Citation
Effects of ionic perturbations and metal ion competition on the binding of a model drug to DNA

Material Information

Title:
Effects of ionic perturbations and metal ion competition on the binding of a model drug to DNA
Creator:
Eisenhardt, Peter Forrest, 1945-
Publication Date:
Language:
English
Physical Description:
xvi, 301 leaves : ill. ; 29 cm.

Subjects

Subjects / Keywords:
Cations ( jstor )
DNA ( jstor )
Electrolytes ( jstor )
Ions ( jstor )
Molecules ( jstor )
pH ( jstor )
Phosphates ( jstor )
Solvents ( jstor )
Teeth ( jstor )
Titration ( jstor )
Acridines ( mesh )
Binding, Competitive ( mesh )
DNA ( mesh )
Dissertations, Academic -- Medicinal Chemistry -- UF ( mesh )
Ions ( mesh )
Macromolecular Systems ( mesh )
Medicinal Chemistry thesis Ph.D ( mesh )
Osmolar Concentration ( mesh )
Genre:
bibliography ( marcgt )
non-fiction ( marcgt )

Notes

Thesis:
Thesis (Ph.D.)--University of Florida, 1977.
Bibliography:
Includes bibliographical references (leaves 295-300).
General Note:
Typescript.
General Note:
Vita.
Statement of Responsibility:
by Peter F. Eisenhardt.

Record Information

Source Institution:
University of Florida
Holding Location:
University of Florida
Rights Management:
Copyright [name of dissertation author]. Permission granted to the University of Florida to digitize, archive and distribute this item for non-profit research and educational purposes. Any reuse of this item in excess of fair use or other copyright exemptions requires permission of the copyright holder.
Resource Identifier:
000898605 ( ALEPH )
25764483 ( OCLC )
AEK7311 ( NOTIS )

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Full Text














EFFECTS OF IONIC PERTURBATIONS AND METAL ION
COMPETITION ON THE BINDING OF A MODEL DRUG TO DNA












By

PETER F. EISENHARDT


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




UNIVERSITY OF FLORIDA


1977




EFFECTS OF IONIC PERTURBATIONS AND METAL ION
COMPETITION ON THE BINDING OF A MODEL DRUG TO DNA
By
PETER F. EISENHARDT
A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1977


DEDICATION
To Dearest Joannie
who, more than any other,
deserves credit for any positive
attributes I may possess.


ACKNOWLEDGEMENTS
I am indebted to my supervisory committee, Dr. Stephen
G. Schulman, Chairman; Dr. B. Andresen; Dr. M. Battiste;
Dr. C.H. Becker; and Dr. R.H. Hammer for their time and
guidance in the preparation of this manuscript.
To three Professors, I owe special thanks. To Dr.
Schulman for the many hours of his time he has given to
share his prodigious scientific knowledge. To Dr. "C" whose
unaffected concern for the well-being of all with whom he is
associated, combined with his keen intellect, makes him with
out peer. And, to Dr. John Baxter for his fine example and
unflagging attempts to maintain high standards within the
academic community.
Space permits mentioning only a few of my friends to
whom I'm indebted: Mum, Dad, Clam, Burns, Sue, Flyman and
Ms. Fly, Teem, Merm, Tony, Roy, Lenny-whats-is-name, Edgar,
and Lysa Chancey Smith.
Robert, Ann, Robbie, and Sarah are four of my life's
mainstays.
Thanks, too, to Carolyn Grantham, expert cryptographer,
who deciphered my cruel first drafts so competently without
once betraying her graceful, elegant mien.
iii


TABLE OF CONTENTS
Page
ACKNOWLEDGEMENTS i;Li
LIST OF FIGURES vi
LIST OF TABLES x
ABSTRACT X1V
CHAPTER I INTRODUCTION 1
Binding of Small Molecules to Biopolymers 2
Thermodynamics of Binding Reactions 34
Rationale for the Selection of the Model
System and Experimental Design 42
CHAPTER II EXPERIMENTAL 47
Materials and Procedures 47
Calculations 55
CHAPTER III RESULTS AND DISCUSSION 62
General Spectral and Titration Characteristics.. 62
Apparent Association Constants for Surface
and Intercalative Binding 84
Association Constants Based on Activities --
Extended Debye-Hckel Considerations 97
Evaluation of a Simple Competitive Binding
Model 102
Thermodynamics of the Binding of 3-amino-
acridinium and 7-aminoquinolinium to DNA 105
APPENDICES
APPENDIX I FIGURES 120
IV


TABLE OF CONTENTS (Continued)
Page
APPENDIX II TABLES 187
APPENDIX III COMPUTER PROGRAMS 281
REFERENCES 29 5
BIOGRAPHICAL SKETCH 301
V


LIST OF FIGURES
Figure
Page
1
2
3
4
5
6
7
8
The structures of selected substituted
acridines and related compounds (Structures
I through XVI) 122
General form of Scatchard plots: R/C vs.
R for one and two classes of binding sites 128
Absorptiometric titration of 3-amino-
acridinium cation with DNA, cesium salt,
pH 5.90, 25.0C. Background electrolyte:
CsH2P04' Htial concentration, 0.15M 130
Absorptiometric titration of 3-amino-
acridinium cation with DNA, cesium salt,
pH 5.90, 25.0C. Background electrolyte:
CsH2PO^, initial concentration, 0.0025M 132
Absorptiometric titration of 3-amino-
acridinium cation with DNA, magnesium salt,
pH 5.90, 25.0C. Background electrolyte:
Mg(09CCH3)2, initial concentration, 0.010N 134
Absorptiometric titration of 3-amino-
acridinium cation with DNA, magnesium salt,
pH 5.90, 25.0C. Background electrolyte:
MgiC^CCH^^' initial concentration, 0.00063N 136
Log([BHP]/[BH]) vs. log([Pt/3l-[BHP]) with
0.0025M CsH2P04 and 6.3 x 10-4N Mg(02CCH3)2,
as supporting electrolytes, pH 5.90, 25.0C 138
Log ( [BHP] / [BH] ) vs. log ( [Pt/31 -.[BHP] ) with
0.15M CsH2P04, 0.010M CSH2P04, and 0.010N
Mg(09CCH9)2 as supporting electrolytes, pH
5.907 25.0C 140
9 Log([BHP]/[BH]) vs. log([Pt/21-2[BHP]) with
0.0025M CsH2P04 and 6.3 x 10~4N Mg(C>2CCH3)2
as supporting electrolytes, pH 5.90, 25.0C 142
vi


LIST OF FIGURES (Continued)
Figure
10
11
12
13
14
15
16
Log([BHP]/[BH]) vs. log([Pt/2^2[BHpl) with
0.15M CsH2P04, 0.01M CsH2P04, and 0.010N
Mg(02CCH3)2 as supporting electrolytes, pH
5.90, 25.06C
Log of the apparent association constant,
Ks, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, I1/2. NaH2P04
and LH2PO4 as supporting electrolytes, pH
5.90, 25.0C
Log of the apparent association constant,
K
s'
for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, I-1-/2. KH2PO4
and KO2CCH3 as supporting electrolytes,
pH 5.90, 25.0C
Log of the apparent association constant,
K
s'
for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, 1^/2 KH2P04
and RbH2P04 as supporting electrolytes,
pH 5.90, 25.0C
Log of the apparent association constant,
Ks, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, ll/2 CsH2P04
and (CH3)4NH2PC>4 as supporting electrolytes,
pH 5.90, 25.0C
Log of the apparent association constant,
Ks, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, li/2. Mg(02CCH3)2
and Ca(02CCH3)2 as supporting electrolytes,
pH 5.90, 25.0C
Log of the apparent association constant,
Kg, for the surface binding of 3-amino-
acridiniura cation to DNA vs. the square
root of the ionic strength, I1/2. Sr(02CCH3)2
and Ba(02CCH3)2 as supporting electrolytes,
pH 5.90, 25.0C
Page
144
146
148
150
152
154
156
Vll


LIST OF FIGURES (Continued)
Figure
17
18
19
20
21
22
23
Log of the apparent association constant,
Ki, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I*-'.
LH2PO4 and NaH2P04 as supporting
electrolytes, pH 5.90, 25.0C
Log of the apparent association constant,
Kj, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I1/2.
KH2PO4 and KO2CCH3 as supporting
electrolytes, pH 5.90, 25.0C
Log of the apparent association constant,
Kj, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I1/2.
KH2PO4 and RbH2PC>4 as supporting
electrolytes, pH 5.90, 25.0C
Log of the apparent association constant,
Kj, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I1/2.
CSH2PO4 and (CH3)4NH2PO4 as supporting
electrolytes, pH 5.90, 25.0C
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, ll/2.
LH2PO4, NaH2PC>4, and KH2PO4 as supporting
electrolytes, pH 5.90, 25.0OC
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, I1/2.
RbH2P04, CsH2P04, and (CH3)4NH2P04 as
supporting electrolytes, pH 5.90, 25.0C
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, ll/2.
Mg(02CCH3)2 and Ca(02CCH3)2 as supporting
electrolytes, pH 5.90, 25.0C
Page
158
160
162
164
166
168
170
viii


LIST OF FIGURES (Continued)
Figure
24
25
26
27
28
29
30
31
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, I^-/2.
Sr(O2CCH3)2 and Ba(02CCH3)2 as supporting
electrolytes, pH 5.90, 25.0C
Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium
salt, pH 5.90, 25.0C. Background
electrolyte: KH2PO4, initial concentra
tion, 0.0050M
Log ( [AHP] / [AH] ) vs. log ( [Pj-/ 3 ] [AHP] ) for
the binding of 7-aminoquinolinium cation
to DNA at 15.0C, 25.0OC, and 35.0C, pH
5.90, 0.010M KH2PO4 as supporting electro
lyte
Log([AHP]/[AH]) vs. log([Pt/2l2[AHP]) for
the binding of 7-aminoquinolinium cation
to DNA at 15.0C, 25.0C, and 35.0C, pH
5.90, 0.010M KH2PO4 as supporting electro
lyte
Log K vs. 1/T(K) for the reaction between
3-aminoacridinium cation and DNA at 15.0C,
25.0C, and 35.0C, pH 5.90, 0.010M KH2P04
as supporting electrolyte
* o
Log Ks vs. 1/T( K) for the reaction between
7-aminoquinolinium cation and DNA at 15.0C,
25.0C, and 35.0OC, pH 5.90, 0.010M KH2P04
as supporting electrolyte
I
Log Kg vs. 1/T(K) for the reaction between
3-aminoacridinium cation and DNA at 15.0C,
25.0C, and 35.0OC, pH 5.90, 0.010M KH2P04
as supporting electrolyte
* o
Log Ks vs. 1/T( K) for the reaction between
7-aminoquinolinium cation and DNA at 15.0C.
25.0C, and 35.0C, pH 5.90, 0.010M KH2P04
as supporting electrolyte
Page
172
174
176
178
180
182
184
186
IX


LIST OF TABLES
Table
1
2
3
4
5
6
7
8
Molar absorptivities of the 3-amino-
acridiniura-DNA complex and the 7-amino-
quinolinium-DNA complex in various
supporting electrolytes
Absorptiometric titration of 3-amino-
acridinium cation with DNA, lithium salt.
Lithium dihydrogen phosphate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, sodium salt.
Sodium dihydrogen phosphate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium
salt. Potassium dihydrogen phosphate as
supporting electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium
salt. Potassium acetate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, rubidium
salt. Rubidium dihydrogen phosphate as
supporting electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, cesium salt.
Cesium dihydrogen phosphate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, tetramethyl-
ammonium salt. Tetramethylammonium
dihydrogen phosphate as supporting
electrolyte
Page
188
189
194
198
202
207
212
216
X


LIST OF TABLES (Continued)
Table Page
9 Absorptiometric titration of 3-amino-
acridinium cation with DNA, magnesium salt.
Magnesium acetate as supporting electrolyte 222
10 Absorptiometric titration of 3-amino-
acridinium cation with DNA, calcium salt.
Calcium acetate as supporting electrolyte 226
11 Absorptiometric titration of 3-amino-
acridinium cation with DNA, strontium salt.
Strontium acetate as supporting electrolyte 230
12 Absorptiometric titration of 3-amino-
acridinium cation with DNA, barium salt.
Barium acetate as supporting electrolyte 234
13 Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium salt,
at 15C 239
14 Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium salt,
at 35C 241
15 Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium salt,
at 15C 243
16 Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium salt,
at 25C 245
17 Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium salt,
at 35C 247
18 Calculated values of [BH], [BHP], "Ks", and
related data for the reaction between 3-
aminoacridinium cation and DNA in 0.15M
CsH2P04 249
19 Calculated values of [BH], [BHP], Ks, Ki, and
related data for the reaction between 3-
aminoacridinium cation and DNA in 0.0025M
CsH2P04 251
xi


LIST OF TABLES (Continued)
Table
20
21
22
Calculated values of [BH], [BHP],
and related data for the reaction
aminoacridinium cation and DNA in
Mg(02CCH3)
Ks, "Kl",
between 3-
0.010M
Calculated values of [BH], [BHP],
and related data for the reaction
aminoacridinium cation and DNA in
Mg(02CCH3)2
Ks, "Ki",
between 3-
6.3 x 10"4N
Percent total 3-aminoacridinium bound to DNA
in the presence of various electrolytes after
addition of excess DNA
23 Initial and final slopes of plots of log([BHP]/
[BH]) vs. log([Pt/m]-q[BHP]) in various con
centrations of CsH2PO^ and Mg(02CCH3)2
24 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj,
binding of 3-aminoacridinium to DNA, LiH^O^
as supporting electrolyte
25 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind
ing of 3-aminoacridinium to DNA, NaH2PO. as
supporting electrolyte
26 Apparent equilibrium association constants
for surface, K and intercalative, K bind
ing of 3-aminoacridinium to DNA, KH^O^ as
supporting electrolyte
27 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind
ing of 3-aminoacridinium to DNA, K02CCH3 as
supporting electrolyte
28 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind
ing of 3-aminoacridinium to DNA, RbH2PO. as
supporting electrolyte
29 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind
ing of 3-aminoacridinium to DNA, CsH2P04 as
supporting electrolyte
Page
253
255
257
258
259
261
262
263
265
267
Xll


LIST OF TABLES (Continued)
Table Page
30 Apparent equilibrium association constants
for the surface, Ks, and intercalative, Kx,
binding of 3-aminoacridinium to DNA, (CH^)^-
Nf^PO^ as supporting electrolyte 269
31 Apparent equilibrium association constants,
Ks, for surface binding of 3-amino-
acridinium to DNA, MgiO^CCH^^ and Ca (0200^)2
as supporting electrolytes 271
32 Apparent equilibrium association constants,
K for the surface binding of 3-amino-
acridinium to DNA, Sr (C^CCH^) 2 and BaiC^CCH^^
as supporting electrolytes 273
33 Logarithms of the apparent surface binding
association constants for the binding of 3-
aminoacridinium cation to DNA, corrected
for ion activities 275
34 Equilibrium association constants for the
binding of 3-aminoacridinium to DNA, Ks,
and for the binding of alkali and alkaline
earth metal cations to DNA, K,. 277
M
35 Apparent association constants for the
surface, Ks, and intercalative, Kj, binding
of 3-aminoacridinium and 7-aminoquinolinium
to DNA at 15.0C, 25.0C, and 35.0C 279
36 Thermodynamic parameters for the surface
and intercalative binding reactions between
3-aminoacridinium and 7-aminoquinolinium
cations to DNA 280
xiii


Abstract of Dissertation Presented to the
Graduate Council of the University of Florida in
Partial Fulfillment of the Requirements for the
Degree of Doctor of Philosophy
EFFECTS OF IONIC PERTURBATIONS AND METAL ION
COMPETITION ON THE BINDING OF A MODEL DRUG TO DNA
By
Peter F. Eisenhardt
August, 1977
Chairman: Stephen G. Schulman
Major Department: Pharmaceutical Chemistry
Absorptiometric titrations of a model drug compound,
3-aminoacridinium cation, with salts of calf thymus DNA
were carried out at pH 5.90 in a series of solutions in
which the supporting electrolyte was the dihydrogen
phosphate salt of one of the following: lithium, sodium,
potassium, rubidium, cesium, or tetramethylammonium ion.
Titrations were also conducted in solutions of the acetate
salts of magnesium, calcium, strontium, and barium. The
ionic strengths of the solutions were varied between 0.25
and 0.00063(M) for each alkali metal electrolyte and between
0.036 and 0.00010 for the alkaline earth electrolytes.
Apparent equilibrium association constants for the surface
xiv


binding and for the intercalative binding of 3-amino-
acridinium cation to DNA were computed for each concen
tration of each alkali metal electrolyte solution and
are reported. Apparent constants for the surface binding
of 3-aminoacridinium in the presence of alkaline earths
are also presented. Values for the intercalative binding
mode in the presence of group IIA metal cations could not
be determined.
The extent of binding of the model drug to DNA was
dependent on the nature and concentration of the electrolyte
Essentially complete binding was attainable for all concen
trations of all alkali metals while as little as 70% binding
was possible in the presence of 0.010N magnesium. Spectral
characteristics of the drug-DNA complex were also dependent
upon the nature of the background electrolyte and are dis
cussed.
The reaction systems did not behave in accordance with
limiting Debye-Hiickel relationships over the range of ionic
strength investigated. After applying extended Debye-
Hiickel corrections, the relationships between the surface
binding equilibrium association constants and the ionic
strength indicated that specific-ion competition between
metal and drug cations, for DNA phosphate sites, was signifi
cant. A simple 1:1 competitive model was evaluated.
At ionic strengths less than about 0.010, experimental
precision was poor, presumably due to alterations in the
conformation of DNA at low ionicities, suggesting that
V\7


evaluations of interactions between small molecules and
DNA be conducted in solutions of ionic strength in excess
of 0.010.
Thermodynamic parameters for surface and intercalative
binding of 3-aminoacridinium cation and its benzolog, 7-
aminoquinolinium, were evaluated from equilibrium association
constants at 15.0C, 25.0C, and 35.0C. The standard
enthalpy change, AH, for the 3-aminoacridinium binding
was less negative than AH for the 7-aminoquinolinium process,
while the standard entropy change, AS, for the former was
more positive than for the latter. Thus, the greater degree
of disordering resulting from the relaxation of the solvent
around the tricyclic 3-aminoacridinium cation as it binds to
DNA appears responsible for the more negative standard free
energy change for the 3-aminoacridinium binding process.
Values of AS for the intercalative binding of the two drugs
suggest that significant ordering of the polymer helix occurs
upon intercalation.
xvi


CHAPTER I
INTRODUCTION
The interaction of small molecules with DNA and
biopolymers, in general, has been the subject of intense
investigation during the past several decades. Interest
in this area has been due, in large part, to the early
realization that such interactions may result in muta
genesis (1) and carcinogenesis (2), especially when the
small molecules are cationic aromatics such as substituted
acridines. Numerous publications deal with various aspects
of the interactions of cationic aromatics with DNA, some
of which will be briefly reviewed here. The specific
systems to be considered in this work include 3-amino-
acridine (I) Hence, a summary of representative and
pertinent research involving aminoacridines is included.
Moreover, we have investigated the effects of metal ions
on the interactions of 3-aminoacridine with DNA and,
therefore, shall review the past research efforts in this
area. It is emphasized that this review is, by no means,
Roman numerals in parentheses after chemical names refer to
their listings in Figure 1 of appendix I.


2
exhaustive, but rather, is intended to provide perspective
and an introduction to the work that is included in this
report.
Binding of Small Molecules to Biopolymers
Probably the first acknowledgement of a significant
interaction between a member of the acridine chemical
family and a biological system was reported in 1913 by
Ehrlich and Benda (3) when they reported their observation
on the antibacterial activity of 3,6-diamino-10-methyl-
acridinium chloride(II) against Trypanosonoma brucei.
Subsequently, evidence accumulated that the antibacterial
activity of substituted aminoacridines is dependent on
their capacity to compete with hydronium ion for a vital
anionic site in a bacterium (4-6, 7: Ch. 2) and that
antibacterial activity and mammalian toxicity were not
directly related (8,9). The antibacterial action of
proflavine(III) is reported to be due to its ability to
inhibit DNA-dependent RNA synthesis (10) which suggests
that DNA may contain the anionic site. However, it is
inadvisable, at this time, to make a direct correlation
since proflavine also binds strongly to a variety of other
cellular structures, any one of which may contain the
vital anionic site (11).
Numerous amino-substituted acridines are known to be
carcinogenic while few simple acridines are cancer inducing
(12) which may be due to the former's greater ability to


3
intercalate into DNA (13). It is worthy of note that
benzacridines(IV) and dibenzacridines(V), first exten
sively studied in 1935, are nearly all highly carcino
genic (2). Two possible modes of action have been
proposed. The benzacridine may be metabolically activated
to a cationic form which then reacts as an electrophile
toward nucleophilic centers of DNA, RNA, and protein (14),
or the unmetabolized species may interact physically with
the DNA (15), perhaps by intercalating into the double
stranded helix.
Mutagenesis, presumably arising from the intercalation
of acridines with nucleic acids, was first reported by
DeMars in 1953 ( 1 ) in relation to his work on proflavine
and T2 bacteriophage (Escherichia coli). The mutagenic
activity of proflavine does not involve its integral
incorporation into the newly synthesized polynucleotide
chain but rather, results from the acridine's capacity to
cause deletions and insertions of nucleotides into the
DNA polynucleotides (16,17). In fact, proflavines' capacity,
in this regard, was used in 1961 to discover the triplet
nature of the genetic code.
In searching for effective anticancer and antitumor
agents, advantage is taken of the fact that compounds
which may interfere with cell replication at the molecular
level may have medicinar utility. As a result, a large
number of acridine derivatives have been investigated, only
to find that, while a select few do have antineoplastic


4
properties, their associated toxicities to the host make
them less desirable than other available compounds. On
the other hand, aminoacridines have enjoyed success as
antimalarial agents and are still in limited use today.
Investigations were begun in the 1920's for a substitute
for quinine, as an antimalarial, which led scientists
first to methylene blue, a phenothiazine dye, then to
pamaquine, both of which had low effectiveness against
acute stages of malaria. Structural similarities in the
above two compounds, with substituted acridines, eventually
led investigators to quinacrine(VI) which became the drug
of choice throughout World War II. Later, the more effec
tive and less toxic chloroquine(VII) was discovered.
Presently, considerable synthetic research is devoted to
acridine-derived antimalarials to supplant chloroquine
and to combat chloroquine-resistant P. falciparum
discovered in 1961 (18). To date, the most promising
compounds of this type have a diamine side chain in the 9-
position(VIII). Structure-activity relationships for a
variety of quinacrine analogues having side chain variants,
terminal amine variants, and aromatic substituent variants,
as well as quinine related acridine antimalarials, are
tabulated by Henry (19).
In light of the broad biological significance of the
binding of aminoacridines to DNA and other biopolymers,
in vitro and in vivo studies aimed at understanding the
exact nature of the interactions have been undertaken.


5
The following paragraphs will briefly review several of
the more important articles germane to our investigation.
Those dealing with the modes of binding will be outlined
first, followed by studies of the effects of the ionic
medium (including possibly competing ions), with a review
of papers involving theoretical discussions, third.
An excellent comprehensive article on the interactions
of aminoacridines with nucleic acids and the methods by
which they may be studied was published in 1956 by A.R.
Peacocke and J.N.H. Skerrett (20). Another, in 1968, by
Blake and Peacocke (21) considered developments in the
field after 1956.
Aminoacridines have long been known to bind to
double-stranded DNA via two distinct modes termed type
I and type II binding. Type I is the stronger and was
first proposed by Lerman (22), to be an internal or
intercalative process. Insertion of a molecular ion is
accomplished by transient rotation of base pairs allowing
the flat, lipophilic, aromatic portion of the molecule to
become situated between them. The charged portion of the
molecule was considered to remain external to the helix.
H.J. Li and D.M. Crothers (23) corroborated this model
when they presented indirect evidence that the base pairs
remain perpendicular to the helix axis upon intercalation
of the molecule. Another model for type I binding was
advanced by N.J. Pritchard, A. Blake, and A.R. Peacocke (24)
in which the small molecule is intercalated between adjacent


6
bases on the same polynucleotide chain. In both of these
models, the cationic heteroatom is associated with an
indeterminate number of anionic phosphate moieties along
the DNA backbone. The viscosity of heat-denatured DNA
is essentially unaffected by binding of acridines (25)
whereas the viscosity of native double-stranded DNA is
increased. An increase in viscosity is indicative of a
lengthening of the polymeric chain such as would result
from intercalation between base pairs or between adjacent
bases. Since the single-stranded DNA undergoes no vis
cosity change, the Lerman model appears more plausible.
Some specific geometrical parameters of substituted
acridine-DNA intercalative complexes were established by
G.R. Kelly and T. Kurucsev (26) through the evaluation of
linear dichroic spectra of stretched DNA films. An advan
tage of films over conventional methods involving aqueous
solutions is that they do not require large ratios of drug
to DNA phosphate. Taking advantage of the dichroism of
two mutually perpendicular polarized transitions of the
9-aminoacridine nucleus, the workers were able to assess
the angles of tilt and twist of the plane of the dye
relative to the polymer helix. The long axis of acri-
flavine(II) was essentially perpendicular to the helix
axis (85 + 5), proflavine(III) was tilted slightly (77 +
2 1/2), while 9-aminoacridine(IX) was tilted significantly
(67 + 3).


7
Type II, surface, or external binding has come to
be generally accepted as an association of the small
molecules with the exterior surface of the DNA polyanion.
This interaction is, consequently, more strongly affected
by alterations of the solvent system than is type I
binding. Li and Crothers (23) determined, on the basis
of temperature-jump kinetic data, that proflavine binds
both internally and externally to calf thymus DNA and
only about 7% of the total amount bound binds externally
in the presence of 0.2M Na+ (sodium phosphate sodium
nitrate). Decreasing the sodium ion concentration to
0.02M results in about 30% of the total being externally
bound.
A careful study of the intercalative binding of a
series of aminoacridines was conducted by Drummond and
co-workers (27) Binding decreases in affinity in the
order: acranil (X) >neomonacrin (XI) >atebrin (VI) ^-amino-
acridine (IX)>proflavine(III) 9-amino-l,2,3,4-tetra-
hydroacridine(XII). This sequence shows that side chain,
and ring, substituents are not predominant in influencing
binding affinity. The effect of greatly decreasing the
planar area of 9-aminoacridine by hydrogenation is
evident upon comparing its position with that of XII. The
authors discuss various geometries of binding and conclude
it is not necessary that there be exact and complete
intercalation for type I (strong) binding. Rather, a
modified Lerman model in which the charged heterocyclic
ring nitrogen interacts with a DNA phosphate while the


8
rings only partially interact with the bases is more in
accord with observed behavior.
It has been shown that the degree of intercalation
may be dependent on the base composition of the DNA.
Heterogeneity of intercalative sites was demonstrated
by J.C. Thornes, G. Weill, and M. Daune (28) using
fluorescence quenching techniques. They demonstrated
that between 2% and 3% of the base pairs form sites where
proflavine is strongly bound while the remainder have
binding constants 3 to 4 times weaker. Proflavine
fluorescence is relatively unaffected in the former case,
but is totally quenched in the latter. The stronger
sites correspond to adenine-thymine rich regions of the
DNA indicating proflavine's specificity for the base
pair. Similarly, Ramstein and Leng (29) showed that the
location of bound proflavine, within the DNA, was dependent
upon the base composition. The absorptiometric and flori-
metric titrations of DNA and partially methylated DNA
7
( 18%, mainly on N of the guanine residues) indicated
acridine was intercalated in both the native and methylated
sites. However, greater apparent equilibrium constants for
the methylated DNA than those for the native type were ob-
4 -1
served at two ionic strengths: 14.25 x 10 M and 43.5 x
10^M ^ in 0.1M NaCl and 40 x 104M ^ and 83.5 x 10^M ^ in
0.01M NaCl, respectively. These differences may be due to
an increase in the distance between base pairs arising from
the presence of the methyl groups which may, in turn, result


9
in the shift in the equilibrium concentrations. S.
Georghiou (30) obtained much the same results as Ramstein
and Leng based on the fluorescence decay of proflavine
when it was bound to DNA at a phosphate to drug ratio of
420. Gabbay and co-workers (31) invoked the "Bookmark"
model of intercalation to postulate ten possible distinct
sites of binding of N-substituted-N-methylphenanthrolinium
cations(XIII) to nucleic acids. They also investigated
the topography of some nucleic acids in solution using
PMR (32).
Binding models for the interactions of acridine
orange(XIV) and proflavine(III) with DNA at ionic strengths
of 0.002M, 0.020M, and 0.20M (tetramethylammonium caco-
dylate buffer, pH 6.5, 22C) were proposed based on
thermodynamic, spectroscopic,and hydrodynamic properties
of the systems (33). The methods precluded consideration
of surface binding. Binding isotherms were generated from
equilibrium association constants which were, in turn,
obtained from theoretical mass action relationships.
Additional information, useful in developing binding
models, was afforded by Goswami, Das, and Das Gupta (34)
who measured the decrease in the static dielectric con
stant of a proflavine-DNA complex while decreasing the
DNA phosphate/proflavine (P/D) ratio. The authors con
firmed that: the bound dye neutralizes the charge on the
DNA polyanion; at high P/D values, intercalation predomi
nates and; that when P/D becomes small (£10), a cooperative


10
increase in the electrostatic binding of phosphate groups
brings about a rapid decrease in the total surface charge
of the polymer. In brief, at P/D values between 10 and
100 most of the cations are intercalated with a concomi
tant elongation of the polyanionic DNA helix, whereas
below P/D of 10, surface binding is predominant. In the
latter case, the spatial extension of the helix is de
creased due to charge neutralization with no accompanying
expansion due to intercalation. These results indicate
that alterations in the dielectric constant of a medium
containing such a system would manifest dissimilar effects
on the two equilibrium association constants.
The specific effects of cations, other than the dye
species, on the reaction between DNA and a dye is an
integral part of this investigation. Therefore, a brief
consideration of articles dealing with the binding of
metal cations to DNA and related topics is in order. A
tabulation of sites and thermodynamic quantities associated
with proton and metal ion interactions with DNA, RNA, and
their constituent bases, nucleosides, and nucleotides has
been published by Izatt et al., (35).. H. Sigel and D.B.
McCormick (36) provide an introduction to the fundamental
theories concerning metal-ligand interactions in biological
systems, including discussions of the Irving-Williams
series, competition of metals and protons for sites, the
formation of ternary complexes, and effects of solvent
polarity and ionic strength. The Pearson theory (37) of


11
hard and soft acids and bases is applied to metals and
polymers in relation to the biological roles of Na+, K+,
Ca++, and Mg++ and to the carcinogenicity of heavy metals
(38) .
The relative effectiveness of ions in altering the
conformation of macromolecules via specific ion inter
actions is termed the Hofmeister, or lyotropic, effect
(39). Generalizations based on observations of simple,
ideal polymers may be confidently extrapolated to more
complex polymers since it has been shown that the
Hofmeister effect of a given ion is remarkably independent
of the nature of the macromolecule. Moreover, while
changes in the entropies of polymer segments are affected
by both anions and cations, as are changes in the partial
molal volume, the partial molal internal energy of a segment
is dependent only on cations. Considering the above,
M.J. Hey, J.M. Clough, and D.J. Taylor (39) reported that
the ion binding strength with protein is (in decreasing
order of strength) NH4+ This series may be useful in separating solvent effects
from ion binding effects where the above cations are
present along with 3-aminoacridinium cation and DNA.
Stabilizing effects of various ions on four biopoly
mers, including DNA, were evaluated by Hippel and Schleich
(40). The effects were discussed in terms of free energy
of transfer for model compounds, activity coefficient
variations, and correlations between ion effects on


12
macromolecules and on water structure. The relative
effectiveness of the ions studied, in increasing the
stability of the native configuration of DNA, paralleled
the Hofmeister series: (CH^)^N+>K+>Na+>Li+ and Cl Br >
CH-.COO >C10. >CNS It was found for a series of
3 4
tetraalkylammonium salts, the greater the length of
the alkyl chains, the lower the stabilization of the
native form of DNA.
Hen and Strauss (41) employed equilibrium dialysis,
dilatometry, and viscometry to assess the counterion
binding of a series of cations with poly(vinylsulfonate).
The Hofmeister series was maintained (in order of de
creasing binding affinity): Ag+>K+>Na+>H+,Li+>(CH^)^N+
and Ba++>Mg++. All of the cations except H+ and Li+
are capable of cross-linking the polyanion through
simultaneous site binding at two sulfonate groups.
The specific interactions of univalent and di
valent cations with calf thymus DNA can be correlated
with their respective unhydrated ionic radii (42,43).
The order of most weakly bound to most strongly bound
is: (CH^)^N+ binding parameters with DNA with polyphosphate as polyanion,
the authors showed that the sites on the biopolymer were
the phosphates rather than bases or base pairs. These
results were in agreement with those of other workers
(44,45).


13
Conductometric titrations of DNA with divalent
cations were first performed by J. Shack, R.J. Jenkins,
and J.M. Thompsett (46) who found a sharp end point cor
responding to 0.8 equilvalents of Mg++ and Ca++ per mole
of polymer phosphate. Similar titration behavior was seen
for poly A and poly U by Felsenfeld and Huang (47) who
developed an improved binding model which showed that the
apparent stoichiometry of 0.8 was in error and that the
true equivalence point corresponded to 1.0 eq of Mg++ or
Ca++ per mole of DNA phosphate. The authors also disproved
the previously held belief (48) that dications are more
strongly bound to purines than pyrimidines in a poly
nucleotide .
It has been reported (49) that there is binding to
base pairs of the DNA helix as well as to phosphates. The
investigators reported that the effects of adding Mg++ to
heat-denatured calf thymus DNA was to produce an absorption
spectrum similar to that of denatured DNA alone, except
for reduced absorbance. However, when Mg++ was added to
native DNA, which was then denatured, a spectral shift to
longer wavelength occurred. The authors concluded, on
these grounds, that the metal was binding only to phos
phates in the former case while in the latter, there was
interaction with the aromatic nitrogens of the bases. The
bonding between Mg++ and the nitrogens is presumably not
able to occur unless the hydrogen bonds between the bases
are broken in the presence of the metal.


14
Mathematical relationships incorporating changes in
activities of ions in solution and deviations in activity
coefficients were developed by Lyons and Kotin (45) to
qualitatively distinguish between specific (site) binding
and non-specific (diffuse electrostatic) binding of
metals to polyanions. The rule of additivity of the
activities of countercations of solutions of polyelectro
lytes and simple salts was re-examined. It was found
that, under certain conditions, there is significant
exchange between the counterions associated with the
polyanion and those of the bulk solution. A decrease in
the activity coefficient of sodium ion, upon dilution of
the system, was observed, indicating increased binding at
lower concentrations. In the case of magnesium the results
were not as simple but suggested that at high magnesium
ion concentrations, specific site binding was predominant
while at low concentrations, non-specific binding obtained.
The differences between Na+ and Mg++ may be due, in part,
to the ability of the latter to bridge between two adjacent
anionic sites on the polymer whereas the monovalent cation
cannot. Note that the distance between adjacent phosphates
o
along the DNA backbone is approximately 7 A. C.Sander and
and O.P. Ts'o (50) believe that the binding of Mg++ to DNA,
RNA, poly A, poly A*U, poly I, poly I*C, and denatured DNA
can be rationalized using a linear Scatchard relationship
(vide infra). Their specific site model was applicable for
up to about 70% saturation.


15
The degree of binding of acridine orange to native
and denatured DNA was determined by equilibrium dialysis
in 0.1M and 0.001M NaCl at 20C (51). Three successive
stages of binding were observed: the first corresponding
to intercalation of monomers; the second, formation of
bound dimers; and the third, external binding of aggre
gates. The appearance of the three modes of binding
compared to two for most aminoacridines is attributed
to acridine orange's proclivity to dimerize. The dialysis
data was augmented with absorption and fluorescence
spectra, electric dichroism, electric birefringence, and
circular dichroic spectra.
Scruggs and Ross (42) found that the intrinsic
viscosity of three types of DNA (salmon sperm, calf thymus,
and T4 phage) decreases in the presence of univalent cations
as the ionic strength increases, attaining separate limiting
values for each. Prior to their work, the apparent insen
sitivity of the viscosity of DNA to variations in ionic
strength was considered anomalous because polyelectrolyte
viscosities are usually highly dependent on salt concentra
tion. The workers asserted that previous invariance was
due to the presence of trace amounts of basic proteins,
polypeptides, and polyamines and that DNA was, in fact,
similar to all other polyelectrolytes.
S. Bram (52) investigated the secondary structure of
DNA in solutions of varying concentrations of Li+, Na+,
Rb+, and Cs+ using X-ray diffraction spectroscopy.


16
Differences between scattering patterns by sodium DNA
in solution and those expected from theoretical calcu
lations based on an unperturbed B form of DNA were due
to structural changes in the B form arising from the
polymer's interaction with the species in the medium.
This is in concert with others'(53-55) conclusions that
the number of base pairs per turn of the DNA helix is
a function of the ionic environment. For cesium DNA
the extrapolated lengths for 10 nucleotide pairs ar
o o .
43.5 A at zero ionic strength and 34.6 A at finite CsCl
o
concentration, whereas in NaCl, the values are 33.5 A
o
and 32.0 A, respectively. (A novel DNA conformation
model which differs markedly from the classical Watson-
Crick B form is proposed by Stig Erlander (56) who argues
that the differences between the values given above cannot
be rationalized using the Watson-Crick model.) A theo
retical treatment of cooperative binding was advanced by
G. Schwarz (57) in which nearest neighbors along a linear
polymer interact with one another. Two types of intrinsic
internal binding are assumed: (1) that of an isolated
ligand (nucleation); (2) that of a ligand to a site
immediately adjacent to one already occupied. The model
also includes considerations for dimer formation and
competitive (non-cooperative) binding at adjacent sites
by unrelated ligands. Matrix methods of calculating
species concentrations under equilibrium conditions are
presented.


17
An understanding of the effects of ionic strength
on the binding of small molecules to DNA requires at
least some knowledge of the charge distribution on the
surface of the polyanion. Electrophoresis and membrane
equilibrium experiments allow the determination of the
apparent fractional charge per DNA phosphorous, i, which
is an index of net charge (58). The value of i increases
from 0.250 to 0.34 on going from a solution containing
0.005m NaCl to one having 0.05m NaCl, showing that the
net charge of the polyanion increases with increasing
ionic strength. Ross (58) considered the above information
in light of Gorin's model of a long rigid cylinder (59)
and applied it to DNA, which is intermediate between a
rigid cylinder and a random-coil polyelectrolyte. Ross'
data substantiates Gorin's assumption that a rigid cylinder
(and, therefore, DNA) obeys Debye-Hiickel approximations.
Manning (60) postulated that, as the ionic strength
approaches zero, the effective charge of a polyelectrolyte
is maintained at a critical constant by condensation of
counterions along its surface as long as the formal charge
does not exceed a critical value. His predictions were
borne out by independent investigators (33) who found
that at low ionic strength (0.002M) the electrophoretic
mobility of DNA was unaffected by bound acridine orange.
Lerman's model for intercalation of small molecules
into DNA satisfies many experimental observations but
leaves some uncertainties unresolved arising from


18
consideration of the energetics involved (61). One of
the most perplexing problems is the strong dependence
of the intercalative mode of binding on the ionic strength
(albeit not as dependent as surface binding). Also, the
existence of a maximum degree of association as a function
of ionic strength requires explanation. Attempts have
been made to resolve these problems by invoking "stacking
energies" rationalized in terms of short range, ionic
strength independent interactions between DNA base pairs
and the aromatic portions of the acridines. These ratio
nalizations, however, are incompatible with the fact that
aromatic hydrocarbons intercalatively bind only very weakly.
M. Gilbert and P. Claverie (61) point out that the solution
energies of the intercalating compounds parallel their
binding affinities and, therefore, may be at least as
important as the stacking energies. Moreover, the limiting
of intercalation processes long before all available sites
have been occupied mandates a limitation mechanism inde
pendent of the heterogeneity of sites. The critical effect
of transferring positive charges from solvent "cavities,"
having dimensions on the order of the small molecules',
into large cavities interior to the lipophilic DNA must
be accounted for in a model of the dye-DNA complex as
it is in their model.
T. Herskovits (62) studied the relationship between
electrolyte concentration and solvent denaturation transi
tions in various media. Both increasing chain length and


19
increasing hydrocarbon content increase the effectiveness
of the denaturant. It was concluded that these observations
demonstrate the importance of hydrophobic forces in main
taining DNA's aqueous configuration. In most cases, the
denaturation was reversible with added electrolyte.
It is Bradley's opinion (63) that by using a combina
tion of statistical and quantum mechanical methods, it
should be possible to relate all absorption, emission,
equilibrium constant, and optical rotation observations
in a unified way. With this aspiration in mind, he
presented a statistical model of dye-polymer binding as
a step toward that goal. He assumed that specific polymer
sites were in a linear array and dye molecules were bound
to them by simple electrostatic forces, with possible
enhancement by neighboring sites. Furthermore, the sites
were sufficiently close to allow bound dye-bound dye
interactions with concomitant changes in their absorption,
fluorescence, and phosphorescence spectra. The distribution
of N consecutive sites occupied by dye molecules, in their
varying states of aggregation, can be computed theoretically
in terms of either the ground state free energy of inter
action of a pair of neighboring dye molecules, AF, or in
terms of a stacking coefficient, K, related to AF,
K = exp(-AF/kT) (1-1)
The AF and K terms may be equated to an experimentally
obtainable ratio of P sites per D dye species,


20
P/D = (1-F1//2) 1 (1+ (K-l) F) (1-2)
Values of K range from 1.25 for DNA to 826 for poly
phosphate (K = 6.2 for denatured DNA). There is
evidence that the value of K is directly related to
the conformation of the polymer. The fact that K is
equal to 1.25 and not to 1.00 for DNA as it would be
for a completely random site occupancy indicates some
non-randomness which could arise from dye-dye interaction
along the polymer. A simplified quantum mechanical
treatment is presented to allow computation of the magni
tudes of the relationships between the dye molecules of
filled units with empty sites.
Kinetics of intercalation of aminoacridines into DNA
were first studied in depth by Li and Crothers (23) in 1969.
Employing data from temperature-jump relaxation methods,
they proposed a relatively simple two-step process for
intercalative binding. It is worth noting that static
experiments will not allow a separation of these two steps
of the intercalative mode. The data fit the mechanism
^12 ^23
P + DNA > (P) (P) (1-3)
v" out ^, in
K21 32
where P is proflavine and (p)Qut and (P) are surface
bound and intercalatively bound dye, respectively. The
overall insertion reaction occurs in the millisecond time
range and is first order from the (P) state. The two-
out
step hypothesis is substantiated by comparing the large


21
relative thermodynamic and spectral changes of the
exterior binding with the small changes for the interior-
bound species that occur upon glucosylation of the DNA
(vide infra). Two limiting mechanisms are advanced and
alternate mechanistic pathways discussed.
Passero, Gabbay, and others (64) used a simplistic
model which included competitive effects of metal ions
in solution to arrive at apparent association constants
of two reporter molecules (Xllla, XHIb) to DNA. Electro
static potentials and interactions between near neighbors
were ignored, as were any forms of surface binding by the
reporter molecules. The overall binding process was
expressed as two separate equilibria
R + P > RP (1-4)
and
M + P > MP (1-5)
where R, M, and P are free reporter, metal, and DNA
phosphate, respectively, and RP and MP the corresponding
bound species. An apparent binding constant, K' may be
defined relative to the metal ion concentration
K'r = [RP] / [R] [P] (1-6)
Finally, assuming the binding of R or M to the polyanion
results in only one phosphate being affected per R or P,
the apparent association constant may be expressed in terms
of association constants for reporter binding, K and


22
metal binding,
K'r = Kr / (1 + Km [M+]) (1-7)
The authors maintain that their assumptions are valid
by virtue of the fact that they obtain reasonably
constant values of and K under a variety of conditions.
However, if the reporter molecules are intercalatively
bound, as the authors assert (and they are probably
correct) then it is not valid to compare the reporter
and metal binding because two entirely different types
of sites are being occupied. The highly hydrophilic,
symmetrically charged metal cations are almost certainly
surface bound. Thus, the constancy of the results are
probably fortuitous or, at best, important deviations
are being masked by experimental error.
Relaxation kinetics of ethidium bromide(XV) binding
to DNA revealed three types of bound species at equilib
rium (65). Aside from the conventional surface-bound and
intercalated species ascribed to type II and type I binding,
a third species was found which involved an ethidium bromide
transiently attached to sites of two separate DNA strands.
The bimolecular rate constant for transferring an ethidium
bromide from one DNA phosphate to another is 3 to 6 times
larger than the rate of intercalation of free molecules.
The more rapid bimolecular mechanism is dependent on the
small molecule's ability to bind to two sites simultaneously.
In the case of ethidium bromide, the tricyclic region of the


23
molecule provides one site while the phenyl moiety is
believed to provide the second, less favorable, site.
Divalent-cation-specific electrodes were used to
study the binding of Mg++ to DNA (50) in a solvent
-3
containing 5 x 10 M pH 7 phosphate buffer. Results
were treated in terms of site-binding to phosphates in
a multiple equilibrium process, employing a Scatchard
treatment, which showed linearity over a range of 30%
to 70% saturation of sites. Approximations to the
Scatchard method yielded an apparent intrinsic binding
constant of 6 x 10^ (AG = 5.2 Kcal/mole). The use of
cation-specific electrodes mandated DNA phosphate concen-
_3
trations in excess of 5 x 10 M, causing aggregation of
DNA and precipitation of MgDNA. Thus, the values of the
constants may include significant error.
Three methods for evaluating equilibrium association
constants for the binding of small molecules to DNA will
be outlined. They are, the classical and most commonly
employed Scatchard method, a modified Scatchard model,
and a mass-action approach which will be used for this
investigation. Scatchard's original treatment was designed
for binding of small molecules to proteins. In Scatchard's
t
words:
If the various groups on a protein molecule
act independently, we can apply the Law of
Mass Action as though each group were on a
separate molecule and the strength of binding
can be expressed as the constant for each
group. Often a single constant will express
the behavior of several groups. (66:660)
If the groups do act independently of one another, changes


24
in free energy for the reaction of the protein with
small molecules is made up of the statistical entropy
terms for v plus a term proportional to v. when the
total number of sites (or groups of sites) is large, the
calculations become tedious if the total number is
not known the calculations cannot be performed. Such
a situation does arise if one attempts to apply the
equations without modification to nucleic acid binding.
Since DNA polyanions commonly have molecular weights
/T
in excess of one million and may contain as many as 10
phosphates (67) Peacocke and Skerrett (20) redefined the
parameters of the Scatchard equations to allow their
application to nucleic acids. Their assignments and an
abbreviated derivation of the Scatchard equations are as
follows.
Let R represent the amount of small molecules, D,
bound per mole of nucleic acid phosphorus (generally a
fraction), and C represent the molar concentration of
free D. Consider a number, n^, of P classes of binding
sites for D on the DNA polyanion. If each class has
associated with it an intrinsic binding constant then
we may write for any general case
J=P
R = I
J=1
n .K .C
1 + K.C
1
(1-8)
This may be simplified when all sites are identical
R/C = Kn Kr
(1-9)


25
If there are two distinct types of binding sites, I
and II, then
R =
ni Ki C
1 + KjC
nll KII C
1 + KIIC
(1-10)
Figures 2a and 2b are representative plots of R/C vs C
for one and two types of sites, respectively. Curvature
may result in the case of only one type of site from
variations in the electrostatic free energy term which
is dependent on R. Such non-linearity can become critical
when the intercalating species bear opposite charges but
may be partially obviated by replacing with the term
K'z exp (AGr/(RT) (1-11)
where K'z is independent of R, and AG0^ is an electro
static free energy term dependent only on the total effec
tive potential. The case in which there are two types of
binding sites is of primary interest in this work. The
stronger, intercalative binding corresponds to type I of
figure 2b whereas surface binding is represented by
region II. With two types of sites, it is sometimes
possible to discern distinct linear portions in the plots
of R/C vs R as suggested in the idealized figure. Only if
the magnitude of one of the constants is about 10 to 100
times greater than the other, can reliable values of the
individual association constants be obtained by extrapola
tion of the data in each of the linear regions. Even then,


26
potentially erroneous assumptions regarding changes in
the electrostatic free energy must be considered. More
over, as previously stated, the proposed model does not
allow for any form of interaction between the two types
of binding sites. This is not generally serious in
protein binding where the sites are often noninteracting
but can be a problem in the case of the nucleic acids.
Factors such as cooperative binding and near-neighbor
exclusion will introduce sigmoidal characteristics into
the plots of R/C vs R (21). Another shortcoming of the
Scatchard method is that extrapolation of the legs of
the hyperbolic plot to obtain slopes and intercepts
often places an inordinately high degree of reliance
on the data at the extrema of titrations. It is at the
extrema that the data, regardless of how it was obtained,
will be the least precise.
For aminoacridines, intercalative binding pre
dominates up to a value of R_<0.20 corresponding to 4 or
more phosphates per bound cation. Surface binding, on
the other hand, is predominant for values of R between
approximately 0.5 and 1 for compounds which do not dimerize
and up to 2 or more for those that are known to aggregate
as free species at moderate concentrations.
Numerous investigators have suggested improvements
of the basic Scatchard method as modified by Peacocke and
Skerrett. Included are corrections for electrostatic
effects (27,31,61,63), competition of metallic ions in


27
solution (45,46,67), near-neighbor effects (23,50,57,
63,65), small molecule aggregation (51,57), and configura
tional effects due to solvent and ionic factors (23,25,28,
39,44,52,62,68). One of these improved models, proposed
by Armstrong, Kurucsev, and Strauss (33) incorporates
some novel features including: (1) the total concentration
of intercalation sites available is fixed, a priori, to
include every slot between successive DNA base pairs; (2)
an intercalated dye inhibits intercalation at its two
immediately adjacent slots; and (3) inclusion of a term
to allow for an intercalated monomer forming a dimer with
a free monomer resulting in a spectroscopically distinct
species. The authors found (3) became important when the
degree of binding, 6, exceeded about 0.2 mole of dye per
mole of DNA phosphate (dyes were acridine orange and pro
flavine) Their expressions for the association constants,
based upon theoretical mass action relationships are, for
intercalation
2(31-3,) (1-231)
(1-12)
and for surface binding
(1-13)
where 6^ = the number of intercalated molecules per DNA
phosphate.


