.
4.24 I, J, K at 100 and 50 µM. The complex induced a highly significant decrease in
relative viability (p˂ 0.001) of the cultures. The IC50 value of (L2Sn) was 5.92 µM,
thus, cells were 18 times more sensitive to this complex in comparison to HL. (L2Zn)
had a strong effect similar to (L2VO) on the relative viability (p˂ 0.001) of the
cultures with an IC50 value of 42.82 (p˂0.001; Fig. 4.24 L, M, N). When compared to
HL, (L2Co) was less found toxic to the cultures with an IC50 value of 107.68 (p˃0.05;
Fig. 4.24 O, P, Q).
4.2.10 Inhibition of alkaline phosphatase of HL and complexes Literature survey reveals that Schiff bases show alkaline phosphatase (ALP)
inhibition activity but their complexes show more enhanced activity. They can block
the active sights of enzyme. Alkaline phosphatase uses a wide variety of
phosphomonoesters to catalyze the transfer of phosphate groups to water [10].
Experimental observations show that inhibition activity of HL is quite insignificant
while the presence of metal ions enhance the deactivation of enzyme. Dibutyl tin
complex has minimum activity due to long alkyl chains having non-polar nature while
Zn complex has shown excellent inhibition of nearly 90%. Inhibition profile can be
seen in Fig. 4.26.
50 100 150 200 2500
20
40
60
80
Inhi
bitio
n of
ALP
(%)
Concentration (µM)
HL L2VO L2Sn L2Zn L2Co
Fig. 4.26 Inhibition profile of alkaline phosphatase (ALP) at different
concentrations of HL, (L2VO), (L2Sn), (L2Zn) and (L2Co) (error less than 3 %)
Chapter 4 Results & Discussion
Page 84
4.2.11 Antimicrobial activities of HL and complexes
Inhibitory action of HL and its metallic complexes was recorded against various
strains of bacteria (Bacillus subtilis, Escherichia coli, Pasturella multocida and
Staphlocuccus aureus) and fungi (Ganoderma lucidum Aspergillus niger, Alternaria
alternate and Penicillium notatum) by disc diffusion method [13]. Streptomycin and
fluconazole were selected for use as positive controls for antibacterial and antifungal
screening tests, respectively. A recommended concentration of a test sample/reference
drug (1mg/1mL of solvent) was introduced into the discs. The biologically active
samples inhibited the bacterial/fungal growth around discs; the zones of inhibition
were measured. The data have been summarized in Table 4.10 and Table 4.11.
Table 4.10 Comparison of antibacterial activities of HL and complexes
Compounds
Average zone of inhibition (mm)
E. coli B. subtilis S. aureus P. multocida
Streptomycin 30a ± 0.27 30ab ± 0.19 30ab ± 0.22 30a ± 0.29
HL 0 0 14c ± 0.18 14c ± 0.15
(L2VO) 18c ± 0.27 15c ± 0.19 20bc ± 0.17 17bc ± 0.22
(L2Sn) 22bc ± 0.17 19bc ± 0.13 22bc ± 0.22 25ab ± 0.13
(L2Zn) 28ab ± 0.12 31a ± 0.14 32a ± 0.18 33a ± 0.11
(L2Co) 25ab ± 0.13 29ab ± 0.19 28ab ± 0.19 31ab ± 0.11
Concentration = 1mg/mL in DMSO; 0 = No activity, 5-10 = Activity present, 11-25 = Moderate
activity, 26-40 = Strong activity; Antibacterial values are mean ± S.D of samples analyzed individually
in triplicate at p < 0.1; Different letters in superscripts indicate significant differences. a = maximum
activity, b = intermediate activity, c = minimum activity, ab = activity between maximum and
intermediate and bc = activity between intermediate and minimum.
The free ligand (HL) was almost inactive against the tested organisms. The
coordination of ligand with a metal (V, Sn, Zn, Co) induced antimicrobial activities in
consequent complexes. Activity of the synthesized complexes was found in good
agreement with their structures. The results demonstrate that each of the coordinated
metals (V, Sn, Zn, Co) plays an important role in biological actions of such
complexes [14]. The zinc complex (L2Zn) was found to be the most potent inhibitor
of bacteria/fungi while the vanadyl product (L2VO) displayed the least activity among
Chapter 4 Results & Discussion
Page 85
all the metal coordinated products. The coordination of ligand with the dibutyltin(IV)
moiety has appreciably enhanced the antibacterial/antifungal activity of product
complex (L2Sn). The inhibitory action of organotin(IV) compounds can be related to
their interaction ability with DNA and proteins. After zinc, the highest antimicrobial
potential was observed for ligand-Co coordination. Complexes 1-4 were found more
potent inhibitors of fungal culture as compared to bacterial growth [15].
Table 4.11 Comparison of antifungal activities of HL and complexes
Compounds
Average zone of inhibition (mm)
A. alternate G. lucidum A. niger P. notatum
Fluconazole 38a ± 0.41 41ab ± 0.52 40ab ± 0.36 35ab ± 0.27
HL 19c ± 0.22 13c ± 0.21 18c ± 0.33 20c ± 0.27
(L2VO) 22bc ± 0.25 25bc ± 0.29 31ab ± 0.28 36ab ± 0.22
(L2Sn) 25bc ± 0.26 39ab ± 0.19 33c ± 0.25 30bc ± 0.21
(L2Zn) 36ab ± 0.12 42a ± 0.14 45a ± 0.08 40a ± 0.18
(L2Co) 32ab ± 0.12 40ab ± 0.18 38ab ± 0.17 34ab ± 0.16
Concentration = 1mg/mL in DMSO; 0 = No activity, 5-10 = Activity present, 11-25 = Moderate
activity, 26-40 = Strong activity; Antifungal values are mean ± S.D of samples analyzed individually in
triplicate at p < 0.1; Different letters in superscripts indicate significant differences. a = maximum
activity, b = intermediate activity, c = minimum activity, ab = activity between maximum and
intermediate and bc = activity between intermediate and minimum.
4.2.12 Computational study of HL and complexes
The DFT optimized structures of the HL and complexes under study were obtained
using 6-311++G** and 6-311G++ basis sets, respectively. Mullikan charges, dipole
moments, optimization energies, spin multiplicity and point groups were calculated
from the optimized structures shown in Fig. 4.27 and Table 4.12. Comparison of bond
lengths obtained from computational studies with that of X-ray analysis showed a
good agreement with each other (see Table 4.6). NMR spectra of HL was calculated
theoretically using 6-311++G** basis set and compared with the experimental findings
(see Table 4.13).
Chapter 4 Results & Discussion
Page 86
Fig. 4.27 Numbering scheme of HL used in computational studies
Table 4.12 Selected parameters of BPIMN and CPIMN calculated by
RB3LYP/3-21G set using DFT method in GUSSIAN03W package
HL (L2VO) (L2Sn) (L2Zn) (L2Co)
Charge 0 0 0 0 0
Spin Singlet Doublet Singlet Singlet Doublet
Total Energy (a.u.) -864.49 -2732.40 -7983.245 -3488.497 -3093.86
Dipole Moment (D) 4.3093 2.6432 3.857 1.4192 2.5232
EHOMO (a.u.) -0.20933 -0.20474 -0.11964 -0.19039 -0.16426
ELUMO (a.u.) -0.08627 -0.06867 0.02566 -0.05793 -0.06152
Point Group C1 C1 C1 C1 C1
Chapter 4 Results & Discussion
Page 87
Table 4.13 Observed and calculated hydrogen chemical shifts (δ, in ppm)
Atoms B3LYPa /6-311++G** Asignment δ (ppm)b C10 191.3 153.1 C8 151.1 171.7 C12 147.3 131.8 C5 145.2 136.8 C2 144.1 140.6 C14 143.2 136.4 C4 138.8 133.3 C3 138.0 130.7 C18 136.7 128.0 C16 135.5 127.0 C6 134.9 127.8 C13 134.1 129.3 C11 133.7 116.9 C17 129.3 122.7 C15 123.2 123.4 C1 121.1 118.6 C9 115.0 108.6 C21 23.2 21.0 C20 21.0 18.1 H29 9.33 9.27 s H28 8.37 7.69 d H24 7.99 8.07 d H25 7.84 7.67 d H23 7.02 7.12 m H33 7.02 2.35 s H37 2.65 2.44 s H34 2.52 2.35 s H38 2.45 2.44 s H35 2.31 2.35 s H36 2.18 2.44 s
a GIAO/ B3LYP/6-311++G** Ref. to TMS b HL dissolved in CDCl3 containing 0.03% vol. TMS as reference; s = singlet; d = doublet;
m = multiplet
Studies reveal that the zwitter ionic form of HL is more stable and feasible
than its enolic form. This structure is also consistent with its x-ray studies. Charge
Chapter 4 Results & Discussion
Page 88
distribution values indicate that O atom of hydroxyl group and N atom of azomethine
group are more negative in charge thus provide oxidizing sites.
4.3 Redox behavior of a novel menadiol derivative
4.3.1 Cyclic voltammetry of HMNA Cyclic voltammetry of 4-hydroxy-5-methoxynaphthalene-1-yl acetate (HMNA) was
carried out using 0.2 mM solution (80% aqueous ethanol), at pH 4.0 using 0.1 M
acetate buffer at a clean GCE between the limits ±1.5. On proceeding from 0 to +1.5
V, a single oxidation signal was observed at +0.37 V. On reversing the scan direction,
a reduction peak was observed at a potential corresponding to -0.193 V with a counter
oxidation signal. Therefore, further cyclic voltammetric experiments were carried out
between +0.6 and −0.5 V using 0 V as a starting potential.
Fig. 4.28 CVs (scan 1–5) of 0.2 mM HMNA obtained at 100 mV s-1 in a
solution of pH 4.0, using 0.1 M acetate buffer
Number of scans ranging from 1 to 5 determined the extent of adsorption/
passivation of the electrode surface by HMNA and its redox products on the electrode
surface. Fig. 4.28 demonstrates the effect of subsequent scans. There is a considerable
decrease of peak current on the 2nd scan showing adsorption of HMNA on GCE
surface but peak current becomes constant after 5th scan and a steady state is
eventually reached. Furthermore, the second scan is accompanied with the appearance
Chapter 4 Results & Discussion
Page 89
of another oxidation peak at a potential less positive to the first oxidation signal.
Interdependency of peaks was ensured using peak clipping experiments. The cyclic
voltammograms obtained in the potential range of 0.0 to +0.6 V established that peak
c in Fig. 4.28 is related to peak a1 while peak a2 is the result of the oxidation of
reduction product obtained from peak c. CVs of 0.2 mM HMNA obtained in different
electrolytes of pH ranging from 1.22 to 12.93 at 100 mV s-1 scan rate is shown in Fig.
4.29 A and B. The effect of pH of the media was used to establish the redox
mechanism of the analyte. With increase in pH of the medium, both peaks a1 and c
were shifted cathodically indicating the involvement of proton during electron
transfer. The peaks were not resolved at pH 12.93 showing non-feasibility of HMNA
oxidation process in strongly alkaline conditions. Epa vs. pH plot with a slope of 51
mV pH-1 according to the equation: Epa = 0.5 + 0.051 V pH suggested the oxidation of
HMNA to occur by the same number of electrons and protons [16, 17].
Fig. 4.29 CVs obtained in Ar saturated solution of 0.2 mM HMNA at a GCE at 100
mVs-1 in different supporting electrolytes of pH ranging from 1.22 to 7.0 (A) and
8.07-12.93 (B)
Effect of increasing scan rate was also studied to evaluate the diffusion
coefficient of HMNA at pH 4.0 and 7.0. Peak current increased with increasing scan
rate with no significant peak potential shift (Fig. 4.30 A and B). At pH 4.0 two
distinct straight line segments were obtained when Ipa was plotted vs. the square root
of scan rate (see Fig. 4.31 A), one being at lower scan rates (up to 100 mV s-1) and
Chapter 4 Results & Discussion
Page 90
other at higher scan rates (up to 400 mV s-1). Therefore, the diffusion coefficient (Dο)
was calculated separately for each linear segment. The slope values at lower and
higher scan rates were 19.82 × 10-6 and 48.8 × 10-6 A/(V s-1)1/2 respectively. From
experimentally obtained Epa −Epa/2 values of 32 and 71 mV at pH 4.0 and 7.0, αan was
calculated using equation:
𝐸𝑝𝑝 − 𝐸𝑝𝑝2
= 47.7/(αan) (4.6)
Fig. 4.30 (A) CVs obtained with the GCE in a solution of 0.2 mM HMNA in pH 4,
0.1 M acetate buffer at different scan rates ranging from 20-400 mV s-1 (B) CVs
obtained with the GCE in a solution of 0.2 mM HMNA in pH 7, 0.1 M Phosphate
buffer at different scan rates
Fig. 4.31 (A) Plot of Ipa(µA) vs. square root of scan rate(Vs-1) at pH 4.0 (B)
Plot of Ipa(µA) vs. square root of scan rate (Vs-1) at pH 7.0
Chapter 4 Results & Discussion
Page 91
Considering n = 1 (from DPV), the diffusion coefficient with values 1.46 ×10-6 and
8.85 ×10-5 cm2s-1 was determined using 𝐼𝑝𝑝 = 2.99 × 105𝑛(𝛼𝑝𝑛)1/2𝑛𝑛0 × 𝐷𝑜1/2𝜈1/2
at pH 4.0 for scan rates below and above 100 mVs-1 [18-20]. Using the slope value
(26.8×10-6 A/(V s-1)1/2) of Ipa vs. ʋ1/2 at pH 7.0 (Fig. 4.31B), diffusion coefficient was
calculated as 1.1 × 10-5 cm2 s-1.
Fig. 4.32 (A) Plots of log Ipa(µA) vs. log ν (V/s) at pH 4.0 and (B) pH 7.0
In order to find out the diffusion-controlled nature of the process, log Ipa was
plotted vs. log ʋ as shown in Fig. 4.32 A and B. The slope values of 0.83 and 1.3
indicated the adsorption controlled process at pH 4.0 while the slope value of 0.57 at
pH 7.0 showed intermediate (diffusion cum adsorption controlled) nature of the
process [21-23]. Adsorption controlled nature of the process is also indicated by the
sharp peaks of CVs (Fig. 4.30A) obtained at pH 4.0 [19]. The reason for their
sharpness is related to the fact that during an adsorption-controlled process, electron
transfer takes place after the analyte is being adsorbed and it does not have to travel to
the electrode surface. The broad shape of cyclic voltammetric peaks at pH 7.0 shown
in Fig. 4.30B demonstrate mixed mode of transport.
Heterogeneous electron transfer rate constant (ks) was evaluated at pH 4.0
from CVs obtained at different concentrations. Effect of varying concentrations has
been shown in Fig. 4.33A. By plotting peak current vs. concentration, a straight line
was obtained with a slope of 26.7×10-6 A/mol m-3 (Fig. 4.33B). Reinmuth’s
expression (4.1) was employed for the determination of ks. The ks with a value of
3.89×10-3 cms-1 characterized the oxidation process to be quasi-reversible in nature.
Chapter 4 Results & Discussion
Page 92
Fig. 4.33 (A) CVs of different concentration of HMNA (B) Plot of Ipa(µA) vs.
concentration(µM/cm3) obtained at 100 mV s-1 scan rate in a medium buffered at pH
4.0 using 0.1 M acetate buffer
4.3.2 Differential pulse voltammetry of HMNA DPVs of 0.2 mM HMNA obtained in different pH media have been depicted in Fig.
4.34. The width at half peak height (W1/2) of 98 mV (close to the theoretical value of
90.4 mV) evidenced one electron oxidation process [24]. The shift of peak potential
with rise in pH showed similar trends as obtained from cyclic voltammetry. Epa was
plotted against pH and a slope of 53 mV per pH unit (close to the theoretical value of
59 mV per pH unit) authenticated the CV results of one electron–one proton oxidation
process. A dramatically different situation from cyclic voltammetry, where oxidation
peak related to the first oxidation product appeared in the second scan was the
appearance of that peak in the first scan of DPVs under strongly acidic conditions.
The appearance of this peak suggests the presence of a small amount of the oxidized
form of HMNA that is detected by DPV due to its more sensitivity as compared to
CV. This attribution is supported by the fact that acid catalyzed hydrolysis of HMNA
can lead to the formation of oxidized product which undergoes two electron–two
proton oxidation process as witnessed by the W1/2 value of 50 mV.
