Physicochemical characterization and stability of doxorubicin in aqueous solutions

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Physicochemical characterization and stability of doxorubicin in aqueous solutions
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Table of Contents
    Title Page
        Page i
    Dedication
        Page ii
    Acknowledgement
        Page iii
        Page iv
    Table of Contents
        Page v
        Page vi
        Page vii
    List of Tables
        Page viii
        Page ix
    List of Figures
        Page x
        Page xi
        Page xii
        Page xiii
    Abstract
        Page xiv
        Page xv
    Chapter 1. Introduction
        Page 1
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    Chapter 2. Materials and methods
        Page 52
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    Chapter 3. Results and discussions
        Page 84
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    Chapter 4. Conclusions
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    Appendix. Buffers
        Page 196
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    Reference list
        Page 198
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    Biographical sketch
        Page 220
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Full Text











PHYSICOCHEMICAL CHARACTERIZATION AND
STABILITY OF DOXORUBICIN IN
AQUEOUS SOLUTIONS



















By

ANUP ZUTSHI


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


1994
































To my parents- who have patiently endured my excuses and

explanations for the delay in preparing this manuscript.













ACKNOWLEDGEMENTS


I am extremely grateful to Dr. Hans Schreier, Associate

Professor of Pharmaceutics, for his continuous academic and

financial support, encouragement, patience and tolerance

during the time it took to complete the work and generate this

document. His infallible confidence in my abilities and his

friendly disposition made working for him an enjoyable

experience.


I am deeply indebted to Dr. Edward R. Garrett, Professor

Emeritus of Pharmaceutics, for giving me the opportunity to

work with him. During my years of association with him, he has

at various times been a 'father figure', persevering teacher

and an advisor. I attribute my scientific background and

knowledge to his unique and effective style of graduate

instruction. I am also grateful to Mrs. Irene B. Garrett, for

her warm and caring disposition and her constant support

especially during some rough times.


I am also in the debt of Dr. Shangxian Chen, Visiting

Professor, University of Florida. His extensive research

experience and fundamental sciences background helped me in

understanding and solving many problems. I am also extremely

grateful to Mrs. Chen for helping with the experiments.

iii







I am also indebted to all my committee members, Drs.

Kenneth B. Sloan, Mark Longer and Kathryn Williams, for their

scientific assistance at various times during my research.


Of the many colleagues that worked with me, Robert

Townsend has become a good friend. His infinite patience

during periods when I would bounce ideas off him and his

intelligent feedback were extremely helpful throughout my

graduate education. Other colleagues who have established long

lasting friendships and who at various times have helped me

unselfishly and unconditionally in preparing and compiling

documents such as these are Jaimini Patel, Janice Cacace,

Victoria Saldajeno, Cary Mobley, Vivek Shenoy, Patricia Khan

and Sharon Fussell-Carter.


I am extremely thankful to my parents NandLal and Shiela

and my brothers Aroop and Ajoy for being a patient and

supportive family at all times.













TABLE OF CONTENTS



ACKNOWLEDGEMENTS . . . . . . . . . iii

LIST OF TABLES . . . . . . . . .. viii

LIST OF FIGURES . . . . . . . . . . x

ABSTRACT . . . . . . . . . . .. xv

CHAPTER 1 INTRODUCTION . . . . . . . 1
History . . . . . . . . . . 3
Structural Configuration of Doxorubicin . . 4
Pharmacology and Mechanism of Action . . . . 6
Interaction with DNA . . . . . . 6
Interaction with Cell Membrane . . . . 7
Pharmacokinetics . . . . . . . . .. 10
Metabolism and Excretion . . . . .. .11
Physical Properties of Doxorubicin . . . .. .12
Partition Coefficient . . . . . .. 13
Metal Ion Complexation . . . . . .. .15
Molecular Spectroscopy . . . . . . .. .17
UV/Vis spectra . . . . . . . .. 17
Fluorescence spectra ..... .............. .18
Dissociation Equilibria of Doxorubicin . . .. .19
Self Association . . . . . . . . .. 22
Adsorption to Glass . . . . . . . .. 26
Analytical Methods ......... ............... . 28
Stability and Kinetics of Degradation . . .. .32
Stability Studies on Doxorubicin . . .. .33
Stability of (I) in Pharmaceutical Formulations . 43
Stability in Biological Media . . . . .. .46
Farmitalia and U.S. Patent # 4946831 . . . .. .48

CHAPTER 2 MATERIALS AND METHODS . . . . .. 52
Materials . . . . . . . . . .. .52
Test Compound . . . . . . . .. 52
Reagents/Solvents . . . . . . .. 53
Instrumentation .... . . . . . .. . 53
Melting Point determination . . . .. 53
Differential Scanning Calorimetry . . .. 53
pH measurements . . . . . . .. 54
Ultraviolet and Visible Spectroscopy . . .. 54
Self association . . . . . . .. 55







HPLC . . . . . . . . . . 55
Hydrolysis . . . . . . . .. 57
Fluorescence Spectroscopy . . . . . .. 57
HPLC . . . . . . . . . .. 57
Self association . . . . . . .. 59
Chromatography . ........ ..................... .61
Mobile Phase(s) . . . . . . .. 61
Solvent Delivery . . . . . . .. 62
Column . . . . . . . . . .. 62
Injector . . . . . . . . .. 63
Detector . . . . . . . . .. 63
Recorder/Integrator . . . . . . .. 63
Temperature controlled studies . . . . .. .64
Water bath(s) . . . . . . . .. 64
Methods . . . . . . . . . . .. 64
Silylating Glassware . . . . . .. 64
Self Association . . . . . . .. 65
High Performance Liquid Chromatography . . .. .69
Assay validation . . . . . . .. 70
Glass Binding . . . . . . . . .. 70
Effect of Temperature . . . . . .. 70
Effect of pH . . . . . . . .. 71
Effect of Concentration . . . . . .. 71
Experimental . . . . . . . .. 71
Sampling . . . . . . . . .. 72
Data Treatment . . . . . . . .. 73
Chemical Kinetics . . . . . . . .. 73
Effect of acids . . . . . . . .. 73
Effect of concentration . . . . . .. 74
Effect of Temperature . . . . . .. 74
Effect of pH . . . . . . . .. 75
Effect of buffers . . . . . . .. 75
Experimental . . . . . . . .. 76
Data Treatment . . . . . . . .. 76
Identification of degradation products . . .. .77
Analysis of the solution and precipitate . 78
Effect of oxygen on the hydrolysis . . .. .79
Detection of product X-COOH by potentiometry 80
Detection of product X-COOH by HPLC . . .. .81
Detection of product X-COOH by
derivatization . . . . . . . .. 82
Precipitate analysis by LC-MS . . . .. .82

CHAPTER 3 RESULTS AND DISCUSSIONS . . . . .. 84
Purity Determination . . . . . . .. 84
Melting Point . . . . . . . .. 84
Differential Scanning Calorimetry . . .. 84
Chromatography . . . . . . .. 86
HPLC Assay . . . . . . . . . .. 89
Internal Standard . . . . . . .. 89
Capacity factor (k') . . . . . .. 89
Calibration curves and statistics . . .. 92







Validation of the Assay .


Self Association . . . . . .
Fluorescence Measurements . . .
Normalization Experiments . . .
Selected Excitation Study . . .
Estimation of Association Constants
Estimation of E[MC-DOX] . . . .
Estimation of OpN-DOXi / MC-DOXI . .


