Electrochemical and photochemical oxidation of tubercidin-5'-monophosphate

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Electrochemical and photochemical oxidation of tubercidin-5'-monophosphate
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xii, 153 leaves : ill. ; 28 cm.
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Peterson, Teresa E., 1961-
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Purines   ( lcsh )
Oxidation, Physiological   ( lcsh )
Oxidation   ( lcsh )
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bibliography   ( marcgt )
theses   ( marcgt )
non-fiction   ( marcgt )

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Thesis:
Thesis (Ph. D.)--University of Florida, 1987.
Bibliography:
Includes bibliographical references (leaves 144-152).
Statement of Responsibility:
by Teresa E. Peterson.
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Typescript.
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Vita.

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University of Florida
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ELECTROCHEMICAL AND PHOTOCHEMICAL OXIDATION OF
TUBERCIDIN-5'-MONOPHOSPHATE






BY






TERESA :E. PETERSON


A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN
PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY


UNIVERSITY OF FLORIDA


1987
















ACKNOWLEDGMENTS


The number of people involved directly or indirectly with this

research project are too numerous to mention in this short space.

Everyone in the chemistry department (i.e., faculty, staff and

graduate students) has provided encouragement, support and technical

advice. Thanks are extended to all of them.

Many thanks go to Dr. Anna Brajter-Toth and the members of the

group for their guidance and friendship. I will miss all of them and

I wish them good luck in their future endeavors.

Heartfelt thanks go to my family for their continued

encouragement and especially to my father who never took life too

seriously.

Most of all I would like to thank my husband, Mark, for showing

me that there is life after graduate school.
















TABLE OF CONTENTS


Page

ACKNOWLEDGMENTS.................................................... ii

LIST OF TABLES..................................................... vi

LIST OF FIGURES...................................................vii

ABSTRACT ...........................................................xi

CHAPTERS

1 INTRODUCTION ............................................. 1

2 ELECTROCHEMICAL STUDIES OF BIOLOGICAL OXIDATIONS.........9

3 INTRODUCTION TO METHODS USED IN THE STUDY OF
BIOLOGICAL OXIDATIONS................................... 18

3.1 Cyclic Voltammetry.................................18
3.1.1 Adsorption, Diffusion and Reversibility.....18
3.1.2 Homogeneous Chemical Reactions..............20
3.1.3 Determination of pKa........................22
3.1.4 Estimation of n.............................23
3.2 Determination of n-Values by Coulometry............ 24
3.3 Formation of Oxidation Products by Constant
Potential Electrolysis.............................25
3.4 Chronocoulometry...................................25
3.4.1 Estimation of n.............................25
3.4.2 Determination of Electrode Area.............27

4 EXPERIMENTAL............................................28

4.1 Cyclic Voltammetry................................. 28
4.2 Coulometry and Constant Potential Electrolysis.....31
4.3 Product Analysis...................................33
4.4 Chronocoulometry...................................37
4.5 Photooxidation Reactions...........................38
4.6 Enzymatic Oxidation Studies........................39
4.7 Electrochemical Thermospray Mass Spectrometry
(EC/TSP/MS)........................................ 42










5 CYCLIC VOLTAMMETRY OF TUBERCIDIN-5'-MONOPHOSPHATE.......44

5.1 Oxidation and Reduction Behavior, Reversibility....44
5.2 Effect of pH.......................................47
5.3 Electrode Surface Effects..........................52
5.4 Behavior in Organic Solvents and Solvent
Mixtures........................................... 58

6 DETERMINATION OF n-VALUE OF TUBERCIDIN-5'-
MONOPHOSPHATE...........................................64

6.1 Coulometry......................................... 64
6.2 Peak Current Comparisons...........................65
6.3 Chronocoulometry................................... 69

7 CONSTANT POTENTIAL ELECTROLYSIS OF TUBERCIDIN-5'-
MONOPHOSPHATE........................................... 73

7.1 Introduction.......................................73
7.2 Cyclic Voltammetry.................................80
7.3 UV Spectra......................................... 81
7.4 HPLC............................................... 83

8 ANALYSIS OF TUBERCIDIN-5'-MONOPHOSPHATE
ELECTROCHEMICAL OXIDATION PRODUCTS...................... 85

8.1 Separation by Gel Permeation Liquid
Chromatography..................................... 85
8.2 Analysis of Products with High Pressure Liquid
Chromatography, Cyclic Voltammetry and
Ultraviolet Spectroscopy...........................86
8.3 Analysis of Products with Gas Chromatography
Mass Spectroscopy, Fast Atom Bombardment Mass
Spectrometry, Fourier Transform Infra-Red
Spectroscopy and Nuclear Magnetic Resonance
Spectroscopy...................................... 90
8.3.1 Product 1...................................95
8.3.2 Product 2...................................96
8.3.3 Product 3..............................104
8.4 Summary of Results................................ 105

9 PHOTOOXIDATION OF TUBERCIDIN-5'-MONOPHOSPHATE.......... 110

9.1 Oxidation with Ultraviolet Light..................110
9.2 Mechanistic Studies...............................113

10 ENZYMATIC OXIDATION OF TUBERCIDIN-5'-MONOPHOSPHATE.....121

11 ELECTROCHEMISTRY THERMOSPRAY MASS SPECTROSCOPY
(EC/TSP/MS) OF TUBERCIDIN-5'-MONOPHOSPHATE............. 128











12 SUMMARY AND FUTURE WORK ................................ 133

APPENDIX. .......................................................... 139

REFERENCES........................................................ 144

BIOGRAPHICAL SKETCH ............................................... 153
















LIST OF TABLES


Table Page

6-1 Estimation of the n-Values of TMP from Cyclic
Voltammetric Peak Currents................................ 68

6-2 Validity of Peak Current Comparisons for Estimation
of n-Values............................................... 68

6-3 Estimation of the n-Value of TMP Using Chronocoulometry...70

6-4 Validity of Chronocoulometric Comparison for
Estimation of n-Values....................................71

8-1 Analysis of Products Eluting Under GPLC Peaks A, B
and C by GC/MS and FABMS.................................. 92

8-2 Analysis of Products Eluting Under GPLC Peaks A, B
and C by FTIR and NMR.....................................93

11-1 Analysis of TMP and Its Oxidation Products by
EC/TSP/MS in 0.1M Ammonium Acetate Buffer pH 7...........130
















LIST OF FIGURES


Figure Page

1-1 Structure of tubercidin and some adenosine analogs.........2

1-2 Structures of bases, nucleosides and nucleotides...........5

2-1 Proposed reaction scheme for the electrochemical
oxidation of uric acid at physiological pH................12

2-2 Proposed reaction scheme for the electrochemical
oxidation of adenine...................................... 14

3-1 Homogeneous chemical reactions that accompany
heterogeneous electron transfer processes.................21

4-1 Schematic of a saturated calomel electrode in an
electrochemical cell with analyte solution................29

5-1 Cyclic voltammogram of TMP at a rough PG electrode,
600UM solution in pH 7 phosphate buffer, y=0.5M...........45

5-2 Effect of concentration and scan rate on the oxidation
peak current of TMP in a pH 7 phosphate buffer,
p=0.5M: a) plot of peak current versus concentration
for the oxidation peak of TMP, scan rate 200mV/s;
b) plot of log peak current versus log scan rate for a
100M TMP solution (c.c. = correlation coefficient)....... 46

5-3 Plot of log peak current versus log scan rate for a
998UM TMP solution in pH 7 phosphate buffer, u=0.5M
(c.c. = correlation coefficient)..........................48

5-4 Cyclic voltammetric behavior of TMP at a PG electrode
in different buffers, i=0.5M: a) pH 2.8 phosphate;
b) pH 7 phosphate; c) pH 9.5 phosphate; d) pH 6.6
ammonium acetate..........................................49

5-5 Plot of peak potential versus pH for the oxidation
peak of TMP in phosphate buffers, 4=0.5M. All
concentrations were ca. 300pM in TMP. Scan rate
5mV/s; c.c. = correlation coefficient.....................51











5-6 Spectral behavior of TMP (ca. 150pM) in phosphate
buffers, y=0.5M, of pH a) 2.8-4.6 and b) 6.8-11...........53

5-7 Cyclic voltammograms in a a) pH 7 phosphate buffer,
4=0.5M, at rough PG; b) 600UM TMP solution in pH 7
phosphate buffer at rough PG; c) same as a, at GC;
d) same as b, at GC....................................... 56

5-8 Cyclic voltammograms of a 575UM TMP solution in pH 7
phosphate buffer, y=0.5M, at a glassy carbon electrode:
a) before pretreatment and b) after pretreatment. Scan
rate 200mV/s.............................................. 59

5-9 Cyclic voltammograms in pH 7 phosphate buffer, p=0.5M,
at a GC electrode: a) before pretreatment and b) after
pretreatment.............................................. 60

5-10 Plot of TMP oxidation peak current at PG and GC
versus % DMF added to a 600pM TMP solution in pH 7
phosphate buffer, P=0.5M.................................. 62

7-1 HPLC during the electrolysis of a 600pM TMP solution
in pH 7 phosphate buffer, y=0.5M, at a rough PG
electrode: a) before; b) 25min; c) 1hr; d) 2hr,
e) 3hr, 22min; f) 7hr.....................................74

7-2 UV spectra during the electrolysis of a 100jM TMP
solution in pH 7 phosphate buffer, 4=0.5M, at a
rough PG electrode........................................ 76

7-3 Cyclic voltammograms of a 6000M TMP solution in a
pH 7 phsophate buffer, y=0.5M, during electrolysis
at a rough PG electrode: a) before; b) 2hr, 2min;
c) 7hr, 2min ..............................................77

7-4 GPLC separation of TMP electrolysis products from a
600pM TMP solution in a pH 9.5, u=0.5M, phosphate
buffer after a) 4-12hr and b) 60hr of electrolysis........78

7-5 UV spectra during the electrolysis of a 100UM TMP
solution in pH 7 phosphate buffer, P=0.5M, at an
unroughened PG surface.................................... 82

8-1 HPLC of compounds under GPLC peaks a) A, b) B and c) C
(flow rate 1ml/min).......................................87

8-2 UV spectrum of separated product fractions under GPLC
peak A.................................................... 88


viii










8-3 UV spectrum of separated product fractions under GPLC
peak B .................................................... 89

8-4 UV spectrum of separated product fractions under GPLC
peak C ....................................................91

8-5 Proposed structures for products 1, 2 and 3...............94

8-6 Standard NMR signals for a) inosine, b) guanosine and
c) adenosine-5'-dihydrogen phosphate......................98

8-7 Key to peak assignments for figures 8-6 and 8-8 to
8-10 ...................................................... 99

8-8 Standard NMR signals for a) 3-aminopyridine,
b) 2-aminopyrimidine and c) 2-aminopyridine.............. 100

8-9 Standard NMR signals for a) 2-pyridyl acetamide,
b) histamine and c) inosine triacetate...................101

8-10 Standard NMR signals for a) p-dimethylamino
benzaldehyde, b) 2-ethyl pyrrole and c) N,N-
dimethylaniline.......................................... 102

8-11 Proposed structure of the second product eluting
under GPLC peak B........................................ 103

8-12 Structure of 5-hydroxyhydantoin 5-carboxamide............108

8-13 Products from the gamma irradiation of adenosine
(80) and the photosensitized oxidation of 3 methyl
indole (81) .............................................. 109

9-1 UV spectra during photooxidation of a 1004M TMP
solution in a pH 7 phosphate buffer, p=0.5M, with a
deuterium lamp ........................................... 111

9-2 HPLC during photooxidation with a deuterium lamp of
a 100uM TMP solution in p=0.5M, pH 7 phosphate
buffer: a) before; b) 2hr, 37min; c) 10hr, 15min........ 112

9-3 Spectral changes of a 67VM TMP solution in pH 7
phosphate buffer, y=0.5M, after 4hr exposure to a
deuterium lamp........................................... 114

9-4 Effect of NaN3 on the photooxidation of TMP with a
deuterium lamp, 67pM TMP/0.1M NaN3 in pH 7 phosphate
buffer................................................... 116










9-5 Effect of DABCO on the photooxidation of TMP with a
deuterium lamp, 80M TMP/1.1mM DABCO in pH 7 phosphate
buffer ................................................... 117

10-1 Cyclic voltammograms during the reaction of TMP with
liver microsomes: a) NADPH and MgCl2 in the absence
of microsomes; b) same as a, plus TMP (450pM); c) same
as b, plus 300ul microsomal pellet, 10min incubation;
d) same as c, 150min incubation..........................126

11-1 Proposed structures for the products from EC/TSP/MS
of TMP ................................................... 132

A-1 FTIR of separated product fractions under GPLC peak A....139

A-2 FTIR of separated product fractions under GPLC peak B....140

A-3 NMR of separated product fractions under GPLC peak B.....141

A-4 FTIR of separated product fractions under GPLC peak C....142

A-5 NMR of separated product fractions under GPLC peak C.....143
















Abstract of Dissertation Presented to the Graduate School
of the University of Florida in Partial Fulfillment of the
Requirements for the Degree of Doctor of Philosophy

ELECTROCHEMICAL AND PHOTOCHEMICAL OXIDATION OF
TUBERCIDIN-5'-MONOPHOSPHATE

BY

TERESA E. PETERSON

December, 1987

Chairperson: Anna Brajter-Toth
Major Department: Chemistry

The purpose of this investigation was to use instrumental

methods to provide insight into the biological degradation pathways

of the purine drug tubercidin-5'-monophosphate (TMP). Tubercidin is

a highly cytotoxic antiviral and chemotherapeutic agent. In this

study the electrochemical and photochemical oxidation of TMP was

investigated.