Chapter 4 Results & Discussion
Page 93
Fig. 4.34 DPVs of 0.2 mM HMNA obtained at 10 mVs-1 scan rate in
supporting electrolytes of pH 1.22 to 12.93
4.3.3 Square wave voltammetry of HMNA The square wave voltammetry (SWV) has an edge over other electrochemical
techniques because of greater speed of analysis, little consumption of the electroactive
species in comparison to DPV and reduced problems with poisoning of the electrode
surface. Another advantage of SWV is that the reversibility of electron transfer
process can be checked during one scan only. Since the current is sampled
simultaneously in both positive and negative going pulses so peaks corresponding to
oxidation and reduction of the electroactive species at the electrode surface can be
obtained in the same experiment. The main objective of SWV is to ensure the
reversibility, irreversibility or quasi-reversibility of the redox process of the analyte.
SWV of 0.2 mM HMNA was performed at GCE which endorsed the quasi-
reversible nature of the oxidation process (Fig. 4.35 A and B). SW voltammogram at
pH 4.24 is shown along with the effect of number of scans at pH 7.4. In cyclic
voltammetry another oxidation peak related to the first oxidation appeared in second
scan similar to DPV. In SWV this peak was observable in the first scan under acidic
conditions as well, owing to the much more sensitivity of SWV in comparison to CV
and DPV [25]. Appearance of this oxidation peak in the first scan suggests that a
small amount of the oxidized form corresponding to this potential already existed in
Chapter 4 Results & Discussion
Page 94
solution as a result of acid catalyzed hydrolysis of HMNA leading to the oxidized
product which could undergo two electron–two proton oxidation as depicted by the
W1/2 value of 50 mV from DPV obtained in supporting electrolytes of pH 1.22.
Moreover, the same peak disappeared from the first scan square wave voltammogram
recorded in a medium buffered at pH 7.4 and observed only in the second scan like
CV indicating the inability of HMNA to get hydrolyzed under these conditions.
B
Fig. 4.35 SW voltammograms obtained in solution of 0.2 mM HMNA in pH
4.2 (A) 1st scan showing total current (It), forward current (If) and backward
current (Ib) (B) SWVs of first three scans of 0.2 mM HMNA recorded at 100
mVs-1 scan rate in pH 7.4
4.3.4 Redox mechanism of HMNA
The oxidation of 4-hydroxy-5-methoxynaphthalene-1-yl acetate was found to occur
by the transfer of a single electron as determined from the average W1/2 value of 98
mV. Epa vs. pH plot with a slope of 53 mV pH-1 indicated the oxidation to occur by
one electron one proton process. Based upon these values it was proposed that
hydroxyl moiety of HMNA after the loss of one electron and proton results in the
formation of oxygen radical, which can exist in three resonance forms. The canonical
form in which electron is residing on carbon atom having attached acetate group gets
converted to dione (5-methoxynaphthalene-1,4-dione) and acetyl radical by the
cleavage of C-O bond as shown in Scheme 4.8. The acetyl radical thus formed can
dimerize to the simplest diketone, butane-2,3-dione. The cathodic peak in the reverse
scan of CV suggested the reduction of 5-methoxynaphthalene-1,4-dione to its
corresponding diol form by two electrons and two protons involvement (Scheme 4.8).
Chapter 4 Results & Discussion
Page 95
Appearance of another anodic peak in the second scan of CV can be corresponded to
the oxidation of diol back to dione by the loss of two electrons and protons in a
quasireversible manner. The oxidation of diol occurs at lower potential as compared
to 4-hydroxy-5-methoxynaphthalene-1-yl acetate because in the latter case acetate
being the electron-withdrawing group makes the oxidation process difficult thus
resulting in more positive oxidation potential of HMNA. This attribution was
supported by performing DFT calculation using B3LYP, 6-31G basis set for the
optimization and energy calculation of HMNA and 5-methoxynaphthalene-1,4-diol
thus authenticating our proposed redox mechanism of HMNA as explained later in the
computational study of the compounds.
OCH3
O
OH
C
O
CH3
OCH3
O
O
C
O
CH3
-e
-H+
OCH3
O
O
C
O
CH3
OCH3
O
O
C
O
CH3
CH3 C
O
OCH3
O
O
+
OCH3
OH
OH
2H+
2eButan-2,3-dione
dimerization
Scheme 4.8 Proposed redox mechanism of HMNA in slightly acidic, neutral
and alkaline conditions
The peak (which is recorded in the second scan of CV) corresponding to the
oxidation of diol is observed in the first scan DP voltammograms obtained in low pH
media. Its appearance in strongly acidic conditions can be related to the reversible
Chapter 4 Results & Discussion
Page 96
acidic hydrolysis of the acetate group that results in the formation of diol as shown in
Scheme 4.9. The absence of second anodic peak in the first scan of DP
voltammograms recorded in media buffered at pH ∼ 7.0 ruled out the possibility of
HMNA hydrolysis under these conditions. However, the broadness of the DPV peak
in high pH conditions may be due to the merging of the signals of HMNA and its
hydrolytic product. For further confirmation of HMNA hydrolysis its high
performance liquid chromatography (HPLC) was carried out in acidic, neutral and
alkaline conditions. Only one signal at 4.283 min was observed in neutral medium
indicating the purity of HMNA while two peaks in acidic and alkaline conditions
were noticed indicating the presence of the starting compound and its hydrolytic
product (see Fig. 4.36). The hydrolysis in acidic medium supports the results obtained
from differential pulse voltammetry. However, the hydrolysis in basic medium is a
dramatically different behavior from that observed in DP voltammograms. The
anomaly can be resolved by considering the broadness of the DPV peak in high pH
conditions to be due to the merging of the signals of HMNA and its hydrolytic
product.
Fig. 4.36 HPLC of HMNA carried out at 300 nm in acidic and neutral conditions
5-Methoxynaphthalene-1,4-dione is structurally related to menadione but its
cathodic potential is much less negative as compared to that reported for menadione
Chapter 4 Results & Discussion
Page 97
thus demonstrating its facile reduction. Moreover, its easy reduction in biological
systems as compared to menadione may result in the generation of reactive oxygen
species (ROS). This is because if the reduction potential is more positive than -0.5 V,
the species fall in the electrochemical range that may sustain electron transfer
reactions in vivo with possible generation of ROS [26].
OCH3
O
OH
C
O
CH3
OCH3
OH
OH
H+
H2O+ CH3COOH
OCH3
O
O
C
O
CH3
H+ -e 2H+ -2e
OCH3
O
O
Adsorption/Dimerization
+ CH3CHOO
Scheme 4.9 Proposed redox mechanism of HMNA and its hydrolyzed product
under strongly acidic conditions
4.3.5 Electronic absorption spectroscopy of HMNA Electronic absorption spectroscopy of 25 µM HMNA was carried out in different pH
media to determine the pKa of phenolic group attached to HMNA molecule. The
Chapter 4 Results & Discussion
Page 98
spectra showed an intense band at 223 nm and a broad band with overlapping peaks in
the range of 275–350 nm. Absorption maxima at 223 nm are thought to be due to
benzene moiety [27]. The fine structure consists of overlapping peaks including n-π*
transitions may be due to acetate group attached to HMNA. Absorption spectra of
naphthalene show a similar maxima and fine structure but at a shorter wavelength
[28]. The shift of peaks to longer wavelengths in case of HMNA (substituted
naphthalene) is due to the electron releasing nature of –OCH3 and –OH groups. Fig.
4.37 and Fig. 4.38A demonstrate the effect of pH on UV–Vis absorption spectra of
HMNA. An increase in pH resulted in slight bathochromic shift of the absorption
maxima at 223 nm while no shift was observed for the peaks of fine structure until pH
9.2. The shift in peak position is attributed to the removal of –OH proton in basic
media with the negative charge on oxygen atom, thus making the transitions
energetically more favorable as represented by the bathochromic shift. The pKa of
HMNA with a value of 9.6 was estimated from the plot of absorbance vs. pH at 300
nm owing to the pronounced change in absorbance and bathochromic effect at this
wavelength (see Fig. 4.38B). For this purpose absorption spectra were taken with
minor increment of pH from 9.2 to 10.2 for the evaluation of more accurate value of
pKa. The pH 9.6 thus corresponds to the deprotonation of phenolic proton of the
naphthalene moiety of HMNA. The pKa of the analyte was not observable in
differential pulse voltammetry because of the probable overlapping of the peaks of
HMNA and its hydrolytic product.
Chapter 4 Results & Discussion
Page 99
Fig. 4.37 Absorption spectra of 25 µM HMNA obtained in pH 4.0-9.2 using 20%
ethanol and 80% aqueous supporting electrolytes
Fig. 4.38 (A)Absorption spectra of 25 µM HMNA obtained in pH 9.2-10.2 using
solvent system of 20% ethanol and 80% aqueous supporting electrolytes (B) Plot of
absorbance of HMNA vs. pH at 300 nm
4.3.6 Computational studies Computational studies of HMNA and 5-methoxynaphthalene-1,4-diol were carried
out to obtain Mulliken charge distribution and energies of HOMO and LUMO.
Molecules were first optimized using 6-31G basis set. Calculations of the optimized
Chapter 4 Results & Discussion
Page 100
structures were then carried out to obtain the charge densities and values of EHOMO
and ELUMO. Charges distribution of HMNA and 5-methoxynaphthalene-1,4-diol are
shown in Fig. 4.39. In the case of HMNA, Mulliken charge distribution shows
negative charge on oxygen atoms with the most negative charge on oxygen of -OH
group indicating its facile oxidation. Similarly, in case of 5-methoxynaphthalene-1,4-
diol, the most negative charge can be seen on the oxygen atoms of the two -OH
groups.
Table 4.14 EHOMO and ELUMO values of compounds
Compound EHOMO ELUMO HMNA (4-hydroxy-5-methoxynaphthalene-1-yl acetate) -0.19813 -0.03838 5-methoxynaphthalene-1,4-diol -0.18028 -0.01635
Chapter 4 Results & Discussion
Page 101
A B
Fig. 4.39 (A) Charge distribution of HMNA and (B) 5-methoxynaphthalene-1,4-diol
A B
Fig. 4.40 (A) Graphical representation of EHOMO of HMNA and (B)
5-methoxynaphthalene-1,4-diol
DFT calculations were performed using 6-31G basis set to verify the proposed
mechanism (see Fig. 4.40). Energies of HOMO and LUMO were calculated for
HMNA and 5-methoxynaphthalene-1,4-diol. The DFT calculations showed 5-
methoxynaphthalene-1,4-diol to have less negative value of EHOMO than HMNA. This
finding is in good agreement with the experimental anodic peak of 5-
methoxynaphthalene-1,4-diol located at less positive potential than HMNA. Less
negative EHOMO value is a manifestation of facile oxidation of the species [29-32]. An
examination of EHOMO values listed in Table 4.14 reveals that the difference in the
EHOMO values match well with the variation of oxidation potentials of the two species.
Chapter 4 Results & Discussion
Page 102
4.4 Redox behavior of two anthracenedione derivatives
4.4.1 Cyclic voltammetry of HAC and DHDN
Cyclic voltammograms (CVs) of HAC obtained at 100 mVs-1 scan rate in the
potential range of ±1.5 V in pH 7.0 are shown in Fig. 4.41. Starting the potential scan
from 0 V and scanning in the positive direction, an oxidation signal appeared around
1.17 V. On the reverse scan of the same voltammogram, a cathodic peak generated at
-0.61 V, was coupled by an anodic wave at -0.41 V. The same number of peaks was
observed in the second scan with a decrease in peak currents can be related to the
adsorption of redox products on the electrode surface. Moreover, the electro-
polymerization products of phenoxy radicals have been reported as the main source of
electrode passivation [33]. Thus, the prominent decrease in current intensity of HAC
at 1.17 V corresponds to the oxidation of the phenolic moiety of HAC which leads to
the formation of polymeric compounds that lowers the sensitivity of the electrode.
The oxidation and reduction of the analyte were found to occur independently. Hence,
for the detailed investigation of HAC, further voltammetric assays were conducted
separately in positive and negative potential domain of glassy carbon electrode.
-1.5 -1.0 -0.5 0.0 0.5 1.0 1.5
-20
-10
0
10
20
I p / µA
E / V vs SCE
(1)(2)
Fig. 4.41 CVs of 0.2 mM solution of HAC scan-I (1) and scan-II (2) recorded
in solution of pH 7.0 at 100 mVs-1 scan rate
All the CVs corresponding to the reduction region of HAC were recorded in
the potential range of -0.2 V to -1.1 V. For examining the effect of pH on the
reduction of HAC, CVs were obtained in different media of pH 1.2 - 12.8 as given in
Chapter 4 Results & Discussion
Page 103
Fig. 4.42A. A regular shift in the location of cathodic peak with change in pH of the
medium indicated the involvement of proton during electron transfer reaction [3, 4,
34, 35]. The dislocation of peak stopped at pH close to10 indicating the reduction of
HAC to occur only by the involvement of electrons in highly alkaline conditions [17].
Intersection of the linear segments of peak potential versus pH plot (Fig. 4.42B)
represented 9.8 as the acid base dissociation constant, pKa of HAC. The Nernstian
slope of 58.2 mVpH-1 indicated the reduction of HAC to occur by the involvement of
the same number of electrons and protons [19]. Similarly, the oxidation of the
corresponding reduction product showed strong medium dependency up to pH 8.4.
Epa as a function of pH suggested the oxidation to occur by the participation of the
same number of electrons and protons.
-1.2 -1.0 -0.8 -0.6 -0.4 -0.2-25
-20
-15
-10
-5
0
5
10
I p /
µA
E / V vs SCE
A
(1)(11)
0 2 4 6 8 10 12 14
-0.9
-0.8
-0.7
-0.6
-0.5
-0.4
-0.3
-0.2E
/ V v
s SC
E
pH
R2 = 0.992Y = -0.0582X -0.264
R2 = 0.990Y = -0.067X -0.167
∇→ anodic peak♦→cathodic peak
B
Fig. 4.42 (A) CVs of 0.2 mM HAC obtained at 100 mVs-1 in different pH media: (1)
pH 2.0 (2) pH 3.0 (3) pH 4.0 (4) pH 5.0 (5) pH 6.0 (6) pH 7.0 (7) pH 8.0 (8) pH 10.0
(9) pH 11.0 (10) pH 12. 8 (B) Plot of Ep vs pH of both anodic and cathodic peaks
For the investigation of oxidation, the potential range was fixed between the
limits of +0.3 V to +1.3 V and CVs were recorded in different pH conditions using
0.2 mM HAC in the presence of 0.1M supporting electrolytes. No signal appeared at
and below pH 3 but in media of pH higher than 3, an oxidation peak was noticed in
the CVs of HAC. However, this oxidation peak showed broadness at pH 10 and 11
indicating the merging of two anodic peaks. The peak changed its position with
variation in pH of the medium thus offering evidence for the involvement of protons
during electron transfer reactions. Moreover, peak current got significantly decreased
in the 2nd scan and disappeared in the 3rd scan. These peculiar CV characteristics
indicated the adsorption of oxidation product on the electrode surface [36].
Chapter 4 Results & Discussion
Page 104
The redox behavior of N2 saturated solution of 1.0 mM DHDN was also
investigated by CV at 100 mVs-1 scan rate in a medium of pH 7.0 using 0.1 M
phosphate buffer. The CVs were initially started at +0.00 V and recorded between
potential limits of ±1.50 V (see Fig. 4.43). Two anodic and two cathodic peaks were
noticed in the CV of DHDN. The peaks position shifted to more negative potentials
with rise in pH. This behavior demonstrated the electron transfer reactions of DHDN
to depend strongly on the pH of the medium [34, 35].
-1.5 -1.0 -0.5 0.0 0.5 1.0 1.5
-140
-105
-70
-35
0
35
E / V vs. SCE
I / µA
a1
a2
c1
c2
(1)
(2)
Fig. 4.43 Cyclic voltammograms of (1) the solvent and (2) 1 mM DHDN showing
anodic peaks (a1 and a2) and cathodic peaks (c1 and c2) in a medium of pH 7.0 at
100 mVs-1
CVs of 1 mM DHDN were also recorded at different scan rates in medium
buffered at pH 7.0. The slight displacement of reduction potentials to more negative
values with increasing scan rate (ν) indicated the quasi-reversibility of cathodic
processes. The signals in the reverse scan with current intensities less than cathodic
peaks offered another evidence for the quasi-reversible nature of DHDN reduction
[19]. The current intensities of both peaks varied linearly with the square root of ν as
consistent with the diffusion-controlled reduction of solution species. The slope of the
plot of logIpc1 vs. logν with a value of 0.54 (close to the theoretical value of 0.50) also
witnessed the reduction process to be diffusion-limited [37]. By plotting Ipc1 vs. ν1/2,
the diffusion coefficient with a value of 3.2 × 10-6 cm2s-1 was evaluated using the
equation [38].