28
82 = the number of externally bound dimers per
DNA phosphate.
C, = the molar concentration of free dye monomer.
M
= the apparent binding constant for the monomer
species intercalating.
K2 = the apparent binding constant for surface
binding.
These relationships imply that the electrostatic potential
of the DNA is unaffected by binding. Values of 8-^ and
may be spectroscopically evaluated knowing the molar
absorptivities of the bound monomer, e^, and the bound
dimer, eD,
£T ~ + (1-14)
where 8^ = 6^-82 = ^^2' anc^ eT as tota^ "mlar
absorptivity" of all bound species. The model was tested
by determining the binding of acridine orange and pro
flavine to DNA (pH 6.5, 22C) at ionic strengths of 0.002m,
0.020m, and 0.200m. The above method is an improvement
over the Scatchard procedure but still demands a priori
selection of an exact mode of reaction. That is, though
the overall stoichiometries are experimentally evaluated,
the thermodynamic manner in which the free species react
is not defined. For example, in the intercalative case,
where 4 phosphates are involved for each 3-aminoacridinium
binding site, are the 4 phosphates to be considered as 4
separate entities, as 2 separately interacting moieties,


29
each containing 2 phosphates, or as one entity containing
all 4 phosphates? In each case, the overall stoichiometry
is 4 to 1, but the equilibrium expression of the first
includes a fourth power term in DNA phosphate concentration,
the second would have a squared phosphate concentration
term, while the third would be linear in phosphate.
To resolve this problem, Capomacchia (69) has proposed
a mass-action approach based on best-fit analyses of
spectroscopically derived concentrations of free and bound
species. Assuming that the binding process obeys the Law
of Mass Action and total concentrations of all species are
sufficiently low that aggregation of like species is in
significant, we may write
BH + qSu ~ > BH(Su) (1-15)
where BH is the free cationic form of the small molecule,
q is the number of DNA binding sites that complex one
small molecule and Su denotes the unbound DNA binding
sites. Each site, Su, acts as one distinct entity in
the equilibrium expression and may contain any number
(1,2,3,. . m) of phosphates. Note that throughout this
text binding sites will be discussed in terms of DNA
phosphates, simply for convenience, because the total
concentration of DNA phosphate is easily obtained ex
perimentally. This terminology does not imply, for
instance, that phosphates are the binding sites for
intercalatively bound drug. The association constant,


30
on a molar basis, for the process is
K = [BHP]/[BH][Su]q
(1-16)
where [BHP] is the molar concentration of bound species,
[BH] is the concentration of free small molecule, and
[Su] is the concentration of the uncomplexed DNA sites.
Unbound site concentration may be determined from the
total site concentration, S^, and bound site concentration,
S
c
(1-17)
Now, since
Sc/q = [BHP] (1-18)
and st = [Pt/m] (1-19)
where Pfc is total DNA phosphate concentration, we may
write
su = [pt/m] qtBHp]
(1-20)
Substituting equation (1-20) into (1-16) yields
K =
[BHP]
[BH]([Pt/m]- q[BHP])q
(1-21)
Taking the log of both sides of (1-21) and rearranging,
we obtain
109 ( !bh]] = log K + qlog ([pt/ml- q[BHP])q (1-22)


31
Molar concentrations of free and bound small molecule
species may be determined from any of a variety of
experimental techniques. The overall stoichiometry of
the process may be obtained from a Scatchard treatment or
a Job's plot. Knowing the overall stoichiometry, initial
values of m and q may be selected which satisfy the simple
relationship mq = n, where n is the total number of DNA
phosphates per bound small molecule. A rapid determination
of the most probable combinations of values may be made by
plotting log ([Pt/m]- q[BHP]) vs.log ( [BHP] / [BH] ) That
pair which yields the most consistent straight line having
a slope of q may then be used to compute values of K for a
series of ratios of [BHP] to [BH]. Again, the pair which
provides the most consistent calculated equilibrium associa
tion constants for all data points, corresponding to either
the surface or intercalative mode, are taken as the exponents
in the binding expression. For interactions between 3-
aminoacridinium cation and calf thymus DNA, best-fit values
of m and q are 3 and 1,respectively for surface binding and
2 and 2 for the intercalative mode (69) To summarize, the
method of Capomacchia allows for the determination of
equilibrium association constants based on mechanistic
considerations. Furthermore, by computing values of K from
data in the midregion of a titration, undue reliance on the
least trustworthy data at the extrema of a titration is
avoided.
Many workers have been satisfied with reporting associ
ation constants calculated from molar quantities rather than


32
from the more correct activities of the various species.
Doing so implies two things: activities of the various
species in the expression are equal to molarities, and
the molarities are equal to molalities. It is unlikely
that approximating molalities with molarities introduces
significant error throughout the usual concentration
ranges employed. But activities, on the other hand,
may differ markedly from molarities. For this reason,
evaluation of association constants in terms of activities
is recommended. Molality is related to activity by
a. = a.m. (1-23)
ill
where a^ is the activity, ou is the activity coefficient,
and itk is the molality of species i. Note that the
activity equals the molality only when a = 1. For a
process in solution
A+ + nB v^ ABn(n+l) (1-24)
the equilibrium constant is rigourously defined as
aAB
n
(A++)(A¡-)n
a M
AB AB
n n
(c++Ma+) ( (1-25)
It is extremely difficult to determine individual activities
since their is no way of separating effects due to positive
ions from those due to the accompanying negative ions. We
may, however, define a mean activity, a, which is the
geometric mean of the individual activities


33
(1-26)
Similarly
(1-27)
For ideal dilute solutions, the mean ionic activity
coefficient will be equal to 1. If deviations from
ideality are caused entirely by electrical interactions
of point charges, it may be shown that, for aqueous
solutions at 25C (70)
log a- = -.509/Z + Z_/l1//2
(1-28)
where Z+ and Z_ are the charges of interacting species
and I is the total ionic strength of the system
2
m. Z7
1=1/2 E
(1-29)
. ill u
ill
Equation (1-28) is an expression of the Debye-Hiickel
limiting law for activity coefficients and is applicable
only to dilute solutions containing point charges. As the
ionic strength becomes large (e.g., >0.01m) and/or the
ions become less ideal (become larger, more polarizable,
etc.) the relationship becomes increasingly approximate.
Under these conditions, an extended Debye-Hiickel equation
may be invoked which introduces a correction for the finite
sizes of ions:
log =
-0.509/Z + Z Jl1/2
1/2
1 + Bdl
(1-30)


34
B is a constant for a given solvent and includes such
terms as solvent viscosity, dielectric, and temperature
while d is an average effective diameter of the ions. A
factor not considered in equation (1-30) is the effect of
ions on solvent molecules and solvent structure. This
may become especially important for solutions containing
large polyanions such as DNA.
Lyons and Kotin (44,45) report a method of assessing
values of a+ and a independently of one another. They
also determined for Na+ and Mg++ when binding to DNA
and three synthetic polyanions. Their somewhat tedious
procedure involving equilibrium dialysis allowed them to
distinguish site-specific and non-specific forms of surface
binding. The activity coefficient of Na+ was shown to
decrease with increasing dilution of metal ion in the
presence of DNA suggesting an increased degree of binding.
Thermodynamics of Binding Reactions
A second topic of this work is concerned with the
thermodynamic parameters associated with the binding of
small molecules to DNA. Hence, a brief review of several
pertinent articles is presented here.
M. Gilbert and P. Claverie (61) measured total energies
for an intercalation process and attributed them to three
types of electrostatic interactions. They were (1) the
attractions between the bound cations and the DNA phosphates,
(2) the repulsive forces between cations, and (3) solvent


35
response to the electric field generated by the charged
species. At high ionic strength the phosphates are
relatively neutralized causing (1) to become insignificant
and, thereby, causing a decrease in both the strength and
extent of binding. Some limiting of the binding process
occurs even at very low ionicities because cation-cation
repulsions are still important with respect to intercalated
species. The third type of interaction serves to reduce
all forms of electrostatic energies and will be most ef
fective for those species most intimately in contact with
the solvent. The authors have shown that phosphates and
complexed acridines are more solvated than free acridines
so that, as the ionic strength is raised, their stabilizing
interactions are more reduced than are acridine-acridine
interactions. In that the forces arising from the latter
are destabilizing, the result of the countervailing trends
is a net destabilization of the complex with increasing
ionic strength. The model used here can also be applied
to other cationic small molecules, including protons.
Similar results were reported for the binding of proflavine
to DNA via fluorimetry (71). It was found that the electro
static contribution to the total free energy decreased with
increasing ionic strength due to screening of the potential
by an ionic atmosphere.
Thermodynamic parameters accompanying the intercalation
of calf thymus DNA were evaluated by Chambron, Daune, and Sadron
(72) using equilibrium dialysis. Temperatures ranged from


36
0C to 70C and ionic strengths of 0.01M, 0.1M, and 1.OM
were maintained using acetate buffer at pH 5.9 in the
presence of EDTA. The number of sites available for
binding decreased upon increasing either the temperature
or the ionic strength. The enthalpy of reaction decreased
with decreasing temperature; AH= -19 Kcal/mole at 70C
and -4.5 Kcal/mole at 0C (I = 0.01). This was ascribed
to a thermally reversible change of state of DNA which
occurs at approximately 40C. Total standard free
energy was considered to be a sum of three terms. The
first of these, AG, characterizes the dipole-induced
dipole interaction between base pairs and cations while
the second, AG, is the electrostatic free energy. The
third, AG, is the solvation free energy corresponding
n
to desolvation of reactants and solvation of the bound
complex. At high ionic strengths, AG and AG are very
small, hence, the total free energy of 6.4 Kcal/mole is
due to AG. The authors' value of AS= 11 Kcal/mole
degree is in agreement with theoretical predictions.
Thermodynamic energies of reaction between proflavine
and two forms of DNA determined from temperature-jump
relaxation methods are summarized below (23).
DNA
EI2
E- E-
^21 L23
E
^32
AB2
AH23
TAS2
TAS23
calf
thymus
4
14 16
14
-9.8
2.0
-5.4
3.2
T2

13
17
-3.4
-3.9
1.6
-2.8
ETj ls
the activation energy for
transformation
from
state i


37
to j (see equation (1-13) and AH?^ and AS?^ are the thermo
dynamic enthalpy and entropy changes for the same reaction.
Measurements were done over the temperature range 10C to
25C, 0.20M Na+ and pH 6.9. Of particular note here is
that the enthalpy change for outside binding is much more
strongly affected by glucosylation of the DNA than it is
for the intercalated form (T2 DNA is glucosylated). For
example, AH2 changes from -9.8 to -3.4 Kcal/mole dye on
going from calf thymus to T2 DNA but for the overall inter
calation process (AH2 + AH2 = ^H13^ t*ie c^an5e i-s only
from -7.8 to -7.3 Kcal/mole dye. It is unlikely that small
differences in base composition between the two forms of
DNA can account for the large differences in the thermo
dynamic values of the binding reaction.
Entropy and enthalpy changes during reaction of
ethidium bromide with calf thymus DNA were computed from
temperature-jump relaxation data by Bresloff and Crothers
(65) over the temperature range 15C to 32C in 1.OM Na+.
Values for the formation of a surface bound species were:
AH = 1.1 Kcal/mole and AS(19C) = 5.6 Kcal/mole, while
for one of a pair of intercalatively bound species: AH =
-7.8 Kcal/mole and TAS(19C) = -2.3 Kcal/mole. Values
for a second intercalated form could not be obtained due
to the rapidity of its binding process.
Bradley (63) determined the ground state free energy
of intercalation of a pair of neighboring acridine orange
molecules bound to DNA to be -0.065 Kcal/mole dye. By


38
comparing this to the spectral shift accompanying aggrega
tion (504 nm to 464 nm) corresponding to 4.8 Kcal/mole dye
separation in energy between the ground and excited states,
he concluded that the spectral shift results from an
increase in energy of the excited state, rather than a
decrease in the ground state.
Enthalpy and entropy changes, AHand ASrespectively,
of reactions may be determined from the relationship be
tween the equilibrium association constants for the processes
at various temperatures and the temperature as set forth
by the van't Hoff relationship.
l0g K 2.303R T
(1-31)
+ Constant
This important equation will be derived below (73).
For ideal systems, any change in the total free
energy, G, can be said to be due to changes in either
temperature or pressure, since the state function, G,
can be defined in terms of pressure, volume, and temperature.
(1-32)
We define
G = H TS
(1-33)
and
H = E + PV
(1-34)
where H is the enthalpy, S is the entropy, and E is the


39
internal energy for the system under consideration. Recall
that, the enthalpy changes are commonly termed "heats of
reaction" for processes, entropy may be used as a measure ,
of the degree of order or disorder, and that the internal
energy is unavailable for work of any form. The change in
the free energy, dG, may now be expressed as
dG = dE + PdV + VdP TdS SdT (1-35)
For a reversible process in which no work other than
expansion occurs,
PdV = dw (1-36)
and
dE dq + dw = 0 (1-37)
The terms q and w refer to the heat absorbed, and work
done, by the system respectively. Equation (1-37) is a
direct consequence of the First Law of Thermodynamics and
states, in effect, that the total change in the internal
energy of a system is the sum (or difference) of the amount
of heat absorbed, dq, by the system and the amount of work
done, dw, by the system. Notice that heat and work are not
functions of state (which are conventionally assigned upper
case symbols). Equation (1-35) now becomes
dG = VdP SdT
(1-38)


40
which upon equating with equation (1-32) yields
(3T}p S
(1-39)
and
<|§) = V (1-40)
Equation (1-39) suggests an experimentally tractable
relationship between free energy (via K) and temperature,
when pressure is held constant. Moreover, we may extend
the relationship to include the changes in free energy as
a function of temperature for a process involving reactants
going to products. We may now equate (1-38) with equation
(1-32), rearrange, and solve for the partial differentials
to obtain
(9AG/9T)p = -AS
(1-41)
where AG = (G G , ) and S = (S ,
products reactants products
S reactants^ The entroPY term of equation (1-41) may be
eliminated by recalling that AS = (-AH + AG)/T and substi
tuting
(9G/9T)p = (-AH + AG)/T
(1-42)
Assume, henceforth, that the pressure will be held constant.
Rearranging (1-42) gives


41
dAG/dT AG/T = -AH/T (1-43)
Equation (1-43) may be expressed as
T (AG/T) = -AH/T (1-44)
since, upon differentiation, it would become
m T(dAG/dT)-AG dAG/dT
1 2
T T
Equation (1-45) may be written
and, in that
AG = -RT In K (1-46)
AG
T
(1-45)
in terms of standard states,
it can be stated as
d In K/dT = AH/RT2
which is equivalent to
d log K
d (1/T)
-AH
2.303R
+ Constant
(1-47)
(1-48)
Assuming that AH is temperature independent, equation (1-48)
can be integrated to yield equation (1-31). The above
relationship may be applied to reactions in solution whose
equilibrium association constants are based on the activities
of the species involved [see equations (1-23) through (1-30)].
Plots of log K vs. 1/T should yield straight lines having
slopes equal to -AH/2.303R. If the relationship between
log K and 1/T is not linear then the enthalpy of binding is


42
not constant with respect to temperature. For systems
containing DNA, enthalpies of binding determined from van't
Hoff plots are often not constant with temperature, but
decrease with increased temperature, presumably due to
changes in DNA's conformation (72). As a result, it should
not be surprising that binding curves obtained spectro-
photometrically and by dialysis equilibrium can agree at
one temperature and disagree at another (21).
Rationale for the Selection of the Model System
and Experimental Design
As this study is concerned with the effects of com
peting metal ions and of alterations of the ionic strength
on the binding of small molecules to biopolymers, the choice
of DNA as the model polymer is evident. The selection of
3-aminoacridine as the small molecule is based on several
factors, two of which are especially important. The first
criterion that the small molecule must meet is that it bind
to DNA. To do so, it is necessary that the species exist
as a cation since neutral compounds do not bind to DNA
or if they do, the binding is too weak for our purposes.
(It should be evident that electrostatic repulsive forces
prevent anions from interacting with DNA.) Furthermore,
the compound must be positively charged in solutions whose
pH is between 4 and 11 so that denaturation of the polymer
is minimal (74). Hence, if the compound is acidic its pK
cl
should be at least 5. The pK of 3-aminoacridine is 8.04
a
(7:Ch 4) meaning it is essentially totally protonated in a


43
solution of pH 6. Moreover, 3-aminoacridine allows con
venient evaluation of both of the principal inodes of binding
of small molecules to DNA, under a given set of conditions,
in a single titration. That is, the range of the ratio of
total drug to total DNA concentration over which the two
modes of binding can be experimentally evaluated is about
0.1 to 10. Obviously, it is mandatory that each type of
binding occur in separate regions of the titration.
The second major criterion that the model compound
must satisfy is that it allow an expedient means of measur
ing the extent of binding to the polyanion. Of the various
methods of determining equilibrium constants one of the best
is absorptiometric spectroscopy. The 3-aminoacridine cation
is particularly amenable to this method for the following
reasons. (1) It absorbs strongly in the visible region of
the spectrum as the free cation having molar absorptivities
of 1.41.3 x 10^M ''"cm ^ at 365 nm and 1.259^ x 10^M ^cm ^ at
454 nm (75). Such magnitudes allow good sensitivity at low
concentrations of the compound. The molar absorptivities
of the bound drug in this region are 8.494 x 10 M cm at
368 nm and 9.07.1 x 10^M "''cm ^ at 463 nm. The large dif
ference between the absorbances of the band maxima of the
free and bound drug provides good sensitivity in following
a titration. The longest wavelength absorption maximum
for DNA is at approximately 260 nm meaning that there is
little spectroscopic interference from the polymer. (2)
Both the free and bound species have two distinct absorption


44
envelopes in the visible region corresponding to and
bands [Platt's nomenclature (76); a and para bands,
cl
respectively, in the Ciar nomenclature (77)]. These arise
from transitions from the ground state to the
(=365 nm), and to the (=460 nm) excited states. The
transition moment of the former is parallel to the long
axis of the acridine ring system while the moment of the
latter is orthogonal to the and parallel to the short
axis of the molecule. By observing shifts in the two
maxima, along with any changes in vibronic character,
during the course of a titration, one may be able to
deduce binding geometries (78) The 3-aminoacridine
molecule also fluoresces in its free cationic form and,
less intensely, as the bound species, thereby providing
another spectroscopic method with which to study its
interaction with DNA.
The more thoroughly studied aminoacridines, proflavine
and acridine orange, meet the criteria mentioned above in
regard to 3-aminoacridine. Both of them, however, begin to
-5
aggregate at concentrations as low as 10 M, thereby
introducing undesireable complications. Also, the mono-
substituted 3-aminoacridine, being a simpler molecule than
either proflavine or acridine orange, is preferable as a
model compound. It is hoped that information gained using
the simple compound may be applied to more complex systems.
Many references to the effects of ionic strength on the
binding of small molecules to biopolymers can be found in


45
the literature. Unfortunately, the majority are ancillary
to other studies and are not sufficiently complete in them
selves to provide much specific information. Too, even for
a given molecule binding to a given polymer, the wide
variety of experimental variables under which the separate
investigations were done precludes any valid comparisons of
results. The present work was undertaken as a result of
this lack of a comprehensive investigation. Our pharmaceu
tical interest led us specifically to DNA as the polyanion,
though other, synthetic polymers may have provided a less
complex system.
Some of the immediate questions which come to mind and
which may be answered by an investigation of the effects of
changing the ionic constitution of a drug-DNA solution are:
are the effects predictable on the basis of electrostatic
theory? That is, may the effects be explained in terms of
the simple Debye-Hiickel limiting law at low ionic strengths
and by extended Debye-Hiickel equations at higher ionicities?
If the above questions cannot be answered affirmatively,
might deviations at low ionic strengths be due to the
polymeric nature of the anion? Are deviations at high
ionic strengths due to specific ion interactions and, if so,
can they be rationalized on the grounds of simple competition
for binding sites or must a less simplified, mixed complex
model be introduced? Other questions are: will altering the
ionic medium affect surface and intercalative binding in the
same manner? If not, how can the differences be reconciled?


46
How will changing the ionic radii, ionic charge, and
charge density of the metallic ions affect the equilibria?
The last question is particularly germane in in vivo
biological binding studies because of the wide variations
in concentrations of Na+, K+, Mg++, and Ca++ in a living
organism (79).
To try to answer these questions, equilibrium associa
tion constants for the surface and intercalative binding
of 3-aminoacridinium cation to DNA in a variety of ionic
media were determined. Background electrolytes included
-4
concentrations between 0.15M and 6.3 x 10 M of the phosphate
salts of Li+, Na+, K+, Rb+, Cs+, and (CH^J^N*. Also con-
++ ++
sidered were the acetates of the alkaline earths: Mg Ca ,
Sr++, and Ba++ at concentrations ranging from 0.025N to 1 x
10 ~*N. The effect of as opposed to CH^COO as the
counteranion of the supporting electrolyte was evaluated.
The pH for all of the titrations was 5.90 and the tempera
ture was 25.0C. A minimum of two titrations were done for
each set of conditions.
Thermodynamic parameters were evaluated for systems
containing 3-aminoacridine and DNA with a background
electrolyte of 0.010M KH2PC>4 at pH 5.90. Similar investiga
tions were conducted using 7-aminoacridinium cation(XVI) in
place of 3-aminoacridinium. By studying the binding of the
linear benzalog of 3-aminoacridine and comparing the thermo
dynamic parameters of the two, the relative importance of
enthalpy and entropy contributions to overall free energies
may be assessed.


CHAPTER II
EXPERIMENTAL
Materials and Procedures
All chemicals were, unless otherwise stated, either
reagent or analytical grade. Water was multiply distilled
in a tin-lined still (Barnstead, Sybron Corp.). It was
redistilled until its specific conductance was less than
1.0 megaohm cm 1 at 25.0C. Stock solutions of supporting
electrolytes were prepared within three weeks of use and
were refrigerated during storage to minimize biological
growth. Approximately 100 to 200 ml of a 0.15M solution
of each electrolyte were prepared as outlined below and
titrated with 0.72M H^PO^, 1M CH^CC^H, or 1M metal hydroxide
solution to pH 5.9, in all cases. Concentrations of the
hydroxide solutions of the various cations were determined
by titration with standardized 1.000N sulfuric acid.
A stock solution of tetramethylammonium phosphate was
made by titrating 13.16 ml of 1.14M tetramethylammonium
hydroxide (Mallinckrodt Organic Reagent, 10% in water,
lot: 15) in 70 ml of water with 0.72M phosphoric acid
(Fisher, lot: 755459) and adjusting the final volume to
100 ml. Lithium phosphate stock solution was prepared by
47


48
titration of 0.476M lithium hydroxide with 0.72M phosphoric
acid. The hydroxide solution was made by reacting stoi
chiometric amounts of anhydrous lithium sulfate (PCR, Inc.,
Gainesville, Florida, lot: 10465) and barium hydroxide
heptahydrate (Mallinckrodt, lot: RTZ). Sodium and potas
sium phosphate stock solutions were prepared using sodium
phosphate monohydrate (Mallinckrodt, lot: ABK) and mono
basic potassium phosphate (Fisher Certified, Reagent, lot:
20175), respectively. The acid level of each solution was
adjusted with 0.72M phosphoric acid to pH 5.9. The stock
solution of rubidium phosphate was prepared by titrating
an aliquot of 0.43M rubidium hydroxide (Pfaltz and Bauer,
no lot number) with 0.72M phosphoric acid. Similarly,
cesium phosphate was prepared from 0.201M cesium hydroxide
(Pfaltz and Bauer, no lot number). Stock solutions of
potassium acetate and tetramethylammonium acetate were
prepared by titrating standardized aliquots of the respec
tive hydroxides with 1M acetic acid (Baker Chemical, lot:
41397). Magnesium acetate solutions were made from
magnesium acetate tetrahydrate (Mallinckrodt, control:
NSB). Solid calcium acetate was made by dissolving calcium
carbonate (Mallinckrodt, no lot number) in excess acetic
acid. The calcium acetate was recrystallized twice from
aqueous ethanol, washed repeatedly with 95% ethanol, and
dried for twenty-four hours at 110C to yield the mono
hydrate of the salt. The sample was found to contain 97.3%
Ca(C2H3O2).H2O, by dissolving a known amount of the acetate


49
in water, precipitating the calcium as calcium carbonate
and titrating the carbonate with hydrochloric acid.
Strontium acetate was prepared from its carbonate (Matheson,
Coleman & Bell, lot: 19) and assayed in a manner similar
to that for calcium acetate. It was determined to contain
95.0% SriC^H^C^). A recrystallized sample of barium
acetate was dried for twenty-four hours at 170C and
assayed by forming barium carbonate. The solid carbonate
was reacted with excess 1.00N sulfuric acid, carbonic acid
removed by gentle heating, and the excess sulfuric acid
determined by titration with standard sodium hydroxide
solution. The assay of the barium acetate was 99.6%
Ba (C2H302).
Solutions of the supporting electrolytes to be used
as solvents were made from the stock solutions. Aliquots
of the stock solutions were diluted to within 90% of their
calculated final volumes and their acidities adjusted to
pH 5.90 + 0.05, if necessary, using the appropriate metal
hydroxide, phosphoric acid, or acetic acid. They were
then diluted to their exact desired final volumes and their
ionic strengths computed, taking into account any additional
ions added to adjust the acid levels.
A commercially available sodium salt of DNA was
dialyzed against a solution of the desired metal phosphate
or acetate to obtain the DNA salts having the same cation as
that of the supporting electrolyte. Solutions containing
approximately 0.01 mole of DNA phosphate per liter were made


50
by placing 3.6 mg of DNA, sodium salt (Calbiochem calf
thymus DNA, A grade, lot: 900007: 8.04% P; 12.21% N;
1%
moisture, 14%; E260nm = Per of solvent
into twenty milliliter volumetric flasks. Between 5 and
20 ml of solvent were then added. The solvent was a 0.01M
solution (pH 5.90) of the metal phosphate or acetate. The
flask and its contents were placed on a vertical rotating
mixer (Scientific Products model 150-V) operating at 6 rpm
for about one hour, after which the sample was placed in
an ultrasonic mixer for two to three minutes. Rotation
and sonication were repeated for a total of at least five
hours or well beyond the time at which the moderately
viscous sample appeared homogeneous. The sodium DNA solu
tion was then transferred to dialysis tubing (Union Carbide
Films Packaging Division, 7.3 mm dia.). The tubing and its
contents were immersed in a volume of 0.01M solution of the
desired salt which was at least ten times greater than the
volume of DNA solution being dialyzed. The external solu
tion was agitated for two hours and then replaced. A
total of six volumes of external solution were used for
each dialysis procedure. The dialyzed DNA was again mixed
for one hour on the rotating mixer and submitted to ultra-
sonication for two to three minutes. Then, the sample was
R
divided among Pyrex culture tubes (6x50 mm) each
receiving approximately one half milliliter of sample.
These were immediately stored at -5C.


51
Determinations of the concentrations of DNA after
dissolution and after dialysis yielded values which were
identical, within experimental error and also within + 5%
of the calculated concentration based on the manufacturer's
data. The agreement of these data were taken as evidence
that any degradation of the biopolymer during the pro
cedures was minimal.
The efficiency of the dialysis method, used here, to
replace the sodium cations of the original DNA salt with
the cations of the external solution was evaluated by
atomic absorption spectroscopy. A Perkin-Elmer model 290B
atomic absorption spectrophotometer outfitted with a combina
tion sodium/potassium lamp was used to measure the residual
concentration of sodium ion in a sample of sodium DNA which
had been dialyzed against a potassium phosphate solution.
Instrument readings were taken using the contents of the
dialysis bag and compared to those obtained for standard
solutions in which the ratio of potassium to sodium ions
was approximately one thousand. Note that in this instance
high relative concentrations of potassium ion must be included
in the standard solution. Less than 0.1% of the original
sodium remained within the dialysis bag when the above
procedure was employed.
Crystalline 3-aminoacridine was prepared by Dr. Timothy
Roy using a standard procedure (80) except for additional
purification of the product by vacuum sublimation. Thin
layer chromatographic analysis of the purified product


52
using a variety of solvents and solid supports failed to
indicate the presence of any impurities. The spectra of
solutions of 3-aminoacridine at pH 11.0 and at pH 2.5 had
the same relative absorptivities as those reported by
Albert (75). Reported (pH 11.0) (log e) : 237 (4.46),
262 (4.83), 321 (3.35), 337 (3.65), 353 (3.92), 410 (3.79).
(pH 2.5) 233 (4.62), 274 (4.65), 349 (4.03), 365 (4.15),
454 (4.10). The melting point agreed with that reported
(81): 218C, uncorr., 223.5C, corr.
-3
Stock solutions of 1 x 10 M 3-aminoacridine were
prepared by dissolving 2 mg of the solid in 7 ml of absolute
ethanol. The solutions were stored at -5C in glass. One
hundred microliter aliquots of the ethanolic solutions were
used for each titration. Periodic tests of the stock
solutions using thin layer chromatography showed no observ
able degradation during the time periods in which the samples
were being employed.
Absorptiometric titrations were conducted by successive
additions of microliter aliquots of DNA stock solution to
10 ml solutions of 3-aminoacridinium cation. After each
addition, the absorbance spectrum between 550 nm and 320 nm
of the solution was recorded. To prepare the solutions, 10 ml
of supporting electrolyte at a given concentration were
measured in a 10 ml volumetric flask. This volume was de
termined when the liquid was at the same temperature as it
was to be during the titrations (e.g., 15.0, 25.0, or 35.0C).
Immediately prior to a titration, 0.1 ml of stock 3-amino-


53
acridine in ethanol was added to the flask, the contents
mixed and then transferred to a 4.00 cm quartz cell (Bolab
Inc., Derry, New Hampshire). The cell and its contents
were placed in the thermostated compartment of the
spectrophotometer (Beckman model DB-GT) and allowed to
become thermally equilibrated. The temperature of the
instrument's cell compartment was maintained at 15.0 +
0.2C, 25.0 + 0.05C or 35.0 + 0.2C by circulating thermo
stated water (Aminco refrigerated water bath) through its
external jacket.
The spectrum of the sample was recorded after which
an aliquot of stock DNA solution was added. Syringes (Uni
metrics, Inc.) were employed to dispense volumes of DNA
solution ranging from 4 yl to 100 ul while volumes greater
than 100 Ul (generally added toward the ends of the
titrations) were measured with a 100 microliter micro
pipette (Centaur Chemical, Danbury, Connecticut). The
individual culture tube of the stock DNA solution from
which aliquots were being withdrawn for a given titration
was maintained at 25C, regardless of the temperature at
which titration was being conducted. After adding the DNA
to the 3-aminoacridinium solution in the cell, the solution
was mixed by withdrawing about 5 ml of the cell's contents
into a serological pipet, then rapidly forcing them back
into the cell. This was done three or four times after
each addition of DNA.


54
Concentrations of the stock DNA solutions were
determined by measuring the absorbance of a solution
containing 50 ml of stock DNA in 10.00 ml of 0.010M
supporting electrolyte at pH 5.90 (1.00 cm pathlength).
The concentrations of at least ten samples, each taken
from separate tubes, were calculated using the molar
4 -1 i
absorptivity at 260 nm of 1.412 x 10 M cm (75). All
glassware was cleaned by immersion into concentrated
sulfuric acid saturated with potassium dichromate. Spec
trophotometer cells were rinsed with water followed by
ethanol and were stored in 50% concentrated hydrochloric
acid in ethanol to prevent accumulative absorption of the
dye or of DNA to their surfaces.
Solid 7-aminoquinoline was a gift from Dr. D. Jackson,
Texas Tech. University, Lubbock, Texas, and was determined
to be at least 98% pure according to molar absorptivity
values (82-84). Reported (pH 12) ^nm (log e): 240 (4.56),
max
275s (3.58), 335 (3.68), 346 (3.64). (pH 2.0): 260 (4.44),
285s (3.57), 392 (3.90). The reported pKa of the compound
_ 3
is 6.65 (84). A stock 3.4 x 10 M solution in absolute
ethanol was prepared and 30 yl aliquots added to 8.00 ml
of 0.10M KH2PO4 buffer solution at pH 5.9. The 7-amino-
quinolinium solutions were titrated at 15.0 + 0.2C, 25.0 +
0.05C, and 35.0 + 0.2C using the same equipment and pro
cedures as for the 3-aminoacridinium titrations, vide
supra.


55
Calculations
The molar absorptivity of the 3-aminoacridinium-DNA
complex, at 368 nm was calculated from absorbance
spectra of solutions in which essentially all of the
compound was in the bound form. Complete binding (>_ 99%
of all drug present) was assumed when at least three suc
cessive additions of excess DNA to the solution resulted
in no change in the total absorbance, corrected for
dilution. It was further presumed that, under these
conditions, all absorbance at 368 nm was due to the complex
and no 3-aminoacridine existed, either as the neutral or
monocationic, unbound form. Also, only solutions containing
monovalent metals (and tertamethylammonium cation) at low
ionicities were used since, for other systems, the large
amounts of DNA needed to drive the equilibrium toward
products caused serious light scattering and, in some
instances, precipitation of DNA salts. Calculated values
of in the presence of each type of electrolyte are
presented in Table 1. The averaged value of 8494^ M ^cm ^
was used for e in the presence of all supporting elec-
trolytes at all concentrations. A similar procedure was
used to evaluate the molar absorptivity of the 7-amino-
quinolinium-DNA complex, at its band maximum (405 nm).
At pH 5.9 the ratio of neutral to protonated 7-amino-
quinoline is 0.178. It was assumed that only the cationic
form of the drug binds to DNA. The averaged molar absorp
tivity of the 7-aminoquinolinium-DNA complex was evaluated


56
using three simultaneous equations and was found to be
5332^ M 'cm ^ at 405 nm.
Concentrations of free 3-aminoacridinium cation,
[BH], and bound monocation, [BHP], were calculated from
absorptiometric data assuming simple additivity of the
absorbances of the two species at the analytical wave
length
A = eBH£[BH] + ?BHptBHP] (2-1)
and considering the mass balance expression
Ct = [BH] + [BHP] (2-2)
where
A = the total absorbance of the solution at a given
wavelength.
e = the molar absorptivity of the free species,
.bn
1.413 x 103 M-1cm_1.
£BHP = mo^ar absorptivity of the bound complex,
8.49£ x 103 M ^cm ^.
H = pathlength of light in the cell.
= the total concentration of free and bound species.
[BH] = the molar concentration of free 3-aminoacridinium
cation.
[BHP] = the molar concentration of bound 3-aminoacridinium-
DNA complex.
Solving equation (2-2) for [BH], substituting into equation
(2-1), and rearranging yields


[BHP]
(2-3)
A-CBHCt*
£BHP£_eBH
Before any DNA has been added, the total absorbance of
the solution will be due entirely to free BH and, since
BH at this point will be equal to C^, the total absorbance
before any titration, Aq, may be expressed as
A0 = £BHCt (2-4)
Similarly, when all of the compound has been bound, the
final absorbance, A is equal to
A
00
eBHPCt*
(2-5)
and for any point in the titration, where the total absor
bance is Afc,
[BHP] = [(At-Ao)/(Aoo-Ao)]Ct (2-6)
and
[BH] = [l-(At-Ao)/(Aoo-AQ) ]Cfc (2-7)
Since an appreciable amount of the total 7-amino-
quinoline exists in the solution in the neutral form at
pH 5.9, the equations representing the protolytic equilibrium,


58
the conservation of mass for 7-aminoquinoline,
Ct = [A] + [AH] + [AHP] (2-9)
and the photometric absorbance at any point in the
titration of the small molecule with DNA,
At = £a[AU + eAH[AH]£ + £AHP[AHP]£ (2-10)
must be employed to determine the concentrations of the
species involved in the binding. Concentrations of drug
species are, explicitly,
[AHP] =
(eA-^- + £ah+> V -V1 + VihV
K
{eA~f7 + eAH eAHP(1 + Ka/[H+])}£
(2-11)
[H' ]
At ~ £AHP Ct£
[AH] =
(2-12)
K
{£ a
A[h+] + £AH £AHP(1 + Ka/[H+])}£
[A] =
(At GAHP Ct£) Ka[H+]
(2-13)
K
[eA~r + eAH eAHP
LH J
K
[H+]
)H
(1 +


59
It was also useful to directly calculate the ratio,
[AHP] (e
+ eAH+) V A(1 + Ka/[H+1)
(2-14)
[AH]
At £AHP Ct£
where
= dissociation constant for the drug or probe,
2.24 x 10~7 M-1.
[H+] = hydronium ion concentration, 1.26^ x 10 ^M.
eAHP = rao-*-ar absorptivity of the bound monocation at
the specified analytical wavelength, 5.332^ x
10"^ M "'"cm at 42 0 nm.
£AH = mo^-ar absorptivity of the free monocation at
the specified analytical wavelength, 7.943^ x
10^ M ^cm at 42 0 nm.
e = molar absorptivity of the free base at the
2 -1
specified analytical wavelength, 3.47 x 10 M
cm at 392 nm.
¡L = absorption cell path, 4.00 cm.
A^ = the absorbance at any point during the titration
at the specified analytical wavelength.
= total concentration of drug or probe.
For each point in a titration the total absorbance was
determined along with appropriate constants which were
substituted into the necessary relationships to yield the
concentrations of free and bound species. Raw data for


60
all titrations are presented in Tables 2 through 17. Note
that volume corrections for the addition of DNA solution
to the sample were made in all cases.
The equilibrium association constants for the
surface and intercalative modes of binding were calculated
for each point in the absorptiometric titrations of the
small molecules with DNA, using the relationships developed
by A.C. Capomacchia and S.G. Schulman (85).
(2-15)
Ks = [BHP]/[BH]([Pt/3]-[BHP])
and
2
(2-16)
KI = [BHP]/[BH]([Pfc/2]2[BHP])
where Kg and are the apparent association constants
for the surface and intercalative modes, respectively, and
P is the total molar DNA phosphate concentration. The
equations for the binding of 7-aminoquinolinium to DNA
are the same except [AHP] and [AH] are substituted for
[BHP] and [BH] respectively. The extended Debye-Hiickel equation
(1-28) was used to calculate activity coefficients at the
various ionic strengths, for charged species involved in
the association equilibrium.
Apparent association constants were corrected for
deviations of molar concentration from activity, by
(2-17)


61
where K1 is the corrected apparent association constant,
and a are activity coefficients for bound
BHP BH p
species, free species, and DNA phosphate, respectively.
The value of n is 1 for surface binding and 2 for inter-
calative binding. For the binding of 7-aminoquinolinium
to DNA, the subscripts BHP and BH are replaced with AHP
and AH corresponding to bound and free 7-aminoquinolinium,
respectively. These, and other calculations, were done
using a Litton-Monroe model Beta 326 Scientist computer
equipped with a tape cassette drive (Litton-Monroe,
Orange, New Jersey). The simultaneous solution of (2-8),
(2-9), and (2-10) to obtain concentrations of A, AH, and
AHP was accomplished using a program supplied by the
computer manufacturer. Programs for specific calculations
used here were designed and are presented in Appendix III.
Thermodynamic parameters for the binding of 3-amino-
acridinium and 7-aminoquinolinium to DNA were evaluated
using equation (1-31). Values of log K at 15C, 25C, and
35C for surface and for intercalative binding of each
species were determined. These values were then plotted
vs. 1/T (K). The slopes of the lines, equal to AH/2.303R,
were used to calculate the standard entalpies of binding.
Values for the standard entropies of the reactions were
evaluated using
AS = (AH AG)/T
(2-18)


CHAPTER III
RESULTS AND DISCUSSION
General Spectral and Titration Characteristics
A number of interesting spectral features and
titration characteristics are common to all of the
systems investigated. Space limitations preclude inclu
sion of all of the absorptiometric titration spectra,
though four which are representative are presented
(Figures 3 through 6). These are spectra of titrations
done in the presence of the background cation having the
lowest charge density, Cs+, and the greatest charge density,
Mg++, at the highest and lowest concentrations of each.
As a rule, the lower the charge and charge density, of
the background cation, the less effect it appears to have
on the two modes of binding. The highest concentration of
the smallest divalent cation represents the other extreme,
as it severely inhibits binding processes. The rest of
the systems are intermediate between these two. In each
of the figures, only 7 curves are drawn for the sake of
clarity, though usually between 18 and 22 spectra were
recorded for any given titration. For the curves ommitted
from figures 3 through 6, the absorbances at of the
62


63
1Lb bands may be obtained from the appropriate tables in
Appendix II.
In all cases, addition of DNA to solutions of 3-
aminoacridinium cation resulted in decreases (hypochromisms)
in the intensities of both the ^L, and bands. It is
D cl
noteworthy that even small amounts of DNA (< 1 mole DNA
per mole of drug) added to 3-aminoacridinium solutions
produced only hypochromic shifts. However, addition of
small amounts of DNA to other compounds such as 7-amino-
quinolinium causes an initial increase in the total intensity
(hyperchromism) of the drug's spectrum. Hypochromism sub
sequently occurs when DNA in excess of about 1 mole per mole
of drug is added. Presently, it is not known what factors
are responsible for the initial hyperchromism of some species'
spectra or why some exhibit initial hyperchromism while
others exhibit only hypochromism throughout their whole
titration, though significant variations in binding behavior
appear highly probable.
It may be seen from Figures 3 through 6 that addition
of increasing increments of DNA to 3-aminoacridinium results
in red (bathochromic) shifts of both bands of the drug's
spectrum. The bathochromism of the ^A and ^L.^' ^A
d b
bands are a result of decreases in the energy separations
between the ground and excited states. This may be due to
a lowering of the '*'L states, a raising of the ground, state,
or a combination of both. If binding of small molecules to
DNA occurs as the reaction proceeds, then it is reasonable


64
to hypothesize that the cationic drug is moving from a
highly polar, high dielectric medium (water) to a less
polar, more lipophilic solvent such as the region in, and
around, the DNA helix. An environmental change of this
type would tend to favor a decrease in charge separation
*
within a molecule and, thus, favor tt > tt transitions
in which the formal positive charge of the heterocyclic
ring nitrogen is decreased. Such a case would mean that
a lowering of the energies of the excited states is the
predominant factor in causing the red shift upon binding
if, indeed, such "solvent effect" arguments are valid when
considering binding processes. Caution must be exercised
in this regard as Capomacchia (69) has found that, for a
number of compounds, spectral shifts upon binding to DNA
are not in qualitative agreement with shifts observed when
the compounds were transferred from polar aqueous solvents
to less polar organic solvents. Maximum shifts of the two
bands are from 362 nm to 368 nm for the ^L, band and from
b
454 nm to 463 nm for the band. These differences (6 nm
a
and 9 nm, respectively) correspond to decreases in the
transition energies of 450^ cm ^ (1.29 Kcal/mole dye) and
428^ cm (1.23: Kcal/mole dye). The similar magnitudes show
that, with excess DNA at low total ionic strength, the two
transitions are equally affected from an energetic stand
point in the presence of excess DNA. However, a large
number of titrations have substantiated the observation
that the moves only slightly, from 362 nm to == 364 nm,


65
over most of the titration and then abruptly shifts from
= 364 nm to 368 nm while the shift of the band is
a
smooth throughout the whole titration. The abrupt shift
of the takes place at approximately the point in the
titration where any additional DNA results in no further
decrease in band intensity. In fact, this phenomena
has been accepted as evidence that all of the protonated
3-aminoacridine has been bound. Moreover, the rapid shift
of the higher energy band always occurs significantly
beyond the point in the titration where very little change
is noted in band intensity or position of the maximum of
the longer wavelength band. These observations suggest
that the environment appears to affect both transitions
equally in the presence of excess DNA, but they are affected
differently during the course of a titration. A cursory
consideration of the data may lead one to postulate that
the abrupt shift is due to intercalation of the molecule.
This is plausible since internal binding would affect the
^"L, more dramatically than the especially if the charged
ring nitrogen remained exterior to the helix. Further
consideration, however, reveals substantial evidence, both
in this work and in the literature, which shows that inter
calation occurs much earlier in the titration and that it is
much less abrupt a process than the bathochromic shift
considered here would require. The abrupt change might
arise from some form of near-neighbor interaction of bound
drug molecules (63). During most of the titration, where


66
the total DNA phosphate to total 3-aminoacridine ratio,
P/D, is low, the bound molecules can interact with one
another. But as the ratio increases, bound species become
isolated from each other to the point of no longer being
able to interact. The hypochromic shift is not incon
sistent with this proposal as both red and blue shifts
are possible upon exciton formation. Another factor to
be considered is that as the DNA concentration is increased
aggregation of the DNA, itself, may begin to occur. At
some critical level, the bound dye may experience a marked
decrease in the hydrophilicity of its environment, perhaps
as a result of the neighboring polymers excluding water
between themselves. Under these circumstances, red shifts
would be expected, though why they would be seen only for
the '*'L, and not for the transition is unanswered,
b a
The onset of the rapid red shift appears to be depen
dent on the total ionic strength of the medium as well as
the type of cations in solution. In the presence of
large, diffusely charged, monovalent cesium a distinct
bathochromism occurred, even in the highest concentration
(0.15M) at the moderate P/D ratio of 65. On the other
4 ++
hand, 6.3 x 10 N Mg was sufficient to prevent the shift
to 368 nm even when phosphate is in excess of drug by a
factor of 65. In 0.0025M CsI^PO^, a P/D ratio of only 14
results in the spectral shift (in fact, it probably occurs
at- about P/D = 11) Charge density of the countercations
influenced the point at which the abrupt red shift was seen


67
in a manner parallel to the degree of completion of a
titration for a given P/D ratio. The overall spectral
behavior of the systems in the presence of the alkaline
earths was the same as for the alkalies except the
maximum concentrations of M++ that allowed at least 90%
binding at P/D < 100 was much lower than for alkali
metals. Concentrations of group Ila cations in excess
of about 0.02N resulted in precipitates forming in the
samples. These were, presumably, the metal salts of DNA
and appeared long before sufficient polymer was added to
appreciably bind the drug.
Throughout this discussion, the absorbances at the
band maxima have been assumed to be directly related to
the areas under the curves. Since it is the areas, and
not peak heights at the maxima which are true measures of
transition probabilities, care must be taken in equating
peak heights with relative transition probabilities. If
the overall geometry of the band were to change during
the course of a titration, peak heights would not be valid
representations of the progress of the titration. In our
studies, inspection of the bands showed that their geome
tries were constant during the titrations, hence, we feel
approximating peak heights with peak areas was valid.
In addition, we have assumed that the 0-0 vibronic
band maximum of the ^L, < ^A transition is coincident
b
with the maximum of the total transition envelope. This
may be misleading if the maximum of the envelope arises


68
from the sums of the intensities of vibronic transitions
which are close enough together to appreciably overlap
at the maximum wavelength. If the vibronic energies are
altered in dissimilar ways during a titration, then the
maximal envelope intensity will reflect this dissimilar
change along with any real shift in the 0-0 band. It is
possible, therefore, that shifts of the envelope
maximum may be due to either a disproportionate change in
the magnitudes of the band maxima of the 0-0 and 0-1 vibronic
bands or a bona fide, pure 0-0 shift.
Variations in the spectral characteristics of the 3-
aminoacridinium-DNA systems in the presence of metal phos
phates and metal acetates were not due to the different
counteranions, as evidenced by the results of titrations
using KH2P0^ and KC^CCH^ as background electrolytes. Re
sults obtained with the two buffer ions were the same in
all respects except at very low ionic strengths. Deviations
in this region are probably due to lack of sufficient pH
control or alterations in the conformation of DNA possibly
arising from mild denaturation.
The envelope of the free 3-aminoacridinium cation
consists of two distinct vibronic bands having maxima at
348 nm and 362 nm, corresponding to the 0-1 and 0-0 vibronic
transitions, respectively. As a titration proceeds, the fine
structure of the envelope is lost as a result of the 348 nm
peak becoming increasingly less pronounced, to the point of
becoming a mere perturbation of the overall envelope. It is


69
difficult to determine the exact position of the shoulder
when the titration is nearly complete, but it appears that
its shift is essentially the same as that of the 362 nm
band. In cases where the abrupt 364 nm to 368 nm shift
of the vnn vibronic band of the ^L, occurs, there is an
accompanying shift of the shoulder maximum from = 350 nm to
about 354 nm with a slight, but definite increase in the
resolution of the two bands. The degree of fine structure
of a compound's spectrum provides qualitative information
regarding the environment of the species. Loss of fine
structure may be rationalized on the basis of increasing
the degrees of vibrational freedom of the molecule. If,
upon binding, the energy separation between the v = 0 and
v = 1 vibrational levels of the ground state were to be
reduced, a coalescing of the two peaks would be anticipated.
Such a situation may arise when a molecule moves from an
environment in which its vibrational motions are restricted
to one in which they are less so. For this to maintain in
our systems, the bound species would have to have greater
vibrational freedom than the wholly water-solvated free
cation. This is feasible considering the strong degree of
interaction between water molecules and any dipole in their
midst compared to the lipophilic, noninteracting "solvent"
of the surface and interior of DNA.
The increase in fine structure which accompanies the
dramatic red shift of the 0-0 band of the envelope at
high DNA concentrations is not well understood but serves


70
to corroborate evidence for a dramatic environmental change
of bound drug at these very high levels.
To summarize, titrations of 3-aminoacridinium cation
with DNA salts in the presence of monovalent and divalent
metal ions (including tetramethylammonium ion) of varying
concentrations result in hypochromism only. In all cases,
both the and envelopes of the free drug's spectrum
red shift, at least to some extent, during titration
the from 454 nm to 463 nm and the ^L, from 362 nm to
a b
about 364 nm. In the presence of all concentrations of
supporting electrolytes having large, monovalent cations,
a further, sharp red shift from 364 nm to 368 nm takes place
when a P/D ratio of 10 to 15 is reached. The shift may be
effected in the presence of smaller alkalimetal ions at
P/D ratios approaching 70. Regardless of the concentration
or charge density of divalent cations, the dramatic
shift did not occur. Loss of vibronic fine structure of
the envelope was observed in all cases as the titrations
progressed, up to the point at which the dramatic batho-
chromic shift occurred. Concomitant with the rapid shift
was a reemergence of the 0-1 vibronic band.
Using relationships developed in the Experimental
section, concentrations of free and bound 3-aminoacridinium
cation were computed from the total absorbances of solutions
which contained varying amounts of DNA. Data for all of
the titrations are not included here, though some, which
include the extremes of experimental conditions, are presented.