The electrochemical oxidation of TMP was monitored by cyclic

voltammetry, ultraviolet spectroscopy and high pressure liquid

chromatography. Oxidation products formed by constant potential

electrolysis were separated by gel permeation liquid chromato-

graphy. The products were analyzed by mass spectrometry, Fourier

transform infra-red spectroscopy, fast atom bombardment mass

spectrometry and nuclear magnetic resonance spectroscopy.

Three oxidation products were isolated and identified. One

product is a hydantoin-type structure resembling a product of the










electrochemical oxidation of uric acid, 2,6-diaminopurine and

hypoxanthine. The other two products resemble products obtained from

the photosensitized oxidation of indoles and the irradiation of

nucleic acid components with ionizing radiation.

To determine if the electrochemical oxidation of TMP provides

any insight into its biological reactivity, the enzymatic oxidation

of TMP was investigated with xanthine oxidase, chloroperoxidase and

cytochrome P450. The results indicate that TMP may be a substrate of

cytochrome P450"

The oxidation of TMP by ultraviolet light indicates that water

and oxygen are necessary for the photooxidation to occur and that

singlet oxygen is produced during photooxidation. Similarities

between the electrochemical and photochemical oxidation are

discussed.

The cyclic voltammetric behavior of TMP was studied in different

environments (i.e., different pH, electrode surfaces and solvents).

The results indicate that the environment (i.e., solvent and surface)

plays a major role in the electrochemical behavior of TMP and

possibly other biological compounds.

A preliminary study of TMP oxidation was also carried out using

electrochemistry thermospray mass spectroscopy. Molecular ions

corresponding to the products generated by bulk constant potential

electrolysis were not detected by this method.
















CHAPTER 1
INTRODUCTION



Studies of the electrochemical oxidation of purines and purine

drugs can provide a great deal of information about the biological

oxidation of these compounds (1-9). In this study we have investi-

gated the electrochemical oxidation of tubercidin-5'-monophosphate

(TMP). The electrochemical oxidation was monitored with ultraviolet

(UV) spectroscopy and high pressure liquid chromatography (HPLC)

which provided additional information about the formation of inter-

mediates and products. On a preparative scale the products were

separated by gel permeation liquid chromatography (GPLC) and were

analyzed by gas chromatography mass spectrometry (GC/MS), fast atom

bombardment mass spectrometry (FABMS), Fourier transform infra-red

spectroscopy (FTIR) and nuclear magnetic resonance spectroscopy

(NMR). Additional studies of the photooxidation of TMP indicate that

most of the electrochemical and photochemical oxidation products are

the same.

Tubercidin is a naturally occurring, highly cytotoxic antiviral

and chemotherapeutic agent. It is structurally similar to adenosine

and belongs to a class of compounds known as 7-deazaadenosines, or

pyrrolopyrimidines (Fig. 1-1). Two other compounds in this class are

sangivamycin and toyocamycin (Fig. 1-1). Although their reactivity

is somewhat different these compounds are all cytotoxic to mammalian





















Tubercidin



NH2 CONH2


N N


^0


HO OH
Sangivomycin


NH2
N N

N




HO OH
Adenosine


Toyocamycin


Figure 1-1 Structure of tubercidin and some adenosine analogs.










cell lines in culture and inhibitory to the growth of bacteria,

fungi, ribonucleic acid (RNA) and deoxyribonucleic acid (DNA) viruses

(10). They are invaluable biochemical tools for studying cellular

and enzyme reactions because they can replace adenosine, adenosine

monophosphate (AMP), adenosine diphosphate (ADP) and adenosine

triphosphate (ATP) in cellular reactions (10). They are also useful

in determining the structural requirements for interaction with

enzymes (10).

Although much is known about the physical effects of these

drugs, little is known about their biological activation or degrada-

tion in the body. Metabolism studies provide a better understanding

of a drug's mode of action, toxicity and biochemical reaction-

mechanism so its safety and effectiveness can be evaluated. Modern

drugs are complex and can be metabolized by several reactions,

producing many metabolites which can be toxic, active or inert.

Drugs can be metabolized by a wide variety of enzymes in the

body. The enzymatic reactions responsible for metabolism are

classified as phase I and phase II reactions (11). Phase I reactions

usually result in oxidation, reduction or hydrolysis. In phase II

reactions the products of phase I reactions are conjugated to natural

compounds such as glucuronic acid. These reactions may transform the

drug into a more polar compound that can be rapidly removed from the

body or transform the drug into a toxic or therapeutically active

metabolite(s). Whether the metabolic reaction produces a toxic,

inert or therapeutically active metabolite is a function of the

drug's structural features (12).










Compounds which result from a slight modification in the

structure of naturally occurring purines or pyrimidines are known as

antimetabolites. Some of these antimetabolites exhibit antiviral and

antineoplastic activity. The most successful modifications usually

involve the base component rather than the sugar component (12). The

following types of modifications have been tried: replacement of

OH-groups with SH or NH2-groups, replacement of a ring C-atom with a

N-atom or a ring N-atom with a C-atom and introduction of halogen

groups at C-2 or C-6 (12). The numbering system for purines is given

in Figure 1-2. The majority of effective drugs are simply structural

modifications of purines and pyrimidines.

Purines and pyrimidines are present in cells as polymerized

nucleotides (DNA, RNA), free nucleotides (ATP, AMP), nucleosides

(adenosine) and bases (adenine) (Fig. 1-2). The numbering system for

purines is illustrated in Figure 1-2. Purines and pyrimidines serve

as precursors for nucleic acid synthesis; are involved in energy

metabolism, group transfer reactions, mediation of hormone action;

and act as metabolic regulators (13).

Antimetabolites can usually interact with the same enzymes as

the parent purine or pyrimidine. If this occurs, the enzyme may

become blocked and will not be able to fulfill its normal biological

functions or the antimetabolite may be converted to another compound

that disturbs the cell's metabolism.

Unlike other groups of therapeutic agents, nucleoside analogs

have the entire cell metabolism as a target. In contrast a folate

















N
3 H
PURINE


NH2

NN

N H

BASE
ADENINE


NH2


NUCLEOSIDE
ADENOSINE


N

N


7 DEAZA
ADENINE


TUBERCIDIN


NH2



NN
0 N
II
HO-P-O-CH2
OH o



HO OH
NUCLEOTIDE
ADENOSINE-S-
MONOPHOSPHATE
AMP


0
HO-P-(


TUBERCIDIN-5-
MONOPHOSPHATE
TMP


Structures of bases, nucleosides and nucleotides.


Figure 1-2











analog can only inhibit dehydrofolate reductase or closely allied

enzymes (14).

Essentially, all of the antimetabolites which act as

chemotherapeutic agents interfere with the synthesis and metabolism

of nucleic acids (12). Inhibition of nucleic acid synthesis

primarily affects those cells which have a high rate of mitosis.

This includes malignant cancer cells as well as other rapidly

dividing cells (e.g., bone marrow, hair and various epithelia

tissues) (12). Fortunately, cancer cells are more susceptible than

normal, rapidly dividing cells, to the actions of the antimeta-

bolites. The exact reasons are not fully understood but they stem

from physiological differences between cancer and normal cells (15).

Antimetabolite drugs are usually administered as nucleosides.

In the process of cell penetration they are converted to nucleotides

by adenosine kinase. An active compound will usually interfere with

nucleic acid synthesis. To interfere with nucleic acid synthesis the

compound must be phosphorylated, since the de novo synthesis of

purines proceeds entirely via phosphorylated intermediates. If the

compound is administered as a nucleotide, dephosphorylation usually

occurs during cell penetration, resulting in an inactive metabolite

(12).

Tubercidin was isolated from streptomyces tubercidus by Anzai,

Nakamura, and Suzuki (16). Subsequent studies by Suzuki and Marumo

determined its structure (17,18). This naturally occurring compound

enters the same anabolic pathways as adenosine, yet it is not

degraded by the enzymes that degrade adenosine. It has been shown to

inhibit mammalian and bacterial cell growth, RNA and DNA viruses, and











glycolysis (19). Glycolysis is one of the pathways by which cells

extract energy in the form of ATP by metabolizing glucose to

pyruvate. In addition to this, tubercidin interferes with

mitochondrial respiration, de novo purine synthesis, protein, RNA and

DNA synthesis, and transfer RNA processing (19).

Tubercidin showed promise as a drug for several reasons. First,

it was converted to the active nucleotide form during cell

penetration and, secondly, it was stable to deamination by adenosine

deaminase and to glycosidic bond cleavage by purine nucleoside

phosphorylases. These are the two major pathways which inactivate

purine drugs (19,20).

The therapeutic effects of tubercidin result from its wide range

of biological activity. Tubercidin was tested against a large number

of potential targets. It was first evaluated as an antitubercular

and antitumor agent. In both cases the drug showed no selectivity

and pronounced toxicity to the host (14). This toxicity manifested

itself in the form of nephrotoxicity and venous thrombosis at the

sites of injection (20). A nephrotoxin is a substance which is toxic

to kidney cells. Thrombosis refers to blood clotting. Toxicity is

not encountered when tubercidin is suspended in a petrolatum base and

applied to the skin. In this application, tubercidin is effective

against basal cell skin carcinomas (14).

Recent reports in the literature have shown that tubercidin,

administered in combination with nitrobenzylthioinosine (a potent

purine transport inhibitor) protects the host against some of the

cytotoxic effects of tubercidin (21,22). Earlier studies showed that

when tubercidin is added to animal or human blood "in vitro" it










enters erythrocytes and is maintained there in nucleotide form. When

the blood is transfused back into the subject nephrotoxicity and

venous thrombosis are avoided (20). The discovery that tubercidin

enters red blood cells led to an effective treatment for

schistosomiasis, a parasitic blood disease (14). A later report

shows that tubercidin, along with other deaza analogues of adenosine,

inhibit blood platelet aggregation (i.e., blood clotting) (23).

A study of the effects of various inhibitors of purine

metabolism showed that tubercidin may be useful against

trichomoniasis in humans (a parasitic infection). This study was

done using two parasites that synthesize nucleotides differently

(24). To understand which cellular reactions contribute to the

cytotoxic effects of tubercidin and toyocamycin genetic and

biochemical studies are being conducted on mutant cells resistant to

toyocamycin and tubercidin (25).

A great deal is known about the effects of tubercidin on

cellular processes such as mitochondrial respiration, synthesis of

purines, RNA and DNA and on biochemical targets, cancer cells,

viruses and parasites (10). However, it is not known which of these

effects is primarily responsible for the inhibition of cell growth

and cytotoxicity (20). Is the compound metabolized to another

compounds) which exerts the vast number of biological effects? Are

these metabolites toxic or inert? These are some of the questions

for which answers are being sought.
















CHAPTER 2
ELECTROCHEMICAL STUDIES OF BIOLOGICAL OXIDATIONS



Carbon electrodes are commonly used in electrochemical studies

of biological oxidations (1-9). These electrodes provide a wide

potential range (+1.2 to -1.7V versus SCE) and a low background

current in aqueous solutions (26,27,28). The mercury electrode

cannot be used at potentials more positive than about +0.4V versus

SCE (27). Platinum electrodes are troublesome because oxygen and

hydrogen adsorb onto the surface and high backgrounds are obtained

(29). The formation of an oxide layer on platinum inhibits many

oxidation processes including those of organic substances. The

potential range for platinum in aqueous solution is between +0.9 and

-0.6V versus SCE.

There are a number of carbon based electrodes in use. Carbon

paste electrodes were introduced by Adams and co-workers (27).

Glassy carbon (GC) electrodes were applied for the first time in

electroanalytical chemistry by Zittel and Miller (cited in 30). They

provide a smooth mirror-like surface with very low residual

currents. The pyrolytic graphite (PG) electrode has a much rougher

surface and usually provides for faster kinetics. It was introduced

by Beilby and co-workers and Miller and Zittel (cited in 27 and 30).

The electrochemical oxidation of a large number of biological

purines has been reviewed (1,2). The most extensively studied purine










is uric acid. Its electrochemical oxidation has been studied using

lead oxide, spectroscopic graphite and PG electrodes (1,2).

The electrochemical oxidations of a large number of purines

besides uric acid have been studied. These include adenine (3),

guanine (7), 6-thiopurine, 2-thiopurine and 2,6 dithiopurine (1).

More recently, there have been studies of 2,6 diaminopurine (4), 9-B-

D ribofuranosyluric acid (5), xanthine (6), 8-oxyguanine (7) and

hypothanthine (8).

In the majority of these studies, both the electrochemical and

enzymatic oxidations were carried out. For the most part, the

electrochemical and enzymatic oxidations produced similar products.

However, more mechanistic information could be obtained from the

electrochemical studies.

The enzymatic oxidation of purines by xanthine oxidase has been

studied by Bergman and Krenitsky and co-workers (31,32). These

studies suggest that oxidation occurs initially at the C-6 position

then at the C-2 or C-8 positions. The numbering system for purines

was given in Figure 1-2. If oxidation occurs at C-2, the C-8

position will also be oxidized. The reverse is also true.