𝐼𝑝𝑝 = −2.99 × 105𝑛(𝛼𝑝𝑛)1/2𝑛𝑛∗𝐷1/2𝜈1/2 (4.7)
Chapter 4 Results & Discussion
Page 105
For the evaluation of diffusion coefficient of HAC, CVs were recorded in media of
pH 5.0 and 7.0 at different scan rates. A regular rise in peak current without a
significant shift in peak potential was observed with increase in scan rate. At pH 7.0,
the plot of Ipc vs square root of scan rate gave a slope of -9.22 × 10-5 A/(Vs-1)1/2. In
case of CV obtained in pH 5.0, two distinct segments appeared when Ipc was plotted
against v1/2 corresponding to scan rate below and above 100 mVs-1. However, the scan
rate lower than 100 was used for diffusion co-efficient and αn calculations. Using the
difference of peak potential (Epc) and half peak potential (Epc/2), diffusion coefficients
with values 3.63 × 10-5 and 6.06 × 10-5 cm2 s-1 were obtained at pH 5.0 and 7.0. Small
value of D under acidic conditions in comparison to neutral medium is suggestive of
thick solvation sphere around HAC that lowers its mobility towards the electrode
surface. The higher D of HAC than DHDN can be explained on the basis of Stokes-
Einstein criteria according to which smaller molecule should move quickly to the
electrode surface than comparatively larger molecule [39]. The less polar nature of
HAC as determined by its computationally determined smaller dipole moment of 3.67
Debye compared to 5.61 Debye of DHDN also supports this finding.
Nicholson equation was used for the determination of heterogeneous electron
transfer rate constant of HAC. This equation is based on correlation between ks and
peak separates on of a reversible/quasi-reversible redox process by a dimensionless
parameter, Ψ, which was calculated from the following equation [40].
Ψ = (𝐷𝑜/𝐷𝑅)𝛼2𝑘𝑠/(𝜋𝛼𝐷𝑜)1/2 (4.8)
Putting Ψ as 8.40 for peak separation of 65 mV at pH 7.0 and substituting the
values of α = 0.71 and Do/Dr = 0.36, ks was calculated as 2.91 × 10-3 cm s-1. The ks of
DHDN with a value of 9.92×10-4 cm s-1 was determined by using Reinmuth’s
equation (4.1), suitably applicable on an irreversible electrode reaction [41]. Thus, the
values of ks support the quasi-reversible and irreversible nature of the redox processes
of HAC and DHDN according to the criteria reported in literature [42].
4.4.2 Differential pulse voltammetry of HAC and DHDN
The voltammetric response of HAC was also studied by the most sensitive
technique, differential pulse voltammetry to ensure the number of electrons involved
in its redox reaction. The differential pulse voltammograms of 0.2 mM HAC recorded
in the pH range 1.2 to 12.8 at 5 mVs-1 scan rate can be seen in Fig. 4.44 A. The pH
Chapter 4 Results & Discussion
Page 106
dependent shift of peak potential supports the results obtained from CV. The pH
dependency of the reduction process is a proof of the involvement of protons. The pKa
of HAC with a value of 9.7 was evaluated from the intersection of the two linear
segments of Ep vs pH plot (see Fig. 4.44B).
-1.2-1.0-0.8-0.6-0.4-0.20.0
036
9
12
-30-25-20
-15
-10
-5
0
pH
E / V vs SCE
I / µ
A
A
0 2 4 6 8 10 12
-0.8
-0.7
-0.6
-0.5
-0.4
-0.3
-0.2
E pc /
V
pH
R2 = 0.994Y = - 0.056x - 0.233
B
Fig. 4.44 (A) DPVs of 0.2 mM HAC showing reduction in various pH media
(B) Plot of Epc vs pH of HAC
The average peak width at half height (W1/2) of 50 mV (close to the theoretical
value of 45 mV) suggested the involvement of 2 electrons in the reduction of HAC.
Like the results obtained from CV, the slope (56 mV pH-1) of the plot of Epc as a
function of pH signified the electrons to be coupled by the same number of protons.
Thus, the reduction of HAC occurs by the gain of 2 electrons and 2 protons. In
contrast to reduction, the DPVs of HAC oxidation demonstrated in Fig. 4.45 indicates
the loss of electrons in two one electron transfer processes as evidenced by the broad
anodic peak upto pH 9 and its splitting into two sub peaks (involving one electron
each) in strongly alkaline conditions. Ep - pH topography indicated the loss of electron
accompanied by proton.
Chapter 4 Results & Discussion
Page 107
0.40.6
0.81.0
4
6
810
12
0.00.51.01.52.02.5
3.0
3.5
4.0
I / µ
A
E / V vs SCE
pH
Fig. 4.45 DPVs representing oxidation of 0.2 mM HAC in different pH media
-1.4 -1.2 -1.0 -0.8 -0.6 -0.4 -0.2 0.0
-30
-25
-20
-15
-10
-5
0
I /µ
A
E / V vs. Ag/AgCl
A
0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.60.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
4.0
I / µ
A
E / V vs. SCE
a1 a2
B
Fig. 4.46 DPVs of 0.5 mM DHDN obtained in solution of pH 7.0 at a scan rate of 5
mV s-1 showing (A) reduction and (B) oxidation labelled by prominent anodic peaks
a1 and a2
Differential pulse voltammetry of DHDN was also performed for determining
the number of electrons involved in its anodic and cathodic processes. For probing the
reduction, differential pulse voltammograms shown in Fig. 4.46 A were obtained in
pH 7.0 at 5 mVs-1. The widths at half peak height, W1/2 were found close to 90 mV
indicating the occurrence of reduction by the involvement of 1e- and 1H+ in each step
[32]. The oxidation supported by the appearance of prominent anodic signals shown
Chapter 4 Results & Discussion
Page 108
in Fig. 4.46B was also investigated under similar conditions. The corresponding
proposed mechanism can be seen in section 4.4.4.
4.4.3 Square wave voltammetry HAC and DHDN
-1.2 -1.0 -0.8 -0.6 -0.4 -0.2 0.0
-40
-30
-20
-10
0
10
I / µ
A
E / V vs SCE
It If Ib
B
-1.2-1.0
-0.8-0.6
-0.4-0.2
0.0
3
6
9
12
-40
-30
-20
-10
0
pH
E / V vs SCE
I / µ
A
A
A
Fig. 4.47 (A) SWVs representing reduction of 0.2 mM HAC in different pH media
(B) SWV of 0.2 mM HAQ in pH 7.0 showing forward and backward components of
the total current
The exact nature of the reversibility of HAC reduction process was
authenticated by square wave voltammetry. The square wave voltammograms
(SWVs) depicted in Fig. 4.47A were recorded in the pH range 1.2 to 12.8 using 0.2
mM HAC solution of the analyte. The shifting of peak potential in a pH dependent
manner was found in very good agreement with the CV and DPV results. The square
wave voltammetric response showed the quasi-reversible nature of the reduction
process because the ratio of backward and forward current components of the total
current was found less than 1 (Fig. 4.47B). The same direction of the peaks appearing
in the forward and backward components of the total square wave voltammogram
recorded in the positive potential window of GCE signified the irreversible nature of
reduction process as supported by the CV results.
Chapter 4 Results & Discussion
Page 109
Table 4.15 Parameters obtained from SWV
Cathodic
peaks S.D.
Sensitivity
(nA ·μM-1)
Adj. R-
Square
LOD
(μM)
LOQ
(μM)
Linearity
(μM)
Peak 1c 1.65×10-3 2.37 0.98437 2.09 6.96 50
Peak 2c 1.47×10-3 3.62 0.98625 1.22 4.06 >50
Fig. 4.48 represents the square wave voltammograms of DHDN. 1st and 2nd
reduction steps showed reversible and quasi-reversible behavior as witnessed by the
equal and unequal forward and backward current components of the total current. In
oxidation region, three anodic peaks were noticed. The first two anodic peaks showed
features of reversibility and the third characterized irreversible electron abstraction.
Effect of different concentration of DHDN was also examined by SWV. Peak current
increased linearly with concentration. Limit of detection (LOD) and limit of
quantification (LOQ) of DHDN were electro-analytically determined from the linear
dependence of cathodic peak currents as a function of concentration in the range of 1
µM to 50 µM. LOD and LOQ with values listed in Table 4.15 were calculated
according to the relations (4.2 and 4.3) [12]. The detection limits of 2nd cathodic
signal were found lower as compared to peak 1c due to its more sensitive character.
Chapter 4 Results & Discussion
Page 110
-1.0 -0.8 -0.6 -0.4 -0.2 0.0-60
-50
-40
-30
-20
-10
0
10
20
I / µ
A
E / V vs. SCE
A
0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4-10
0
10
20
30
40
I / µ
A
E / V vs. SCE
B
Fig. 4.48 SWVs obtained at100 mVs-1 showing the nature of cathodic (A) and
anodic (B) peaks of 0.5 mM DHDN in pH 7.0
4.4.4 Proposed redox mechanism HAC and DHDN
The W1/2 value of 50 mV, unequal forward and backward current components and Epc
versus pH plot with a slope of 58 mV pH-1 indicate one step quasi-reversible HAC
reduction involving the gain of two electrons and two protons as shown in Scheme
4.10. The appearance of two signals in pH dependent DPVs suggested the oxidation
of HAC to occur by a mechanism depicted in Scheme 4.11. The absence of oxidation
signal at pH ≤ 3 is attributed to the fact that in highly acidic medium, the electrophore
of HAC gets blocked for oxidation due to high concentration of H+.
O
O
OH
CH2OHO
OH
O
CH2OH
H
2e- + 2 H+
Scheme 4.10 Proposed reduction mechanism of HAC
Chapter 4 Results & Discussion
Page 111
O
O
OH
CH2OHO
O
O
-1e- - 1 H+ CH2OH
O
O
O
CH2OH-1e- - 1 H+
O
O
O
CH2OH
Scheme 4.11 Suggested oxidation mechanism of HAC
From the appearance of two cathodic signals with W1/2 values corresponding
to the transfer of one electron per reduction step, ratio of forward and backward
components of the total current and shift of peak potential with rise in pH, DHDN is
suggested to reduce in two 1e- and 1H+ reversible and quasi-reversible reduction
steps, thus, resulting in the conversion of dione to diol. The mechanism is presented in
Scheme 4.12. The nitro groups present in the analyte should reduce to hydroxyl amine
and nitroso products but the experimental observations showing no signs of NO2
electro-reduction due to their probable involvement in intramolecular hydrogen
bonding with -OH groups present in their close vicinity. DHDN can exist in keto and
enol forms so its three steps oxidation can be attributed to the electron abstraction
from -OH groups. The oxidation product of enol form of the analyte can lose another
electron from 2nd -OH group. The irreversibility of the third step can be related to the
dimerization or adsorption of the product at the electrode surface. The oxidation
mechanism of DHDN is presented in Scheme 4.13.
Chapter 4 Results & Discussion
Page 112
12
34
56
O 7
89
10
1112
13 14
O15
16
N+ 17O18 O-19
N+20O 21-O
22
OH23
OH24
1H+, 1e-
OH
O
N+O O
N+O-O
OH OH
OH
OH
N+O O
N+O-O
OH OH
1H+, 1e-
Scheme 4.12 Proposed reduction mechanism of DHDN
O
O
N+O O
N+O-O
OH O
O
O
N+O O
N+O-O
O O
O
O
N+O O
N+O-O
O OH
O
O
N+O O
N+O-O
O OGCEGCE
O
O
N+O O
N+O-O
O O
O
O
N+O O
N+O-O
OH OH
-1e-
-1H+-1e-
-1H+
-1e-
-1H+
Scheme 4.13 Suggested oxidation mechanism of DHDN
Chapter 4 Results & Discussion
Page 113
4.5 Redox mechanism, kinetic and thermodynamic
parameters of a novel anthraquinone
4.5.1 Cyclic voltammetry of HCAQ
To obtain a general picture of the redox behavior of the compound, cyclic
voltammetry (CV) measurements of HCAQ were recorded in media of different pH
using a potential window of ±1.3 V. The voltammograms obtained for pH 3.0 and 7.0
can be seen in Fig. 4.49. The CV response under neutral conditions (pH = 7) depicts a
sharp oxidation signal, 1a, at +0.69 V, followed by a broad peak, 2a, at +1.02 V. The
absence of reduction peaks corresponding to 1a and 2a in the backward scan at pH 7.0
demonstrates the irreversible nature of these electrochemical oxidation processes. The
compound shows two successive one-electron reduction peaks in the negative
potential domain of the GCE. The reduced species resulting in peaks 1c and 2c gets
re-oxidized at nearly the same potential as evidenced by a single broad oxidation
peak, 2a*. For confirming 2a* to be the combination of two overlapping peaks, a
more sensitive electrochemical technique, DPV, was employed. The differential pulse
voltammograms of the analyte shown in Fig. 4.50 unequivocally represent the
splitting of peak 2a* into two sub-peaks at lower concentration. The anodic and
cathodic peak potential differences witness the first and second reductions to follow
reversible and quasi-reversible processes. Unlike the behavior of the analyte for pH
7.0, the backward CV scan for pH 3.0 shows an oxidation-dependent reduction peak
at +0.33 V. The different voltammetric signatures in acidic and neutral media indicate
the redox behavior of HCAQ to depend strongly on the pH of the medium.
Chapter 4 Results & Discussion
Page 114
Fig. 4.49 1st scan CVs of 1 mM HCAQ obtained in pH 3.0 and 7.0 at 100 mV s-1
-1.2 -1.0 -0.8 -0.6 -0.4 -0.2 0.068
101214161820222426
b
I/µA
E / V vs. Ag/AgCl
a 0.4 mMb 0.6 mM
a
Fig. 4.50 DPVs of 0.4 and 0.6 mM HCAQ recorded in pH 7.0 at 5 mVs-1 after
cathodic scan without cleaning the electrode surface
An experiment based on the variation of peak current Ip, with scan rate ʋ, was
used for the determination of diffusion coefficient and heterogeneous electron transfer
rate constant. The scan rate effect of HCAQ was monitored for pH 7.0 at different
temperatures. The voltammograms at 323 K shown in Fig. 4.51A demonstrate a linear
increase in Ip with the square root of scan rate. Peak 1a shows a large potential shift
with scan rate while the reduction peaks exhibit a slight shift in potential. These
voltammetric features are characteristics of irreversible (peak 1a) and quasi-reversible
redox processes (peak 2c).
Chapter 4 Results & Discussion
Page 115
Fig. 4.51 Effect of (A) scan rate at 323±1K and (B) temperature at 80 m Vs-1
on the cyclic voltammetric behaviour of 0.7 mM HCAQ
The values of diffusion coefficients at different temperatures were obtained
from the plots of Ip (current intensity of peak 1a) versus square root of scan rate as
given in Fig. 4.52 using following form of Randles–Sevcik equation [43].
𝐼𝑝 = 5.16×106𝛼(𝛼𝑝𝛼)1/2𝐴𝐴𝐷1/2𝜈1/2
𝑅1/2 (4.9)
Where n is the number of electrons transferred, αa is the anodic charge transfer
coefficient, and A, D, C and ʋ are the area of the electrode in cm2, diffusion
coefficient in cm2s-1, bulk concentration of the species in molcm-3, scan rate in Vs-1,
and T is temperature in kelvin scale respectively. The values of αan using equation
(4.6) were determined for the irreversible process at different temperatures ranging
from 308 to 328 K. An examination of Fig. 4.51B reveals that the effect of
temperature on peak potential of the irreversible process is more pronounced as
compared to the quasi-reversible process due to their respective kinetic and
thermodynamic-controlled natures. The values of D show an increasing trend with
increasing temperature due to the probable decrease in viscosity of the media. The
slope values of logIp vs. logʋ fall in the range of 0.50–0.53 at all temperatures, thus
indicating the redox process to be limited by diffusion [44].
Chapter 4 Results & Discussion
Page 116
Fig. 4.52 Plots of peak current of (A) 1a and (B) 2c against the square root of
scan rate using 0.7 mM solution of HCAQ
Heterogeneous rate constant (ks) is an important diagnostic criterion for
predicting the nature of a redox process. Large values of ks suggest that, following the
application of an applied potential, equilibrium between the oxidized and reduced
species gets re-established quickly. In contrast, small values of ks indicate slow
kinetics and a longer time required for equilibrium [45, 46]. The values of ks listed in
Table 4.16 were calculated on the basis of peak 1a (Fig. 4.49) at different
temperatures using equation (4.1) [5, 47, 48].
Table 4.16 Diffusion coefficient and heterogeneous rate constant values of HCAQ
evaluated on the basis of oxidation peak 1a at different temperatures
Peak Temperature (K)
1/T×10-3 (K-1) αna
D/10-7 (cm2 s-1)
ks/10-4 (cm s-1) log ks
1a
308 3.25 0.99 0.16 2.01 (0.02) 3.698
313 3.19 1.12 0.48 2.31 (0.02) 3.636
318 3.14 1.25 1.17 3.66 (0.04) 3.437
323 3.09 1.32 1.34 6.15 (0.07) 3.211
328 3.05 1.40 4.04 7.27 (0.08) 3.139
The heterogeneous electron transfer rate constant of HCAQ was evaluated
from the peaks corresponding to the quasi-reversible redox process using the
following form of the Nicholson equation [49].