Discussion . . . . . . . . .
Structure of the Dimer . . . . . .
Glass Binding . . . . . . . . .
Data Analysis . . . . . . . .
Types of Binding Sites . . . . . .
Capacity of the Binding Sites . . . .
Effect of Temperature . . . . . .
Effect of Concentration . . . . . .
Amounts Bound to Glass . . . . . .
Effect of pH . . . . . . . .
Effect of Silylation . . . . . .
Kinetics of Hydrolysis . . . . . . .
Rate Constants . . . . . . .
Statistics of the Kinetic Data . . . .
Effect of Concentration . . . . . .
Effect of Mineral Acids . . . . . .
Effect of pH . . . . . . . .
pH of Maximum Stability . . . . . .
Effect of Buffers and General Acid-Base
Catalysis . . . . . . . . .
Determination of the Catalytic Species of the
Buffer . . . . . . . . . .
Effect of Temperature . . . . . .
Degradation of (I) in the pH range 0.8 to
3 .0 . . . . . . . . . . .
Degradation of (I) in the pH range 4.0 to 8.0
~~~~~~ . . . .
Identification of the Degradation Products
Mechanism of degradation of (I) . . . .
Effect of oxygen and nitrogen on the
degradation of (I) . . . . . .
Identification of Glycollic Acid and
Glycolaldehyde . . . . . . . .

CHAPTER 4 CONCLUSIONS . . . . . . . .


APPENDIX . . . .

REFERENCE LIST . ..

BIOGRAPHICAL SKETCH .


vii


156

161
164

173

175
178
178

180

188

189


. . 196


198

220


. . . . 92


93
93
S99
101
104
108
109
112
117
119
119
120
128
128
134
134
137
140
141
141
143
143
144
148
154














LIST OF TABLES


Table 1.1



Table 1.2


Table 3.1


Table 3.2


Table 3.3


Table 3.4


Table 3.5



Table 3.6




Table 3.7


Arrhenius parameters for the
hydrolysis of (I) . . .


Farmitalia patent data fitted
to the Arrhenius function . .


Capacity factors (k') and
Retention times (tr) of various
analytes studied by HPLC . .


Statistics of a typical
calibration curve for all
analytes in the concentration
range 25-700 ng/ml . . . .


Analysis of Variance of the
calibration curves for the
various analytes (I) to (V) .


Molar absorptivity (Mc-DOx) of
the monocation form of (I) at
ambient temperature . . .


Association constants estimated
at different ionic strengths .


Hybrid constants obtained from
the fitting of the glass
binding data . . . .


Derived microscopic rate
constants obtained from the
hybrid constants . . . .


. . . 90


. 94


. . . 95


109


. 117


. . 124




. . 126


viii







Table 3.8




Table 3.9


Table 3.10




Table 3.11


Table 3.12




Table 3.13






Table 3.14


Table 3.15


Table 3.16





Table 3.17


Equilibrium constants derived
from the microscopic rate
constants . . . . .


Binding capacities of the type
1 (A/a) and type 2 (B/9) sites
and the predicted amounts bound
to the type 1 site per unit
area . . . . . . .


Predicted amounts bound
(ig/cm2) to the type 1 and type
2 binding sites . . . .


Validation of y, . . .


Effect of mineral acids on the
specific acid catalysis of (I)
at 50C . . . . . .


Second order rate constants
obtained from log kobs vs pH
profiles corrected for buffer
effects at various
temperatures . . . . .


pH of maximum stability as a
function of temperature . .


Values of kobs and kcat for the
general acid-base catalysis of
(I ) . . . . . . .


Second order catalytic rate
constants for the hydrolysis of
(I) and their corresponding
Arrhenius parameters . . .


Activation Energy Ea + Standard
Error and Frequency Factor (Ln
A) for the hydrolysis of (I) at
different pH . . . . .


. . 129






. . 130


138


147


. . 150






. . 154



. . 156


160


. . 165





. . 168














LIST OF FIGURES


Figure 1.1


Figure 1.2


Figure 2.1


Chemical structures of
Anthracyclines (I) to (VI) and
Rhodomycin . . . . . .


Twisted boat conformation of
the cyclohexyl ring of (I) . .


UV spectra for analytes (I) -
(VI) dissolved in mobile phase


C . . . .


S . 2


. 5


. . . . 56


Figure 2.2




Figure 2.3



Figure 2.4



Figure 3.la




Figure 3.lb




Figure 3.2


Fluorescence spectra for
analytes (I) (VI) in mobile
phase C . . . . . .


Front surface and right angled
illumination setup . . . .


Typical fluorescence spectra of
(I) in aqueous solution . .


A typical DSC endotherm of the
melting of (I) for purity
evaluation . . . . . .


Plot of Ts vs 1/F for (I),
based on equation 3.1 in text
. . . . . . . . .


A typical chromatogram for
analytes (I) (VI) . . .


S. 58



S. 60


. 85


. 87


. 88







Figure 3.3




Figure 3.4



Figure 3.5



Figure 3.6


Figure 3.7




Figure 3.8





Figure 3.9




Figure 3.10




Figure 3.11




Figure 3.12



Figure 3.13


Calibration curves (25-700
ng/ml) for analytes (I) (VI)
prepared in mobile phase C . .


Relative Fluorescence Intensity
vs Concentration at 590 nm . .


Representative example of a
normalization spectra . . .


Diffrential spectra of (I) . .


Emission spectra of (I),
selectively excited at 580 nm



Fluorescence spectra used to
determine E[N-DOX] /([MC-DOX] at
boundary conditions of pH as
per equation 3.15 in text . .


Non-Linear fitted plots for the
self-association of (I) at pH
5 .72 . . . . . . .


Non-Linear fitted plots for the
self-association of (I) at pH
6 .72 . . . . . . .


Non-Linear fitted plots for the
self-association of (I) at pH
7 .69 . . . . . . .


Effect of ionic strength on the
association constant . . .


Representative plot of percent
Remaining vs time for glass
binding of (I) at pH 3.18,
temp: 5C and Conc: 1 p.g/ml .


. . . 91



. . . 98



. . 100


. . 102




. . 103





. . 111




. . 113




. . 114




. . 115



. . 116





. . 121







Figure 3.14




Figure 3.15




Figure 3.16


Figure 3.17


Figure 3.18


Figure 3.19



Figure 3.20


Figure 3.21


Figure 3.22




Figure 3.23


Figure 3.24




Figure 3.25


Representative plot of percent
Remaining vs time for glass
binding of (I) at pH 5.01,
temp: 5C and Conc: 10 gg/ml .

Representative plot of percent
Remaining vs time for glass
binding of (I) at pH 7.00,
temp: 5C and Conc: 100 ig/ml

Temperature dependence of glass
binding; pH 3.18; Conc: 10
jig/ml . . . . . . .

Effect of concentration on
glass binding; pH 5.01 . . .

Amount bound to glass at pH
5 .01 . . . . . . .

Typical first order plot for
the degradation of (I) in HCl
at 60C . . . . . .

Effect of Mineral Acids on the
hydrolysis of (I) at 50C . .

Log kobs vs -Log a[H+] plot
showing no influence of the
acid type on the activity of
[H'] ions . . . . . .


Log kobs vs pH plot for the
hydrolysis of (I) at zero
buffer concentration . . .


General Acid-Base catalysis-
Effect of Formate buffer at
60C . . . . . . .


General Acid-Base catalysis-
Effect of Acetate buffer at
60C . . . . . . .


General Acid-Base catalysis-
Effect of Phosphate buffer at
60C . . . . . . .


. 122




123



133


. . 135


. . 136



. . 142


. . 145




. . 149


. 152


S157


. . 158




. . 159


xii







Figure 3.26


Figure 3.27


Figure 3.28



Figure 3.29



Figure 3.30


Figure 3.31




Figure 3.32


Figure 3.33




Figure 3.34


General Base catalysis-Effect
of CH3COO on the hydrolysis of
(I ) . . . . . . .


General Base catalysis-Effect
of HP042 on the hydrolysis of
(I ) . . . . . .


Arrhenius plot for the
hydrolysis of (I) in HC1 . .


Arrhenius plot for the
hydrolysis of (I) at pH 3 & 4


Arrhenius plot for the
hydrolysis of (I) at pH 4.5, 5
& 5 .5 . . . . . .


Arrhenius plot for the
hydrolysis of (I) at pH 6, 7 &
8 . . . . . .


Degradation of (I) and the
formation of (IV) in the pH
range 0.8 to 3 . . . .


Sigma minus plot showing first
order formation of (IV) in the
pH range 0.8 to 3 . . .


Representative chromatogram of
a degraded solution of (I) at
pH values from 4 to 8 . .


Figure 3.35


Degradation of (I)
presence of oxy
nitrogen . . .


in the
gen and


. . 183


Figure 3.36


Zero order formation of [H]
when (I) is degraded in the
presence of oxygen . . .


xiii


. . 163




. . 166


169


170


171


. . . 172




. . . 174


176


179


185














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

PHYSICOCHEMICAL CHARACTERIZATION AND
STABILITY OF DOXORUBICIN IN
AQUEOUS SOLUTIONS


By

Anup Zutshi

April 1994

Chairperson: Hans Schreier Ph.D.
Major Department: Pharmaceutics

Doxorubicin (I), an anthracycline antitumor antibiotic,

is unstable in aqueous solution, resulting in loss of

activity. A sensitive, reproducible HPLC assay to quantify (I)

and its major metabolites/degradation products in the low

ng/ml range was developed.

Literature reports on the value of the equilibrium

constant for the self-association of (I) are conflicting.

Using a 'front surface illumination' fluorescence technique,

a mathematical model was developed which uniquely estimates

the influence of pH and ionic strength on the association

process and recognizes the ability of the various ionic

species of (I) to self-associate.

Nonspecific glass binding, contributes to losses of (I)

from aqueous solutions. This phenomenon, studied as a function

xiv







of concentration, temperature and pH, found that lowering the

temperature or concentration or increasing the pH increased

the amount of (I) bound to glass. The percent loss of (I) from

solution was fitted to a biexponential equation, suggesting

the possibility of two independent binding sites, each of

which could be saturated at higher concentrations.

The effect of pH, temperature and buffers on the

hydrolysis of (I) was studied. Specific acid catalysis of (I)

was observed from pH 0.8 to 3.0. The pH of maximum stability

was between 4.5 and 5.0. At pH>3 the hydrolysis was influenced

by general base catalysis. The contribution of each of these

buffer species was calculated. Hydrolysis was carried out

under accelerated conditions at 70, 65, 60 and 55C (50C for

the mineral acids). Arrhenius parameters showed the activation

energy to depend on pH, indicating different mechanisms of

degradation. At pH<3, the aglycone (V) was formed as a 1:1

degradation product. At pH>3, the chromatograms showed

unidentified peaks which appeared irrespective of the buffer

species, indicating that degradation might be pH specific.

Aqueous solutions of (I), degraded in the presence of oxygen,

showed a lowering of pH, indicating the formation of an acidic

product. A structure for this product is proposed.








CHAPTER 1
INTRODUCTION



Cancer (1) is a generic term for a variety of malignant

neoplasms, arising in all human and animal tissues composed

of dividing cells, which have adverse effects on the host

tissue by invasive growth and metastases. The most rapidly

developing therapy in cancer treatment uses chemical agents

with activity against human neoplastic disease.

Doxorubicin (I) (Adriamycine, Adriblastina) of the

anthracycline antitumor antibiotics has been shown to be

active against a wide spectrum of human malignancies

including soft tissue, osteogenic and other sarcomas like

Hodgkin's disease; non-Hodgkin's lymphomas; breast,

genitourinary, thyroid, lung, and stomach cancers;

neuroblastoma and acute leukeamias (acute granulocytic and

acute lymphocytic) (2).

The anthracyclines used as part of the studies

conducted in this thesis are (Figure 1.1): doxorubicin (I);

doxorubicinol (II), doxorubicinolone (III), doxorubicinone

(IV), 7-deoxy doxorubicinone (V) and daunorubicin (VI).










0 OH


0 OH

A: AGLYCONE


Figure 1.1:


Chemical structures of Anthracyclines (I) to (VI) (see text) and
Rhodomycin.


S: SUGAR


COMPOUND R1 R2 R3 R4 R'

RHODOMYCIN B OH S H CH3 (NCH3)2
DOXORUBICIN (I) OCH3 S -0 CH2OH NH2

DOXORUBICINOL (10) OCH3 S OH CH2OH NH2

DOXORUBICINOLONE (111) OCH3 OH OH CH2OH

DOXORUBICINONE (IV) OCH3 OH = 0 CH2OH -

7-DEOXY DOXORUBICINONE (V) OCH3 H = O CH2OH -
DAUNORUBICIN (VI) OCH3 S O CH3 NH2
(INTERNAL STANDARD) 3









History


In the late 1950's a pigmented compound, rhodomycin B

(Figure 1.1), originating from a strain of the Streptomyces

species found in soil samples in India, was characterized

and found to exhibit antitumor properties (3,4). In the

early sixties (VI), a precursor to (I), was isolated from a

strain of Streptomyces peucetius (5). Daunomycin was shown

to have rapid and excellent activity, with no cross

resistance against acute leukeamias (6). Major side effects

included severe aplasia, immunodepression and cardiac

complications, the latter requiring a limitation of the

total dose to 25 mg/kg with consequent impossibility of use

for maintenance treatment. In light of the pharmacological

results with (VI), subsequent research was developed along

a) elucidation of the structure and stereochemistry of the

antibiotic, and b) search for new biosynthetic analogs in

cultures derived from the (VI) producing micro organisms

Streptomyces peucetius. The latter approach, which resulted

in the isolation and characterization of (I), was based on

the premise that variations in the anthracycline structure

could induce a remarkable improvement of pharmacological

properties.

Doxorubicin was isolated from the cultures of one of

the strain varieties derived from a mutant of the original

Streptomyces peucetius, namely Streptomyces peucetius var.

caesius (7). Doxorubicin had better activity against tumors







4

than (VI) (8), but as with all anthracycline antibiotics its

use is restricted due to a dose limiting cardiac toxicity,

resulting in an acute cardiac myopathy (8). In clinical

practice the cumulative dose of (I) should not exceed 450-

550 mg/m2. Attempts to circumvent the problem of

cardiotoxicity have dealt with alternative routes of

administration (9-11), alternative dosage schedules (12-18),

pre- and co-administration of free radical scavengers (19-

26) and encapsulation in liposomes (27-35).

Structural Configuration of Doxorubicin


Most of the work in this field was first conducted on

(VI) and carminomycin using X-ray crystallography. Since

(VI) and (I) are similar in structure, these studies also

apply to (I). X-ray crystallographic studies (36,37,38) on

(VI) and carminomycin hydrochloride (37,38) showed that the

anthracycline part of the molecule was flat. This was also

indicated by the resonance stabilization that led to an

extended conjugation and a colored compound with a strong

fluorescence. The NMR spectra of (VI) confirmed that the

cyclohexyl ring is in the half chair conformation (or

twisted boat conformation) (Figure 1.2) and the sugar moiety

is nearly perpendicular to the plane of the chromophore.

Studies (39) on the solution conformation of (I) indicated

the same structure described by X-ray crystallography. The

distance between the 0-5 and 0-6 atoms (2.45 A) was found to

be much lower than between the 0-11 and 0-12 (2.67 A)







5

indicating a different charge distribution within the

quinonic chromophore. The relevance of this structural

property for antitumor activity has not been established.

The half-chair conformation of the cyclohexyl ring

stabilizes the molecule due to intramolecular hydrogen

bonding (Figure 1.2).

The 0-9 and 0-7 oxygen atoms can hydrogen bond via the

-OH at 0-9. Alternatively there could be a binding

interaction between the OH at C-9 and the oxygen of the

sugar moiety.

The positively charged NH2 group of the molecule in

(VI) has been shown to be hydrogen bonded with the C-4' OH

group.

Another important finding indicated the axial H-8 and

H-10 protons had a long range coupling in all the

derivatives of (VI) studied (40) which suggested the

preferred half chair conformation of the cyclohexyl ring.








//
H >iH10a







90


R


Figure 1.2: Twisted boat conformation of the cyclohexyl
ring of (I)









Pharmacology and Mechanism of Action


Interaction with DNA.

Most of the (I) in the cell accumulates in the nucleus

(41,42) up to a maximum of 1 drug molecule per 9 base pairs

(43). Doxorubicin binds to DNA by intercalation between

successive base pairs of the helix. The affinity constant,

determined by Scatchard analysis, is about 2 X 106 M-i and

the number of binding sites is 0.2 per DNA phosphate.

The biological significance of this interaction is not

clear. Both DNA and RNA syntheses are inhibited, with the

effect on initiation being greater than on elongation (44)

and the effect on repair being less than on replication

(45,46). Additionally, DNAase (47,48), reverse transcriptase

(49,50) and DNA polymerase II (51) are also inhibited. Drug

binds to chromatin with the same affinity as DNA but with

lower number of binding sites (52,53). The drug binds to the

internucleosomal regions, leading to compaction of the

chromatin (54). This is probably the lethal event, rather

than the inhibition of nucleic acid syntheses. Cytotoxicity

does correlate with nuclear drug content (55).

Additionally, covalent binding to DNA can occur (56,57,58)

through reductive activation, as well as photoactivation

(59,60) resulting in the degradation of the DNA (61). This

damage was thought to occur due to the generation of

reactive oxygen species like the OH radical (56,57).

However, breakage in the DNA also occurred in the absence of









oxygen, implicating an enzymic mechanism (62).