71
Tables 18 through 21 contain tabulations of [BH] and
[BHP] along with log ([BHP]/[BH]) values for reactions in
0.15M and 0.0025M Cs^PO^ and in 0.010N and 6.3 x 10
MgiC^CCH^^- In lower concentrations of alkali metals
and tetramethylammonium ion, the addition of as little
-3
as 5 yl of 9 x 10 M DNA phosphate solution results in a
significant degree of binding of the drug. For instance,
nearly 10% of the total 3-aminoacridinium present in
0.0025N CsI^PO^ is bound after addition of only 5.0 yl of
DNA. Since volumes of DNA less than 5 yl cannot be
measured with much reliability due to its viscosity, the
ratio of bound to free drug of 0.1 should be considered
as the effective minimum for titrations of this sort. When
lower ratios were possible, such as in solutions containing
divalent cations or >_ 0.010M monocations, random experi
mental error was sufficiently high to make the data unreliable.
As a result, our studies indicate that absorptiometric
titrations of small molecules with DNA, under conditions
similar to ours, are probably invalid whenever the ratio of
bound to total drug is less than 0.1. It should be recalled
that our use of 4.00 cm cells maximized sensitivity, thereby
allowing minimal increments of DNA solution to be used.
Smaller amounts of DNA could be added in 5 yl increments by
employing a more dilute stock solution, but the errors arising
from other sources would likely be unimproved. At the other
titration extreme, ratios of bound to free drug in excess of
about 0.95 also result in unacceptable randomness in most


72
cases. Our procedures are designed to minimize the impact
of these sources of imprecision on the calculated association
constants.
The maximum extent of total binding which could be
achieved varied with the character and concentration of
the supporting electrolyte. Table 22 lists the percent
of total drug bound to DNA after the final increment of
DNA was added in titrations done in various media. The
P/D values listed are the maximum that could be obtained
due to solubility and pH-control limitations. No effort
was made to determine how the drug was bound or what per
centage of the total was surface bound compared to inter-
calatively bound. As discussed previously, for those titra
tions in which an abrupt bathochromic shift of the band
was observed, the binding was considered to be complete
(100 + 2%). The averaged molar absorptivity, entJT_, determined
from these titrations was used to compute concentrations of
free and bound species for all systems. In light of the
experimental error incurred in evaluating and the
Dflr
simple fact that thermodynamic equilibrium demands that
there cannot be total binding, percentages greater than
about 96% are taken as indicative of complete binding. From
Table 22, it may be seen that at least 97% of the total 3-
aminoacridinium present is bound in solutions of < 0.010M
MI^PO^. To attain essentially complete binding in these
systems only about 10 to 20 times as much DNA phosphate as
total drug need be present (see P/D column of Table 22). As


73
expected, larger and larger ratios of DNA to drug are
necessary for complete binding as the concentration of
a given electrolyte is increased. Similarly, it becomes
increasingly more difficult to shift the equilibrium wholly
toward bound product in the order (of electrolyte cations):
(CH^J^N* £ Cs+ required for 97% binding in 0.0025N Cs+; 17-fold was needed
for 97% binding in 0.010M Cs; while, even with a 65-fold
excess of DNA, only 90% of the drug was bound in 0.15M
CsI^PO^. It was relatively easy to totally bind the 3-
aminoacridinium cation in low concentrations of monovalent
cations but very difficult to attain greater than 90%
binding in even the lowest concentrations of the divalent
4 ++
electrolytes. In 6.3 x 10 N Ba (the dication having
the lowest charge density of those investigated) only 90%
of the drug was bound at a P/D ratio of 47 and, for Mg++
at the same concentration, 90% was bound at P/D = 65.
Similarly, in 0.010N ca++ and 0.010N Mg++ the maximum
degrees of binding were 77% and 70%, respectively, with
P/D ratios of 103 and 104. The choices of the upper limits
of total phosphate concentrations reported here were de
pendent on the solubility of species in the concentrated
electrolytes and on pH considerations for the dilute systems
Visible precipitation or erratic spectral behavior were
evident in solutions having P/D values greater than those
listed in Table 22 for barium and magnesium. The maximum
concentration of the divalent electrolytes was set at 0.010N


74
since obvious precipitation occurred in solutions of
0.025N M well before the titrations were complete (i.e.,
at 30% to 50% binding). Acetates of the alkaline earths
were used instead of their phosphates due to the latter's
very low solubilities. Despite dialysis of DNA solutions
against pH 5.90 buffer, their pH levels were still ap
parently close to 7. Thus, addition of moderate volumes
of DNA to the very poor acetate buffer systems caused
excessive increases in the pH levels of the dilute buffer
solutions. Consequently, acetate solutions having more
-4
than approximately 5 x 10 M DNA phosphate had pH values
considerably in excess of 5.90.
Upon perusing the second column of Table 22, it becomes
quickly evident that the extent of binding is not predomi
nately and directly related to ionic strength. Total ionic
strengths were computed for the systems having the P/D
values listed using equation (1-29) and Program 4 of
Appendix III. Notice that while the ionic strengths of
0.0025M CsH2P04 and 6.3 x 10 4N BaiC^CCH^^ are the same,
the binding is complete in the former and only 90% complete
in the latter. This shows that the effect of the formal
charge on the cations of the electrolyte is greater than
predicted on the basis of the simple electrostatic effects
incorporated in the limiting Debye-Hiickel equation. That
the contribution of excess DNA to the total ionic strength
is not responsible for these large differences may be seen
by comparing 0.010N barium and 0.010N magnesium solutions


75
which both contain the same amount of DNA and have, there
fore, the same total ionic strengths. The extent of
binding in the former is 77% while in the latter it is
only 70%.
Rapid estimations of the values of m and q of equation
(1-21) may be obtained by selecting reasonable terms such
that mq = n (n is total phosphate per drug). It has been
shown that (85) in a medium containing 0.002M to 0.005M
Ki^PO^ (pH 5.9) the binding of 3-aminoacridinium cation to
DNA is best expressed by
, ->
BH + (P04)3
bh(po4) 3
(3-1)
for
the surface mode, and
BH+ + 2(P04)3
BH((P04)2 > 2~
(3-2)
for
the intercalative mode.
In equation (3-1),
three DNA
phosphates act as a single moiety, reacting with the small
molecular cation to form the surface-bound species. In
equation (3-2) 2 phosphates act as a single reacting
entity with two of these, in turn, interacting with one
cation to form the intercalated complex. The values of m
and q are 3 and 1, respectively, for the first case and 2
and 2 for the second. Substituting these into equation (1-21)
yields
Ks = [BHP]/[BH] ( [Pt/3]-[BHP])) (3-3)
and


76
Kj = [BHP]/([BH]([Pt/2]-2[BHP])2) (3-4)
Ks and are the apparent association constants for the
surface-bound and intercalatively-bound drug-DNA complexes,
respectively. These equations were tested to certify their
validity under our experimental conditions. Recall that
plots of log ( [BHP]/ [BH] ) vs. log ( [P^m] -q [BHP] ) yield lines
whose slopes are equal to q and whose intercepts are equal
to log K. Tables 18 through 21 include calculated values
of log ( [BHP]/[BH]), log([Pt/3]-[BHP]), and log([Pfc/2]-2[BHP])
for representative titrations. Figures 7 and 8 are plots
of log([BHP]/[BH]) vs. log([Pty3]-[BHP]) for the reactions
in the presence of 0.0025M Cs+ and 6.3 x 10 Mg++; and in
0.15M Cs+, 0.010N Mg++, and 0.010M Cs+, respectively.
Figures 9 and 10 are plots of log ( [BHP] / [BH] ) vs. logUP^.^]-
2 [BHP]) for the same systems. Three general features are
particularly significant. One, the shapes and slopes of
the curves are essentially the same whether log ([P^^] [BHP] )
or log( IP32[BHP]) is plotted against log([BHP]/[BH]).
(Calculated slopes are tabulated in Table 23.) Only in the
latter portions of the titrations done in the more dilute
alkali metal electrolytes does this rule begin to break down.
Two, except for 0.15M Cs^PO., all plots of the data in
volving the presence of monovalent cations exhibit definite
curvature the lower the concentration, and further down
a series of the periodic chart, the greater the curvature.
2+
Conversely, all concentrations of all M salts yield


77
essentially straight lines over the whole titration
region. Three, in those cases in which there is no
curvature throughout the titration range, the slopes of
the lines are all nearly 1 and parallel the initial
regions of the other titrations.
That the values of q evaluated from both types of
plots are much the same in most media is evidence of the
validity of this technique. It appears that going from
q = 1 to q = 2 or 3 does not substantially affect the
values of log(£Pt]-q[BHP]) except at the extrema of the
titrations. With our procedure, however, data very early
and very late in a titration need not be relied upon when
estimating lines, unlike, for instance, the Scatchard
method. For cases where the background electrolyte is
especially weak (i.e., low concentrations of cations having
low charge density) differences in the slopes of the two
plots become appreciable for the later portions of the
titrations. For example, the final slopes where 0.010M
Cs+ was employed as the background electrolyte are 1.68 and
1.88 for q = 2 and q = 1, respectively, and for 0.0025M Cs+,
the values are 1.71 and 2.16. Several explanations may be
advanced for this lack of agreement, including the fact
that the relatively short overall range of the amount of
DNA needed to complete the titrations magnifies the error
inherent at the extrema. Also, the high degree of curvature
makes the determination of the two separate straight line
regions more subjective. Perturbations of the reaction


78
system which are significant only in low ionic strength
media (e.g., mild polymer denaturation) would be expected
to affect both types of plots in the same manner. Re
gardless of the sources of the discrepancies, the pro
cedure is still useful, as both plots yield the same
integer values of q. Final selection of m and q is made
on the basis of the constancy of the computed values of
the equilibrium association constant (vide infra).
The presence of at least two distinct linear regions
of the plots for lower concentrations of alkali metals
strongly indicates two binding stoichiometries. Numerous
investigators, as outlined in the Introduction, have shown
the presence of two types of binding processes which likely
account for the different values of q. What has not been
previously reported is the apparent change in the ratio of
phosphate to 3-aminoacridinium cation on going from one
type of binding to the other as evidenced by these plots.
Also, as the ionic strength is increased, the extent of
intercalative binding apparently decreases. This is mani
fested in Figures 7 through 10 by a progressive trend
toward a wholly linear plot as the ionic strength increases.
That is, compare the shapes of the curves for systems con
taining 0.0025M, 0.010M, and 0.15M CsI^PO^. Similarly, only
. 2+
one stoichiometry is evident from the plots for M -containing
systems, which are linear throughout their entire titration
regions. Slopes of 1 + .1 for all of these, including that
for the reaction in the presence of 1 x 10 N Ba(02CCH3)2,


79
indicates that only surface binding occurs in the
presence of the divalent metal cations over the titration
range investigated.
Previously, it was brought out that, in the presence
of low concentrations of monovalent cations as background
electrolytes, plots of log([BHP]/[BH]) vs. log(-q[BHP])
exhibit two regions of linearity. During the early part of
a titration the slope of the lines is approximately 1 while
in the later region the slope is about 2. As the ionic
strength is increased, or as the charge density of the
cation of the electrolyte becomes greater, the slope of
the later region decreases, eventually becoming equal to
about 1. Slopes of the initial portions are essentially
unaffected by alterations in the supporting media in which
the titrations are conducted.
The ramifications of these relationships may be explored
by first defining the binding interaction in terms of con
ventional metal-ligand chemistry. We shall analogize the
cationic 3-aminoacridinium to the metal center and the DNA
phosphates to ligands which bind to the center. For clarity,
we shall further define two types of ligands, having indepen
dent and variable affinities for the cation centers; a weakly
bound ligand (corresponding to surface binding) which forms
a 1:1 complex, and a more strongly bound ligand (corresponding
to intercalative binding) which forms a 1:2 metal:ligand
complex. These definitions are based on the values of q for
surface-bound and intercalatively-bound ligand types being 1


80
and 2, respectively. Finally, we shall assume that the
ligands compete for metal binding sites such that if a
cation is bound in the surface mode, then it is unavailable
for intercalative binding, and vice versa.
During the initial stages of a titration there is a
large excess of 3-aminoacridinium over DNA phosphate or,
by analogy, an excess of metal over ligand. In such a
situation, there would be minimal competition for metal
sites and both forms of binding would be expected to occur
to their fullest extent, based on their respective association
constants. We may view the titration of the cation with DNA
ligand as a decrease in the metal center concentration as
the process proceeds. Addition of aliquots of DNA serves
to limit the number of sites available for binding, causing
the more weakly bound surface species to be affected first.
At this stage, the strongly bound intercalative species is
relatively unaffected. This argument presupposes that the
binding affinity of the intercalative mode is substantially
greater than that of the surface type. Hence, the first
region of the titration reflects surface binding of 3-
aminoacridinium to DNA. As the titration progresses, the
relative number of available binding sites becomes limited
to such an extent that the equilibrium involving the inter
calated complex is the one predominantly affected. Now the
plot of log([BHP]/[BH]) vs. log([Pt^m]-q[BHP]) is a reflection
of equilibrium concentrations governed by the stronger mode
which, in turn, results in the slope of the line being 2


81
instead of 1. Note that the difference between the
relative binding affinities of the two modes must be
sufficiently great to allow one to predominate over the
other in the initial and final portions of the titrations.
On the other hand, if they are too disparate, only one
would be seen throughout the experimentally accesible
titration range. Increasing the ionic strength of a
reaction medium causes electrostatic contributions to the
total free energy to become less exothermic for a process in
which oppositely charged moieties coalesce. Based on this,
decreases in the apparent association constants for both
surface and intercalative binding would be anticipated. It
is assumed, at this stage of the argument, that changes in
electrostatic free energies are the predominant affects
accompanying alterations in ionic strength. Extrapolation
of the straight lines of Figures 7 through 10 to the vertical
axis shows that, as the ionic strength is increased, both
Kg and do become less positive in agreement with the above
prediction. However, such extrapolations may be misleading
in the intercalative cases. Recall that, upon altering the
ionic strength, the slopes of the initial, surface binding
regions of the plots are essentially unchanged and that they
form a family of lines. This constancy of slope indicates
that there are no changes in the surface binding reaction
(aside from equilibrium shifts) as the ionic medium is
altered. However, two factors contribute to the decrease in
the intercept values of intercalative plots: (1) lower values


82
of log([BHP]/[BH]) for any given point along the abcissa
as ionic strength is increased, and (2) the decrease in
the slopes of the lines as the ionic strength is increased.
Two explanations for the change in slope may be advanced.
The value of q may be indeed changing from 2 to 1, for the
intercalation reaction, as the ionic strength is increased.
If this is true, then either the overall stoichiometry is
changing such that mq ^ 4 or, the stoichiometry is remaining
constant and m is becoming equal to 4 while q is decreasing
to 1. A more likely explanation can be advanced on the basis
of relative changes of K and K If K and KT do not change
in the same manner as the ionic strength is varied, then
what we may be observing is an apparent change in q for
intercalative binding when, in reality, the apparent change
is due to an overlapping of surface binding manifestations
with intercalative binding. Assume that the values of m and
q remain unchanged for each type of binding, then, the real
slopes of the two line segments will remain constant and
their point of intersection will be dependent on the ratio
of Ks to K. For example, in one extreme case there is
essentially no intercalative binding, K /K is very large,
b -L
and only the line having slope 1 will be observed. At the
other extreme, Ks/KI would be very small when only inter
calative binding occurs and a single line having slope 2
would be seen. If the system is perturbed in a manner which
caused K to decrease more rapidly than KT the point of
intersection of the two lines will reflect the change in


83
K /K In plots such as Figures 7 through 10, this would
S X
be manifested as a shift toward the right, or to less
negative log([BHP]/[BH]) values. The total number of
sites which are involved in surface and intercalative
binding are, of course, directly related to the magnitudes
of the equilibrium constants. In other words, an increase
in Kg/KI indicates that a greater percentage of the total
occupied sites are occupied by surface-bound species. The
ratio of K to K is seen to increase for the binding
S _L
reactions between 3-aminoacridinium and DNA in the presence
of alkali metal electrolytes as the ionic strength is raised'.
5
The following data are representative of Ks/K.j. x 10 at
ionic strengths of 0.0025, 0.0050, 0.025, 0.050, and 0.10(M),
respectively. With (CH^^NI^PO^ as supporting electrolyte:
0.44, 0.56, 1.28, 1.83, and 2.9; whereas with Lil^PO^ as
supporting electrolyte; 0.43, 10.7, 1.2, 5.0, and 10.7,
-5
respectively. Excepting the value of 10.7 x 10 for 0.005
LiH2P04,the ratios show that as the ionic strength is
increased, the total percentage of bound drug which is
surface-bound increases. Moreover, for any given ionic
strength, the ratio is greater when tetramethy1ammonium
dihydrogen phosphate is the electrolyte than when the
lithium salt is present. It is obvious from the decreases
in both Kg and K^. that increasing the dielectric of the
medium does not enhance surface binding relative to inter
calative binding. Rather, K does not decrease as rapidly
s
as Kj. The change in Kg/Kj may result from specific ion


Full Text
EFFECTS OF IONIC PERTURBATIONS AND METAL ION
COMPETITION ON THE BINDING OF A MODEL DRUG TO DNA
By
PETER F. EISENHARDT
A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1977

DEDICATION
To Dearest Joannie
who, more than any other,
deserves credit for any positive
attributes I may possess.

ACKNOWLEDGEMENTS
I am indebted to my supervisory committee, Dr. Stephen
G. Schulman, Chairman; Dr. B. Andresen; Dr. M. Battiste;
Dr. C.H. Becker; and Dr. R.H. Hammer for their time and
guidance in the preparation of this manuscript.
To three Professors, I owe special thanks. To Dr.
Schulman for the many hours of his time he has given to
share his prodigious scientific knowledge. To Dr. "C" whose
unaffected concern for the well-being of all with whom he is
associated, combined with his keen intellect, makes him with¬
out peer. And, to Dr. John Baxter for his fine example and
unflagging attempts to maintain high standards within the
academic community.
Space permits mentioning only a few of my friends to
whom I'm indebted: Mum, Dad, Clam, Burns, Sue, Flyman and
Ms. Fly, Teem, Merm, Tony, Roy, Lenny-whats-is-name, Edgar,
and Lysa Chancey Smith.
Robert, Ann, Robbie, and Sarah are four of my life's
mainstays.
Thanks, too, to Carolyn Grantham, expert cryptographer,
who deciphered my cruel first drafts so competently without
once betraying her graceful, elegant mien.
iii

TABLE OF CONTENTS
Page
ACKNOWLEDGEMENTS i;Li
LIST OF FIGURES vi
LIST OF TABLES x
ABSTRACT X1V
CHAPTER I - INTRODUCTION 1
Binding of Small Molecules to Biopolymers 2
Thermodynamics of Binding Reactions 34
Rationale for the Selection of the Model
System and Experimental Design 42
CHAPTER II - EXPERIMENTAL 47
Materials and Procedures 47
Calculations 55
CHAPTER III - RESULTS AND DISCUSSION 62
General Spectral and Titration Characteristics.. 62
Apparent Association Constants for Surface
and Intercalative Binding 84
Association Constants Based on Activities --
Extended Debye-Hückel Considerations 97
Evaluation of a Simple Competitive Binding
Model 102
Thermodynamics of the Binding of 3-amino-
acridinium and 7-aminoquinolinium to DNA 105
APPENDICES
APPENDIX I - FIGURES 120
IV

TABLE OF CONTENTS (Continued)
Page
APPENDIX II - TABLES 187
APPENDIX III - COMPUTER PROGRAMS 281
REFERENCES 29 5
BIOGRAPHICAL SKETCH 301
V

LIST OF FIGURES
Figure
Page
1
2
3
4
5
6
7
8
The structures of selected substituted
acridines and related compounds (Structures
I through XVI) 122
General form of Scatchard plots: R/C vs.
R for one and two classes of binding sites 128
Absorptiometric titration of 3-amino-
acridinium cation with DNA, cesium salt,
pH 5.90, 25.0°C. Background electrolyte:
CsH2P04' ÍHÍtial concentration, 0.15M 130
Absorptiometric titration of 3-amino-
acridinium cation with DNA, cesium salt,
pH 5.90, 25.0°C. Background electrolyte:
CsH2PO^, initial concentration, 0.0025M 132
Absorptiometric titration of 3-amino-
acridinium cation with DNA, magnesium salt,
pH 5.90, 25.0°C. Background electrolyte:
Mg(09CCH3)2, initial concentration, 0.010N 134
Absorptiometric titration of 3-amino-
acridinium cation with DNA, magnesium salt,
pH 5.90, 25.0°C. Background electrolyte:
MgiC^CCH^^' initial concentration, 0.00063N 136
Log([BHP]/[BH]) vs. log([Pt/3l-[BHP]) with
0.0025M CsH2P04 and 6.3 x 10-4N Mg(02CCH3)2,
as supporting electrolytes, pH 5.90, 25.0°C 138
Log ( [BHP] / [BH] ) vs. log ( [Pt/31 -.[BHP] ) with
0.15M CsH2P04, 0.010M CSH2P04, and 0.010N
Mg(09CCH9)2 as supporting electrolytes, pH
5.907 25.0°C ‘ 140
9 Log([BHP]/[BH]) vs. log([Pt/21-2[BHP]) with
0.0025M CsH2P04 and 6.3 x 10~4N Mg(C>2CCH3)2
as supporting electrolytes, pH 5.90, 25.0°C 142
vi

LIST OF FIGURES (Continued)
Figure
10
11
12
13
14
15
16
Log([BHP]/[BH]) vs. log([Pt/2^“2[BHpl) with
0.15M CsH2P04, 0.01M CsH2P04, and 0.010N
Mg(02CCH3)2 as supporting electrolytes, pH
5.90, 25.06C
Log of the apparent association constant,
Ks, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, I1'2. NaH2P04
and LÍH2PO4 as supporting electrolytes, pH
5.90, 25.0°C
Log of the apparent association constant,
K
s'
for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, I-1-/2. KH2PO4
and KO2CCH3 as supporting electrolytes,
pH 5.90, 25.0°C
Log of the apparent association constant,
K
s'
for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, 1^/2 . KH2P04
and RbH2P04 as supporting electrolytes,
pH 5.90, 25.0°C
Log of the apparent association constant,
Ks, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, ll/2 . CsH2P04
and (CH3)4NH2PC>4 as supporting electrolytes,
pH 5.90, 25.0°C
Log of the apparent association constant,
Ks, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the ionic strength, li/2. Mg(02CCH3)2
and Ca(02CCH3)2 as supporting electrolytes,
pH 5.90, 25.0°C
Log of the apparent association constant,
Kg, for the surface binding of 3-amino-
acridiniura cation to DNA vs. the square
root of the ionic strength, I1/2. Sr(02CCH3)2
and Ba(02CCH3)2 as supporting electrolytes,
pH 5.90, 25.0°C
Page
144
146
148
150
152
154
156
Vll

LIST OF FIGURES (Continued)
Figure
17
18
19
20
21
22
23
Log of the apparent association constant,
Ki, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I-*-/ .
LÍH2PO4 and NaH2P04 as supporting
electrolytes, pH 5.90, 25.0°C
Log of the apparent association constant,
Kj, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I1/2.
KH2PO4 and KO2CCH3 as supporting
electrolytes, pH 5.90, 25.0°C
Log of the apparent association constant,
Kj, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I1/2.
KH2PO4 and RbH2PC>4 as supporting
electrolytes, pH 5.90, 25.0°C
Log of the apparent association constant,
Kj, for the intercalative binding of 3-
aminoacridinium cation to DNA vs. the
square root of the ionic strength, I1/2.
CSH2PO4 and (CH3)4NH2PO4 as supporting
electrolytes, pH 5.90, 25.0°C
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, ll/2.
LÍH2PO4, NaH2PC>4, and KH2PO4 as supporting
electrolytes, pH 5.90, 25.0OC
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, I1/2.
RbH2P04, CsH2P04, and (CH3)4NH2P04 as
supporting electrolytes, pH 5.90, 25.0°C
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, ll/2.
Mg(02CCH3)2 and Ca(02CCH3)2 as supporting
electrolytes, pH 5.90, 25.0°C
Page
158
160
162
164
166
168
170
viii

LIST OF FIGURES (Continued)
Figure
24
25
26
27
28
29
30
31
Log of the corrected association constant,
Kg, for the surface binding of 3-amino-
acridinium cation to DNA vs. the square
root of the activity of solution, I^-/2.
Sr(O2CCH3)2 and Ba(02CCH3)2 as supporting
electrolytes, pH 5.90, 25.0°C
Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium
salt, pH 5.90, 25.0°C. Background
electrolyte: KH2PO4, initial concentra¬
tion, 0.0050M
Log ( [AHP] / [AH] ) vs. log ( [Pj-/ 3 ] - [AHP] ) for
the binding of 7-aminoquinolinium cation
to DNA at 15.0°C, 25.0OC, and 35.0°C, pH
5.90, 0.010M KH2PO4 as supporting electro¬
lyte
Log([AHP]/[AH]) vs. log([Pt/2l“2[AHP]) for
the binding of 7-aminoquinolinium cation
to DNA at 15.0°C, 25.0°C, and 35.0°C, pH
5.90, 0.010M KH2PO4 as supporting electro¬
lyte
Log K¿ vs. 1/T(°K) for the reaction between
3-aminoacridinium cation and DNA at 15.0°C,
25.0°C, and 35.0°C, pH 5.90, 0.010M KH2P04
as supporting electrolyte
* o
Log Ks vs. 1/T( K) for the reaction between
7-aminoquinolinium cation and DNA at 15.0°C,
25.0°C, and 35.0OC, pH 5.90, 0.010M KH2P04
as supporting electrolyte
I
Log Kg vs. 1/T(°K) for the reaction between
3-aminoacridinium cation and DNA at 15.0°C,
25.0°C, and 35.0OC, pH 5.90, 0.010M KH2P04
as supporting electrolyte
* o
Log Ks vs. 1/T( K) for the reaction between
7-aminoquinolinium cation and DNA at 15.0°C.
25.0°C, and 35.0°C, pH 5.90, 0.010M KH2P04
as supporting electrolyte
Page
172
174
176
178
180
182
184
186
IX

LIST OF TABLES
Table
1
2
3
4
5
6
7
8
Molar absorptivities of the 3-amino-
acridiniura-DNA complex and the 7-amino-
quinolinium-DNA complex in various
supporting electrolytes
Absorptiometric titration of 3-amino-
acridinium cation with DNA, lithium salt.
Lithium dihydrogen phosphate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, sodium salt.
Sodium dihydrogen phosphate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium
salt. Potassium dihydrogen phosphate as
supporting electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium
salt. Potassium acetate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, rubidium
salt. Rubidium dihydrogen phosphate as
supporting electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, cesium salt.
Cesium dihydrogen phosphate as supporting
electrolyte
Absorptiometric titration of 3-amino-
acridinium cation with DNA, tetramethyl-
ammonium salt. Tetramethylammonium
dihydrogen phosphate as supporting
electrolyte
Page
188
189
194
198
202
207
212
216
X

LIST OF TABLES (Continued)
Table Page
9 Absorptiometric titration of 3-amino-
acridinium cation with DNA, magnesium salt.
Magnesium acetate as supporting electrolyte 222
10 Absorptiometric titration of 3-amino-
acridinium cation with DNA, calcium salt.
Calcium acetate as supporting electrolyte 226
11 Absorptiometric titration of 3-amino-
acridinium cation with DNA, strontium salt.
Strontium acetate as supporting electrolyte 230
12 Absorptiometric titration of 3-amino-
acridinium cation with DNA, barium salt.
Barium acetate as supporting electrolyte 234
13 Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium salt,
at 15°C 239
14 Absorptiometric titration of 3-amino-
acridinium cation with DNA, potassium salt,
at 35°C 241
15 Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium salt,
at 15°C 242
16 Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium salt,
at 25°C 245
17 Absorptiometric titration of 7-amino-
quinolinium cation with DNA, potassium salt,
at 35°C 247
18 Calculated values of [BH], [BHP], "Ks", and
related data for the reaction between 3-
aminoacridinium cation and DNA in 0.15M
CsH2P04 249
19 Calculated values of [BH], [BHP], Ks, Ki, and
related data for the reaction between 3-
aminoacridinium cation and DNA in 0.0025M
CsH2P04 251
xi

LIST OF TABLES (Continued)
Table
20
21
22
Calculated values of [BH], [BHP],
and related data for the reaction
aminoacridinium cation and DNA in
Mg(02CCH3)
Ks, "Kl",
between 3-
0.010M
Calculated values of [BH], [BHP],
and related data for the reaction
aminoacridinium cation and DNA in
Mg(02CCH3)2
Ks, "Ki",
between 3-
6.3 x 10"4N
Percent total 3-aminoacridinium bound to DNA
in the presence of various electrolytes after
addition of excess DNA
23 Initial and final slopes of plots of log([BHP]/
[BH]) vs. log([Pt/m]-q[BHP]) in various con¬
centrations of CsH2PO^ and Mg(02CCH3)2
24 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj,
binding of 3-aminoacridinium to DNA, LiH^O^
as supporting electrolyte
25 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind¬
ing of 3-aminoacridinium to DNA, NaH2PO. as
supporting electrolyte
26 Apparent equilibrium association constants
for surface, K , and intercalative, K , bind¬
ing of 3-aminoacridinium to DNA, KH^O^ as
supporting electrolyte
27 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind¬
ing of 3-aminoacridinium to DNA, K02CCH3 as
supporting electrolyte
28 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind¬
ing of 3-aminoacridinium to DNA, RbH2PO. as
supporting electrolyte
29 Apparent equilibrium association constants
for surface, Ks, and intercalative, Kj, bind¬
ing of 3-aminoacridinium to DNA, CsH2P04 as
supporting electrolyte
Page
253
255
257
258
259
261
262
263
265
267
Xll

LIST OF TABLES (Continued)
Table Page
30 Apparent equilibrium association constants
for the surface, Ks, and intercalative, Kt,
binding of 3-aminoacridinium to DNA, (CH^)^-
Ni^PO^ as supporting electrolyte 269
31 Apparent equilibrium association constants,
Ks, for surface binding of 3-amino-
acridinium to DNA, MgiO^CCH^^ and CaiC^CCI^^
as supporting electrolytes 271
32 Apparent equilibrium association constants,
K , for the surface binding of 3-amino-
acridinium to DNA, Sr (0200113)2 and Ba (0200113)2
as supporting electrolytes 273
33 Logarithms of the apparent surface binding
association constants for the binding of 3-
aminoacridinium cation to DNA, corrected
for ion activities 275
34 Equilibrium association constants for the
binding of 3-aminoacridinium to DNA, Ks,
and for the binding of alkali and alkaline
earth metal cations to DNA, K,. 277
M
35 Apparent association constants for the
surface, Ks, and intercalative, Kj, binding
of 3-aminoacridinium and 7-aminoquinolinium
to DNA at 15.0°C, 25.0°C, and 35.0°C 279
36 Thermodynamic parameters for the surface
and intercalative binding reactions between
3-aminoacridinium and 7-aminoquinolinium
cations to DNA 280
xiii

Abstract of Dissertation Presented to the
Graduate Council of the University of Florida in
Partial Fulfillment of the Requirements for the
Degree of Doctor of Philosophy
EFFECTS OF IONIC PERTURBATIONS AND METAL ION
COMPETITION ON THE BINDING OF A MODEL DRUG TO DNA
By
Peter F. Eisenhardt
August, 1977
Chairman: Stephen G. Schulman
Major Department: Pharmaceutical Chemistry
Absorptiometric titrations of a model drug compound,
3-aminoacridinium cation, with salts of calf thymus DNA
were carried out at pH 5.90 in a series of solutions in
which the supporting electrolyte was the dihydrogen
‘phosphate salt of one of the following: lithium, sodium,
potassium, rubidium, cesium, or tetramethylammonium ion.
Titrations were also conducted in solutions of the acetate
salts of magnesium, calcium, strontium, and barium. The
ionic strengths of the solutions were varied between 0.25
and 0.00063(M) for each alkali metal electrolyte and between
0.036 and 0.00010 for the alkaline earth electrolytes.
Apparent equilibrium association constants for the surface
xiv

binding and for the intercalative binding of 3-amino-
acridinium cation to DNA were computed for each concen¬
tration of each alkali metal electrolyte solution and
are reported. Apparent constants for the surface binding
of 3-aminoacridinium in the presence of alkaline earths
are also presented. Values for the intercalative binding
mode in the presence of group IIA metal cations could not
be determined.
The extent of binding of the model drug to DNA was
dependent on the nature and concentration of the electrolyte
Essentially complete binding was attainable for all concen¬
trations of all alkali metals while as little as 70% binding
was possible in the presence of 0.010N magnesium. Spectral
characteristics of the drug-DNA complex were also dependent
upon the nature of the background electrolyte and are dis¬
cussed.
The reaction systems did not behave in accordance with
limiting Debye-Hiickel relationships over the range of ionic
strength investigated. After applying extended Debye-
Hiickel corrections, the relationships between the surface¬
binding equilibrium association constants and the ionic
strength indicated that specific-ion competition between
metal and drug cations, for DNA phosphate sites, was signifi
cant. A simple 1:1 competitive model was evaluated.
At ionic strengths less than about 0.010, experimental
precision was poor, presumably due to alterations in the
conformation of DNA at low ionicities, suggesting that

evaluations of interactions between small molecules and
DNA be conducted in solutions of ionic strength in excess
of 0.010.
Thermodynamic parameters for surface and intercalative
binding of 3-aminoacridinium cation and its benzolog, 7-
aminoquinolinium, were evaluated from equilibrium association
constants at 15.0°C, 25.0°C, and 35.0°C. The standard
enthalpy change, AH°, for the 3-aminoacridinium binding
was less negative than AH° for the 7-aminoquinolinium process,
while the standard entropy change, AS°, for the former was
more positive than for the latter. Thus, the greater degree
of disordering resulting from the relaxation of the solvent
around the tricyclic 3-aminoacridinium cation as it binds to
DNA appears responsible for the more negative standard free
energy change for the 3-aminoacridinium binding process.
Values of AS° for the intercalative binding of the two drugs
suggest that significant ordering of the polymer helix occurs
upon intercalation.
xvi

CHAPTER I
INTRODUCTION
The interaction of small molecules with DNA and
biopolymers, in general, has been the subject of intense
investigation during the past several decades. Interest
in this area has been due, in large part, to the early
realization that such interactions may result in muta¬
genesis (1) and carcinogenesis (2), especially when the
small molecules are cationic aromatics such as substituted
acridines. Numerous publications deal with various aspects
of the interactions of cationic aromatics with DNA, some
of which will be briefly reviewed here. The specific
systems to be considered in this work include 3-amino-
acridine (I) . Hence, a summary of representative and
pertinent research involving aminoacridines is included.
Moreover, we have investigated the effects of metal ions
on the interactions of 3-aminoacridine with DNA and,
therefore, shall review the past research efforts in this
area. It is emphasized that this review is, by no means,
Roman numerals in parentheses after chemical names refer to
their listings in Figure 1 of appendix I.

2
exhaustive, but rather, is intended to provide perspective
and an introduction to the work that is included in this
report.
Binding of Small Molecules to Biopolymers
Probably the first acknowledgement of a significant
interaction between a member of the acridine chemical
family and a biological system was reported in 1913 by
Ehrlich and Benda (3) when they reported their observation
on the antibacterial activity of 3,6-diamino-10-methyl-
acridinium chloride(II) against Trypanosonoma brucei.
Subsequently, evidence accumulated that the antibacterial
activity of substituted aminoacridines is dependent on
their capacity to compete with hydronium ion for a vital
anionic site in a bacterium (4-6, 7: Ch. 2) and that
antibacterial activity and mammalian toxicity were not
directly related (8,9). The antibacterial action of
proflavine(III) is reported to be due to its ability to
inhibit DNA-dependent RNA synthesis (10) which suggests
that DNA may contain the anionic site. However, it is
inadvisable, at this time, to make a direct correlation
since proflavine also binds strongly to a variety of other
cellular structures, any one of which may contain the
vital anionic site (11).
Numerous amino-substituted acridines are known to be
carcinogenic while few simple acridines are cancer inducing
(12) which may be due to the former's greater ability to

3
intercalate into DNA (13). It is worthy of note that
benzacridines(IV) and dibenzacridines(V), first exten¬
sively studied in 1935, are nearly all highly carcino¬
genic (2). Two possible modes of action have been
proposed. The benzacridine may be metabolically activated
to a cationic form which then reacts as an electrophile
toward nucleophilic centers of DNA, RNA, and protein (14),
or the unmetabolized species may interact physically with
the DNA (15), perhaps by intercalating into the double
stranded helix.
Mutagenesis, presumably arising from the intercalation
of acridines with nucleic acids, was first reported by
DeMars in 1953 ( 1 ) in relation to his work on proflavine
and T2 bacteriophage (Escherichia coli). The mutagenic
activity of proflavine does not involve its integral
incorporation into the newly synthesized polynucleotide
chain but rather, results from the acridine's capacity to
cause deletions and insertions of nucleotides into the
DNA polynucleotides (16,17). In fact, proflavines' capacity,
in this regard, was used in 1961 to discover the triplet
nature of the genetic code.
In searching for effective anticancer and antitumor
agents, advantage is taken of the fact that compounds
which may interfere with cell replication at the molecular
level may have medicinar utility. As a result, a large
number of acridine derivatives have been investigated, only
to find that, while a select few do have antineoplastic

4
properties, their associated toxicities to the host make
them less desirable than other available compounds. On
the other hand, aminoacridines have enjoyed success as
antimalarial agents and are still in limited use today.
Investigations were begun in the 1920's for a substitute
for quinine, as an antimalarial, which led scientists
first to methylene blue, a phenothiazine dye, then to
pamaquine, both of which had low effectiveness against
acute stages of malaria. Structural similarities in the
above two compounds, with substituted acridines, eventually
led investigators to quinacrine(VI) which became the drug
of choice throughout World War II. Later, the more effec¬
tive and less toxic chloroquine(VII) was discovered.
Presently, considerable synthetic research is devoted to
acridine-derived antimalarials to supplant chloroquine
and to combat chloroquine-resistant P. falciparum —
discovered in 1961 (18). To date, the most promising
compounds of this type have a diamine side chain in the 9-
position(VIII). Structure-activity relationships for a
variety of quinacrine analogues having side chain variants,
terminal amine variants, and aromatic substituent variants,
as well as quinine related acridine antimalarials, are
tabulated by Henry (19).
In light of the broad biological significance of the
binding of aminoacridines to DNA and other biopolymers,
in vitro and in vivo studies aimed at understanding the
exact nature of the interactions have been undertaken.

5
The following paragraphs will briefly review several of
the more important articles germane to our investigation.
Those dealing with the modes of binding will be outlined
first, followed by studies of the effects of the ionic
medium (including possibly competing ions), with a review
of papers involving theoretical discussions, third.
An excellent comprehensive article on the interactions
of aminoacridines with nucleic acids and the methods by
which they may be studied was published in 1956 by A.R.
Peacocke and J.N.H. Skerrett (20). Another, in 1968, by
Blake and Peacocke (21) considered developments in the
field after 1956.
Aminoacridines have long been known to bind to
double-stranded DNA via two distinct modes — termed type
I and type II binding. Type I is the stronger and was
first proposed by Lerman (22), to be an internal or
intercalative process. Insertion of a molecular ion is
accomplished by transient rotation of base pairs allowing
the flat, lipophilic, aromatic portion of the molecule to
become situated between them. The charged portion of the
molecule was considered to remain external to the helix.
H.J. Li and D.M. Crothers (23) corroborated this model
when they presented indirect evidence that the base pairs
remain perpendicular to the helix axis upon intercalation
of the molecule. Another model for type I binding was
advanced by N.J. Pritchard, A. Blake, and A.R. Peacocke (24)
in which the small molecule is intercalated between adjacent

6
bases on the same polynucleotide chain. In both of these
models, the cationic heteroatom is associated with an
indeterminate number of anionic phosphate moieties along
the DNA backbone. The viscosity of heat-denatured DNA
is essentially unaffected by binding of acridines (25)
whereas the viscosity of native double-stranded DNA is
increased. An increase in viscosity is indicative of a
lengthening of the polymeric chain such as would result
from intercalation between base pairs or between adjacent
bases. Since the single-stranded DNA undergoes no vis¬
cosity change, the Lerman model appears more plausible.
Some specific geometrical parameters of substituted
acridine-DNA intercalative complexes were established by
G.R. Kelly and T. Kurucsev (26) through the evaluation of
linear dichroic spectra of stretched DNA films. An advan¬
tage of films over conventional methods involving aqueous
solutions is that they do not require large ratios of drug
to DNA phosphate. Taking advantage of the dichroism of
two mutually perpendicular polarized transitions of the
9-aminoacridine nucleus, the workers were able to assess
the angles of tilt and twist of the plane of the dye
relative to the polymer helix. The long axis of acri-
flavine(II) was essentially perpendicular to the helix
axis (85 + 5°), proflavine(III) was tilted slightly (77 +
2 1/2°), while 9-aminoacridine(IX) was tilted significantly
(67 + 3°).

7
Type II, surface, or external binding has come to
be generally accepted as an association of the small
molecules with the exterior surface of the DNA polyanion.
This interaction is, consequently, more strongly affected
by alterations of the solvent system than is type I
binding. Li and Crothers (23) determined, on the basis
of temperature-jump kinetic data, that proflavine binds
both internally and externally to calf thymus DNA and
only about 7% of the total amount bound binds externally
in the presence of 0.2M Na+ (sodium phosphate - sodium
nitrate). Decreasing the sodium ion concentration to
0.02M results in about 30% of the total being externally
bound.
A careful study of the intercalative binding of a
series of aminoacridines was conducted by Drummond and
co-workers (27) . Binding decreases in affinity in the
order: acranil (X) >neomonacrin (XI) >atebrin (VI) ^-amino-
acridine (IX)>proflavine(III) 9-amino-l,2,3,4-tetra-
hydroacridine(XII). This sequence shows that side chain,
and ring, substituents are not predominant in influencing
binding affinity. The effect of greatly decreasing the
planar area of 9-aminoacridine by hydrogenation is
evident upon comparing its position with that of XII. The
authors discuss various geometries of binding and conclude
it is not necessary that there be exact and complete
intercalation for type I (strong) binding. Rather, a
modified Lerman model in which the charged heterocyclic
ring nitrogen interacts with a DNA phosphate while the

rings only partially interact with the bases is more in
accord with observed behavior.
It has been shown that the degree of intercalation
may be dependent on the base composition of the DNA.
Heterogeneity of intercalative sites was demonstrated
by J.C. Thornes, G. Weill, and M. Daune (28) using
fluorescence quenching techniques. They demonstrated
that between 2% and 3% of the base pairs form sites where
proflavine is strongly bound while the remainder have
binding constants 3 to 4 times weaker. Proflavine
fluorescence is relatively unaffected in the former case,
but is totally quenched in the latter. The stronger
sites correspond to adenine-thymine rich regions of the
DNA indicating proflavine's specificity for the base
pair. Similarly, Ramstein and Leng (29) showed that the
location of bound proflavine, within the DNA, was dependent
upon the base composition. The absorptiometric and flúori-
metric titrations of DNA and partially methylated DNA
7
( 18%, mainly on N of the guanine residues) indicated
acridine was intercalated in both the native and methylated
sites. However, greater apparent equilibrium constants for
the methylated DNA than those for the native type were ob-
4 -1
served at two ionic strengths: 14.25 x 10 M and 43.5 x
10^M in 0.1M NaCl and 40 x 104M ^ and 83.5 x 10^M ^ in
0.01M NaCl, respectively. These differences may be due to
an increase in the distance between base pairs arising from
the presence of the methyl groups which may, in turn, result

9
in the shift in the equilibrium concentrations. S.
Georghiou (30) obtained much the same results as Ramstein
and Leng based on the fluorescence decay of proflavine
when it was bound to DNA at a phosphate to drug ratio of
420. Gabbay and co-workers (31) invoked the "Bookmark"
model of intercalation to postulate ten possible distinct
sites of binding of N-substituted-N-methylphenanthrolinium
cations(XIII) to nucleic acids. They also investigated
the topography of some nucleic acids in solution using
PMR (32).
Binding models for the interactions of acridine
orange(XIV) and proflavine(III) with DNA at ionic strengths
of 0.002M, 0.020M, and 0.20M (tetramethylammonium caco-
dylate buffer, pH 6.5, 22°C) were proposed based on
thermodynamic, spectroscopic,and hydrodynamic properties
of the systems (33). The methods precluded consideration
of surface binding. Binding isotherms were generated from
equilibrium association constants which were, in turn,
obtained from theoretical mass action relationships.
Additional information, useful in developing binding
models, was afforded by Goswami, Das, and Das Gupta (34)
who measured the decrease in the static dielectric con¬
stant of a proflavine-DNA complex while decreasing the
DNA phosphate/proflavine (P/D) ratio. The authors con¬
firmed that: the bound dye neutralizes the charge on the
DNA polyanion; at high P/D values, intercalation predomi¬
nates and; that when P/D becomes small (£10), a cooperative

10
increase in the electrostatic binding of phosphate groups
brings about a rapid decrease in the total surface charge
of the polymer. In brief, at P/D values between 10 and
100 most of the cations are intercalated with a concomi¬
tant elongation of the polyanionic DNA helix, whereas
below P/D of 10, surface binding is predominant. In the
latter case, the spatial extension of the helix is de¬
creased due to charge neutralization with no accompanying
expansion due to intercalation. These results indicate
that alterations in the dielectric constant of a medium
containing such a system would manifest dissimilar effects
on the two equilibrium association constants.
The specific effects of cations, other than the dye
species, on the reaction between DNA and a dye is an
integral part of this investigation. Therefore, a brief
consideration of articles dealing with the binding of
metal cations to DNA and related topics is in order. A
tabulation of sites and thermodynamic quantities associated
with proton and metal ion interactions with DNA, RNA, and
their constituent bases, nucleosides, and nucleotides has
been published by Izatt et al., (35).. H. Sigel and D.B.
McCormick (36) provide an introduction to the fundamental
theories concerning metal-ligand interactions in biological
systems, including discussions of the Irving-Williams
series, competition of metals and protons for sites, the
formation of ternary complexes, and effects of solvent
polarity and ionic strength. The Pearson theory (37) of

11
hard and soft acids and bases is applied to metals and
polymers in relation to the biological roles of Na+, K+,
Ca++, and Mg++ and to the carcinogenicity of heavy metals
(38) .
The relative effectiveness of ions in altering the
conformation of macromolecules via specific ion inter¬
actions is termed the Hofmeister, or lyotropic, effect
(39). Generalizations based on observations of simple,
ideal polymers may be confidently extrapolated to more
complex polymers since it has been shown that the
Hofmeister effect of a given ion is remarkably independent
of the nature of the macromolecule. Moreover, while
changes in the entropies of polymer segments are affected
by both anions and cations, as are changes in the partial
molal volume, the partial molal internal energy of a segment
is dependent only on cations. Considering the above,
M.J. Hey, J.M. Clough, and D.J. Taylor (39) reported that
the ion binding strength with protein is (in decreasing
order of strength) NH4+ This series may be useful in separating solvent effects
from ion binding effects where the above cations are
present along with 3-aminoacridinium cation and DNA.
Stabilizing effects of various ions on four biopoly¬
mers, including DNA, were evaluated by Hippel and Schleich
(40). The effects were discussed in terms of free energy
of transfer for model compounds, activity coefficient
variations, and correlations between ion effects on

12
macromolecules and on water structure. The relative
effectiveness of the ions studied, in increasing the
stability of the native configuration of DNA, paralleled
the Hofmeister series: (CH^)^N+>K+>Na+>Li+ and Cl , Br >
CH-.COO >C10. >CNS . It was found for a series of
3 4
tetraalkylammonium salts, the greater the length of
the alkyl chains, the lower the stabilization of the
native form of DNA.
Hen and Strauss (41) employed equilibrium dialysis,
dilatometry, and viscometry to assess the counterion
binding of a series of cations with poly(vinylsulfonate).
The Hofmeister series was maintained (in order of de¬
creasing binding affinity): Ag+>K+>Na+>H+,Li+>(CH^)^N+
and Ba++>Mg++. All of the cations except H+ and Li+
are capable of cross-linking the polyanion through
simultaneous site binding at two sulfonate groups.
The specific interactions of univalent and di¬
valent cations with calf thymus DNA can be correlated
with their respective unhydrated ionic radii (42,43).
The order of most weakly bound to most strongly bound
is: (CH^)^N+ binding parameters with DNA with polyphosphate as polyanion,
the authors showed that the sites on the biopolymer were
the phosphates rather than bases or base pairs. These
results were in agreement with those of other workers
(44,45).

