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

i

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.

ii

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

iii

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

iv

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

v

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

vi

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

vii

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

viii

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

ix

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


5 mVs-1 showing reduction of HL (A) and (L2Sn) (B) ……………… 72


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


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


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.


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-


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


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,


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


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

http://en.wikipedia.org/wiki/Anthraquinone

http://en.wikipedia.org/wiki/Dye

http://en.wikipedia.org/wiki/Thermoplastic

http://en.wikipedia.org/wiki/Polystyrene

http://en.wikipedia.org/wiki/Synthetic_fiber

http://en.wikipedia.org/wiki/UGT1A8

http://en.wikipedia.org/wiki/Anthraquinone


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.


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,


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.


Page 11

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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

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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)


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)


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:


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)


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)


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


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

https://en.wikipedia.org/wiki/Voltammetry


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


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.


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.


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.

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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,


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).


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.


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.


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


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.


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


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


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].


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).


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


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.


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


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)


Page 46

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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)


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


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


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


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


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


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.


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.


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


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


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.


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


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


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


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


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.


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


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).


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


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


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


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)


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)


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)


5 mVs-1 showing reduction of HL (A) and (L2Sn) (B)


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)


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


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)


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)


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.


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


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).


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


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


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,


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)


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 %)


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


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).


Page 86

Fig. 4.27 Numbering scheme of HL used in computational studies


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


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


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


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


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


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.


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.


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


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).


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


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


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


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.


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


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


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.


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


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].


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)


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


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.


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


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.


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.


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


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.


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


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.


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).


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].


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].


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)


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


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.


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


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].


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


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].


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


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


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


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.


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


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].


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].


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.


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


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


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


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


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.


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


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.


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.


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.


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



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




Peak a*

1H

1H

1e

1e


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


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


O

OOH

O

O

O

OOH

O

O

GCE

GCE


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


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.


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


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


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


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


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.


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-


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


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.


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


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


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).


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


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


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.


Page 160

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