Interaction with Cell Membrane.

Doxorubicin undergoes a physical interaction with

membranes (63,64,65). An initial ionic attraction is

followed by the insertion of the hydrophobic region of (I)

into the phospholipid region of the membrane (64,66). The

affinity of the binding is equivalent to that for DNA (64).

The interaction is noncovalent unless the lipid is

peroxidized, in which case Schiff's base formation can occur

(65). The effect of the insertion into the membrane is an

increase in membrane fluidity (67,68,69) which may be the

cause of experimentally noted effects such as stimulation of

membrane NADH oxidase (70), inhibition of membrane ascorbic

oxidase (71), histamine release from mast cells (72),

increased glycosylation of the cell surface (73) and

deacylation of phospholipid (74). Whilst (I) will insert

into membranes of any phospholipid composition, there is a

specific binding interaction with cardiolipin, a

phospholipid generally only found in the mitochondria

(65,67,75). The specialized cardiolipin organization within

the membrane is destroyed, the (I)-cardiolipin complex

segregates, (76) there is an increase in the rigidity of the

membrane, (77) and the complex can act as an electron

transport system (77,78). These effects, along with lipid

peroxidation, could be the cause of damage to heart cells

and muscle.









Another membrane effect associated with (I) is the

effect on the Ca"2 uptake and transport. This effect could

be related to the lipid peroxidation, (I)-cardiolipin or

(I)-membrane interactions (79,80). Calcium transport is

affected (81,82), leading to a reduction of fast exchanging

calcium in heart cell mitochondria and sarcolemmal vesicles

(83,84). Calcium linked processes such as the Na'/K pump

are impaired (82,85,86), giving a prolongation of the action

potential (87,88). Other consequences include inhibition of

Ca2 dependent protein kinases (89), inhibition of actin

polymerization (90) and leukotriene formation via

phospholipase A2 activation (91). Support for the hypothesis

that Ca2 control impairment is related to the cardiotoxic

effects is given by the reversal of the cardiotoxicity by

the chelating agent ICRF-159 and its levo isomer ICRF-187

(92,93) and by verapamil, propranolol and hydralazine (94).

It is apparent that (I) has multiple effects and the

effect on cardiac mitochondria is probably the major cause

of the cardiotoxicity. This may be mediated by lipid

peroxidation (due to free radical formation) and/or physical

interaction with cardiolipin coupled with a direct or

indirect action on the control of Ca2 flux. A recent

hypothesis seems to implicate (II), a metabolite of (I), in

the cardiac toxicity. This hypothesis is supported by the

fact that cardiotoxicity is time delayed and is unrelated to

the plasma or cardiac levels of (I). The terminal half life









of (I) in plasma is 24 hours, and thus (I) is virtually

eliminated from plasma and the heart in 4 to 5 days (95,96).

The pharmacokinetics are similar in multiple as well as

single dose studies in rats (97). In contrast, the

cardiotoxicity takes weeks to develop and seems to be

correlated with the appearance of (II), which appears in

detectable levels after 24 hours (95,97,98). Doxorubicinol

accumulates in cardiac tissue and appears much more toxic

and more potent in isolated cardiac preparations than (I)

(98,99). However, direct comparisons of cardiotoxicity

between (I) and (II) are confounded by other factors.

Cardiac dysfunctions correlate to time dependency with (I)

whereas the cardiotoxicity of (II) is time independent.

Thus, critical evaluation of the metabolite hypothesis and

the use of in vitro vs in vivo models to assess the

cardiotoxicity should help to resolve and elucidate the

mechanism of cardiotoxicity (100).

Some of the other toxic manifestations (common to most

anticancer agents) are immunosuppression, myelosuppression,

reversible alopecia, diarrhoea, nausea and vomiting (101). A

serious side effect is severe dermatological toxicity if (I)

is extravasated during infusion. Severe pain is noted

immediately upon extravasation, followed in a few hours by

swelling and reddening. Skin lesions can become severe over

a period of several months resulting in necrosis of the skin

and the underlying tissues (102-104) down to the bone.









Spontaneous healing of these ulcers is rare, and surgical

excision of the involved tissues has become the recommended

treatment (102,103).

Pharmacokinetics


It is generally accepted that plasma pharmacokinetics

of (I) following intravenous administration is best

described by a sum of three exponentials in both man and

animals (105-114). Earlier studies (115-119) showed fits to

a sum of two exponentials, probably due to a lack of

analytical sensitivity. The terminal half life is 14 to 30

hours in human patients suffering from cancers other than

hepatic tumors (110,114,120), 21-30 hours in the dog (113),

40 hours in the guinea pig (121), and 24-61 hours in the rat

(116,117,122,123). Plasma protein binding is about 59% in

rabbits, 90% in humans, 66% in rats and 48% in guinea pigs

(121,122). Hepatic enzyme induction and enzyme inhibition in

rats (122,125) pretreated with phenobarbital (an enzyme

inducer) reduced the plasma half life of (I) from 26 to 16

hours (125) and increased the excretion of (I) by 27% (122).

There was also an increase in the plasma levels of the

metabolite (II). Pretreatment with carbon tetrachloride (an

enzyme inhibitor) resulted in significantly higher levels of

(I) in plasma by reducing the total elimination by about 30%

(122). A similar observation made in humans being treated

for various cancers with impaired hepatic functions resulted

in a corresponding reduction in dose (126,127). In order to







11

expose tumors to prolonged levels of (I), and to reduce the

accompanying cardiotoxcity, a general consensus seems to be

to deliver the drug as a slow i.v. infusion in periods

ranging from 10 to 96 hours (110,120,128-131).

Metabolism and Excretion.

The principal metabolite in all species appears to be

(II) and the 7-deoxyaglycones of (I) and (II) (132). The

conversion to the alcohol is catalyzed by widely distributed

cytoplasmic aldo-keto reductases (133,134) and requires the

use of NADPH as a cofactor. Doxorubicin was reduced about 20

times more slowly than its analog (VI) (135), suggesting a

possible reason for its higher efficacy. The reductive

splitting of the amino sugar moiety of (I) and its principal

metabolite (II) to their corresponding deoxyaglycones was

demonstrated to occur in the microsomal fraction of the rat

liver (136). This catalysis was inhibited by oxygen,

required NADPH as a cofactor, was induced by

phenobarbital,and was not affected by enzyme inhibitors such

as carbon monoxide, j-diethylaminoethyldiphenylpropyl

acetate or Mg2, Mn+2, Fe2 and Ni+2, although it was inhibited

by Cu2 and Zn2. In the dog, i.v. administered (II) had a

half life of 3.7 hours but demonstrated a half-life similar

to the parent compound when analyzed after administration of

(I). This indicated the rate-determining nature of the

metabolic step (113). Over a 24-hour period only 5.2% of the

unchanged drug and 0.13% of the derived (II) was excreted in









the urine of dogs (113). In rats, between 6-8% of the dose

was recovered in the urine. About 35% was excreted in the

bile after 10 hours, of which 70% was as unchanged drug and

30% was as (II), with small amounts of the corresponding

deoxyaglycones (137). A mass balance study of radiolabelled

(I) in rats showed that after 96 hours 7.5% of the

radioactivity was in the urine, 65% in the feces and 6.75%

in expired CO2, with the remaining 21% in other tissues,

where a higher concentration in cardiac muscle than skeletal

muscle, liver, lymphatic and glandular tissues was found

(138). In man, the metabolism and excretion was

qualitatively similar to that of the rat. However, only 60%

of the dose could be accounted for, the rest probably being

present as nonfluorescent metabolites. Of this, only 10%

appeared in the urine over a period of 5 days and about 50%

in the bile (118,139). The principal products excreted in

the bile and the urine were about 40% of (I) and 29% of the

major metabolite (II) with 9% of the combined aglycones, 10%

of the 4-demethoxysulfate and about 12% of the 4-

demethoxyglucouronide (140).


Physical Properties of Doxorubicin(141)

Chemical Name.

The chemical name for (I) is (8S,10S)-10-(3 amino-

2,3,6-trideoxy-alpha-L-lyxo-hexopyranosyl)oxy-7,8,9,10-

tetrahydro-6,8,11-trihydroxy-8-hydroxyacetyl-l-methoxy-5,12

napthacenedione (CAS # 23214-92-8).









Appearance and Color.

The hydrochloride salt is a free flowing crystalline

powder, and the freeze dried formulation containing lactose

is a red cake.

Empirical Formula.

The empirical formula for (I) is C27 H29 01 N.HCl

Molecular Weight

The molecular weight of (I) is 579.98 (of HCl salt).

Melting Point.

The melting point of (I) has been determined as 205C

(with decomposition).

Solubility.

The hydrochloride salt is readily soluble in water,

physiological saline and methanol. It is slightly soluble or

practically insoluble in organic non polar solvents.

Optical Rotation.

Strongly dextrorotatory [a], = +255 at 589 nm in 0.1%

methanol.

Partition Coefficient.

The determination of the partition coefficient of any

drug between an aqueous and organic phase(s) is an important

parameter. In addition to revealing important

physicochemical information like ionization behavior,

lipophilicity and solubility, one can also extrapolate this

information to indicate binding to tissues, cells and other

macromolecules, uptake into cells and transference across







14

biological membranes. In addition, partition behavior can be

utilized to establish sample handling strategies such as

extraction of the drug from interfering biological or

chemical matrices prior to detection or quantification by

various analytical techniques. The apparent partition

coefficient of (I) between 1-octanol and Tris buffer at pH

7.0 with constant ionic strength (1:0.1) at room temperature

after shaking for 15 hours was determined to be 0.52 (141).

Eksborg (142) studied the partitioning behavior of a series

of anthracyclines as a function of pH in order to study the

self association of (I). Using a chloroform:l-pentanol (9:1)

system as the organic phase he found that the distribution

coefficient was strongly dependent on the pH of the aqueous

phase with optimal extraction occurring between pH 8.0 to

8.6. The apparent dissociation constant (pKj) of (I) was

determined to be 7.20, resulting in the isoelectric point

(isoelectric point is that pH where 50% or more of the

species is in the zwitterion form) of the molecule lying

between 8.0 and 8.6 (pK2 = 9.3). It can be argued that the

value of the pK2 has been underestimated from the

spectroscopic data (see discussion on dissociation

equilibria) and that the isoelectric point and the pH range

of maximum extraction lies between 9.0 and 9.2. The

extraction efficiency decreased as concentration increased,

presumably due to the formation of dimers and tetramers of

doxorubicin in the aqueous phase. Nakazawa et al.(143)









employed a countercurrent extraction method to study the

effect of various inorganic salts on the partition

coefficients of (I) and (VI) using three different organic

phases. The highest extraction was obtained between a 1-

butanol:1.0 M Na2HPO4 (1:1) system. However, in this study

the effect of pH on the extraction efficiency was not

reported. In studies dealing with the analytical development

of (I) from biological matrices, extraction with

chloroform:isopropanol (4:1 v/v) (144) and

chloroform:methanol (9:1 v/v) (145) have been considered

satisfactory (>90% extraction efficiency). The pH of the

aqueous phase in most cases is adjusted between 8.0 and 9.5

prior to extraction.

Metal Ion Complexation.

Aluminum, copper, magnesium, iron, calcium, gallium and

a host of other metal ions bind to (I) in solution (146).

The metal binding seems to affect the activity of (I) with

respect to its interaction with DNA. Thus, physiological

levels of Cu2 and Fe"3 ions interact with (I) resulting in

complexes that intercalate to a larger degree with DNA (147,

148-153). The Fe"3 complex intercalates with the DNA and the

metal ion is released (154). Both Cu2 and Fe"3 form ternary

complexes with (I), and these complexes are capable of

producing the superoxide free radical, which in turn is

capable of inducing DNA damage (147,151,155,156). Studies

dealing with the interaction of these ternary complexes with







16

phospholipids, erythrocyte ghost membranes and ADP show that

all of these systems initiate the destruction of these

biomembranes (157-159). Other metal ions like Ca2, Cd2,

Mg2, Pb+2, and Zn+2 do not interact with the (I)-DNA complex.

Complexation studies using UV-Vis, circular dichroism, NMR,

Raman, Mossbauer and IR spectroscopy indicate the

involvement of the deprotonated hydroxyanthraquinone moiety,

which binds with the metal ion forming a six membered

chelate (154,147,160,161,162,163-165). Initially the C-lI

hydroxy and C-12 carbonyl functions are the preferred sites

for chelation. The second metal ion is introduced at the C-5

ketone and C-6 hydroxy site. Complexation occurs in the pH

range from 2 to 8 and the optimum pH largely depends on the

metal ion involved (161). Binding does not take place at pH

values below 2 and above 8 (where metal hydroxide formation

predominates). Metal binding is accompanied by a change in

spectral characteristic and a drop in pH (147,160-165)

indicating the displacement of the proton from the C-ll or

C-6 hydroxy groups on the chromophore.

A polymeric form of Fe1-(I) complex, called

Quelamycin, is undergoing clinical trials. The complex was

first described by Gosalvez et al.(166), and is believed to

be less cardiotoxic than (I). It exists only in concentrated

solutions and hydrolyses to Fe-(I)-OH upon dilution (167).

Studies have shown that Quelamycin is in fact a mixture of

Fe(I)3 and polymeric ferric hydroxide (154).









Molecular Spectroscopy



UV/Vis spectra.

Absorption maxima occur at 233, 253, 290, 477, 495 and

530 nm. The UV spectra shows the characteristic peaks of

extended conjugation of an aromatic nucleus (141). The broad

hump starting at 420 nm is indicative of the highly

conjugated anthraquinone moiety. This gives the compound its

red color. Even though the molar absorptivity between 420

and 540 nm is low, the lack of interference from other

components makes this region of the spectrum an attractive

possibility for chromatographic analysis. On adding alkali

(pH>9), the UV-Vis spectrum shifts towards longer wavelength

due to the characteristic indicator-like properties of

quinones. The color change associated with this spectral

shift is from orange-red to violet-blue. The absorption

spectrum of (I) is also dependent on the presence of metal

ions (147,149,157,158), solvent (168,169), drug

concentration (168,161-174) and ionic strength (146,169).

The red shift associated with the deprotonation of the

phenolic proton from the chromophore is analogous to the one

observed when (I) interacts with metal ions. The spectra of

the monomer and the dimer of (I) can be established by

differential measurements (172). Increase in the ionic

strength has a similar effect as the concentration

indicating the possibility that increased ionic strength









promotes aggregation of (I) in solution (169). UV-Vis

spectroscopy has been extensively used to study the

interaction of (I) with nucleic acids, as this interaction

is accompanied by changes in the spectra (146,150,175-178).

Fluorescence Spectra.

Excitation at the longest wavelength absorption maximum

at 485 un gives an emission band between 520-620 nm. Since

the absorption and emission bands overlap, certain inner-

cell effects, such as self-absorption and concentration

quenching, can affect the spectral behavior. The

fluorescence emission has a high quantum yield, and is used

as a sensitive method for the analysis of (I) from solutions

or biological matrices (179-186). However, to ensure

specificity, a technique like chromatography is necessary to

separate the principal metabolites and degradation products

of (I) which have identical spectral characteristics.



In some cases of fluorescence spectroscopy the apparent

fluorescence intensity and spectral distribution can be

dependent on the optical density of the sample and the

precise geometry of sample illumination by the excitation

light (187). The most commonly used geometry is the right

angle observation of the center of a centrally illuminated

cuvette. Frequently, however, off-center illumination

techniques are used, especially in samples having a high

optical density or turbidity, in order to reduce the path