Dryhurst and co-workers made a number of generalizations about

the electrochemical oxidation of purines (1).

1. Purine itself is not electrochemically oxidizable. To be

oxidized, a purine must contain an amino or oxy group in the

pyrimidine ring of the purine molecule.

2. Susceptibility to oxidation increases as the purine ring is

oxygenated.












3. Oxidation appears to begin initially at any unoxidized or

unsubstituted N=C bonds. The final site of electrochemical

attack is the C4=C5 bond.

4. An unstable diimine type structure appears to be the initial

oxidation product. This can occasionally be detected by fast

sweep voltammetry at a PG electrode.

5. If the purine is substituted at both N-3 and N-7 or N-9 by

methylation, a diiminium ion forms which is more reactive

than the diimine. This is usually not detected by cyclic

voltammetry.

6. Hydrolysis of the diimine or diminium ion leads to a 4,5-

diol, which undergoes a number of secondary reactions to give

final products.

7. A large number of purines are adsorbed at PG.

These generalizations are supported by the electrochemical studies of

uric acid (33,34).

The primary reaction in the electrochemical oxidation of uric

acid involves a quasi-reversible 2e-2H+ process, leading to an

anionic diimine intermediate, II (Fig. 2-1) (33,34). This diimine,

II, undergoes a rapid (pseudo) first order hydration reaction, giving

an anionic imine-alcohol, III. This intermediate may be detected by

UV spectroscopy during thin layer electrochemical oxidation. It is

also the species responsible for a reduction peak at -0.9V versus SCE

in the cyclic voltammogram of uric acid. A (pseudo) first order

hydration of this imine alcohol, III, gives a diol, IV, which


















-2e-,-2H+-
+2e-,+2H*


H20
H+


H20


(111)


H20


H
H2N O > + CO2


H H H
(VII)


Figure 2-1


Proposed reaction scheme for the electrochemical
oxidation of uric acid at physiological pH. Modified
from reference (34).


(IV)


0
II H
C I
it 0
N

O N
H










decomposes via an isocyanate to allantoin, VII. Cyclic voltammetry,

thin layer spectroelectrochemistry and GC/MS provided information to

support this mechanism (33,34).

The enzyme (peroxidase) oxidation of uric acid yields

intermediates that are identical, spectrally and electrochemically,

to the UV-absorbing intermediates generated upon electrochemical

oxidation (33,34). The ultimate reaction product formed by

electrochemical and enzymic oxidation is allantoin. These results

support the conclusion that the electrochemical and enzymic reactions

are, in a chemical sense, the same (34).

Since TMP is structurally similar to the purine nucleoside,

adenosine (Fig. 1-1), a similar type of electrochemical behavior may

be expected for both compounds. There are no reports on the

electrochemical oxidation of adenosine but there is a report on the

electrochemical oxidation of the parent purine adenine (3). Since

the electrochemical behavior of uric acid and uric acid riboside is

very similar (i.e., the same types of products are formed in both

oxidations) and the ribose group is not electroactive, the electro-

chemical behavior of adenosine may be similar to that of adenine (5).

The primary electrochemical oxidation of adenine is shown in

Figure 2-2. It was shown that six electrons were involved in the

oxidation (3). The number of electrons was determined using

coulometry and the majority of products were analyzed by paper and

thin layer chromatography. The proposed mechanism indicates a series

of 2e-2H+ oxidations to 2-hydroxyadenine, II, then 2,8-dihydroxy

adenine, III. Finally, oxidation of the 4,5 double bond leads to a














2N.
3


NH2

+ H20 ON 5 + 2H + 2e
H H
(I1)


NH2 H
N NX
H20 ---> O N C=0 2H* 2e .___

H H
(111)


:=0 + 2e


* H
(IV)


H20
(IV) -


0
HN

H


NH2 COOH

O N JO
H
(VI)


NH2
* H* + CO2 + C=0 + NH3 + e
NH2


0


O N NH2
H
(VII)


H2Ny NH
jH J-.,
HO N'^
H
(VIII)


+ NH3 + CO2


0 0
S11 II

0 N N2
H H
(IX)


(IV) H20


Figure 2-2


+ H' e H20


H

ON C=0 +" NH3 CO2. H*
H HH
(X)


Proposed reaction scheme for the electrochemical
oxidation of adenine. Modified from reference (3).


(IV)


H20 le











dicarbonium ion, IV, similar to that formed in the electrochemical

oxidation of uric acid (3).

Following the primary electrochemical oxidation, three distinct

chemical and electrolytic reactions occur (Fig. 2-2). Further

electrochemical oxidation of the dicarbonium ion, IV, leads to

parabanic, V, and oxaluric acids, VI, urea and ammonia. Some of IV

is also reduced to 4-amino-2,5,6 trihydroxy-, VII, and 5 amino-2,4,6-

trihydroxy pyrimidine, VIII, which condenses to form 4-aminopurpuric

acid, IX. Allantoin, X, forms from acid hydrolysis of IV, liberating

NH3 and CO2 (3).

Based on what is known about the electrochemical and enzymatic

oxidation of purines, we can hypothesize what might occur during the

electrochemical oxidation of TMP. Since the C-6 position contains an

NH2 group, oxidation will occur at the C-2 or C-8 positions or

both. The C4=C5 double bond may also be oxidized. It is already

known that TMP is resistant to deamination and glycosidic bond

cleavage (19,20). The oxidation products identified confirm this.

The purpose of this investigation was to provide insight into

the biological degradation of the purine drug tubercidin. The

nucleotide form of this drug (i.e., TMP) was investigated because it

is more soluble in water and because tubercidin is probably in the

nucleotide form inside the cell (12). Although much is known about

its biological activity, little is known about the metabolites which

may be the source of drug toxicity. For the most part, purines are

biologically degraded via oxidation reactions and purine drugs are











often activated in the process (1, p. 127). In this study electro-

chemistry was used as a tool for evaluating the oxidative degradation

of TMP.

A number of reports in the literature, including the above

description of the oxidation of uric acid, illustrate that the

electrochemical and enzymatic oxidation of purines can be chemically

the same (1,2,4,5,6). The experimental methods used in these types

of studies, which were described earlier, were also used in this

study. In addition, the electrolysis of TMP was also monitored by

HPLC and products were further analyzed by FTIR and NMR.

To determine if the results from electrochemical methods provide

any insight into the biological redox reaction of TMP the enzymatic

oxidation of TMP was also investigated. Tubercidin-5'-monophosphate

was not significantly oxidized by any of the enzymes that were

investigated although experimental results indicated that TMP may be

a substrate of cytochrome P450"

Studies of TMP oxidation as a function of pH, at different

electrode surfaces and in polar and nonpolar solvents were carried

out to provide more information on the reactivity of TMP. These

studies indicate that the environment (i.e., solvent and surface)

plays a major role in the oxidation of TMP.

The photooxidation of TMP, which has not been previously

reported, was also investigated. Similarities between the

electrochemical and photochemical oxidation are discussed.

Oxidation of TMP was also studied using electrochemistry

thermospray mass spectrometry (EC/TSP/MS). This technique allows

oxidation products to be analyzed on-line by mass spectrometry







17



without the need for separations. Preliminary results do not confirm

the products identified from bulk constant potential electrolysis.
















CHAPTER 3
INTRODUCTION TO METHODS USED IN THE STUDY
OF BIOLOGICAL OXIDATIONS



3.1 Cyclic Voltammetry

3.1.1 Absorption, Diffusion and Reversibility

Cyclic voltammetry is used quite often in the mechanistic

studies of redox systems (26, p. 86). The experimental set up

usually consists of a potentiostat, waveform generator and a three-

electrode cell containing counter, working and reference

electrodes. The current is measured at the working electrode as the

potential is scanned at different rates (26, p. 86). The resulting

electrochemical process can be adsorption or diffusion controlled and

the electron transfer may be reversible or irreversible. The peak

current equations for the different process are given below (29, pp.

218, 222, 522, 525).

Diffusion controlled systems:

reversible i = (2.69x105)n3/2 A D/2 v2 C (3-1)

5 1/2 1/2 1/2
irreversible i = (2.99x10 )n(an ) A C D v (3-2)
p a o o
Adsorption controlled systems:

2 2
reversible i vA (3-3)
p 4RT o

nan F 2Av *
irreversible i = a (3-4)
p 2.718RT










where

i = peak current, A

n = number of electrons transferred per molecule or ion,
eq mol-

a = transfer coefficient, measure of the symmetry of the
energy barrier

na = number of electrons involved in the rate determining
step, eq mol1

A = electrode area, cm2

D = diffusion coefficient, cm2sece

Co* = bulk concentration, mol cm-3

v = scan rate, V/sec-1

F = surface concentration of species before redox reaction,
mol cm

T = temperature, K

R = gas constant, J mol1K-1

F = Faraday, 96,496 coul mole-1 of electrons


For diffusion controlled processes, ip is proportional to C *
1 /2
and v/2 For adsorption controlled processes ip is proportional to

v. A broad peak usually indicates a diffusion process while a sharp

symmetrical peak indicates an adsorption controlled process (35).

In a diffusion controlled process reactant moves from one region

to another as the result of a concentration gradient. As the

concentration of an electroactive species, X, is depleted at the

electrode a concentration gradient is set up and more X diffuses from

the bulk solution to the electrode (26, pp. 9-49).











Adsorption of a species onto the electrode surface can occur as

a result of electrostatic attraction or covalent bonding between the

species and the electrode. Adsorption can also result from

hydrophobic interactions between a species in solution and the

electrode. For example, neutral organic molecules in aqueous

solutions may adsorb onto the electrode surface as a result of

hydrophobic interactions (26, pp. 43-48).

For reversible and irreversible adsorption and reversible and

irreversible diffusion controlled systems the equations for ip are

similar. One of the diagnostic criteria for a reversible diffusion

controlled system is a .059/n V separation between oxidation and

reduction peaks (26, pp. 90-92). This is observed when the voltage

scan rate is small compared to the electrode reaction rate con-

stant. If the reaction is quasi-reversible, the peak potentials will

be separated by more than .059/n volts (26, pp. 90-92). Typically, a

system is irreversible if only one peak, either cathodic or anodic,

is present in the cyclic voltammogram or if cathodic and anodic peaks

are present and they are separated by more than 0.10 volts.

3.1.2 Homogeneous Chemical Reactions

In many cases the heterogeneous electron transfer reaction at

the electrode is affected by homogeneous chemical reactions.

Consider the reaction 0 + ne -- R, where 0 is the oxidized form of

the standard system and R is the reduced form of the standard

system. Figure 3-1 summarizes some of the reactions that could

occur. The oxidized form of the standard system could be produced

from a preceding chemical reaction. In a following reaction R












a) Preceding


Y0O
0+ne R


b) Following Reaction (EC)
0+ ne -R
R X

c) Catalytic Reaction
0+nee R
R + Z->0+ Y





Figure 3-1 Homogeneous chemical reactions that accompany
heterogeneous electron transfer processes.


Reaction (CE)










could chemically react (e.g., with solvent) to form an electro-

inactive species X at potentials where 0 is reduced. If the product

of a following chemical reaction is electroactive where 0 is reduced,

a second electron transfer could take place. Additionally, a

catalytic reaction could occur where the product R undergoes a

chemical reaction with a nonelectroactive species Z in solution to

regenerate 0 (29, p. 431).

3.1.3 Determination of pKa

Cyclic voltammetry at slow scan rates allows application of the

Nernst equation to the determination of pKa's for compounds involved

in acid/base reactions. The Nernst equation is valid at slow scan

rates because there the system may be assumed to be at equilibrium.


mR <=> pP+ + rH+ + ne

0.0591 [P]P [H]r
Ep = E + log (3-5)
Sln [R]m
H.0591 log[P]P
E = E 59(rpH -) (3-6)
p o n R]m

where

n = number of electrons transferred per molecule or ion,
eq mol-

Ep = measured peak potential, V

E = standard potential, V


The compound exists in different forms in the pH region above

and below the pKa. This should be reflected in the measured r/n

values for different acid/base forms (equation 3-6). Thus a plot of

Ep versus pH may be used to estimate the pKa. Since the slope of

this plot equals -.0591r/n the number of protons, r, and electrons,











n, involved in the oxidation or reduction of the electroactive

species can be estimated.

3.1.4 Estimation of n

The equations of ip for reversible and irreversible processes

(equations 3-1 through 3-4) show that ip is proportional to n.

Therefore, the n-value for a compound can be estimated by comparing

its ip to the ip of a compound with a known n-value. Equation 3-7

can be used to compare the i values of two electrochemically

irreversible diffusion controlled systems, where k is the system with

known n-value and u is the system with unknown n-value. The terms A,

v1/2, D 1/2 and Co* do not have to be included in the calculation if

these values are the same for both compounds. Since Do depends on

the size of the diffusing molecules in liquids, it will be similar

for structurally similar compounds under the same conditions (36,37).

i n (an )1/2
p uk \ak
n = (3-7)
u i 1/2
P (ana)

If the ip values for two irreversible processes of similar compounds

are compared, the ana values have to be estimated independently for

both compounds. The ana values can be estimated from equation 3-8 by

measuring Ep at two sweep rates v1 and v2 (27, pp. 135-137).