13
Conductometric titrations of DNA with divalent
cations were first performed by J. Shack, R.J. Jenkins,
and J.M. Thompsett (46) who found a sharp end point cor¬
responding to 0.8 equilvalents of Mg++ and Ca++ per mole
of polymer phosphate. Similar titration behavior was seen
for poly A and poly U by Felsenfeld and Huang (47) who
developed an improved binding model which showed that the
apparent stoichiometry of 0.8 was in error and that the
true equivalence point corresponded to 1.0 eq of Mg++ or
Ca++ per mole of DNA phosphate. The authors also disproved
the previously held belief (48) that dications are more
strongly bound to purines than pyrimidines in a poly¬
nucleotide .
It has been reported (49) that there is binding to
base pairs of the DNA helix as well as to phosphates. The
investigators reported that the effects of adding Mg++ to
heat-denatured calf thymus DNA was to produce an absorption
spectrum similar to that of denatured DNA alone, except
for reduced absorbance. However, when Mg++ was added to
native DNA, which was then denatured, a spectral shift to
longer wavelength occurred. The authors concluded, on
these grounds, that the metal was binding only to phos¬
phates in the former case while in the latter, there was
interaction with the aromatic nitrogens of the bases. The
bonding between Mg++ and the nitrogens is presumably not
able to occur unless the hydrogen bonds between the bases
are broken in the presence of the metal.

14
Mathematical relationships incorporating changes in
activities of ions in solution and deviations in activity
coefficients were developed by Lyons and Kotin (45) to
qualitatively distinguish between specific (site) binding
and non-specific (diffuse electrostatic) binding of
metals to polyanions. The rule of additivity of the
activities of countercations of solutions of polyelectro¬
lytes and simple salts was re-examined. It was found
that, under certain conditions, there is significant
exchange between the counterions associated with the
polyanion and those of the bulk solution. A decrease in
the activity coefficient of sodium ion, upon dilution of
the system, was observed, indicating increased binding at
lower concentrations. In the case of magnesium the results
were not as simple but suggested that at high magnesium
ion concentrations, specific site binding was predominant
while at low concentrations, non-specific binding obtained.
The differences between Na+ and Mg++ may be due, in part,
to the ability of the latter to bridge between two adjacent
anionic sites on the polymer whereas the monovalent cation
cannot. Note that the distance between adjacent phosphates
o
along the DNA backbone is approximately 7 A. C.Sander and
and O.P. Ts'o (50) believe that the binding of Mg++ to DNA,
RNA, poly A, poly A*U, poly I, poly I*C, and denatured DNA
can be rationalized using a linear Scatchard relationship
(vide infra). Their specific site model was applicable for
up to about 70% saturation.

15
The degree of binding of acridine orange to native
and denatured DNA was determined by equilibrium dialysis
in 0.1M and 0.001M NaCl at 20°C (51). Three successive
stages of binding were observed: the first corresponding
to intercalation of monomers; the second, formation of
bound dimers; and the third, external binding of aggre¬
gates. The appearance of the three modes of binding
compared to two for most aminoacridines is attributed
to acridine orange's proclivity to dimerize. The dialysis
data was augmented with absorption and fluorescence
spectra, electric dichroism, electric birefringence, and
circular dichroic spectra.
Scruggs and Ross (42) found that the intrinsic
viscosity of three types of DNA (salmon sperm, calf thymus,
and T4 phage) decreases in the presence of univalent cations
as the ionic strength increases, attaining separate limiting
values for each. Prior to their work, the apparent insen¬
sitivity of the viscosity of DNA to variations in ionic
strength was considered anomalous because polyelectrolyte
viscosities are usually highly dependent on salt concentra¬
tion. The workers asserted that previous invariance was
due to the presence of trace amounts of basic proteins,
polypeptides, and polyamines and that DNA was, in fact,
similar to all other polyelectrolytes.
S. Bram (52) investigated the secondary structure of
DNA in solutions of varying concentrations of Li+, Na+,
Rb+, and Cs+ using X-ray diffraction spectroscopy.

16
Differences between scattering patterns by sodium DNA
in solution and those expected from theoretical calcu¬
lations based on an unperturbed B form of DNA were due
to structural changes in the B form arising from the
polymer's interaction with the species in the medium.
This is in concert with others'(53-55) conclusions that
the number of base pairs per turn of the DNA helix is
a function of the ionic environment. For cesium DNA
the extrapolated lengths for 10 nucleotide pairs aré
o o . .
43.5 A at zero ionic strength and 34.6 A at finite CsCl
o
concentration, whereas in NaCl, the values are 33.5 A
o
and 32.0 A, respectively. (A novel DNA conformation
model which differs markedly from the classical Watson-
Crick B form is proposed by Stig Erlander (56) who argues
that the differences between the values given above cannot
be rationalized using the Watson-Crick model.) A theo¬
retical treatment of cooperative binding was advanced by
G. Schwarz (57) in which nearest neighbors along a linear
polymer interact with one another. Two types of intrinsic
internal binding are assumed: (1) that of an isolated
ligand (nucleation); (2) that of a ligand to a site
immediately adjacent to one already occupied. The model
also includes considerations for dimer formation and
competitive (non-cooperative) binding at adjacent sites
by unrelated ligands. Matrix methods of calculating
species concentrations under equilibrium conditions are
presented.

17
An understanding of the effects of ionic strength
on the binding of small molecules to DNA requires at
least some knowledge of the charge distribution on the
surface of the polyanion. Electrophoresis and membrane
equilibrium experiments allow the determination of the
apparent fractional charge per DNA phosphorous, i, which
is an index of net charge (58). The value of i increases
from 0.250 to 0.34 on going from a solution containing
0.005m NaCl to one having 0.05m NaCl, showing that the
net charge of the polyanion increases with increasing
ionic strength. Ross (58) considered the above information
in light of Gorin's model of a long rigid cylinder (59)
and applied it to DNA, which is intermediate between a
rigid cylinder and a random-coil polyelectrolyte. Ross'
data substantiates Gorin's assumption that a rigid cylinder
(and, therefore, DNA) obeys Debye-Hiickel approximations.
Manning (60) postulated that, as the ionic strength
approaches zero, the effective charge of a polyelectrolyte
is maintained at a critical constant by condensation of
counterions along its surface as long as the formal charge
does not exceed a critical value. His predictions were
borne out by independent investigators (33) who found
that at low ionic strength (0.002M) the electrophoretic
mobility of DNA was unaffected by bound acridine orange.
Lerman's model for intercalation of small molecules
into DNA satisfies many experimental observations but
leaves some uncertainties unresolved arising from

18
consideration of the energetics involved (61). One of
the most perplexing problems is the strong dependence
of the intercalative mode of binding on the ionic strength
(albeit not as dependent as surface binding). Also, the
existence of a maximum degree of association as a function
of ionic strength requires explanation. Attempts have
been made to resolve these problems by invoking "stacking
energies" rationalized in terms of short range, ionic
strength independent interactions between DNA base pairs
and the aromatic portions of the acridines. These ratio¬
nalizations, however, are incompatible with the fact that
aromatic hydrocarbons intercalatively bind only very weakly.
M. Gilbert and P. Claverie (61) point out that the solution
energies of the intercalating compounds parallel their
binding affinities and, therefore, may be at least as
important as the stacking energies. Moreover, the limiting
of intercalation processes long before all available sites
have been occupied mandates a limitation mechanism inde¬
pendent of the heterogeneity of sites. The critical effect
of transferring positive charges from solvent "cavities,"
having dimensions on the order of the small molecules',
into large cavities interior to the lipophilic DNA must
be accounted for in a model of the dye-DNA complex — as
it is in their model.
T. Herskovits (62) studied the relationship between
electrolyte concentration and solvent denaturation transi¬
tions in various media. Both increasing chain length and

19
increasing hydrocarbon content increase the effectiveness
of the denaturant. It was concluded that these observations
demonstrate the importance of hydrophobic forces in main¬
taining DNA's aqueous configuration. In most cases, the
denaturation was reversible with added electrolyte.
It is Bradley's opinion (63) that by using a combina¬
tion of statistical and quantum mechanical methods, it
should be possible to relate all absorption, emission,
equilibrium constant, and optical rotation observations
in a unified way. With this aspiration in mind, he
presented a statistical model of dye-polymer binding as
a step toward that goal. He assumed that specific polymer
sites were in a linear array and dye molecules were bound
to them by simple electrostatic forces, with possible
enhancement by neighboring sites. Furthermore, the sites
were sufficiently close to allow bound dye-bound dye
interactions with concomitant changes in their absorption,
fluorescence, and phosphorescence spectra. The distribution
of N consecutive sites occupied by dye molecules, in their
varying states of aggregation, can be computed theoretically
in terms of either the ground state free energy of inter¬
action of a pair of neighboring dye molecules, AF, or in
terms of a stacking coefficient, K, related to AF,
K = exp(-AF/kT) . (1-1)
The AF and K terms may be equated to an experimentally
obtainable ratio of P sites per D dye species,

20
P/D = (1-F1//2) 1 (1+ (K-l) F) (1-2)
Values of K range from 1.25 for DNA to 826 for poly¬
phosphate (K = 6.2 for denatured DNA). There is
evidence that the value of K is directly related to
the conformation of the polymer. The fact that K is
equal to 1.25 and not to 1.00 for DNA — as it would be
for a completely random site occupancy — indicates some
non-randomness which could arise from dye-dye interaction
along the polymer. A simplified quantum mechanical
treatment is presented to allow computation of the magni¬
tudes of the relationships between the dye molecules of
filled units with empty sites.
Kinetics of intercalation of aminoacridines into DNA
were first studied in depth by Li and Crothers (23) in 1969.
Employing data from temperature-jump relaxation methods,
they proposed a relatively simple two-step process for
intercalative binding. It is worth noting that static
experiments will not allow a separation of these two steps
of the intercalative mode. The data fit the mechanism
^12 ^23
P + DNA -y.. > (P) (P) . (1-3)
£ out ^—, —■■ in
K21 32
where P is proflavine and (p)out and (P) are surface
bound and intercalatively bound dye, respectively. The
overall insertion reaction occurs in the millisecond time
range and is first order from the (P) . state. The two-
out
step hypothesis is substantiated by comparing the large

21
relative thermodynamic and spectral changes of the
exterior binding with the small changes for the interior-
bound species that occur upon glucosylation of the DNA
(vide infra). Two limiting mechanisms are advanced and
alternate mechanistic pathways discussed.
Passero, Gabbay, and others (64) used a simplistic
model which included competitive effects of metal ions
in solution to arrive at apparent association constants
of two reporter molecules (Xllla, XHIb) to DNA. Electro¬
static potentials and interactions between near neighbors
were ignored, as were any forms of surface binding by the
reporter molecules. The overall binding process was
expressed as two separate equilibria
R + P —> RP (1-4)
and
M + P —> MP (1-5)
where R, M, and P are free reporter, metal, and DNA
phosphate, respectively, and RP and MP the corresponding
bound species. An apparent binding constant, K' may be
defined relative to the metal ion concentration
K'r = [RP] / [R] [P] . (1-6)
Finally, assuming the binding of R or M to the polyanion
results in only one phosphate being affected per R or P,
the apparent association constant may be expressed in terms
of association constants for reporter binding, K , and

22
metal binding,
K'r = Kr / (1 + Km [M+]) . (1-7)
The authors maintain that their assumptions are valid
by virtue of the fact that they obtain reasonably
constant values of and K„ under a variety of conditions.
However, if the reporter molecules are intercalatively
bound, as the authors assert (and they are probably
correct) then it is not valid to compare the reporter
and metal binding because two entirely different types
of sites are being occupied. The highly hydrophilic,
symmetrically charged metal cations are almost certainly
surface bound. Thus, the constancy of the results are
probably fortuitous or, at best, important deviations
are being masked by experimental error.
Relaxation kinetics of ethidium bromide(XV) binding
to DNA revealed three types of bound species at equilib¬
rium (65). Aside from the conventional surface-bound and
intercalated species ascribed to type II and type I binding,
a third species was found which involved an ethidium bromide
transiently attached to sites of two separate DNA strands.
The bimolecular rate constant for transferring an ethidium
bromide from one DNA phosphate to another is 3 to 6 times
larger than the rate of intercalation of free molecules.
The more rapid bimolecular mechanism is dependent on the
small molecule's ability to bind to two sites simultaneously.
In the case of ethidium bromide, the tricyclic region of the

23
molecule provides one site while the phenyl moiety is
believed to provide the second, less favorable, site.
Divalent-cation-specific electrodes were used to
study the binding of Mg++ to DNA (50) in a solvent
-3
containing 5 x 10 M pH 7 phosphate buffer. Results
were treated in terms of site-binding to phosphates in
a multiple equilibrium process, employing a Scatchard
treatment, which showed linearity over a range of 30%
to 70% saturation of sites. Approximations to the
Scatchard method yielded an apparent intrinsic binding
constant of 6 x 10^ (AG° = 5.2 Kcal/mole). The use of
cation-specific electrodes mandated DNA phosphate concen-
_3
trations in excess of 5 x 10 M, causing aggregation of
DNA and precipitation of MgDNA. Thus, the values of the
constants may include significant error.
Three methods for evaluating equilibrium association
constants for the binding of small molecules to DNA will
be outlined. They are, the classical and most commonly
employed Scatchard method, a modified Scatchard model,
and a mass-action approach which will be used for this
investigation. Scatchard's original treatment was designed
for binding of small molecules to proteins. In Scatchard's
t
words:
If the various groups on a protein molecule
act independently, we can apply the Law of
Mass Action as though each group were on a
separate molecule and the strength of binding
can be expressed as the constant for each
group. Often a single constant will express
the behavior of several groups. (66:660)
If the groups do act independently of one another, changes

24
in free energy for the reaction of the protein with
small molecules is made up of the statistical entropy
terms for v plus a term proportional to v. when the
total number of sites (or groups of sites) is large, the
calculations become tedious — if the total number is
not known the calculations cannot be performed. Such
a situation does arise if one attempts to apply the
equations without modification to nucleic acid binding.
Since DNA polyanions commonly have molecular weights
/T
in excess of one million and may contain as many as 10
phosphates (67) Peacocke and Skerrett (20) redefined the
parameters of the Scatchard equations to allow their
application to nucleic acids. Their assignments and an
abbreviated derivation of the Scatchard equations are as
follows.
Let R represent the amount of small molecules, D,
bound per mole of nucleic acid phosphorus (generally a
fraction), and C represent the molar concentration of
free D. Consider a number, n^, of P classes of binding
sites for D on the DNA polyanion. If each class has
associated with it an intrinsic binding constant , then
we may write for any general case
J=P
R = I
J=1
n .K .C
1 + K.C
1
(1-8)
This may be simplified when all sites are identical
R/C = Kn - Kr
(1-9)

25
If there are two distinct types of binding sites, I
and II, then
R =
ni Ki C
1 + KjC
nll KII C
1 + KIIC
(1-10)
Figures 2a and 2b are representative plots of R/C vs C
for one and two types of sites, respectively. Curvature
may result in the case of only one type of site from
variations in the electrostatic free energy term which
is dependent on R. Such non-linearity can become critical
when the intercalating species bear opposite charges but
may be partially obviated by replacing with the term
K'z exp (AG°r/(RT) (1-11)
where K'z is independent of R, and AG0^ is an electro¬
static free energy term dependent only on the total effec¬
tive potential. The case in which there are two types of
binding sites is of primary interest in this work. The
stronger, intercalative binding corresponds to type I of
figure 2b whereas surface binding is represented by
region II. With two types of sites, it is sometimes
possible to discern distinct linear portions in the plots
of R/C vs R as suggested in the idealized figure. Only if
the magnitude of one of the constants is about 10 to 100
times greater than the other, can reliable values of the
individual association constants be obtained by extrapola¬
tion of the data in each of the linear regions. Even then,

26
potentially erroneous assumptions regarding changes in
the electrostatic free energy must be considered. More¬
over, as previously stated, the proposed model does not
allow for any form of interaction between the two types
of binding sites. This is not generally serious in
protein binding where the sites are often noninteracting
but can be a problem in the case of the nucleic acids.
Factors such as cooperative binding and near-neighbor
exclusion will introduce sigmoidal characteristics into
the plots of R/C vs R (21). Another shortcoming of the
Scatchard method is that extrapolation of the legs of
the hyperbolic plot to obtain slopes and intercepts
often places an inordinately high degree of reliance
on the data at the extrema of titrations. It is at the
extrema that the data, regardless of how it was obtained,
will be the least precise.
For aminoacridines, intercalative binding pre¬
dominates up to a value of R_<0.20 corresponding to 4 or
more phosphates per bound cation. Surface binding, on
the other hand, is predominant for values of R between
approximately 0.5 and 1 for compounds which do not dimerize
and up to 2 or more for those that are known to aggregate
as free species at moderate concentrations.
Numerous investigators have suggested improvements
of the basic Scatchard method as modified by Peacocke and
Skerrett. Included are corrections for electrostatic
effects (27,31,61,63), competition of metallic ions in

27
solution (45,46,67), near-neighbor effects (23,50,57,
63,65), small molecule aggregation (51,57), and configura¬
tional effects due to solvent and ionic factors (23,25,28,
39,44,52,62,68). One of these improved models, proposed
by Armstrong, Kurucsev, and Strauss (33) incorporates
some novel features including: (1) the total concentration
of intercalation sites available is fixed, a priori, to
include every slot between successive DNA base pairs; (2)
an intercalated dye inhibits intercalation at its two
immediately adjacent slots; and (3) inclusion of a term
to allow for an intercalated monomer forming a dimer with
a free monomer resulting in a spectroscopically distinct
species. The authors found (3) became important when the
degree of binding, 6, exceeded about 0.2 mole of dye per
mole of DNA phosphate (dyes were acridine orange and pro¬
flavine) . Their expressions for the association constants,
based upon theoretical mass action relationships are, for
intercalation
2 (3-, — 35) (1-28-^
(1-12)
and for surface binding
(1-13)
where 6^ = the number of intercalated molecules per DNA
phosphate.

28
82 = the number of externally bound dimers per
DNA phosphate.
C, = the molar concentration of free dye monomer.
M
= the apparent binding constant for the monomer
species intercalating.
K2 = the apparent binding constant for surface
binding.
These relationships imply that the electrostatic potential
of the DNA is unaffected by binding. Values of 8-^ and $2
may be spectroscopically evaluated knowing the molar
absorptivities of the bound monomer, e^, and the bound
dimer, eD,
£T “ + (1-14)
where 8^ = = ^^2' anc^ £T as t^e tota^ "m°lar
absorptivity" of all bound species. The model was tested
by determining the binding of acridine orange and pro¬
flavine to DNA (pH 6.5, 22°C) at ionic strengths of 0.002m,
0.020m, and 0.200m. The above method is an improvement
over the Scatchard procedure but still demands a priori
selection of an exact mode of reaction. That is, though
the overall stoichiometries are experimentally evaluated,
the thermodynamic manner in which the free species react
is not defined. For example, in the intercalative case,
where 4 phosphates are involved for each 3-aminoacridinium
binding site, are the 4 phosphates to be considered as 4
separate entities, as 2 separately interacting moieties,

29
each containing 2 phosphates, or as one entity containing
all 4 phosphates? In each case, the overall stoichiometry
is 4 to 1, but the equilibrium expression of the first
includes a fourth power term in DNA phosphate concentration,
the second would have a squared phosphate concentration
term, while the third would be linear in phosphate.
To resolve this problem, Capomacchia (69) has proposed
a mass-action approach based on best-fit analyses of
spectroscopically derived concentrations of free and bound
species. Assuming that the binding process obeys the Law
of Mass Action and total concentrations of all species are
sufficiently low that aggregation of like species is in¬
significant, we may write
BH + qSu ~ > BH(Su) (1-15)
where BH is the free cationic form of the small molecule,
q is the number of DNA binding sites that complex one
small molecule and Su denotes the unbound DNA binding
sites. Each site, Su, acts as one distinct entity in
the equilibrium expression and may contain any number
(1,2,3,. . . m) of phosphates. Note that throughout this
text binding sites will be discussed in terms of DNA
phosphates, simply for convenience, because the total
concentration of DNA phosphate is easily obtained ex¬
perimentally. This terminology does not imply, for
instance, that phosphates are the binding sites for
intercalatively bound drug. The association constant,

30
on a molar basis, for the process is
K = [BHP]/[BH][Su]q
(1-16)
where [BHP] is the molar concentration of bound species,
[BH] is the concentration of free small molecule, and
[Su] is the concentration of the uncomplexed DNA sites.
Unbound site concentration may be determined from the
total site concentration, S^, and bound site concentration,
S
c
(1-17)
Now, since
Sc/q = [BHP] (1-18)
and st = [Pt/m] (1-19)
where Pfc is total DNA phosphate concentration, we may
write
Su = [Pt/m]- qtBHP]
(1-20)
Substituting equation (1-20) into (1-16) yields
K =
[BHP]
[BH]([Pt/m]- q[BHP])q
(1-21)
Taking the log of both sides of (1-21) and rearranging,
we obtain
l0g ( !b£]]) = log K + qlog ([Pt/m]- q[BHP])q (1-22)

31
Molar concentrations of free and bound small molecule
species may be determined from any of a variety of
experimental techniques. The overall stoichiometry of
the process may be obtained from a Scatchard treatment or
a Job's plot. Knowing the overall stoichiometry, initial
values of m and q may be selected which satisfy the simple
relationship mq = n, where n is the total number of DNA
phosphates per bound small molecule. A rapid determination
of the most probable combinations of values may be made by
plotting log ([Pt/m]- q[BHP]) vs.log ( [BHP] / [BH] ) . That
pair which yields the most consistent straight line having
a slope of q may then be used to compute values of K for a
series of ratios of [BHP] to [BH]. Again, the pair which
provides the most consistent calculated equilibrium associa¬
tion constants for all data points, corresponding to either
the surface or intercalative mode, are taken as the exponents
in the binding expression. For interactions between 3-
aminoacridinium cation and calf thymus DNA, best-fit values
of m and q are 3 and 1,respectively for surface binding and
2 and 2 for the intercalative mode (69) . To summarize, the
method of Capomacchia allows for the determination of
equilibrium association constants based on mechanistic
considerations. Furthermore, by computing values of K from
data in the midregion of a titration, undue reliance on the
least trustworthy data at the extrema of a titration is
avoided.
Many workers have been satisfied with reporting associ¬
ation constants calculated from molar quantities rather than

32
from the more correct activities of the various species.
Doing so implies two things: activities of the various
species in the expression are equal to molarities, and
the molarities are equal to molalities. It is unlikely
that approximating molalities with molarities introduces
significant error throughout the usual concentration
ranges employed. But activities, on the other hand,
may differ markedly from molarities. For this reason,
evaluation of association constants in terms of activities
4
is recommended. Molality is related to activity by
a. = a.m. (1-23)
ill
where a^ is the activity, ou is the activity coefficient,
and itk is the molality of species i. Note that the
activity equals the molality only when a = 1. For a
process in solution
A+ + nB v^ ABn(n+l) (1-24)
the equilibrium constant is rigourously defined as
aAB
n
(A++)(A¡-)n
a M
AB AB
n n
(c++Ma+) ( (1-25)
It is extremely difficult to determine individual activities
since their is no way of separating effects due to positive
ions from those due to the accompanying negative ions. We
may, however, define a mean activity, a—, which is the
geometric mean of the individual activities

33
(1-26)
Similarly
(1-27)
For ideal dilute solutions, the mean ionic activity
coefficient will be equal to 1. If deviations from
ideality are caused entirely by electrical interactions
of point charges, it may be shown that, for aqueous
solutions at 25°C (70)
log a— = -.509/Z + Z_/l1//2
(1-28)
where Z+ and Z_ are the charges of interacting species
and I is the total ionic strength of the system
2
m. Z7
1=1/2 E
(1-29)
. m . ¿j .
Ill
Equation (1-28) is an expression of the Debye-Hiickel
limiting law for activity coefficients and is applicable
only to dilute solutions containing point charges. As the
ionic strength becomes large (e.g., >0.01m) and/or the
ions become less ideal (become larger, more polarizable,
etc.) the relationship becomes increasingly approximate.
Under these conditions, an extended Debye-Hiickel equation
may be invoked which introduces a correction for the finite
sizes of ions:
log =
-0.509/Z + Z Jl1/2
1/2
1 + Bdl
(1-30)

34
B is a constant for a given solvent and includes such
terms as solvent viscosity, dielectric, and temperature
while d is an average effective diameter of the ions. A
factor not considered in equation (1-30) is the effect of
ions on solvent molecules and solvent structure. This
may become especially important for solutions containing
large polyanions such as DNA.
Lyons and Kotin (44,45) report a method of assessing
values of a+ and a independently of one another. They
also determined for Na+ and Mg++ when binding to DNA
and three synthetic polyanions. Their somewhat tedious
procedure involving equilibrium dialysis allowed them to
distinguish site-specific and non-specific forms of surface
binding. The activity coefficient of Na+ was shown to
decrease with increasing dilution of metal ion in the
presence of DNA suggesting an increased degree of binding.
Thermodynamics of Binding Reactions
A second topic of this work is concerned with the
thermodynamic parameters associated with the binding of
small molecules to DNA. Hence, a brief review of several
pertinent articles is presented here.
M. Gilbert and P. Claverie (61) measured total energies
for an intercalation process and attributed them to three
types of electrostatic interactions. They were (1) the
attractions between the bound cations and the DNA phosphates,
(2) the repulsive forces between cations, and (3) solvent

35
response to the electric field generated by the charged
species. At high ionic strength the phosphates are
relatively neutralized causing (1) to become insignificant
and, thereby, causing a decrease in both the strength and
extent of binding. Some limiting of the binding process
occurs even at very low ionicities because cation-cation
repulsions are still important with respect to intercalated
species. The third type of interaction serves to reduce
all forms of electrostatic energies and will be most ef¬
fective for those species most intimately in contact with
the solvent. The authors have shown that phosphates and
complexed acridines are more solvated than free acridines
so that, as the ionic strength is raised, their stabilizing
interactions are more reduced than are acridine-acridine
interactions. In that the forces arising from the latter
are destabilizing, the result of the countervailing trends
is a net destabilization of the complex with increasing
ionic strength. The model used here can also be applied
to other cationic small molecules, including protons.
Similar results were reported for the binding of proflavine
to DNA via fluorimetry (71). It was found that the electro¬
static contribution to the total free energy decreased with
increasing ionic strength due to screening of the potential
by an ionic atmosphere.
Thermodynamic parameters accompanying the intercalation
of calf thymus DNA were evaluated by Chambron, Daune, and Sadron
(72) using equilibrium dialysis. Temperatures ranged from

36
0°C to 70°C and ionic strengths of 0.01M, 0.1M, and 1.OM
were maintained using acetate buffer at pH 5.9 in the
presence of EDTA. The number of sites available for
binding decreased upon increasing either the temperature
or the ionic strength. The enthalpy of reaction decreased
with decreasing temperature; AH°= -19 Kcal/mole at 70°C
and -4.5 Kcal/mole at 0°C (I = 0.01). This was ascribed
to a thermally reversible change of state of DNA which
occurs at approximately 40°C. Total standard free
energy was considered to be a sum of three terms. The
first of these, AG°, characterizes the dipole-induced
dipole interaction between base pairs and cations while
the second, AG°, is the electrostatic free energy. The
third, AG°, is the solvation free energy corresponding
n
to desolvation of reactants and solvation of the bound
complex. At high ionic strengths, AG° and AG° are very
small, hence, the total free energy of 6.4 Kcal/mole is
due to AG°. The authors' value of AS°= 11 Kcal/mole
degree is in agreement with theoretical predictions.
Thermodynamic energies of reaction between proflavine
and two forms of DNA determined from temperature-jump
relaxation methods are summarized below (23).
DNA
EI2
E- E-
zl L23
E±
^32
AB°2
AH23
TAS°2
TAS53
calf
thymus
4
14 16
14
-9.8
2.0
-5.4
3.2
T2
—
13
17
-3.4
-3.9
1.6
-2.8
4jis
the activation energy for
transformation
from
state i

37
to j (see equation (1-13) and AH?^ and AS?^ are the thermo¬
dynamic enthalpy and entropy changes for the same reaction.
Measurements were done over the temperature range 10°C to
25°C, 0.20M Na+ and pH 6.9. Of particular note here is
that the enthalpy change for outside binding is much more
strongly affected by glucosylation of the DNA than it is
for the intercalated form (T2 DNA is glucosylated). For
example, AH°2 changes from -9.8 to -3.4 Kcal/mole dye on
going from calf thymus to T2 DNA but for the overall inter¬
calation process (AH°2 + AH°2 = ^H13^ t*ie c^an5e i-s only
from -7.8 to -7.3 Kcal/mole dye. It is unlikely that small
differences in base composition between the two forms of
DNA can account for the large differences in the thermo¬
dynamic values of the binding reaction.
Entropy and enthalpy changes during reaction of
ethidium bromide with calf thymus DNA were computed from
temperature-jump relaxation data by Bresloff and Crothers
(65) over the temperature range 15°C to 32°C in 1.OM Na+.
Values for the formation of a surface bound species were:
AH° = 1.1 Kcal/mole and AS°(19°C) = 5.6 Kcal/mole, while
for one of a pair of intercalatively bound species: AH° =
-7.8 Kcal/mole and TAS°(19°C) = -2.3 Kcal/mole. Values
for a second intercalated form could not be obtained due
to the rapidity of its binding process.
Bradley (63) determined the ground state free energy
of intercalation of a pair of neighboring acridine orange
molecules bound to DNA to be -0.065 Kcal/mole dye. By

38
comparing this to the spectral shift accompanying aggrega¬
tion (504 nm to 464 nm) corresponding to 4.8 Kcal/mole dye
separation in energy between the ground and excited states,
he concluded that the spectral shift results from an
increase in energy of the excited state, rather than a
decrease in the ground state.
Enthalpy and entropy changes, AH°and AS°respectively,
of reactions may be determined from the relationship be¬
tween the equilibrium association constants for the processes
at various temperatures and the temperature as set forth
by the van't Hoff relationship.
l0g K 2.303R T
(1-31)
+ Constant
This important equation will be derived below (73).
For ideal systems, any change in the total free
energy, G, can be said to be due to changes in either
temperature or pressure, since the state function, G,
can be defined in terms of pressure, volume, and temperature.
(1-32)
We define
G = H - TS
(1-33)
and
H = E + PV
(1-34)
where H is the enthalpy, S is the entropy, and E is the

39
internal energy for the system under consideration. Recall
that, the enthalpy changes are commonly termed "heats of
reaction" for processes, entropy may be used as a measure ,
of the degree of order or disorder, and that the internal
energy is unavailable for work of any form. The change in
the free energy, dG, may now be expressed as
dG = dE + PdV + VdP - TdS - SdT . (1-35)
For a reversible process in which no work other than
expansion occurs,
PdV = dw (1-36)
and
dE - dq + dw = 0 . (1-37)
The terms q and w refer to the heat absorbed, and work
done, by the system respectively. Equation (1-37) is a
direct consequence of the First Law of Thermodynamics and
states, in effect, that the total change in the internal
energy of a system is the sum (or difference) of the amount
of heat absorbed, dq, by the system and the amount of work
done, dw, by the system. Notice that heat and work are not
functions of state (which are conventionally assigned upper
case symbols). Equation (1-35) now becomes
dG = VdP - SdT
(1-38)

40
which upon equating with equation (1-32) yields
(3T}p S
(1-39)
and
<|§) = V . (1-40)
Equation (1-39) suggests an experimentally tractable
relationship between free energy (via K) and temperature,
when pressure is held constant. Moreover, we may extend
the relationship to include the changes in free energy as
a function of temperature for a process involving reactants
going to products. We may now equate (1-38) with equation
(1-32), rearrange, and solve for the partial differentials
to obtain
(9AG/9T)p = -AS
(1-41)
where AG = (G , , - G , , ) and S = (S ,
products reactants products
S reactants^ * T*ie entroPY term of equation (1-41) may be
eliminated by recalling that AS = (-AH + AG)/T and substi¬
tuting
(9G/9T)p = (-AH + AG)/T
(1-42)
Assume, henceforth, that the pressure will be held constant.
Rearranging (1-42) gives

41
dAG/dT - AG/T = -AH/T . (1-43)
Equation (1-43) may be expressed as
T (AG/T) = -AH/T (1-44)
since, upon differentiation, it would become
m T(dAG/dT)-AG _ dAG/dT
1 2
T T
Equation (1-45) may be written
and, in that
AG° = -RT In K (1-46)
AG
T
(1-45)
in terms of standard states,
it can be stated as
d In K/dT = AH°/RT2
which is equivalent to
d log K
d (1/T)
-AH°
2.303R
+ Constant
(1-47)
(1-48)
Assuming that AH° is temperature independent, equation (1-48)
can be integrated to yield equation (1-31). The above
relationship may be applied to reactions in solution whose
equilibrium association constants are based on the activities
of the species involved [see equations (1-23) through (1-30)].
Plots of log K vs. 1/T should yield straight lines having
slopes equal to -AH°/2.303R. If the relationship between
log K and 1/T is not linear then the enthalpy of binding is

42
not constant with respect to temperature. For systems
containing DNA, enthalpies of binding determined from van't
Hoff plots are often not constant with temperature, but
decrease with increased temperature, presumably due to
changes in DNA's conformation (72). As a result, it should
not be surprising that binding curves obtained spectro-
photometrically and by dialysis equilibrium can agree at
one temperature and disagree at another (21).
Rationale for the Selection of the Model System
and Experimental Design
As this study is concerned with the effects of com¬
peting metal ions and of alterations of the ionic strength
on the binding of small molecules to biopolymers, the choice
of DNA as the model polymer is evident. The selection of
3-aminoacridine as the small molecule is based on several
factors, two of which are especially important. The first
criterion that the small molecule must meet is that it bind
to DNA. To do so, it is necessary that the species exist
as a cation since neutral compounds do not bind to DNA —
or if they do, the binding is too weak for our purposes.
(It should be evident that electrostatic repulsive forces
prevent anions from interacting with DNA.) Furthermore,
the compound must be positively charged in solutions whose
pH is between 4 and 11 so that denaturation of the polymer
is minimal (74). Hence, if the compound is acidic its pK
cl
should be at least 5. The pK of 3-aminoacridine is 8.04
a
(7:Ch 4) meaning it is essentially totally protonated in a

43
solution of pH 6. Moreover, 3-aminoacridine allows con¬
venient evaluation of both of the principal inodes of binding
of small molecules to DNA, under a given set of conditions,
in a single titration. That is, the range of the ratio of
total drug to total DNA concentration over which the two
modes of binding can be experimentally evaluated is about
0.1 to 10. Obviously, it is mandatory that each type of
binding occur in separate regions of the titration.
The second major criterion that the model compound
must satisfy is that it allow an expedient means of measur¬
ing the extent of binding to the polyanion. Of the various
methods of determining equilibrium constants one of the best
is absorptiometric spectroscopy. The 3-aminoacridine cation
is particularly amenable to this method for the following
reasons. (1) It absorbs strongly in the visible region of
the spectrum as the free cation having molar absorptivities
of 1.41.3 x 10^M ''"cm ^ at 365 nm and 1.259^ x 10^M ^cm ^ at
454 nm (75). Such magnitudes allow good sensitivity at low
concentrations of the compound. The molar absorptivities
of the bound drug in this region are 8.494 x 10 M cm at
368 nm and 9.07.1 x 10^M "''cm ^ at 463 nm. The large dif¬
ference between the absorbances of the band maxima of the
free and bound drug provides good sensitivity in following
a titration. The longest wavelength absorption maximum
for DNA is at approximately 260 nm meaning that there is
little spectroscopic interference from the polymer. (2)
Both the free and bound species have two distinct absorption

44
envelopes in the visible region corresponding to and
bands [Platt's nomenclature (76); a and para bands,
cl
respectively, in the Ciar nomenclature (77) ]. These arise
from transitions from the ground state to the
(=365 nm), and to the (=460 nm) excited states. The
transition moment of the former is parallel to the long
axis of the acridine ring system while the moment of the
latter is orthogonal to the and parallel to the short
axis of the molecule. By observing shifts in the two
maxima, along with any changes in vibronic character,
during the course of a titration, one may be able to
deduce binding geometries (78) . The 3-aminoacridine
molecule also fluoresces in its free cationic form and,
less intensely, as the bound species, thereby providing
another spectroscopic method with which to study its
interaction with DNA.
The more thoroughly studied aminoacridines, proflavine
and acridine orange, meet the criteria mentioned above in
regard to 3-aminoacridine. Both of them, however, begin to
-5
aggregate at concentrations as low as 10 M, thereby
introducing undesireable complications. Also, the mono-
substituted 3-aminoacridine, being a simpler molecule than
either proflavine or acridine orange, is preferable as a
model compound. It is hoped that information gained using
the simple compound may be applied to more complex systems.
Many references to the effects of ionic strength on the
binding of small molecules to biopolymers can be found in

45
the literature. Unfortunately, the majority are ancillary
to other studies and are not sufficiently complete in them¬
selves to provide much specific information. Too, even for
a given molecule binding to a given polymer, the wide
variety of experimental variables under which the separate
investigations were done precludes any valid comparisons of
results. The present work was undertaken as a result of
this lack of a comprehensive investigation. Our pharmaceu¬
tical interest led us specifically to DNA as the polyanion,
though other, synthetic polymers may have provided a less
complex system.
Some of the immediate questions which come to mind and
which may be answered by an investigation of the effects of
changing the ionic constitution of a drug-DNA solution are:
are the effects predictable on the basis of electrostatic
theory? That is, may the effects be explained in terms of
the simple Debye-Hiickel limiting law at low ionic strengths
and by extended Debye-Hiickel equations at higher ionicities?
If the above questions cannot be answered affirmatively,
might deviations at low ionic strengths be due to the
polymeric nature of the anion? Are deviations at high
ionic strengths due to specific ion interactions and, if so,
can they be rationalized on the grounds of simple competition
for binding sites or must a less simplified, mixed complex
model be introduced? Other questions are: will altering the
ionic medium affect surface and intercalative binding in the
same manner? If not, how can the differences be reconciled?

46
How will changing the ionic radii, ionic charge, and
charge density of the metallic ions affect the equilibria?
The last question is particularly germane in in vivo
biological binding studies because of the wide variations
in concentrations of Na+, K+, Mg++, and Ca++ in a living
organism (79).
To try to answer these questions, equilibrium associa¬
tion constants for the surface and intercalative binding
of 3-aminoacridinium cation to DNA in a variety of ionic
media were determined. Background electrolytes included
-4
concentrations between 0.15M and 6.3 x 10 M of the phosphate
salts of Li+, Na+, K+, Rb+, Cs+, and (CH^J^N*. Also con-
++ ++
sidered were the acetates of the alkaline earths: Mg , Ca ,
Sr++, and Ba++ at concentrations ranging from 0.025N to 1 x
10 ^N. The effect of as opposed to CH^COO as the
counteranion of the supporting electrolyte was evaluated.
The pH for all of the titrations was 5.90 and the tempera¬
ture was 25.0°C. A minimum of two titrations were done for
each set of conditions.
Thermodynamic parameters were evaluated for systems
containing 3-aminoacridine and DNA with a background
electrolyte of 0.010M KH2PC>4 at pH 5.90. Similar investiga¬
tions were conducted using 7-aminoacridinium cation(XVI) in
place of 3-aminoacridinium. By studying the binding of the
linear benzalog of 3-aminoacridine and comparing the thermo¬
dynamic parameters of the two, the relative importance of
enthalpy and entropy contributions to overall free energies
may be assessed.

CHAPTER II
EXPERIMENTAL
Materials and Procedures
All chemicals were, unless otherwise stated, either
reagent or analytical grade. Water was multiply distilled
in a tin-lined still (Barnstead, Sybron Corp.). It was
redistilled until its specific conductance was less than
1.0 megaohm cm 1 at 25.0°C. Stock solutions of supporting
electrolytes were prepared within three weeks of use and
were refrigerated during storage to minimize biological
growth. Approximately 100 to 200 ml of a 0.15M solution
of each electrolyte were prepared as outlined below and
titrated with 0.72M H3P04, 1M CH^CC^H, or 1M metal hydroxide
solution to pH 5.9, in all cases. Concentrations of the
hydroxide solutions of the various cations were determined
by titration with standardized 1.000N sulfuric acid.
A stock solution of tetramethylammonium phosphate was
made by titrating 13.16 ml of 1.14M tetramethylammonium
hydroxide (Mallinckrodt Organic Reagent, 10% in water,
lot: 15) in 70 ml of water with 0.72M phosphoric acid
(Fisher, lot: 755459) and adjusting the final volume to
100 ml. Lithium phosphate stock solution was prepared by
47

48
titration of 0.476M lithium hydroxide with 0.72M phosphoric
acid. The hydroxide solution was made by reacting stoi¬
chiometric amounts of anhydrous lithium sulfate (PCR, Inc.,
Gainesville, Florida, lot: 10465) and barium hydroxide
heptahydrate (Mallinckrodt, lot: RTZ). Sodium and potas¬
sium phosphate stock solutions were prepared using sodium
phosphate monohydrate (Mallinckrodt, lot: ABK) and mono¬
basic potassium phosphate (Fisher Certified, Reagent, lot:
20175), respectively. The acid level of each solution was
adjusted with 0.72M phosphoric acid to pH 5.9. The stock
solution of rubidium phosphate was prepared by titrating
an aliquot of 0.43M rubidium hydroxide (Pfaltz and Bauer,
no lot number) with 0.72M phosphoric acid. Similarly,
cesium phosphate was prepared from 0.201M cesium hydroxide
(Pfaltz and Bauer, no lot number). Stock solutions of
potassium acetate and tetramethylammonium acetate were
prepared by titrating standardized aliquots of the respec¬
tive hydroxides with 1M acetic acid (Baker Chemical, lot:
41397). Magnesium acetate solutions were made from
magnesium acetate tetrahydrate (Mallinckrodt, control:
NSB). Solid calcium acetate was made by dissolving calcium
carbonate (Mallinckrodt, no lot number) in excess acetic
acid. The calcium acetate was recrystallized twice from
aqueous ethanol, washed repeatedly with 95% ethanol, and
dried for twenty-four hours at 110°C to yield the mono¬
hydrate of the salt. The sample was found to contain 97.3%
Ca(C2H3O2).H2O, by dissolving a known amount of the acetate

49
in water, precipitating the calcium as calcium carbonate
and titrating the carbonate with hydrochloric acid.
Strontium acetate was prepared from its carbonate (Matheson,
Coleman & Bell, lot: 19) and assayed in a manner similar
to that for calcium acetate. It was determined to contain
95.0% SriC^H^C^). A recrystallized sample of barium
acetate was dried for twenty-four hours at 170°C and
assayed by forming barium carbonate. The solid carbonate
was reacted with excess 1.00N sulfuric acid, carbonic acid
removed by gentle heating, and the excess sulfuric acid
determined by titration with standard sodium hydroxide
solution. The assay of the barium acetate was 99.6%
Ba (C2H302).
Solutions of the supporting electrolytes to be used
as solvents were made from the stock solutions. Aliquots
of the stock solutions were diluted to within 90% of their
calculated final volumes and their acidities adjusted to
pH 5.90 + 0.05, if necessary, using the appropriate metal
hydroxide, phosphoric acid, or acetic acid. They were
then diluted to their exact desired final volumes and their
ionic strengths computed, taking into account any additional
ions added to adjust the acid levels.
A commercially available sodium salt of DNA was
dialyzed against a solution of the desired metal phosphate
or acetate to obtain the DNA salts having the same cation as
that of the supporting electrolyte. Solutions containing
approximately 0.01 mole of DNA phosphate per liter were made

50
by placing 3.6 mg of DNA, sodium salt (Calbiochem calf
thymus DNA, A grade, lot: 900007: 8.04% P; 12.21% N;
1%
moisture, 14%; E260nm = Per of solvent
into twenty milliliter volumetric flasks. Between 5 and
20 ml of solvent were then added. The solvent was a 0.01M
solution (pH 5.90) of the metal phosphate or acetate. The
flask and its contents were placed on a vertical rotating
mixer (Scientific Products model 150-V) operating at 6 rpm
for about one hour, after which the sample was placed in
an ultrasonic mixer for two to three minutes. Rotation
and sonication were repeated for a total of at least five
hours or well beyond the time at which the moderately
viscous sample appeared homogeneous. The sodium DNA solu¬
tion was then transferred to dialysis tubing (Union Carbide
Films Packaging Division, 7.3 mm dia.). The tubing and its
contents were immersed in a volume of 0.01M solution of the
desired salt which was at least ten times greater than the
volume of DNA solution being dialyzed. The external solu¬
tion was agitated for two hours and then replaced. A
total of six volumes of external solution were used for
each dialysis procedure. The dialyzed DNA was again mixed
for one hour on the rotating mixer and submitted to ultra-
sonication for two to three minutes. Then, the sample was
R
divided among Pyrex culture tubes (6x50 mm) each
receiving approximately one half milliliter of sample.
These were immediately stored at -5°C.

51
Determinations of the concentrations of DNA after
dissolution and after dialysis yielded values which were
identical, within experimental error and also within + 5%
of the calculated concentration based on the manufacturer's
data. The agreement of these data were taken as evidence
that any degradation of the biopolymer during the pro¬
cedures was minimal.
The efficiency of the dialysis method, used here, to
replace the sodium cations of the original DNA salt with
the cations of the external solution was evaluated by
atomic absorption spectroscopy. A Perkin-Elmer model 290B
atomic absorption spectrophotometer outfitted with a combina¬
tion sodium/potassium lamp was used to measure the residual
concentration of sodium ion in a sample of sodium DNA which
had been dialyzed against a potassium phosphate solution.
Instrument readings were taken using the contents of the
dialysis bag and compared to those obtained for standard
solutions in which the ratio of potassium to sodium ions
was approximately one thousand. Note that in this instance
high relative concentrations of potassium ion must be included
in the standard solution. Less than 0.1% of the original
sodium remained within the dialysis bag when the above
procedure was employed.
Crystalline 3-aminoacridine was prepared by Dr. Timothy
Roy using a standard procedure (80) except for additional
purification of the product by vacuum sublimation. Thin
layer chromatographic analysis of the purified product

52
using a variety of solvents and solid supports failed to
indicate the presence of any impurities. The spectra of
solutions of 3-aminoacridine at pH 11.0 and at pH 2.5 had
the same relative absorptivities as those reported by
Albert (75). Reported (pH 11.0) (log e): 237 (4.46),
262 (4.83), 321 (3.35), 337 (3.65), 353 (3.92), 410 (3.79).
(pH 2.5) 233 (4.62), 274 (4.65), 349 (4.03), 365 (4.15),
454 (4.10). The melting point agreed with that reported
(81): 218°C, uncorr., 223.5°C, corr.
-3
Stock solutions of 1 x 10 M 3-aminoacridine were
prepared by dissolving 2 mg of the solid in 7 ml of absolute
ethanol. The solutions were stored at -5°C in glass. One
hundred microliter aliquots of the ethanolic solutions were
used for each titration. Periodic tests of the stock
solutions using thin layer chromatography showed no observ¬
able degradation during the time periods in which the samples
were being employed.
Absorptiometric titrations were conducted by successive
additions of microliter aliquots of DNA stock solution to
10 ml solutions of 3-aminoacridinium cation. After each
addition, the absorbance spectrum between 550 nm and 320 nm
of the solution was recorded. To prepare the solutions, 10 ml
of supporting electrolyte at a given concentration were
measured in a 10 ml volumetric flask. This volume was de¬
termined when the liquid was at the same temperature as it
was to be during the titrations (e.g., 15.0, 25.0, or 35.0°C).
Immediately prior to a titration, 0.1 ml of stock 3-amino-

53
acridine in ethanol was added to the flask, the contents
mixed and then transferred to a 4.00 cm quartz cell (Bolab
Inc., Derry, New Hampshire). The cell and its contents
were placed in the thermostated compartment of the
spectrophotometer (Beckman model DB-GT) and allowed to
become thermally equilibrated. The temperature of the
instrument's cell compartment was maintained at 15.0 +
0.2°C, 25.0 + 0.05°C or 35.0 + 0.2°C by circulating thermo¬
stated water (Aminco refrigerated water bath) through its
external jacket.
The spectrum of the sample was recorded after which
an aliquot of stock DNA solution was added. Syringes (Uni¬
metrics, Inc.) were employed to dispense volumes of DNA
solution ranging from 4 yl to 100 ul while volumes greater
than 100 Ul (generally added toward the ends of the
titrations) were measured with a 100 microliter micro¬
pipette (Centaur Chemical, Danbury, Connecticut). The
individual culture tube of the stock DNA solution from
which aliquots were being withdrawn for a given titration
was maintained at 25°C, regardless of the temperature at
which titration was being conducted. After adding the DNA
to the 3-aminoacridinium solution in the cell, the solution
was mixed by withdrawing about 5 ml of the cell's contents
into a serological pipet, then rapidly forcing them back
into the cell. This was done three or four times after
each addition of DNA.