19

length of illumination of the sample and thereby reduce the

possibility of attenuation of the emitted signal due to

self-absorption and inner-cell effects (187).

In one case of the off-center illumination technique,

front face illumination is performed using triangular

cuvettes or square cuvettes oriented at 45 to the incident

beam of light. A drawback of this system is that a large

portion of the incident light is reflected off the front

surface of the cuvette into the emission monochromator

increasing the chances of interference from stray light. A

better position (187) is to orient the cuvette at 30 to the

incident beam, thereby reducing interference from stray

light.

Dissociation Equilibria of Doxorubicin.

Knowledge of the protolytic behavior of (I) and related

compounds in aqueous solution is of obvious importance for

the evaluation of the biochemical interactions with

biologically significant macromolecules and receptors, and

for the interpretation of their degradation kinetics in

biological fluids and aqueous solutions.

Sturgeon and Schulman (188) investigated the protolytic

equilibria of (I) using UV-Vis and fluorescence spectroscopy

(Scheme 1.1). As shown, they concluded that (I) was present

in aqueous solutions at pH<7.0 as the monocation MC-DOX.

However, at higher pH values MC-DOX could lose a proton to

form a neutral species N-DOX or a zwitterion Z-DOX. Further









removal of a proton from N-DOX or Z-DOX will give rise to

the anion MA-DOX. The negative charge on Z-DOX or MA-DOX

could be placed on either the C-6 or C-lI structurally

equivalent oxygen atoms. The doubly charged cation can be

observed when the drug is placed in concentrated sulfuric

acid. A doubly charged anion can be seen at pH>13. The

values for the microscopic constants were pK1=8.22,

pK2=10.10, pK3=9.01l, pK4=9.36. The macroscopic constants were

pK1=8.15 and pK2=10.16. In the above study self-association

was avoided by maintaining concentrations at less than 10-6

M.

In another study using isoelectric focusing techniques,

Righetti et al.(189) reported that the isoelectric point of

the compound was 8.76 Spectrophotometric titration of the

phenolic OH group gave a pKa of 9.6. On this basis, the pKa

of the NH3' group was estimated at 7.92. This estimate did

not take into account the presence of the zwitterion and the

apparent equilibria that could exist between the zwitterion

and neutral species. Direct titration of (I) with 0.05 M

NaOH (141) gave a PKa value of 8.22. The pKi and pK2

estimated by Eksborg (142) using solvent extraction and

spectroscopic techniques described before were 7.20 and

9.30, respectively.

The ionization of (I) has been shown to be dependent on

the degree of self-association in aqueous solutions

(132,190), configurational changes in the daunosamine moeity


















pK1 = 8.1


pK3=9.01


MONOCATION [MC-DOX]


SCH2OH
'OH


K1

pK =8.22


5




K=[Z = 0.16
Kt =[Zl/[N] =0.16


pK2 = 10.16


i T OH- CH2OH


UOHn



CH3


p 0OHl NEUrH2 L [N-OX
pe(=1 0.10 NEUTRAL [N.Doxi


K4

pK =9.36


ZWTTERION [Z-DOX]


MONOANION [MA-DOX]


Protolytic equilibria of (I) in aqueous solution


Scheme 1.1:









(191,192) and the number of hydroxyl groups on the sugar

moeity (191,193) (lowering the number of hydroxy groups on

the sugar increases the pKa of the amino group). In

addition, the interaction of the amino sugar with the

quinoidal moeity (161) can influence the pKas. There is a

suggestion that the side chain of (I) might be involved in

some stabilizing interactions with the amino sugar via

hydrogen bonds. This could also conceivably influence the

ionization of the NH3 group (142).

Self Association.

Many dyes, including compounds endowed with biological

activity such as acridines (194-196), purines (197), and

actinomycin D (198), can form molecular aggregates by

vertical stacking, even in dilute solutions. The tendency of

anthracycline molecules to self-associate to dimeric or even

polymeric aggregates was first revealed by X-ray diffraction

(36) of freshly crystallized (VI) from acidified methanol.

Studies conducted using circular dichroism and NMR

demonstrated that for low concentrations (<5 mM) the

association process could be represented by a dimerization

model with an association constant of 570-700 M-'1, whereas

at higher concentrations further association to polymers was

possible. Spectroscopic studies (199) reported somewhat

higher values of 3.0 X 103 M-1, 1.83 X 103 M-1 and 1.16 X 103

M1 at 25C, 35C and 45C, respectively, at pH 7.0 (0.01 M

phosphate buffer). Thermodynamic parameters estimated by









microcalorimetric measurements (199) gave the association

constant of (I) as 1.1 X 103 M-1, with a A Ha = -9.6

kcal/mole and A Sa = -18.3 eu (at 25C). These values are

similar to enthalpy values of other dyes undergoing

association (199). An enthalpy decrease of 8-9 kcal/mole

suggests that the dimerization process is based on the

stacking interactions of the planar chromophores. Eksborg

(142) studied the dimerization process by absorption

spectroscopy and solvent extraction techniques using a

chloroform:l-pentanol (9:1) system as the organic phase. He

estimated association constants of 15,848 and 19,952 M-'

with the former technique and 30,199 and 30,902 M-1 from the

latter for (VI) and (I), respectively. Spectroscopic studies

carried out in the pH range 9.8-10.5, where the effect of

MCDOX could be ignored, yielded values that were attributed

to the neutral species. He further concluded that the

interaction of the molecules was due to their ring systems

since no dimerization of the negatively charged phenolate

ion was found. He also reported tetramerization constants of

2 X 1012 and 1 X 1012 M-1 for 'I) and (VI), respectively.

Martin (173), while working on the interaction of (VI) with

calf thymus DNA, studied the phenomenon with respect to

temperature, ionic strength and solvent composition, using

visible absorption, fluorescence and circular dichroism

spectroscopy. He estimated an association constant of 6,400

M-1 and found that increasing the ionic strength increased







24

the aggregation of (VI) almost twofold. This was similar to

behavior reported by Blauer (200), and was thought to be a

modification of the solvent structure by the salt which had

been observed in the case of actinomycin D (198). Chaires et

al.(171) studied the self association of (VI) using

absorbance, sedimentation and NMR measurements and used an

indefinite association model to postulate a dimerization

constant of 1,500 M-1. Thermodynamic values from NMR data

indicated stacking of the planar aromatic structures.

However, similar studies on actinomycin D (201) did not

favor the indefinite association model. NMR data showed that

the ring protons and the 0-CH3 on Ring D (Figure 1.1) was

the major site of interaction. The side chain COCH3 and

sugar moieties did not affect the aggregation process.

In an attempt to resolve the discrepancies in the

literature, Arcamone et al.(172) studied the self-

association of (I), (VI) and related compounds in aqueous

solutions using absorption spectroscopy. Their method

involved the measuring of the difference spectra of these

compounds by using cells of different path lengths in the

sample and reference beams. They were able to differentiate

between the monomer and the dimer spectra. They estimated

the association constant for (I) to be 1.29 X 104 M-' which

was dependent on buffer composition and ionic strength. An

increase in the ionic strength resulted in an increase in

the association constant. Since the association constant of









(VI) was more affected by increase in ionic strength than

that of (I), it was concluded that the OH group at the C-14

position of (I) was involved in the association process

either directly via hydrogen bond formation, or indirectly

via dipole effects. McLennan et al. (202) used NMR

spectroscopy, and particularly the shift of the OCH3 protons

to estimate the self-association. Using the rotational

correlation time of the associated species as compared to

the monomer species, they were able to estimate a value of

4.0 X 10 M-'1 at 22C for the association constant.

The effect of formulation additives like NaCl and

methyl p-hydroxybenzoate on the self-association process

have been studied (203) using viscosity and NMR

measurements. The addition of higher concentrations of NaCi

increases the viscosity of (I) solutions indicating that an

increase in the ionic strength facilitates self-association.

By contrast, the addition of various amounts of methyl p-

hydroxybenzoate reverses the effect of the NaCl. Most

studies (142,171,199,202,203) conclude that association

takes place by the vertical stacking of (I) molecules at the

chromophoric part.The forces responsible for this

association are not completely understood. A R-n interaction

between the quinone chromophores is thought to occur, but

the contribution of hydrophobic or electrostatic

interactions cannot be ruled out (132). In any case, the

aggregation of (I) molecules should not be important at the









biologically active concentration of 106 M and less. This

may be due partly to the high protein binding of the drug

(90%) resulting in very low free drug concentrations.

The self-association literature discussed above shows

conflicting values of the association constant and is

inconclusive about the effect of pH and ionic strength on

the association process. To this end, a systematic study of

the association process of (I) in aqueous solutions was

carried out, using fluorescence with a 'front surface

illumination' technique. A mathematical model allowed for

the prediction of the association constants of each

individual ionic species generated due to pH changes and the

influence of ionic strength on these processes.

Adsorption to Glass.

During studies involving uptake of (I) into tumor cells

in vitro, Tomlinson and Malspeis found significant loss of

(I) from Hank's culture medium (204) to the walls of the

glass container. A detailed study carried out in different

containers at 37C showed that about 7.3% of (I) was bound

to glass, 4.6% to polyethylene Petri dishes, and 45% to

polytetrafluoroethylene (PTFE). However, there was no

binding to polypropylene and to siliconized glass. In order

to minimize losses due to adsorption, they recommended that

a) the number of wall surface contacts should be reduced, b)

an appropriate container material like polypropylene or

siliconized glass should be chosen, c) the walls of the









container should be presaturated with the drugs and d)

cosolvents should be used. Adsorption of hydrophobic drugs

to the walls of a container have been reported by various

authors (205-212). Invariably the use of cosolvents

(containing a more nonpolar solvent than water) or

protecting the glass surface by siliconization helps to

alleviate the problem. Similar effects were observed for (I)

during analysis (213) and when sampling from solution (214).

Conventional approaches like cosolvent extraction

techniques, have been used to resolve the difficulty

(213,215,216).

For HPLC analysis desipramine has been added to mobile

phases (217,227) and to extraction media (218-221) to

prevent the adsorption of (I) to HPLC hardware and to glass.

Desipramine appears to occupy the same adsorptive sites as

(I) on these surfaces.

In complete contrast to the study of Tomlinson and

Malspeis (204), Schutz et al.(174) found that (VI) binds to

siliconized glass below concentrations of 2 x 10-' M.

Andrews et al. (222) found that silylation of glassware had

no effect on the extraction of (I) from aqueous solutions.

In studies (223) carried out to compare the stability of

antitumor agents in glass and polyvinylchloride bags (I) was

found to be more stable in the plastic bags than in glass

bottles. Binding to end line (0.2 pm) infusion filters was

found to be negligible (224), as greater than 96% of the









drug was recovered when a saline solution of the drug was

passed through the filters at 80-100 ml/hour for 2.5 hours.

The apparent complexity of the nonspecific adsorption

of (I) to glass prompted the design of a study to quantify

and explain the phenomenon as a function of concentration,

temperature and pH of aqueous solutions of (I).

Analytical Methods.

A fluorometric method of anthracycline analysis

developed by Schwartz (225) involved the extraction of

tissue and plasma samples with isoamyl alcohol and

measurement of the solution's fluorescence with excitation

at 490 nm and emission at 560 nm. Calibration curves were

linear in the range 0.05-10 Lg/ml. Prior to the extraction,

0.2 ml of a 33% w/v solution of silver nitrate was added to

precipitate the proteins, nucleotides, etc. and to aid in

the subsequent release and extraction of the anthracyclines.

This method was nonspecific for (I) and its major

metabolites. The oxidative effects of the silver ion were

not evaluated.

Cummings et al.(226) determined the concentrations of

(I), (II) and their 7-deoxy aglycones in human serum by

HPLC. A reverse phase C-18 column with fluorescence

detection (excitation 480 nm and emission 560 nm) was used.

The mobile phase was 62.5% orthophosphoric acid, pH 3.2 (5mM

final concentration) and 37.5% of a mixture of methanol,

acetonitrile and 2-propanol (either 12.5:12.5:12.5 or









15:15:7.5). Elution was isocratic with a flow rate of 2.5

ml/min. The limits of detection of (I), (II) and (IV) were

4.4 ng/ml, 2.1 ng/ml and 4.2 ng/ml, respectively. The

internal standard was (VI). Serum samples were extracted

with chloroform:2-propanol (2:1). Unfortunately, the binding

of (I) to PTFE (206) tubes used in the study was not

evaluated. Thus, their report that 9.1% of (I) degraded per

hour at 25C cannot be accepted.

HPLC was also used by Oosterbaan et al.(144) to analyze

(I), (II), carminomycin and the 4'-epi and 4-demethoxy

derivatives of (I). The analytical system consisted of a 15

cm reverse phase C-8 column (8 im) with a dual pump solvent

delivery system. The first pump delivered a 10 Lg/ml

solution of desipramine in demineralized water to a 25 cm

silica column used to concentrate the sample. The

desipramine prevented the adsorption of the anthracyclines

to column and tubing materials. The second pump delivered

the mobile phase of acetonitrile:citric acid buffer, pH

2.2:water (35:10:55% v/v) and transferred the concentrated

anthracyclines from the silica column onto the analytical

column. Calibration curves were established in the range 5-

500 ng/ml. The limit of detection for (I) was 0.5 ng/ml and

for (II) was 0.4 ng/ml. Detection was by fluorescence

(excitation at 474 nm and emission at 590 nm).

In an attempt to prevent the adsorption of (I) and its

major metabolite (II) on to the chromatographic system,









Ichiba et al.(227) used normal phase HPLC with a

fluorescence detector (excitation at 480 nm and emission at

590 nm). The stationary phase was a 25 cm silica gel column

and the mobile phase was a mixture of methylene chloride,

methanol, glacial acetic acid and 0.01 M MgCl2 solution

(200:50:7:5 v/v). The addition of Mg2 ions was purported to

reduce adsorption and give sharper peaks. The calibration

curves for (I) and (II) were linear down to 1.5 ng/ml, when

(VI) was used as an internal standard.

Watson et al.(228) studied the effects of different

surfactants on the HPLC capacity factors of (I) and its

major metabolites. They used a reverse phase 10 cm C-18

column (5 Jm), and a mobile phase of acetonitrile:phosphoric

acid (0.01 M) (35:65% v/v). Detection was by fluorescence

(excitation 450 nm; emission 550 nm). They found that 6mM

Brij-35, polyoxyethylene lauryl ether, a non ionic

surfactant, in the mobile phase reduced the analysis time

from 18 to 12 minutes. The limit of detection for

doxorubicin was 1.5 ng/ml.

Weenan et al.(229) successfully separated and

quantified doxorubicin, 4'-epidoxorubicin and their

metabolites using a 25 cm reverse phase C-18 column (3 pim).

The mobile phase consisted of 70 parts of a 0.06 M potassium

phosphate monobasic buffer (pH 4) and 32.5 parts (v/v) of

acetonitrile. A fluorescence detector (excitation at 470 nm;

emission at 580 nm) was employed. Plasma samples were









treated with chloroform:isopropanol (4:1 v/v) and the

organic phase evaporated to dryness. The total analysis time

was about 30 minutes. Calibration curves were linear from 3

ng/ml to 3 jg/ml.

Riley et al.(230) used electrochemical detection to

quantify a series of closely related anthracycline

antibiotics, including (I). Taking advantage of the

oxidation-reduction behavior of the phenolic groups and

quinone moieties of (I), they were able to effectively

quantify (I) to 1 ng/ml in plasma. The chromatographic

conditions consisted of a reverse phase C-18 column with an

optimum mobile phase of acetonitrile:isopropanol:0.1M