Chapter 4 Results & Discussion
Page 117
𝑘𝑠 = 𝛹�𝜋𝐷𝑜𝑎𝜈 (4.10)
where Ψ is the charge transfer parameter, Do the diffusion coefficient of the
oxidized species, ʋ the scan rate and 𝛼 = 𝑛𝑛/𝑅𝑅. The values of dimensionless
parameter Ψ, were determined from the difference of anodic and cathodic peak
potential [50]. The ks values (Table 4.17) of a quasi-reversible redox process of
HCAQ are greater than those of irreversible oxidation as expected [51]. The values of
∆Ep at the experimental temperatures were converted to ∆𝐸𝑝298 by using equation
(4.11) [52].
Δ𝐸𝑝298 = �298𝑅� �Δ𝐸𝑝𝑅� (4.11)
The ks values (listed in Table 4.16) corresponding to the redox process labeled
as 1a were used for the determination of thermodynamic parameters. ∆G# was
evaluated by equation (4.12) [53].
Δ𝐺# = 5778.8(5.096 − 𝑙𝑙𝑙𝑘𝑠) (4.12)
With an increase in temperature, heterogeneous rate constant increases and
∆G# decreases. The Arrhenius equation, 𝑘𝑠 = 𝑛𝑒−𝐸𝑝/𝑅𝑅 was used for the calculation
of enthalpy and entropy changes of activation. Ea obtained from the slope of a plot of
logks against 1/T (Fig. 4.53) allowed the evaluation of ∆H# and ∆S# according to
equations (4.13) and (4.14). An examination of Table 4.18 reveals an increase in ∆S#
and a slight decrease in ∆H# values with increasing temperature.
Δ𝑑# = 𝐸𝑝 − 𝑅𝑅 (4.13)
Δ𝐺# = Δ𝑑# − 𝑅Δ𝑆# (4.14)
Table 4.17 Heterogeneous electron transfer rate constant values of HCAQ
evaluated on the basis of quasi-reversible redox peaks at different temperatures
Temperature (K) ks/10-4 (cm s-1)
313 1.21 (0.04)
318 7.60 (0.19)
323 7.79 (0.16)
Chapter 4 Results & Discussion
Page 118
Fig. 4.53 Plot of log ks as a function of 1/T
Table 4.18 Thermodynamic parameters of HCAQ
Peak Temperature
(K) ∆G#/(kJ/mol) Ea/(kJ/mol) ∆H#/(kJ/mol) ∆S#/(J/K.mol)
1a
308 51.00
62.66
60.09 29.54
313 50.37 60.06 30.95
318 49.31 60.01 33.67
323 48.09 59.97 36.77
328 47.51 59.93 37.87
Although estimated errors in ∆G# and Ea are ± 2%, additional significant figures are given to
allow more accurate calculation of ∆H# and ∆S#
The positive ∆G# value is responsible for the non-spontaneity of the oxidation
reaction. The slight decrease in ∆G# with increasing temperature shows the tendency
of the electrode process towards spontaneity. The trend of ∆S# shown in Table 4.18
indicates that, at higher temperature values, the oxidized product of HCAQ diffuses
Chapter 4 Results & Discussion
Page 119
quickly from the electrode surface towards the solution. The trend is the same as
predicted by transition state theory �𝒌 = 𝜿 𝒌𝑩𝑻𝒉𝒆−𝚫𝑯
#𝑹𝑻� 𝒆𝚫𝑺
#𝑹� � [54]. Thus, at higher
temperature the increase in ∆S# and decrease in ∆H# are responsible for higher rate
constant values.
4.5.2 Differential pulse voltammetry of HCAQ
The DPV technique is associated with the minimizing of charging current and
consequent enhanced sensitivity [55]. Therefore, the redox behavior of HCAQ was
investigated by DPV in media of different pH. The DPVs shown in Fig. 4.54A and
Fig. 4.55 indicate a shift in peak potentials with changing pH. This behavior
corresponds to the involvement of protons during the electron transfer processes.
Moreover, the shift of Epa towards lower potentials with increasing pH shows facile
electron abstraction under basic conditions. The differential pulse voltammograms in
the potential domain of 0 to +1.3 V shows two oxidation peaks at low pH and three
peaks at higher pH values. DPV was also employed to study the reduction of HCAQ.
The potential scan in the negative direction has a couple of peaks corresponding to the
two-step reduction of HCAQ. The DPV scan in the potential domain of -1.3 to 0 V
following scanning from 0 to -1.3 V without cleaning the electrode surface shows a
single reverse oxidation peak. Peak 1a with a width at half peak height (W1/2) of 54
mV indicates the involvement of two electrons in this oxidation step [56]. For the
determination of the number of protons involved in the oxidation process, peak
potential was plotted against pH.
Chapter 4 Results & Discussion
Page 120
Fig. 4.54 (A) DPVs of 0.7 mM HCAQ recorded at 5 mVs-1 in 50% ethanol and 50%
BR buffer of pH 3-11(B) Plot of Ep versus pH using DPV data of peak 1a
Fig. 4.55 (A) 1st and (B) 2nd scan DPVs showing reduction of 0.7 mM HCAQ and
oxidation of the reduction product in 50% ethanol and 50% BR buffer of pH 3-11 at
5 mVs-1
Chapter 4 Results & Discussion
Page 121
The slope of the plot shown in Fig. 4.54B, shows that the oxidation
(corresponding to peak 1a) occurs with the involvement of two protons and two
electrons. The intersection of the two linear segments of the Ep vs. pH plot indicates
HCAQ to have an apparent acid dissociation constant value of 10.4. In contrast to
peak 1a, peak 2a is broad indicating one electron loss during this oxidation step. The
plot of peak potential versus pH shown in Fig. 4.56 demonstrates the removal of an
electron to be accompanied by a proton. Similar behavior was also observed for peak
2a’.
2 3 4 5 6 7 8 9 10 11 12 13
-0.75
-0.70
-0.65
-0.60
-0.55
-0.50
-0.45
-0.40
-0.35
E p / V
vs. A
g/Ag
Cl
pH
Ep = -0.055 pH -0.25R2 = 0.98
pKa = 9.1
Fig. 4.56 Ep as a function of pH using DPV data of peak 1c
4.5.3 Square wave voltammetry of HCAQ Square wave voltammetry (SWV), being one of the most sensitive and fast
electrochemical techniques [57], was employed to determine the limit of detection
(LOD) and limit of quantification (LOQ) of HCAQ. From the calibration curves
shown in Fig. 4.57A, LOD and LOQ with values of 0.08 and 0.26 mmolL-1 were
determined according to the previously reported equations (4.2) and (4.3) [12].
Chapter 4 Results & Discussion
Page 122
0.0 0.1 0.2 0.3 0.4 0.5 0.60
1
2
3
4
5
6B
I p/µ
A
Concentration (µmol/cm3)
Slope = 27.325
0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8
0
2
4
6
8
10
12
14
16
E / V vs. Ag/AgCl
I/µA
a 1.00 mMb 0.80 mMc 0.60 mMd 0.50 mMe 0.40 mMf 0.20 mMg 0.10 mMh 0.08 mMi 0.06 mMj 0.04 mM
a
j
A
Fig. 4.57 (A) SWV of different concentrations of HCAQ obtained at 100 mV
s-1 in a medium of pH 10 (B) Plot of peak current versus concentration
4.5.4 Electronic absorption spectroscopy of HCAQ
Fig. 4.58 (A) UV-Vis spectra of different concentration of HCAQ obtained in
pH 7 (B) Absorption spectra of 16 µM HCAQ recorded in pH 3-11
UV-visible spectroscopy was carried out for the purpose of photometric
characterization and confirmation of the apparent acid dissociation constant obtained
from electrochemical techniques. Simple anthracenedione has been reported to give
four electronic absorption signals [58]. The signal at 207 nm is attributed to n–σ*
transition of the carbonyl group while the others are assigned to π– π* transition. The
signal at 252 nm is allocated to π– π* transition of the benzenoid system and those at
Chapter 4 Results & Discussion
Page 123
272 and 326 nm are assigned to π– π* transition of the quinonoid system. The
assignment of signals is made on the basis of comparison with the UV spectrum of
anthracene [58]. The UV-visible spectra of HCAQ show five signals (Fig. 4.58A).
The signal at 207 nm (n– σ* transition of carbonyl group) is absent in the spectrum of
HCAQ due to blocking of the lone pair on -C=O because of possible hydrogen
bonding between -OH at 1 and 4 positions. The signals at 232 and 252 nm are
assigned to π– π* transitions of the benzenoid system while that at 287 nm is assigned
to π– π* transition of the quinonoid system [58, 59]. The splitting of the benzenoid
band may be due to the existence of mesomeric structures or conjugation breakage.
The broad band (at 425–550 nm) lacking in the absorption spectrum of simple
anthraquinone [60] is present in the spectrum of HCAQ. As the hydroxyl group is
electron donating and gives an absorption band in the visible region as a result of a
charge transfer (CT) mechanism [59, 61] so the spectral band displacements are
related to the increase in length of the chromophore by additional ring formation due
to intermolecular hydrogen bonding [61, 62].
The absorption spectra of HCAQ in media of different pH are shown in Fig.
4.58B. Observation of the spectra reveals that the absorption band due to CT
mechanism is strongly affected by the pH as compared to other peaks. The band
corresponding to CT mechanism of the orange-colored solution appears in the range
of 470–500 nm at pH < 9 and gets shifted to 552–590 nm at pH ≥ 9 due to
deprotonation of HCAQ. The drastic color variation (see Fig. 4.59) of HCAQ from
pH 10 to 11 unequivocally indicates that HCAQ ionizes at pH-10. The isosbestic
point at 299 nm indicates that HCAQ and its mesomeric form exist in equilibrium
with the same value of molar extinction coefficient. The presence of -OH groups at
positions 1 and 4 is responsible for the mesomeric form of HCAQ. Another isosbestic
point at 516 nm represents HCAQ and its deprotonated form existing in equilibrium
and absorbing the same amount of light at this wavelength. Isosbestic point typically
indicates the presence of only two species [63-65]. The occurrence of isosbestic
points for multicomponent solutions is almost impossible because the likelihood of
identical absorbance of three or more compounds at any wavelength is negligibly
small [63].
Chapter 4 Results & Discussion
Page 124
Fig. 4.59 Color variation of 16 µM HCAQ in pH 3-12
From the slopes of the plots of absorbance versus concentration, the molar
extinction coefficients with values shown in Fig. 4.60A were calculated at the
respective wavelengths. On the basis of the peaks at 475 and 554 nm, the apparent
acid dissociation constant of HCAQ with a value of 10.1 was obtained from the plot
of absorbance versus pH as shown in Fig. 4.60B. This apparent acid dissociation
constant value is in close agreement with that obtained from DPV and the drastic
color variation shown in Fig. 4.59.
Fig. 4.60 (A) Plots of absorbance versus concentration using UV-Vis spectroscopic
data obtained in pH 7 (B) Absorbance of 16 µM HCAQ as a function of pH
Chapter 4 Results & Discussion
Page 125
4.5.5 Redox mechanism of HCAQ
The oxidation of HCAQ exhibited two peaks in acidic and three in basic
media (see Fig. 4.54). The oxidation corresponding to peak 1a was found to occur by
the transfer of two electrons as determined from the average W1/2 value of 54 mV.
The Epa vs. pH plot indicated the oxidation to occur with the involvement of two
protons. On the basis of the results obtained from DPV and SWV it is proposed that
the hydroxyl groups of HCAQ, after the loss of two electrons and two protons,
presumably form oxygen radicals [47] that may lead to conversion to 2-(3-hydroxy-3-
(trichloromethyl)pentyl)-8-methoxyanthracene-1,4,9,10(4aH,9aH)-tetraone [66] as
shown in Scheme 4.14.
O
OOCH3
OH
OH OH
CCl3
O
OOCH3
O
O OH
CCl3 -2H+
O
OOCH3
O
O OH
CCl3-2e-
Scheme 4.14 Proposed oxidation mechanism of HCAQ corresponding to peak 1a
Fig. 4.54 shows that peak 2a is pH dependent in acidic conditions while
independent of pH in basic media and thus the mechanistic pathway gets switched
from an EC to a CE mechanism (see Scheme 4.15 and Scheme 4.16). The average
W1/2 value and slope of Epa vs. pH plot of peak 2a show a one-electron and one proton
oxidation process. This peak is assigned to the hydroxyl group of the pentyl
substituent because the trichloromethyl is an electron-withdrawing group, and
therefore this hydroxyl group is oxidized at high potential.
O
OOCH3
OH
OH OH
CCl3
O
OOCH3
OH
OH O
CCl3 -1H+
O
OOCH3
OH
OH O
CCl3-1e
GCE
Dimerization Scheme 4.15 Proposed oxidation mechanism corresponding to peak 2a showing
proton-coupled electron loss in acidic conditions
Chapter 4 Results & Discussion
Page 126
O
OOCH3
OH
OH OH
CCl3
O
OOCH3
OH
OH O
CCl3-1H+
O
OOCH3
OH
OH O
CCl3-1e
Adsorption Dimerization
Scheme 4.16 Suggested CE type oxidation mechanism of HCAQ corresponding to
peak 2a in neutral and alkaline conditions
The DPVs in basic media show the appearance of a third pH-dependent peak
2a’. The half peak potential value and shifting of peak position with pH show the
involvement of one electron and one proton in this process. On the basis of these
results, a CEC mechanism is proposed as shown in Scheme 4.17.
OH
OH
O
OOCH3
O
CCl3
H
H
HO
H OH
OH
O
OOCH3
O
CCl3
OH
H-1H+
OH
OH
O
OOCH3
O
CCl3
O
H
OH
OH
O
OOCH3
O
O
H
OH
OH
O
OCH3OCH3
O
H
O
-1e-
Adsorption
Dimerization
Scheme 4.17 Proposed oxidation mechanism of HCAQ corresponding to peak 2a’ in
neutral and alkaline conditions
The results of peak width at half peak current and Epc versus pH plots revealed
the two-step reduction of HCAQ to occur with the involvement of one electron and
one proton in each step. The mechanism proposed is shown in Scheme 4.18.
OH
OH
O
OOCH3
CCl3
OH
1e
1H+
OH
OH
OH
OOCH3
CCl3
OH
1e
1H+
OH
OH
OH
OHOCH3
CCl3
OH
Scheme 4.18 Proposed redox mechanism of HCAQ corresponding to peaks appearing
in the negative potential window of GCE
Chapter 4 Results & Discussion
Page 127
4.6 Temperature responsive redox behavior of
hydroxyanthracenediones
4.6.1 Cyclic voltammetry
Cyclic voltammetry studies of the selected anthracenediones was first carried out
between the potential limits of ±1.5 V in a medium buffered at pH 7.0. Consecutive
scans of the analytes were recorded to know the extent of adsorption of the
compounds on the electrode surface. Fig. 4.61A demonstrates the effect of successive
scans on the peak current of the derivative 4,8-dihydroxy-9,10-dioxo-9,10-
dihydroanthracen-1-yl acetate (HACAD). The decrease of the current intensity after
the first scan indicates the adsorption of the analyte at the electrode surface but the
peak current becomes constant after several scans and a steady state is finally reached.
The CV response reveals a sharp oxidation signal, 1a, of the analyte at +0.55 V,
followed by a broad peak, 2a, at +1.02 V. The absence of reduction peaks
corresponding to 1a and 2a in the reverse scan points to the irreversible nature of
these electrochemical oxidation processes [2]. In the negative potential range, the CV
of HACAD shows two reduction signals labeled as 1c and 2c at a potential of -0.63
and -0.90 V respectively. On the reverse scan, the emergence of a single anodic peak
is related to the oxidation of the reduced product of HACAD. The cyclic
voltammograms of 1,4,5-trihydroxy-2-methyl-3-(3-oxobutyl)anthracene-9,10-dione
(HOAD) and 1,4,5-trihydroxyanthracene-9,10-dione (HAD) were also obtained in
medium of pH 7.0 and both compounds exhibited quite closely related
electrochemical behavior to HACAD.