54
Concentrations of the stock DNA solutions were
determined by measuring the absorbance of a solution
containing 50 ml of stock DNA in 10.00 ml of 0.010M
supporting electrolyte at pH 5.90 (1.00 cm pathlength).
The concentrations of at least ten samples, each taken
from separate tubes, were calculated using the molar
4 -1 _ i
absorptivity at 260 nm of 1.412 x 10 M cm (75). All
glassware was cleaned by immersion into concentrated
sulfuric acid saturated with potassium dichromate. Spec¬
trophotometer cells were rinsed with water followed by
ethanol and were stored in 50% concentrated hydrochloric
acid in ethanol to prevent accumulative absorption of the
dye or of DNA to their surfaces.
Solid 7-aminoquinoline was a gift from Dr. D. Jackson,
Texas Tech. University, Lubbock, Texas, and was determined
to be at least 98% pure according to molar absorptivity
values (82-84). Reported (pH 12) ^nm (log e): 240 (4.56),
max
275s (3.58), 335 (3.68), 346 (3.64). (pH 2.0): 260 (4.44),
285s (3.57), 392 (3.90). The reported pKa of the compound
_ 3
is 6.65 (84). A stock 3.4 x 10 M solution in absolute
ethanol was prepared and 30 yl aliquots added to 8.00 ml
of 0.10M KH2PO4 buffer solution at pH 5.9. The 7-amino-
quinolinium solutions were titrated at 15.0 + 0.2°C, 25.0 +
0.05°C, and 35.0 + 0.2°C using the same equipment and pro¬
cedures as for the 3-aminoacridinium titrations, vide
supra.

55
Calculations
The molar absorptivity of the 3-aminoacridinium-DNA
complex, at 368 nm was calculated from absorbance
spectra of solutions in which essentially all of the
compound was in the bound form. Complete binding (>_ 99%
of all drug present) was assumed when at least three suc¬
cessive additions of excess DNA to the solution resulted
in no change in the total absorbance, corrected for
dilution. It was further presumed that, under these
conditions, all absorbance at 368 nm was due to the complex
and no 3-aminoacridine existed, either as the neutral or
monocationic, unbound form. Also, only solutions containing
monovalent metals (and tertamethylammonium cation) at low
ionicities were used since, for other systems, the large
amounts of DNA needed to drive the equilibrium toward
products caused serious light scattering and, in some
instances, precipitation of DNA salts. Calculated values
of in the presence of each type of electrolyte are
presented in Table 1. The averaged value of 8494^ M ^cm ^
was used for e in the presence of all supporting elec-
trolytes at all concentrations. A similar procedure was
used to evaluate the molar absorptivity of the 7-amino-
quinolinium-DNA complex, at its band maximum (405 nm).
At pH 5.9 the ratio of neutral to protonated 7-amino-
quinoline is 0.178. It was assumed that only the cationic
form of the drug binds to DNA. The averaged molar absorp¬
tivity of the 7-aminoquinolinium-DNA complex was evaluated

56
using three simultaneous equations and was found to be
5332^ M ‘'’cm ^ at 405 nm.
Concentrations of free 3-aminoacridinium cation,
[BH], and bound monocation, [BHP], were calculated from
absorptiometric data assuming simple additivity of the
absorbances of the two species at the analytical wave¬
length
A = eBH£[BH] + ?BHp¿tBHP] (2-1)
and considering the mass balance expression
Ct = [BH] + [BHP] (2-2)
where
A = the total absorbance of the solution at a given
wavelength.
e = the molar absorptivity of the free species,
tín
1.413 x 103 M-1cm_1.
£BHP = mo^ar absorptivity of the bound complex,
8.49£ x 103 M ^cm ^.
H = pathlength of light in the cell.
= the total concentration of free and bound species.
[BH] = the molar concentration of free 3-aminoacridinium
cation.
[BHP] = the molar concentration of bound 3-aminoacridinium-
DNA complex.
Solving equation (2-2) for [BH], substituting into equation
(2-1), and rearranging yields

[BHP]
(2-3)
A-CBHCt*
£BHP£_eBH¿
Before any DNA has been added, the total absorbance of
the solution will be due entirely to free BH and, since
BH at this point will be equal to C^, the total absorbance
before any titration, Aq, may be expressed as
Ao = EBhV • (2-4)
Similarly, when all of the compound has been bound, the
final absorbance, A , is equal to
A
00
eBHPCt*
(2-5)
and for any point in the titration, where the total absor¬
bance is A^_,
[BHP] = [(At-Ao)/(Aoo-Ao)]Ct (2-6)
and
[BH] = [l-(At-Ao)/(Aoo-AQ) ]Cfc . (2-7)
Since an appreciable amount of the total 7-amino-
quinoline exists in the solution in the neutral form at
pH 5.9, the equations representing the protolytic equilibrium,

58
the conservation of mass for 7-aminoquinoline,
Ct = [A] + [AH] + [AHP] , (2-9)
and the photometric absorbance at any point in the
titration of the small molecule with DNA,
At = £a[AU + £AH[AHU + £AHP[AHP]£ (2-10)
must be employed to determine the concentrations of the
species involved in the binding. Concentrations of drug
species are, explicitly,
[AHP] =
(eA-^- + £ah+> V -V1 + VihV
K
{eA~TT + eAH " eAHP(1 + Ka/[H+])}£
(2-11)
[H' ]
At ~ £AHP Ct£
[AH] =
(2-12)
K
{£ a
A[h+] + £AH eAHP(1 + Ka/[H+])}£
[A] =
(At GAHP Ct£) Ka[H+]
(2-13)
K
{eA~+7 + eAH " eAHP
LH J
K
[H+]
â– )H
(1 +

59
It was also useful to directly calculate the ratio,
[AHP] (e
+ eAH+) V - A(1 + Ka/[H+1)
(2-14)
[AH]
At " £AHP Ct£
where
= dissociation constant for the drug or probe,
2.24 x 10~7 M-1.
[H+] = hydronium ion concentration, 1.26. x 10 ^M.
eAHP = ra°lar absorptivity of the bound monocation at
the specified analytical wavelength, 5.332^ x
10"^ M "'"cm ^ at 42 0 nm.
£AH = mo^ar absorptivity of the free monocation at
the specified analytical wavelength, 7.943^ x
10^ M ^cm at 42 0 nm.
e = molar absorptivity of the free base at the
2 -1
specified analytical wavelength, 3.47 x 10 M
cm at 392 nm.
SL = absorption cell path, 4.00 cm.
At = the absorbance at any point during the titration
at the specified analytical wavelength.
= total concentration of drug or probe.
For each point in a titration the total absorbance was
determined along with appropriate constants which were
substituted into the necessary relationships to yield the
concentrations of free and bound species. Raw data for

60
all titrations are presented in Tables 2 through 17. Note
that volume corrections for the addition of DNA solution
to the sample were made in all cases.
The equilibrium association constants for the
surface and intercalative modes of binding were calculated
for each point in the absorptiometric titrations of the
small molecules with DNA, using the relationships developed
by A.C. Capomacchia and S.G. Schulman (85).
(2-15)
Ks = [BHP]/[BH]([Pt/3]-[BHP])
and
2
(2-16)
KI = [BHP]/[BH]([Pfc/2]“2[BHP])
where Kg and are the apparent association constants
for the surface and intercalative modes, respectively, and
P is the total molar DNA phosphate concentration. The
equations for the binding of 7-aminoquinolinium to DNA
are the same except [AHP] and [AH] are substituted for
[BHP] and [BH] , respectively. The extended Debye-Hiickel equation
(1-28) was used to calculate activity coefficients at the
various ionic strengths, for charged species involved in
the association equilibrium.
Apparent association constants were corrected for
deviations of molar concentration from activity, by
(2-17)

61
where K1 is the corrected apparent association constant,
and a are activity coefficients for bound
BHP BH p
species, free species, and DNA phosphate, respectively.
The value of n is 1 for surface binding and 2 for inter-
calative binding. For the binding of 7-aminoquinolinium
to DNA, the subscripts BHP and BH are replaced with AHP
and AH corresponding to bound and free 7-aminoquinolinium,
respectively. These, and other calculations, were done
using a Litton-Monroe model Beta 326 Scientist computer
equipped with a tape cassette drive (Litton-Monroe,
Orange, New Jersey). The simultaneous solution of (2-8),
(2-9), and (2-10) to obtain concentrations of A, AH, and
AHP was accomplished using a program supplied by the
computer manufacturer. Programs for specific calculations
used here were designed and are presented in Appendix III.
Thermodynamic parameters for the binding of 3-amino-
acridinium and 7-aminoquinolinium to DNA were evaluated
using equation (1-31). Values of log K at 15°C, 25°C, and
35°C for surface and for intercalative binding of each
species were determined. These values were then plotted
vs. 1/T (K). The slopes of the lines, equal to AH°/2.303R,
were used to calculate the standard entalpies of binding.
Values for the standard entropies of the reactions were
evaluated using
AS° = (AH° - AG°)/T
(2-18)

CHAPTER III
RESULTS AND DISCUSSION
General Spectral and Titration Characteristics
A number of interesting spectral features and
titration characteristics are common to all of the
systems investigated. Space limitations preclude inclu¬
sion of all of the absorptiometric titration spectra,
though four which are representative are presented
(Figures 3 through 6). These are spectra of titrations
done in the presence of the background cation having the
lowest charge density, Cs+, and the greatest charge density,
Mg++, at the highest and lowest concentrations of each.
As a rule, the lower the charge and charge density, of
the background cation, the less effect it appears to have
on the two modes of binding. The highest concentration of
the smallest divalent cation represents the other extreme,
as it severely inhibits binding processes. The rest of
the systems are intermediate between these two. In each
of the figures, only 7 curves are drawn for the sake of
clarity, though usually between 18 and 22 spectra were
recorded for any given titration. For the curves ommitted
from figures 3 through 6, the absorbances at of the
62

63
1Lb bands may be obtained from the appropriate tables in
Appendix II.
In all cases, addition of DNA to solutions of 3-
aminoacridinium cation resulted in decreases (hypochromisms)
in the intensities of both the ^L, and bands. It is
D cl
noteworthy that even small amounts of DNA (< 1 mole DNA
per mole of drug) added to 3-aminoacridinium solutions
produced only hypochromic shifts. However, addition of
small amounts of DNA to other compounds such as 7-amino-
quinolinium causes an initial increase in the total intensity
(hyperchromism) of the drug's spectrum. Hypochromism sub¬
sequently occurs when DNA in excess of about 1 mole per mole
of drug is added. Presently, it is not known what factors
are responsible for the initial hyperchromism of some species'
spectra or why some exhibit initial hyperchromism while
others exhibit only hypochromism throughout their whole
titration, though significant variations in binding behavior
appear highly probable.
It may be seen from Figures 3 through 6 that addition
of increasing increments of DNA to 3-aminoacridinium results
in red (bathochromic) shifts of both bands of the drug's
spectrum. The bathochromism of the ^A and ^L.^' ^A
b b
bands are a result of decreases in the energy separations
between the ground and excited states. This may be due to
a lowering of the '*'L states, a raising of the ground, state,
or a combination of both. If binding of small molecules to
DNA occurs as the reaction proceeds, then it is reasonable

64
to hypothesize that the cationic drug is moving from a
highly polar, high dielectric medium (water) to a less
polar, more lipophilic solvent such as the region in, and
around, the DNA helix. An environmental change of this
type would tend to favor a decrease in charge separation
*
within a molecule and, thus, favor ir —> tt transitions
in which the formal positive charge of the heterocyclic
ring nitrogen is decreased. Such a case would mean that
a lowering of the energies of the excited states is the
predominant factor in causing the red shift upon binding
if, indeed, such "solvent effect" arguments are valid when
considering binding processes. Caution must be exercised
in this regard as Capomacchia (69) has found that, for a
number of compounds, spectral shifts upon binding to DNA
are not in qualitative agreement with shifts observed when
the compounds were transferred from polar aqueous solvents
to less polar organic solvents. Maximum shifts of the two
bands are from 362 nm to 368 nm for the ^L, band and from
b
454 nm to 463 nm for the band. These differences (6 nm
a
and 9 nm, respectively) correspond to decreases in the
transition energies of 450^ cm 1 (1.29 Kcal/mole dye) and
428^ cm (1.23: Kcal/mole dye). The similar magnitudes show
that, with excess DNA at low total ionic strength, the two
transitions are equally affected from an energetic stand¬
point in the presence of excess DNA. However, a large
number of titrations have substantiated the observation
that the moves only slightly, from 362 nm to == 364 nm,

65
over most of the titration and then abruptly shifts from
= 364 nm to 368 nm while the shift of the band is
a
smooth throughout the whole titration. The abrupt shift
of the takes place at approximately the point in the
titration where any additional DNA results in no further
decrease in band intensity. In fact, this phenomena
has been accepted as evidence that all of the protonated
3-aminoacridine has been bound. Moreover, the rapid shift
of the higher energy band always occurs significantly
beyond the point in the titration where very little change
is noted in band intensity or position of the maximum of
the longer wavelength band. These observations suggest
that the environment appears to affect both transitions
equally in the presence of excess DNA, but they are affected
differently during the course of a titration. A cursory
consideration of the data may lead one to postulate that
the abrupt shift is due to intercalation of the molecule.
This is plausible since internal binding would affect the
^"L, more dramatically than the especially if the charged
ring nitrogen remained exterior to the helix. Further
consideration, however, reveals substantial evidence, both
in this work and in the literature, which shows that inter¬
calation occurs much earlier in the titration and that it is
much less abrupt a process than the bathochromic shift
considered here would require. The abrupt change might
arise from some form of near-neighbor interaction of bound
drug molecules (63). During most of the titration, where

66
the total DNA phosphate to total 3-aminoacridine ratio,
P/D, is low, the bound molecules can interact with one
another. But as the ratio increases, bound species become
isolated from each other to the point of no longer being
able to interact. The hypochromic shift is not incon¬
sistent with this proposal as both red and blue shifts
are possible upon exciton formation. Another factor to
be considered is that as the DNA concentration is increased
aggregation of the DNA, itself, may begin to occur. At
some critical level, the bound dye may experience a marked
decrease in the hydrophilicity of its environment, perhaps
as a result of the neighboring polymers excluding water
between themselves. Under these circumstances, red shifts
would be expected, though why they would be seen only for
the '*'L, and not for the 1L transition is unanswered,
b a
The onset of the rapid red shift appears to be depen¬
dent on the total ionic strength of the medium as well as
the type of cations in solution. In the presence of
large, diffusely charged, monovalent cesium a distinct
bathochromism occurred, even in the highest concentration
(0.15M) at the moderate P/D ratio of 65. On the other
— 4 ++
hand, 6.3 x 10 N Mg was sufficient to prevent the shift
to 368 nm even when phosphate is in excess of drug by a
factor of 65. In 0.0025M CsI^PO^, a P/D ratio of only 14
results in the spectral shift (in fact, it probably occurs
at- about P/D = 11) . Charge density of the countercations
influenced the point at which the abrupt red shift was seen

67
in a manner parallel to the degree of completion of a
titration for a given P/D ratio. The overall spectral
behavior of the systems in the presence of the alkaline
earths was the same as for the alkalies except the
maximum concentrations of M++ that allowed at least 90%
binding at P/D < 100 was much lower than for alkali
metals. Concentrations of group Ila cations in excess
of about 0.02N resulted in precipitates forming in the
samples. These were, presumably, the metal salts of DNA
and appeared long before sufficient polymer was added to
appreciably bind the drug.
Throughout this discussion, the absorbances at the
band maxima have been assumed to be directly related to
the areas under the curves. Since it is the areas, and
not peak heights at the maxima which are true measures of
transition probabilities, care must be taken in equating
peak heights with relative transition probabilities. If
the overall geometry of the band were to change during
the course of a titration, peak heights would not be valid
representations of the progress of the titration. In our
studies, inspection of the bands showed that their geome¬
tries were constant during the titrations, hence, we feel
approximating peak heights with peak areas was valid.
In addition, we have assumed that the 0-0 vibronic
band maximum of the ^L, < ^A transition is coincident
b
with the maximum of the total transition envelope. This
may be misleading if the maximum of the envelope arises

68
from the sums of the intensities of vibronic transitions
which are close enough together to appreciably overlap
at the maximum wavelength. If the vibronic energies are
altered in dissimilar ways during a titration, then the
maximal envelope intensity will reflect this dissimilar
change along with any real shift in the 0-0 band. It is
possible, therefore, that shifts of the envelope
maximum may be due to either a disproportionate change in
the magnitudes of the band maxima of the 0-0 and 0-1 vibronic
bands or a bona fide, pure 0-0 shift.
Variations in the spectral characteristics of the 3-
aminoacridinium-DNA systems in the presence of metal phos¬
phates and metal acetates were not due to the different
counteranions, as evidenced by the results of titrations
using KH2P0^ and KC^CCH^ as background electrolytes. Re¬
sults obtained with the two buffer ions were the same in
all respects except at very low ionic strengths. Deviations
in this region are probably due to lack of sufficient pH
control or alterations in the conformation of DNA possibly
arising from mild denaturation.
The envelope of the free 3-aminoacridinium cation
consists of two distinct vibronic bands having maxima at
348 nm and 362 nm, corresponding to the 0-1 and 0-0 vibronic
transitions, respectively. As a titration proceeds, the fine
structure of the envelope is lost as a result of the 348 nm
peak becoming increasingly less pronounced, to the point of
becoming a mere perturbation of the overall envelope. It is

69
difficult to determine the exact position of the shoulder
when the titration is nearly complete, but it appears that
its shift is essentially the same as that of the 362 nm
band. In cases where the abrupt 364 nm to 368 nm shift
of the v.. vibronic band of the ^L, occurs, there is an
accompanying shift of the shoulder maximum from = 350 nm to
about 354 nm with a slight, but definite increase in the
resolution of the two bands. The degree of fine structure
of a compound's spectrum provides qualitative information
regarding the environment of the species. Loss of fine
structure may be rationalized on the basis of increasing
the degrees of vibrational freedom of the molecule. If,
upon binding, the energy separation between the v = 0 and
v = 1 vibrational levels of the ground state were to be
reduced, a coalescing of the two peaks would be anticipated.
Such a situation may arise when a molecule moves from an
environment in which its vibrational motions are restricted
to one in which they are less so. For this to maintain in
our systems, the bound species would have to have greater
vibrational freedom than the wholly water-solvated free
cation. This is feasible considering the strong degree of
interaction between water molecules and any dipole in their
midst compared to the lipophilic, noninteracting "solvent"
of the surface and interior of DNA.
The increase in fine structure which accompanies the
dramatic red shift of the 0-0 band of the envelope at
high DNA concentrations is not well understood but serves

70
to corroborate evidence for a dramatic environmental change
of bound drug at these very high levels.
To summarize, titrations of 3-aminoacridinium cation
with DNA salts in the presence of monovalent and divalent
metal ions (including tetramethylammonium ion) of varying
concentrations result in hypochromism only. In all cases,
both the and envelopes of the free drug's spectrum
red shift, at least to some extent, during titration —
the from 454 nm to 463 nm and the from 362 nm to
a b
about 364 nm. In the presence of all concentrations of
supporting electrolytes having large, monovalent cations,
a further, sharp red shift from 364 nm to 368 nm takes place
when a P/D ratio of 10 to 15 is reached. The shift may be
effected in the presence of smaller alkalimetal ions at
P/D ratios approaching 70. Regardless of the concentration
or charge density of divalent cations, the dramatic
shift did not occur. Loss of vibronic fine structure of
the envelope was observed in all cases as the titrations
progressed, up to the point at which the dramatic batho-
chromic shift occurred. Concomitant with the rapid shift
was a reemergence of the 0-1 vibronic band.
Using relationships developed in the Experimental
section, concentrations of free and bound 3-aminoacridinium
cation were computed from the total absorbances of solutions
which contained varying amounts of DNA. Data for all of
the titrations are not included here, though some, which
include the extremes of experimental conditions, are presented.

71
Tables 18 through 21 contain tabulations of [BH] and
[BHP] along with log ([BHP]/[BH]) values for reactions in
0.15M and 0.0025M Csí^PO^ and in 0.010N and 6.3 x 10
MgiC^CCH^^- In lower concentrations of alkali metals
and tetramethylammonium ion, the addition of as little
-3
as 5 yl of 9 x 10 M DNA phosphate solution results in a
significant degree of binding of the drug. For instance,
nearly 10% of the total 3-aminoacridinium present in
0.0025N CsI^PO^ is bound after addition of only 5.0 yl of
DNA. Since volumes of DNA less than 5 yl cannot be
measured with much reliability due to its viscosity, the
ratio of bound to free drug of 0.1 should be considered
as the effective minimum for titrations of this sort. When
lower ratios were possible, such as in solutions containing
divalent cations or >_ 0.010M monocations, random experi¬
mental error was sufficiently high to make the data unreliable.
As a result, our studies indicate that absorptiometric
titrations of small molecules with DNA, under conditions
similar to ours, are probably invalid whenever the ratio of
bound to total drug is less than 0.1. It should be recalled
that our use of 4.00 cm cells maximized sensitivity, thereby
allowing minimal increments of DNA solution to be used.
Smaller amounts of DNA could be added in 5 yl increments by
employing a more dilute stock solution, but the errors arising
from other sources would likely be unimproved. At the other
titration extreme, ratios of bound to free drug in excess of
about 0.95 also result in unacceptable randomness in most

72
cases. Our procedures are designed to minimize the impact
of these sources of imprecision on the calculated association
constants.
The maximum extent of total binding which could be
achieved varied with the character and concentration of
the supporting electrolyte. Table 22 lists the percent
of total drug bound to DNA after the final increment of
DNA was added in titrations done in various media. The
P/D values listed are the maximum that could be obtained
due to solubility and pH-control limitations. No effort
was made to determine how the drug was bound or what per¬
centage of the total was surface bound compared to inter-
calatively bound. As discussed previously, for those titra¬
tions in which an abrupt bathochromic shift of the band
was observed, the binding was considered to be complete
(100 + 2%). The averaged molar absorptivity, entJT_, determined
from these titrations was used to compute concentrations of
free and bound species for all systems. In light of the
experimental error incurred in evaluating and the
simple fact that thermodynamic equilibrium demands that
there cannot be total binding, percentages greater than
about 96% are taken as indicative of complete binding. From
Table 22, it may be seen that at least 97% of the total 3-
aminoacridinium present is bound in solutions of < 0.010M
MI^PO^. To attain essentially complete binding in these
systems only about 10 to 20 times as much DNA phosphate as
total drug need be present (see P/D column of Table 22). As

73
expected, larger and larger ratios of DNA to drug are
necessary for complete binding as the concentration of
a given electrolyte is increased. Similarly, it becomes
increasingly more difficult to shift the equilibrium wholly
toward bound product in the order (of electrolyte cations):
(CH^J^N* £ Cs+ required for 97% binding in 0.0025N Cs+; 17-fold was needed
for 97% binding in 0.010M Cs; while, even with a 65-fold
excess of DNA, only 90% of the drug was bound in 0.15M
CsI^PO^. It was relatively easy to totally bind the 3-
aminoacridinium cation in low concentrations of monovalent
cations but very difficult to attain greater than 90%
binding in even the lowest concentrations of the divalent
— 4 ++
electrolytes. In 6.3 x 10 N Ba (the dication having
the lowest charge density of those investigated) only 90%
of the drug was bound at a P/D ratio of 47 and, for Mg++
at the same concentration, 90% was bound at P/D = 65.
Similarly, in 0.010N ca++ and 0.010N Mg++ the maximum
degrees of binding were 77% and 70%, respectively, with
P/D ratios of 103 and 104. The choices of the upper limits
of total phosphate concentrations reported here were de¬
pendent on the solubility of species in the concentrated
electrolytes and on pH considerations for the dilute systems
Visible precipitation or erratic spectral behavior were
evident in solutions having P/D values greater than those
listed in Table 22 for barium and magnesium. The maximum
concentration of the divalent electrolytes was set at 0.010N

74
since obvious precipitation occurred in solutions of
0.025N M well before the titrations were complete (i.e.,
at 30% to 50% binding). Acetates of the alkaline earths
were used instead of their phosphates due to the latter's
very low solubilities. Despite dialysis of DNA solutions
against pH 5.90 buffer, their pH levels were still ap¬
parently close to 7. Thus, addition of moderate volumes
of DNA to the very poor acetate buffer systems caused
excessive increases in the pH levels of the dilute buffer
solutions. Consequently, acetate solutions having more
-4
than approximately 5 x 10 M DNA phosphate had pH values
considerably in excess of 5.90.
Upon perusing the second column of Table 22, it becomes
quickly evident that the extent of binding is not predomi¬
nately and directly related to ionic strength. Total ionic
strengths were computed for the systems having the P/D
values listed using equation (1-29) and Program 4 of
Appendix III. Notice that while the ionic strengths of
0.0025M CsH2P04 and 6.3 x 10 4N BaiC^CCH^^ are the same,
the binding is complete in the former and only 90% complete
in the latter. This shows that the effect of the formal
charge on the cations of the electrolyte is greater than
predicted on the basis of the simple electrostatic effects
incorporated in the limiting Debye-Hiickel equation. That
the contribution of excess DNA to the total ionic strength
is not responsible for these large differences may be seen
by comparing 0.010N barium and 0.010N magnesium solutions

75
which both contain the same amount of DNA and have, there¬
fore, the same total ionic strengths. The extent of
binding in the former is 77% while in the latter it is
only 70%.
Rapid estimations of the values of m and q of equation
(1-21) may be obtained by selecting reasonable terms such
that mq = n (n is total phosphate per drug). It has been
shown that (85) in a medium containing 0.002M to 0.005M
Ki^PO^ (pH 5.9) the binding of 3-aminoacridinium cation to
DNA is best expressed by
, ->
BH + (P04)3
bh(po4) 3
(3-1)
for
the surface mode, and
BH+ + 2(P04)3
BH((P04)2)^-
(3-2)
for
the intercalative mode.
In equation (3-1),
three DNA
phosphates act as a single moiety, reacting with the small
molecular cation to form the surface-bound species. In
equation (3-2) , 2 phosphates act as a single reacting
entity with two of these, in turn, interacting with one
cation to form the intercalated complex. The values of m
and q are 3 and 1, respectively, for the first case and 2
and 2 for the second. Substituting these into equation (1-21)
yields
Ks = [BHP]/[BH] ( [Pt/3]-[BHP])) (3-3)
and

76
Kj = [BHP]/([BH]([Pt/2]-2[BHP])2) . (3-4)
Ks and are the apparent association constants for the
surface-bound and intercalatively-bound drug-DNA complexes,
respectively. These equations were tested to certify their
validity under our experimental conditions. Recall that
plots of log ( [BHP]/ [BH] ) vs. log ( [P^m] -q [BHP] ) yield lines
whose slopes are equal to q and whose intercepts are equal
to log K. Tables 18 through 21 include calculated values
of log ( [BHP]/[BH]), log([Pt/3]-[BHP]), and log([Pfc/2]-2[BHP])
for representative titrations. Figures 7 and 8 are plots
of log([BHP]/[BH]) vs. log([Pty3]-[BHP]) for the reactions
in the presence of 0.0025M Cs+ and 6.3 x 10 Mg++; and in
0.15M Cs+, 0.010N Mg++, and 0.010M Cs+, respectively.
Figures 9 and 10 are plots of log ( [BHP] / [BH] ) vs. logUP^.^]-
2 [BHP]) for the same systems. Three general features are
particularly significant. One, the shapes and slopes of
the curves are essentially the same whether log ([P^^] - [BHP] )
or log( IP1“2[BHP]) is plotted against log([BHP]/[BH]).
(Calculated slopes are tabulated in Table 23.) Only in the
latter portions of the titrations done in the more dilute
alkali metal electrolytes does this rule begin to break down.
Two, except for 0.15M Cs^PO., all plots of the data in¬
volving the presence of monovalent cations exhibit definite
curvature — the lower the concentration, and further down
a series of the periodic chart, the greater the curvature.
2+
Conversely, all concentrations of all M salts yield

77
essentially straight lines over the whole titration
region. Three, in those cases in which there is no
curvature throughout the titration range, the slopes of
the lines are all nearly 1 and parallel the initial
regions of the other titrations.
That the values of q evaluated from both types of
plots are much the same in most media is evidence of the
validity of this technique. It appears that going from
q = 1 to q = 2 or 3 does not substantially affect the
values of log([P^y^]-q[BHP]) except at the extrema of the
titrations. With our procedure, however, data very early
and very late in a titration need not be relied upon when
estimating lines, unlike, for instance, the Scatchard
method. For cases where the background electrolyte is
especially weak (i.e., low concentrations of cations having
low charge density) differences in the slopes of the two
plots become appreciable for the later portions of the
titrations. For example, the final slopes where 0.010M
Cs+ was employed as the background electrolyte are 1.68 and
1.88 for q = 2 and q = 1, respectively, and for 0.0025M Cs+,
the values are 1.71 and 2.16. Several explanations may be
advanced for this lack of agreement, including the fact
that the relatively short overall range of the amount of
DNA needed to complete the titrations magnifies the error
inherent at the extrema. Also, the high degree of curvature
makes the determination of the two separate straight line
regions more subjective. Perturbations of the reaction

78
system which are significant only in low ionic strength
media (e.g., mild polymer denaturation) would be expected
to affect both types of plots in the same manner. Re¬
gardless of the sources of the discrepancies, the pro¬
cedure is still useful, as both plots yield the same
integer values of q. Final selection of m and q is made
on the basis of the constancy of the computed values of
the equilibrium association constant (vide infra).
The presence of at least two distinct linear regions
of the plots for lower concentrations of alkali metals
strongly indicates two binding stoichiometries. Numerous
investigators, as outlined in the Introduction, have shown
the presence of two types of binding processes which likely
account for the different values of q. What has not been
previously reported is the apparent change in the ratio of
phosphate to 3-aminoacridinium cation on going from one
type of binding to the other as evidenced by these plots.
Also, as the ionic strength is increased, the extent of
intercalative binding apparently decreases. This is mani¬
fested in Figures 7 through 10 by a progressive trend
toward a wholly linear plot as the ionic strength increases.
That is, compare the shapes of the curves for systems con¬
taining 0.0025M, 0.010M, and 0.15M CsI^PO^. Similarly, only
. 2+
one stoichiometry is evident from the plots for M -containing
systems, which are linear throughout their entire titration
regions. Slopes of 1 + .1 for all of these, including that
for the reaction in the presence of 1 x 10 N Ba(02CCH3)2,

79
indicates that only surface binding occurs in the
presence of the divalent metal cations over the titration
range investigated.
Previously, it was brought out that, in the presence
of low concentrations of monovalent cations as background
electrolytes, plots of log([BHP]/[BH]) vs. log(-q[BHP])
exhibit two regions of linearity. During the early part of
a titration the slope of the lines is approximately 1 while
in the later region the slope is about 2. As the ionic
strength is increased, or as the charge density of the
cation of the electrolyte becomes greater, the slope of
the later region decreases, eventually becoming equal to
about 1. Slopes of the initial portions are essentially
unaffected by alterations in the supporting media in which
the titrations are conducted.
The ramifications of these relationships may be explored
by first defining the binding interaction in terms of con¬
ventional metal-ligand chemistry. We shall analogize the
cationic 3-aminoacridinium to the metal center and the DNA
phosphates to ligands which bind to the center. For clarity,
we shall further define two types of ligands, having indepen¬
dent and variable affinities for the cation centers; a weakly
bound ligand (corresponding to surface binding) which forms
a 1:1 complex, and a more strongly bound ligand (corresponding
to intercalative binding) which forms a 1:2 metal:ligand
complex. These definitions are based on the values of q for
surface-bound and intercalatively-bound ligand types being 1

80
and 2, respectively. Finally, we shall assume that the
ligands compete for metal binding sites such that if a
cation is bound in the surface mode, then it is unavailable
for intercalative binding, and vice versa.
During the initial stages of a titration there is a
large excess of 3-aminoacridinium over DNA phosphate or,
by analogy, an excess of metal over ligand. In such a
situation, there would be minimal competition for metal
sites and both forms of binding would be expected to occur
to their fullest extent, based on their respective association
constants. We may view the titration of the cation with DNA
ligand as a decrease in the metal center concentration as
the process proceeds. Addition of aliquots of DNA serves
to limit the number of sites available for binding, causing
the more weakly bound surface species to be affected first.
At this stage, the strongly bound intercalative species is
relatively unaffected. This argument presupposes that the
binding affinity of the intercalative mode is substantially
greater than that of the surface type. Hence, the first
region of the titration reflects surface binding of 3-
aminoacridinium to DNA. As the titration progresses, the
relative number of available binding sites becomes limited
to such an extent that the equilibrium involving the inter¬
calated complex is the one predominantly affected. Now the
plot of log([BHP]/[BH]) vs. log(-q[BHP]) is a reflection
of equilibrium concentrations governed by the stronger mode
which, in turn, results in the slope of the line being 2

81
instead of 1. Note that the difference between the
relative binding affinities of the two modes must be
sufficiently great to allow one to predominate over the
other in the initial and final portions of the titrations.
On the other hand, if they are too disparate, only one
would be seen throughout the experimentally accesible
titration range. Increasing the ionic strength of a
reaction medium causes electrostatic contributions to the
total free energy to become less exothermic for a process in
which oppositely charged moieties coalesce. Based on this,
decreases in the apparent association constants for both
surface and intercalative binding would be anticipated. It
is assumed, at this stage of the argument, that changes in
electrostatic free energies are the predominant affects
accompanying alterations in ionic strength. Extrapolation
of the straight lines of Figures 7 through 10 to the vertical
axis shows that, as the ionic strength is increased, both
Ks and do become less positive in agreement with the above
prediction. However, such extrapolations may be misleading
in the intercalative cases. Recall that, upon altering the
ionic strength, the slopes of the initial, surface binding
regions of the plots are essentially unchanged and that they
form a family of lines. This constancy of slope indicates
that there are no changes in the surface binding reaction
(aside from equilibrium shifts) as the ionic medium is
altered. However, two factors contribute to the decrease in
the intercept values of intercalative plots: (1) lower values

82
of log([BHP]/[BH]) for any given point along the abcissa
as ionic strength is increased, and (2) the decrease in
the slopes of the lines as the ionic strength is increased.
Two explanations for the change in slope may be advanced.
The value of q may be indeed changing from 2 to 1, for the
intercalation reaction, as the ionic strength is increased.
If this is true, then either the overall stoichiometry is
changing such that mq ^ 4 or, the stoichiometry is remaining
constant and m is becoming equal to 4 while q is decreasing
to 1. A more likely explanation can be advanced on the basis
of relative changes of K and K . If K and KT do not change
in the same manner as the ionic strength is varied, then
what we may be observing is an apparent change in q for
intercalative binding when, in reality, the apparent change
is due to an overlapping of surface binding manifestations
with intercalative binding. Assume that the values of m and
q remain unchanged for each type of binding, then, the real
slopes of the two line segments will remain constant and
their point of intersection will be dependent on the ratio
of Ks to K. For example, in one extreme case there is
essentially no intercalative binding, K /K is very large,
b -L
and only the line having slope 1 will be observed. At the
other extreme, Ks/KI would be very small when only inter¬
calative binding occurs and a single line having slope 2
would be seen. If the system is perturbed in a manner which
caused K to decrease more rapidly than KT the point of
intersection of the two lines will reflect the change in

83
K /K . In plots such as Figures 7 through 10, this would
S X
be manifested as a shift toward the right, or to less
negative log([BHP]/[BH]) values. The total number of
sites which are involved in surface and intercalative
binding are, of course, directly related to the magnitudes
of the equilibrium constants. In other words, an increase
in Kg/Kj indicates that a greater percentage of the total
occupied sites are occupied by surface-bound species. The
ratio of K to K is seen to increase for the binding
S _L
reactions between 3-aminoacridinium and DNA in the presence
of alkali metal electrolytes as the ionic strength is raised'.
5
The following data are representative of Ks/K.j. x 10 at
ionic strengths of 0.0025, 0.0050, 0.025, 0.050, and 0.10(M),
respectively. With (CH^^Ni^PO^ as supporting electrolyte:
0.44, 0.56, 1.28, 1.83, and 2.9; whereas with Lil^PO^ as
supporting electrolyte; 0.43, 10.7, 1.2, 5.0, and 10.7,
-5
respectively. Excepting the value of 10.7 x 10 for 0.005
LiH2P04,the ratios show that as the ionic strength is
increased, the total percentage of bound drug which is
surface-bound increases. Moreover, for any given ionic
strength, the ratio is greater when tetramethy1ammonium
dihydrogen phosphate is the electrolyte than when the
lithium salt is present. It is obvious from the decreases
in both Kg and that increasing the dielectric of the
medium does not enhance surface binding relative to inter¬
calative binding. Rather, K does not decrease as rapidly
s
as Kj. The change in Kg/Kj may result from specific ion

84
competition of the electrolyte cations with the drug
cations. As the ionic strength increases, the concen¬
tration of cations which may directly compete for the DNA
phosphate ligands also increases. Moreover, an increase
in the effectiveness of competition by the metals for the
ligands would be expected as the charge density on the
metal increases. Therefore, specific ion interactions may
strongly affect surface binding. It is important to note
that no evidence from the experiments conducted with the
alkali and alkaline earth metals, as counterions, indicated
any alterations in surface binding stoichiometry resulting
from changes in the medium (aside from the magnitude of K ).
O
That is, all initial slopes of figures such as Figure 7 are
equal to 1. Apparent changes in q for the intercalative
region are not likely to be due to intercalation of metal
cations because intercalation of the symmetrically charged,
highly hydrophilic metal cations is improbable. Specific
ion interferences occurring at the surface may, however,
affect the overall intercalation of the drug cations.
Apparent Association Constants for
Surface and Intercalative Binding
Tables 24 through 32 include the apparent association
constants for surface binding, Kg, of 3-aminoacridinium
cation to DNA in the presence of various concentrations of
background electrolytes. For a given titration, values of
Kg were computed at each data point as discussed in the

85
Experimental section. By definition, an equilibrium
constant is independent of relative and absolute con¬
centrations (specifically, activities) of reactants and
products. In fact, this requirement may be used as a
test of the validity of an expression postulated to
represent an equilibrium situation. Inspection of
representative data, included in Tables 18 through 21,
reveals that the values of Ks for points within a specific
titration are reasonably constant throughout the initial
regions. However, in many cases, notably when low con¬
centrations of alkali metal salts are present, the values
calculated from the later points of the titrations are
significantly larger than those calculated from the initial
portions. For example, in the presence of 0.0025M CSH2PO4,
K computed using points from the initial portion of the
s
5-1 .
titration is about 2 x 10 M while toward the end it is
5 -1
about 7 x 10 M . It may be recalled, however, that the
plots of log ([P^/m]- q[BHP]) for this system, indicate
that the value of q is different for the two regions and,
therefore, the binding reactions are different. Thus, when
testing various values of m and q in search of the pair
which results in the most consistent Kg, it is necessary
that Ks be constant only throughout the region over which
it applies. Conversely, upon establishing that specific
values of m and q are optimal, the test of constancy may
be used to define titration regions involving only one
binding mode. The region of the titration from which

86
averaged, reported Kg values were computed, was the same
for all titrations. It has been shown that, as the ionic
strength, I, is decreased, the ratio of bound to free
drug at which intercalative binding becomes significant
increases. The ratio also increases as the charge density
of the countercation increases. For the most dilute
solutions of rubidium, cesium, and tetramethy1ammonium
dihydrogen phosphate, intercalative effects begin to
become important when the logarithms of the ratios approach
+0.2. In solutions having high ionic strengths, or which
contain divalent cations, surface binding predominates even
at the highest ratios of bound to free drug. To allow re¬
liable comparisons of calculated surface binding association
constants between high and low ionic strength solutions, only
data where log ([BHP]/[BH]) was less than +0.2 were used.
The lower limit of the ratio which provides for reliable
values of K was estimated to be -0.8. Below -0.8, data
s
often deviated significantly from values obtained above
this limit. Using the above guidelines, the averaged values
of Kg listed in Tables 24-32 were computed. It is readily
acknowledged that the method of limiting the range of data
is somewhat arbitrary and may be eventually replaced with
more specific, but as yet unknown, guidelines.
The number of individual data points, n, used to
calculate Ks, for each titration, is given in the tables.
Specific points which were used are indicated by double
asterisks (**) in Tables 2 through 12 and 18 through 21.

87
In most instances, n was greater than, or equal to 3.
Standard deviations from the mean, a, are also included.
Deviations were generally within 5% to 10% of the averaged
values of Kg, which shows that the constancy over the
intervals selected is within acceptable limits for these
types of systems. As a rule, the lower the concentration
of background electrolyte, the greater the standard devia¬
tion. Similarly, the lower the charge density of the
cation of the supporting electrolyte, the greater the
standard deviation. For divalent metal electrolytes the
values of Kg were constant over nearly the entire accessible
range. In 0.010N MgfC^CCH^^' Kg varied from 1.32^ x 104 M 1
4 _i
to 1.21 x 10 M over the entire titration for an average
of 1.27 x 104 M-1 + 5%. In 6.3 x 10_4N Mg (0„CCH_,) „, K
— — 2 3 2 s
4-1 4
went from a low of 4.71 x 10 M to a high of 5.09 x 10
-1 4-1
M (average: 4.9() x 10 M + 4%). The essentially in¬
variant apparent equilibrium association constants through¬
out the entire titrations, along with the single straight
line plots of log([BHP]/[BH]) vs. log([Pt/m]- q[BHP]), is
convincing evidence that only the effects of the surface
bound species are observed in these systems and that,
either intercalative binding is not taking place, or it
is being masked. On the other hand, in 0.15M CsH„PO , K
^ ~r S
varies from an average over the whole titration by + 25%
and by + 55% in 0.0025M CsI^PO^. The magnitudes of the
standard deviations over the range -0.8 £ log([BHP]/[BH] <
0.2 are proportional to those given above for the entire

88
range. The higher deviations at lower ionic strengths
may be due to the greater extent of intercalative binding
during the initial parts of the titrations at the low
ionic strengths. Also, there may be a change in the
nature of the surface binding arising from mild denatura-
tion of the DNA. In solutions of low ionic strength, the
configuration of the polyanion may be sufficiently altered
from that of its native state to affect the surface
binding. Addition of even small quantities of DNA to
solutions having very low initial ionicities causes large
relative changes in their ionic strengths. Therefore, if
DNA configuration is dependent on ionic strength and K is
s
dependent on DNA configuration, a substantial variation
may be seen over a narrow titration region.
Agreement between apparent surface binding association
constants obtained from replicate titrations was generally
within + 10% in the presence of all concentrations of
divalent metals and in solutions containing monovalent
cations in excess of about 0.010M. In light of the com¬
plexity of the systems, errors of this magnitude are
considered wholly acceptable. For the more dilute solutions
of monovalent electrolytes, agreement between replicate runs
was considerably poorer, ranging from + 35% to + 3%. It is
believed that the lack of precision, under these conditions,
stems from uncontrolled perturbations of the configuration of
DNA resulting from the low ionic strength. Changes in solu¬
tion pH may be postulated as an aggravating factor contributing

89
to imprecision in the poorly buffered, dilute solutions
in that Schulman and Capomacchia (85) have shown that
pH changes from 5.9 to 7.4 may substantially alter
binding of 3-aminoacridinium to DNA. However, this is
not likely a major contributor because of the relatively
small amount of DNA polyanion added to the solutions.
Even if DNA were a strong pH 7 buffer, the amount added
would not affect the pH as much as the far greater volumes
used for titrations in the presence of dilute alkaline
earths where good precision was obtained. Also, at pH
5.9, the H2PO^ anion is a far better buffer at any given
concentration than is the acetate anion used with the
alkaline earth electrolytes.
The limiting Debye-Hückel law predicts a linear
relationship between the logarithms of equilibrium associa¬
tion constants and the square roots of the total ionic
strengths of the solutions. Figures 11 through 16 are
plots of the log of the apparent surface binding associa¬
tion constant, Ks, vs. the square root of the total ionic
1/2
strength, I . Calculations of Ks and I, from which the
data for these plots are derived, were based on simple
molar concentrations of all species in solution. We have
included this data here so that it may be compared with
literature data and also to allow comparisons with more
refined results presented later. The values of Ks are the
averaged terms reported in Tables 24 through 32 derived
1/2
from molar concentrations. Similarly, the I '
values are

90
averages of the ionic strength at the beginning and at
the end of the regions over which Kg was obtained and
include contributions of all species toward the total
ionic strength (all concentrations in molar terms). It
is immediately apparent, upon inspection of the figures,
that the systems do not obey the limiting Debye-Hiickel
relationship. Most significantly, all plots exhibit
decided curvature over the ionic range investigated. Such
non-linearity is not, however, particularly surprising,
nor is it necessarily unique to the 3-aminoacridinium-DNA
system under consideration since the limiting case is
designed only for dilute, ideal solutions of electrolytes.
Ideal systems are defined as being infinitely dilute solu¬
tions consisting of point charges - ions having no dimen¬
sions and engaging in none other than simple electrostatic
interactions. Real solutions which approximate the ideal
generally consist of simple inorganic ions having high
charge densities with total ionic strengths less than 0.01M.
Even if our systems were ideal in all respects other than
concentration, linearity would be unlikely over the broad
ionic strength range studied. The curvature should, if
simply a result of deviations from electrostatic ideality,
be obviated using corrections included in extended Debye-
Hiickel equations. These will be considered below.
Reconsidering portions of Figures 11 through 16 in
which the total ionic strength is less than 0.01 (I^^ < o.l),
we notice that there is approximate linearity, most easily

91
recognized in the cases of divalent metal electrolytes.
Precision of these data may be estimated from their repro¬
ducibility at a given ionic strength or from the tendency
of data, at various ionic strengths, to form a smooth
1/2
curve when plotted against I . Using either of these
criteria, it appears that the use of alkaline earth
acetates as background electrolytes provides more repro¬
ducible results than alkali phosphates. This suggests
that pH control may not be the most critical factor in
maintaining reproducibility because, for the same ionic
strength, the buffer capacity of the phosphates is markedly
greater than that of the acetates used for the alkaline
earth solutions.
If the conformation of the polyanion is an important
factor in surface binding, then perhaps it is maintained
more effectively by the divalent cations. Their greater
charge and charge density may provide for increased inter¬
action between the polyanion and the metal cations thus
generating a disproportionately higher local ionic field
around the polymer. Hence, though the bulk ionicities of
the two types of solutions are similar, the local electro¬
static potential surrounding the polyanion binding sites
may be much greater in the presence of the alkaline earths,
thereby stabilizing the native DNA. If stabilization in
the divalent metal acetate systems were due to acetate
rather than phosphate, then the data for KC^CCH^ as electro¬
lyte would be expected to be more precise than that for
K^PO^. Figure 12 shows that such is not the case.