phosphate buffer, pH 4.5 (25:3:72 v/v). The electrochemical

detector was set at +0.8 V for the best signal-to-noise

ratio.

In similar studies (231-233) (I) and its major

metabolites were analyzed in biological samples using

reverse phase HPLC systems with a C-18 column (232), a

phenyl column (231) or a cyano column (233). In all cases,

mobile phases had varying concentrations of acetonitrile and

aqueous format or phosphate buffers, and fluorescence

detection was used (excitation 470-475 nm; emission 565-580

nm).

All authors using HPLC claimed a high sensitivity,

specificity and reproducibility, whereas in prior assays of

(I) specificity was a major problem, since fluorescence and









absorbance of metabolites and degradation products had

identical spectral characteristics as the parent compound.

Most studies did not mention whether the assayed (I)

solutions were protected from light, oxygen or glass

binding, processes which would affect the concentration of

(I) solutions.

The assay presented herein for (I) and its major

metabolites and/or degradation products (e.g.(II),(III),

(IV) and (V)) were developed in preparation for kinetic

studies on the degradation of (I).


Stability and Kinetics of Degradation


In order to establish appropriate storage and handling

conditions, all drugs are evaluated for their stability

characteristics as a function of temperature, solution

composition (pH, ionic strength, solvents, effects of

buffers and additives), light and oxygen. These studies are

usually carried out under accelerated, conditions wherein

the drug is subjected to extreme levels of the conditions

described above. The results thus obtained are

mathematically interpreted using simple laws of

thermodynamics, chemical kinetics and statistics, and are

extrapolated to yield information about the stability

(chemical and consequently biological) at room or ambient

temperatures, refrigerator and freezer conditions, in the

presence or absence of light, oxygen and moisture and in the







33

solid or solution form. These studies are carried out on the

unformulated as well as formulated drug. A logical extension

to these studies is the degradation behavior of the drug,

the rate and extent of formation of the degradation

products, the chemical pathways by which this degradation

occurs, the pharmacological usefulness or uselessness of

such products and the conditions under which one could

inhibit or accelerate the formation of these products. This

information helps all branches of the scientific community

who may handle the drug, and would therefore need this

knowledge, to design better studies.

Stability Studies on Doxorubicin.

The anthracyclines, especially (I) and (VI), have been

extensively studied for their stability in the solid and

solution state. However as will be detailed below, most

studies did not attempt to evaluate the whole spectrum of

conditions that can affect the stability of (I). Some

studies indicated that the conditions were considered, but

for some reason the authors failed to report the results in

their publications.

Stability in the solid state. In the crystalline state

(I).HC1 was stable (<10% degradation) for more than five

years at room temperature, without chemical modification and

without loss of activity (141,234). The lyophilized

formulation containing lactose as a diluent was stable for

more than two years under the same conditions.









Stability in solutions. The solvolysis of (I) is

catalyzed by hydrogen and hydroxide ions (specific acid-base

catalysis). This catalysis is enhanced in the presence of

other additive ions (ionic strength), buffer components

(general acid-base catalysis) and an increase in

temperature. In addition, (I) is photolabile in solution and

the photolysis is catalyzed by oxygen. The latter also plays

a role in the absence of light. Metal ions that can interact

with (I) may also catalyze the oxidative degradation of (I)

in solution.

Effect of pH. The hydrolysis of (I) increases with

higher acidity, temperature and ionic strength (235,236).

The log kobs vs pH (236,237) (kobs is the pseudo-first-order

degradation rate constant) profile shows a slope of -1 up to

a pH of 3 which is indicative of a specific hydrogen ion

catalyzed reaction. In this region hydrolysis is solely

catalyzed by hydrogen ions with no influence of the

conjugate base (237). Different authors report that the pH

rate profile approaches a minimum (pH of maximum stability)

between pH 3 and 4 (238) or between 4 and 5 (237). An

earlier study (239) indicated considerably greater stability

at a pH of 2.6 compared to pH 4.8 or 5.5. In this case the

authors failed to do a detailed pH study of (I) and relied

on studies at two pH values to claim their result. At pH

values greater than 5 the overall rate of degradation

increases with increasing pH (234,237,240,241). The slope of









the alkaline phase (pH>6) of the pH rate profile gives a

value of + 0.5 (237). This slope holds true up to pH 10,

where there is an indication of a slight inflection in the

curve, from ionization of the phenolic proton of the

molecule resulting in the differential rates of hydrolysis

of the two species formed. The authors (237) considered all

four species of (I) formed from the protolytic equilibria of

the amino group (of the amino sugar), and the two phenolic

protons on the chromophore. Each one of these four species

could have its C-9 side chain in the keto or the enol form

especially at pH ranges greater than 10 (the pKa of C-17 a-

ketol side chain in hydrocortisone is 11.05 at 25C) (242).

This factor was determined to be inconsequential in the pH

range studied: 1-11. The overall observed rate constant was

fitted to the following equation:

kObs= (kO+koH- [OH-] +kH.H] ) .[H (1.1)



Here kobs was the overall pseudo first order rate constant;

ko was the first order rate constant for the solvent

catalyzed degradation; koH- and k, represented the specific

second order rate constant for the hydroxide and hydrogen

ion catalyzed degradation; and fi was the fraction of

species involved at each pH. At pH values less than 4, ko

was determined to be negligible (not different from zero).







36

Effect of ionic strength. In the range from 0.1 to 0.4

M, the ionic strength does not seem to play any significant

role in the hydrolysis of (I) at pH values above 4 (237).

However, there does not seem to be any agreement from

observations made below pH 4. Studies carried out in acidic

conditions (236) using HC1 as the catalyzing acid showed no

influence of the ionic strength at values up to 0.5 M. In

contrast, studies carried out in the same pH range, but

using perchloric acid as the source of protons, revealed a

linear dependance of the log of the overall pseudo first

order rate constant on the square root of the ionic strength

(slope: 0.365; intercept: -4.53) in the range from 0.058 to

0.56 M (237). At ionic strength values greater than 0.56 M

no linearity existed.

In the present studies of the hydrolysis of (I) the

ionic strength was maintained at 0.2 M throughout by adding

appropriate amounts of KC1.

Effect of solvent. There is no detailed study of the

effect of organic cosolvents (methanol, ethanol etc) on the

degradation of (I). Methanolic solutions of (I) degrade

rapidly (246). In water a 2 mg/ml solution of (I) has a pH

of around 5.5 at room temperature and is stable (10%

degradation) for at least a month (141). However, the

general consensus is that water is an inappropriate solvent

for (I) (243,244,245,246). In sterile water for injection

(USP), (I) can be stored at -20C for at least a month







37

(247). Eksborg (248) recommended storage of (I) solutions in

0.01-0.1 M phosphoric acid, which is ludicrous in light of

the results of other studies (236,237).

Effect of light. In solutions (I) is photolabile

(59,249-252). This decomposition is dependent on pH (251),

the nature of the solvent and the concentration of (I). The

rate of degradation was shown to be inversely proportional

to the concentration (250), but the data were fitted to a

first-order kinetic equation (250) in obvious contradiction

to its concentration dependency. Explanations given to

rationalize this anomaly were related to the intensely

strong color of concentrated solutions of (I), which in turn

self-protect themselves from photolytic degradation (250).

The fact, that there was negligible photolytic degradation

of (I) in bile supported the hypothesis, since the strong

color of bile protected the sample from photolysis. The free

radical scavenger butylated hydroxy toluene (BHT), was found

to reduce photodegradation (249), and its presence was given

as an alternate explanation for the protection properties of

bile, since animal feed is a source of BHT which gets

excreted primarily through the bile and might be present in

it (253). Williams and Tritton (252) found that (I) was

photolytically (with 366 nm light) converted to a product

that did not exhibit any cytotoxicity against Sarcoma 180

cells, probably because these products (probably polymeric

in nature) were not taken up by the cells. In photolytic









studies of (VI) (58) under anaerobic conditions only the

aglycones were observed as photoproducts, including a

product with an aromatized A-ring (see Figure 1.1) and with

the C-9 acetyl group replaced by the C-9 hydroxyl group.

Spin trapping and direct electron spin resonance studies

(254) concluded that, upon irradiation at 310 nm, (I) and

(VI) generate the superoxide anion radical and at least two

carbon centered radicals which could degrade (I) to yield

the A-ring aromatized deacetylated aglycone (58). This was

not observed when the samples were irradiated at 490 nm,

indicating that the formation of the superoxide free radical

was not the photolytic pathway in the visible region of the

spectrum. Additional studies (255) confirmed the formation

of the superoxide free radical ion via a one electron

transfer mechanism in both aprotic and protic media. The

semiquinone formed as a consequence yields a diamagnetic

dimer which further decomposes to yield the 7-deoxyaglycone

(V) of (I). Glutathione (GSH) added to buffered solutions of

(I) in clear glass vials appreciably enhanced the stability

of (I) (t1,2 with GSH = 462 hours vs 38 hours without GSH)

from photolytic degradation (256).

In the present case, care was taken to ensure that the

solutions of (I) were protected from light (wrapped in

aluminum foil) at all times and each solution was purged

with nitrogen before and during sampling to exclude oxygen

from the reaction mixture.









Effect of temperature. In general an increase in the

temperature increases the overall rate of hydrolysis of (I).

This increase follows the Arrhenius equation (257).

Ea
kobs = Ae RT (1.2)



where kobs is the overall degradation rate constant, A is the

frequency factor, Ea is the activation energy, R is the gas

constant, and T is the absolute temperature. Using elevated

temperatures that result in accelerated conditions of

hydrolysis and the Arrhenius equation, one can predict the

half-life and shelf-life at room temperature and under

refrigerated conditions (4-5C).

The Arrhenius parameters have been estimated for the

hydrolysis of (I) by various authors (236,237,238) and are

summarized in Table 1.1. At pH values less than 3.0 one can

see that the estimate of the activation energy is not

consistent among the various authors (even though specific

acid catalysis was shown) indicating that the mechanism of

hydrolysis depends on the mineral acid used.

At pH values greater than 4.0 the activation energy

varies inconsistently with pH. This suggests a different

hydrolytic pathway (and a different rate limiting step) for

the degradation of (I). The formation of more than one

degradation product above pH 4.0 accounts for the varying

estimates of Ea. It is also possible that the different pH









dependence of the Ea may be due to the varying degrees of

protonation of (I) in this pH region (237).

Janssen et al.(238) found that the Arrhenius plot

deviated from linearity between 61 and 72C and concluded

that there may be a change in the rate-controlling mechanism

in that temperature region. This appears to be unlikely, and

experimental error seems to be a more plausible cause since

the authors used a kinetic study at 4C as one of the points

to establish linearity in the Arrhenius plots. It can be

shown that the half life of degradation of (I) is very large

and prone to error at 4C. Moreover, at 4C glass binding

accounts for about 3-8% of loss of (I) from solution in the

concentration range 1-100 ig/ml (see present study) and

should be considered in the interpretation of the results.










Table 1.1: Arrhenius parameters for the hydrolysis of (I).

Ea A CONDITIONS REFERENCE
kJ/mole sec'1


8.2 x i07


3.3 x 107


4.0 x 108


2.6 x 1011


HC1 (0.01-0.5 M)
pH = 0.43-2.13
. = 0.2 or 0.5

HC104
Ho/pH = 0.43-3.0
S= 0.4

Acetate (0.01 M)
pH = 4.0

Acetate (0.01 M)
pH = 5.0

Phosphate (0.01 M)
pH = 6.0

Phosphate (0.01 M)
pH = 7.0


Tris (0.01 M)
pH = 7.4


1.2 x 1012


8.5 x 1010


2.3 x 10"


5.0 x 107


Phosphate
pH = 7.4

Phosphate
pH = 7.4

Phosphate
pH = 8.0

Carbonate
pH = 9.0

Carbonate
pH = 10.0


(0.01 M)


(0.01 M)


(0.01 M)


(0.01 M)


(0.01 M)


* In these studies g = 0.3 and [EDTA] = 5 x 10-4 M


92.0



114-128


(236)


(237)*


89.7


86.6


90.4


104.1


(237)


(237)


(237)


(237)


(238)


(238)


(237)


(237)


(237)


(237)


105.5


95.7


69.1


67.3









Effect of concentration. In studies on the photolytic

degradation of (I), there was no evidence of a concentration

dependent hydrolysis of (I), even though there was a

concentration-dependent photolytic degradation (250). This

was confirmed by Poochkian et al.(258) for 10 and 20 jig/ml

samples and by Beijnen et al.(237) in the concentration

range 1-20 jig/mi. Janssen et al.(238) claimed that (I)

degrades faster in solutions containing 500 jig/ml as

compared to 50 jig/ml at pH 7.4. A possible explanation for

this discrepancy was provided by Beijnen et al.(237), who

noticed the formation of precipitates when (I) was degraded

at high concentrations. Analysis of the precipitates

revealed the presence of considerable amounts of undegraded

(I). This could disrupt the kinetics and make it difficult

to interpret the results.

The present studies on the concentration dependency

were conducted at 10 and 100 jig/ml in 0.2 M HCl.

Effect of buffers. The overall rate of hydrolysis

depended on the concentration of the buffer used to maintain

a certain pH (141,237,238). The buffer species could be a

general acid and/or base and catalyze the degradation of

(I). At pH values less than 10, catalysis by acetate,

phosphate and carbonate were demonstrated (237). There was

no buffer catalysis at pH values greater than 10. The kobs vs

buffer concentration plots deviated from linearity at higher

buffer concentrations. The authors failed to use their data









to estimate the specific second-order rate constants for

each of the conjugate acids and bases of the buffers used in

their study.

In order to evaluate the kinetics of hydrolysis of (I)

in solution, a carefully designed study was established to

consider the effect of pH, temperature, and to determine the

catalytic contribution of each buffer species on the

hydrolysis of (I). Ionic strength was maintained constant

throughout these studies to avoid interference from this

variable. Each hydrolysis study was conducted at four

different temperatures to give statistically (t-values < 3

at a = 0.05; where a is the type 1 statistical error) better

and more reliable estimates of the Arrhenius parameters

rather than the fewer than 4 conducted to justify the claims

of a Farmitalia patent (U.S. Patent # 4946831, August 7,

1990). The pH of maximum stability (from the log kobs vs pH

profile) is evaluated. Possible mechanisms and degradation

pathways are proposed.


Stability of (I) in Pharmaceutical Formulations.

The commercial preparation of (I) in its hydrochloride

form (Adriamycin, Adriablastina) is a lyophilized powder

containing 5 parts by weight of lactose as a diluent. The

compound in this form is quite stable when its moisture

content is less than 1% (234,259). These preparations might

contain trace quantities of the aglycones as impurities

which are inherent to the mode of preparation (260). In a









recent study (261) it was shown that dissolution

(reconstitution with an appropriate infusion fluid) prior to

administration was facilitated when the freeze-dried mixture

contained sub-preservative amounts of hydroxy benzoates

(methyl and propyl parabens). Once dissolved, the

anthracyclines were susceptible to hydrolytic and photolytic

degradation.

The manufacturer (Adria labs, Cincinnati, Ohio) states

that the stability (>10% degradation) of the prepared