Chapter 4 Results & Discussion
Page 128
Fig. 4.61 (A) Consecutive CVs of 0.6 mM HACAD obtained in pH 7.0 at 100 mVs-1
(B) CVs of 0.6 mM HAD and (C) HACAD at different scan rates
4.6.2 Evaluation of kinetic parameters from Cyclic voltammetry
As electrochemistry mainly focuses on the kinetics of reactions at the electrode
surface [5] therefore, the values of different kinetic parameters such as diffusion
coefficient (D) and heterogeneous electron transfer rate constant (ks) were obtained
from CV by monitoring the variation of the peak current with the scan rate. The
Randles-Sevcik equation relates the magnitude of the peak current (Ip) to the
temperature (T), the concentration (C), the electrode area (A), the number of electrons
involved in oxidation or reduction (n), the diffusion coefficient (D), and the scan rate
[5, 67]. The temperature effects of scan rate on the redox properties of the three
anthracenediones were monitored at pH 7.0 at different temperatures. The
voltammograms at 323 K are shown in Fig. 4.61 B and C and show an increase of Ip
with increasing scan rate. The values of the diffusion coefficients are summarized in
Chapter 4 Results & Discussion
Page 129
Table 4.19 and were obtained from the slope of the plots shown in Fig. 4.62. An
examination of Table 4.19 reflects that the diffusion coefficients are affected by
molecular mass and temperature. The values of D decrease with increasing molecular
mass because heavier molecule of larger radius moves slowly towards the electrode
according to the Stokes-Einstein equation [5]. As the temperature increases, the
solvent viscosity decreases, which in turn increases the mobility of analytes.
0.12 0.16 0.20 0.24 0.28 0.322
4
6
8
10
12
14
16
18
20 323 [Ip = 5.63×10-5ν1/2] 318 [Ip = 4.56×10-5ν1/2] 313 [Ip = 3.07×10-5ν1/2] 308 [Ip = 1.84×10-5ν1/2]
I pa/µ
A
ν1/2 (V/s)1/2
A
0.12 0.16 0.20 0.24 0.28 0.32 0.363
6
9
12
15
18
21
24B
328K [Ip = 6.49×10-5ν1/2] 323K [Ip = 5.45×10-5ν1/2] 318K [Ip = 4.94×10-5ν1/2] 313K [Ip = 4.61×10-5ν1/2] 308K [Ip = 3.99×10-5ν1/2]
I pa/µ
A
ν1/2 (V/s)1/2
0.10 0.15 0.20 0.25 0.30 0.351
2
3
4
5
6
7
8
9
A
323 [Ip = 2.559×10-5ν1/2] 318 [Ip = 1.826×10-5ν1/2] 313 [Ip = 1.313×10-5ν1/2] 308 [Ip = 9.635×10-6ν1/2]
I pa/µ
A
ν1/2 (V/s)1/2
Fig. 4.62 Plots of peak current corresponding to peak 1a against square root of scan
rate using 0.6 mM solution of (A) HACAD, (B) HAD and (C) HOAD
The CV experiments also showed a variation of peak potentials (Ep) with scan
rate and temperature. Fig. 4.61B and C shows a prominent shift of the peak potentials
towards more positive potentials with increasing sweep rate for the oxidation
corresponding to peak 1a, suggesting the irreversible nature of the process [68]. The
cathodic shift of Ep indicates facile oxidation at higher temperature. The value of ks,
based on peak 1a was evaluated at different temperatures using equation (4.15) [69].
𝐸𝑝 = 𝐸𝑜 + 𝑅𝑅(1−𝛼)𝛼𝑝𝛼
�0.78 + 2.32
𝑙𝑙𝑙 �𝐷(1−𝛼)𝛼𝑝𝛼𝑘𝑠2𝑅𝑅
�� + 2.32
𝑅𝑅(1−𝛼)𝛼𝑝𝛼
𝑙𝑙𝑙𝜈 (4.15)
The value of the standard electrode potential, Eo was taken as the potential
corresponding to 0.85 Ip [70]. From the value of intercept, ks was determined at
different temperatures and our results clearly indicate that ks increases with increasing
temperature showing a tendency towards reversibility at higher temperature as shown
in Table 4.19. The ks of all the three compounds are quite similar and slight
differences can be related to variations in their molecular masses because the
processes are diffusion controlled and the sluggish diffusion of a comparatively
heavier molecule limits the redox processes [71].
Chapter 4 Results & Discussion
Page 130
Table 4.19 Diffusion coefficient and heterogeneous rate constant values of the
selected compounds at different temperatures
Compounds Mol. Wt (g/mol)
Temperature (K) D cm2/s Ksh cm/s
HACAD 298.05
308 2.82×10-6 2.58×10-6
313 8.23×10-6 5.56×10-6
318 1.90×10-5 1.38×10-5
323 3.04×10-5 2.80×10-5
HOAD 340.10
308 1.47×10-7 1.13×10-7
313 2.50×10-7 3.17×10-7
318 4.68×10-7 5.94×10-7
323 8.75×10-7 1.16×10-6
HAD 256.04
308 3.10×10-6 1.56×10-6
313 4.08×10-6 2.17×10-6
318 4.61×10-6 6.33×10-6
323 5.53×10-6 1.61×10-5
4.6.3 Thermodynamic parameters from cyclic voltammetry
Information about the spontaneity/non-spontaneity of the redox reactions can be
obtained from thermodynamic parameters [68]. The rate constant evaluated from scan
rate studies at different temperatures was used for the determination of the
thermodynamic parameters listed in Table 4.20. ∆G# was evaluated from ks using
equation (4.12) [72].
Chapter 4 Results & Discussion
Page 131
3.1x10-3 3.1x10-3 3.2x10-3 3.2x10-3 3.2x10-3 3.2x10-3
-6.6
-6.0
-5.4
-4.8
-4.2
-3.6
c
b
a HOACAQ logko = -6969.77×1/T+17.0b HAQ logko = -5968.14×1/T+13.3c HMOAQ logko = -6574.35×1/T+14.4
logk
o
1/T (K)-1
a
Fig. 4.63 Plot of logko as a function of 1/T
The Arrhenius equation, 𝑘𝑠 = 𝑛𝑒− 𝐸𝑝𝑅𝑅 was used for the calculation of the
enthalpy (∆H#) and entropy (∆S#) changes of activation. The value of the activation
energy (Ea) obtained from the slope of log ks versus 1/T plot (Fig. 4.63) was inserted
into equation (4.13) for the evaluation of ∆H#. The value of ∆𝑆# was obtained using
equation (4.14) [2]. An inspection of Table 4.20 reveals a decrease in ∆H# and an
increase of ∆S# with an increase in temperature. Similarly, ks increases and ∆G#
decreases with increasing temperature. The overall result shows a trend toward
reversibility and spontaneity at higher temperatures.
Chapter 4 Results & Discussion
Page 132
Table 4.20 Thermodynamic parameters of the redox processes of the selected
compounds
Compounds Temperature (K)
∆G# (kJ/mol)
Ea (kJ/mol)
∆H# (kJ/mol)
∆S# (J/K. mol)
HACAD 308 61.75 133.28 130.72 223.93 313 59.81 130.68 226.39 318 57.54 130.63 229.87 323 55.76 130.59 231.68
HOAD 308 313 318 323
69.59 67.00 65.43 63.76
125.72 123.16 173.91 123.11 179.27 123.07 181.26 123.03 183.50
HAD 308 63.63 114.12 111.56 155.62 313 63.01 111.52 155.00 318 62.18 111.48 155.04 323 59.49 111.44 160.83
4.6.4 Differential pulse voltammetry
Fig. 4.64 (A) DPV’s of HACAD obtained in 50% ethanol and 50% BR buffer of
pH 3-11 at 5 mVs-1 (B) Plots of peak potential as function of pH
Chapter 4 Results & Discussion
Page 133
Differential pulse voltammetric technique is associated with an enhanced
sensitivity due to its ability of minimizing charging current [2]. Hence, the pH
dependent redox behaviour of HACAD, HAD, and of HOAD was investigated by
DPV. The DPVs HOAD exhibit the same two oxidation peaks in the potential
window of 0 to +1.5 V as expected from the CV results discussed earlier. However,
the responses of HACAD and HAD are different as an additional sharp peak (labelled
as a*) appears under acidic conditions. The DPVs of HACAD are shown in Fig.
4.64A and show a shift of the peak potentials with a change of the pH of the medium,
indicating proton-coupled electron transfer processes [73]. Oxidation becomes more
facile at higher pH as indicated by a shift of the Ep towards more negative potentials.
This suggests a stabilization of oxidized species by hydroxyl ions present in the
solution [5]. The DPVs of HACAD recorded in the potential range of 0 to +1.5 V
show three independent signals in acidic media and two peaks at neutral and under
alkaline conditions. The appearance of peak a* is attributed to the oxidation of OH
group at position 4 formed as a result of possible acid hydrolysis of the acetate group
[74]. The number of hydrolyzed species decreases with increasing pH of the medium,
and consequently the peak current gets decreases. The reason for different redox
potentials of OH groups at position 1 and 5 may be the different extent of their
hydrogen bond formation. Peak a* of HACAD with a half peak width (W1/2) value of
65 mV using equation W1/2 = 3.52RT/nF indicates the loss of one electron in this
oxidation step. The same calculations for the processes corresponding to peak 1a and
2a show that both processes are in fact one electron processes [75]. The number of
electrons and protons involved in oxidation of the selected anthracenediones can be
seen in Table 4.21. Ep-pH plots shown in Fig. 4.64B were used for the determination
of protons accompanying electron transfer processes. The slope of plots of Ep versus
pH of a*, 1a and 2a shows the involvement of a single proton transfer during the
electron transfer reactions [5]. The presence of a third peak under strongly acidic
conditions in case of HAD (see Fig. 4.65A) offers evidence of hydrolysis of the
acetate group. A decrease in W1/2 values and an increase in peak current suggests the
merging of two peaks (a* and 1a) of HAD under basic pH conditions. The acid
dissociation constants, pKa, of these compounds listed in Table 4.21 were evaluated
from the intersection of the two linear segments of Ep versus pH plots. The DPVs of
HOAD show two peaks (1a and 2a) in the entire pH range studied (Fig. 4.65B). The
absence of the third peak is presumably related to the two hydroxyl groups at position
Chapter 4 Results & Discussion
Page 134
1 and 4 being in the same chemical environment, i.e., both can exhibit hydrogen
bonding to the same extent (Scheme 4.19). On the basis of half peak width values, it
can be concluded that peak 1a corresponds to two electrons and 2a to one electron
transfer mechanism.
Fig. 4.65 DPV’s of (A) HAD and (B) HOAD in media of different pH at 5 mVs-1 (C)
plot of peak potentials of HOAD against pH
O
OOH O
OH
OH
Scheme 4.19 Chemical structure of HOAD showing hydrogen bonding
Chapter 4 Results & Discussion
Page 135
Table 4.21 Values of half peak width, slope of Ep – pH plots, number of electrons
and protons and LOD and LOQ
Compounds Peaks Number of electrons
Number of protons
pKa LOD (mmolL-1)
LOQ (mmolL-1)
HACAD a* 1 1
0.9 0.05 0.16 1a 1 1 2a 1 1
HOAD 1a 2 2 9.2 0.09 0.31 2a 1 1
HAD
a* 1 1
8.9 0.08 0.27 1a 1 at pH ≤ 4.0
and 2 at pH ≥ 5.0
2
2a 1 1
4.6.5 Square wave voltammetry
-0.2 -0.1 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8
0
1
2
3
4
5
6
0.02 0.04 0.06 0.08 0.100.5
1.0
1.5
2.0
2.5
3.0
3.5
Slope = 5.09×10-5
SD = 8.27×10-7
I p/µA
Concentration (mmol/L)
a 0.022 mMb 0.031 mMc 0.045 mMd 0.078 mMe 0.112 mMf 0.278 mM
I/µA
E / V vs. Ag/AgCl
f
aA
-0.2 0.0 0.2 0.4 0.6 0.8123456789
0.00 0.04 0.08 0.12 0.16 0.200
1
2
3
4
Slope = 3.167E-5
SD = 8.49E-7
I p/µA
Concentration (mmol/L)
E / V vs. Ag/AgCl
a 0.0122 mMb 0.0183 mMc 0.0274 mMd 0.0617 mMe 0.1112 mMf 0.1667 mM
I/µA
a
f
B
Fig. 4.66 SWVs of different concentrations of (A) HACAD and (B) HAD obtained at
100 mVs-1 in a medium of pH 10.0. Insets show plots of peak current versus
concentration
Square wave voltammetry has a very low detection limit due to minor
contributions of the capacitive charging current [2]. The problems such as higher
analyte consumption and poisoning of the electrode can be minimized by the
application of SWV [76]. Based on these considerations, SWV was employed for the
determination of the limit of detection and the limit of quantification of the selected
Chapter 4 Results & Discussion
Page 136
anthracenediones HACAD, HAD, and HOAD. An expected increase in peak current
with concentration was observed for all three compounds (Fig. 4.66 and Fig. 4.67).
-0.1 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.823456789
1011
0.0 0.1 0.2 0.3 0.4 0.5 0.60
1
2
3
4
5
6
7
Slope = 2.89×10-5
SD = 8.9×10-7
I p/µA
Concentration (mmol/L)f
I/µA
E / V vs. Ag/AgCl
a 0.6667 mMb 0.5556 mMc 0.3215 mMd 0.0872 mMe 0.0491 mMf 0.0271 mM a
Fig. 4.67 Square wave voltammograms of different concentrations of HOAD at 100
mVs-1 in a medium buffered at pH 10.0. Inset shows plot of peak current versus
concentration
The plot of Ip versus concentration was used for calculation of LOD and LOQ
according to equations (4.2) and (4.3). The LOD and LOQ values of anthracenediones
listed in Table 4.21 are comparable to those of closely related compounds [2]. The
ability of SWV to record forward and backward current components of the total
current in a single scan makes it valuable among all other electrochemical methods.
The same direction of forward and backward current components of peak 1a and 2a of
the three anthracenediones shown in Fig. 6.68 indicates the irreversible nature of their
redox processes, thus, providing useful insights about the redox mechanisms of the
selected compounds.
Chapter 4 Results & Discussion
Page 137
-0.2 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.60123456789
10
Ib
If
I/µA
E / V vs. Ag/AgCl
a It
b If
c Ib
It
Fig. 4.68 SWVs of 0.6 mM HACAD showing forward, backward and total current
4.6.6 Electronic absorption spectroscopy
Anthracenediones are coloring agents and their electronic states are assigned to So,
Tn,π*, Tπ,π*, Sn,π* and Sπ,π* on the basis of prior theoretical and experimental studies [77].
According to the literature, the UV-visible spectrum of simple anthracenedione
exhibits four signals at 207, 252, 272 and 326 nm. The signal at 207 nm corresponds
to the n − σ∗ transition of the carbonyl group while the later three signals are assigned
to π − π∗ transitions of benzenoid and quinonoid systems.
250 300 350 400 450 500 550 600 650
0.0
0.3
0.6
0.9
1.2
1.5
1.8
c
b
Abso
rban
ce
wavelength/nm
a HACADb HADc HOAD
a
Fig. 4.69 UV-Vis spectra of 0.02 mM solutions of the selected
anthracenediones in pH 7.0
Chapter 4 Results & Discussion
Page 138
The optical characteristics of substituted ACDs show differences from simple
anthracenedione presumbaly due to the presence of substituents, formation of
hydrogen bonds, and other inter-molecular interactions. Unlike simple
anthracenedione, the UV-visible spectra of HACAD (Fig. 4.69) lacks the signal
corresponding to a -C=O centered n − σ∗ transition, presumably due to the -C=O
group being involved in hydrogen bonding with OH groups at positions 1 and 4. The
peaks appearing around 226, 250 and 281 nm are assigned to π − π∗ transitions of
benzenoid and quinonoid systems [59, 78]. The broad band at 435 nm which is absent
in the absorption spectrum of simple anthracenedione, appears in the spectra of the
three ACDs studied here due to charge transfer (CT) mechanism [59] of the electron
donating effect of the hydroxyl groups attached at the aromatic rings. The difference
in absorption spectra of HAD and HOAD is due to presence of a third -OH group at
position 4 having the ability of forming another ring. The band corresponding to the
CT exhibits a bathochromic shift due to the electron donating effect of the -OH group
compared to the acetate group located at the same position. An observation of the
absorption spectra of the selected ACDs recorded at a different pH (see Fig. 4.70)
reveals that the absorption band due to CT mechanism is strongly affected by pH as
compared to the other signals.