RT v2
E E n (3-8)
p p2 ana- Fv










3.2 Determination of n-Values by Coulometry

Coulometry is the standard method used to determine n-values.

This experiment is usually carried out in a three-compartment, three-

electrode cell. The working electrode is maintained at a constant

potential where the reaction occurs at a maximum rate (i.e., ca.

100mV past the cyclic voltammetric peak). In coulometry the current

is integrated until the redox reaction is completed.

As the electrolysis proceeds the current decays exponentially as

a function of time (26, pp. 119-121). A coulometer integrates this

current over time to give the total charge (Q) passed during the

electrolysis. The total charge can be related to n through Faraday's

Law (equation 3-9).


Q = nFN (3-9)


where

Q = total charge passed, C

n = number of electrons transferred per molecule or ion, eq mol"1

F = 96,496 C eq-1

N = number of moles of substance being electrolyzed


Ideally the coulometer stops collecting charge when the electroactive

starting material is completely electrolyzed.

If the charge increases after the starting material has been

completely electrolyzed it is not possible to calculate an exact

value of n. Under such conditions products may be further oxidized

at the electrolysis potential or a catalytic oxidation or reduction











of water may be occurring (38). The presence of oxygen containing

functional groups such as carboxyl, hydroxyl, carbonyl, lactone and

quinones on carbon electrodes has been suggested (39). These groups

could participate in catalyzing the oxidation or reduction of water

(38,40).



3.3 Formation of Oxidation Products by Constant
Potential Electrolysis

Constant potential electrolysis can be used to generate products

of the electrode reaction which can be isolated and further

analyzed. The experimental set up is the same as that for

coulometry. When the electrode is held at a constant potential the

intermediates and products from the electrode reaction are generated

in solution. By periodically stopping the electrolysis, the solution

can be analyzed for these intermediates and products by a number of

methods. Cyclic voltammetry and UV spectroscopy are most commonly

used.

In this investigation an HPLC method was developed and used to

monitor the electrolysis to show when TMP was completely

electrolyzed. Cyclic voltammetry and UV spectroscopy were not useful

in this respect because the intermediates and products have cyclic

voltammetric peaks and UV absorbance bands in the same region as TMP.



3.4 Chronocoulometry

3.4.1 Estimation of n

The n-value can also be estimated using chronocoulometry. In a

chronocoulometric experiment the current associated with an













electrochemical process is integrated over time. The time scale of

the experiment is typically 50-500ms. The behavior can be analyzed

by plotting charge versus the square root of time (29, p. 200). This

method was popularized by Anson and co-workers and is sometimes

referred to as an Anson plot (41,42). The experiment is carried out

in an unstirred solution at a planar electrode held at an initial

potential (Ei) where no significant electrolysis takes place. The

potential is then stepped to Ef, a potential where electrolysis

occurs at a maximum diffusion controlled rate. This current is

described by the Cottrell equation (equation 3-10). The integrated

form of this equation gives the cumulative charge passed, Qd

(equation 3-11). The terms have their usual meaning.


nFAD 1/2C
i(t) = id(t) = 1/2 0 (3-10)



2nFAD 1/2C *t1/2
0 0
Qd = 1/2 (3-11)



The cumulative charge, however, is due only to the diffusing

material. Additional charge arises from double layer charging Qd1

and from the oxidation or reduction of adsorbed reactant

molecules (nFAr ). The total charge measured is
0

2nFAD 1/2C *t1/2
Q 0 1/2 + Qdl + nFAro (3-12)
iT










The slope of a Q vs t1/2 plot is


Qd 2nFAD C 0
slope = /-- 1/2 (3-13)
t 1

Since the slope is proportional to n, the n-value of a compound can

be estimated by comparing its slope to that of a compound whose

n-value is known (equation 3-14). In equation 3-14 k is the known

and u is the unknown system.


slopes (n C o*)k
slope (n C *) (3-14)
u o u

In order to use this equation A must be the same and Do must be

known. If structurally similar compounds are chosen, Do is expected

to be the same for both compounds (36,37).

3.4.2 Determination of Electrode Area

From a chronocoulometric experiment using a standard solution of

K 3Fe(CN)3 (Do 7.6x10-6cm 2sec-1) electrode area can be determined

(equation 3-15). The terms have their usual meaning. This equation


1/2
A slope 2 (3-15)
2nFD /C
0 0


was derived from equation 3-13.
















CHAPTER 4
EXPERIMENTAL



4.1 Cyclic Voltammetry

All cyclic voltammetric experiments were carried out with an IBM

EC225 2A Voltammetric Analyzer (IBM Instruments Inc., Danbury,

Connecticut) with a Houston Instruments 2000 X-Y recorder (Bausch and

Lomb, Austin, Texas) or a model 173 Potentiostat and model 175

Universal Programmer from EG&G Princeton Applied Research (Princeton,

New Jersey). A 10 ml glass beaker contained the analyte solution and

working, counter, and reference electrodes. The PG working electrode

was prepared by sealing a piece of PG (1.5mm x 10mm) (Pfizer) into

glass tubing with inert epoxy (1C white, Dexter Corporation, Hysol

Division). In some experiments a GC working electrode was used.

These were prepared by sealing a rod (3mm diam x 10mm length) into a

glass tube with epoxy. Vitreous (glassy) carbon was purchased from

Electrosynthesis Company, Inc. (Amherst, New York). The platinum

wire counter electrode was encased in glass. Electrical contact for

these electrodes was made with a copper wire and mercury.

The saturated calomel (SCE) reference electrode was prepared

according to a standard procedure (43). The SCE was connected to the

analyte solution by a salt bridge prepared by filling one side of a

U-shaped glass tube with KC1 agar and the other side with analyte

solution (Fig. 4-1). The KC1 agar is prepared by warming 4g of agar


































analyte
solution


Figure 4-1


Schematic of a saturated calomel electrode in an
electrochemical cell with analyte solution.


SCE










and 90ml of water on a steam bath, then adding 30g of KC1. When the

salt has dissolved the gel is pipetted into the glass connection.

In order to obtain a reproducible surface the working electrodes

were resurfaced before each measurement. For the PG electrode,

resurfacing was accomplished by grinding the electrode on 600 grit

SiC paper using a Buehler Ecomet 1 Polisher-Grinder (Evanstown,

Illinois). The polishing paper was a Carbimet Abrasive disc (Buehler

Ltd., Lake Bluff, Illinois). GC electrodes were resurfaced by

polishing on Buehler billiard cloth with Gammal gamma alumina (Fisher

Scientific). Once polished the GC electrode was rinsed and placed in

an ultrasound bath for 10min to remove any residual alumina. Both

glassy carbon and pyrolytic graphite electrodes were rinsed

thoroughly with deionized water and carefully wiped with a Kimwipe

tissue to remove excess water and graphite particles.

The supporting electrolytes used for cyclic voltammetric

experiments were phosphate and ammonium acetate buffers of varying

pH, ionic strength 0.5M, prepared with doubly distilled deionized

water. All solutions were deaerated 5min with N2 before

measurement. A blanket of N2 was passed over the solution throughout

the experiment to keep the solution free of dissolved 02.

For the cyclic voltammetric experiments done in N,N-

dimethylformamide (DMF) and dimethylsulfoxide (DMSO) (Fisher

Scientific) the supporting electrolyte was 0.5M tetrabutylammonium

perchlorate (TBAP) (Kodak, Rochester, New York). When a cyclic

voltammogram was run in organic solvent the SCE was connected to the

analyte solution by a vicor tip and not by the salt bridge described










earlier. The vicor tip prevented the organic solvent from mixing

with the solution inside the SCE better than the salt bridge.

In some experiments the GC electrode was electrochemically

pretreated before use. Electrochemical pretreatment was carried out

by applying a constant positive potential of +1.8V for 8min to the

electrode in a solution of supporting electrolyte (44). This is

followed by 2-5 cyclic scans towards negative potentials first (+1.8

to -1.7V) in the supporting electrolyte at a sweep rate of 20mV/s.

These cyclic scans are carried out until the background peaks at ca.

-1 and 1V have disappeared.

Tubercidin-5'-monophosphate was obtained from Sigma Chemicals

(St. Louis, Missouri) and was used without purification.



4.2 Coulometry and Constant Potential Electrolysis

All coulometric and constant potential electrolysis experiments

were carried out with a model 173 potentiostat and model 179 digital

coulometer (EG&G and Princeton Applied Research). The current range

was 10mA.

The experiments were carried out in a three-compartment, three-

electrode cell with compartments separated by cation exchange

membranes type P-1010 (RAI Research Corporation, Hauppauge, Long

Island, New York). The middle compartment contained the working

electrode and the electroactive species in a solution of supporting

electrolyte. The outer compartments contained a reference SCE and Pt

mesh counter electrode in equal volumes of supporting electrolyte.

Working electrodes were two rectangular plates of PG (6.3 x 1.8cm)










(Pfizer) which were cleaned by polishing with 600 grit SiC paper

followed by copious washing. All three-cell compartments contained

10ml of solution. Typically, 600iM TMP solutions in phosphate or

ammonium acetate buffers, ionic strength 0.5M, were electrolyzed.

Nitrogen (N2) was passed during the electrolysis through a

continuously stirred solution. Stirring was done because it

increased the magnitude of the current during electrolysis and

consequently decreased the time of an electrolysis by minimizing the

diffusion layer (26, p. 199).

To see the formation of intermediates and products the

electrolysis was monitored with cyclic voltammetry, UV spectroscopy

and HPLC. This was done by stopping the electrolysis every 10-

30min. Cyclic voltammograms were recorded using a PG electrode.

A Tracor Northern TN 6500 diode array spectrophotometer

(Middleton, Wisconsin) with a Hewlett Packard ink jet printer or a

Hewlett Packard 8450A diode array spectrophotometer with a Hewlett

Packard 7470A plotter were used for measuring UV spectra during

electrolysis. For UV spectra, the electrolysis solution was pipetted

into a 1cm quartz cuvette (4ml volume) and the absorbance spectrum

was recorded from 200 to 400nm. The reference quartz cuvette

contained the supporting electrolyte used for electrolysis (i.e.,

phosphate or ammonium acetate buffer).

For HPLC analysis an Altex 110A pump and solvent programmer were

coupled to a Kratos Spectroflow variable wavelength UV detector. A

Fisherbrand Resolvex C-18 (4.6mm x 25cm) column was used. The C-18

column was equilibrated with the aqueous mobile phase for 1hr at a











flow rate of 2ml/min. When the experiment was complete acetonitrile

was flushed through the column for 1hr at a flow rate of 2ml/min to

remove the aqueous buffer. The sample injection volume was 20pl.

The separation was isocratic with a 0.02M potassium hydrogen

phosphate solution, pH 4.7-5.1, as the mobile phase. This mobile

phase was used because TMP was retained for a reasonable length of

time on the column. Doubly distilled deionized water exposed to UV

light for 24hr was used to prepare the mobile phase. The mobile

phase was filtered through a 0.45pm filter before use. Exposure to

UV light photodecomposes many organic compounds which would interfere

with UV detection. The flow rate was 1ml/min. Separations of

electrolysis products were monitored at 220, 237, 260 and 305nm.



4.3 Product Analysis

Once the bulk electrolysis was complete the solutions were

quantitatively removed from the working electrode compartment with a

pipet and transferred to a clean glass vial. The solution in the

vial was immediately frozen in a dry ice acetone bath (-700C) and

lyopholized using a Freeze Dry-8 (model 75040, Labconco Corporation,

Kansas City, Missouri). The dry residue was dissolved in 1-2ml of

doubly distilled deionized water and separated on a liquid

chromatographic column (38cm x 3.0cm) packed with G-10 Sephadex size

exclusion gel (molecular weight cut off was 700) using doubly

distilled deionized water as the mobile phase. The G-10 resin was

obtained from Pharmacia Fine Chemicals (Piscataway, New Jersey). The

flow rate was maintained at approximately 0.065ml/min. An LC 200










Fraction Collector (HaakeBuchler Instruments, Saddlebrook, New

Jersey), was used to collect separated fractions. Ca. 80 drops

(3.4ml) were collected per fraction. Sixty fractions were

collected. All products eluted within this volume. Some of the

products coeluted approximately with each other.

It was found that in the separation of electrolysis products, as

the pH of the electrolysis solution increased, the region where

phosphate eluted in the separation decreased. For optimum separation

of products from phosphate, electrolysis was carried out in pH 9.5

phosphate buffer. It was verified that the same products formed at

every pH. Separation from phosphate was important because phosphate

can be an interference in GC/MS analysis of the silylated

derivatives.

Electrolysis was also attempted in ammonium acetate buffer

(NH4Ac) (pH 6.6). This buffer was chosen because it can be removed

by lyopholization which eliminates the need to separate products from

phosphate. However, it was still necessary to separate the products

from each other by GPLC so that they could be analyzed further by NMR

and FTIR. The results from such separations indicated that the major

products formed in phosphate buffers also formed in NH4Ac. However,

the GPLC retention times of the products formed in NH4Ac were so

similar that this method of product separation was not pursued

further.

The presence of free phosphate in each fraction was detected by

a wet chemical test. Five drops of the sample were acidified with 2

drops of 0.01M nitric acid and a few crystals of ammonium molybdate










(Fisher Scientific) were added. The formation of a yellow

precipitate indicated the presence of phosphate. Phosphate bound to

the ribose group was not detected by this method. Only fractions

free of phosphate were lyopholized and analyzed.