infusion solution is 24 hours at room temperature and 48

hours in the refrigerator (4-10C). The manufacturer advises

against freezing (262), as this might degrade the drug. In

contrast, some of the literature (247) suggests that, when

reconstituted and filtered, (I) is stable for 6 months in a

refrigerator (4C) and 30 days when frozen (247). Repeated

freezing and thawing does not have any effect on the

stability (247). In 0.9% sodium chloride and 5% dextrose

solution (co-formulated with methotrexate) (I) was stable

(10% degradation) for at least 1 month at -20C (263). The

stability of (I) in four different infusion fluids was

studied by HPLC at ambient temperature (21C) (258). The to.9

(shelf life: defined as the time it takes for 10%

degradation) values for (I) in 5% dextrose injection (USP)

at pH 4.5 was estimated at 100 hours; about 63 hours in 0.9%

sodium chloride injection (USP), pH 6.2; about 28 hours in

lactated Ringer's injection (USP), pH 6.3; and 24 hours in









Normosol-R, pH 7.4. This indicated a pH dependence on the

stability of (I) in the various infusion fluids with a 5%

dextrose solution being the most stable. In contrast,

Ketchum et al. (264) reported that (I) formulations

reconstituted in 0.9% sodium chloride showed no loss in

potency (determined spectrophotometrically by a non-specific

method) during storage for 28 days at 5 or 25C. Beijnen et

al.(265) studied the degradation of (I) in four different

infusion fluids in the absence of light. Using a sensitive

and specific HPLC assay they found that (I) had a shelf-life

(10% degradation) at 25C of at least 4 weeks in 5%

dextrose, pH 4.7 and a mixture of 3.3% dextrose and 0.3%

NaCI, pH 4.4. In lactated Ringer's solution, pH 6.8 the

shelf life was 1.7 days and in 0.9% sodium chloride, pH 7

the shelf life was 6 days. The stability of (I) in the

presence of vincristine (266) was investigated in three

different infusion solutions at three different

temperatures. In 0.9% sodium chloride injection solution and

0.45% sodium chloride and 2.5% dextrose injection solution,

(I) was stable (<10% degradation) in the presence of the

vinca alkaloid vincristine for at least 7 days at 15 and 4C

(<5% loss). However, when placed in a 0.45% NaCI and

Ringer's acetate injection solution at the same

temperatures, (I) was not as stable (>20% loss after 7

days). When delivered by an implanted battery operated

infusion device, (I) degraded by more than 10% in 14 days at









37C (267). In portable pump reservoirs (268) (I) at a

concentration of 2 mg/ml was stable (<10% degradation) for

up to 14 days at 3 or 23C and for an additional 28 days at

30C. In a hydrophilic ointment (I) was stable (<10%

degradation) for four weeks at 5C and at room temperature

(269). There was no significant difference in the stability

of (I) when encapsulated in liposomes (238).

Stability in Biological Media.

The analysis and stability evaluation of (I) and its

metabolites and degradation products in biological

preparations is crucial for pharmacokinetic studies, tissue

binding and uptake studies, and for binding studies to

macromolecules like DNA, plasma proteins, etc.

Heparin an anticoagulant used in blood collecting tubes

binds (I) to form a complex (132). The formation of the

complex interfered with the bioanalysis of (I) from blood

(226), and hence heparin was avoided. In studies carried out

on implantable devices in dogs, heparin and (I) could be

mixed provided the concentration of heparin was below 1.3

units/ml (270). Doxorubicin was also incompatible with 5-

fluorouracil, dexamethasone, sodium phosphate (271) and

NaHCO3 (272). With the latter, the unknown degradation

product retained its anti-proliferative property but lost

its lethality to Colo 320 (colon cancer) and CCL2 (cervical

cancer) tumor cell lines. Red blood cells degrade (I)

rapidly by a cytoplasmic aldo-keto reductase enzyme







47

(248,273,274). Thus, blood should be rapidly centrifuged and

the RBC removed from the collected blood (275).

In plasma (I) is stable (<10% degradation) for long

periods of time at -80C (248). However the reported

stability at -20C in plasma is conflicting (144,273).

Oosterbaan et al. (144) found (I) to be stable (<10%

degradation) for 14 days at -20C even if repeatedly frozen

and thawed, while Eksborg et al.(273) came to the opposite

conclusions. At 4C plasma samples of (I) were found to have

unchanged concentrations for at least 24 hours (273,276).

Urine samples can be preserved at -20C (244) or, after

acidification, at 4C in the dark for at least a month

(277). The biological activity of (I) (tumor cell kill) is

a) reduced by 50% in vitro in the presence of 25 mg/mL of

human albumin (278); b) affected by cell density in a

monolayer system (279); and c) dependent on temperature

(280-283). Pavlik et al.(272) showed that (I) degraded in

the dark in a tissue culture medium (not containing serum)

to a product that was cytostatic but not cytotoxic to Colo

320 and CCL2 tumor cell lines. Doxorubicin was stable under

the conditions of a clonogenic assay (284) and under storage

conditions of -40 and -196C with a t,/2 greater than 6

weeks. Under conditions of the assay at 37C (I) had a

degradation t1/2 of 29 hours.

In the conventional organic solvents that extract drugs

from biological matrices, no detailed stability study of (I)









has been reported.

Farmitalia and U.S.Patent # 4946831 (August 7, 1990).

The Italian pharmaceutical company Farmitalia Carlo

Erba, Milan, Italy filed a patent for a new ready-to-use

injectable solution of (I). It was contended that, prior to

administration (I) did not have to be reconstituted from

lyophilized powders, where spillage and loss during the

reconstitution process could occur. The formulation was a

hydrochloride salt of (I) dissolved in nitrogen-purged

sterile water for injection with the pH of the resulting

solution adjusted to 3.0 using HC1 and stored at 4C.

Kinetic and statistical analysis of the data presented

does not appear to support the claims of the patent. The

'successful' formulation presented (example 2 in the patent)

was analyzed by fitting the Arrhenius function (equation

1.2), at temperatures of 55, 45 and 35C (Table 2 in the

patent). The results of the analysis are presented in the

Table 1.2. The R2 values of the least squares plots of in

concentration vs time were 0.9976, 0.9166 and 0.9699 at 55,

45 and 35C, respectively. However, when the obtained rate

constants kobs were plotted in accordance with the Arrhenius

equation (in kobs vs l/T, where kobs are slopes of the in

concentration vs time plots and T is the absolute

temperature) an R2 value of 0.999995 was obtained. This

suggests an almost perfect (R2=1.0000) Arrhenius fit of the

data and hence an error-free estimate of the rate constants









for the corresponding half-lives and shelf-life (to.9) at 4

and 8C given in their table. It is suspicious that kinetic

studies carried out at only three higher temperatures (55,

45 and 35C) yielded statistically poor correlation

coefficients (notwithstanding that errors in rate constants

obtained at elevated temperatures are smaller than those

obtained at lower temperatures) and yet the values of the

shelf life (to.9), estimated from the Arrhenius equation were

predicted with such precision. Estimates of rate constants

at room or refrigerated temperatures extrapolated from only

three elevated temperatures have high standard errors since

they have only two degrees of freedom, in which case

statistically significant (a = 0.05) results are difficult

to obtain.

The claims of Farmitalia that the stability maximum was

observed at pH 3, based on their Arrhenius and kinetic data,

disagreed with the pH profile of Beijnen et al. (237), whose

log kobs vs pH profile, corrected for buffer effects, had a

pH of maximum stability between 4 and 5.









Table 1.2: Farmitalia patent data fitted to the
Arrhenius function (equation 1.2 in text).

Temp. kobs In 1/T R2* Ea#
K (wk-1) kobs K-1 kJ/mole

328 0.33769 -1.08563 0.003049
(0.9976)
318 0.09192 -2.38684 0.003145 0.999995 112.19
(0.9166) (0.36)
308 0.02335 -3.75716 0.003247
(0.9699)

Numbers in parentheses are the R2 values of the
regressions in the estimates of the rate constants.
Absolute temperature (273.15 + tC)
* Correlation coefficient
# Activation energy


Farmitalia's scientific rationale for developing this

formulation was flawed due to a misinterpretation of the log

kobs-pH rate profile. The pH of maximum stability in the

presence of buffers that catalyze degradation (general acid-

base catalysis) shifts towards the acid region as the buffer

concentration increases. Since studies in buffer solutions

showed the pH of maximum stability to be about 3, Farmitalia

assumed that, if pH 3 was attained by a non-buffering

species such as HC1, one would have a 'ready-to-use' product

with a better shelf life. Additionally, as demonstrated in

this thesis, (I) degrades in solution in the presence of

oxygen. This in turn results in a pH drop from 4.95 to 2.97.

Thus a solution of (I) would inherently come to a pH of

about 3 after degradation.

This dissertation attempts to: 1) eliminate some of the

discrepancies present in the literature concerning self-







51

association of (I) in aqueous solution, 2) evaluate the

glass binding of (I) and the effects of pH, temperature and

silylation on the binding, and 3) establish the kinetics of

hydrolysis of (I) in aqueous solution to explain unresolved

issues concerning pH of maximum stability, degradation

behavior of (I) and possible mechanisms of degradation.








CHAPTER 2
MATERIALS & METHODS

Materials

Test Compound.

Doxorubicin hydrochloride (I) was initially purchased

from Sigma Chemical Co. St.Louis, MO. It was subsequently

received as a gift from Adria Laboratories, Columbus, OH and

from Farmitalia Carlo Erba, Milan, Italy. Metabolic and

degradation products like doxorubicinol (II), doxorubicinone

(III), doxorubicinolone (13, dihydroxydoxorubicinone) (IV),

and 7-deoxydoxorubicinone (V) were received as a gift from

Farmitalia Carlo Erba, Milan, Italy, through the auspices of

Dr.Federico Arcamone. Daunomycin (VI), used as an internal

standard during chromatographic experiments, was obtained

from Sigma Chemical Co. St. Louis, MO. All samples received

were assumed to be pure and used as received, excepting (I),

which was subjected to multiple tests to evaluate its

purity. These included melting point determination,

chromatographic analysis and differential scanning

calorimetry.

Each of the samples received was stored in glass tubes

placed in larger tubes containing an excess of anhydrous

calcium sulfate as a desicant. These tubes were protected

from light and stored in a freezer maintained from 0-5C.









Reagents/Solvents.

All chemical reagents, pH standards and solvents used

were of HPLC or A.C.S. analytical grade and were used as

supplied (Fisher Scientific, Fairlawn, NJ).

Hexamethyldisilazane in hexane (PCR Inc., Gainesville,

FL) was used as a glass silylating agent. Specific reagents

used for specific purposes will be described in the

appropriate sections.

Water used for preparing mobile phases, buffers,

reaction media etc. was deionized, distilled water which was

boiled, cooled and filtered through a 0.2 pinm nylon membrane

filter (Rainin). Water used for any other purpose was

deionized, distilled water.


Instrumentation


Melting Point determination.

Approximately 10 mg of (I) was powdered in a dry glass

mortar. A portion of this sample was packed into a melting

point capillary tube (0.8 X 1.10 X 90 mm; KIMAX 51, Kimble

products, USA). The tube was heated in a melting point

apparatus (Thomas Hoover Capillary melting point apparatus,

A.H. Thomas Co., Philadelphia, PA) at an average rate of

22C/min from room temperature to 220C.

Differential Scanning Calorimetry (DSC).

All samples of (I) received had their purity evaluated

using DSC. In order to reduce the degradation of (I) at









temperatures lower than its melting point, samples were

heated at a rate of 20C/min rather than the 2-5C/min

recommended for thermal equilibration. Using a micro balance

(Cahn electrobalance, Ventron Corpn., Cerritos, CA),

approximately 1 mg of (I) was accurately weighed in an

aluminum pan and crimped. This pan was placed in the DSC (PE

DSC 7, Perkin Elmer Corpn, Norwalk, CT) against a blank pan

and heated at a rate of 20C/min up to 220C. The resulting

endotherm was recorded on a graphics plotter (Perkin Elmer

2, Perkin Elmer Corp., Norwalk, CT) and the data processed

with the supplied software on the interfaced microprocessor.

pH measurements.

To measure pH (Corning 190 pH meter, Corning Scientific

Products, England) at elevated hydrolysiss and glass binding

studies) or low temperatures (glass binding studies), the pH

meter was calibrated using standard pH buffers (pH 4,7 and

10) maintained at the temperature of the studies. These

calibrating buffers were brought to the appropriate

temperature in a water bath and using the electrode the pH

was measured and calibrated using the mV control. An

"apparent" pH was recorded in mixed aqueous-organic solvents

for use in chromatography.

Ultraviolet and Visible Spectroscopy (UV-Vis).

All UV-Vis measurements were made at ambient

temperature on a Cary 219 spectrophotometer (Varian

Associates, Palo Alto, CA). The instrument was calibrated









absorbancee at different wavelengths) using a 14.2 g/l

solution of potassium chromate in 0.1 M potassium hydroxide.

Self association. In order to assess the non-linearity

of the plots of absorbance vs concentration of (I) in

solutions at various pH's, the UV-Vis spectra of (I) were

measured in different buffers (Appendix) from pH 5 to 10.

Concentrations ranging from 5 X 10-8M to 1 X 10-4M were

recorded at each pH using longer path length quartz cuvettes

(10 and 5 cm) for the dilute solutions (<5 X 10-6M) and

regular path length (1 cm) for the more concentrated

solutions. The reference compartment holder for the longer

path length cell was designed in-house.

The molar absorptivities (E) of the monomer species of

(I) at various pH's, were determined from the slopes of the

absorbance vs concentration plots obtained in the dilute

concentration range of (I) (<5 X 10-6M).

HPLC. UV-Vis spectroscopy was used to establish the

longest wavelength absorption band for compounds (I to VI)

to be analyzed by HPLC. Appropriate amounts of the analytes

were dissolved in mobile phase 'C' (see page 61 for

composition) and the resulting solutions were scanned from

650-200 nm versus a reference cell containing blank mobile

phase (Figure 2.1). The longest wavelength absorption band

with the highest molar absorptivity was used as the

excitation wavelength in the fluorescence detector.








(IV)
0.7 -


0.6 -



0.5 -



0.4-



0.3 -



0.2 -


0.1 -




400 450 500 550 600
WAVELENGTH (nm)


Figure 2.1: UV spectra for analytes (I)-(VI) dissolved in mobile phase C
(see page 61 of text).







57

Hydrolysis. UV-Vis scans were obtained from 750-200 nm,

for solutions of (I) that were degraded under various

conditions of pH, in order to detect the formation of highly

conjugated degradation products appearing red shifted

compared to (I).

Fluorescence Spectroscopy.

The fluorescence spectrophotometer was a Perkin Elmer

MPF-44 A (Hitachi Ltd., Tokyo, Japan). Typical settings for

fluorescence measurements were: excitation monochromator at

495 nm (as determined from absorption studies); emission

monochromator scanned from 500-650 nm at 120 nm/min. Chart

speed (recorder) was 60 mm/min. When necessary, the

sensitivity setting on the fluorimeter was altered during

the self-association study to record the spectra of dilute

solutions of (I). Repeat spectra on both scales were run to

estimate the conversion from one scale to the other. The

full scale setting on the chart recorder was 10 mV.

Excitation and emission bandwidths were set at 10 nm. The

instrument was calibrated absorbancee at different

wavelengths) using a 1% quinine bisulfate solution in 0.1 M

sulfuric acid. All measurements were made at ambient

temperature.

HPLC. Spectral scans of compounds (I to VI) in mobile

phase 'C' (see page 61 for composition) were obtained from

500-650 nm with excitation at 495 nm (established through UV

spectral studies). Based on the maximum emission wavelength





































525 550 575 600

WAVELENGTH (nm)


Figure 2.2:


Fluorescence spectra for analytes (I)-(VI) in mobile phase C
(see page 61 of text).


625


650







59
for all compounds (Figure 2.2) an optimum wavelength of 590

nm was selected for the HPLC fluorescence detector.

Self-association. The self-association of (I) was

studied using fluorescence spectroscopy. In order to reduce

the effects due to 'inner cell effects' (concentration