Chapter 4 Results & Discussion
Page 139
250 300 350 400 450 500 550 600 650 700-0.2
0.0
0.2
0.4
0.6
0.8
1.0
1.2
Isosbestic Point
Abso
rban
ce
wavelength/nm
base linea pH3b pH4c pH6d pH7e pH8f pH9g pH10h pH11i pH12
a
i
475 nm
A
250 300 350 400 450 500 550 600 650 700
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
B
Isosbestic Point
Abso
rban
ce
wavelength/nm
Base line a pH3b pH4c pH5d pH6e pH7f pH8g pH9h pH10i pH11j pH12
500 nm
a
j
Fig. 4.70 Absorption spectra of 0.02 mM solutions of (A) HACAD and (B)
HAD in media of pH 3-12
2 3 4 5 6 7 8 9 10 11 12 13 14 15 160.05
0.10
0.15
0.20
0.25
0.30
4 6 8 10 12
-4
0
4
8
Deriv
ativ
e/10
-2
pH
a b
pKa = 8.9
a λ = 462nmb λ = 489nm
Abso
rban
ce
pH
A
2 3 4 5 6 7 8 9 10 11 12
0.00
0.04
0.08
0.12
0.16 a λ = 490nmb λ = 554nm
2 4 6 8 10 12-0.03-0.02-0.010.000.010.020.030.04
Deriv
ativ
e
pH
a b
Abso
rban
ce
pH
pKa = 9.2
B
Fig. 4.71 Absorbance of 20 µM (A) HAD and (B) HOAD as a function of pH
The band corresponding to the CT band of the orange colored solution appears
in the range of 375-480 nm in media of pH ≤ 9 and is shifted to 475-610 nm at pH > 9
due to the expected deprotonation of HACAD. The isosbestic point suggests that
HACAD and its deprotonated form exist in equilibrium with the same value of molar
extinction coefficient. An isosbestic point also appears in the electronic absorption
spectra of the other two ACDs. The absorbance - pH plots shown in Fig. 4.71 were
used to evaluate the pKa of theses compounds and they are identical to those obtained
from the DPV studies. From the slopes of the plots of absorbance versus
concentration, the molar exinction coeffients were obtained and are listed in Table
4.22.
Chapter 4 Results & Discussion
Page 140
Table 4.22 Molar extinction coefficient and pKa values of the analytes
Compounds pKa Wavelength/
nm Ɛ/104
Lmol-1cm-1
HACAD 9.4 437 1.19
HOAD 9.2 233 491
2.79 0.96
HAD 8.9 266 489
2.35 0.63
4.6.7 Redox Mechanism
On the basis of results obtained from the electrochemical experiments conducted at
different pH, redox mechanistic pathways of the ACDs are proposed. The appearance
of a third peak, a*, in the DPV of HACAD at pH ≤ 6.0 suggests hydrolysis of the
acetate group [74]. Based on the peak-width-at-half-peak current and slope value of
the plot of peak potential versus pH, a CE mechanism is proposed (see Scheme 4.20).
Peak 1a and 2a are assigned to the oxidation of -OH at position 1 and 5. In media of
pH ≥ 7.0, the peak related to the hydrolyzed form of the analyte does not appear as
expected. The proposed redox mechanism under these conditions is presented in
Scheme 4.21.
Chapter 4 Results & Discussion
Page 141
Scheme 4.20 Proposed redox mechanism of HACAD in media having pH ≤ 6.0
Like HACAD, the oxidation of HAD exhibits three peaks under acidic and
two under basic conditions. The peak a* appearing in strongly acidic conditions could
not emerge in solutions of pH ≥ 5.0. The signals a* and 1a merged with a rise in pH.
The redox response of HAD shows features similar to those of HACAD in highly
acidic media as shown in Scheme 4.22. At pH ≤ 4.0, the oxidation corresponding to
peaks a*, 1a and 2a indicates the occurance of an oxidation involving one electron
and one proton in each of the steps. At pH higher than 4.0, peak 1a represents the
oxidation process to occur by the loss of 2e, 2H+ resulting in the formation of a
radical [79] that may get converted to 5-hydroxyanthracene-1,4,9,10-tetraone [80] or
adsorb at the GCE as shown in Scheme 4.23. The redox pathways for HOAD is
summarized in Scheme 4.24.
O
OOH
OH
OC
CH3
O
H
H2O
O
OOH
OH
OH
+ CH3COOH
1e1H 1e1H
O
OO
OH
OC
CH3
O
O
OOH
OH
O
Peak 2a Peak a*
O
OOH
O
OC
CH3
O
Peak 1a 1H1e
12
3456
78
10
9
AdsorptionDimerization
Adsorption
Dimerization
Adsorption Dimerization
Chapter 4 Results & Discussion
Page 142
Scheme 4.21 Proposed redox mechanism of HACAD in media having pH ≥ 7.0
Scheme 4.22 Proposed redox mechanism of HAD in media having pH ≤ 4.0
O
OOH
OH
OAc
1e
1H
O
OO
OH
OAc Peak 2a
O
OOH
O
OAcPeak 1a
1e
1H
12
345
6
78
9
10
Adsorption Dimerization Adsorption Dimerization
O
OOH
OH
OH
O
OO
OH
OH
O
OOH
OH
O
Peak 2a
O
OOH
O
OH
Peak 1a 1H
1e
12
3456
78
10
9
AdsorptionDimerization
Adsorption Dimerization
Adsorption Dimerization
Peak a*
1H
1H
1e
1e
Chapter 4 Results & Discussion
Page 143
Scheme 4.23 Suggested redox mechanism of HAD in media having pH ≥ 5.0
O
OOH
OH
OH
1e1H
2e
2H
O
OO
OH
OH
O
OOH
O
O
Peak 2a
Peak 1a1
2
3456
78
10
9
AdsorptionDimerization
O
OOH
O
O
O
OOH
O
O
GCE
GCE
CH2CH2COCH3CH2CH2COCH3
CH2CH2COCH3
CH2CH2COCH3
CH2CH2COCH3
Scheme 4.24 Proposed redox mechanism of HOAD in pH 3-12
O
OOH
OH
OH
1e1H
2e
2H
O
OO
OH
OH
O
OOH
O
O
Peak 2a
Peak 1a1
2
3456
78
10
9
AdsorptionDimerization
O
OOH
O
O
O
OOH
O
O
GCE
GCE
Chapter 4 Results & Discussion
Page 144
Our electrochemical results reveal that the studied compounds undergo a two-
step reduction. On the basis of peak separation, the ratio of Ipa and Ipc obtained by CV,
and forward and backward current components of the total current from SWV, it can
be concluded that the 1st reduction process corresponding to 1c is reversible and that
the 2nd reduction wave related to 2c is quasi-reversible. The results of the peak width
at half peak current and plots of Epc versus pH reveal that the two step reduction
processes in these compounds occur by the participation of 1e- and 1H+ in each step.
The reduction processes are accompanied by the formation of a semiquinone radical
as an intermediate and a hydroquinone as the fully reduced species. The proposed
mechanism is shown in Scheme 4.25.
Scheme 4.25 Proposed redox mechanism corresponding to peak 1c and 2c of
all the selected compounds
O
O
-1e
OH
O
OH
OH
-1H -1H
-1e
Chapter 4 Results & Discussion
Page 145
4.7 Biological activity, pH dependent redox behavior and
UV–Vis spectroscopic studies of naphthalene
derivatives
4.7.1 Electrochemical studies
As antioxidant activity is related to the electron donating ability of an anti-oxidant
[81], so, cyclic voltammetry was carried out to investigate the electron transfer
behavior of the selected naphthalene derivatives. Starting scanning the potential from
zero to positive potential values, two oxidation peaks of DMN at 0.53 and 1.07 V
were observed in the forward scan in a solution of pH 7.0 as shown in Fig. 4.72A. In
the reverse scan, a small cathodic peak at 0.94 V corresponding to the reduction of the
oxidation product formed at 1.07 V was noticed. The unequal current intensity of both
peaks suggests the redox process to occur in a quasi-reversible manner. Two cathodic
peaks corresponding to the reduction of DMN appeared at -0.74 and -1.0 V.
Differential pulse voltammetric technique was used for recording the voltammetric
behavior of DMN from 0 V toward negative potential values. The appearance of both
the reduction peaks shown in Fig. 4.72B excluded the possibility of these signals to be
due to the reduction of DMN oxidation products. The comparison of cyclic
voltammograms in pH 7.0 and 10.0 revealed the reduction peaks of DMN to depend
strongly on the pH of the medium. The shift of peak position with change in pH
indicates the participation of proton during electron capture (reduction) process [34].
In contrast to reduction, Fig. 4.72A demonstrates that the position of anodic peaks of
DMN remains unchanged with change in pH of the medium. Thus, the oxidation of
DMN occurs by the loss of electron without involvement of proton. NDA registered
two pH dependent oxidation peaks and one reduction signal as shown in Fig. 4.73A
and B. The absence of peaks corresponding to oxidation and reduction signals of
NDA indicates the irreversibility of its redox processes.
Chapter 4 Results & Discussion
Page 146
Fig. 4.72 (A) CVs at 100 mVs-1 (B) DPVs showing reduction and (C) oxidation of
1 mM solution of DMN at 5 mVs-1
Differential pulse voltammograms were recorded for both compounds in a
wide range pH range in order to ensure the involvement of protons during electron
transfer processes. Differential pulse voltammograms of the oxidation peak(s) of
DMN and NDA are presented in Fig. 4.72C and Fig. 4.73C. It can be seen that there is
no shift in peak potential by changing pH of the medium in case of DMN, thus,
evidencing the electron abstraction from DMN to occur without the involvement of
protons. While, the oxidation peaks of NDA change their position to less positive
potential values with increase in pH, thus demonstrating protons coupled electron
transfer reactions and decrease in ionization potential with change in medium from
acidic to neutral and basic conditions [34, 81]. The oxidation of NDA occurs at lower
potential than DMN; hence, NDA acts as a stronger reducing agent (antioxidant
Chapter 4 Results & Discussion
Page 147
agent). This electrochemical finding supports the biological antioxidant results.
Literature survey reveals that antioxidants exert their role either by the transfer of
electron or hydrogen [81]. Based on this report, the pH dependent oxidation of NDA
indicates its antioxidant role to be exerted both by the donation of electron and proton.
In contrast to it, the pH independent oxidation of DMN suggests no involvement of
protons in its mechanistic role as antioxidant. Although the oxidation potential of
NDA is greater than the widely used natural antioxidant, ascorbic acid, yet it is
capable of donating two electrons (confirmed by the appearance of two peaks) as
compared to one electron donating ability of ascorbic acid [82].
Fig. 4.73 (A,B) CVs of 1 mM solution of NDA obtained at 100 mVs-1 (C) DPVs at
5 mVs-1 in different pH media and (D) plots of peak potentials vs. pH
Chapter 4 Results & Discussion
Page 148
OCH3 OCH3
Electrochemical reaction
-1e-
OCH3 OCH3
Chemical reaction
OCH3 OCH3
H+
1,8-dimethoxynaphthalene
OCH3 OCH3
2
OCH3 OCH3
OCH3OCH3
-1e-
OCH3 OCH3
OCH3 OCH3
OCH3 OCH3 OCH3 OCH3 OCH3 OCH3
+1e-
+1H+
+1e-
+1H+
Scheme 4.26 Proposed oxidation and reduction mechanism of DMN
On the basis of CV and DPV results, redox mechanistic pathways of
naphthalene derivatives were proposed. The number of electrons was determined
from half peak width values of differential pulse voltammograms. The protons
involved in electrochemical reactions were evaluated from the shift in potential of the
signals with change in pH of the media. The DPVs of DMN shown in Fig. 4.72C
demonstrates no shift in the position of oxidation peaks with increase in pH. Thus,
DMN can oxidize by the lose electron without the transfer of proton. However, the
absence of corresponding cathodic peak of oxidation signal of DMN at 0.53 V shown
in Fig. 4.72A suggests an EC reaction i.e. electron transfer reaction followed by a
chemical reaction [83]. The overall suggested oxidation mechanism of DMN has been
presented in Scheme 4.26. In contrast to oxidation, a regular shift in the reduction
peaks of DMN by change in pH indicates protons coupled electron transfer reactions
as displayed in Scheme 4.26. Similarly using CV and DPV results, the redox
mechanism of NDA shown in Scheme 4.27 was also proposed. For the 1st oxidation
peak of NDA appearing at lower potential, the slope value of Ep - pH plot as shown
Fig. 4.73D is 39.9 mV per pH unit which indicates the involvement of 0.67 protons
per molecule. Thus, the loss of two protons is suggested to occur per three molecules
of the analyte. The involvement of a single electron per molecule was determined
Chapter 4 Results & Discussion
Page 149
from the average width at half peak height of the peaks obtained in different pH
media. The Ep - pH plot (Fig. 4.73D) of the 2nd oxidation peak of NDA with a slope
value of 55.7mV/pH suggests the loss of one electron accompanied with one proton
[32].
OH
OH
O
O
OH
OH
O
O
OH
OH
O
O
OH
OH
O
O
Reduction Peak+1e- , +1H+
GCE
COO-
COO-
COO-
COO-
-OOC
-OOC-OOC
-OOC
1st Oxidation Peak-3e- , -2H+
2nd O
xi dat ion Peak
-1e - , -1H+
COO-
COO-
-OOC
C-OOC
-OOC
O
O
Scheme 4.27 Proposed reduction and oxidation mechanism of NDA
Chapter 4 Results & Discussion
Page 150
4.7.2 Electronic absorption spectroscopy
250 300 350 400 450 5000.0
0.2
0.4
0.6
0.8
Abso
rban
ce
Wavelength (nm)
10µM NDA 20µM DMN
Fig. 4.74 Electronic absorption spectra of 10 μM NDA and 22 μM DMN
As the ease of electron abstraction (oxidation) is related to facile excitation/transition,
so UV–Vis spectroscopy of the selected naphthalene derivatives was carried out. UV–
Vis spectroscopic investigations were done in different pH in order to know whether
the potential antioxidants and anti-diabetic agents maintain their integrity in acidic,
neutral and basic media. Fig. 4.74 displays the UV–Visible spectra of NDA and
DMN. The spectrum of NDA shows an intense peak at 241 nm and a less intense
broad band having two peaks at 341 nm and 325 nm. In case of DMN, an intense
signal is observed at 227 nm while a broad band is observed in the wavelength range
of 287–330 nm. Absorption spectrum of DMN is quite similar to that of naphthalene
but both the bands are at longer wavelengths. The shift of peaks to longer wavelength
is due to the electron donating nature of -OCH3 groups making the transitions more
favorable. Literature survey of the absorbance spectroscopy shows that intense peaks
of both compounds are due to π-π* transition of the naphthalene moiety [84]. A clear
bathochromic shift in maximum absorbance of NDA (as compared to DMN) can be
attributed to the interactions of polar solvent with the polar part of NDA to form
hydrogen bonds, thus reducing the energy of electronic transitions [85]. Less intense
bands at longer wavelengths are due to n-π* transitions of the substituents [85]. The
band at 241 and 227 nm were used to calculate the molar extinction coefficient of
NDA and DMN with values 9.64 × 104 M-1cm-1 and 2.69 × 104 M-1cm-1 respectively.
Chapter 4 Results & Discussion
Page 151
The absorption intensity of NDA is significantly higher than DMN due to its
obvious extended chromophore. The absorption peak of NDA is pH dependent in
correspondence to its pH dependent electrochemical oxidation. While the absorption
signal of DMN shows pH independent behavior like its electrochemical response.
Thus, UV–Vis spectroscopy supports electrochemical results.
200 250 300 350 4000.0
0.2
0.4
0.6
0.8
Abso
rban
ce
Wavelength (nm)
pH-2.0 pH-4.0 pH-6.0 pH-7.0 pH-8.0 pH-10.0 pH-12.0
240 280 320 360 4000.00
0.15
0.30
0.45
0.60
0.75
Abs
orba
nce
Wavelength (nm)
pH-3.0 pH-5.0 pH-7.0 pH-9.0 pH-11.0
Fig. 4.75 Electronic absorption spectra of 15 µM NDA (A) and 30 µM DMN (B)
recorded in different pH media using 50% ethanol and 50% BR buffer
Fig. 4.75shows the spectra of NDA and DMN in different pH media. The
intense peak of NDA exhibits hypsochromic shift with increase in pH up to 7.0. At
pH higher than 7.0, there is no appreciable change in absorbance as well as in peak
position. Thus, pH dependent UV–Vis spectroscopic study suggests pH 7.0 as acid-
base dissociation constant, pKa of NDA. In case of DMN, spectra are almost pH
independent. This is consistent with its chemical structure where no ionizable proton
is available which could result in the shift of λmax. No change in the spectra was
observed when the compounds were exposed to sunlight for two hours. These
investigations conclude that NDA and DMN are stable toward pH and sunlight.
4.7.3 Anti-hyperglycemic activity
The selected naphthalene derivatives NDA and DMN were administered to alloxan-
induced diabetic adult male Sprague–Dawley rats to test if they possess anti-diabetic
properties. In general all animals remained healthy and active and did not show any
signs of lethargy or inactivity. Compared to glibenclamide which is a known anti-
Chapter 4 Results & Discussion
Page 152
diabetic agent and is on the scene since 1970s [86], application of compound NDA
led to a significant increase in plasma glucose levels (P < 0.001) with a concomitant
significant increase in plasma cholesterol and triglyceride concentrations (P < 0.001
for both).