The absorbance of each fraction was measured separately in a 1cm

quartz cuvette with the spectrophotometer described earlier. The

reference quartz cuvette contained doubly distilled deionized

water. A spectrum was recorded from 200 to 400nm for each

fraction. All products absorbed at 220nm.

The fractions which contained significant amounts of products as

indicated by UV spectroscopy were analyzed by HPLC. The HPLC

analyses were monitored at 237, 260 and 305nm.

Gas chromatographic/mass spectral analysis of the oxidation

products was carried out on a Finnigan Model 4021 GC/MS system with a

DB-5 (0.32mm id x 15m) capillary column (J & W Scientific, Rancho

Cordova, California). Electron impact, EI, spectra using a 70eV beam

and chemical ionization, CI, spectra using methane gas were

obtained. The injection volume was 5pl.

Since purines and their metabolites are often too polar and

thermally unstable to be volatilized, the oxidation products were

silylated before GC/MS analysis to insure volatility. The dry

samples were derivatized for 10-20min at 120C with 70ul

bis(trimethylsilyl)trifluoroacetamide (BSTFA) from Supelco Inc.

(Bellefonte, Pennsylvania) and 70pl silylation grade acetonitrile

from Pierce Chemical Co. (Rockford, Illinois) in 3ml reaction vials

(45). The silylations were carried out in a Reacti-Therm heating











module (Pierce Chemical Company). The reaction vials were tightly

sealed with teflon lined plastic caps (Supelco, Inc.).

Methane CI mass spectra give characteristic M+1, M+29 and M+41

ions corresponding to the addition of H+, C2H5+ and C H5+ to the

molecular ion, M+ (46). In a silylation acidic protons are replaced

by Si(CH3)3 groups (47). Typically, silylated samples give large

M-15 peaks corresponding to the loss of CH3 from the Si(CH3)3 group

(46, pp. 295-305;47). Silylation can result in overlapping mass

spectra of partially silylated derivatives which may not be separated

by GC. The resulting mass spectra then give sets of M and M-15 peaks

separated by 72 mass units; 72 corresponds to the molecular weight of

Si(CH3)3 minus the H it would replace. During silylation

decomposition of the oxidation products may occur (7).

A computer program developed by Toth et al. was used to analyze

the mass spectra of the silylated derivatives (48). This program

allowed determination of the number of silyl groups and the molecular

weight of the compound before silylation.

Fast atom bombardment mass spectrometry (FABMS) was also used to

confirm the molecular weight assignments. In FABMS there is little

fragmentation and polar samples can be analyzed without the need for

derivatization (49,50,51). The spectra were obtained with a Kratos

MS-50 mass spectrometer equipped with an RF magnet and DS-55 data

system. Fast xenon atoms (8eV) were used. The compounds were

dissolved in a thioglycerol matrix. Both positive and negative ions

were detected. The sensitivity of this technique is dependent on the

sample concentration and its susceptibility to ionization (50,51).











This analysis was carried out at the Middle Atlantic Mass

Spectrometry Facility at Johns Hopkins University School of Medicine.

To obtain additional structural information about the oxidation

products FTIR spectra were run. A Nicolet 5-DX spectrometer was

used. The samples were prepared in a Nujol mull and analyzed on KC1

or NaCl plates. Data were obtained after 126-256 scans. Nujol

absorbs in the region from 2850-3000cm-1 and 1400-1500cm-1 and the

spectra were background corrected. Because of the very small amount

of sample available it was not possible to correct for the background

100%; therefore, a spectrum of the background was overlayed on each

FTIR to distinguish between sample and Nujol absorptions.

Analysis of products by proton NMR was also used to obtain

structural information. The NMR spectra were run on a Nicolet NT-300

system. The samples were 0.2-0.5mg dissolved in l00l of deuterated

dimethyl sulfoxide (dmso-d6) (99.9 atom %D, MSD Isotopes, Montreal,

Canada) and pipetted into microcapillary NMR tubes. The NMR tubes

were rinsed with acetone and dried in an oven for 2 days at 900C to

remove water. The tubes were then flushed with N2 before capping.

Data were acquired for about 24hr. Since the oxidation products were

hygroscopic water was present in the samples and was evident in the

NMR spectra between 2-4.5 ppm. Presaturation of the water peak was

done to enhance resolution (52).



4.4 Chronocoulometry

To determine the electrode area a chronocoulometric experiment

was carried out using a Bioanalytical Systems BAS 100 Electrochemical







38



analyzer (West Lafayette, Indiana). A 1mM solution of potassium

ferricyanide, K3Fe(CN)6 (Mallinkrodt) was prepared in a 0.50M KC1

solution and was used to determine the electrode area. The potential

was stepped from 359 to -172mV and the pulse width was 250ms. The

electrode area was determined from the slope of an Anson plot (i.e.,

charge versus square root of time) using equation 3-15 (41,42).



4.5 Photooxidation Reactions

The Tracor Northern spectrophotometer was used to record the

spectral changes during irradiation of various TMP solutions in a 1cm

quartz cell with a deuterium lamp (type L1637, Hamamatsu TV Co.,

Ltd., Japan). The operating current of the lamp was 0.3amps. All

solutions were less than 100UM in TMP. For solutions at TMP

concentrations greater than 100M the detector was saturated and

spectral changes were difficult to see. The spectral changes were

monitored every 15-30min. The photooxidation of TMP was also

monitored by HPLC using the same HPLC conditions described to monitor

the constant potential electrolysis of TMP.

To determine if oxygen is necessary for the oxidation to proceed

the photooxidation of TMP was carried out in a pH 7 phosphate buffer,

y=0.5M, exposed to air and bubbled with N2. The photooxidation was

also carried out in a nonaqueous solvent, DMSO, to determine if water

is necessary for the oxidation to proceed. The photooxidation of TMP

was also performed in water alone and the results indicate that the

presence of phosphate does not affect the spectral changes during

photooxidation in a pH 7 phosphate buffer.











Photooxidation experiments were also carried out to determine if

singlet oxygen was involved in the photooxidation. The singlet

oxygen quenchers sodium azide (0.01M) (Fisher Scientific) and

diazabicyclo[2.2.2]octane (DABCO) (.0011M) (Sigma Chemicals) were

dissolved in 60-80M solutions of TMP. Supporting electrolyte was a

pH 7 phosphate buffer, u10.5M. The UV spectrum in the region from

200 to 400nm was recorded after 4hr of irradiation with the deuterium

lamp. Photooxidations of sodium azide and DABCO solutions with no

TMP present were also carried out to determine if these compounds

were degraded by the UV light after 48hr.

A similar photooxidation experiment was carried out with Trolox

(6-hydroxy-2,5,7,8-tetramethylchroman-2-carboxylic acid), a water

soluble antioxidant (Aldrich). This antioxidant is a quencher of

radical species (53). Spectral changes observed in a solution of

Trolox indicate that it is decomposed by UV light thus further

experiments were not performed.



4.6 Enzymatic Oxidation Studies

Enzymatic oxidations were carried out with xanthine oxidase

(E.C. 1.1.3.22, grade I from buttermilk, 0.42units/mg prot, 24mg

prot/ml, Sigma Chemical), chloroperoxidase (E.C. 1.11.1.10, purified

grade, 1556units/mg prot, 15mg prot/ml, Sigma Chemical) and the

cytochrome P450 monooxygenase system (E.C. 1.6.2.4) in a rat liver

microsomal pellet (54). All enzymatic oxidations were carried out in

pH 7 phosphate buffer, ionic strength 0.5M. The E.C. number for each

enzyme consists of four numbers separated by periods. The first










number defines the class (one of six reactions) to which the enzyme

belongs. The next two numbers indicate subclass and sub-subclass.

The fourth number is a serial number given to each enzyme in its sub-

subclass (55).

In the chloroperoxidase oxidations 3ml of a 480pM TMP solution

in pH 7 phosphate buffer contained 1.6E-3M H202 and ca. 8.9E-8M of

chloroperoxidase (MW=42,000g/mole). In xanthine oxidase oxidations

the enzyme substrate molar ratio was 1:40. The typical concentration

of xanthine oxidase was 1pM. The molecular weight of xanthine

oxidase is ca. 260,000g/mol. All oxidations were monitored

spectrally and the oxidations were carried out in a 1cm quartz

cuvette.

Microsomal oxidations were performed with microsomal pellets

(obtained from Dr. James in the pharmacology department at the

University of Florida) which were prepared according to described

procedures (56). Briefly, a male rat liver is homogenized in a

buffer and a microsomal pellet is obtained after 1hr centrifugation

at 100,000 x g. The pellets were obtained from the livers of male

rats dosed with phenobarbitol to increase the cytochrome P450

concentration. The cytochrome P450 concentration in each pellet as

determined by standard procedures ranged from 2-44nmol/mg liver (57).

A UV difference spectrum was obtained to determine if TMP was a

substrate of cytochrome P450 (58,59). A difference spectrum reflects

changes in the electronic configuration of the iron porphyrin

prosthetic group of cytochrome P450 resulting from specific types of

interactions with a compound (58,59).













Since the microsomal fractions had been stored at -70C before

use it was necessary to verify that the enzymes were still active.

It was also important to test if a cyclic voltammetric assay, using a

rough PG working electrode, was practical. A cyclic voltammetric

assay was chosen because the microsomal solution is too opaque to

measure small absorbance changes. This was done by incubating (380C)

a 10ml solution containing 300l1 microsomal pellet, 9.8mM MgCl2,

564yM NADPH and 2004M uric acid in a pH 7 phosphate buffer, u=0.5M,

and recording a cyclic voltammogram every 5-10min of incubation. The

solution was shaken vigorously to dissolve the added uric acid. The

uric acid oxidation peak at 0.28V in the cyclic voltammogram

disappeared after 1hr incubation. This showed that the microsomal

enzymes were still active and that the cyclic voltammetric assay

could be used to monitor the oxidation of an electroactive species.

A similar experiment was performed on a 10ml solution containing

450PM TMP, 9.8mM MgCl2 and 280pM NADPH in phosphate buffer pH 7,

y=0.5M. Cyclic voltammograms were recorded every 10min of

incubation, up to 160min. After incubation the solution was

centrifuged for 10min and a UV spectrum of the supernatant was

recorded, from 200 to 400nm, to determine if any UV absorbing

products had been produced. The presence of any UV absorbing

products could not be determined because the solution was still too

opaque to give a clear spectrum.










4.7 Electrochemical Thermospray Mass Spectrometry (EC/TSP/MS)

Using the technique of EC/TSP/MS the electrolysis products can

be analyzed by mass spectrometry immediately after they are generated

electrochemically (60,61). In this technique the analyte is pumped

through an electrochemical cell held at the desired potential.

Electrolysis occurs and the products are forced into a heated

capillary tube of a thermospray ion source. Complete vaporization of

the liquid occurs to produce a superheated mist. This mist contains

molecules which are randomly positively and negatively charged. The

mass to charge ratio of these ions is then analyzed by a mass

analyzer.

The EC/TSP/MS experiments were carried out with a system

containing a coulometric cell (ESA, Inc.) which was coupled via a

Vestec thermospray interface to a Finnegan MAT TSQ45 triple

quadrupole mass spectrometer with an INCOS data system. Both counter

and reference electrodes were palladium. The working electrode was a

reticulated vitreous carbon block of large area (12cm 2), insuring ca.

100% electrochemical conversion efficiency. The cell volume was

5pl. The thermospray probe tip and source block temperature were

2500C and 2900C, respectively.

The conditions for TSP/MS were scan range m/z 125-300, electron

multiplier voltage 1000V and preamplifier gain 108 V A-1. For MS/MS,

the scan range was m/z 15 to 300, electron multiplier voltage 1600V,

preamplifier gain 108V A-1 collision energy 25eV and N2 collision gas

pressure 1.8mTorr. The mobile phase was 0.1M ammonium acetate






43



(pH 7), and the flow rate was 2.0ml/min. Twenty-five microliters of

a 300 yM solution was injected. The working electrode voltage was

varied from 0.8 to 1.2V. Both positive and negative ions were

analyzed by mass spectrometry.
















CHAPTER 5
CYCLIC VOLTAMMETRY OF TUBERCIDIN-5'-MONOPHOSPHATE



5.1 Oxidation and Reduction Behavior, Reversibility

The typical cyclic voltammogram of a 600M TMP solution in pH 7

phosphate buffer at a PG electrode shows an oxidation peak at ca.

0.90V and a broad shoulder at ca. 1.0V versus SCE (Fig. 5-1). Three

reduction peaks are evident at -0.56, -1.09, and -1.22V on the

reverse scan. These reduction peaks are present in the cyclic

voltammogram only after the oxidation peak has been scanned.

The sharp symmetrical peak shape of the oxidation peak suggests

an adsorption process. A study of the effect of concentration (Co)

of TMP on the oxidation peak current (i p) can give information about

the presence or absence of adsorption (section 3.1.1). In pH 7

phosphate buffer, P=0.5M, the plot of ip versus Co (Fig. 5-2a) shows

a change in the slope at 156gM. The slope changes from 1.17E+6 to

2.00E+5PA/M. Since the ip versus Co plot shows isotherm-like

behavior adsorption of TMP onto the electrode surface is indicated.