92
The imprecision at low ionic strengths is sufficiently
great as to preclude determining straight lines having
slopes agreeing within, say, + 20% for replicate runs.
The large magnitudes of the surface binding constants
for our model compound, to DNA, allow nearly optimal
conditions of reactant concentrations, compared to many
systems reported in the literature. The lack of precision
evidenced at ionic strengths less than 0.010N for our
investigations indicates that conclusions, based on data
obtained at only a single ionic strength below 0.01N, may
be suspect.
Figures 11 through 16 reveal a second manner in which
these systems disobey the limiting Debye-Hückel law in
that, for any given ionic strength, the magnitude of Ks
is a function of the type of background electrolyte. Again,
the actual case is not surprising since we are dealing with
real ions having finite and differing charge densities.
These differences mean that the electrostatic fields around
the ions will be dissimilar and, in turn, the dielectric
properties of the media will be affected. Extrapolation
1/2
of plots of log Ks vs. I ' for dilute solutions of the
several different types of electrolytes would, in theory,
yield lines intersecting at the ordinate. The y-intercept
would correspond to the log of the association constant at
zero ionic strength. To extrapolate the lines, in our case,
is not feasible due to the imprecision of the data. However,
instructive qualitative information may be obrained by

93
comparing values of Kg for solutions of intermediate ionic
strength. For 0.010M solutions of the following cations,
the apparent association constant for the reaction between
3-aminoacridinium cation and DNA decreases in the order:
(CH^)4N+>Rb+>K+>Na+>(Cs+)>Li+. The order is the same in
0.025M solutions (Rb+ was not evaluated at this concentra¬
tion) and also in 0.050M In 0.10M solutions, the
order was (CH^)^N+>>(Cs+)>K+ Rb+>Na+>Li+. In 0.010M
electrolytes Kg ranged from 17 x 10^ M ^ to 9 x 10^ M \
in 0.025M electrolytes, from 11 x 10^ M ^ to 5.4 x 10^ M ^,
in 0.050M electrolytes, from 7.3 x 10^ M ^ to 3.9 x 10^ M ^,
and in 0.10M electrolytes, from 6.1 x 10^ M ^ to 2.9 x 10^ M ^.
With the exception of Cs+, the order of the association
constants parallels the order of the sizes of the electro¬
lyte cations. Ionic radii for the cations are (86): Cs+(1.68
A), (CH3)4N+(>1.3 A), Rb+(1.48 A), K+(1.33 A), Na+(0.97 A),
. + o +
and Li (0.64 A). The exact size of the hydrated (CH^)
ion is not known but it is likely that it is at least as
large as Rb+. It has no available empty orbitals through
which any form of dative binding may occur and thereby,
provides a means of assessing possible effects of available
unfilled metal orbitals on the metal's role in the inter¬
action between the drug and DNA. Qualitatively, the
behavior of the system in the presence of the totally satu¬
rated tetramethylammonium cation is the same as in the
presence of any of the alkali metal cations. Quantitatively,
the apparent surface binding association constants are

94
greatest for the reaction in (CH^)^N+ solutions, in accord
with its large size and low charge density. There is no
evidence that the empty orbitals impart any special
properties to the alkali metals in terms of their effects
on the binding of 3-AA to DNA.
The seemingly strong suppression of the binding process
by cesium ion is not logical in view of the effects of
(CH^)^N+ and Rb+. Because Cs+ is larger than Rb+ it would
be expected to exert a weaker or, at least, similar effect
than rubidium. Despite the fact that Cs+ is in the ex¬
pected order in the 0.10M series, the otherwise anomalous
behavior, in its presence, is likely due to some undefined
variable such as its original purity. The 20%, by weight,
impurity of the solid CsOH used to prepare the phosphate,
as determined by acid-base titration, was considered to be
entirely water of hydration.
At any given ionic strength, based on formal molarities,
the alkaline earths more effectively suppress the binding
of 3-aminoacridinium to DNA than do the alkali metals.
This behavior is not, of course, anticipated on the basis
of the limiting Debye-Huckel law but may be partially
rationalized with extended equations. The order of the
magnitudes of log Kg for the binding of 3-AA to DNA in
the presence of the alkaline earths, decreases in the
order: Ba++>Sr++>Mg++>Ca++ over the entire range of ionic
strength studied (for 0.010N electrolyte, log K is slightly
s
less with Mg++ than with Ca++). In decreasing size of

95
++ ° ++
unhydrated cations (86) the order is: Ba (1.35 A)>Sr
(1.13 A)>Ca++(0.96 A)>Mg++(.65 A). Ion size parameters
for the metal ions are 5, 5, 6, and 8, respectively (87).
The order: Ba++>Sr++>Ca++ is logical, both in terms of
ion size and in terms of ion size parameters. However,
calcium more strongly inhibiting the binding of 3-AA to
DNA than does Mg++ is in opposition to charge density and
ion size arguments. This reversed order indicates that
the effect of the metal ions encompasses more than simple
bulk solution electrostatic factors which can be correlated
with ion sizes and charge densities. The most apparent
additional effect is specific ion competition in which
the overall energetics of the competition for phosphate
sites must be evaluated.
Figures 17 through 20 are plots of the apparent
intercalative association constants, , against the square
1/2
root of the ionic strength, I ' , for the alkali metals.
They are similar to Figures 11 through 16 as. they show the
1/2
relationship between log Ks and I ' based on conventionally
employed, uncorrected molar concentrations. In the presence
of the alkali metals, the relationships between the apparent
intercalative association constants, K^, and the ionic
strength were essentially the same as for K and I discussed
s
above. That is, limiting Debye-Hiickel behavior is not
obtained over the entire concentration region. Curvature
is especially apparent at higher concentrations, and the
nature of the background cation affects the degree of

96
binding. Except in the presence of 0.10M electrolytes,
the order of the magnitudes of the intercalative binding
constants for 3-aminoacridinium and DNA in the presence
of the monovalent cations is the same as for surface
binding: (CH^)^N+>Rb+>K+>Na+>(Cs+)>Li+. Again, the
sequence can be rationalized on the basis of charge
density, excepting cesium. In 0.10M solutions, the
order is (CH^)^N+>>Rb+ ~K+>(Cs+)>Na+>Li+. This series
differs from the one for Kg only in the position of Cs+
which, again, is probably an experimental artifact. We
were unable to observe intercalative binding of 3-amino¬
acridinium to DNA in the presence of alkaline earth cations
at any concentrations. Data were generated and plotted
for these systems similar to Figures 17 through 20 but
are not included here. The order of the "K " values so
determined were entirely random with respect to either
unhydrated ionic radii or ion size parameters. Possible
reasons why the intercalative binding was not seen in
these media were presented previously.
To summarize, the relationship between apparent
association constants for both surface and intercalative
binding of 3-aminoacridinium to DNA over the concentration
range of 0.10N to 0.00063M alkali metal background electro¬
lyte and 0.025N to 0.00063N alkaline earth electrolytes do
not obey the Debye-Hiickel limiting law. In all cases, the
plots are non linear and the magnitudes of the association
constants are dependent on the nature of the cations. With

97
the exception of Cs+, the relative effects of the alkali
metals parallel their respective charge densities. The
same is true for surface binding in the presence of di¬
valent cations, except that Ca++ and Mg++ are reversed.
Plausible comparisons between charge density and relative
values of Kj in the presence of divalent electrolytes
cannot be made because of the inability to evaluate inter-
calative binding in solutions containing the divalent
cations. The implications by other investigators that
there is a straight-forward relationship between ionic
composition of solutions in terms of molarities and binding
constants must be considered with reservations.
Association Constants Based on Activities —
Extended Debye-Huckel Considerations
The procedures outlined above for the determination of
binding constants of small molecule to DNA are employed
by the majority of investigators in this field. That is,
calculations are made with the assumption, either explicit
or implied, that molar (or molal) concentrations of species
are equal to their activities. Only when activity coeffi¬
cients are equal to 1 is this assumption strictly valid —
a condition met only by dilute solutions of "well-behaved"
systems such as fully ionized, monovalent, inorganic salts.
Ideally, evaluation of the relationship between ionic
strength and association constants should be based on
solutions having total ionic strengths approaching zero, and

98
including low concentrations of reactants and products.
Unfortunately, the conformational instability of DNA in
low ionic strength media (<_ 0.0050 (M)) prevents such
investigations. Furthermore, limits of instrumental
detectability and the nature of the equilibrium demand
moderately high concentrations of the biopolymer. But,
while concentrations of background electrolyte and
reactants must be higher than desired, a more accurate
reflection of the relationship between ionic strength
and the association of small molecules to DNA may be
made using species' activities as opposed to molar
concentrations.
Activities of each electrolyte at the various molar
concentrations were estimated by determining the mean
activity coefficients of the constituent ions from
equation (2-18). Ion size parameters, d, were obtained
from the literature (87) and are included in Table 33.
The activities of the electrolytes were used, in turn,
in equation (2-18) as the values of I when estimating the
activity coefficients of the species involved in the binding
processes. Ion size parameters for BHP, BH, and DNA phos¬
phate were not available and had to be estimated. DNA
_ g
phosphates were assigned a d value of 4 x 10 , being
approximately the same as the value for free I^PO. . Free
and bound 3-aminoacridinium cation were given d values of
“8 — s
3 x 10 and 4 x 10 , respectively. The assignment of ion
size parameters for these species is, admittedly, rather

99
arbitrary as there are few guidelines by which to estimate
them. However, it is the effects of varying the back¬
ground electrolyte which is of principal concern here.
So, as long as the values of d for the species involved
in the equilibrium are held constant, relative effects
of changing the medium can be estimated though absolute
constants may be incorrect. After calculating activity
coefficients for reactants and products, the correction
factor, £ , was computed from equation (2-17) where =
^Dun/(^u^ )• Table 33 lists the correction factors for
rSHir rJn p
the surface binding along with logarithms of the corrected
association constants, log Ks, in the presence of varying
concentrations of the electrolytes. For any given initial
molarity of alkali metal electrolyte, only minor variations
in E are observed for the differing types of cations.
These small differences are reflections of the minor
effect d has on the correction factor. For instance,
values for 0.10M LiH2PO^ and 0.10M CsH-^PO^ are 4.28 and
4.20, respectively, where d is 8 x 10 8 and 2.5 x 10 8 for
the respective cations. Below about 0.010M electrolyte
the correction factor is independent of the character of
the monovalent cation. For ionic strengths greater than
I
about 0.01(M), the order of increasing log Ks is the same
as for log Ks, namely: (CH3)4N+>Rb+>K+>Na+>(Cs+)>Li+.
(Cesium's position is not the same, but data pertaining
to it are probably not reliable, vide supra.) The parallel
order for the corrected values with those based on molar

100
concentrations is expected in light of the relatively
minor effect that the ion size parameter has on £ .
Below I = 0.010(M), the data is too scattered to make
definitive statements. Too, at low ionic strengths, the
assumption that DNA phosphates can be treated as if they
were species independent of one another becomes especially
approximate. Free species may be diluted to any degree
whereas, below a given value, further addition of bulk
solvent will have little effect on altering the ionic
environment of a site on the polymeric surface since its
environment is predominantly determined by near-neighbor
charges along the polymer and not by charged species in
the bulk solvent. Hence, the greater the degree of dilution,
the greater the disparity between the calculated ionicity
and the true ionic strength local to the binding site.
I
There is a substantial decrease in the range of Kg
compared to K values for the range of concentration (activity)
s
studied. At higher ionic strengths the value of B, is maximal
s
f
while Kg is minimal while at low ionic strength .£ is low
with log K being at its maximum. The product of the two,
s
I
therefore, tends to decrease the range of Kg relative to Kg.
1 1/2
As a result, plots of log K vs. I are more linear as
s
evidenced in Figures 21 and 22. Having corrected for
variations in activities of species in solution, and also
having presumably eliminated the effects of differing elec-
I
trolyte ion sizes, theory predicts that plots of log Kg vs.
1/2
I would be straight lines which are superimposable.

101
Obviously, while data for I > 0.010(M) for most of the
systems do form reasonably straight lines, they do not
superimpose. Moreover, there is still some curvature in
the plots which is most pronounced for lithium, becoming
less apparent as one goes from Na+ to (CH^).N . These
two features likely result from specific-ion competition
by metals for the phosphate sites on the DNA.
Figures 23 and 24 are plots of the corrected surface
• 1/2
binding equilibrium association constants, Ks vs. I
for reactions in the presence of alkaline earth electro¬
lytes. It can be seen from the figures and from data in
I
Table 33 that the order of the magnitudes of Kg in the
presence of the alkaline earths is dependent on background
electrolyte concentration. This is in contrast to the
I
trends for K for alkali metal solutions and also, for
s
the uncorrected values of both alkali and alkaline earth
electrolytes. The orderings are not random but are due
to the unequal slopes of the distinctly straight lines
for each type of cation. The lack of superposition of
the lines shows that specific ion competition probably is
operative.
The above treatment has shown that application of
extended Debye-Hiickel equations to surface binding data is
sufficient to rationalize the nonlinear relationship between
Kg and I. The remaining deviations are probably due to
specific ion effects which will be briefly considered below.
Similar attempts to explain the intercalative data are not

102
so fruitful. The correction factor for the intercalative
binding expression, £ , is essentially a constant, equal
to 1, for all ionic strengths investigated (the highest
value, for I = 0.15(M), is 1.05). Of course, the wide
1/2
range of Kj values and the nonlinearity of log vs. I
are unaltered by application of the correction for activi¬
ties. The implications here are uncertain at this time,
but it appears that, whatever it is that predominantly
affects the intercalative binding, it is not explainable
in terms of simple concentrations or ionicities. One
possibility is that the intercalative mode is highly
susceptible to minor changes in the pitch of the DNA
which is, in turn, a function of ionic strength.
Evaluation of a Simple Competitive Binding Model
A thorough investigation of specific ion competition
between electrolyte metals and 3-aminoacridinium cations,
for DNA phosphate sites, would include complex equilibria
involving mixed complex systems which are beyond the scope
of this work. Such a study would include the selection of
probable stoichiometries of metal-DNA, or metal-drug-DNA
mixed complexes using procedures similar to those for the
drug-DNA systems. These models would then be evaluated to
determine the one most probable, on the basis of a "best-fit
The most elementary of these will be briefly considered here
This system would be one in which there are no mixed
complexes and in which the binding stoichiometries of the

103
metal, and 3-AA, to DNA are the same. It has been shown
that 3 DNA phosphates, acting as a single tridentate
ligand, bind to each 3-AA, which we shall assume is the
same for the metal binding. Considering the 3 phosphates
as constituting each site, P, in equations (1-4) and (1-5),
we may apply equation (1-7) to our system. Then, the
relationship between the observed association constant,
I
K , for the 3-aminoacridinium-DNA reaction, the actual
s
II
association constant for drug-DNA binding, Ks, and the
association constant for the metal-DNA interaction, K^,
as developed by Passero and Gabbay (64) would be
k' = k"/(1 + K„ M) (3-5)
s s M
where [M] is the molar concentration of the electrolyte
I
metal cation. Substituting values of Kg and their cor¬
responding [M] values into equation (3-5) allows for the
If
simultaneous solution of K and Kw. Table 34 lists calculated
S M
II
values of log Kg and for the competitive reactions of
alkali and alkaline earth cations and 3-AA for DNA phosphate
sites. The concentration ranges of electrolytes listed in
Table 34 were selected on the basis of the straight line
I
regions of plots of K against [M]. For the simple systems
assumed here, there would be a linear relationship between
the two parameters because of the first order concentration
terms for all species in both equilibrium expressions. Data
for lithium as an electrolyte were not included in Table 34
since there was no region of linearity throughout the whole

104
range of metal ion concentration. This may indicate
that the interaction between DNA and Li+ is substantially
different than between DNA and the other metals. Caution
must be advised in drawing conclusions about trends ob¬
served in the values of the metal association constants
for each series. These values are obtained as the result
I
of small differences between large numbers (K ) which, in
s
themselves, have large error limits. It is evident,
though, that values for the divalent metals are approxi¬
mately an order of magnitude greater than the corresponding
values for the alkali metals. The logarithm of the equilib¬
rium association constant for the surface binding of 3-
aminoacridium cation to DNA, having been corrected for
ionic strength, ion sizes of species in solution, and
specific ion interactions by alkali metals appears to be
5.41 + 0.10. While not as large as for the alkali metal
II II
value of K , the value of 4.71 + 0.08 for K in the presence
s s
of divalent metal competitors appears reasonable. Having
ostensibly corrected for ionic charges, along with the
It
above-mentioned parameters, the values of log for the
alkali and alkaline earth series should be equal. That
they are substantially different proves that the model
needs further refinement. It is logical, though, that
the binding stoichiometries of divalent cations and mono¬
valent cations with DNA phosphates would be different.

105
Thermodynamics of the Binding of 3-aminoacridiniuin
and 7-aminoquinolinium to DNA
The standard free energies, AG°, enthalpies, AH°,
and entropies, AS°, for the two binding reactions of 3-
aminoacridinium and 7-aminoquinolinium to DNA were
evaluated from the temperature dependencies of their
equilibrium association constants. By comparing the
relative magnitudes of AH° and AS° for 3-aminoacridinium
with those of its benolog, 7-aminoquinolinium, specific
effects resulting from variations in the aromatic portion
of the molecule may be evaluated. Differences in chemical
properties, especially acidities, must be borne in mind
when making such comparisons.
The general spectral characteristics of the 3-amino-
acridinium-DNA system, along with the regions in the titra¬
tions corresponding to surface and intercalative binding
and the evaluation of their constants have all been
previously discussed for reactions at 25.0°C. Except
for small changes in the magnitudes of the equilibrium
constants, the systems at 15.0°C and 35.0°C behave in
essentially the same manner. These topics will be briefly
considered for the reactions between 7-aminoquinolinium and
DNA at 25.0°C.
Figure 25 contains representative spectra of the
-5
absorptiometric titration of a 2 x 10 M solution of 7-
• • • —2
aminoquinolinium with the potassium salt of DNA (10 M in
DNA phosphate). The total absorbances of the 7-AMQ solutions

106
are not as great as for the 3-AA solutions due to the
former's lower molar absorptivities (free, 794() M ^cm ^
at 392 nm; bound, 533.2 M ^cm ^ at 407 nm) . Drug concen¬
trations could have been used which would have allowed
absorbances greater than 0.4. However, it was decided
to use the same formal concentrations of 7-aminoquinoline
as 3-aminoacridine. Then, any differences observed in
the titration behavior of the two could not be due to
differences in their concentrations. Also, the lower
binding affinity of 7-aminoquinolinium for DNA makes it
desirable to keep its concentration at a minimum to allow
maximum ratios of DNA to drug without introducing compli¬
cating concentration problems.
The visible spectra of both the free and bound 7-
aminoquinolinium cation consist of only one envelope,
having no fine structure, unlike the spectra of the 3-
aminoacridinium species which have well differentiated
^La and bands. Presumably, substitution of the amino
group in the 6 position of the molecule results in the
transition having increased probability over the . Thus,
cl
the observed spectra consist predominantly of the ' ^A
transition. The band may be buried beneath the
though no attempts were made to ascertain its position.
The lack of distinct and bands and the absence of
a b
fine structure in the spectra of either the free or bound
species preclude discussions concerning possible molecular
orientations upon binding. There is no evidence to indicate,

107
however, that 7-aminoquinolinium does not interact with
DNA in essentially the same manner as does 3-aminoacridinium.
Similar reaction stoichiometries and values of m and q in
equation (1-21) are in agreement with this presumption.
Significant shifts in the baseline of the spectra occur
with the addition of moderate to large amounts of DNA.
Figure 25B is a reproduction of a titration spectrum in
which it may be seen that the absorbance at wavelengths
greater than 475 nm increases by approximately 0.03 absor-
-2
bance units upon addition of 1400 yl of 10 M DNA phosphate
to 8.04 ml of solution (resultant P/D = 105). Magnitudes
of baseline shifts for 7-aminoquinolinium titrations are
particularly significant in view of the relatively small
total changes in absorbance, at the analytical wavelength,
throughout the titrations. Reverse titrations in which
stock DNA solution was used to titrate pH 5.90 buffer in a
4.00 cm cell showed that an increase in the baseline absor¬
bance of about 0.03 units may be anticipated from light
scattering arising from the presence of relatively high
concentrations of DNA polymer. Correcting the absorbances
at 500 nm of each individual spectra to zero results in
absorptiometric titration spectra similar to the one shown
in Figure 25A. Here it becomes evident that what appeared
to be an isobestic point at « 420 nm was, in fact, an arti¬
fact and, thus, the system may involve more than two species
which are in equilibrium. It is also evident that, even
after addition of 1400 yl of DNA solution to the 8.04 ml of

108
drug, (P/D = 105) the titration was not complete. Adding
greater amounts of DNA is counterproductive because of
increasingly severe light scattering and the dilution of
the sample. Baseline absorbance at the analytical wave¬
length was determined at the point of intersection of a
line drawn between the absorbances at 475 nm and 325 nm, and
a perpendicular at the analytical wavelength. The analytical
wavelength was that of the band maximum for the total solu¬
tion .
The band maximum for free, protonated 7-aminoquinolinium
is at 392 nm. As DNA is added to the drug, there is both a
decrease in the intensity of the band and a red shift of its
maximum. The hypochromism and bathochromism are both quali¬
tatively similar to those seen in the titrations of 3-amino-
acridinium. However, the 7-AMQ red shift indicates a
significantly greater energy difference between the free
and bound 7-aminoquinolinium compared to the difference
between free and bound 3-aminoacridinium. In the presence
of a 100-fold excess of DNA phosphate to drug, the band
maximum for bound 7-aminoquinolinium appears at 407 nm cor¬
responding to a difference in the transition energies between
the ground and excited states of the free vs. the bound
complexes of 940 cm ^ (2.70 Kcal/mole 7-AMQ). Recall that
the transition energy difference for free and bound 3-amino¬
acridinium amounted to only about 1.3 Kcal/mole 3-AA.
There were no titrations of 7-aminoquinolinium in which
an abrupt shift in the spectra, similar to that seen in the

109
band of the 3-aminoacridinium system, was observed.
The rbsence of a dramatic shift at DNA phosphate levels
in excess of those causing shifts in the 3-aminoacridinium
systems shows that shifts in the latter instances were due
to something other than simple changes in the bulk solvent's
composition. It is not known if greater concentrations of
DNA would have resulted in dramatic shifts in the 7-amino-
quinolinium spectra.
Figures 26 and 27 are plots of log([AHP]/[AH]) vs.
log([Pfc/3]-[AHP]) and log([P /2]-2[AHP]), respectively for
the titrations of 7-aminoquinolinium with DNA at 15.0°C,
25.0°C, and 35.0°C. In all instances, two distinct regions
of linearity are seen, indicative of two separate modes of
binding. In Figure 26 the averaged slopes of the initial
portions of the curves, corresponding to the surface binding
process, is 1.03 + 0.13. This value is nearly the same as
for 3-aminoacridinium and shows that equation (2-15) also
applies to 7-aminoquinolinium surface binding. The averaged
values for the same region of Figure 27 is 1.02 + 0.12. The
value for q for the intercalative binding, obtained from the
slopes of the later regions of the lines in Figure 27, is
1.64 + 0.09 (from Figure 26, q = 1.8 + 0.4). These are
somewhat ambiguous in terms of the actual integral value of
q but there is little doubt that it is greater than 1 and
is presumed to be less than 3, based on the lack of constancy
of the computed values of when q = 3. Because of the
relatively weak binding of 7-aminoquinolinium to DNA, it was

110
not possible to titrate the system in a region in which
the surface binding did not overlap intercalative binding.
As a result, the later points in Figures 26 and 27 probably
include some contributions from surface binding which would
cause the slopes to be less than 2. On the bases of best-
fit analyses and the similarity of the binding with that of
3-aminoacridinium, values of m = q - 2 were selected for
the intercalative mode of 7-aminoquinolinium binding to
DNA.
Table 35 includes the calculated apparent association
constants for the surface and intercalative binding of 3-
aminoacridinium and 7-aminoquinolinium to DNA at 15.0°C,
25.0°C, and 35.0°C. Constants were computed using the
estimated activities of the species (equations (1-30) and
_ g
(2-17)). An ion size parameter equal to 3.5 x 10 was
used for the free drug cations and bound complexes. For
any given temperature, the apparent surface binding associa¬
tion constant of 7-aminoquinolinium is approximately an order
of magnitude lower than for 3-aminoacridinium while the
constants for the intercalative process are about two orders
of magnitude lower for the former. Magnitudes of Kg and
for both compounds decrease with increasing temperature in¬
dicating a negative free energy for the binding process.
* n
Figures 28 and 29 are plots of log Ks vs. 1/T ( K) for 3-
aminoacridinium and 7-aminoquinolinium, respectively, and
1 o
Figures 30 and 31 are similar plots of log Kg vs. 1/T ( K).
With the possible exception of Figure 31, the relationship

Ill
between log K and 1/T is linear as predicted by the
van't Hoff relationship (equation (1-31)). Values of
the standard enthalpy for the reactions, included in Table
36, were calculated from the slopes of the lines of Figures
28 through 31. Standard entropies of binding were computed
from values of AH° and AG° at 298°K. The standard state
upon which the thermodynamic parameters are based is
related to activities of the species estimated from their
molar concentrations using equation (1-30). While the
total free energy of binding of 3-aminoacridinium is more
negative than for 7-aminoquinolinium (AG° = -7.35 Kcal/mole
3-AA vs. -6.07 Kcal/mole 7-AMQ) the standard enthalpy change
or heat of reaction involving 3-aminoacridinium is less than
that involving 7-aminoquinolinium (AH° = -4.90 Kcal/mole 3-AA
vs. -5.13 Kcal/mole 7-AMQ). Enthalpy is predominantly a
measure of the overall energy changes resulting from the
disruption of bonds associated with reactants and the
formation of new bonds among product species. Considered
in this vein, it is evident that the greater binding affinity
of 3-aminoacridinium over 7-aminoquinolinium for DNA cannot
be attributed to electrostatic bonding energies because,
based solely on AH°, the binding of 7-AMQ to DNA would be
greater than that of 3-AA. Two types of electrostatic inter¬
actions are probably most significant here: the ion-dipole
forces between drug cations and water molecules of the
solvent, and the ion-ion interactions of the charged ring
nitrogens with the DNA phosphate anions. Ion-dipole forces

112
between solvent molecules and the phosphate anions would
be less important than the above two. The two drug cations
are similar in terms of the structural characteristics which
would affect the attractive forces between them and water
molecules. Both have hydrogen-bond acceptor amino groups
and protonated hydrogen-bond ring nitrogens. Thus, to a
first approximation, the total energy necessary to disrupt
the ion-dipole interactions between water and 3-amino-
acridinium cations would be expected to be about the same
as that for water and 7-aminoquinolinium. More important,
though, is a consideration of the relative acidities and
basicities of the groups. The extent of hydrogen bonding
between the solvent and a hydrogen-bond acceptor is related
to the acceptor's basicity while the extent between the
solvent and a hydrogen-bond donor can be correlated with
the donor's acidity. The protonated heterocyclic ring
nitrogens are hydrogen-bond donors. Therefore, 7-amino-
quinoline (pK = 6.65), being a stronger acid than 3-amino-
acridine (pK = 8.04), will form stronger hydrogen bonds
with water making the AH° for its binding reaction less
positive than that of 3-AA. Likewise, the basicity of the
substituted amino group of 7-aminoquinoline is somewhat
greater (pK = 0) than the amino group on 3-aminoacridinium
(pKa = -1.5), meaning the hydrogen-bond acceptor of the former
would interact more strongly with water than the latter, thus,
generating a less negative AH° for reactions involving 7-AMQ
as a reactant. The above factors tend to make the change in

113
enthalpy for the reaction of 7-aminoquinolinium with DNA
less negative than that of 3-aminoacridinium with DNA.
It is clear that simple arguments based solely on ion-
dipole energies of the reactant cations are not sufficient,
in themselves, to explain the more negative AH° for the
3-AA system.
The amount of energy evolved upon formation of the
electrostatic ion-ion bonds between the cationic drug
molecules and the anionic phosphates of the DNA may explain
the slightly more negative AH° for 7-aminoquinolinium
binding over that for 3-aminoacridinium. The cation having
the greater charge density at the heterocyclic nitrogen
would bind more strongly. 7-aminoquinolinium, being the
stronger acid, would generate the stronger electrostatic
interaction of the two, in agreement with the experimentally
observed values. The effect of the third ring of 3-amino-
acridine as a steric barrier to hydrogen bonding of the ring
nitrogen would serve to make AH° for the reaction involving
3-AA more negative than for 7-AMQ, but this effect would
likely be small.
Since the binding sites are presumed to be identical
for the two cations, the energies involved in disrupting
bonds associated with the solvent and the DNA binding sites
would be the same. Likewise, enthalpic contributions arising
from the formation of the solvent cages surrounding the bound
species would be expected to be similar if the binding geo¬
metries of the two are alike.

114
Intercalation of 7-aminoquinolinium and 3-amino-
acridinium may occur by their "slipping" between DNA
base pairs. Such a mechanism would allow for interactions
between the aromatic base pairs and the ring systems of
the cations and thus the 3-AA, having a larger aromatic
surface area, might be expected to interact more strongly.
Changes in the standard enthalpies for the intercala-
tive binding processes of the two cations to DNA are more
negative than for the surface binding, but are qualitatively
similar in that AH° for 7-aminoquinolinium is slightly more
negative (-15.1 Kcal/mole) than for 3-aminoacridinium (-13.7
Kcal/mole). The absolute magnitude of AH° for 7-AMQ is 1.06
times greater than AH° for 3-AA while the absolute magnitude
s
of AH°for 7-AMQ is 1.10 times greater than AH° for 3-AA.
These similar ratios show that in terms of the enthalpic
contributions to the free energy of binding, intercalation
of 3-aminoacridinium and 7-aminoquinolinium are the same.
That is, the larger, more lipophilic aromatic portion of
3-AA does not impart any particular advantage to the com¬
pound in terms of the enthalpy of the intercalation process.
The entropy of surface binding of 3-aminoacridinium is
considerably more positive than for the similar process with
7-AMQ. For 3-AA, TAS° at 298°K equals 2.45 Kcal/mole and
for 7-AMQ, TAS° at 298°K is 0.94 Kcal/mole. In fact, it is
this relatively large difference in entropy which obviates
the enthalpic advantage 7-AMQ has over 3-AA, resulting in
the more negative free energy of binding of 3-aminoacridinium.

115
Entropy change is a measure of the extent to which the
randomness of a system is altered during the course of a
reaction, being positive when the final state is less
ordered than the initial state. In terms of the stoichio¬
metries of reactants and products, AS° for the two would
be expected to be the same. Obviously, the solvent and
the solute's effect on its structure are of great importance
in the present instance. The role of solvent structure
may be understood by envisaging the bulk solvent as having
a relatively random, fluid structure. Dispersed throughout
the solvent are the organic cations, 3-AA and 7-AMQ, which
both disrupt the bulk solvent structure by their presence.
But the tricyclic 3-aminoacridinium, being more lipophilic
and larger in size, will demand a greater extent of solvent
structuring to accommodate it. We have already considered
the effects that the polar regions of the cations would have
on the water via hydrogen bonding. Therefore, since the
DNA sites of the two species are presumed to be the same,
the binding of 3-aminoacridinium will result in a more
positive entropy change as the cation is "removed" from
the aqueous solvent. This would be true even if surface
binding effected only partial removal of the cationic
species from the water.
Large error limits for the values of AS° for the inter-
calative binding of the two cations to DNA mandate caution
in comparing them, though it appears that AS° for the binding
of 7-aminoacridinium is considerably less positive than for

116
3-aminoacridinium. This relationship is in parallel with
the surface binding AS° values and may be due to the same
factors as discussed above. The entropy changes given
in Table 36 are per mole of cation but, since the surface
and intercalative modes involve different ratios of DNA
phosphate to cation, comparisons of the entropies of the
two processes can be misleading. However, AS° terms for
surface binding are significantly positive whereas for
the intercalative mode they are, within experimental error,
nearly zero or less than zero. Intercalative binding entropy
changes include a positive contribution arising from the
release of water molecules as the cations are removed from
the bulk solvent. These would likely be about the same as
the AS° terms for surface binding meaning that the entropy
associated with the insertion of a 3-AA cation already on
the surface of the polymer would be approximately -10.4
(-2.2 - 8.2) e.u. or -3.10^ Kcal/mole at 298°K. For 7-AMQ
the value would be about -17 e.u. (-14 - 3) or -5.0^7 Kcal/
mole at 298°K. These negentropic terms indicate that it is the
highly negative enthalpies of the intercalative reactions
which result in the overall negative free energy of binding.
Surface bound cations will apparently spontaneously bind
intercalatively due to the small enthalpic advantage afforded
the latter process. It appears that, while the enthalpy
arising from intercalation is decisive, it is only slightly
more negative for 3-aminoacridinium than for 7-aminoquinolinium.

117
Earlier, we discussed the binding reactions in terms
of ligand-metal models. Thus, it is instructive to continue
the analogy to specifically include the chelate effect.
The well known chelate effect (88) refers to the more
negative free energy of formation of polycoordinate metal
complexes formed from polydentate chelates over similar
complexes formed from monodentate ligands of chelates having
2+
fewer coordination sites. For example, log 84 for Cu(NH2)^
is 11.9 while log 82 for Cu(H2NCH2CH2NH2)is 20.0 (89) in
which the former is formed from 4 monodentate ammonia ligands
and the latter from 2 bidentate ethylenediamine chelates.
The chelate effect for non-transition metals is almost
entirely an entropy effect (90). For transition metals,
greater crystal field stabilization of chelates over mono¬
dentate complexes introduces an enthalpy contribution (91).
The more positive entropy for chelate binding means, in
terms of probability, that it is easier to form an aggregate
from a few molecules than from many molecules. The three
DNA phosphates involved in surface binding could conceivably
act as three monodentate ligands, a monodentate ligand and a
bidentate chelate, or as a single tridentate chelate. The
case in which the phosphates act as a single chelate gives
the most consistent calculated values of K and is also the
s
one which would be expected based on the chelate effect. On
the other hand, the 4 phosphates involved in intercalative
binding could react as 4 monodentates; 1 monodentate and 1
tridentate; 2 monodentates and 1 bidentate; 2 bidentates; or

118
as 1 tetradentate ligand. Clearly, the chelate effect
predicts that the tetradentate ligand would generate the
4-
most positive entropy. However, the use of a (PO^)^ term
in the equilibrium association expression yields values of
Kj which are not nearly as consistent throughout the inter-
2- .
calative titration region as those when ^(PO^^ is
employed. Thus, an antichelate effect is seen for the
intercalative process in which the greater AH° for the 2-
bidentate-ligand mechanism appears to be sufficient to
compensate for any positive entropy gain which might arise
from a reaction involving a single tetradentate chelate.
When discussing the forces which account for the
behavior of polymers and biologically important systems,
such as lipid membranes in aqueous media, the "hydrophobic
bond" is often invoked. The avoidance here of any reference
to this term as an aid to explaining the energetics of the
binding reactions has been intentional. In most cases,
"hydrophobic forces" are merely manifestations of well
established fundamental forces arising from energies of
adhesion and cohesion and from entropic factors. It perhaps
could be argued that the "driving force" for the intercalative
binding of the compounds to DNA is the formation of hydrophobic
bonds between the aromatic rings of the cations and the
aromatic portions of the DNA bases. But as Hildebrand (92)
has pointed out, it is rare where the energy involved in the
interaction of two organic species in an aqueous system is
any greater than it is in a corresponding non-aqueous
environment. For example, the alkyl chains in soap micelles

119
are not bonded together by a phobia for water — they are
just as adhesive in the absence of water. In other words,
there is not necessarily any phobic response of the
species toward water, but rather the water molecules
would "prefer" to be hydrogen bonded together rather than
separate to admit the immiscible molecules.

APPENDIX I
FIGURES

Figure 1. The structures of selected substituted
acridines and related compounds
(Structures I through XVI)

I.
3-AMIN0ACRIDINIUM CATION (3-AA)
II.3,6- DIAMINO - IO-METHYLACRIDINIUM
CATION (ACRIFLAVIN)
III.3,6-DI AMINOACRI DI NIUM CATION
(PROFLAVINE)
IV.BENZ [c]ACRIDINIUM CATION

123
V. D1BENZ [qc] ACRIDINIUM CATION
NHCH(CH3)(CH2)3NEt2
OCH,
VI. QUINACRINE CATION (ATEBRIN.MEPACRIN)
eNHCH2CH(CH3)(CH2)2 N E»2
VII.CHLOROQUIN CATION
NH(Y)NR,R2
VIII.GENERAL STRUCTURE OF SOME
ANTIMALARIALS

124
NHCH2CH(OH)CH2N ei2
0CH3
X.ACRANIL CATION
XI.NEOMONA
XII.9-AMINO- 1,2,3,4-TETR AH YDROACRID INI UM
CATION
Nh2
IN CATION

12 5
NHR
a) R= (CH2)2NH(CH3)2
b) R « (CH2)2N(CH3)2CH2N(CH3)3
XIII. REPORTER MOLECULES FROM REFERENCE
(3 I )
XIV. 3,6-DIMETHYLAMINOACRIDINIUM
CATION (ACRIDINE ORANGE)

126
XV. ETHIDIUM BROMIDE CATION
XVI. 7-AMINOQUINOLINIUM CATION (7-AMQ)

Figure 2. General form of Scatchard plots: R/C vs. R for one and two classes of
binding sites
A One class of binding site (K,n).
B Two classes of binding sites
Kji , nIX) where Kj >>Ki;i
(K j , n j ;

crio
(a)
(b)
128

Figure 3. Absorptiometric titration of 3-aminoacridinium
cation with DNA, cesium salt, pH 5.90, 25.0°C.
Background electrolyte: C H2PO , initial
concentration, 0.15M
a P/D = — ([P] = 0)
b P/D = 2.89
c P/D =5.97
d P/D =9.65
e P/D =17.0
f P/D = 26.7
g P/D = 65.3
P/D is the ratio of total DNA phosphate to total
3-aminoacridinium cation, D = 1.35 x 10“^ moles.

130

Figure 4. Absorptiometric titration of 3-aminoacridinium
cation with DNA, cesium salt, pH 5.90, 25.0°C.
Background electrolyte: C H„PO
concentration
, 0.0025M
a
P/D =
— ([P] = 0)
b
P/D =
0.652
c
P/D =
1.63
d
P/D =
2.94
e
P/D =
4.57
f
P/D =
5.87
g
P/D =
13.7
P/D is the ratio of total DNA phosphate to total
3-aminoacridinium cation, D = 1.33 x 10-^ moles.

132

Figure 5. Absorptiometric titration of 3-aminoacridinium
cation with DNA, magnesium salt, pH 5.90, 25.0°C.
Background electrolyte: Mg (C^CCH-^) 2 / initial
concentration, 0.010N
a P/D = — ( [P] = 0)
b P/D =3.14
c P/D = 9.57
d P/D =18.9
e P/D =37.7
f P/D =69.1
g P/D = 104
P/D is the ratio of total DNA phosphate to total
3-aminoacridinium cation, D = 1.31 x 10“^ moles.

.8-
.6-
UJ
o
z
<
m
or
o
(/)
CD
<
.2-
.0.
300
400 500
WAVELENGTH (i\m)

Figure 6. Absorptiometric titration of 3-aminoacridinium
cation with DNA, magnesium salt, pH 5.90, 25.0°C.
Background electrolyte: Mg(0?CCH-J ~, initial
concentration, 0.00063N
a P/D = — ( [P] = 0)
b P/D =1.63
c P/D = 3.59
d P/D = 6.49
e P/D = 13.0
f P/D = 26.0
g P/D =64.9
P/D is the ratio of total DNA phosphate to total
3-aminoacridinium cation, D = 1.27 x 10-^ moles.

136

Figure 7
Log([BHP]/[BH]) vs. log([Pt/3l“[BHP])
CsH2P04 and 6.3 x 10“^N Mg(C>2CCH3)2/
electrolytes, pH 5.90, 25.0°C
A 0.0025M CsH2P04
6.3 x 10-4N Mg(02CCH3)2
with 0.0025
as supportin
O

133

Figure 8. Log([BHP]/[BH]) vs. log([Pt/3]~[BHP]) with 0.15M
CsH2P04, 0.010M CsH2P04, and 0.010N Mg(02CCH3)2
as supporting electrolytes, pH 5.90, 25.0°C
A 0.15M C H„PO.
S 2 4
• 0.01M C HoP0.
s 2 4
O 0.010N Mg(02CCH3)

UOG ( [bhpJ/IbhI)
140

Log([BHP]/[BH]) vs. log([Pt/2l“2[BHP]) with 0.00
CsH2P04 and 6.3 x 10-4N Mg(02CCH3)2 as supportin
electrolytes, pH 5.90, 25.0°C
A 0.0025M C H„PO.
S 2 4
0 6.3 x 10"4N Mg(02CCH3)2
Figure 9.

([dH| 2-g/j] ) 901
+
ro
LOG ([BHP]/[BH])
142

Figure 10. Log([BHP]/[BH]) vs. log([Pt/21“2[BHP]) with 0.1
CsH2PC>4, 0.01M CsH2P04, and 0.010N Mg(C>2CCH3)2
as supporting electrolytes, pH 5.90, 25.0°C
A 0.15M CsH2P04
• 0.010M CsH2P04
O 0.010N Mg(02CCH3)2

144

Figure 11. Log of the apparent association constant, Ks, for the surface binding of
3-aminoacridinium cation to DNA vs. the square root of the ionic strength,
I1/2# NaH2P04 and LÍH2PO4 as supporting electrolytes, pH 5.90, 25.0°C
X NaH2P04
0 LiH2P04

LOG
146

Figure 12. Log of the apparent association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the ionic
strength, I^/^. KH2PO4 and KO2CCH3 as supporting electrolytes, pH 5.90,
25.0°C
X KH2P04
0 ko2cch3

LOG
550-
«/>
X.
5.00-
4.50
V\
sS
0.10
143

Figure 13. Log of the apparent association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the ionic
strength, 1^-/2. KH2PO4 and RbH2PC>4 as supporting electrolytes, pH 5.90,
25.0°C
X kh2P04
0 RbH2P04

150

Figure 14. Log of the apparent association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the ionic
strength, I-*-/2. CSH2PC>4 and (CH3) 4NH2PO4 as supporting electrolytes,
pH 5.90, 25.0°C

O G
152

Figure 15. Log of the apparent association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the ionic
strength, Mg(C>2CCH3)2 and Ca(C>2CCH3)2 as supporting electrolytes,
pH 5.90, 25.0°C
X Mg(02CCH3)2
0 Ca(02CCH3)2

5.0-
154

Figure 16. Log of the apparent association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the ionic
strength, I1/2. Sr(C>2CCH3)2 and Ba (O2CCH3) 2 as supporting electrolytes,
pH 5.90, 25.0°C
X Sr(02CCH3)2
0 Ba(02CCH3)2

156

Figure 17. Log of the apparent association constant, Kj, for the intercalative
binding of 3-aminoacridinium cation to DNA vs. the square root of
the ionic strength, LÍH2PO4 and NaH2PC>4 as supporting electro¬
lytes, pH 5.90, 25.0°C
0 LiH2P04
X NaH2P04

LOG
153

Figure 18. Log of the apparent association constant, Kj, for the intercalative
binding of 3-aminoacridinium cation to DNA vs. the square root of the
ionic strength, KH2PO4 and KO2CCH3 as supporting electrolytes,
pH 5.90, 25.0°C
X kh2po4
0 ko2cch3

160

Figure 19.
Log of the apparent association constant, Kj, for the intercalative
binding of 3-aminoacridinium cation to DNA vs. the square root of the
ionic strength, 11/2. KH2PO4 and RbH2PC>4 as supporting electrolytes,
pH 5.90, 25.0°C
kh2po4
X
0
RbH2P04

162
z/\l

Figure 20.
Log of the apparent association constant, Kj, for the intercalative
binding of 3-aminoacridinium cation to DNA vs. the square root of
the ionic strength, I-*-/2. CSH2PC>4 and (CH3)4NH2P04 as supporting
electrolytes, pH 5.90, 25.0°C
X CsH2P04
(ch3)4nh2po4

v'1
O'
s*

Figure 21. Log of the corrected association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the activity
of solution, 1^-/2. LÍH2PO4, NaH2PC>4, and KH2PO4 as supporting electro¬
lytes, pH 5.90, 25.0°C
O LiH2P04
• NaH2P04
â–¡
KH2p°4


Figure 22. Log of the corrected association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the activity
of solution, 1-*-/^. RbH2PC>4, CSH2PC>4, and (CH3) 4NH2PO4 as supporting
electrolytes, pH 5.90, 25.0°C
O RbH2P04
• CsH2P04
(CH3>4NH2P04

891

I
Figure 23. Log of the corrected association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the activity
of solution, I1/2. Mg (O2CCH3) 2 and Ca(C>2CCH3)2 as supporting electro¬
lytes, pH 5.90, 25.0°C
O Mg(02CCH3)2
Ca (02CCH3)

4.9-

Figure 24. Log of the corrected association constant, Ks, for the surface binding
of 3-aminoacridinium cation to DNA vs. the square root of the activity
of solution, I1/2. Sr(C>2CCH3)2 and Ba(02CCH3)2 as supporting electro¬
lytes, pH 5.90, 25.0°C
O Sr(02CCH3)2
• Ba(02CCH3)2

LOG
ZLI

Figure 25. Absorptiometric titration of 7-aminoquinolinium cation with DNA,
potassium salt, pH 5.90, 25.0°C. Background electrolyte: KH_PO.,
initial concentration, 0.0050M
Absorbances of individual spectra normalized to 0.00 at 500 nm
B
Individual
spectra as
a
[P/D]
=
-([P] = 0)
b
[P/D]
=
7.56
c
[P/D]
=
18.9
d
[P/D]
=
37.8
e
[P/D]
=
75.6
f
[P/D]
=
105.8
P/D is
the
ratio
of
total DNA
cation,
D ¡
= 1.08
X
10“^ moles

ABSORBANCE
350 450
WAVELENGTH (nm)
I T
350 450
WAVELENGTH (nm)
A
B
-*n> Qo

Figure 26. Log([AHP]/[AH]) vs. log([Pt/3]-[AHP]) for the binding of 7-amino-
quinolinium cation to DNA at 15.0°C, 25.0°C, and 35.0°C, pH 5.90,
0.010M KH2P04 as supporting electrolyte
15.0°C
25.0°C
35.0°C

♦1.5H
■—^i.ch
E1
*0.5-1
©
O
OH
-0.51
v
L 0 G(Q>t/|] - |AH0 )
176

Figure 27. Log([AHP]/[AH]) vs. log([Pt/2]-2[AHP]) for the binding of 7-amino
quinolinium cation to DNA at 15.0°C, 25.0°C, and 35.0°C, pH 5.90,
0.010M KH2PO4 as supporting electrolyte.
15.0°C
25.0°C
35.0°C

LOG(|AHB/iag)
173

* o ....
Figure 28. Log K vs. 1/T( K) for the reaction between 3-aminoacridiniuui
and dSa at 15.0°C, 25.0°C, and 35.0°C, pH 5.90, 0.010M KH2PC>4
supporting electrolyte
cation
as

3.30 3.40
LOG K¿
en en tn
fr q»
081

1 o
Figure 29. Log K vs. 1/T( K) for the reaction between
and D§A at 15.0°C, 25.0°C, and 35.0°C, pH 5
supporting electrolyte
7-aminoquinolinium cation
90, 0.010M KH2P04 as

4.7-1
182

* o ...»
Figure 30. Log Ki vs. 1/T( K) for the reaction between 3-aminoacndinium
and DNA at 15.0°C, 25.0°C, and 35.0°C, pK 5.90, 0.010M KH2P04
supporting electrolyte
cation
as

10.40-
10.20-
° 10.00-
o
9.80-
330
l/T (°K x I 03)
3.40
184

* o . . . . •
Figure 31. Log Kj vs. 1/T( K) for the reaction between 7-aminoguinolinium cation
and DNA at 15.0°C, 25.0°C, and 35.0°C, pH 5.90, 0.010M KH2P04 as
supporting electrolyte

Xo) 1/1
^>1
I
LOG k',
00
981
8.5 A

APPENDIX II
TABLES

a« l
Table l
. Molar absorptivities of the 3-aminoacridinium-DNA complex and the 7-amino-
quinolinium-DNA complex in various supporting electrolytes3
3-aminoacridinium-DNA complex
Supporting
electrolyte
Cone. range
of electro¬
lyte (M)
Number of
determina¬
tions (n)
368 ,n-3
£BHP X 10
-i -1
M cm
Standard
deviation
(a) x 10 3
463 ._-3
£BHP X 10
M-1cm-1
Standard
deviation
(a) x 10
LiH2P04
.025-.00063
8
8.52 7
0.111
9.9097
0.0804
NaH.PO.
2 4
0.025-0.0025
6
8.477
0.073
9.090
0.039
kh2po4
0.025-0.00063
11
8.557
0.082
9.082^
0.119
K02CCH3
0.025-0.00063
8
8.515
0.059
9.104
0.040
RbH2P04
0.025-0.0025
10
8.456
0.133
9.067
0.145
CsH»P0.
2 4
0.025-0.0025
6
8.763
0.069
9.353
0.071
(ch3)4nh2po4
0.10-0.0025
18
8.429
0.125
8.98 8^
0.123
£bhp
6
8.4 94
0.048
9.071
0.043
7-aminoquinolinium-
DNA complex
kh2p°4
0.010
10
5.332
0.094
-420
£AHP
10
5.332
aAll at pH 5.90 and 25.0°C. Initial concentration of drug species was 1 to 2 x 10 5M.