quenching) the cuvette holder in the cell compartment was

modified from the conventional right angle arrangement

(where the source of incident light and the emission light

detector are at right angles to each other) to a 'front

surface illumination' set up (where the emitted light is

detected from the same surface as the incident light).

An unused cuvette holder was detached from its base

plate and then glued back on to another plate in a manner

where the incident beam of light from the excitation

monochromator struck the front surface of the cuvette at an

angle of 30 to the normal at the surface (Figure 2.3). This

orientation gave the maximum signal response for any

solution ( e.g. 1 X 10'6 M) of (I), at a fixed setting. This

also allowed for the reflected beam (also at 30) from the

front surface to be dissipated into the back of the sample

compartment and away from the emission monochromator,

avoiding interference with the emission signal (Figure 2.3).

This orientation was used throughout the self-

association study and permitted the measurement of dilute (1

X 10-M) as well as high concentrations of (I) (1 X 10-'M)

without a significant loss of signal due to inner cell

















EXCITATION MONOCHROMATOR








REFLECTED BEAM
.II
I I
Ie I




/ \-,* \
1 -< -~ FRONT SURFACE
R E B \ ILLUMINATION


i I RIGHT ANGLED
ILLUMINATION


Figure 2.3:


Front surface and right angled illumination
setup. Shaded area represents incomplete
penetration by excitation light.









effects.

Chromatography.

A High Performance Liquid Chromatography (HPLC) method

using a fluorescence detection system was developed that was

specific for (I), its possible degradation products (II to

V) and the internal standard (VI).

Mobile Phase(s). Three different mobile phases were

used during chromatographic studies. Mobile phases (A) and

(B) were used for purity assessment and mobile phase (C) was

used for routine analysis of compounds (I to VI). Mobile

phase (A) consisted of 70 volumes acetate buffer (0.05 M, pH

3.98) + 30 volumes acetonitrile containing 5% v/v

tetrahydrofuran (THF). Mobile phase (B) consisted of 70

volumes of phosphate buffer (0.05 M, pH 3.86 adjusted with

orthophosphoric acid) + 30 volumes of acetonitrile

containing 3% v/v THF. Mobile phase (C) consisted of 70

volumes of 0.05 M phosphate buffer (7.8 g/l of sodium

phosphate monobasic dihydrate in water), 30 volumes of

acetonitrile and 0.4 ml/l of tetrabutyl ammonium hydroxide

(1 M solution in water). The "apparent" pH of the resulting

solution was adjusted to 3.55 with orthophosphoric acid. All

mobile phases were filtered through a 0.2 gm nylon membrane

filter (Rainin) under vacuum using a Millipore filtration

unit (Millipore Corp., Milford, MA). Each mobile phase was

recycled (except during evaluation of the reproducibility of

the calibration curve) and replaced if chromatographic









properties of the analytes changed or after about 2-3 days.

Fresh mobile phase was equilibrated with the column

overnight. Between mobile phases the HPLC system was rinsed

overnight with a "wash solvent" system consisting of

water:methanol:isopropanol in the ratio 2:1:1. This washed

out potentially damaging buffer and analyte residues in the

HPLC system, thereby increasing system lifespan.

Solvent Delivery. The mobile phase was isocratically

delivered initially by a Waters' pump (Model M-6000 A,

Waters Associates, Millipore Corp., Milford, MA) and later

by a LDC/Milton Roy pump (Model CM 4000, LDC/Milton Roy,

Riviera Beach, FL). No significant difference was observed

in the retention times or capacity factors of the various

analytes on replacing the pumps. Flow rate for mobile phase

(A) and (B) was optimized at 1.0 ml/min and for (C) at 1.2

ml/min.

Column. The stationary phase consisted of a Whatman

Partisil (Whatman Inc., Clifton, NJ) ODS (C-18) column (25

cm X 0.45 cm i.d.; 5 pm particle size) for mobile phases (A)

and (B). The column for mobile phase (C) was a Whatman

Partisil (Whatman Inc., Clifton, NJ) ODS (C-18) (10 cm X

0.45 cm i.d.; 5 pm particle size). Both analytical columns

were protected by a 1 cm guard column packed in-house with a

Permaphase ODS 10 pm material (Dupont Instruments,

Wilmington, DE) using a traditional tap and fill technique.







63

Injector. All injections were made with a 50 pl syringe

(Hamilton Co., Reno, NV) through a Rheodyne 7125 high

pressure injector (Rheodyne Corp., Cotati, CA) fitted with a

20 p.1 loop.

Detector. Analytes eluting from the column were

detected by a Perkin Elmer 650S fluorescence detector

(Hitachi Ltd., Tokyo, Japan) with excitation and emission

wavelengths set at 470 nm and 553 nm, respectively for

mobile phases (A) and (B). The excitation and emission

wavelengths were optimized to 495 nm and 590 nm,

respectively, in the case of mobile phase (C). Excitation

and emission bandwidths were set at 10 nm. Power to the lamp

was supplied by a Perkin Elmer 150 B Xenon Power supply unit

(Hitachi Ltd., Tokyo, Japan).

Recorder/Integrator. Data was recorded on a HP 3394 A

integrator (Hewlett Packard Cc-p., San Fernando, CA) and a

Fisher recordall series 5000 strip chart recorder (Bausch

and Lomb, Houston Instrument division, Austin, TX). Peak

heights recorded on the strip chart recorder were measured

for data analysis. The integrator data was maintained as a

means of peak identification (printed retention times),

column performance (Area/Peak height ratios) and for

estimation of the peak heights that were "off" scale on the

strip chart recorder (using peak area ratios).









Temperature-controlled studies.

In studies involving glass binding and hydrolysis, the

temperature was rigorously controlled ( 0.1C) throughout

the study.

Water bath(s). Conditions designed to mimic those in a

refrigerator were maintained by a Lauda K-2/R refrigeration

unit (Brinkmann Instruments, Germany). Water at 5.0 + 0.1C

was delivered by gravity assist to a reaction bath and was

recycled back to the refrigeration unit by a pump. Ambient

and high temperature settings ( 0.1C) were maintained by a

Haake water bath and alternately an Isotemp Immersion

Circulator (Model 70; Fisher Scientific, Pittsburgh, PA).


Methods


Silylating Glassware.

All glassware in contact with (I) was pre-silylated.

Each item was thoroughly washed with detergent, rinsed with

water and methanol and dried in a hot air oven at 160C for

30 minutes. The dried glassware was allowed to cool and

about 8-20 ml of hexamethyldisilazane (silylating agent) in

hexane was added. The liquid was swirled in the glassware so

that all exposed surfaces of the glass would be coated. The

silylating liquid was used sequentially for up to four

times. Silylated glassware were allowed to drain and dry in

air, rinsed with methanol and dried in a hot air oven for 2-

3 hours. This glassware was used for handling of (I)









solutions. Care was taken to avoid scratching the inside

surface of the treated glassware in order that the

unprotected surface would not be reexposed.

Self-Association.

Stock solutions. All stock solutions of (I) (10-s, 10-4,

10-3 M) used for the self association study were made in

deionized, distilled, deoxygenated (with nitrogen) and

filtered water. These solutions were protected from light

and stored in silylated glass volumetric flasks in the

refrigerator at 5C without freezing, and allowed to warm to

ambient temperature before use. All stock solutions were

discarded after one week.

Sample preparation. In order to minimize losses that

might occur due to hydrolysis, precalculated aliquots of (I)

from stock solutions were pipetted directly into 1 cm square

quartz cuvettes (fluorescence studies) and 5 or 10 cm

cylindrical quartz cuvettes absorbancee measurements).

Immediately prior to spectral recording, sufficient volume

of the appropriate buffer was pipetted into the cuvettes to

yield the desired final concentrations of (I). Between

successive samples the cuvette was thoroughly rinsed with

acetone and distilled water.

Normalization and Selective Excitation Measurements. In

order to determine the presence of a longer wavelength

emission band (if the dimer fluoresces) than a 'monomer

spectrum', the latter was recorded at a concentration of 5 X









10'7 M (it was assumed that dimerization was negligible at

this concentration, since concentrations of dimer formed

will be < 10% of the total concentration of (I) used if the

association constant is assumed to be 4 X 10' M-1 the

largest value cited in the literature (202)). This spectrum

had two maxima, at 560 and 590 nm (cf. Figure 3.5). Each of

these maxima was selected as a 'Normalizing Wavelength'. The

concentration of (I) solution was progressively increased in

the cuvette and at each increase the emission spectra of (I)

was scanned till 560 nm (or 590 nm) when the recorder pen

would show a higher fluorescence intensity than the 'monomer

spectrum' for the 5 X 10'7 M solution. Using the 'coarse'

and 'fine' sensitivity selectors on the spectrophotometer,

the emission signal at 560 nm (or 590 nm) was attenuated

such that the recorder pen overlay the spectral recording

for the 'monomer spectrum'. Using this 'new' instrument

setting the spectrum of the concentrated solution was

recorded and overlayed on the 'monomer spectrum'. If the

dimer, fluoresces, the overlayed spectra at higher

concentrations would be shifted towards longer wavelengths

than the 'monomer spectrum' (cf. Chapter 3 for explanation).

Using the technique of El-Sayed et al. (285), and the

results published by Arcamone et al. (172), the excitation

monochromator was set at 580 nm (a wavelength at which the

dimer spectra does not overlap the monomer spectra) and the

emission scanned from 580 to 710 nm. Since the absorbance of
















































WAVELENGTH


Figure 2.4:


Typical fluorescence spectra of (I) in
aqueous solution, a: 1 X 10-4M; b: 1 X 10-6M.


100 -r


50 -+-


25 -+


500


550


600


Xnm


-l-
650







68

the dimer at 580 nm is low, the sensitivity settings of the

instrument were maximized to allow for the weakest signal to

be detected.

Sample Spectral Measurements. Using the modified cell

holder, spectral recordings of (I) from 1 X 10-7 to 1 X 10-4

M were obtained at three different pH's (5.72, 6.72, 7.69)

with each maintained at three different ionic strengths

(0.01, 0.05, 0.1 M). For estimating the ratios of the

quantum efficiencies of the monocation (MC-DOX) and neutral

(N-DOX) species of (I) (ON-DOX/MC-DOX) fluorescence

measurements were recorded for 1 X 10-6 and 1 X 10-4M

solutions, each maintained at pH's of 5 and 10 to bracket

the pH range of the first ionization (pKI) constant of (I)

in solution (cf. Chapter 3: Results and Discussion for

rationale). To determine the molar absorptivity of the

monomer species of (I), spectral scans of (I) solutions in

concentration ranges from 8 X 10' to 5 X 10-M were obtained

in buffers of pH's 6, 7, 8 and 9 at ionic strengths of 0.01,

0.05, 0.1 and 0.2 M each. Since the molar absorptivity of

the monomeric monocation (MC-DOX) and neutral (N-DOX)

species of (I) were assumed to be identical (cf. Chapter 3:

Results and Discussion), they were averaged across the pH

range studied. Compositions of buffers for the above studies

are detailed in the Appendix.









Data Treatment. Typical fluorescence spectra are

shown in Figure 2.4. Peak heights were measured at 590 nm in

terms of relative fluorescence intensity (relative to a

blank) and plotted versus concentration. The absorbance at

495 nm was recorded versus the concentration and used to

calculate the molar absorptivity. The fluorescence data was

used to develop an equilibrium model to explain self-

association. The model development is outlined in Chapter 3:

Results and Discussion.

High Performance Liqcruid Chromatoqraphy.

Samples used to generate calibration curves for HPLC

studies were prepared in mobile phase (C) for all analytes.

Appropriate volumes from stock solutions (= 10 .g/ml) of the

various analytes (I-V) were placed in silylated 3 ml glass

tubes. An appropriate volume of (VI) as the internal

standard was added to yield a final concentration of 500

ng/ml. Calibration curves were generated in the

concentration range from 25-700 ng/ml for all analytes. To

establish the ruggedness and the reproducibility of the

chromatographic system, calibration curves were generated on

various days using the same and/or different mobile phase.

Additional calibration curves were generated prior to and

after each study to ensure consistency of the

chromatographic system. In cases where a study proceeded

over a period of days, calibration samples were run at the

beginning of each day and also during the study. The slopes









and intercepts obtained from the calibration curves

generated for a study were averaged over the number of days

of the study and used for the estimation of concentration.

Assay validation. Validation was performed by first

measuring peak heights of the HPLC calibration samples using

a standard ruler (with 1 mm rulings). The peak height ratio

was calculated by dividing the measured peak height of the

analyte by that of the internal standard. Plots of the peak

height ratio vs the formal concentration of the analytes

gave rise to calibration curves. Simple linear regression of

the calibration curves gave the slope, intercept,

correlation coefficient, and the standard errors of the

slope and regression. Analysis of Variance (ANOVA) of the

slope and the intercept between days and with different

mobile phases gave interday variability. Based on the

signal-to-noise ratio (S/N), and the regression parameters

of the calibration curves, information on Limits of

Detection (LOD), Sensitivity and Minimum Quantifiable Limit

(MQL) were obtained. The results and interpretation of these

analysis are in Chapter 3: Results and Discussion.


Glass Binding.

The non-specific binding of (I) to glass was studied as

a function of concentration, pH and temperature.

Effect of Temperature. Glass binding was studied at

three different temperatures: 5C, 25C and 50C. The

temperatures were selected to mimic the binding of (I) to









glass in the refrigerator, at ambient temperature (room),

and at elevated temperature (e.g., in a hydrolysis

experiment), respectively. All buffer solutions used were

pre-equilibrated at the appropriate temperature before start

of the study.

Effect of pH. The influence of pH and the ionization

of (I) was achieved using buffers at pH 3.18 (acetate 0.05

M), 5.01 (acetate 0.05 M) and 7.00 (phosphate 0.05 M) (see

Appendix). Each buffer was thoroughly deoxygenated with

nitrogen and equilibrated at the temperature of the study.

Effect of Concentration. The percent loss of (I) due to

glass binding from dilute solutions was expected to be

greater than from concentrated ones. To evaluate the extent

of binding as a function of concentration 1, 10 and 100

pg/ml solutions of (I) were studied at each of the pH's and

temperatures mentioned above. Increasing concentrations

should lead to saturation of the binding sites that would

show as a negligible loss due to glass binding.

Experimental. Corning 50 ml (15 ml for 100 gJg/ml

solutions) test tubes with ground glass stoppers and a

tapered bottom were used for the study. Appropriate aliquots

of (I) from stock solutions (2 mg/ml) were pipetted into

temperature equilibrated deoxygenated buffer solutions to

yield 30 mL (10 mL for 100 gg/mL) of the appropriate

concentrations of (I). This volume allowed for a limited

head space to be maintained above the liquid level and thus








72

reduced the diffusion of air into (and subsequent oxidation

of) the sample. To further prevent oxidation the head space

was constantly flushed with nitrogen before, during, and

after every sample was taken. The nitrogen was humidified at

the temperature of the study. Each study was conducted in

silylated and unsilylated tubes as a means of comparison and

to study the effectiveness of the silylation process. Due to

the large surface area of contact/unit volume of the tubes,

there was negligible change (<5%) in the contact area at the

end of a study (after multiple sampling). All tubes were

wrapped with aluminum foil to protect samples from

photolytic oxidation. Each study was run in triplicate.

Sampling. Sample volumes for analysis depended on the

concentration of (I) used in the study. In a clean

previously silylated glass tube 200, 50 and 10 gl samples

were withdrawn from the 1, 10 and 100 gg/ml reaction

solutions respectively and plunged into an ice bath. To each

solution an appropriate volume of (VI) was added as an

internal standard to yield a final concentration of 500

ng/ml. The mobile phase was used to make up volumes to 1 ml

for the 1 and 10 ug/ml solutions and 2 ml for the 100 gg/ml

solution. Solutions were sometimes stored on ice for a