To further verify if the process is adsorption controlled the

dependence of ip on scan rate (v) was examined. For the TMP

oxidation peak in pH 7 phosphate buffer, y=0.5M, at a Co of 1001M

(before the change of slope of the ip versus Co plot), the plot of

logi versus logy has a slope of 0.8280.123 (Fig. 5-2b). At a C of
P o















scan rate: 200mV/s


-1.2


20 LA


I


0.5 0
Volts vs


-0.5
SCE


-1.0


-1.5


Figure 5-1 Cyclic voltammogram of TMP at a rough PG electrode, 600UM
solution in pH 7 phosphate buffer, y=0.5M.


1.0


I I I I


I


I


I


I











a.
. 400-

0300-

3200-

o 100-

SI I I I


4
Conc.


6
Mx 104


c.c.= .9904
slope = .828


1-


Log Scan


Figure 5-2


2
Rate, mV/s


Effect of concentration and scan rate on the oxidation peak
current of TMP in a pH 7 phosphate buffer, u=0.5M: a) plot of
peak current versus concentration for the oxidation peak of TMP,
scan rate 200mV/s; b) plot of log peak current versus log scan
rate for a 100pM TMP solution (c.c. = correlation coefficient).


c



L
(M
3


0

a-
0)
01
-I










998yM (after the change in slope of the ip versus Co plot) the same

plot has a slope of 0.5120.0998 (Fig. 5-3). Thus as indicated by

the slope values, at low Co the TMP oxidation peak is primarily

adsorption controlled and at high C it is primarily diffusion

controlled.

The oxidation appears to be irreversible since no corresponding

reduction peak is observed on the reverse scan. However, the peak

potential dependence (E ) on pH indicates that the oxidation is not a
simple irreversible process. For simple irreversible reactions Ep is

independent of pH. The absence of the reduction peak may be due to a

fast following chemical reaction occurring after electron transfer.



5.2 Effect of pH

The cyclic voltammetric behavior of TMP varies with buffer pH.

The cyclic voltammograms in Figure 5-4 illustrate this. As the

phosphate buffer pH increases from 2.8 to 9.5, the oxidation and

reduction peaks shift to more negative potentials and become

broader. At pH 9.5 the oxidation peak is a combination of two peaks

(Fig. 5-4c). In a pH 6.6 ammonium acetate buffer (Fig. 5-4d) the

oxidation peak is sharper than in a pH 7 phosphate buffer (Fig. 5-4b)

but the peak potentials are comparable. It is clear that the cyclic

voltammetric behavior of the oxidation peak is affected by buffer

composition.

The increased broadness of the oxidation peak in phosphate

buffer as the pH increases indicates that either the kinetics of the
























OOU
0)
oo
.)4)
O .
u2
U U)


. | I -


. I I I I I I I


0
cm





ipl
oE




0)


0


Q.
4)


0
0


0.





0
co


E-
-1
00
0N
0











0 C
L









0






0 1,

0)
to






0
L.




0.
,0 -
0OC
CL .0


(Vri)di 6o| m
I
Lt1
0

300
*-4
rz.


|


i











scan rate: 200mV/s


ao


I 50 IA


1.0 0.5 0
Volts vs


-0.5
SCE


-1.0 -1.5


Figure 5-4 Cyclic voltammetric behavior of TMP at a PG electrode in
different buffers, p=0.5M: a) pH 2.8 phosphate; b) pH 7
phosphate; c) pH 9.5 phosphate; d) pH 6.6 ammonium acetate.


-1.09


-1.2


-.53










reaction decreases or adsorption has less of an effect on the

oxidation. It seems contradictory that the kinetics of the reaction

should decrease when the oxidation peak shifts to less positive

potentials which indicates that the rate of the reaction slows down

when the oxidation becomes easier. However, the shift of the peak

potential to less positive values may be a result of different

acid/base forms of TMP being involved in the redox reaction at

different pH (62). The dependence of Ep on pH indicates that protons

may be involved in the rate determining step or in a preceding

chemical reaction. The complex nature of this reaction prevented

application of existing theories to the elucidation of the reaction

kinetics and the mechanism of the first e/H+ transfer steps.

The dependence of Ep on pH is defined by the Nernst equation

(equation 3-6) for a fast electrochemical process in which protons

are involved in the rate determining step. The Ep versus pH

measurements were carried out at scan rates of 5mV/s because at this

slow scan rate the system may be assumed to be at equilibrium and the

Nernst equation applies.

From the slope of a plot of Ep versus pH (Fig. 5-5) the number

of protons (r) and electrons (n) involved in the oxidation of TMP at

a PG electrode were estimated. In the pH region from 2 to 5, the

slope is 68mV/pH suggesting an r/n ratio of 1:1 (Fig. 5-5). At a pH

of ca. 5.7 the peak potential dependence on pH changes slope. This

corresponds closely to the pKa of 5.3 for tubercidin reported in the

literature (10, p. 316). Above pH 5.7 the slope is 31mV/pH,

suggesting an r/n ratio of 1:2.























C')
0)O


0) "
0> .


I* a
3.2
U n


I I I p I I I I


A'd3


*

04)
0





*C
OS
I 0


-4
0 L.
-4



o

00
0 *O




.a




0.4)
4-





(0 0
O1





44










0
o











CId


Cd
5-.. C
) 9
0 *






00
0






0-S
-4-c











The pKa value was also verified spectrally. In the pH region

from 2 to 4.6 the UV spectrum of TMP gives three absorption bands at

205, 225 and 270nm (Fig. 5-6a). In the pH region from 6.8 to 11

(Fig. 5-6b) the UV spectrum of TMP gives only two absorption bands at

210 and 270nm. The purine base of TMP probably exists in a

predominantly protonated form below the pH=pKa and a predominantly

neutral form above the pH=pKa since structurally similar adenine is

predominantly protonated below pH=pkal (4.2) and predominantly

neutral above pH=pKal (63). Crystallographic data show that the

cationic proton of the aminopurine is attached to the 1-nitrogen of

the purine ring (Fig. 1-2) (64). This may also be the case for the

structurally similar purine base of TMP. At a pH below the pKa of

TMP the phosphate group on TMP is negatively charged. The net result

may be a zwitterionic TMP molecule at pH below 5.7.



5.3 Electrode Surface Effects

The cyclic voltammetric behavior of TMP was studied at PG, GC

and pretreated GC electrodes. The cyclic voltammetric behavior was

affected by the nature and condition of the carbon-electrode surface.

Pyrolitic graphite is formed by depositing carbon from the vapor

phase on the surface of a substrate, usually a metal. This process

results in a highly oriented polycrystalline form of carbon. The

substrate will always contain minute flaws which serve as nucleation

sites for the growth of cones. These cones are the main structural

characteristics of PG (26, p. 303; 65,66). To activate the PG


















3.5 a(


200 220


240 260 280 300 320 340 360 380 400
WAVELENGTH,nm


220 240 260 280 300 320 340 360 380
WAVELENGTH,nm


Figure 5-6 Spectral behavior of TMP (ca. 150pM) in phosphate
buffers, P=0.5M, of pH a) 2.8-4.6 and b) 6.8-11.


3.0

2.5


0.0










surface it is usually polished by grinding on 600 grit silicon

carbide paper before voltammetry.

Glassy carbon is formed by slowly heating a polymeric resin in

an inert atmosphere. A carbonization process starts as the

temperature increases above 3000C where oxygen, nitrogen, hydrogen

and anything else is eliminated until only carbon is left. The

material obtained by this process is referred to as vitreous or

glassy carbon. The structure of GC is believed to be a network of

tangled aromatic ribbon molecules which are cross-linked by highly

strained carbon-carbon bonds (26, p. 308; 30). It is isotropic,

impermeable to gas, electrically conductive and resistant to chemical

attack (26, p. 308). The surface is polished to a smooth mirror-like

finish with polishing alumina before use as an electrode.

Figures 5-7b and d show the cyclic voltammetric behavior of a

600uM TMP solution in pH 7, p=0.5M, phosphate buffer at a PG and GC

electrode. Figures 5-7a and c show the cyclic voltammograms in

buffer at these electrodes. The TMP oxidation peak is sharp and

symmetrical at PG and broad at GC. This suggests that the surface

plays a role in the oxidation of TMP. It was determined that the TMP

oxidation is an adsorption controlled process at PG from studies of

the behavior of peak current on concentration and scan rate (Figs.

5-2 and 5-3).

The diffusion nature of the TMP oxidation peak at a GC electrode

was confirmed by a linear plot of ip versus concentration for TMP

solutions of 63-575gM in pH 7 phosphate buffer, u=0.5M.































Figure 5-7 Cyclic voltammograms in a a) pH 7 phosphate buffer,
0=O.5M, at rough PG; b) 600vM TMP solution in pH 7
phosphate buffer at rough PG; c) same as a, at GC; d)
same as b, at GC.









scan rate: 200mV/s


50 PA



b









C


I I


Volts vs


I I -1.0 I
0 -0.5 -1.0 -1.5


d





1.5


1.0


SCE


I


I I


0.5










To provide additional information about the electrochemical

behavior of TMP at GC the dependence of i on v at GC was also

examined. Scan rates ranged from 20 to 500mV/s. The plot of logip

versus logv for a 106iM TMP solution in pH 7 phosphate buffer has a

slope of 0.5630.754. At a concentration of 575pM, the same plot has

a slope of 0.5110.231. Thus at both concentrations the TMP

oxidation peak at GC occurs predominantly by diffusion.

Cyclic voltammetry at fast scan rates shows no corresponding

reduction peak indicating that the oxidation peak at GC is also

irreversible.

Since ip is proportional to electrode area and the area of the

GC and PG electrodes are not the same; the peak currents cannot be

directly compared. The normalized peak current, pA/cm2, however, can

be directly compared. To determine the normalized peak current the

electrode area was estimated using chronocoulometry (section

3.4.2). The results indicate that the GC electrode is about twice as

large as the PG electrode (84E-4cm2 versus 49E-4cm2). Using these

values, the normalized peak current for the oxidation peak at GC and

PG in Figure 5-7 is 501 and 1708IA/cm respectively.

Electrochemical pretreatment in many cases improves the

performance of GC electrodes (i.e., enhances the peak currents and

improves reversibility) (44,67,68). It has been suggested that

electrochemical pretreatment introduces or alters functional groups

on the electrode surface. These functional groups may facilitate

electrode reactions in some way. Specific chemical interactions such

as adsorption or changes in hydration of the electroactive species










may also be occurring as a result of pretreatment (69,70).

Electrochemical pretreatment of GC was carried out to observe the

effects on the TMP oxidation peak.

Electrochemical pretreatment was performed by applying a

constant positive potential of +1.8V versus SCE for 8min to the GC

electrode, followed by 2-5 cyclic scans towards negative potentials

first (+1.8 to -1.7V) at 20mV/s (68). These cyclic scans are carried

out until the background peaks at ca. -1 and +1V disappear (section

4.1).

Figure 5-8 shows the cyclic voltammograms of a TMP solution in

pH 7 phosphate buffer, y=0.5M, at GC before and after electrochemical

pretreatment. The cyclic voltammograms in a pH 7 phosphate buffer

alone (Fig. 5-9) illustrate how the background is affected by

pretreatment. The cyclic voltammogram of TMP in Figure 5-8b

illustrates that pretreatment does not improve the appearance of the

oxidation peak at ca. 1V which is difficult to discern from the

background. Thus electrochemical pretreatment did not enhance the

TMP oxidation peak current or improve the kinetics of the oxidation.



5.4 Behavior in Organic Solvents and Solvent Mixtures

The cyclic voltammetric behavior of TMP is also affected by the

composition of the solvent. TMP gives no oxidation or reduction

peaks in a solution of dimethylformamide (DMF) with 0.05M tetrabutyl

ammonium perchlorate (TBAP) added as the supporting electrolyte.

Adding successive amounts of organic solvent (DMF) to a solution of











T
20A


-1.200


E (UOLT I


T
50A
SOIf


-1.200


E (UOLT)


Figure 5-8


Cyclic voltammograms of a 575uM TMP solution in pH 7
phosphate buffer, P=0.5M, at a glassy carbon electrode:
a) before pretreatment and b) after pretreatment. Scan
rate 200mV/s.











ZOi


-1.200


E (UOLTI


T
oif


-1.200


E(UOLT)


Figure 5-9


Cyclic voltammograms in pH 7 phosphate buffer, P=0.5M, at
a GC electrode: a) before pretreatment and b) after
pretreatment.











TMP in pH 7 phosphate buffer, decreases the oxidation peak at a PG

electrode (Fig. 5-10). No shift in peak potential is observed.

After addition of 40% organic solvent, the peak has decreased 56%.

After addition of 30% organic solvent, the shape and the magnitude of

the TMP oxidation peak at PG is similar to that at GC for the same

concentration of TMP in pH 7 phosphate buffer (Fig. 5-7d).

Mixed water organic solvent systems are commonly used to prevent

adsorption of surface active organic compounds (70). This type of

medium may lead to an increase in solvent-solute interactions over

electrode solution interactions. The peak due to TMP at a PG

electrode may decrease in this medium because the organic solvent may

prevent adsorption of TMP onto the PG electrode surface.