Table 2
Absorptiometric titration of 3-aminoacridinium cation with DNA, lithium salt.
Lithium dihydrogen phosphate as supporting electrolyte a
0.10Mb
0.05Mb
Run
1
Run
2
Run
1
Run
2
Vol
Abs^
Vol _
Absd
Vol
-. d
Abs
Volf
Abs^
DNAC
DNA1
DNAC
DNA
0
0.980
0
0.645
0
0.976
0
0.670
10.0
0.957
15.0
0.623
10.0
0.945
10.0
0.662
20.0
0.936
30.0**
0.600
20.0**
0.920
20.0
0.640
40.0**
0.900
45.0**
0.581
30.0**
0.893
35.0**
0.613
60.0**
0.865
65.0**
0.554
40.0**
0.870
50.0**
0.589
80.0**
0.835
85.0**
0.538
60.0**
0.822
65.0**
0.567
100**
0.809
105
0.517
80.0
0.783
85.0
0.540
125
0.779
130
0.497
100
0.750
105
0.518
150
0.751
155
0.480
125
0.712
125
0.501
200
0.710
180
0.468
150*
0.683
150
0.482
250*
0.680
220*
0.452
175*
0.661
175*
0.467
300*
0.658
260*
0.438
200*
0.642
210*
0.452
400*
0.627
310*
0.426
250*
0.615
250*
0.437
500*
0.603
410*
0.410
300
0.600
295*
0.430
600
0.591
560
0.399
350
0.588
345
0.423
700
0.582
780
0.392
450
0.572
445
0.414
A e
CO
0.551
A 6
OO
0.360
Q)
8
0.556
A e
CO
0.375
189

Table 2 (Continued)
0.025Mb
Run
1
Vol
DNAC
d
Abs
0
0.983
10.0
0.951
20.0**
0.919
30.0**
0.888
40.0**
0.860
60.0
0.806
80.0
0.758
100*
0.718
125*
0.679
150*
0.649
200*
0.612
250
0.593
300
0.582
400
0.577
500
0.571
A e
CO
0.563
Run
1
Vol
DNAg
Abs
0
0.963
10.0
0.922
20.0**
0.885
30.0**
0.849
40.0**
0.812
60.0
0.750
80.0*
0.697
100*
0.653
125*
0.617
150
0.594
175
0.582
200
0.578
225
0.572
275
0.568
A e
OO
0.564
010Mb 0.0050Mb
Run
2
Run
1
Volf
Abs^
Vol
Abs^
DNA
DNAC
0
0.690
0
0.040
5.0
0.673
10.0
0.986
12.5
0.649
20.0**
0.945
20.0**
0.626
30.0**
0.904
30.0**
0.596
40.0**
0.865
40.0**
0.570
50.0
0.829
50.0
0.546
60.0
0.797
60.0
0.523
70.0*
0.767
70.0*
0.503
80.0*
0.742
80.0*
0.488
100*
0.697
95.0*
0.465
120*
0.665
110*
0.449
140
0.643
130*
0.433
160
0.632
155
0.423
210
0.619
205
0.415
260
0.614
305
0.414
A e
OO
0.610
A e
OO
0.403
190

Table 2 (Continued)
0.0050Mb
Run 2
Vol
-m. . d
Abs
DNAC
0
0.973
10.0
0.910
20.0**
0.868
30.0**
0.823
40.0
0.787
50.0
0.752
60.0*
0.721
80.0*
0.667
100*
0.624
125
0.589
150
0.571
175
0.568
A e
0.575
Run 3
Volf
d
Abs
DNA
0
0.657
7.5
0.630
15.0**
0.604
22.5**
0.583
30.0**
0.561
37.5
0.542
45.0
0.523
55.0
0.499
65.0*
0.475
75.0*
0.455
85.0*
0.441
95.0*
0.429
110
0.410
130
0.400
180
0.394
280
0.393
A e
0.384
0.0025Mb
Run 1
Vol
DNAC
.. d
Abs
0
0.990
10.0
0.925
20.0**
0.878
30.0**
0.831
40.0
0.791
50.0
0.752
60.0*
0.719
80.0*
0.660
100
0.618
125
0.591
150
0.581
200
0.575
A e
0.584
Run 2
Vol
DNAC
Abs^
0
0.981
10.0**
0.932
20.0**
0.883
30.0**
0.840
40.0**
0.800
50.0
0.765
60.0
0.741
70.0*
0.707
80.0*
0.678
90.0*
0.655
100*
0.638
120
0.610
140
0.592
190
0.584
240
0.579
A 6
0.576
191

Table 2 (Continued)
0.00063Mb
Run 1
Run 2
Volf Absd
j
Volf Absa
DNA DNA
0
0.772
0
0.068
5.0
0.702
5.0
0.654
10.0
0.683
10.0
0.633
15.0*
0.662
15.0**
0.615
22.5**
0.636
22.5**
0.587
30.0**
0.607
30.0**
0.558
37.5
0.581
37.5
0.530
45.0
0.554
45.0
0.508
52.5*
0.530
52.5*
0.483
60.0*
0.508
60.0*
0.464
70.0*
0.483
70.0*
0.438
80.0*
0.464
80.0
0.422
90.0
0.451
90.0
0.413
110
0.439
110
0.407
160
0.435
210
0.405
260
0.434
310
0.404
>
8
fl>
0.423
A e
oo
0.397
aT = 25.0 + 0.1°C,
pH = 5.90
192

Table 2 (Continued)
Initial total concentration of supporting electrolyte.
'Total volume of solution of calf thyrm
moles of DNA phosphate per microliter,
c — 9
Total volume of solution of calf thymus DNA added (yl) . Titrant contained 9.300 x 10
aTotal absorbance of sample at 365 nm, 4.00 cm pathlength. Data marked (**) used to
compute K , those marked (*) were used to compute K,.
S _L
G
Calculated absorbance of solution if all 3-aminoacridinium cation is bound. Based on
molar absorptivity of bound species of 84 94M"-'-cm“1.
f -9
Same as c above except titrant contained 8.373^ x 10 moles of DNA phosphate per
microliter.
Q — 9
3Same as c above except titrant contained 9.416^ x 10 moles of DNA phosphate per
microliter.

Table 3
Absorptiometric titration of 3-aminoacridinium cation with DNA, sodium
salt. Sodium dihydrogen phosphate as supporting electrolyte a
0.100M
b
0.050Mb
Run 1
Vol
Absd
Vold
DNAC
DNA
0
0.940
0
10.0
0.912
10.0
20.0
0.891
20.0
30.0**
0.872
40.0
40.0**
0.851
60.0
60.0**
0.812
80.0
80.0**
0.780
100**
100
0.752
125
150
0.691 •
150
200*
0.648
175
250*
0.619
200*
300*
0.602
250*
400
0.578
300*
500
0.564
350*
600
0.556
400*
700
0.550
500
A 6
00
0.528
650
800
888
A e
Run 2 Run
Abs Vol
DNAC
0.953
0
0.929
10.0
0.908
20.0
**
0.870
30.0**
**
0.832
40.0**
* *
0.802
60.0**
0.776
80.0**
0.745
100
0.720
125*
0.698
150*
0.679
200*
0.650
250
0.626
300
0.610
350
0.598
400
0.577
500
0.560
600
0.550
A e
oo
0.543
0.526
1
Run
2
.. d
Abs
Vol
DNAC
Absd
0.948
0
0.961
0.919
10.0
0.930
0.889
20.0
0.900
0.862
30.0**
0.871
0.835
40.0**
0.848
0.791
60.0**
0.800
0.750
80.0
0.758
0.713
100
0.720
0.677
125*
0.683
0.648
150*
0.652
0.608
175*
0.637
0.583
200*
0.610
0.570
250
0.585
0.561
300
0.572
0.558
350
0.563
0.552
400
0.561
0.551
500
0.556
0.536
600
0.555
A e
OO
0.545
fr6T

Table 3 (Continued)
0.025Mb
Run 1
Vol
Absd
DNAC
0
0.958
10.0
0.920
20.0**
0.887
30.0**
0.853
40.0**
0.828
60.0
0.772
80.0*
0.722
100*
0.686
125*
0.644
150*
0.615
175
0.597
200
0.583
225
0.578
250
0.570
275
0.567
325
0.564
375
0.560
425
0.560
A 6
00
0.553
Run 2
Vol
Absd
DNAC
0
0.973
10.0
0.934
20.0**
0.900
30.0**
0.868
40.0**
0.837
60.0
0.781
80.0*
0.730
100*
0.689
125*
0.650
150*
0.621
175
0.601
200
0.587
225
0.578
250
0.572
300
0.569
350
0.565
A 6
0.565
0.010Mb
Run 1
Run 2
Vol Absd
. c
Vol Absd
c
0
0.952
0
0.955
10.0
0.905
10
. 0
0.910
20.0**
0.866
20
.0**
0.868
30.0**
0.829
30
.0**
0.831
40.0**
0.795
40
.0
0.795
60.0
0.730
60
.0**
0.730
80.0*
0.675
80
.0**
0.675
100*
0.631
100
* ★
0.631
125
0.593
125
0.595
150
0.572
150
0.574
175
0.562
175
0.564
200
0.558
200
0.561
225
0.558
250
0.559
275
0.552
300
0.558
e
A
0.557
A 6
0.558
195

Table 3 (Continued)
0.005 0Mb 0.0025M
Run 1
Run 1
Vol
Absd
Volc
_ â– , d
Abs
DNAC
DNA
0
0.951
0
0.948
10.0
0.887
10.0**
0.890
20.0**
0.843
20.0**
0.842
30.0**
0.800
30.0**
0.800
40.0**
0.762
40.0
0.761
60.0
0.691
50.0
0.727
80.0*
0.634
60.0*
0.694
100*
0.597
80.0*
0.636
125
0.571
100
0.595
150
0.562
125
0.567
175
0.560
150
0.559
200
0.557
175
0.555
A S
0.561
e
A
0.560
oo
OO
aT = 25.
0 + 0.1°C, pH = 5.90.
^Initial
total concentration of
supporting
electrolyte
c
Total volume of solution of calf thymus DNA added (yl) . Titrant contained 8.86£ x 10
moles of DNA phosphate per microliter.

Table 3 (Continued)
aTotal absorbance of sample at 365 ran, 4.00 cm pathlength. Data points marked
used to compute K , those marked (*) were used to compute K, .
'Calculated absorbance of solution if all 3raminoacridinium cation is bound,
molar absorptivity of bound drug of
8494M“^cm_-1-,
(**)
Based on

Table 4
Absorptiometric titration of 3-aminoacridinium cation with DNA, potassium
salt. Potassium dihydrogen phosphate as supporting electrolyte a
0.10Mb
0.050Mb
Run 1
Run 2
Vol Absd
c
Vol Absd
n
0
0.700
0
0.703
10.0
0.688
10.0
0.675
20.0
0.665
20.0**
0.659
30.0**
0.649
35.0**
0.632
45.0**
0.624
50.0**
0.612
60.0**
0.602
65.0**
0.591
75.0**
0.584
85.0
0.568
90.0
0.567
105
0.547
110
0.544
125
0.527
130
0.528
145
0.513
150
0.511
170*
0.494
175*
0.494
195*
0.480
200*
0.480
220*
0.470
225*
0.470
270*
0.452
275*
0.452
320*
0.439
325*
0.441
370
0.430
400*
0.430
470
0.420
500
0.419
670
0.411
700
0.415
870
0.404
1000
A 6
m
0.409
0.383
A e
OO
0.389
Run
1
Run
2
Vol
DNAC
Absd
Vol
DNAC
Absd
0
0.703
0
0.699
5.0
0.687
7.5
0.682
10.0
0.673
15.0**
0.663
20.0**
0.652
25.0**
0.640
30.0**
0.632
35.0**
0.621
40.0**
0.612
45.0**
0.602
50.0**
0.596
60.0**
0.572
65.0**
0.569
75.0
0.548
80.0
0.547
90.0
0.528
95.0
0.526
105*
0.510
110*
0.508
125*
0.489
130*
0.489
145*
0.473
150*
0.473
170*
0.458
175*
0.459
195*
0.447
200*
0.447
235
0.434
250
0.431
285
0.427
300
0.422
385
0.418
400
0.415
465
0.415
600
A e
0.414
0.399
A e
OO
0.402
oo
198

Table 4 (Continued)
.025Mb
Run
1
Run
2
Vol
DNAC
a
Abs
Vol
DNAC
Absb
0
0.718
0
0.713
10.0
0.687
10.0
0.679
20.0**
0.656
20.0**
0.654
30.0**
0.630
30.0**
0.627
40.0**
0.605
40.0**
0.605
50.0
0.584
50.0
0.586
60.0
0.565
65.0
0.554
70.0
0.546
80.0*
0.528
85.0*
0.519
95.0*
0.504
100*
0.498
110*
0.486
115*
0.480
125*
0.470
130*
0.467
145*
0.452
150*
0.452
170
0.442
175
0.441
220
0.425
200
0.435
270
0.418
300
0.427
370
0.416
400
A e
OO
0.425
0.415
A e
OO
0.414
.010Mb
Run
1
Run
2
Vol
DNAC
d
Abs
Vol
DNA
d
Abs
0
0.667
0
0.716
5.0
0.649
7.5
0.687
10.0**
0.631
15.0**
0.662
17.5**
0.606
22.5**
0.635
25.0**
0.583
30.0**
0.614
32.5**
0.562
40.0
0.587
40.0
0.542
50.0
0.563
50.0
0.517
60.0
0.537
60.0*
0.497
70.0*
0.516
70.0*
0.477
80.0*
0.496
80.0*
0.460
90.0*
0.480
90.0*
0.443
100*
0.467
100*
0.431
115
0.450
115*
0.419
135
0.435
130
0.410
160
0.426
180
0.399
200
0.422
280
0.396
300
0.420
A e
0.390
A 6
0.418
199

Table 4 (Continued)
0.0050Mb
Run 1
Vol
Abs
DNAC
0
0.714
5.0
0.697
10.0**
0.679
15.0**
0.664
22.5**
0.641
30.0**
0.618
40.0
0.592
50.0
0.565
60.0
0.547
70.0*
0.522
85.0*
0.496
100*
0.470
115*
0.453
130
0.442
155
0.434
230
0.426
330
0.424
A e
0.416
Run 2
Vol
DNAC
Absb
0
0.666
7.5
0.627
15.0**
0.600
22.5**
0.578
30.0
0.554
37.5
0.532
45.0
0.512
52.5
0.497
60.0*
0.479
70.0*
0.460
80.0*
0.443
90.0*
0.428
110
0.407
130
0.397
170
0.389
270
0.388
A e
0.390
T =
25.0 + 0.1°C
pH = 5.90.
0.00063Mb
Run 1
Vol
Abs^
DNAC
0
0.669
5.0
0.637
10.0**
0.612
15.0**
0.593
20.0**
0.576
27.5**
0.548
35.0
0.521
42.5
0.500
50.0*
0.477
57.5*
0.454
65.0*
0.439
75.0
0.421
85.0
0.410
100
0.407
125
0.405
A e
0.397
Run 2
Vol
Abs^
DNAC
0
0.698
5.0
0.668
10.0**
0.649
15.0**
0.629
20.0**
0.612
27.5**
0.583
35.0
0.560
42.5
0.535
50.0*
0.510
57.5*
0.489
65.0*
0.469
75.0*
0.449
85.0
0.431
95.0
0.423
115
0.416
165
0.415
A 6
0.413
200

Table 4 (Continued)
-9
Initial total concentration of supporting electrolyte.
**
'Total volume of solution of calf thymus DNA added (yl) . Titrant contained 8.09 x 10
moles of DNA phosphate per microliter.
Total absorbance of sample at 365 nm, 4.00 cm pathlength. Data marked (**) used to
compute K , those marked (*) were used to compute K. .
b -L
*
'Calculated absorbance of solution if all 3-aminoacridinium cation is bound. Based on
molar absorptivity of bound species of 8494M-lcm--1-.
201

Table 5
Absorptiometric titration of 3-aminoacridinium cation with DNA, potassium
salt. Potassium acetate as supporting electrolyte a
0.15Mb
0.05 0Mb
Run 1
Vol
Abs^
DNAC
0
0.721
10.0
0.706
20.0
0.690
40.0**
0.662
60.0**
0.638
80.0**
0.606
100**
0.597
125
0.572
150
0.553
175
0.537
200
0.523
250*
0.500
300*
0.483
350*
0.467
400*
0.458
450*
0.449
550*
0.437
650
0.429
850
0.422
1050
0.415
1300
0.414
A e
0.384
Run 2
Vol
DNAC
Abs^
0
0.747
20.0
0.716
40.0**
0.686
60.0**
0.659
80.0**
0.638
100**
0.619
125
0.597
150
0.578
175
0.562
200*
0.547
250*
0.522
300*
0.505
350*
0.492
450
0.470
550
0.459
700
0.449
900
0.441
A e
0.412
Run 1
Vol
DNAC
Abs^
0
0.724
10.0
0.698
20.0**
0.677
30.0**
0.653
40.0**
0.631
50.0**
0.611
60.0
0.593
70.0
0.577
85.0
0.553
100
0.533
115*
0.516
130*
0.500
150*
0.485
175*
0.467
200*
0.455
225
0.448
275
0.437
325
0.430
425
0.423
625
0.422
A 6
0.410
oo
Run 2
Vol
Abs^
DNAC
0
0.731
10.0
0.705
20.0**
0.683
30.0**
0.663
40.0**
0.642
50.0**
0.623
65.0
0.594
85.0
0.571
95.0
0.549
110*
0.531
130*
0.510
150*
0.493
175*
0.477
200*
0.463
225
0.454
275
0.443
375
0.433
575
0.427
a/
0.416
202

Table 5 (Continued)
0.025Mb
Run 1
Run 2
Vol Absd
c
Vol Absd
c
DNA DNA
0
0.713
0
0.749
10.0
0.684
10.0
0.718
20.0**
0.655
20.0**
0.693
35.0**
0.613
30.0**
0.668
50.0**
0.578
40.0**
0.644
65.0
0.546
50.0
0.622
80.0*
0.517
65.0
0.591
95.0*
0.496
80.0
0.565
110*
0.477
95.0*
0.542
125*
0.461 •
110*
0.523
145*
0.445
130*
0.504
165
0.433
150*
0.487
185
0.426
175*
0.474
210
0.420
210
0.461
260
0.414
260
0.449
310
0.411
360
0.438
510
A e
no
0.409
0.408
A e
00
0.435
0.0050M
b
Run 1
Vol
DNAC
Absd
0
0.704
10.0**
0.656
15.0**
0.636
20.0**
0.619
27.5**
0.592
35.0
0.569
42.5
0.547
50.0
0.527
60.0*
0.501
70.0*
0.480
80.0*
0.465
90.0*
0.450
100
0.439
120
0.424
145
0.418
190
0.419
A e
0.415
Run 2
Vol
DNAC
Abs
0
0.739
7.5
0.695
15.0
0.670
22.5**
0.648
30.0**
0.620
37.5
0.600
45.0
0.581
55.0
0.552
65.0*
0.529
75.0*
0.513
85.0*
0.496
A 6
0.441
00
203

Table 5 (Continued)
0.0050M
0.0025M
Run
3
Run
1
Run
2
Vol
Absd
Vol
Absd
Vol
Absd
DNAC
DNAC
DNAC
0
0.720
0
0.721
0
0.727
7.5
0.689
5.0
0.700
7.5
0.691
15.0**
0.662
10.0**
0.679
15.0**
0.665
22.5**
0.637
17.5**
0.649
22.5**
0.637
30.0**
0.610
25.0**
0.623
30.0
0.612
37.5
0.588
32.5**
0.598
37.5
0.588
45.0
0.565
40.0
0.572
45.0
0.564
52.5
0.547
47.5
0.550
55.0*
0.533
60.0*
0.526
55.0*
0.528
65.0*
0.509
70.0*
0.504
62.5*
0.508
75.0*
0.486
80.0*
0.485
70.0*
0.490
85.0*
0.467
90.0*
0.466
80.0*
0.469
95.0
0.453
100
0.454
90.0
0.452
105
0.443
110
0.447
100
0.440
130
0.430
140
0.432
110
0.432
180
0.429
240
0.429
130
0.424
A e
oo
0.429
A e
00
0.423
180
0.422
A e
0.426
204

Table 5 (Continued)
0.00063Mb
Run
1
Run
2
Vol
Abs^
Vold
Abs^
DNAC
DNA
0
0.715
0
0.713
5.0
0.693
7.5
0.688
10.0**
0.675
15.0**
0.665
15.0**
0.656
22.5**
0.640
20.0**
0.638
30.0**
0.617
27.5**
0.610
37.5
0.593
35.0**
0.587
45.0
0.573
42.5
0.560
55.0
0.543
50.0
0.536
65.0*
0.517
57.5*
0.514
75.0*
0.496
65.0*
0.495
85.0*
0.474
75.0*
0.468
105
0.442
85.0*
0.449
125
0.429
95.0
0.437
150
0.427
115
0.427
A e
OO
0.422
165
0.424
265
0.424
A e
oo
0.419
aT = 25.0 + 0.1°C
/
pH = 5.90
205

Table 5 (Continued)
T_
^Initial total concentration of supporting electrolyte.
cTotal volume of solution of calf thymus DNA added (yl) . Titrant contained 8.16^3 x 10
moles of DNA phosphate per microliter.
aTotal absorbance of sample at 365 nm, 4.00 cm pathlength. Data marked (**) used to
compute Kg, those marked (*) were used to compute .
'Calculated absorbance of solution if all 3-aminoacridinium cation is bound,
molar absorptivity of bound species of 84 94M'
n â– 'â– cm -'â– 
Based on

Table 6
Absorptiometric titration of 3-aminoacridinium cation with DNA, rubidium
salt. Rubidium dihydrogen phosphate as supporting electrolyte a
0.10Mb
0.050Mb
Run 1
Vol
Abs^
DNAC
0
0.760
20.0**
0.723
40.0**
0.691
60.0**
0.662
80.0**
0.637
100
0.613
125
0.588
150
0.569
175
0.550
200*
0.535
225*
0.526
250*
0.515
300*
0.498
350*
0.485
400*
0.476
500
0.462
600
0.456
700
0.450
A e
0.427
00
Run 2
Vol
DNAC
-. a
Abs
0
0.738
20.0**
0.705
40.0**
0.673
60.0**
0.643
80.0**
0.617
100
0.595
125
0.569
150
0.548
175*
0.531
200*
0.517
225*
0.502
250*
0.493
275*
0.484
300*
0.480
350*
0.467
400*
0.458
450
0.450
550
0.440
650
0.437
750
0.432
A 6
OO
0.413
Run 1
Vol
DNAC
Abs^
0
0.720
10.0
0.698
20.0**
0.678
40.0**
0.638
60.0**
0.603
80.0
0.571
100
0.546
120
0.525
140*
0.506
160*
0.490
180*
0.478
205*
0.465
255
0.455
280
0.450
305
0.438
405
0.427
505
0.422
A 6
0.412
oo
Run 2
Vol
Abs^
DNAC
0
0.748
25.0**
0.689
50.0**
0.641
75.0
0.598
100
0.564
125*
0.536
150*
0.513
175*
0.495
200*
0.482
215*
0.478
230*
0.471
250
0.464
270
0.459
295
0.454
345
0.447
445
0.439
545
0.435
A 6
0.423
207

Table 6 (Continued)
0.025Mb
Run 1
Vol
Abs^
DNAC
0
0.778
10.0
0.744
20.0**
0.719
30.0**
0.695
40.0**
0.673
60.0
0.630
80.0
0.593
100*
0.561
120*
0.535
140*
0.517
160*
0.499
185*
0.485
210
0.475
260
0.462
310
0.458
410
0.454
A e
0.445
Run 2
Vol
DNAC
d
Abs
0
0.735
20.0**
0.674
40.0**
0.627
60.0
0.585
80.0
0.549
100*
0.520
115*
0.502
130*
0.483
145*
0.474
160*
0.463
175
0.457
200
0.445
225
0.440
275
0.437
350
0.432
A e
0.427
0.010Mb
Run 3
Vol Abs^
DNAC
0
0.719
20.0**
0.662
40.0**
0.613
60.0
0.571
80.0
0.536
100*
0.506
115*
0.490
130*
0.476
145*
0.462
160
0.454
175
0.445
200
0.435
225
0.429
275
0.422
350
0.418
8
c
0.418
Run 1
Vol
Abs^
DNAC
0
0.701
10.0
0.673
20.0**
0.644
30.0**
0.613
40.0
0.585
60.0
0.537
80.0*
0.496
100*
0.467
120*
0.444
140
0.431
160
0.420
210
0.413
260
0.411
A e
0.411
OO
208

Table 6 (Continued)
0.010Mb
Run 2
Vol
Absb
DNAC
0
0.702
10.0**
0.667
30.0**
0.600
50.0
0.549
70.0*
0.507
80.0*
0.489
90.0*
0.471
100*
0.458
110*
0.447
115
0.442
125
0.434
135
0.428
155
0.419
205
0.407
A e
0.414
Run 1
Vol
Abs^
DNAC
0
0.703
10.0**
0.671
20.0**
0.641
30.0**
0.610
40.0**
0.584
50.0
0.560
60.0
0.537
70.0*
0.518
85.0*
0.488
100*
0.467
115*
0.452
130*
0.440
150
0.429
175
0.420
225
0.416
A e
0.413
0.0050Mb
Run 2
Vol
DNAC
Abs^
0
0.742
10.0
0.687
20.0**
0.651
30.0**
0.619
40.0
0.589
50.0
0.562
65.0*
0.528
80.0*
0.500
95.0*
0.477
110
0.456
125
0.443
140
0.434
155
0.429
180
0.427
230
0.421
A e
0.436
Run 3
Vol
DNA
.. d
Abs
0
0.745
10.0**
0.709
20.0**
0.678
30.0**
0.647
40.0**
0.619
50.0
0.591
65.0*
0.556
80.0*
0.524
95.0*
0.499
110*
0.478
125*
0.463
140
0.452
155
0.448
170
0.441
210
0.438
A e
0.439
209

Table 6 (Continued)
0.0025Mb
Run 1
Run 2
Vol Absd
c
Vol Absd
c
0
0.752
0
0.725
10.0
0.714
10.0
0.688
20.0**
0.683
20.0**
0.656
30.0**
0.653
30.0**
0.627
40.0**
0.624
40.0**
0.600
50.0
0.597
50.0
0.572
60.0
0.574
60.0
0.549
70.0*
0.550
70.0*
0.528
80.0*
0.531
80.0*
0.508
90.0*
0.510
90.0*
0.491
110*
0.479
100*
0.478
130
0.459
110*
0.465
150
0.446
120
0.455
175
0.439
135
0.443
225
0.431
150
0.436
A e
00
0.442
175
0.4 30
225
0.428
A e
0.426
aT = 25.0 + 0.1°C, pH =
5.90
o

Table 6 (Continued)
Initial total concentration of supporting electrolyte.
'Total volume of solution of calf thymu
moles of DNA phosphate per microliter.
c t — 9
Total volume of solution of calf thymus DNA added (yl). Titrant contained 7.209^ x 10
1Total absorbance of sample at 365 nm, 4.00 cm pathlength. Data marked (**) used to
compute K , those marked (*) were used to compute K-, .
'Calculated absorbance of solution if all 3-aminoacridinium cation is bound. Based on
molar absorptivity of bound species of 8 4 94M~-'-cm_1.
211

Table 7
Absorptiometric titration of 3-aminoacridinium cation with DNA, cesium salt.
Cesium dihydrogen phosphate as supporting electrolyte a
0.15Mb
0.10Mb
Run
1
Run
2
Vol
DNAC
Abs^
Vol
DNAC
Abs^
0
0.868
0
0.763
10.0
0.844
15.0
0.738
25.0**
0.829
30.0**
0.716
40.0**
0.803
45.0**
0.692
55.0**
0.783
60.0**
0.672
70.0**
0.763
75.0**
0.654
85.0**
0.747
90.0**
0.635
100**
0.729
110**
0.613
120
0.708
130
0.596
140
0.685
150
0.582
160
0.669
170
0.570
180
0.654
195
0.552
205
0.633
225*
0.540
230*
0.622
265*
0.521
280*
0.596
315*
0.507
330*
0.578
365*
0.497
380*
0.564
415*
0.488
480*
0.544
515*
0.474
580*
0.533
715*
0.460
760
0.525
1015
0.454
A 6
oo
0.485
>
8
(D
0.417
Run 1
Run 2
Vol Abs
c
_q
Vol Absa
c
0
0.782
0
0.749
20.0**
0.742
10.0
0.715
35.0**
0.713
20.0** •
0.691
50.0**
0.687
30.0**
0.676
60.0**
0.673
40.0**
0.659
70.0**
0.661
50.0**
0.640
85.0*
0.639
65.0
0.621
100
0.627
80.0
0.597
120
0.600
95.0
0.581
140
0.584
110
0.565
160
0.569
130
0.545
185*
0.550
150
0.530
210*
0.539
170*
0.516
235*
0.527
205*
0.501
285*
0.508
240*
0.485
335*
0.498
290*
0.469
385*
0.488
340
0.458
485
0.479
440
0.443
785
0.471
640
0.437
A e
OO
0.436
840
0.436
A e
on
0.416
212

Table 7 (Continued)
0.05 0Mb
Run 1
Vol
d
Abs
DNAC
0
0.763
5.0
0.752
10.0
0.742
17.5
0.727
25.0**
0.712
32.5**
0.697
40.0**
0.677
50.0**
0.660
60.0**
0.640
70.0
0.627
85.0
0.600
100
0.581
115
0.562
130
0.548
145*
0.533
165*
0.519
190*
0.505
240*
0.485
290*
0.477
390
0.467
590
0.462
A e
0.433
Run 2
Vol
DNAC
,,, d
Abs
0
0.740
7.5
0.714
15.0**
0.699
22.5**
0.679
30.0**
0.659
37.5**
0.648
45.0**
0.631
55.0**
0.612
65.0
0.594
75.0
0.580
85.0
0.560
100
0.542
115*
0.528
130*
0.511
150*
0.495
170*
0.479
195*
0.470
230
0.460
280
0.448
380
0.442
580
0.441
A e
0.421
0.025Mb
Run 1
Vol
Abs^
DNAC
0
0.767
7.5
0.734
15.0**
0.713
22.5**
0.692
30.0**
0.670
37.5**
0.651
45.0**
0.630
55.0
0.610
65.0
0.587
75.0
0.570
90.0*
0.545
105*
0.520
120*
0.506
135*
0.490
155*
0.477
180
0.466
230
0.453
330
0.451
630
0.450
A e
0.434
oo
Run 2
Vol Abs
DNA
0
0.778
7.5
0.752
15.0**
0.733
22.5**
0.719
30.0**
0.702
40.0**
0.680
50.0**
0.661
65.0
0.645
75.0
0.619
90.0
0.594
105
0.577
120*
0.558
140*
0.539
160*
0.522
185*
0.504
210*
0.494
245
0.483
295
0.470
395
0.464
695
0.457
A e
0.437
213

Table 7 (Continued)
0.010Mb
Run 1
Vol
_ d
Abs
DNAC
0
0.749
5.0
0.727
10.0**
0.710
17.5**
0.683
25.0**
0.662
32.5**
0.638
40.0**
0.620
47.5
0.598
55.0
0.580
65.0*
0.557
75.0*
0.539
85.0*
0.518
100*
0.496
120*
0.478
140
0.467
165
0.458
265
0.453
A 6
00
0.439
Run 2
Vol
d
Abs
DNAC
0
0.778
5.0
0.746
10.0
0.730
15.0**
0.713
22.5**
0.691
30.0**
0.669
37.5**
0.648
45.0
0.629
52.5
0.608
60.0*
0.590
70.0*
0.572
80.0*
0.552
90.0*
0.530
105*
0.510
120*
0.493
140
0.477
190
0.461
290
0.455
590
0.454
A e
0.442
0.0025Mb
Run
1
Run
2
Vol
DNAC
a
Abs
Vol
DNAC
Abs^
0
0.752
0
0.773
5.0
0.728
5.0
0.739
10.0
0.701
10.0
0.719
15.0**
0.681
15.0**
0.701
20.0**
0.666
20.0**
0.686
25.0**
0.649
25.0**
0.672
30.0**
0.631
32.5**
0.648
37.5
0.605
40.0
0.626
45.0
0.581
47.5
0.604
52.5*
0.561
55.0
0.584
60.0*
0.539
62.5*
0.569
70.0*
0.515
70.0*
0.551
80.0*
0.497
80.0*
0.534
90.0*
0.481
90.0*
0.518
100
0.467
100*
0.506
110
0.462
130
0.475
160
0.456
230
0.466
210
A e
0.453
0.443
A e
oo
0.454
214

Table 7 (Continued)
aT = 25.0 + 0.1°C, pH = 5.90,
initial total concentration of supporting electrolyte.
'Total volume of solution of calf thymus DNA added (yl)
moles of DNA phosphate per microliter.
compute , those marked (*) were used to compute K
Titrant contained 8.68(3 x 10
Total absorbance of sample at 365 nm, 4.00 cm pathlength. Data marked (**) used to
Based on
-9
'Calculated absorbance of solution if all 3-aminoacridinium cation is bound
molar absorptivity of bound species of 84 94M"-*-cm-^.
215

Table 8
Absorptiometric titration of 3-aminoacridinium cation with DNA, tetra-
methy1ammonium salt. Tetramethylammonium dihydrogen phosphate as supporting
electrolyte a
0.10Mb
0.050M
b
Run
1
Run
2
Vol
DNAC
d
Abs
Vol
DNAC
Abs^
0
0.786
0 '
0.788
20.0**
0.726
15.0*
0.746
40.0**
0.677
30.0*
0.712
60.0**
0.637
45.0*
0.680
80.0
0.600
60.0*
0.650
100*
0.569
80.0
0.614
120*
0.546
100**
0.584
140*
0.524
125**
0.553
160*
0.510
150**
0.530
180*
0.498
200**
0.498
200
0.487
250
0.477
250
0.472
300
0.468
300
0.464
400
0.460
400
0.458
500
0.458
A e
0.455
A e
0.451
Run
1
Run
2
Vol
DNA
Abs^
Vol
DNAC
.. d
Abs
0
0.773
0
0.773
20.0**
0.714
10.0**
0.745
40.0**
0.662
20.0**
0.716
60.0
0.613
30.0**
0.690
80.0
0.575
40.0**
0.664
100*
0.544
60.0
0.621
110*
0.530
80.0
0.581
120*
0.518
100*
0.549
140*
0.498
125*
0.520
160
0.485
150*
0.498
180
0.473
175
0.483
220
0.460
200
0.472
270
0.453
250
0.464
320
0.452
300
0.462
A 6
0.450
A e
0.454
oo oo
216

Table 8 (Continued)
Run 1
Vol
Abs^
DNAC
0
0.778
20.0
0.692
40.0**
0.637
60.0**
0.588
70.0
0.569
80.0
0.550
90.0*
0.532
100*
0.520
115*
0.500
130*
0.486
150
0.473
200
0.458
250
0.456
A e
00
0.456
Run 2
Vol
Abs^
DNAC
0
0.794
10.0**
0.758
20.0**
0.727
30.0**
0.697
40.0**
0.669
60.0
0.613
80.0*
0.569
100*
0.533
120*
0.506
140
0.488
160
0.478
180
0.470
200
0.465
250
0.462
A e
0.466
02 5 0Mb
Run 3
Vol
DNAC
Abs^
0
0.749
20.0**
0.671
40.0**
0.613
60.0
0.548
70.0
0.529
80.0
0.513
90.0*
0.498
95.0*
0.495
100*
0.489
105*
0.485
115
0.474
125
0.467
140
0.457
155
0.450
180
0.442
A e
0.442
CO
Run 4
Vol
Abs
DNA
0
0.762
10.0**
0.723
20.0**
0.692
30.0**
0.663
40.0
0.636
60.0
0.583
80.0*
0.541
100*
0.508
120*
0.481
140
0.465
160
0.453
180
0.449
200
0.445
250
0.441
A e
0.447
217

Table 8 (Continued)
0.010Mb
Run 1
Vol
.. d
Abs
DNAC
0
1.010
10.0**
0.963
20.0**
0.922
30.0**
0.879
40.0
0.840
60.0
0.772
80.0*
0.711
100*
0.668
120*
0.634
140
0.610
160
0.598
200
0.589
250
0.587
A e
00
0.593
Run 2
Vol
DNAC
a
Abs
0
0.672
10.0**
0.630
20.0**
0.592
30.0
0.553
40.0
0.523
60.0*
0.468
80.0*
0.429
100
0.405
120
0.397
140
0.390
160
0.389
A e
0.398
Run 3
Vol
DNAC
Abs*^
0
0.768
20.0**
0.681
40.0**
0.608
60.0*
0.549
70.0*
0.527
80.0*
0.507
90.0*
0.490
100
0.475
110
0.465
130
0.454
155
0.450
205
0.449
A e
0.453
218

Table 8 (Continued)
0.0050Mb
Run 1
Vol
Absd
DNAC
0
0.758
10.0
0.699
20.0**
0.660
30.0**
0.622
40.0**
0.588
50.0
0.558
60.0*
0.531
80.0*
0.488
100
0.462
125
0.449
150
0.445
A e
00
0.449
Run 2
Vol
Absd
DNAC
0
0.756
20.0
0.649
40.0
0.571
50.0
0.540
60.0*
0.518
65.0*
0.507
70.0*
0.499
75.0*
0.489
85.0
0.473
95.0
0.461
115
0.445
140
0.439
A e
0.448
Run 3
Vol
DNAC
—. d
Abs
0
0.764
10.0**
0.713
30.0**
0.629
40.0
0.593
50.0
0.558
60.0*
0.531
65.0*
0.524
70.0*
0.512
80.0*
0.493
90.0
0.476
100
0.462
115
0.450
140
0.445
165
0.442
A e
0.452
219

Table 8 (Continued)
0.0025Mb
Run 1 Run 2
Vol Absd Vol Absd
DNAC DNAC
0
0.768
10.0**
0.719
20.0**
0.678
30.0**
0.638
40.0
0.600
50.0*
0.568
60.0*
0.537
70.0*
0.512
80.0*
0.491
100
0.462
120
0.451
140
0.446
A e
00
0.456
0
0.768
10.0**
0.703
30.0**
0.619
40.0
0.583
50.0*
0.552
60.0*
0.524
70.0*
0.503
75.0*
0.494
80.0
0.487
90.0
0.471
100
0.459
120
0.447
170
0.443
A e
0.454
aT = 25.0 + 0.1°C, pH = 5.90.
Initial total concentration of supporting
Run 3
Vol
DNAC
Absd
0
0.780
10.0**
0.719
30.0**
0.640
40.0
0.605
50.0
0.573
60.0*
0.545
70.0*
0.522
75.0*
0.515
80.0*
0.505
85.0
0.500
90.0
0.492
100
0.479
125
0.460
150
0.456
A e
0.462
electrolyte

Table 8 (Continued)
CTotal volume of solution of calf thymus DNA added (yl) . Titrant contained 8.92£ x 10
moles of DNA phosphate per microliter.
Total absorbance of sample at 365 nm, 4.00 cm pathlength. Data marked (**) used to
compute Kg, those marked (*) were used to compute .
'Calculated absorbance of solution if all 3-aminoacridinium cation is bound,
molar absorptivity of bound species of 8494M"
"“l cm~l
Based on

Table 9
Absorptiometric titration of 3-aminoacridinium cation with DNA, magnesium
salt. Magnesium acetate as supporting electrolyte a
0.010Nb
.0050Nb
Run 1
Vol
DNAC
Absa
0
0.752
25.0
0.727
50.0
0.703
75.0**
0.685
100**
0.664
150**
0.637
200**
0.611
250**
0.590
300
0.575
350
0.560
400
0.548
500
0.527
600
0.513
700
0.500
800
0.490
1000
0.476
1300
0.463
1500
0.460
A e
00
0.393
Run 2
Vol
DNAC
Abs
0
0.742
25.0
0.716
50.0
0.693
75.0**
0.673
100**
0.657
150**
0.625
200**
0.600
250
0.581
300
0.565
350
0.551
400
0.541
500
0.520
600
0.506
700
0.495
800
0.485
900
0.478
1100
0.467
1300
0.459
1650
0.449
A e
OO
0.383
Run 1
Vol
DNA
c
Abs^
0
0.748
10.
0
0.729
20.
0
0.714
35.
0
0.698
50.
0**
0.679
70.
0**
0.661
90.
0**
0.642
115**
0.622
140
0.606
165
0.591
215
0.567
265
0.547
315
0.530
365
0.517
415
0.505
465
0.490
565
0.478
665
0.467
765
0.459
965
0.446
1165
0.440
1465
0.430
& e
0.392
Run 2
Vol
DNAC
Abs
0
0.739
20.0
0.709
40.0
0.686
60.0**
0.669
80.0**
0.649
100**
0.630
125**
0.613
150
0.597
175
0.568
200
0.549
250
0.530
300
0.516
350
0.505
400
0.496
450
0.489
550
0.475
650
0.466
750
0.458
950
0.447
A e
Hoo
0.406
222

Table 9 (Continued)
0.0025Nb
Run 1 Run 2
Vol Absd Vol Absd
DNAC DNAC
0
0.735
20.0
0.698
40.0**
0.670
60.0**
0.646
80.0**
0.624
100**
0.605
125
0.587
150
0.570
200
0.542
250
0.523
300
0.509
350
0.495
400
0.485
500
0.470
600
0.460
800
0.448
1000
0.439
1200
0.435
A e
0.395
0
0.754
20.0
0.717
40.0**
0.687
60.0**
0.664
80.0**
0.643
100**
0.623
125
0.605
150
0.587
175
0.573
200
0.560
250
0.540
300
0.525
350
0.514
400
0.504
500
0.489
600
0.481
700
0.473
900
0.463
1100
0.458
A e
0.409
oo
0.0013 Nb
Run 1
Vol
DNA
Absd
0
0.734
15.0
0.699
30.0**
0.672
45.0**
0.648
60.0**
0.627
80.0**
0.604
100
0.582
125
0.561
150
0.545
175
0.530
200
0.518
250
0.498
300
0.483
350
0.474
400
0.465
450
0.459
550
0.450
A e
0.418
Run 2
Vol
DNA
Absd
0
0.734
15.0**
0.704
30.0**
0.677
45.0**
0.653
60.0**
0.631
75.0**
0.613
90.0
0.597
110
0.577
135
0.557
160
0.539
185
0.528
225
0.509
275
0.491
325
0.479
375
0.469
425
0.461
475
0.457
575
0.448
775
0.435
1075
0.427
A e
0.399
223

Table 9 (Continued)
0.00063Nb
Run 1 Run 2
Vol
Absd
Vol
Absd
DNAC
DNAC
0
0.718
0
0.733
10.0
0.692
20.0
0.687
25.0
0.660**
40.0**
0.649
40.0
0.632**
60.0**
0.616
55.0
0.608
80.0
0.590
70.0
0.587
100
0.569
85.0
0.569
125
0.547
100
0.554
150
0.528
125
0.533
175
0.513
150
0.517
200
0.502
175
0.503
250
0.485
200
0.492
300
0.472
250
0.475
350
0.464
300
0.464
400
0.458
350
0.455
500
0.447
400
0.448
600
0.442
500
0.439
A e
OO
0.416
600
0.434
800
0.426
1000
0.426
0.393
224

Table 9 (Continued)
aT = 25.0 + 0.1°C, pH = 5.90.
Initial total concentration of supporting electrolyte.
'Total volume of solution of calf thynu
moles of DNA phosphate per microliter,
c , — 9
Total volume of solution of calf thymus DNA added (yl). Titrant contained 8.246^ x 10
Total absorbance of sample at 365 nm, 4.00 cm pathlength. Data points marked (**)
used to compute Kg.
^Calculated absorbance of solution if all 3-aminoacridinium cation is bound. Based on
molar absorptivity of bound drug of 8494M“^cm-l.
225

Table 10 . Absorptiometric titration of 3-aminoacridinium cation with DNA, calcium
salt. Calcium acetate as supporting electrolyte a
0.025Nb 0.010Nb 0.0050Nb
Run
1
Run
1
Run
1
Run
2
Vol
Abs^
Vol
,, d
Abs
Vol
_. d
Abs
Vol
—. d
Abs
DNAC
DNAC
DNA
c
DNAC
0
0.694
0
0.717
0
0.735
0
0.704
30.0
0.669
20.0
0.696
10.
0
0.723
20.0
0.673
60.0**
0.652
40.0**
0.679
20.
0
0.713
40.0**
0.651
90.0**
0.637
60.0**
0.660
50.
0**
0.681
60.0**
0.634
140**
0.614
80.0**
0.647
80.
0**
0.652
80.0**
0.616
190**
0.595
110**
0.627
110**
0.629
110**
0.595
240**
0.578
140
0.608
140**
0.610
140
0.575
290
0.564
170
0.593
170**
0.594
180
0.555
365
0.546
210
0.576
200
0.577
220
0.537
450
0.527
250
0.559
250
0.557
270
0.519
550
0.512
300
0.543
300
0.542
320
0.506
650
0.498
350
0.530
350
0.524
400
0.488
800
0.482
400
0.520
400
0.515
500
0.473
950
0.470
500
0.501
500
0.496
600
0.460
1250
0.455
600
0.485
600
0.483
800
0.446
1450
0.449
700
0.478
700
0.472
1100
0.432
0.365
900
0.460
900
0.459
>
8
CD
0.381
1100
0.450
1200
0.447
1475
0.436
1600
0.437
A e
00
0.376
A e
oo
0.381
226

Table 10 (Continued)
0.0025Nb
Run 1
Vol
Abs
DNA
0
0.728
20.0
0.695
40.0**
0.671
65.0**
0.643
90.0**
0.618
115**
0.598
140**
0.580
175
0.559
210
0.541
260
0.521
310
0.507
385
0.491
460
0.478
560
0.464
760
0.458
860
0.447
1160
0.437
1360
0.431
A e
0.385
Run
2
Vol
Absd
DNAC
0
0.712
20.0
0.682
40.0**
0.657
60.0**
0.634
80.0**
0.614
100**
0.598
130**
0.577
160
0.556
190
0.541
240
0.520
290
0.503
340
0.490
440
0.471
540
0.460
690
0.449
890
0.439
1190
0.429
1390
0.427
A e
00
0.376
0.0013Nb
Run
1
Run
2
Vol
c
Absd
Vol
c
Absd
DNA
DNA
0
0.708
0
0.729
15.0
0.665
20.
0
0.698
30.0**
0.646
40.
o**
0.672
45.0**
0.630
60.
0**
0.648
60.0**
0.609
80.
0**
0.630
80.0**
0.600
100**
0.612
100
0.581
125**
0.592
125
0.565
150
0.577
150
0.549
175
0.563
200
0.528
200
0.550
500
0.459
250
0.529
600
0.449
300
0.513
800
0.436
350
0.499
1000
0.429
400
0.489
1330
0.420
500
0.476
A e
oo
0.376
600
0.465
700
0.456
900
0.444
1200
0.437
A e
0.391
227

Table 10 (Continued)
0.00063Nb
Run 1
Vol
-. d
Abs
DNAC
0
0.733
10.0
0.714
20.0
0.697
35.0**
0.675
50.0**
0.653
65.0**
0.634
85.0**
0.612
105**
0.593
125
0.578
150
0.559
175
0.545
475
0.464
575
0.453
675
0.447
975
0.435
1375
0.428
A e
0.388
Run 2
Vol Absu
DNAC
0
0.740
15.0
0.709
30.0**
0.684
45.0**
0.663
60.0**
0.643
80.0**
0.612
100**
0.600
125
0.580
150
0.562
175
0.547
225
0.522
275
0.503
325
0.491
425
0.471
525
0.461
725
0.448
925
0.440
A e
0.406
aT = 25.0 + 0.1°C
/
pH = 5.90
0.000 010Nb
Run 1
Vol
DNA
Abs^
0
0.667
10.0
0.632
20.0**
0.610
30.0**
0.587
40.0**
0.566
55.0**
0.539
70.0
0.520
85.0
0.500
105
0.483
125
0.469
145
0.458
170
0.445
220
0.431
270
0.419
370
0.409
590
0.398
A e
0.379
228

Table 10 (Continued)
Initial total concentration of supporting electrolyte,
'Total volume of solution of calf thymi
moles of DNA phosphate per microliter,
c — 9
Total volume of solution of calf thymus DNA added (ul). Titrant contained 7.935 x 10
Total absorbance of sample at 365 nm, 4.00 cm pathlength. Data points marked (**)
used to compute Kg.
'Calculated absorbance of solution if all 3-aminoacridinium cation is bound. Based on
molar absorptivity of bound drug of 84 94M“-*-cm--'-.
229

Table 11 . Absorptiometric titration of 3-aminoacridinium cation with DNA, strontium
salt. Strontium acetate as supporting electrolyte a
0.025Nb
0.010Nb
Run 1
Run 2
Run
1
Run
2
Vol
DNAC
Absd
Vol
DNAC
Absd
Vol
DNAC
Absd
Vol
DNAC
Absd
0
0.668
0
0.660
0
0.678
0
0.677
25.0
0.651
25.0
0.639
20.0
0.659
20.0
0.654
50.0
0.636
50.0
0.624
40.0
0.640
40.0
0.636
75.0
0.620
75.0
0.609
60.0**
0.623
60.0**
0.619
100**
0.604
100**
0.596
80.0**
0.608
80.0**
0.607
150**
0.583
150**
0.574
100**
0.593
100**
0.594
200**
0.562
200**
0.555
125**
0.577
125**
0.580
250**
0.547
250**
0.537
150**
0.564
150**
0.564
300
0.532
300
0.523
200
0.542
200
0.540
400
0.509
400
0.502
250
0.524
250
0.523
500
0.494
500
0.484
300
0.508
300
0.507
600
0.481
600
0.473
350
0.497
350
0.494
800
0.463
700
0.465
400
0.485
400
0.483
1000
0.450
900
0.450
500
0.469
500
0.468
1300
0.440
1100
0.440
600
0.457
600
0.457
1525
0.433
1400
0.433
800
0.444
800
0.442
A S
00
0.349
1725
A e
oo
0.430
0.339
1100
A e
OO
0.430
0.367
1000
A e
OO
0.430
0.370
230

Table 11 (Continued)
0.0050Nb
Run 1 Run 2
Vol Abs^ Vol Abs^
DNAC DNAC
0
0.704
15.0
0.657
30.0**
0.639
50.0**
0.617
75.0**
0.597
100**
0.578
125
0.559
150
0.544
175
0.529
200
0.518
250
0.497
300
0.481
350
0.470
400
0.461
500
0.449
700
0.432
900
0.423
1200
0.417
A e
OO
0.380
0
0.648
20.0
0.629
40.0**
0.602
60.0**
0.583
80.0**
0.567
100**
0.552
125**
0.536
150
0.520
175
0.509
200
0.497
250
0.479
300
0.464
350
0.453
400
0.445
500
0.433
600
0.425
700
0.419
1000
0.408
A e
0.354
Run 3
Vol
DNAC
.. d
Abs
0
0.633
20.0
0.604
40.0**
0.583
60.0**
0.562
80.0**
0.546
100**
0.529
125**
0.514
150
0.497
175
0.488
200
0.477
250
0.460
300
0.447
350
0.436
400
0.429
500
0.415
700
0.400
900
0.395
A 6
0.349
oo
231

Table 11 (Continued)
0.0025Nb
Run 1
Run 2
Vol Absd
c
Vol Abs^
_ r.
0
0.660
0
0.632
20.0
0.632
20.0
0.605
40.0**
0.607
40.0**
0.579
60.0**
0.583
50.0**
0.567
80.0**
0.551
60.0**
0.554
100*
0.545
80.0**
0.532
125
0.526
100
0.517
150
0.508
125
0.494
175
0.495
150
0.479
200
0.484
175
0.465
250
0.467
200
0.453
300
0.454
250
0.438
400
0.437
300
0.424
500
0.424
350
0.417
700
0.414
400
0.413
1000
0.407
500
0.400
A e
OO
0.361
700
0.391
900
0.387
A e
0.349
oo
0.0013Nb
Run 1
Vol
Abs
DNA
0
0.701
20.0
0.662
40.0**
0.629
60.0**
0.600
80.0**
0.578
100
0.557
125
0.536
150
0.519
175
0.504
200
0.494
250
0.479
300
0.468
400
0.452
500
0.443
800
0.432
A 6
OO
0.390
Run 2
Vol
Absd
DNA
0
0.577
20.0
0.539
30.0**
0.523
40.0**
0.508
50.0**
0.498
60.0**
0.487
80.0
0.467
100
0.450
125
0.433
150
0.420
175
0.411
200
0.402
250
0.390
300
0.380
350
0.376
400
0.373
500
0.368
A e
0.331
232

Table 11 (Continued)
aT = 25.0 + 0.1°C, pH = 5.90.
b
Initial total concentration of supporting electrolyte.
C — Q
Total volume of solution of calf thymus DNA added (yl) . Titrant contained 7.906^ x 10
moles of DNA phosphate per microliter.
dTotal absorbance of sample at 365 nm, 4.00 cm pathlength. Data points marked (**) used
to compute Kg.
0
Calculated absorbance of solution if all 3-aminoacridinium cation is bound. Based on
molar absorptivity of bound drug of 8 4 94M--^cm--*-.