1.0 1.2 1.4 1.6 1.8 2.010121416182022
B
A
B
0 2 4 6
100
200
300
400
500
600
**
Gluc
ose
(Con
cent
ratio
n m
g/dl
)
Time (hr)
Negative Control Positive Control NDA DMN
**A
172
343 340
394
0
100
200
300
400
500
600
Chol
estro
lCo
ncen
tratio
n (m
g/dl
)
Negative Control Positive Control NDA DMN
* ** *
B
0
100
200
300
400
500
* *
Trig
lyce
rides
(Con
c. m
g/dl
)
Negative Control Positive Control NDA DMN * *
C
Fig. 4.76 Glucose (A), cholesterol (B) and triglycerides (C) concentrations following
treatment with naphthalene derivates NDA and DMN; control groups were treated
with glibenclamide, an antidiabetic drug, whereas negative control was treated with
Alloxan Monohydrate
In contrast the administration of compound DMN to diabetic rats led to a
significant lowering of plasma glucose concentration (P < 0.001). This decrease was
almost similar to glibenclamide (Fig. 4.76A). Furthermore, this compound on the
contrary led to a significant elevation of cholesterol concentration (P < 0.001; Fig.
4.76B) but caused significant reduction in triglyceride concentration (P < 0.030; Fig.
4.76C).
Our experiments demonstrate that it is the compound DMN that possess anti-diabetic
potential. Presently, compound DMN appears to be similar to glibenclamide as
regards glucose lowering. Glibenclamide, which is also known as glyburide is an anti-
diabetic drug that belongs to the sulphonylurea group of antihyperglycemic agents.
One of its major side-effects is cholestasis jaundice [87]. Cholestasis occurs due to
Chapter 4 Results & Discussion
Page 153
excessive cholesterol. The present results also demonstrate that compound DMN
cause elevation of plasma cholesterol similar to that caused by glibenclamide.
Data regarding the effect of anti-diabetic drugs like glibenclamide on
cholesterol and triglycerides is scarce, although a few studies do indicate that
glibenclamide administration improves postprandial hypertriglyceridemia acutely by
reducing postprandial triglycerides of intestinal origin [88]. Since lipid metabolism is
affected by anti-diabetic drugs [89, 90], it is not clear at present as to why and how
our compound could have altered lipid metabolism but nonetheless raise the need for
evaluation of these compounds in a greater depth to justify physiological alteration of
bodily lipids. Moreover, until then it is difficult to state that a compound that may
possess glucose lowering properties can in actual act as an anti-diabetic agent.
4.7.4 Cytotoxicity analysis on HeLa
The compounds were tested for their cytotoxic effect against HeLa cell line
(Fig. 4.77 and Table 4.23). The cells were treated with different dilutions of the
compounds ranging from 0 to 150 µM for 24 h, and then submitted to the MTT assay
and percent viabilities were measured relative to the NTC sample. Data showed that
NDA was not toxic to the cells up to concentrations of 100 µM and showed slight
toxicity with percent viability of 84.4 ± 3.3 at 150 µM, hence, doses under this
concentration i.e., 10, 50 and 100 µM were used for the biological antioxidant activity
investigations. Whereas, DMN significantly inhibited the cell proliferation in a dose
dependent manner with percent viability of 103.9 ± 4.1, 75.4 ± 3.1, 58.0 ± 2.6 and
34.6 ± 4.2 at 10, 50, 100 and 150 µM, respectively. An IC50 of 132.92 µM was
calculated for DMN.
Chapter 4 Results & Discussion
Page 154
Fig. 4.77 Percent viability of HeLa cells treated with different concentrations
(10, 50, 100 and 150 µM) of NDA and DMN.
Cultures were exposed to the compounds for 24 h and relative % viabilities
(mean ± SD) were measured using MTT Assay. *(P ˂ 0.01) and **(P ˂ 0.001)
significantly different from NTC (two tailed t-test)
Table 4.23 Percent viability of HeLa cells following treatment with NDA and DMN
No. Treatment Percent viability (mean ± SD)
1. NTC 100.0 ± 3.1
2. Abs. ethanol 99.2 ± 3.4
3. NDA
i. 10 µM 98.2 ± 0.8
ii. 50 µM 98.6 ± 2.7
iii. 100 µM 102.1 ± 4.6
iv. 150 µM 84.4 ± 3.3
4. DMN
i. 10 µM 103.9 ± 4.1
ii. 50 µM 75.4 ± 3.1
iii. 100 µM 58.0 ± 2.6
iv. 150 µM 34.6 ± 4.2
Chapter 4 Results & Discussion
Page 155
4.7.5 Biological antioxidant activity
To investigate the biological antioxidant activity of naphthalene derivatives, cells
were cultured with and without addition of the compounds for 24 h followed by the
induction of oxidative stress by adding 100 µ M Fe++ solution (as FeSO4) for 1 h.
MDA (µ M) levels were calculated relative to the NTC sample. The oxidative
treatment resulted in at least 2-fold increases in MDA levels compared with control
cells. The average MDA level of NTC was 2.59 ± 0.10 µM and for Fe++ alone treated
cells the level was 6.13 ± 0.32 µM. The increase was highly significant (P 6 0.001).
For NDA as shown in Fig. 4.78A, protection against ROS induced oxidative stress
was obtained and the effect was dose dependent. A significant (P < 0.001) decrease in
MDA levels was observed at the concentrations of 50 and 100 µM as compared to
Fe++ alone treated cells. For DMN no significant decrease in MDA levels was
observed at all the concentrations tested (Fig. 4.78B).
Fig. 4.78 MDA (µM, mean ± SD) levels relative to the NTC sample in (A) NDA (B)
DMN treated HeLa cell line
Cells were exposed to the different concentrations (10, 50 and 100 µM) of the
compounds for 24 hours. Oxidative stress was induced by addition of Fe++ (100 μM)
for 1 h and TBARs were measured relative to the NTC sample. Absolute ethanol (AE)
treated, Fe++ alone and Fe++ with absolute ethanol (AE + Fe++) treated samples were
also included. *significantly (P ˂ 0.001) different from NTC, **significantly
(P ˂ 0.001) different from Fe++ alone treated sample (two tailed t-test)
To our knowledge, this is the first report about the cytotoxicity and biological
antioxidant activity of NDA and DMN. It is important to note that NDA is not toxic to
HeLa cells and it shows good antioxidant activity. Several naphthalene derivatives
Chapter 4 Results & Discussion
Page 156
have been reported to exhibit antioxidant activity by scavenging ROS [91, 92]. The
ROS scavenging activity is higher in compounds containing OH groups attached to
quinone moiety [93] and can be attributed to the donation of electron or hydrogen
radical [94]. To get insights about the proton coupled electron transfer mechanism of
naphthalene derivatives, we carried out detailed electrochemical investigations as
discussed above (section 4.7.1).
Chapter 4 Results & Discussion
Page 157
Conclusions Several Schiff bases and metal complexes were successfully synthesized and
characterized by different analytical techniques. The existence of molecules in
different tautomeric forms was confirmed from FTIR, UV–Vis spectroscopy,
computational, electrochemical and single crystal X-ray studies. The redox and
photometric behavior of Schiff bases were found to depend strongly on the pH of the
medium. Computational studies showed high negative charge density on oxygen
atoms, thus, the ligands characteristics are due to these donor atoms. The Schiff base
HL and its metal complexes were found to decrease the blood glucose, triglyceride
and cholesterol levels. However, the oxovanadium complex of HL showed an
increase in the levels of glucose, triglyceride and cholesterol. All the complexes
exhibited good inhibition against alkaline phosphatase enzyme. The zinc complex was
found as the most potent inhibitor of bacteria/fungi.
4-Hydroxy-5-methoxynaphthalene-1-yl acetate (HMNA) registered robust
electrochemical signatures over the surface of glassy carbon electrode in media of pH
2-12. Its oxidation resulted in the formation of an electroactive product that gave
reduction signal with a counter oxidation peak. The resulting reduced product was
quite similar to vitamin K3 as investigated previously in aprotic media using gold
electrode. DPV and SWV supported the proposed mechanism and established that
acid catalyzed hydrolysis of HMNA could lead to the formation of the same diol as
produced from the reduction of dione in the 2nd CV scan of HMNA. HPLC and
computational results also supported the electrochemical investigations. Furthermore,
UV–vis spectroscopy was employed to determine the photometric response and pKa
of HMNA.
The electrochemical signatures of 1-hydroxy-2(hydroxymethyl)anthracene-
9,10-dione (HAC) and 1,8-dihydroxy-4,5-dinitro anthracenedione (DHDN) in
buffered aqueous ethanol media were found quite different from those obtained in
aprotic solvents. Both the anthraquinones followed independent oxidation and
reduction at a glassy carbon electrode. DHDN was found to oxidize resulting in the
formation of an electroactive product that gives another oxidation signal with
irreversible characteristics. The results revealed two step pH dependent reduction of
DHDN with reversible behavior of the first cathodic peak and quasi-reversible
characteristic of the second cathodic wave. The overall reduction led to the
Chapter 4 Results & Discussion
Page 158
conversion of dione to diol. The appearance of a pH dependent sharp cathodic peak of
HAC revealed its 2e− and 2H+ reduction. A broad anodic peak in acidic conditions
and its splitting in to two sub peaks in alkaline media helped in concluding the
hydroxyl groups of HAC to oxidize in two one electron transfer processes.
Other selected anthracenediones 1,4-dihydroxy-2-(3-hydroxy-3-
(trichloromethyl)pentyl)-8-methoxyanthracene-9,10-dione (HCAQ), 4,8-dihydroxy-
9,10-dioxo-9,10-dihydroanthracen-1-yl acetate (HACAD), 1,4,5-
trihydroxyanthracene-9,10-dione (HAD) and 1,4,5-trihydroxy-2-methyl-3-(3-
oxobutyl)anthracene-9,10-dione (HOAD) were found to oxidize in a pH-dependent
irreversible manner as ensured by the absence of corresponding reduction peak in the
reverse scan of cyclic voltammograms, very low value of rate constant, shift in peak
potential with sweep rate and the same direction of forward and backward current
compnents of square wave voltammogram. Various kinetic parameters of the
compounds were evaluated at different temperatures and employed to know about the
thermodynamics of the redox processes. Thermodynamic parameters revealed the
endothermic nature of the redox processes. The temperature-dependent peak shift was
found more pronounced for the irreversible process as compared to the reversible
process. The pH-dependent results of three electroanalytical techniques helped in the
proposal of a redox mechanism of the compound. HACAD and HAD showed the
same redox behavior in media of more acidic character due to acid hydrolysis of
acetate group as expected. The electronic absorption spectroscopy results showed the
charge transfer mechanism to depend strongly on the pH of the medium. The
differene in UV-Vis spectra of the the selected compounds from that of parent
anthracenedione was due to the possible hydrogen bonding of hydroxyl groups
present at different positions. The values of apparent acid dissociation constant
determined by electronic absorption spectroscopy, DPV and drastic change in solution
color were found to be in very good agreement.
The redox and electronic absorption spectroscopic behavior of biologically
important naphthalene derivatives, naphthalene-2,3-dicarboxylic acid (NDA) and 1,8-
dimethoxynaphthalene (DMN) was found to depend on the pH of the media. UV–Vis
spectroscopic results evidenced NDA to have a pKa of 7.0. DMN showed higher anti-
diabetic activity than NDA. The antioxidant studies revealed NDA as a better
antioxidant than DMN. Both the electrochemical and biological antioxidant results
were found in good agreement and revealed NDA as a stronger antioxidant than
Chapter 4 Results & Discussion
Page 159
DMN. Since many anti-diabetic drugs also possess antioxidant properties through
which they actually cure the disease, it is very important to assess the status of
antioxidants in further studies with these compounds so that the anti-diabetic status of
naphthalene derivatives could be clearly demonstrated.
Chapter 4 Results & Discussion
Page 160
References
[1] Karabıyık H, Ocak-Iskeleli N, Petek H, Albayrak Ç, Agar E, An intermediate
structure trapped in solid-state tautomerization process of (E)-4-[(4-
chlorophenylimino) methyl] benzene-1,2,3-triol, Journal of Molecular Structure 2008;
873: 130-136.
[2] Ahmad K, Shah AH, Adhikari B, Rana UA, Vijayaratnam C, Muhammad N, pH-
dependent redox mechanism and evaluation of kinetic and thermodynamic parameters
of a novel anthraquinone, RSC Advances 2014; 4: 31657-31665.
[3] Munir A, Shah A, Shah AH, Rana UA, Adhikari B, Khan SB, Detailed
electrochemistry of the environmental toxin ethylene diamine, Journal of The
Electrochemical Society 2014; 161: H370-H374.
[4] Munir S, Shah A, Rana UA, Shakir I, Shah SM, Probing of the pH-dependent redox
mechanism of a biologically active compound, 5,8-dihydroxynaphthalene-1,4-dione,
Australian Journal of Chemistry 2014; 67: 206-212.
[5] Shah AH, Zaid W, Shah A, Rana UA, Hussain H, Ashiq MN, pH dependent
electrochemical characterization, computational studies and evaluation of
thermodynamic, kinetic and analytical parameters of two phenazines. Journal of the
Electrochemical Society 2015; 162: H115-H123.
[6] Latha B, Gunasekaran S, Srinivasan S, Ramkumaar G, Computation and
interpretation of vibrational spectra on the structure of losartan using ab initio and
density functional methods, Spectrochimica Acta Part A: Molecular and Biomolecular
Spectroscopy 2014; 132: 375-386.
[7] Tamas L, Huttova J, Mistrk I, Kogan G, Effect of carboxymethyl chitin-glucan on the
activity of some hydrolytic enzymes in maize plants, Chemical Papers 2002; 56:
326-329.
[8] Kim EE, Wyckoff HW, Structure of alkaline phosphatases, Clinica Chimica Acta
1990; 186: 175-187.
[9] Shephard M, Peake M, Quantitative method for determining serum alkaline
phosphatase isoenzyme activity I. Guanidine hydrochloride: new reagent for
selectively inhibiting major serum isoenzymes of alkaline phosphatase, Journal of
Clinical Pathology 1986; 39: 1025-1030.
[10] Rauf A, Shah A, Abbas S, Rana UA, Khan SU-D, Ali S, Synthesis, spectroscopic
characterization and pH dependent photometric and electrochemical fate of Schiff
Chapter 4 Results & Discussion
Page 161
bases, Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 2015;
138: 58-66.
[11] Shah AH, Shah A, Khan SU-D, Rana UA, Hussain H, Khan SB, Probing the pH
dependent electrochemistry of a novel quinoxaline carboxylic acid derivative at a
glassy carbon electrode, Electrochimica Acta 2014; 147: 121-128.
[12] Shah AH, Shah A, Rana UA, Khan SUD, Hussain H, Khan SB, Redox mechanism
and evaluation of kinetic and thermodynamic parameters of 1, 3‐dioxolo [4, 5‐g]
pyrido [2, 3‐b] quinoxaline using electrochemical techniques, Electroanalysis 2014;
26: 2292-2300.
[13] Hussain S, Hussain Bukhari I, Ali S, Shahzadi S, Shahid M, Shahzad Munawar K.
Synthesis, spectroscopic and thermogravimetric characterization of heterobimetallic
complexes with Sn (IV) and Pd (II); DNA binding, alkaline phosphatase inhibition
and biological activity studies, Journal of Coordination Chemistry 2014: 1-23.
[14] Bhalodia NR, Shukla V, Antibacterial and antifungal activities from leaf extracts of
Cassia fistula l.: An ethnomedicinal plant, Journal of Advanced Pharmaceutical
Technology & Research 2011; 2: 104-109.
[15] Huynh Q, Borgmeyer J, Smith C, Bell L, Shah D, Isolation and characterization of a
30 kDa protein with antifungal activity from leaves of engelmannia pinnatifida,
Biochemical Journal 1996; 316: 723-727.
[16] Shah A, Diculescu VC, Qureshi R, Oliveira-Brett AM, Electrochemical behaviour of
dimethyl-2-oxoglutarate on glassy carbon electrode, Bioelectrochemistry 2010; 77:
145-150.
[17] Shah A, Diculescu VC, Muhammad N, Qureshi R, Ali S, Oliveira‐Brett AM.
Electrochemical Investigation of Na‐Salt of 2‐Methyl‐3‐(4‐nitrophenyl) acrylate on
Glassy Carbon Electrode, Electroanalysis 2010; 22: 121-127.
[18] Shah A, Diculescu VC, Qureshi R, Oliveira-Brett AM, Electrochemical reduction
mechanism of camptothecin at glassy carbon electrode, Bioelectrochemistry 2010; 79:
173-178.