At a GC electrode the oxidation is predominantly controlled by

diffusion (section 5.3). If adding organic solvent prevents

adsorption of TMP onto the PG surface, then it should not affect the

electrochemical behavior of TMP at GC. Figure 5-10 shows that

addition of organic solvent (DMF) to an aqueous solution of TMP has

little effect on the ip of TMP at GC. These results indicate that

addition of organic solvent may be preventing adsorption of TMP onto

the PG electrode surface. It is clear that TMP behaves differently

at GC and PG.

TMP gives no cyclic voltammetric peaks in an organic solvent.

If water is added to a solution of TMP in DMF with TBAP added as a

supporting electrolyte no oxidation or reduction peaks appear in the

cyclic voltammogram when 40% water is added. If phosphate buffer is



























S5C

4C

3C

2C

1C






Figure 5-10


A PG
o GC


/o Organic


Plot of TMP oxidation peak current at PG and GC versus %
DMF added to a 600pM TMP solution in pH 7 phosphate
buffer, P=0.5M.






63


added to a similar solution of TMP instead of water, the poor

solubility of phosphate causes it to precipitate out of solution.
















CHAPTER 6
DETERMINATION OF THE n-VALUE OF TUBERCIDIN-5'-MONOPHOSPHATE



6.1 Coulometry

When the constant potential electrolysis of TMP was carried out

the coulometer continued to accumulate charge after TMP was

completely electrolyzed as shown by HPLC. Cyclic voltammograms and

UV spectra did not clearly show when the electrolysis of TMP was

complete because products which formed had oxidation peaks and UV

absorbance bands that overlapped with those of TMP.

When TMP is completely electrolyzed in a pH 7 buffer, u=0.5M, at

rough PG, as indicated by the disappearance of the TMP peak in the

HPLC, the n-value was determined to be 11 by coulometry (section

3.2). This n-value is independent of concentration and depends on

pH. When TMP is electrolyzed in pH 9.5 buffer, u=O0.5M, the n-value

was determined to be 5.2 by the same method. A background correction

was carried out at a typical electrolysis potential of ca. 0.97V

versus SCE. The values of charge that were obtained were ca. 1% of

the total charge value obtained during electrolysis of TMP. The

background corrected values are reported.

The accumulation of charge which is observed after TMP has been

completely electrolyzed indicates that there may be a catalytic

oxidation of water occurring at the typical electrolysis potentials

in solution in presence of TMP and/or its oxidation products in










solution. There are reports which suggest that functional groups on

the electrode surface can cause the catalytic oxidation or reduction

of species in solution (38,40). It is also possible that the initial

electrolysis products of TMP are being further oxidized. This is

supported by the product analysis (section 8). Thus the n-value

obtained by this method may be inflated.



6.2 Peak Current Comparisons

Since the n-value of TMP obtained from coulometry may be

inflated the n-value was also estimated by comparing the cyclic

voltammetric peak current of TMP to that of structurally similar

compounds. The ip for reversible and irreversible systems is defined

in equations 3-1 through 3-4. The cyclic voltammetric i values of

compounds with known n were compared to the ip value for TMP. By

assuming that both systems were diffusion controlled, the Do values

were the same, using the same electrode, concentration and scan rate,

equations 6-1 and 6-2 were used to determine the n-value of TMP.

Equation 6-1 compares the ip of a reversible (r) standard to the ip

of an irreversible (ir) system (TMP).

diffusion controlled systems


i (ir)u 1.11[C *n(na) /2](ir) u
u (6-1)
i p(r)k [C *n3/2(r)k



i (ir) C *n(na) 1/2(ir)
i (ir)u 1na)/2(r) (6-2)
p u [C0*n(ana) ](ir)
o a. u










If both systems were assumed to be adsorption controlled (r *
o
values and electrode area the same), equations 6-3 and 6-4 were used

to determine the n-value of TMP.

adsorption controlled systems


i (ir) 1.47[nan a](ir)u
p u a-- u(6-3)
i (r) 2
p k n (r)k


i (ir) [nan ](ir)
p(ua u 6-4)
i (ir) [nan ](i r) k6
p k a k

Equations 6-1 and 6-3 compare the ip of a reversible (r)

standard to the ip of an irreversible (ir) system (TMP). From cyclic

voltammetry TMP oxidation appears to be irreversible and adsorption

controlled (section 5.1). Equations 6-2 and 6-4 compare the ip of

two irreversible systems (r), where k and u represent the known and

unknown systems, respectively.

Equations 6-5 and 6-6 could be used to compare the ip values of

an adsorption and diffusion controlled system but the values for ro*

and Do are not known. Therefore, to estimate n by comparing


Pads (nanaF 2o *)ads
S ( 731,142RT(n3/2 D012 C */2)


ipads (ir) (n 1/2 r 1/2 1/2V)ads (6-6)
idiff (r 812,682RT[n(an ) 1/2C *D v 2]


cyclic voltammetric ip values, both systems were assumed to be either

adsorption or diffusion controlled.










The compounds that were used for oxidation peak comparisons were

xanthine, guanine and uric acid. The oxidations of guanine and

xanthine are four-electron irreversible processes and the oxidation

of uric acid is a reversible two-electron process (6,7,33,34). These

oxidations are primarily adsorption controlled at a rough PG surface

(1,2,6,7,33,34). Equations 6-1 and 6-3 were used to determine the n

of TMP by comparison to uric acid and equations 6-2 and 6-4 were used

to determine the n of TMP by comparison to xanthine and guanine. The

ana values were estimated by measuring the shift in Ep at two sweep

rates using equation 3-8. Table 6-1 illustrates the results that

were obtained which estimate the n-value of TMP to be between 4.57

and 7.32 if both systems are assumed to be diffusion controlled and

between 4.68 and 10.90 if both systems are assumed to be adsorption

controlled.

In order to determine the validity of this approach n-values of

compounds whose n-values have been reported in the literature were

determined in a similar manner. In this study, uric acid, guanine

and xanthine were compared with each other. The results in Table 6-2

show that estimating the n of xanthine by comparing it to uric acid

using equation 6-1 gives an n of 4.32 for xanthine (+8% error).

Estimating the n of uric acid by comparing it to xanthine using

equation 6-1 gives an n of 1.90 for uric acid (-5% error). The

errors of other comparisons range from -32 to +132%. It appears that

the % error is less when equations 6-1 and 6-2 (which assume a

diffusion controlled system) are used to estimate n.










Table 6-1


Estimation of the n-Value of TMP from Cyclic Voltammetric
Peak Currents.


1 2 3 4
compound ana conc,pM 1p, A n nTMP


Guanine (ir) .546 101.6 24.453.97 4(7) 4.575,4.687
Xanthine (ir) .512 101.3 15.942.4 4(6) 6.775,6.747
Uric acid (r) 101.0 13.09.89 2(33,34) 7.326,10.908
TMP (ir) .515 101.0 27.02.33 ? --

1 an determined from equation 3-8. The scan rates were 100 and
208mV/s.
2 All solutions were prepared in pH 9.5 phosphate buffer, u=0.5M,
working electrode was rough PG.
4 Determined at a scan rate of 100mV/s.
5 Literature value, references in ( ).
5 Determined from equation 6-2, u=TMP.
6 Determined from equation 6-1, u=TMP.
8 Determined from equation 6-4.
8 Determined from equation 6-3.
* Uric acid is a reversible reaction; therefore, it has no ana value.


Table 6-2


Validity of Peak Current Comparisons of Estimation of
n-Values.


comparison1 ncalc. n6 %error ncalc. n6 %error


U vs G G=6.402,9.303 4(7) +61,+132 U=1.462,1.313 2(33) -28,-34
U vs X X=4.322,6.473 4(6) +8,+62 U=1.902,1.573 2(33) -5,-21
G vs X X=2.70 ,2.785 4(6) -32,-30 G=5.924,5.755 4(7) +48,+44

1 G-guanine, U=uric acid, X-xanthine.
2 Determined using equation 6-1.
3 Determined using equation 6-3.
Determined using equation 6-2.
4 Determined using equation 6-2.
5 Determined using equation 6-4.
6 Literature value, references in ( ).
Note: Values for i and an were given in Table 6-1.
Experimental conditions were the same as in Table 6-1.










The calculations summarized in Tables 6-1 and 6-2 were made

using equations 6-1 to 6-4. The calculations are based on the

assumption that both systems were diffusion controlled or both

systems were adsorption controlled. Uric acid, xanthine, guanine and

TMP are primarily adsorption controlled systems at the PG surface.

However, the results in Table 6-2 indicate that the % error is less

than when the systems are assumed to be diffusion controlled.

Diffusion coefficients and electrode area were assumed to be the same

which may not be valid. For example, the electrode area changes

every time the electrode is resurfaced. Typical precision of

measurements at a PG electrode is ca. 5-10%. The bulk concentration

and the concentration at the electrode surface may also be different.



6.3 Chronocoulometry

Because of the large errors associated with the determination of

n-values by cyclic voltammetric peak current comparisons, the n-value

of TMP was also estimated using chronocoulometry. In a chronocoulo-

metric experiment the electrode potential is stepped from a potential

where no electrolysis takes place (Ei) to a potential where electro-

lysis occurs at a maximum diffusion controlled rate (Ef) (section

3.4.1). Chronocoulometry can be used to estimate the n-value of TMP

by comparing its slope to that of a compound whose n-value has been

reported. By using the same electrode and assuming the diffusion

coefficients are the same equation 3-14 can be used for this

purpose. Adenosine-5'-monophosphate (AMP) was used for this










comparison because it is structurally similar to TMP thus their

diffusion coefficients should be very similar.

Although the n of AMP has not been determined it was assumed

that the n of AMP equals that of adenine, which is 6 (3). This

assumption was based on the report that the presence of a ribose

group did not affect the value of n (5). The n-value of TMP that was

estimated using this approach was 3.32. The conditions and results

of this experiment are given in Table 6-3.




Table 6-3 Estimation of the n-Value of TMP Using Chronocoulometry


slopes conc,pM2 Ei,mV Ef,mV value


AMP 7.9167 1017 300 1250 6(3)3

TMP 4.3107 998 200 925 3.324


1 Determined from a chronocoulometric experiment (section 4.4).
2 All solutions were prepared in pH 7 phosphate buffer, y=0.5M.
3 Literature value, reference in ( ).
Determined from equation 3-14.





To validate this approach a similar comparison was carried out

to estimate the n of uric acid, xanthine, and guanine. Table 6-4

shows that this approach leads to an error of -36 to +57%. The error

may be due to the assumption that Do and A are the same for all of

these compounds when using equation 3-14 for the calculations.










Table 6-4 Validity of Chronocoulometric Comparison of Estimation of
n-Values.


compound slope2 conc,M3 k4 u5 nucalc6 nu7 %error


G .7517 101.9 G X 3.86 4(6) -3.5
X .7164 100.6 X G 4.14 4(7) +3.5
U .2286 96.96 U G 6.26 4(7) +57
G U 1.28 2(33) -36
U X 6.04 4(6) +51
X U 1.32 2(33) -34


2 G-guanine, X=xanthine, U=uric acid.
2 Determined from chronocoulometry (section 3.3) pulse width 500ms
for all compounds, for guanine and xanthine Ei=500mV and Ef=675mV,
for uric acid Ei=150mV and Ef=325mV.
3 All solutions were prepared in pH 7 phosphate buffer, u=0.5M.
4 k-known system in equation 3-14.
5 u=unknown system in equation 3-14.
6 Determined using equation 3-14.
Literature value, reference in ( ).





When the n-value is obtained by comparing the slopes of a plot

of charge versus t1/2 for xanthine and guanine (irreversible

processes) there is a small error, 3.5%, between estimated and true

n-values. When comparing an irreversible and reversible process

(e.g., xanthine to uric acid and guanine to uric acid), the error

ranges from -36 to +57%.

Using coulometry the n-value of TMP was estimated to be 11 when

the TMP peak disappeared in the HPLC. Under these conditions further

oxidation of initial electrolysis products or catalytic oxidation of

solvent may also be contributing to the n-value. The n-value of TMP









that was determined from cyclic voltammetric peak current comparisons

ranged from 4.57 to 7.32 when both systems were assumed to be

diffusion controlled. The n-value ranged from 4.68 to 10.90 when

both systems were assumed to be adsorption controlled. However, this

method of determination could be in error by as much as +132% as

indicated in Table 6-2. Chronocoulometric comparisons suggest that n

of TMP is ca. 3 assuming n of AMP is 6. This method of n

determination gives a small error (3.5%) when two irreversible

processes are compared and a significant error (-36 to +57%) when an

irreversible and a reversible process are compared. The discrepancy

between n-values determined by coulometry, cyclic voltammetry and

chronocoulometry probably results from the different time scale of

these experiments.
















CHAPTER 7
CONSTANT POTENTIAL ELECTROLYSIS OF TUBERCIDIN-5'-MONOPHOSPHATE



7.1 Introduction

Using the technique of constant potential electrolysis the

stable intermediates and products from an electrode reaction can be

generated for further analysis (section 4.2). The electrolysis of

TMP was monitored by cyclic voltammetry, UV spectroscopy and HPLC to

provide information on the formation of intermediates and products.