period not exceeding 24 hours prior to analysis. Twenty

microliters of each solution was injected into the HPLC for

analysis.








73

Data Treatment. The Peak Height Ratio (PHR) of (I)

was calculated. Using the most recent, (generated before the

study) calibration curve the PHR were converted to the

corresponding concentrations. The percent of the initial

concentration remaining and the amount of (I) bound to the

glass were determined as a function of time. The kinetic

data curves generated were fitted to a sum of exponentials

which in turn yielded important binding parameters.

Chromatograms for studies undertaken at 50C showed the

appearance of the aglycone (IV), making it difficult to

quantitate the extent of binding at this temperature.




Chemical Kinetics.

The kinetics of hydrolysis of (I) was studied under

various conditions of pH, temperature and buffer

concentrations in order to establish log kobs-PH profiles, pH

of maximum stability, the Arrhenius activation energy (Ea)

and frequency factor (A), the existence of general and

specific acid-base catalysis, the possible mechanisms,

pathways and structures of degradation products.

Effect of acids. In order to evaluate the effect of the

conjugate bases of different mineral acids, hydrochloric,

perchloric and sulfuric acids were used to study the

hydrolysis of (I) at 50C. Further, to compare the activity

of the hydrogen ion obtained from these acids, their

concentrations (0.01, 0.05, 0.1, 0.2 M) and ionic strengths










(0.2 M with KC1) were kept constant. The activity of the

hydrogen ion (as determined by its activity coefficient)

varied as a function of the temperature, ionic strength and

the conjugate species present in solution (289). In the

present case measuring the pH at the temperature of the

study (i.e., 50C) for concentrated acid solutions was not a

reliable estimate of the activity of the hydrogen ions in

solution, since significant errors develop in glass

electrodes under these conditions. Estimates of the activity

coefficients of the hydrogen ions under the influence of the

above mentioned combination of conditions were obtained from

Harned and Owens (289). In cases where the literature did

not provide the relevant information, estimates of the

activities of the hydrogen ions were extrapolated from the

available data using the extended Debye Huckel equation

(257). These extrapolations and the rationale for their use

are presented in detail in Chapter 3: Results and

Discussions.

Effect of concentration. The concentration dependence

of the hydrolysis, was studied by evaluating the reaction

rate of 10 and 100 gg/ml solutions of (I) in 0.2 M HC1 at

70C.

Effect of Temperature. To obtain Arrhenius parameters

and to predict rate constants of degradation at ambient

storage conditions (room temperature or refrigerated

temperature), accelerated studies were carried out at four










elevated temperatures in each of the buffer systems. Four

different concentrations of each buffer system at a given

temperature and pH were studied. The temperature studies

were carried out at 70C, 65C, 60C and 55C for all buffers

and 50C for HC1 studies instead of 55C.

Effect of pH. The hydrolysis study focused

predominantly on the acid-catalyzed hydrolysis of (I) in

solution. The pH from 0.70 to 2.90 was generated by using

different concentrations of HC1. The pH range from 3.00 to

8.00 was maintained by various buffers starting with format

buffer for pH 3.00 and 4.00; acetate for pH 4.50, 5.00 and

5.50 and phosphate for pH 6.00, 7.00 and 8.00 (see

Appendix). Each pH study was conducted at the temperatures

mentioned above. The pH of each solution was measured at the

temperature of the study using an electrode that had been

calibrated using standard buffers at the same temperature.

Effect of buffers. The general acid/base catalysis of

(I) was studied using each buffer at four different

concentrations, taking care to maintain the pH and the ionic

strength constant. The format and acetate buffers were made

at concentrations of 0.01, 0.05, 0.1 and 0.2 M at a constant

ionic strength of 0.2 M (with KCl). In case of the phosphate

buffers the concentrations were reduced to 0.01, 0.05, 0.1

and 0.12 M for pH 6.00; 0.01, 0.025, 0.05 and 0.075 M for pH

7.00 and 0.01, 0.025, 0.05 and 0.065 M for pH 8.00 (see

Appendix). This compensated for the presence of the










monobasic and dibasic phosphate ion, each of which

contributes to the ionic strength. Each buffer solution was

deoxygenated with nitrogen before the pH was adjusted.

Experimental. A fixed volume of each of the buffers

was equilibrated to the desired temperature in 50 ml

silylated test tubes protected from light. Each buffer was

thoroughly deoxygenated with nitrogen. An appropriate

aliquot of a stock solution of (I) was added to each tube

(time, t=0) to yield a final concentration of 10 pg/ml (100

gg/ml for concentration dependence studies). At various time

intervals during the study (selected on the basis of the

predicted or previously reported half-lives at that

temperature, pH and buffer concentration) 50 gl aliquots of

the test sample were transferred to previously silylated

test tubes and cooled in an ice bucket (to arrest any

further reaction). Appropriate amounts of (VI) were added as

an internal standard to yield a final concentration of 500

ng/ml. The remaining volume was made up to 1 ml with the

mobile phase. Twenty microliters of this solution was

injected into the HPLC. Some solutions (that could not be

analyzed on the day of sampling) were stored in an ice

bucket without freezing for no longer than 24 hours prior to

analysis without any detrimental effects. Glass binding did

not interfere with this phase, probably due to the presence

of the organic component acetonitrilee) in the mobile phase.








77

Data Treatment. As described for the glass binding

studies, the peak heights of (I) and its degradation

products, as well as the internal standard (VI), were

measured and the corresponding peak height ratios were

calculated. Using the most recently generated calibration

curves or the mean parameters of a series of calibration

curves generated prior to, during and at the end of each

study, the peak height ratios were converted to

concentrations. These data were converted to percent of (I)

remaining at the corresponding times and fitted to a

monoexponential decay equation (in the ln C vs t form). The

resulting slope gave the overall pseudo first order observed

rate constant kobs for hydrolysis. This rate constant was

kinetically processed to yield Arrhenius parameters (plots

of ln kobs vs 1/T), second order rate constants (when plotted

versus [H']) for the various catalytic species ([H], [OH-]

and [H20]) and second order rate constants (when plotted

versus buffer concentration) for the general acid base

catalysis by the various species that constituted the

buffers. The respective kinetic treatments are presented in

detail in Chapter 3: Results and Discussion.

Identification of degradation products.

During buffer catalyzed hydrolysis of (I), sample

chromatograms showed various peaks which did not match

(based on retention times) the chromatogram of the standards

(I-VI). Therefore an attempt was made to isolate and, if










possible, to elucidate the structure of these degradation

products. In addition, all solutions undergoing hydrolysis

(from pH 0.79 to 8) showed the appearance of a precipitate.

This precipitate was also collected and an attempt was made

to analyze it.

Analysis of the solution and precipitate (obtained

after hydrolysis) in the pH range from 0.8 to 3. The

hydrolysis study in the pH range from about 0.8 to 3 showed

the appearance of one additional peak on the chromatograms,

having the same retention time as the aglycone (IV). The

aglycone was quantitatively generated from (I) in this range

as shown in Chapter 3: Results and Discussion. The

precipitate was collected, washed with cold water and then

dissolved in the mobile phase. HPLC analysis of the

resultant solution showed the presence of only one peak

having the same retention time as the aglycone (IV). Thus,

the aglycone was the only degradation product formed in this

pH range.

Analysis of the solution and precipitate (obtained

after hydrolysis) in the pH range from 4 to 8. In all

solutions at pH>3, the formation of the aglycone was

negligible. Instead unknown peaks were observed (see Chapter

3 for details), that did not fit the chromatograms of the

known analytes (I-VI). At the end of a study the test

solution acquired a deep pink color (instead of the orange

red color of (I)) indicating the formation of a product that









appeared to be more conjugated than (I).

The solutions and precipitates obtained at the end of

the hydrolysis were subjected to UV-Vis spectroscopy as

described in 'Ultra violet and Visible Spectroscopy:

Hydrolysis'. The results are discussed in Chapter 3: Results

and Discussion.

Solutions of the precipitate prepared in mobile phase

and injected into the HPLC showed the same chromatograms as

the solutions themselves. Isolation of the various peaks by

manual fraction collection did not yield sufficient

quantities for characterization and could not be separated

from the buffer components of the mobile phase.

Effect of oxygen on the hydrolysis. In order to study

the effect of oxygen on the hydrolysis of (I), three 50 ml

silylated tubes A, B and C each protected from light were

filled with approximately 10 ml of deionized, distilled,

boiled and filtered (2pm filter) water. Each tube was

preequilibrated to a temperature of 70C in a water bath.

Nitrogen was continuously bubbled through the water in tubes

A and B, and tube C was equilibrated with oxygen. Just prior

to the start of the study, 8 mg/ml stock solution of (I) was

prepared in deionized, distilled and deoxygenated water. An

aliquot of the stock solution was pipetted into each of the

equilibrated tubes to yield a final concentration of 2 mg/ml

(concentration of the reconstituted formulation of (I) used

in clinical practice), and the tubes were stoppered. The pH








80

of the solutions were measured throughout the study using an

electrode calibrated at 70C. The appropriate gas was

bubbled through each tube for 1 minute to displace any air

that might have entered the tubes during pH measurements. At

a predetermined time (based on a similarly run pilot study)

the nitrogen gas that was bubbled through tube B was

replaced with oxygen until the conclusion of the experiment

(no change in pH).

Due to a drop in the pH noted (from about 4.8 to 2.9)

in the tube equilibrated with oxygen, the formation of an

acidic product was thought to have formed. Kinetic treatment

of the resulting pH vs time data suggested the formation of

a product having a pKa around 3 (see Chapter 3: Results and

Discussion for kinetic development). Since this pH drop was

not observed in solutions equilibrated with nitrogen, the

formation of the acidic product appeared to be due to an

oxidation process. Beijnen et al. (237) had suggested that

the side chain of (I) could break apart and tautomerize to

form glycolaldehyde (hydroxy-acetaldehyde). The logical

explanation to the formation of an acidic product was the

oxidation of the aldehyde to the corresponding hydroxy

acetic acid (X-COOH), which could be responsible for

lowering the pH of the oxygenated solution.

Detection of product X-COOH by potentiometry. An

experiment similar in design to the one above was set up to

generate the acidic product and titrate it with dilute NaOH








81

to determine its pKa. Consequently about 20 ml of a 2 mg/ml

solution of (I) was prepared in deionized, distilled and

filtered water which had been equilibrated at 70C. The

solution was allowed to react with oxygen bubbled

continuously through the solution. A pH electrode calibrated

at 70C was placed in the solution to record the pH of the

solution. At the end of the experiment (no further change in

pH took place), the flow of oxygen was stopped and a

magnetic stirrer bar was dropped into the solution. Aliquots

of 10, 50 and 100 gil of a 0.01 M or 0.1 M solution of NaOH

were added to the solution at 70C with continuous stirring.

The pH after each addition was recorded until no further

change in pH could be noticed (pH>12). A blank solution was

titrated in a similar manner and the difference between the

sample solution and the blank was plotted using the Parke-

Davis method (290) as a function of the measured pH.

Detection of product X-COOH by ion-pair extraction and

HPLC. Aliquots of the hydrolyzedd' solutions from each tube

(A, B and C) were adjusted to a pH between 6 and 7 with

NaOH. This aqueous mixture was saturated with NaCI and

extracted repeatedly with small volumes of

chloroform:isopropanol (4:1 v/v) until colorless. The NaCI

helped to salt out the colored component into the organic

phase. To the remaining aqueous phase, 0.5 ml

tetrabutylammonium hydroxide (1 M) was added as an ion pair

agent. The resulting solution was extracted three times










using small volumes of diethyl ether. The ether extracts

were pooled and evaporated to dryness. The residue obtained

was dried under vacuum.

A HPLC system was set up consisting of a 25 cm reverse

phase C-18 column. The mobile phase was a 85:15% v/v of 0.05

M tetrabutylammonium hydroxide in water: acetonitrile.

Detection was by UV with a X. set at 210 nm. Flow rate was

1.0 ml/min. The residue obtained from the ether extracts was

reconstituted in the mobile phase and 20 gl of the sample

was injected into the HPLC. Standard solutions of known two

carbon carboxylic acids glycollicc, glyoxylic, oxalic etc.)

were prepared and injected into the HPLC as a means of

comparison.

Detection of product X-COOH by derivatization of the

aldehyde. Since the acidic product X-COOH was theorized to

have formed from the oxidation of an aldehyde, another

method of confirming the existence of X-COOH would be to

detect its precursor. Consequently a 2 mg/ml solution of (I)

in water was heated to 70C. Nitrogen gas was continuously

bubbled through this solution and the gas outflow was led

into a methanolic solution of 2,4-dinitrophenylhydrazine.

The appearance of an 'osazone' precipitate would confirm the

presence of an aldehyde.

Precipitate analysis by LC-MS: The precipitate

generated in the oxygen study showed the same

chromatographic profile as that obtained in the hydrolysis









study. This precipitate was collected, washed with cold

water and dried in a vacuum desiccator. A Kratos LC-MS with

a thermospray attachment (vaporizes the LC column effluent

prior to mass spectral analysis) was set up. In order to

maximize the efficiency of the thermospray assembly, the

mobile phase of the LC was devoid of non-volatile buffers

like phosphate, carbonate, etc. In addition sodium and

potassium salts and organic modifiers like

tetrabutylammonium hydroxide that can create occlusion of

the exit pore of the thermospray assembly, were avoided. A

fresh mobile phase was prepared consisting of a 70% solution

of 0.05 M ammonium acetate (pH 3.5, adjusted with glacial

acetic acid) and 30% acetonitrile. This mobile phase also

separated and resolved all the degradation peaks as mobile

phase C (see 'Methods: Mobile Phase(s)'). The only

difference in the chromatograms was that the aglycone (IV)

peak appeared later than the internal standard (VI). The

precipitate obtained was dissolved in 1 ml of the mobile

phase and about 100 gl of this solution was injected into

the LC-MS. The high concentration of the acetate buffer

interfered with the analysis and no satisfactory results

could be obtained.







CHAPTER 3
RESULTS AND DISCUSSION


Purity Determination


Melting Point.

The melting point measured for each of the lots of (I)

hydrochloride received, showed similar characteristics. In

all cases (I) melted at 205C with decomposition (sample

decolorized from a reddish orange to a charred ashy

residue). If the samples melted and discolored at a

temperature of 205C, as mentioned in the literature (141),

they were considered pure. All the samples received passed

the melting point test.

Differential Scanning Calorimetry (DSC).

A typical DSC endotherm of the melting of (I) is shown

in Figure 3.la. The sample was rapidly heated to its melting

point, the endotherm recorded and the purity estimated using

a derived form of the van't Hoff equation (3.1).

RT^x, i
T = T T02 .1 (3.1)




where Ts (K) is the sample temperature, To (K) is the

melting point of a 100% pure sample, R is the gas constant

(8.314 J/mole.K), AHf0 is the molar heat of fusion (J/mole),

X2 is the mole fraction of the impurity and F is the
























F, 10












0 a a

190 200 210 220 230

TEMPERATURE C

Figure 3.1 a: A typical DSC endotherm of the melting of (I) for purity evaluation.