[19] Diculescu VC, Militaru A, Shah A, Qureshi R, Tugulea L, Brett AMO, Redox
mechanism of lumazine at a glassy carbon electrode, Journal of Electroanalytical
Chemistry 2010; 647: 1-7.
[20] Shah A, Nosheen E, Qureshi R, Yasinzai MM, Lunsford SK, Dionysiou DD,
Electrochemical characterization, detoxification and anticancer activity of
Chapter 4 Results & Discussion
Page 162
didodecyldimethylammonium bromide, International Journal of Organic Chemistry
2011; 1: 183-190.
[21] Bard AJ, Faulkner LR, Fundamentals and applications: Electrochemical Methods,
Wiley: New York 2001.
[22] Stevic MC, Ignjatovic LM, Ciric-Marjanovic G, Stanisic SM, Stankovic DM, Zima J.
Voltammetric behaviour and determination of 8-Hydroxyquinoline using a glassy
carbon paste electrode and the theoretical study of its electrochemical oxidation
mechanism, International Journal of Electrochemical Science 2011; 6: 2509-2525.
[23] Rosspeintner A, Koch M, Angulo G, Vauthey E, Spurious observation of the marcus
inverted region in bimolecular photoinduced electron transfer. Journal of the
American Chemical Society 2012; 134: 11396-11399.
[24] Holze R. CMA Brett, AMO Brett, Electrochemistry—Principles, methods and
applications, OXford 1994.
[25] Ramaley L, Krause Jr MS. Theory of square wave voltammetry. Analytical Chemistry
1969; 41: 1362-1365.
[26] Kovacic P, Cooksy AL, Electron transfer as a potential cause of diacetyl toxicity in
popcorn lung disease, reviews of environmental contamination and toxicology:
Springer; 2010: 133-148.
[27] Ewing D, Vandenbelt J, Kamm O, The ultraviolet absorption of vitamins K1, K2, and
some related compounds, Journal of Biological Chemistry 1939; 131: 345-356.
[28] Friedel RA, Orchin M, Ultraviolet spectra of aromatic compounds: John Wiley; 1951.
[29] Marinov M, Minchev S, Stoyanov N, Ivanova G, Spassova M, Enchev V, Synthesis,
spectroscopic characterization and ab initio investigation of thioanalogues of
spirohydantoins, Croatica Chemica Acta 2005; 78: 9-16.
[30] Rahman A, Qureshi R, Kiran M, Ansari FL, Electron affinities, solvation energies and
redox potentials of some chalcones: substituentseffect and correlation with semi-
empirical MO energies, Turkish Journal of Chemistry 2007; 31: 25-34.
[31] D Andrade BW, Datta S, Forrest SR, Djurovich P, Polikarpov E, Thompson ME,
Relationship between the ionization and oxidation potentials of molecular organic
semiconductors, Organic Electronics 2005; 6: 11-20.
[32] Nosheen E, Shah A, Badshah A, Hussain H, Qureshi R, Ali S, Electrochemical
oxidation of hydantoins at glassy carbon electrode, Electrochimica Acta 2012; 80:
108-117.
Chapter 4 Results & Discussion
Page 163
[33] Xavier J, Ortega E, Ferreira J, Bernardes A, Perez-Herranz V, An electrochemical
study of phenol oxidation in acidic medium, International Journal of Electrochemical
Science 2011; 6: 622-636.
[34] Shah A, Ullah A, Rauf A, Rehman Z-u, Shujah S, Shah SM, Detailed electrochemical
probing of a biologically active isoquinoline. Journal of the Electrochemical Society
2013; 160: H597-H603.
[35] Ullah I, Naveed A, Shah A, Badshah A, Khan GS, Nadeem A, High yield synthesis,
detailed spectroscopic characterization and electrochemical fate of novel cationic
surfactants, Journal of Surfactants and Detergents 2014; 17: 243-251.
[36] Garrido EM, Lima JLC, Delerue-Matos CM, Brett AMO, Electrochemical oxidation
of bentazon at a glassy carbon electrode: application to the determination of a
commercial herbicide, Talanta 1998; 46: 1131-1135.
[37] Shah A, Ullah A, Nosheen E, Rana UA, Shakir I, Badshah A, Detailed
electrochemical probing of the ph dependent redox behavior of 1-methoxyphenazine.
Journal of the Electrochemical Society 2013; 160: H765-H769.
[38] Brett CM, Brett AO, Electrochemistry P, Methods and applications. Oxford; 1993.
[39] Swiergiel J, Bouteiller L, Jadżyn J, Compliance of the stokes–einstein model and
breakdown of the stokes–einstein–debye model for a urea-based supramolecular
polymer of high viscosity, Soft Matter 2014; 10: 8457-8463.
[40] Nicholson RS, Shain I, Theory of stationary electrode polarography for a chemical
reaction coupled between two charge transfers, Analytical Chemistry 1965; 37:
178-190.
[41] Reinmuth W, Theory of stationary electrode polarography, Analytical Chemistry
1961; 33: 1793-1794.
[42] Velasco JG, Determination of standard rate constants for electrochemical irreversible
processes from linear sweep voltammograms. Electroanalysis 1997; 9: 880-882.
[43] Randles JEB, A cathode ray polarograph. part ii.—the current-voltage curves,
Transactions of the Faraday Society 1948; 44: 327-338.
[44] Kalanur SS, Seetharamappa J, Katrahalli U, Kandagal PB, Electrochemical studies of
buzepide methiodide and their analytical applications, International Journal of
Electrochemical Science 2008; 3: 711-720.
[45] McMahon BH, Muller JD, Wraight CA, Nienhaus GU, Electron transfer and protein
dynamics in the photosynthetic reaction center, Biophysical Journal 1998; 74: 2567-
2587.
Chapter 4 Results & Discussion
Page 164
[46] Kirchhoff H, Schottler MA, Maurer J, Weis E, Plastocyanin redox kinetics in spinach
chloroplasts: evidence for disequilibrium in the high potential chain, Biochimica et
Biophysica Acta (BBA)-Bioenergetics 2004; 1659: 63-72.
[47] Munir S, Shah A, Rauf A, Badshah A, Lunsford SK, Hussain H, Redox behavior of a
novel menadiol derivative at glassy carbon electrode, Electrochimica Acta 2013; 88:
858-864.
[48] Nadjo L, Saveant J, Linear sweep voltammetry: kinetic control by charge transfer
and/or secondary chemical reactions: I. formal kinetics, Journal of Electroanalytical
Chemistry and Interfacial Electrochemistry 1973; 48: 113-145.
[49] Schmitz J, Van der Linden J, Temperature-dependent determination of the standard
heterogeneous rate constant with cyclic voltammetry, Analytical Chemistry 1982; 54:
1879-1880.
[50] Nicholson RS, Theory and application of cyclic voltammetry for measurement of
electrode reaction kinetics, Analytical Chemistry 1965; 37: 1351-1355.
[51] Zuckerman JJ, Inorganic reactions and methods, cumulative index, part 1: John Wiley
& Sons; 2009.
[52] Monk P, The viologens: synthesis, physicochemical properties and applications of the
salts of 4, 4’-bipyridine, Wiley UK 1998.
[53] Yasin SA, Electrochemical studies and thermodynamic parameters of citrazinic acid
azo dye and its nitro derivatives in dm f-aqueous solutions, Portugaliae
Electrochimica Acta 2006; 24: 23-36.
[54] Laidler KJ, Symbolism and terminology in chemical kinetics. Pure and Applied
Chemistry 1981; 53: 753-771.
[55] Michael AC, Borland LM, John CE, Jones SR, Fast scan cyclic voltammetry of
dopamine and serotonin in mouse brain slices, 2007.
[56] Enache TA, Oliveira-Brett AM, Phenol and para-substituted phenols electrochemical
oxidation pathways, Journal of Electroanalytical Chemistry 2011; 655: 9-16.
[57] Altınoz S, Uyar B, Electrochemical behaviour and voltammetric determination of
rosuvastatin calcium in pharmaceutical preparations using a square-wave
voltammetric method, Analytical Methods 2013; 5: 5709-5716.
[58] aEl Ezaby M, Spectral studies of some hydroxy-derivatives of anthraquinones,
Journal of the Chemical Society B: Physical Organic 1970: 1293-1296.
Chapter 4 Results & Discussion
Page 165
[59] Etaiw SH, Abou Sekkina MM, El-Hefnawey GB, Assar SS, Spectral behaviour and
solvent effects of some aminoanthraquinone dyes, Canadian Journal of Chemistry
1982; 60: 304-307.
[60] Markovic Z, Manojlovic N, Zlatanovic S, Electronic absorption spectra of substituted
anthraquinones and their simulation using ZINDO/S method, Journal of the Serbian
Society for Computational Mechanics, 2008; 2: 73-79.
[61] Diaz AN, Absorption and emission spectroscopy and photochemistry of 1,10-
anthraquinone derivatives: a review, Journal of Photochemistry and Photobiology A:
Chemistry 1990; 53: 141-167.
[62] Capps RN, Vala M, Luminescence studies of quinizarin and daunorubicin,
Photochemistry and Photobiology 1981; 33: 673-682.
[63] Sauer M, Hofkens J, Enderlein J, Handbook of fluorescence spectroscopy and
imaging: from ensemble to single molecules, John Wiley & Sons; 2010.
[64] Doose S, Neuweiler H, Sauer M, A close look at fluorescence quenching of organic
dyes by tryptophan, ChemPhysChem 2005; 6: 2277-2285.
[65] Heinlein T, Knemeyer J-P, Piestert O, Sauer M, Photoinduced electron transfer
between fluorescent dyes and guanosine residues in DNA-hairpins, The Journal of
Physical Chemistry B 2003; 107: 7957-7964.
[66] Brett AMO, Ghica ME, Electrochemical oxidation of quercetin, Electroanalysis 2003;
15: 1745-1750.
[67] Shah AH, Shah A, Khan SU-D, Rana UA, Hussain H, Khan SB, Probing the pH
dependent electrochemistry of a novel quinoxaline carboxylic acid derivative at a
glassy carbon electrode, Electrochimica Acta 2014; 147: 121-128.
[68] Shah AH, Shah A, Rana UA, Khan SU-D, Hussain H, Khan SB, Redox mechanism
and evaluation of kinetic and thermodynamic parameters of 1,3-dioxolo[4,5-
g]pyrido[2,3-b]quinoxaline using electrochemical techniques, Electroanalysis 2014;
26: 2292-2300.
[69] Zhao Y-D, Zhang W-D, Chen H, Luo Q-M. Anodic oxidation of hydrazine at carbon
nanotube powder microelectrode and its detection, Talanta 2002; 58: 529-534.
[70] Fry AJ, Synthetic Organic Electrochemistry: Wiley; 1989.
[71] El Fallah J, Boujana S, Dexpert H, Kiennemann A, Majerus J, Touret O, Redox
processes on pure ceria and on Rh/CeO2 catalyst monitored by X-ray absorption (fast
acquisition mode), The Journal of Physical Chemistry 1994; 98: 5522-5533.
Chapter 4 Results & Discussion
Page 166
[72] Yasin SA, Electrochemical studies and thermodynamic parameters of citrazinic acid
azo dye and its nitro derivatives in DMF-aqueous solutions, Portugaliae
Electrochimica Acta 2006; 24: 23-36.
[73] Parveen N, Shah A, Khan SZ, Khan SU-D, Rana UA, Fathi F, Synthesis,
spectroscopic characterization, ph dependent electrochemistry and computational
studies of piperazinic compounds. Journal of the Electrochemical Society 2015; 162:
H32-H9.
[74] Rodriguez R, Angels Estiarte M, Diez A, Rubiralta M, Colell A, Garcia-Ruiz C,
Conformationally restricted analogues of methionine: synthesis of chiral 3-Amino-5-
methylthio-2-piperidones, Tetrahedron 1996; 52: 7727-7736.
[75] Smith ET, Feinberg B, Redox properties of several bacterial ferredoxins using square
wave voltammetry, Journal of Biological Chemistry 1990; 265: 14371-14376.
[76] Munir S, Shah A, Zafar F, Badshah A, Wang X, Rehman Z-u, Redox behavior of a
derivative of vitamin k at a glassy carbon electrode, Journal of the Electrochemical
Society 2012; 159: G112-G116.
[77] Navas Diaz A, Absorption and emission spectroscopy and photochemistry of 1,10-
anthraquinone derivatives: a review, Journal of Photochemistry and Photobiology A:
Chemistry 1990; 53: 141-167.
[78] Peters RH, Sumner HH, 430 Spectra of anthraquinone derivatives, Journal of the
Chemical Society (Resumed) 1953: 2101-2110.
[79] Si W, Lei W, Zhang Y, Xia M, Wang F, Hao Q, Electrodeposition of graphene oxide
doped poly (3, 4-ethylenedioxythiophene) film and its electrochemical sensing of
catechol and hydroquinone, Electrochimica Acta 2012; 85: 295-301.
[80] Brett AMO, Ghica M-E, Electrochemical oxidation of quercetin, Electroanalysis
2003; 15: 1745-1750.
[81] Prior RL, Wu X, Schaich K, Standardized methods for the determination of
antioxidant capacity and phenolics in foods and dietary supplements, Journal of
Agricultural and Food Chemistry 2005; 53: 4290-4302.
[82] Ngai KS, Tan WT, Zainal Z, Zawawi RM, Zidan M, Voltammetry detection of
ascorbic acid at glassy carbon electrode modified by single-walled carbon
nanotube/zinc oxide, Int J Electrochem Sci 2013; 8: 10557-10567.
[83] Weinhold E, Marquez O, Marquez J, Veratrole anodic oxidation in presence of
perchlorate, Avances en Quimica 2007; 2: 9-14.
Chapter 4 Results & Discussion
Page 167
[84] Berlman I, Handbook of fluorescence spectra of aromatic molecules, Acadamic.
Press, New York 1965.
[85] Selvarani V, Annaraj B, Neelakantan M, Sundaramoorthy S, Velmurugan D.
Synthesis and crystal structure of hydroxyacetophenone Schiff bases containing
propargyl moiety: solvent effects on UV–visible spectra, Spectrochimica Acta Part A:
Molecular and Biomolecular Spectroscopy 2012; 91: 329-337.
[86] Marble A, Glibenclamide, a new sulphonylurea: whither oral hypoglycaemic agents,
Drugs 1971; 1: 109-115.
[87] Gangji AS, Cukierman T, Gerstein HC, Goldsmith CH, Clase CM, A systematic
review and meta-analysis of hypoglycemia and cardiovascular events a comparison of
glyburide with other secretagogues and with insulin, Diabetes Care 2007; 30:
389-394.
[88] Skrapari I, Perrea D, Ioannidis I, Karabina SA, Elisaf M, Tselepis A, Glibenclamide
improves postprandial hypertriglyceridaemia in type 2 diabetic patients by reducing
chylomicrons but not the very low‐density lipoprotein subfraction levels, Diabetic
Medicine 2001; 18: 781-785.
[89] Waysbort J, Regitz G, Chaimowitz D, Tuval M, Nakash I, Brunner D, Effects of
glibenclamide on serum lipids, lipoproteins, thromboxane, beta-thromboglobulin, and
prostacyclin in non-insulin-dependent diabetes mellitus, Clinical Therapeutics 1988;
10: 358-371.
[90] Keidan B, Hsia J, Katz R, Plasma lipids and antidiabetic agents: a brief overview.
British Journal of Diabetes and Vascular Disease 2002; 2: 40-43.
[91] Hamdy NA, Anwar MM, Abu-Zied KM, Awad HM, Synthesis, tumor inhibitory and
antioxidant activity of new polyfunctionally 2-substituted 5,6,7,8-
tetrahydronaphthalene derivatives containing pyridine, thioxopyridine and
pyrazolopyridine moieties, Acta Poloniae Pharmaceutica 2013; 70: 987-1001.
[92] Matsushita Y, Jang I-C, Imai T, Fukushima K, Lee J-M, Park H-R, Antioxidant and
cytotoxic activities of naphthalene derivatives from diospyros kaki, Journal of Wood
Science 2011; 57: 161-165.
[93] Berghot M, Kandeel E, Abdel-Rahman A, Abdel-Motaal M, Synthesis, antioxidant
and cytotoxic activities of novel naphthoqui-none derivatives from 2, 3-dihydro-2, 3-
epoxy-1, 4-naphthoquinone, Medicinal Chemistry 2014; 4: 381-388.
Chapter 4 Results & Discussion
Page 168
[94] Alkan M, Yuksek H, Gürsoy-Kol O, Calapoglu M, Synthesis, acidity and antioxidant
properties of some novel 3,4-disubstituted-4,5-dihydro-1H-1,2,4-triazol-5-one
derivatives, Molecules 2008; 13: 107-121.