It was concluded that TMP was completely electrolyzed when the

TMP peak disappeared in the HPLC (Fig. 7-1e). The UV spectrum at

this point resembles trace 4 in Figure 7-2 and the cyclic voltammo-

gram (Fig. 7-3c,d) shows a broad oxidation peak at potentials of TMP

oxidation. Separation of the products by GPLC gives two peaks (Fig.

7-4a). From the spectral changes observed (Figure 7-2) it is clear

that after TMP is completely electrolyzed oxidation of products must

continue. This is supported by cyclic voltammetry (Fig. 7-3) which

shows an increase in the reductions peaks at -1.09 and -1.2V and HPLC

(Fig. 7-1f) and GPLC (Fig. 7-4b) which show the formation of new

products after continued electrolysis at TMP oxidation potentials.

It is clear that cyclic voltammetry, UV spectroscopy and GPLC cannot

be used to determine when TMP is completely electrolyzed. Complete

electrolysis is easily determined using HPLC. If constant potential

electrolysis is continued after TMP is completely electrolyzed,

















0.20-


0.15.


E 0.25-
C
in
( 2
" 0.20-
k-

w 0.15-
U
z
m 0.10-
0
C 0.05-
4


TMP


BEFORE
a














25 MIN.
b










1 HOUR
C


TMP


0 5


Figure 7-1


15 20


HPLC during the electrolysis of a 600pM TMP solution in
pH 7 phosphate buffer, u=0.5M, at a rough PG electrode:
a) before; b) 25min; c) 1hr; d) 2hr, e) 3hr, 22min; f)
7hr.
















0.30-
2 HOURS
0.25 d
w
uE_
z 0.20-
C 0 0.15-
m<

0.10-

0.05- n

\I I TMP i i


0.20- I
w 3HR. 22MIN.
UE e
z 0.15-
am
0
( 0.10-
'D<
< <
0.05- 3

II
w i
u E 0.10-
Z 7 HOURS
tM E f
Ix 0.05-
O0
< i ,, i i
I I I ( I I ) I
0 5 10 15 20
(MI N)


Figure 7-1--continued.


















. ... .,.'- "-.4H
**,, b .d A
....- l N H


*, 'I '. .
"Z /" / .











solution in pH 7 phosphate buffer, =05M, at a rough PG






i electrode.
solution in pH 7 phosphate buffer, v=0.5M, at a rough PG
electrode.













a

501 A


b



c





d


scan rat


Figure 7-3 Cyclic voltammograms of a 600VM TMP solution in a pH 7
phsophate buffer, 4=0.5M, during electrolysis at a rough
PG electrode: a) before; b) 2hr, 2min; c) 7hr, 2min.


e: 200mV/s -1.-1.2 before









2 hr.








7hr.














1.0 0.5 0 -0.5 -1.0 -1.5
Volts vs SCE

















a

E FREE PHOSPHATE
0
N 3.0-
(I

020
U 2.0-
z
0
O A B
m 1.0-




8 24 40 56 72 88 104 120 136 152 168



b

FREE PHOSPHATE

o 3.0-

C
I-

W 2.0 -
z D

0 1.0


A

8 24 40 56 72 88 104 120 136 152 168
ML COLLECTED



Figure 7-4 GPLC separation of TMP electrolysis products from a 600pM
TMP solution in a pH 9.5, P=0.5M, phosphate buffer after
a) 4-12hr and b) 60hr of electrolysis.









further oxidation of products occurs (Fig. 7-1f). This is clearly a

slow process since the GPLC (Fig. 7-4b) shows that after 60hr of

electrolysis only two new product peaks formed and that one of the

first formed products finally disappeared.

The time required to completely electrolyze TMP was a function

of the condition of the PG electrode surface. At a rough surface the

time ranged between 1-4hr and at a surface which was not roughened

the time ranged from 8-12hr. The electrolysis times at these

different surfaces varied because the degree of roughness was not

well controlled when the electrode was manually polished on silicon

carbide paper. Changes in the UV spectra show that the oxidation

pathway was also affected by the condition of the electrode surface

(section 7.3).

Electrolysis of TMP was carried out in phosphate buffer (pH 2.8,

7 and 9.5), u=0.5M, to optimize the separation of products from

phosphate by GPLC (section 4.3). The optimum pH for separation of

products from phosphate was 9.5. Since the objective of this

investigation was to provide insight into TMP reactivity at physio-

logical pH it was important to verify if the same products were

produced during electrolysis in both pH 9.5 and 7. Monitoring the

electrolysis of TMP in pH 9.5 and 7 phosphate buffers with HPLC

(section 7.4) verified that the same products were produced in both

pH 7 and 9.5.

The specific cyclic voltammetric, spectral and HPLC behavior of

TMP during electrolysis in pH 7 and 9.5 phosphate buffers will be

discussed.










7.2 Cyclic Voltammetry

Figure 5-4 showed that the cyclic voltammetric behavior of TMP

before electrolysis depended on pH. The cyclic voltammetric behavior

also changed with pH during electrolysis. However, the same-major

products which were separated and analyzed formed at both pH 7 and

9.5. This was verified by HPLC (section 7.4).

Figure 7-3 shows typical cyclic voltammograms obtained during

electrolysis of TMP in a pH 7 phosphate buffer. The reduction peak

at -1.2V increases throughout the electrolysis and increase further

after TMP has been completely electrolyzed, according to HPLC. The

reduction peak at -0.56V disappears before TMP is completely

electrolyzed. A broad oxidation peak forms at 0.89V when the amount

of TMP has decreased by 77% as determined by HPLC. This new broad

oxidation peak does not increase further after TMP is completely

electrolyzed. The products) oxidized at this potential decreases

very slowly. The formation of this broad oxidation peak obscures the

TMP oxidation making it difficult to determine if TMP is completely

electrolyzed using cyclic voltammetry. Electrolysis at pH 7 was

carried out until no significant changes were observed by cyclic

voltammetry and UV spectroscopy.

Figure 5-4c showed the cyclic voltammogram of TMP in a pH 9.5

phosphate buffer before electrolysis. As the electrolysis proceeds

the reduction peak at -1.32V increases. Before TMP has been

completely electrolyzed a broad oxidation peak forms at slightly more

positive potentials than the TMP oxidation peak. This oxidation peak

does not increase as the electrolysis proceeds but rather slowly










decreases. The reduction peaks at -.66V and -1.22V disappear before

TMP is completely electrolyzed. This is similar to the behavior at

pH 7. When electrolysis in pH 9.5 is carried out for a very long

time, 60hr, all oxidation and reduction peaks in the cyclic

voltammogram disappear and the product absorbance at 305nm in the UV

spectrum disappears. The UV spectrum at this point shows no

characteristic peaks.


Figures 7-2 and

electrolysis at a PG

has not. Since this

TMP at physiological

surface was confined

has been shown to be

(71).


7.3 UV Spectra

7-5 illustrate the spectral changes during

surface which has been roughened and one which

investigation was concerned with the behavior of

pH, the comparison of effect of electrode

to pH 7. The chromophore circled in Figure 7-2

responsible for the absorption peak at 270nm


The general behavior during electrolysis at a rough PG electrode

as determined by UV spectroscopy was similar regardless of pH. The

spectral changes are a function of electrolysis time as indicated in

Figure 7-2. During electrolysis the band of TMP at approximately

210nm decreases and a new band grows in at 230nm. The band due to

TMP at 270nm decreases and shifts to 265nm. Simultaneously a new

band forms at 305nm. The shift to 265nm is observed in both pH 7 and

9.5.

The spectral changes during electrolysis at an unroughened PG

electrode (Fig. 7-5) are similar to those observed at a rough PG




















































400


280 300 320
WAVELENGTH (nm)


Figure 7-5


UV spectra during the electrolysis of a 100pM TMP
solution in pH 7 phosphate buffer, p=0.5M, at an
unroughened PG surface.


w 2.
z

~2.
0
V)









electrode (Fig. 7-2) except that during electrolysis at an

unroughened surface isosbestic points are held at 217, 247 and

292nm. The presence of isosbestic points suggests that the reaction

proceeds through a simultaneous (A + B + C) rather than a consecutive

reaction pathway (A + B + C) (72).



7.4 HPLC

Many of the results which were described are based on a reverse

phase HPLC assay that was developed and used to monitor the

electrolysis of TMP. This HPLC method is based on that developed by

P. Brown and co-workers for the separation of nucleosides,

nucleotides and bases (73,74). The conditions of this experiment

were described in section 4.2. Typically, in reverse phase HPLC

neutral compounds are retained while ionic compounds are only

minimally retained (73). As shown by our results, TMP was relatively

strongly retained using a pH 4.7-5.1 mobile phase and a nonpolar C-18

column. The purine base of TMP is believed to exist in a

predominantly protonated form in this pH region (section 5.2) but the

phosphate group of the nucleotide is negatively charged in this pH

region. The net result may be zwitterionic and possibly neutral

molecule in this pH region. For this reason the molecule may be

retained under these conditions (73).

A typical electrolysis in pH 7 phosphate buffer monitored by

HPLC is shown in Figure 7-1. Immediately after initiation of

electrolysis a major product with very short retention time formed,

peak I (Fig. 7-1b). In addition, three other peaks formed, peaks II,










III and IV (Fig. 7-1c). As electrolysis proceeds, peaks I, II, III,

and IV increase and a new peak, V, forms late in the electrolysis

when TMP is ca. 90% oxidized (Fig. 7-1d). When TMP is completely

oxidized (Fig. 7-ie) peak I has decreased ca. 41%, and peak III has

decreased ca. 66% from their maximum absorbance values. Peak II

continues to increase and peak VI forms after TMP is completely

oxidized. This indicates that peak VI must form from other products.

When electrolysis in a pH 9.5 phosphate buffer is monitored by

HPLC, peak V does not form but another product peak forms instead

which has a retention time between that of peak III and TMP. This

peak is the same height as peak II and decreases when TMP is

completely electrolyzed.

The HPLC results indicate that at least 6 UV absorbing products

form. Five peaks form before TMP is completely electrolyzed (peaks

I-V) and one forms after TMP is completely electrolyzed (peak VI).

The products corresponding to HPLC peaks I, II, IV and VI are the

major UV absorbing products at 225 nm after TMP has been completely

electrolyzed.
















CHAPTER 8
ANALYSIS OF TUBERCIDIN-5'-MONOPHOSPHATE ELECTROCHEMICAL
OXIDATION PRODUCTS


8.1 Separation by Gel Permeation Liquid Chromatography

Typical GPLC separations were shown in Figure 7-4 for different

electrolysis times. The experimental conditions were discussed in

section 4.3. Figure 7-4a shows a typical GPLC separation after 4-8hr

electrolysis of TMP at a rough PG electrode in pH 9.5 phosphate

buffer. At this point TMP has been completely electrolyzed and two

product peaks are present, A and B. Figure 7-4b shows a separation

after 60hr of electrolysis in a pH 9.5 phosphate buffer. The

fractions eluting under A and B in Figure 7-4a and 7-4b are the same

compounds) based on spectral analysis. The products) eluting under

peak A is still present while those eluting under peak B have

decreased considerably. Two new product peaks C and D have formed.

Peak D is a mixture of the product eluting under C and another minor

component based on analysis by GC. Since the other component under

peak D is a minor component in the mixture it was not analyzed

further.

Separation of the solutions electrolyzed at intermediate times

(12-20hr) between the separations in Figure 7-4 showed one new very

broad product peak which eluted between A and B and coeluted with

phosphate thus it was not analyzed further. Since HPLC analysis of

the products eluting under GPLC peaks A and B verified that the










products did not decompose on the column it was concluded that the

products under GPLC peaks C and D were not decomposition products but

formed as a result of further electrolysis. Since GPLC peak B has

decreased considerably after 60hr electrolysis and peaks C and D have

formed it is quite possible that peak C or D formed from the

compounds) eluting under GPLC peak B.



8.2 Analysis of Products with High Pressure Liquid
Chromatography, Cyclic Voltammetry and
Ultraviolet Spectroscopy

Analysis of the fractions eluting under GPLC peak A in Figure

7-4a using HPLC is shown in Figure 8-1a. This analysis indicates

that primarily one component elutes under this peak with a retention

time similar to HPLC peak I formed during electrolysis of TMP in

Figure 7-1b. A UV spectrum of the fractions under GPLC peak A is

shown in Figure 8-2. This product exhibits an absorption at 225nm.

Cyclic voltammetry of this compound shows no oxidation or reduction

peaks.

A UV spectrum of the fractions eluting under GPLC peak B is

shown in Figure 8-3. The fractions under this GPLC peak give an

absorption maximum at 235 and 305nm. Analysis of the fractions

eluting under GPLC peak B using HPLC is shown in Figure 8-1b. This

analysis shows primarily two HPLC peaks which correspond to HPLC

peaks II and VI formed during the electrolysis of TMP in Figure

7-1f. Analysis of both peaks in Figure 8-1b, at 225nm, 260nm and

305nm, using HPLC indicates that the products corresponding to both

peaks have a significant absorbance at 225 and 305nm and a minimum

absorbance at 260nm. Thus the two UV absorbing components which


















w
z E 0.02

0 0.01
<






0.03-
w
UE
z


m0 0.02-
0
In -
CO,<
'< 0.01-







0.02-
zE

0
0C 0.01-










Figure 8-1


a ~


10
(M IN)


HPLC of compounds under GPLC peaks a) A, b) B and c) C
(flow rate Iml/min).




























































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