Table 12 . Absorptiometric titration of 3-aminoacridinium cation with DNA, barium salt.
Barium acetate as supporting electrolyte a
0.010Nb
0.0050N
b
Run
1
Run
2
Vol
DNAC
Absd
Vol
DNAC
Abs^
0
0.734
0
0.739
20.0
0.707
20.0
0.713
40.0
0.683
40.0
0.688
60.0**
0.664
60.0**
0.669
80.0**
0.647
80.0**
0.650
100**
0.630
100**
0.636
125**
0.613
125**
0.617
150**
0.596
150**
0.605
200**
0.568
175
0.592
250
0.548
200
0.582
300
0.531
250
0.560
350
0.519
300
0.539
400
0.508
350
0.528
450
0.497
400
0.517
550
0.484
500
0.499
650
0.472
600
0.489
750
0.465
700
0.479
950
0.454
900
0.470
1150
0.449
1100
0.465
1350
0.444
1300
0.464
1650
0.443
1700
0.463
A e
00
0.379
QJ
8
<
0.380
Run 1
Run 2
Vol Absa
ji
Vol Absa
c
0
0.740
0
0.734
15.0
0.718
20.0
0.701
30.0
0.695
40.0**
0.673
50.0**
0.672
60.0**
0.649
70.0**
0.646
80.0**
0.627
90.0**
0.624
100**
0.608
110**
0.607
125
0.586
130
0.591
150
0.570
150
0.579
175
0.555
175
0.564
200
0.542
200
0.551
250
0.520
250
0.531
300
0.505
300
0.513
350
0.492
350
0.501
400
0.480
400
0.493
500
0.469
500
0.477
600
0.460
600
0.466
700
0.453
700
0.459
900
0.447
800
0.455
1200
0.441
1000
0.452
A e
00
0.394
1300
0.448
A e
0.394
234

Table 12 (Continued)
0.0025Nb
Run 1 Run 2
Vol Abs^ Vol Abs^
DNAC DNAC
0
0.718
15.0
0.688
30.0**
0.663
45.0**
0.641
60.0**
0.620
80.0**
0.599
100**
0.578
120
0.562
140
0.547
165
0.532
190
0.517
225
0.502
275
0.486
325
0.474
375
0.463
425
0.456
500
0.449
600
0.441
750
0.435
950
0.430
1250
0.429
A e
0.384
CO
0
0.733
20.0
0.698
40.0**
0.659
60.0**
0.632
80.0**
0.609
100**
0.588
125
0.567
150
0.549
175
0.534
200
0.521
225
0.512
275
0.494
325
0.483
375
0.473
425
0.465
525
0.455
625
0.449
725
0.445
1025
0.438
1350
0.437
A e
0.388
CO
0.0013Nb
Run 1
Vol
Abs^
DNA
0
0.729
15.0
0.693
30.0**
0.661
45.0**
0.633
60.0**
0.608
75.0
0.587
90.0
0.569
110
0.550
130
0.535
155
0.517
180
0.504
205
0.494
255
0.478
305
0.465
355
0.456
455
0.446
555
0.442
855
0.437
A e
0.404
oo
Run 2
.q
Vol Absa
DNA
0
0.735
10.0
0.717
20.0**
0.692
30.0**
0.673
40.0**
0.656
60.0**
0.623
80.0**
0.597
100
0.576
125
0.552
150
0.535
175
0.520
200
0.509
250
0.491
300
0.479
350
0.471
400
0.468
500
0.460
650
0.457
950
0.450
1150
0.450
A e
0.396
235

Table 12 (Continued)
0.00063N
b
Run 1 Run 2
Vol Abs^ Vol Abs^
DNA DNAC
0
0.726
10.0
0.692
20.0**
0.670
30.0**
0.650
45.0**
0.621
60.0
0.598
75.0
0.576
90.0
0.559
110
0.538
130
0.520
155
0.503
180
0.493
230
0.476
280
0.463
330
0.455
430
0.447
530
0.442
730
0.438
A e
0.407
oo
0
0.728
10.0
0.701
20.0**
0.676
30.0**
0.653
40.0**
0.633
50.0**
0.617
65.0
0.591
80.0
0.570
100
0.549
120
0.531
140
0.516
160
0.502
190
0.488
220
0.476
250
0.467
300
0.458
350
0.451
450
0.443
550
0.440
850
0.438
A e
0.404
oo
1.0 X 10~5Nb
Run
1
Run
2
Vol
Absd
Vol .
DNAC
—. d
Abs
DNAC
0
0.680
0
0.698
10.0**
0.644
10.0
0.674
20.0**
0.610
20.0**
0.648
30.0**
0.582
30.0**
0.626
40.0**
0.558
40.0**
0.607
50.0
0.535
55.0**
0.581
65.0
0.517
70.0
0.559
80.0
0.491
85.0
0.539
95.0
0.474
105
0.516
115
0.450
125
0.499
135
0.437
145
0.488
160
0.425
170
0.475
210
0.413
195
0.465
260
0.404
245
0.450
360
0.400
295
0.441
520
0.400
395
0.429
A e
OO
0.389
595
0.405
A e
0.396
OO
236

Table 12 (Continued)
1.0 x 10 5Nb
Run 3
Vol
Absd
DNAC
0
0.683
10.0**
0.642
20.0**
0.610
30.0**
0.583
40.0**
0.557
50.0
0.539
65.0
0.510
80.0
0.489
95.0
0.474
115
0.457
135
0.443
160
0.434
210
0.420
260
0.417
360
0.411
560
0.409
A e
00
0.389
aT = 25.0 + 0.1°C, pH = 5.90.
^Initial total concentration of supporting electrolyte.
237

Table 12 (Continued)
c — y
Total volume of solution of calf thymus DNA added (yl) . Titrant contained 8.18£ x 10
moles of DNA phosphate per microliter.
aTotal absorbance of sample at 365 nm, 4.00 cm pathlength. Data points marked (**) used
to compute Kg.
Calculated absorbance of solution if all 3-aminoacridinium cation is bound. Based on
molar absorptivity of bound drug of 8494M-1cm~ .

Table 13. Absorptiometric titration of 3-aminoacridinium cation with DNA, potassium
salt, at 15°C.a
Run
1
Run
2
Run
3
Volh
DNA
AbsC
Volh
DNA
AbsC
Volh
DNA
AbsC
0
0.717
0
0.653
0
0.668
7.5
0.689
7.5**
0.622
7.5
0.640
15.0**
0.654
15.0**
0.596
15.0**
0.615
22.5**
0.628
22.5**
0.573
22.5**
0.590
30.0**
0.601
30.0**
0.549
30.0**
0.568
40.0**
0.573
40.0
0.519
40.0
0.539
50.0
0.544
50.0
0.492
50.0
0.511
60.0*
0.520
60.0*
0.470
60.0*
0.489
70.0*
0.500
70.0*
0.451
70.0*
0.468
80.0*
0.482
80.0*
0.437
80.0*
0.452
95.0*
0.463
95.0
0.428
95.0*
0.429
110
0.450
110
0.407
110
0.414
130
0.441
130
0.398
130
0.401
180
0.431
230
0.393
230
0.395
280
A d
0.429
0.
A d
00
0.
A d
OO
0.
3.
Potassium dihydrogen phosphate as supporting electrolyte, 0.010M initial total
concentration of supporting electrolyte: 0.010M, pH = 5.90, temperature,15.0 + 0.2°C.
^Total volume of solution of calf thymus DNA added ( y 1) . Titrant contained 8.132^ x 10
moles of DNA phosphate per microliter.

Table 13 (Continued)
0
Total absorbance of sample, 4.00 cm pathlength. Data marked (**) used to compute K ,
those marked (*) were used to compute K^.
Calculated absorbance of solution if all 3-aminoacridinium cation is bound,
molar absorptivity of bound species of
8494M â– 'â– cm ,
Based on

Table 14 . Absorptiometric titration of 3-aminoacridinium cation with DNA, potassium
salt, at 35°C a
Run
1
Run
2
Volb
AbsC
Volb
AbsC
DNA
DNA
0
0.683
0
0.671
7.5
0.658
7.5
0.638
15.0**
0.636
15.0**
0.621
22.5**
0.617
22.5**
0.600
30.0**
0.593
30.0**
0.583
40.0**
0.570
40.0
0.561
50.0
0.547
50.0
0.540
60.0
0.528
60.0
0.521
70.0
0.509
70.0
0.502
80.0
0.495
80.0*
0.488
95.0*
0.477
95.0*
0.467
110*
0.459
110*
0.450
130*
0.443
130*
0.430
155*
0.432
155*
0.419
205
0.427
205
0.405
305
0.420
305
0.403
405
0.420
A d
oo
0.390
A d
0. 393
Potassium dihydrogen phosphate as supporting electrolyte, 0.010M initial total
concentration of supporting electrolyte, pH 5.90, temperature 35.0 + 0.2°C.
j-
Total volume of solution of calf thymus DNA added (yl) . Titrant contained 8.132^ x 10
moles of DNA phosphate per microliter.
241

Table 14 (Continued)
CTotal absorbance of sample, 4.00 cm pathlength. Data marked (**)
those marked (*) were used to compute K^.
Calculated absorbance of solution if all 3-aminoacridinium cation
molar absorptivity of bound species of 8494M“-1-cm--*-.
used to compute Kg,
is bound.
Based on

Table 15 . Absorptiometric titration of 7-aminoquinolinium cation with DNA, potassium
salt, at 15°C a
Run
1
Run
2
Volh
DNA*3
AbsC
Vol
DNA*3
AbsC
0
0.341
0
0.331
50.0**
0.323
25.0
0.325
75.0**
0.314
50.0**
0.315
100**
0.309
75.0**
0.312
150
0.298
125
0.299
200*
0.289
175
0.290
250*
0.281
225*
0.282
300*
0.275
275*
0.273
400
0.264
375*
0.263
500
0.255
475
0.252
700
0.247
675
0.241
1000
0.233
875
0.231
A d
00
0.240
A d
00
0. 236
£
Potassium dihydrogen phosphate as supporting electrolyte, 0.010M initial total
concentration, pH = 5.90, temperature, 15.0 + 0.2°C.
Id “
Total volume of solution of calf thymus DNA added. Titrant contained 8.163^ x 10
moles DNA phosphate per microliter.
0
Total absorbance of sample, 4.00 cm pathlength. Data marked (**) used to compute
Kg, those marked (*) were used to compute .
243

Table 15 (Continued)
aCalculated absorbance of solution if all 7-aminoquinolinium originally involution
were bound. Based on a molar absorptivity if bound species of 5332M-Xcm

Table 16
Absorptioinetric titration of 7-aminoquinolinium cation with DNA, potassium
salt, at 25°C a
Run
1
Run
2
Volh
DNA
AbsC
Volh
DNA
AbsC
0
0.362
0
0.360
50.0
0.353
50.0
0.352
75.0**
0.344
75.0**
0.346
100**
0.338
100**
0.337
150**
0.327
150**
0.324
200
0.318
200
0.316
250*
0.311
250*
0.303
300*
0.302
300*
0.297
400*
0.294
400*
0.285
500*
0.286
500*
0.287
700*
0.271
600*
0.278
1000
0.259
900
0.264
1400
0.243
1200
0.255
A d
0.244
1680
0.238
CO
A d
0.235
aPotassium dihydrogen phosphate as supporting electrolyte, 0.010M initial total
concentration, pH = 5.90, temperature, 25.0 + 0.1°C.
^Total volume of solution of calf thymus DNA added (yl) . Titrant contained 8.16.3 x 10
moles of DNA phosphate per microliter.
245

Table 16(Continued)
cTotal absorbance of sample, 4.00 cm pathlength. Data marked (**) used to compute
Kg, those marked (*) were used to compute K-^.
calculated absorbance of solution if all 7-aminoquinolinium originally in solution
were bound. Based on a molar absorptivity if bound species of 5332M_^-cm_1

Table 17 . Absorptiometric titration of 7-aminoquinolinium cation with DNA, potassium
salt, at 35°C a
Run
1
Run 2
Run
3
Run
4
Volb
AbsC
Volb
AbsC
Volb
AbsC
Volh
AbsC
DNA
DNA
DNA
DNA
0
0.342
0
0.326
0
0.326
0
0.316
50.0
0.329
50.0
0.317
50.0
0.317
50.0
0.307
75.0**
0.324
100**
0.309
75.0**
0.311
100
0.297
100**
0.321
150**
0.300
100**
0.308
150
0.289
150**
0.309
200**
0.293
150**
0.297
200
0.279
200
0.307
250
0.285
200
0.289
250
0.273
250
0.300
300
0.280
250
0.284
300
0.269
300*
0.293
400*
0.268
300
0.278
400
0.261
400*
0.282
600*
0.261
400*
0.271
600
0.253
500*
0.276
900*
0.242
500*
0.265
800
0.243
700*
0.259
1300
0.229
700*
0.251
1100
0.229
900*
0.251
1800
0.215
1000*
0.243
1380
0.225
1200
0.242
2400
0.195
1400
0.226
1880
0.208
1650
0.227
3000
0.185
1900
0.215
A d
oo
0.224
A d
CO
0. 159
2580
0.196
A d
oo
0.203
£
Potassium dihydrogen phosphate as supporting electrolyte, 0.010M initial total
concentration, pH = 5.90, temperature, 35.0 + 0.2°C.
Total volume of solution of calf thymus DNA added (yl) . Titrant contained 8.163 x 10
moles of DNA phosphate per microliter.
247

Table 17 (Continued)
Total absorbance of sample, 4.00 cm pathlength. Data marked (**) used to compute
K , those marked (*) were used to compute K..
S
Calculated absorbance of solution if all 7-aminoquinolinium originally in solution
were bound. Based on a molar absorptivity if bound species of 5332M-^-cm-l.

Table 18. Calculated values of [BH],
between 3-aminoacridinium
Vol
DNA
(Ul)
Abs
[BH]
[BHP]
M
M
5
5
x 10
x 10
0
0.763
1.350
0
15.0
0.738
1.251
.097
30.0**
0.716
1.164
.183
45.0**
0.692
1.069
.275
60.0**
0.672
.990
.353
75.0**
0.654
.919
.422
90.0**
0.635
.844
.494
110**
0.613
.757
. 578
130
0.596
.691
.642
150
0.582
.635
. 695
170
0.570
.588
.740
195
0.552
.518
.807
225*
0.540
.471
.850
265*
0.521
. 397
.919
315*
0.507
.342
.967
365*
0.497
.303
1.000
[BHP], "K ", and related data for the reaction
cation and DNA in 0.15M CsI^PC^3
ii k ii ^
, .
K
,—,
i i
I
cm
K
s
CM
-2
K
CQ
-1
K
M
CQ
1—1
M
m
x 10-8
CN
1
\
Pm
x 10-4
i
CN
K
\
\
P3
i—i
r—i
-P
4->
—
Pi
Pi
CP
*—11
l 1
0
—'
r—1
CP
CP
0
0
iH
1—1
1
i
37.5
5.34
-1
.11
2.31
5.47
18.0
5.03
-
.804
2.30
5.17
13.3
4.86
-
.589
2.52
4.99
10.0
4.73
-
.448
2.59
4.86
8.05
4.62
-
. 338
2.65
4.76
7.05
4.54
-
.232
2.81
4.68
6.00
4.45
-
.117
2.97
4.59
5.07
4.37
-
.031
3.03
4.51
4.33
4.30
+
.039
3.05
4.45
3.77
4.24
+
.099
3.07
4.39
3.48
4.17
+
.192
3.29
4.33
2.93
4.11
+
.257
3.27
4.26
2.64
4.03
+
.365
3.54
4.18
2.21
3.95
+
.452
3.60
4.10
1.87
3.88
+
.519
3.60
4.04

Vol
Abs
[BH]
[BHP]
"Kt"
DNA
M
CL
M
_ 9
(ui)
M 2
x 10
x 10
o
i—i
X
PM
X
m
CM
I
CM
\
PM
CP
o
—I
I
X
m
PM
w
m
CP
o
K
M
s
-1
x 10
-4
x
X
ffl
co
X
CP
o
t—!
I
415*
0.488
.268
1.029
1.66
3.82
+ .585
3.66
3.98
515*
0.474
.213
1.071
1.38
3.72
+ .701
3.84
3.88
715*
0.460
.158
1.102
.973
3.57
+ .843
3.83
3.74
1015
0.454
.133
1.093
.577
3.42
+ .916'
3.22
3.59
aRefer to Table 7 for reaction conditions.
bNot a valid association constant as the intercalative process was not seen under these
conditions. Included for illustrative purposes only. See text.
250

Table 19
. Calculated values of [BH], [BHP], Ks, Kj, and related data for the reaction
between 3-aminoacridinium cation and DNA in 0.0025M Csi^PO^a
Vol
DNA
(yi)
Abs
[BH]
[BHP]
KI
M
M
-2
5
, „5
M
x 10
x 10
x 10
r—i
1—1
K
i—i
tc
S
04
K
m
ffi
m
M ^
CQ
■—>
\
ivi
*—■*
(N
r—1
1
|
Ck
n 5
ro
CM
K
X 10
\
\
0Q
4—1
r—«
•—1
â– P
-P
—-
04
ft
Cn
1—1
U—J
0
-
r—1
d
0
0
I—1
0
0.752
1.331
0
-
-
-
-
-
5.0
0.728
1.227
.103
793
6.99
-1.07
2.04
6.38
10.0
0.701
1.110
.219
-
-
-
-
-
15.0**
0.681
1.024
. 305
191
6.40
- .526
2.33
5.89
20.0**
0.666
. 959
.370
23.9
5.90
- .414
1.85
5.68
25.0**
0.649
.885
.442
12.8
5.70
- .301
1.79
5.55
30.0**
0.631
.808
.519
9.58
5.59
- .192
1.86
5.46
37.5
0.605
.695
.631
6.99
5.44
- .042
2.01
5.35
45.0
0.581
.592
.733
5.42
5.32
+ .093
2.20
5.25
52.5*
0.561
.506
.818
4.07
5.20
+ .209
2.33
5.16
60.0*
0.539
.411
.912
3.79
5.12
+ .346
2.72
5.09
70.0*
0.515
. 308
1.013
3.35
5.00
+ .517
3.29
5.00
80.0*
0.497
.231
1.089
2.94
4.90
+ .673
3.91
4.92
90.0*
0.481
.163
1.157
2.93
4.81
+ .852
4.99
4.85
100
0.467
.103
1.215
3.39
4.73
+1.07
7.17
4.78
110
0.462
. 0814
1.235
2.99
4.65
+1.18
7.94
4.72
160
0.456
.0555
1.255
1.21
4.36
+1.35
6.84
4.48
210
0.453
.0426
1.261
0.722
4.19
+1.47
6.31
4.33
251

a
Refer to Table 7 for reaction conditions.

Table 20. Calculated values of [BH], [BHP], Ks, "Kj", and related data for the reaction
between 3-aminoacridinium cation and DNA in 0.010N MgiC^CCH^^3
Vol
DNA
(yi)
Abs
t)
[BH]
[BHP]
KI
r—i
r—i
K
r y
a
b
a
N
N
-2
a
m
~1
a
r
N
CQ
•—•
N
a
x 10
x 105
x 10-7
(N
1
\
a
x 10“4
i
ro
a
\
a
f—1
r—i
-P
-p
—-
a
a
Cn
■—i
o
•—
■—
i—i
Cn
o
O «H
fH |
I
0
0.742
1.313
0
-
-
-
-
-
25.0
0.716
1.215
.095
Ill
5.08
-1
.11
1.32
5.23
50.0
0.693
1.128
.178
55.0
4.77
-
.801
1.33
4.92
75.0**
0.673
1.053
.251
36.0
4.59
-
. 624
1.32
4.75
100**
0.657
. 992
.308
25.8
4.46
-
.508
1.28
4.62
150**
0.625
. 872
.422
17.5
4.28
-
.316
1.33
4.44
200**
0.600
.778
.509
13.1
4.15
-
.184
1.34
4.31
250
0.581
.707
.575
10.2
4.05
-
. 090
1.33
4.21
300
0.565
. 646
. 629
8.40
3.97
-
.012
1.32
4.13
350
0.551
.594
.675
7.16
3.90
+
. 056
1.32
4.06
400
0.541
.556
.707
6.09
3.84
+
.105
1.29
4.01
500
0.520
.477
.773
4.95
3.74
+
.210
1.32
3.91
600*
0.506
.424
.814
4.07
3.66
+
.283
1.30
3.83
700*
0.495
.383
. 844
3.45
3.60
+
.343
1.29
3.77
800*
0.485
.346
. 871
3.03
3.54
+
.401
1.29
3.71
900*
0.478
.319
.886
2.67
3.49
+
.444
1.27
3.66
253

Vol
Abs
[BH]
[BHP]
R b
r—“i
DNA
N
c.
N
R
1
Om
K
CP
(yl)
N 2
â– >
o
i—t
X
x 10
x 10"7
CM
CM
Pm
Cn
O
r—i
K
s
Pm
CQ
-1
SC
N
PC
\
-4
1—1
PM
x 10
1
ro
\
m
i—i
-P
—■
Pm
cn
1—1
0
'—
1—1
O
■—I
I
1100*
0.467
.277
.906
2.15
3.41
+ .515
1.24
3.58
1300*
0.459
.246
.916
1.79
3.34
+ .571
1.21
3.51
1650
0.449
.207
.920
1.39
3.25
+ .647
1.17
3.42
Refer to Table 9 for reaction conditions.
uNot a valid association constant as the intercalative process was not seen under these
conditions. Included for illustrative purposes only. See text.

Table 21. Calculated values of [BH], [BHP], Ks, "Kj", and related data for the^reaction
between 3-aminoacridinium cation and DNA in 6.3 x 10 4 N Mg(C^CCH^)^
Vol
Abs
[BH]
[BHP]
KIb
f—i
CM
K
K
s
DNA
(yi)
N
R
N
5
N-
CQ
U—.4
m
\
N-1
x 10
x 10
x 10-8
(N
1
x 10"4
m 1 1
-p
cu
o
tjl
o
I—I
I
0 0.718
10.0
0.692
1.271
0
198
5.68
-1
.06
5.01
5.76
25.0
0.660
1.168
.101
65.1
5.24
-
. 664
4.71
5.34
40.0
0.632
1.042
.226
37.8
5.01
-
. 449
4.72
5.12
55.0
0.608
.931
.334
26.0
4.85
-
.292
4.74
4.97
70.0
0.587
. 837
.427
19.6
4.73
-
.172
4.80
4.85
85.0
0.569
.750
.508
15.6
4.63
-
. 074
4.85
4.76
100
0.554
. 683
.577
12.8
4.55
+
.007
4.86
4.68
125
0.533
. 624
.634
9.80
4.44
+
.119
4.91
4.57
150
0.517
.542
.713
7.80
4.34
+
.208
4.91
4.48
175
0.503
.479
.773
6.56
4.26
+
.289
4.98
4.41
200
0.492
.424
.825
5.62
4.20
+
.356
5.02
4.34
250
0.475
. 381
.865
4.37
4.09
+
.469
5.09
4.24
300
0.464
.314
.926
3.49
4.00
+
.551
5.04
4.15
350
0.455
.271
.963
2.94
3.92
+
.624
5.07
4.08
400
0.448
.236
.992
2.54
3.86
+
.687
5.09
4.02

\ Lun Lj.ii ut;u;
Vol
Abs
[BH]
[BHP]
Kib
DNA
N
N
c
(yi)
N-2
x 10
x 10
x 10
CM
X
£Q
I
fN
\
,_+J
X
cn
O
i—i
i
X
s
i i
X
m
N_1
X
m
\
-4
•—■
r—»
x 10
i
0^
n
X
\
X
r—t
U—J
4J
—'
X
'1
0
■—'
rH
Cn
o
! 1
I
500
0.439
.208
1.014
1.94
3.76
+
.778
4.98
3.92
600
0.434
.173
1.037
1.51
3.67
+
.835
4.71
3.84
800
0.427
.153
1.046
1.04
3.55
+
.926
4.37
3.71
1000
0.426
.125
1.052
0.69
3.45
+
.940
3.64
3.62
aRefer to Table 9 for reaction conditions.
Not a valid association constant as the intercalative process was not
conditions. Included for illustrative purposes only. See text.
seen under these

257
Table 22. Percent total 3-aminoacridinium bound to DNA
in the presence of various electrolytes after
addition of excess DNAa
Electrolyte0 Ionic strength P/Da % of total
(I)c drug bound
0.15M CsH2P04
0.14
65
90
0.10M C HoP0.
S 2 4
0.094
55
94
0.010M CsH2P04
0.010
17
97
0.0025M C H„PO.
s 2 4
0.0028
14
97
0.10M LiH2P04
0.094
57
89
0.010M LiH2P04
0.010
15
99
6.3 x 10“4M LiH2P04
0.0012
11
98
0.010N Ba(02CCH3)
0.017
103
77
6.3 x 10-4N Ba(02CCH3)2
0.0027
47
90
0.010N Mg(02CCH3)
0.017
104
70
6.3 x 10"4N Mg(02CCH3)2
0.0034
65
90
apH 5.90, T = 25.0°C.
^Concentrations before any DNA added.
c . .
In equilvalents per liter for the solution containing the
volume of DNA necessary for the P/D ratio listed.
Total moles of DNA phosphate per total moles of 3-AA.

258
Table 23. Initial and final slopes of plots of log([BHP]/
[BH]) vs. log([Pt/m]-q[BHP]) in various concen¬
trations of CsH2PO^ and MgiC^CCH^^
Electrolyte
Initial
slope
Final
slope
b
_c
b
T _c
I
II
i
II
0.15M C H„PO.
1.14
1.16
1.14
1.16
s 2 4
0.010M C H„PO.
1.07
1.23
1.68
1.88
s 2 4
0.0025M C H„PO.
1.10
1.14
1.71
2.16
s 2 4
0.010N Mg(02CCH3)2
0.96
0.97
0.96
0.97
6.3 x 10"4N Mg(02CCH3)2
0.98
1.0
0.98
1.0
aTitrations done at 25.0°C, pH 5.90.
blog( [BHP]/[BH]) vs. log([P /2]-2 [BHP]).
Clog ( [BHP] / [BH] ) vs. log ( [P^] - [BHP] ) .

Table 24. Apparent equilibrium association constants for surface, Ks, and inter-
calative, Kj, binding of 3-aminoacridinium
electrolyte.a
to DNA,
LiH2P0^ as supporting
I
(ave)
K M_1 o(n)b I
S _4 _4 (ave)
x 10 x 10
Kj M-2
x 10~9
a (n) b
x 10~9
.099
2.66
0.112 (4)
.097
0.233
0.044(4)
.10
3.03
0.181 (4)
.098
0.299
0.066 (4)
2.85
0.266
.050
4.28
0.333(4)
.049
0.989
0.118 (4)
050
3.54
0.325(3)
. 049
0.565
0.116 (4)
3.91
0.777
. 025
5.40
0.309(3)
. 025
2.41
0.366(4)
. 010
9.01
0.888(3)
.010
8.35
.661(3)
. 010
11.0
0.938(3)
. 010
8.89
.974 (5)
10.0
8.62
. 0050
13.5
0.404 (3)
.0051
15.8
3.73(4)
. 0050
21.7
0.071(2)
.0051
34.6
10.8(3)
.0050
12.4
0.416 (3)
.0051
13.8
1.06(4)
15.9
14.8
.0025
27.1
0.990 (2)
.0026
78.2
29.3 (2)
. 0025
16.0
1.56 (4)
. 0026
22.8
2.81(4)
21.6
50.5
259

I
(ave)
K M-1
s
x 10-4
o (n) b
x 10-4
I
(ave)
-2
KI M
x 10-9
o (n)b
x 10"9
.00066
16.4
1.53(3)
.0074
53.7
4.18(4)
. 00066
22.7
1.50(3)
.0073
80.5
7.00(3)
19.6
67.1
aConstants based on molar concentrations of species, T = 25.0°C,
pH 5.90.
fa • 4-V,
a is the
standard deviation from the
mean for the
population, n
260

. Apparent equilibrium association constants for surface, Ks, and mter-
calative, Kt , binding of 3-aminoacridinium to DNA, NaH2PC>4 as supporting
electrolyte®
I
(ave)
Ks M'1
x 10-4
o (n)
-4
x 10
I
(ave)
Kj M 2
x 10"9
a (n)
x 10~9
.100
3.49
0.221(4)
. 098
0.518
0.0825(3)
.099
3.31
3.40
0.155(4)
.098
0.363
0.441
0.0817(5)
. 050
5.23
0.268 (3)
. 050
1.67
0.234 (4)
.050
5.42
5.33
0.258(3)
. 050
1.56
1.62
0.223 (3)
.025
7.51
0.312(3)
. 025
4.59
0.797(4)
. 025
7.96
7.74
0.284 (3)
. 025
5.55
5.07
0.875(4)
. 010
12.7
0.265(3)
. 010
16.4
1.20(2)
. 010
12.8
12.8
0.071 (2)
. 010
17.0
16.7
1.27 (2)
. 0050
27.8
2.59(3)
.0051
59.8
10.3 (2)
. 0025
31.2
7.61(3)
. 0026
75.0
22.7(2)
aConstants
based on molar
concentrations
of species.
T = 25.0°C,
pH 5.90.
b
a is the standard deviation from the mean for the population,
n.
261

Apparent equiiiuriuiu dasucidiiun cunsLdias
calative, Kj, binding of 3-aminoacridinium
electrolyte3
ror surxace, as , ana muer¬
to DNA, KH2PC>4 as supporting
I
KM1
a (n)b
I
Kt M-2 a(n)b
(ave)
s
-4
-4
(ave)
-9 -9
x 10
x 10
x 10 x 10 3
.10
3.48
0.252 (4)
.098
.478
0.122(6)
.10
4.33
3.91
0.147 (4)
. 098
.644
.561
0.135 (6)
. 050
6.12
0.266(5)
. 050
1.84
0.285(5)
. 050
5.97
6.05
0.552 (5)
. 050
2.07
1.96
0.306 (5)
.025
9.48
0.348 (3)
.025
5.44
0.502(5)
. 025
8.85
9.17
0.350 (3)
. 025
5.25
5.35
0.389 (5)
.010
13.6
0.594 (4)
.010
14.7
1.14 (6)
. 010
13.6
13.6
0.793 (3)
. 010
15.0
14.9
0.997 (4)
. 0050
11.0
0.583 (4)
. 0051
11.8
0.707 (4)
. 0050
23.0
17.0
3.75(3)
.0051
26.4
19.1
3.06 (4)
.00066
43.1
8.54 (3)
.00072
176
29.1(3)
.00066
30.0
36.6
4.43(4)
.00073
112
144
21.0 (5)
a
Constants based on molar concentrations of species.
T
25.0°C, pH 5.90.
ka is the standard deviation from the mean for the population, n.
262

Apparent equilibrium association constants
calative, Ki, binding of 3-aminoacridinium
electrolyte3
for surface, Kg, and mter-
to DNA, KC^CCH.^ as supporting
I
(ave)
K M-1
s
x 10“4
a(n)b
x 10"4
I
(ave)
K M-2 a(n)b
-9 -9
x 10 3 x 10 3
.15
2.64
0.249 (4)
.14
.204
0.047(6)
.15
2.76
2.70
0.080(4)
.15
. 247
.226
0.044(4)
.050
5.70
0.547 (4)
. 050
1.96
0.201(5)
. 050
5.37
5.54
0.269 (4)
. 050
1.86
1.91
0.217(5)
.025
8.92
1.12(3)
. 025
5.75
0.589(5)
. 025
7.40
8.16
0.379(3)
. 025
3.64
4.70
0.481(5)
.0050
26.4
3.33(4)
.0051
39.8
7.39(4)
. 0050
25.5
7.14 (2)
. 0051
29.6
6.20(3)
. 0050
15.8
22.6
1.13(3)
. 0051
27.2
32.2
2.59(4)
. 0025
19.8
1.58(4)
.0026
51.8
3.69(4)
. 0025
19.2
19.5
0.494(2)
.0026
49.6
50.7
6.21(4)
.00066
17.0
1.51(5)
.00075
55.7
5.74 (4)
.00066
11.2
14.1
1.25(3)
.00075
23.0
39.4
1.36 (3)
a
Constants based on molar concentrations of species.
T = 25.0°C,
pH 5.90.
263

'Ule L¡ ILUlillIiUtídJ
b
O
is the standard deviation from the mean for the population, n.

Table 28. Apparent equilibrium association constants for surface, Ks, and inter-
calative, Kj, binding of 3-aminoacridinium to DNA, RbH-PO. as supporting
electrolyte3 4
I
K M-1
a (n)b
I
Kj M-2
o(n)b
(ave)
s
x 10-4
x 10"4
(ave)
x 10-9
x 10"9
.10
3.95
0.162 (4)
.098
0.566
0.118(6)
.10
3.73
3.84
0.397 (4)
. 098
0.629
0.598
0.134 (8)
. 050
5.66
0.619 (4)
.049
1.98
0.167 (4)
. 050
6.45
6.06
0.205(2)
.049
2.11
2.05
0.317 (6)
.025
9.17
0.197 (3)
.025
5.28
0.606 (5)
.025
10.9
0.283 (2)
.025
6.78
0.445(5)
.025
10.2
10.1
0.601 (2)
.025
6.70
6.25
0.605 (4)
. 010
11.5
1.70(2)
. 010
17.8
1.31 (3)
. 010
17.5
14.5
2.83 (2)
.010
26.4
22.1
2.02(5)
. 0050
13.8
1.67(4)
.0051
18.5
1.65(5)
. 0050
6.74C
2.23 (3)
.0051
67.4*
22.3 (3)
. 0050
15.7
14.8
0.902 (4)
. 0051
27.3
22.9
2.94(5)
265

Table 28 (Continued)
I
(ave)
K M-1
s
x 10~4
a (n)b
x 10-4
I
(ave)
Kx M-2
x 10~9
o(n)n
x 10-9
. 0025
16.2
0.850(3)
.00026
28.2
2.08 (4)
. 0025
16.6
0.379 (3)
.00026
24.8
2.54 (5)
16.4
26.5
Constants based on molar concentrations of species.
O
0
o
•
in
CM
II
Eh
pH 5.90.
a is the standard deviation from the mean for the population, n.
c
Value not included in average.

Table 29. Apparent equilibrium association constants for surface, Ks, and inter-
calative, Kj, binding of 3-aminoacridinium to DNA, CsI^PO^ as supporting
electrolyte3
I
(ave)
K M-1
a (n)
s
-4
x 10
x 10
I
(ave)
-2
K M
0 (n)
-Q
X 10 "
x 10
.15
2.30
0.154(6)
.15
2.59
2.45
0.275(6)
.10
3.49
0.212(6)
.10
4.99
4.24
0.610(5)
. 050
4.21
0.534 (5)
. 050
5.99
5.10
0.320(6)
.025
8.88
0.384(5)
. 025
5.40
7.14
0.384 (5)
. 010
11.7
0.390(5)
. 010
12.3
12.0
1.37 (4)
. 0025
19.6
2.50(4)
.0025
18.5
19.1
3.44(4)
c
. 098
0.414
0.097(6)
. 098
0.627
0.521
0.115(4)
.049
0.965
0.022 (5)
. 050
1.71
1.34
0.176 (5)
. 025
4.13
0.429 (5)
. 025
1.34
2.74
0.172 (5)
. 010
9.97
1.21(5)
. 010
7.55
8.76
0.588 (4)
.0026
34.2
5.09 (5)
. 0026
15.9
25.1
2.91 (5)
267

Table 29 (Continued)
a , . o
Constants based on molar concentrations of species. T = 25.0 C, pH
a is the standard deviation from the mean for the population, n.
C . "I" .
Intercalative binding not seen due to high Cs concentration.
.90.

Table 30.
Apparent equilibrium association constants
calative, Kj, binding of 3-aminoacridinium
electrolyte^
for the
to DNA,
surface, Ks, and inter-
(CH^^NI^PO^ as supporting
I
K M_1
a (n)b
I
-2
Kt m
a (n) b
(ave)
x 10-4
x 10-4
(ave)
-9
x 10 y
x 10"9
.10
6.75
0.317(3)
. 099
2.39
0.235(5)
.10
5.35
6.05
0.296 (4)
. 099
1.75
2.07
0.226(4)
. 050
6.92
0.578 (2)
.050
3.86
0.059 (4)
. 050
7.58
7.25
0.288 (4)
. 050
4.06
3.96
0.169 (3)
.025
10.5
0.495 (2)
.025
4.19
0.577(4)
. 025
8.80
4.65(4)
. 025
8.33
0.417 (3)
. 025
13.3
0.0(2)
.025
12.0
1.07 (4)
.025
10.1
10.7
0.630(3)
. 025
8.78
8.33
0.849(3)
.010
13.1
0.785 (3)
.010
21.7
3.01(3)
. 010
17.6
2.50 (3)
. 010
36.3
0.905 (2)
. 010
19.0
16.6
2.34 (2)
.010
30.9
29.6
4.06(4)
.0050
27.0
0.712(3)
. 0051
51.7
6.59 (2)
.0050
25.3
26.2
0.488 (2)
. 0051
.0051
74.4
53.3
14.0 (4)
10.8 (4)
59.8
269

Table 30 (Continued)
I
(ave)
K M-1
s
x 10~4
a (n) b
x 10-4
I
(ave)
Kx M~2
x 10~9
o(n)b
x 10~9
.0025
21.3
1.18 (3)
.0025
100
9.00(2)
. 0026
63.9
7.77(4)
.0025
50.4
32.0(2)
. 0026
109
4.48 (4)
57.2
.0026
55.1
11.5(4)
76.0
aConstants based on molar concentrations of species.
T = 25.0°C
, pH 5.90.
a is the standard deviation from the mean for the population, n.

Table 31. Apparent equilibrium association constants, Ks, for surface binding of
3-aminoacridinium to DNA, Mg(0„CCHn)9 and Ca (O-CCH-J ~ as supporting
electrolytes3
Mg(02CCH3)2
Ca(02CCH3)2
I
(ave)
K N-1
s
x 10-4
0 (n)
_4
x 10
I
(ave)
K N-1
s
x 10"4
o (n)
x 10-4
. 015
1.32
0.031
(5)
.037
.968
0.035 (5)
. 015
1.32
1.32
0.026
(4)
. 015
1.43
0.032 (4)
. 0077
2.08
0.054
(4)
.0077
1.68
0.042 (5)
. 0077
2.04
2.06
0.073
(4)
.0076
2.15
1.92
0.107 (4)
. 0039
2.76
0.026
(4)
.0040
2.41
0.034 (5)
.0039
2.76
2.76
0.079
(4)
.0039
2.34
2.38
0.024(5)
.0021
4.27
0.025
(4)
.0021
3.48
0.520 (4)
.0021
3.50
3.89
0.082
(4)
.0022
2.45
2.97
0.036 (5)
.0011
4.74
0.040
(4)
.0011
2.98
0.078 (5)
.0011
4.90
4.82
0.170
(2)
.0011
3.54
3.26
0.201(5)

Table 31 (Continued)
aConstants based on molar concentrations of species. T = 25.0°C, pH 5.90.
a is the standard deviation from the mean for the population, n.

Table 32. Apparent equilibrium association constants, Ks, for the surface binding
of 3-aminoacridinium to DNA, Sr (O-CCH.,) „ and Ba(0»CCH?)9 as supporting
electrolytes3
I
(ave)
Ks Kf1
x 10_4
a (n)
x 10"4
I
(ave)
Ks S'1
x 10-4
a (n)
x 10~4
. 037
1.03
0.018(4)
.037
1.03
1.03
0.013(4)
'
.015
1.64
0.044(5)
. 015
1.78
0.021(5)
.015
1.65
0.052 (5)
. 015
1.74
0.042(5)
1.65
1.76
. 0076
3.64
0.752 (4)
.0077
2.42
0.144 (4)
. 0077
2.15
0.021(5)
.0076
2.60
0.044 (4)
. 0077
2.57
2.36
0.048(5)
2.51
. 0039
2.90
0.381 (4)
.0039
3.26
0.047 (5)
. 0039
2.94
0.209 (4)
. 0039
3.34
0.037(4)
2.92
3.30
.0020
4.05
0.087 (3)
. 0020
4.99
0.161(3)
. 0020
4.91
0.142 (4)
.0020
3.93
0.132(4)
4.48
4.46
. 0010
6.30
0.312 (3)
.0010
5.73
0.091(4)
6.02
273

Table 32 (Continued)
I
(ave)
x 10
a (n)
x 10
I
(ave)
K N 1 o(n)
3
x 10-4 x 10
X
10 l
12.1
X
10
6.03
X
10"5
13.7
12.9
0.469(4)
0.160 (4)
1.24 (4)
aConstants based on molar concentrations of species. T = 25.0°C, pH 5.90.
a is the standard deviation from the mean for the population, n.

Table 33
Logarithms of the apparent surface binding association constants for the binding
of 3-aminoacridinium cation to DNA, corrected for ion activities3
Init. LiH
Cone.
of
Elec- £
trolvte
(M)
2P04
log x's
NaH
?s
2P04
log K¡
kh2
^s
P04
iog K¡
ko2
^s
CCH3
iog K¡
RbH
^s
2P°4
iog K¡
CsH
^s
2P°4
iog Ks
(CK
^s
3>4NH2P04
1
log Ks
0.15
.
5.31
5.156
_
_
5.20
5.104
_
.10
4.28
5.086
4.24
5.158
4.22
5.217
-
-
4.20
5.207
4.20
5.250
4.25
5.410
. 050
3.12
5.085
3.10
5.21‘8
3.09
5.271
3.12
5.236
3.09
5.272
3.09
5.198
3.10
5.352
. 025
2.51
5.132
2.38
5.264
2.38
5.339
2.38
5.289
2.38
5.377
2.38
5.231
2.39
5.407
. 010
1.82
5.259
1.82
5.789
1.82
5.393
-
-
1.82
5.420
1.82
5.338
1.82
5.478
. 0050
1.56
5.392
1.56
5.637
1.56
5.423
1.56
5.546
1.56
5.362
1.56
-
1.56
5.610
. 0025
1.38
5.473
1.38
5.634
-
-
1.39
5.432
1.38
5.355
1.38
5.420
1.38
5.898
.0013
1.26
5.392
-
-
-
-
-
-
1.26
-
1.26
-
1.26
-
6.3 x
10“ 4
1.19
5.601
1.19
5.224
1.19
1.19
1.19
275

Table 33 (Continued)
Init. Mg(02CCH3)2
Cone.
of
Elec- £ log K
trolyte
(N)
Ca(02CCH3)
log Ks
Sr(02CCH3)2
^s log Ks
Ba(02CCH3)
log
0.025
-
-
0.010
1.94
4.409
0.0050
1.66
4.534
0.0025
1.47
4.608
0.0013
1.32
4.711
6.3 x
10-4
1.24
4.776
2.51
4.386
2.50
1.93
4.441
1.93
1.66
4.501
1.65
1.46
4.542
1.47
1.32
4.594
1.31
1.24
4.606
-
4.410
-
-
4.502
1.93
4.532
4.592
1.65
4.619
4.631
1.47
4.685
4.784
1.31
4.768
-
1.233
4.870
9. 8 __J_ _1_
Ion size parameters, d x 10 , from reference (87) are: Li (8); Na (4); K (3); Rb ,Cs (
(CH3)4N+, H2PO4 (4.5); Sr4"4, Ba1-+(5) ; Caf+(6); and M<^+(8). Values estimated for BH, BHP,
P are 3, 4, and 4 (x 108), respectively.
2.5) ;
and
276

277
Table 34. Equilibrium association constants for the binding
of 3-aminoacridinium to DNA, Ks, and for the
binding of alkali and alkaline earth metal cations
to DNA, KMa
Alkali metals^
metal
Na+
K+
Rb+
Cs+
(ch3)4n+
metal
cone.(M)
0.10-
0.025
0.050-
0.010
0.050-
0.010
0.050-
0.025
0.10-
0.025
leg k's
5.158-
5.264
5.271-
5.393
5.272-
5.420
5.198-
5.420
5.410-
5.407
km
4.06
8.82
11.3
14.6
-0.09
log K
3 s
5.31
5.43
5.47
5.44
5.41
Q
Alkaline earth metals
metal
++
Mg
Ca++
Sr++
++
Ba
metal
cone.(N)
0.010-
0.0025
0.0050-
0.0013
0.010-
0.0025
0.010-
0.0025
lo9 Ks
4.409-
4.608
4.501-
4.594
4.410-
4.631
4.532-
4.685
km
96.1
70.5
H4.
65.5
leg k"
4.70
4.63
4.74
4.75
Calculated using equation (3-5). All reactions at 25.0 +
0.1°C, pH 5.9. For DNA and 3-AA reactant concentrations
and other experimental parameters, refer to the appropriate
tables for the specific metals.

278
Table 34 (Continued)
^Counteranion, H2P0^
cCounteranion, CH^COO

I I
Table 35. Apparent association constants for the surface, Ks, and intercalative, Ki,
binding
and 35.
of 3-aminoacridinium and 7-aminoquinolinium to
o°ca
DNA at 15.
0°C, 25.0'
3-Aminoacridinium
7-Aminoquinolinium
15.0°C
1 -1
K M x
S -5
x 10 3
a (n)
' _2
KT M
, n-10
x 10
o (n)
' -i
KM
S -4
x 10
a (n)
k' m~2
x 10"8
a (n)
3.60
.129 (4)
2.75
.606 (4)
5.50
. 306(3)
2.84
.163(3)
3.15
.211 (4)
2.54
. 386 (3)
4.00
.819 (2)
1.98
.096(2)
2.62
.145(3)
2.14
.181(4)
25.0°C
2.48
.108(4)
1.47
.114(6)
2.70
.755(3)
1.08
.202(5)
2.48
.114(3)
1.50
.100(5)
2.95
.242 (3)
1.34
.562 (5)
35.0°C
1.78
.149
0.479
.0754 (4)
2.49
.039 (2)
0.604
.0104 (5)
1.98
.072
0.588
.0524(5)
2.44
.291(3)
0.494
.0213 (3)
2.17
.169(3)
0.434
.0157 (4)
2.66
.339(3)
0.395
.0187 (4)
aBased on activities of all species, pH 5.90.
279

Table 36. Thermodynamic parameters for the surface and intercalative binding reactions
between 3-aminoacridinium and 7-aminoquinolinium cations to DNA
cation
3-aminoacridinium
7-aminoquinolinium
3-aminoacridinium
7-aminoquinolinium
Log K AG° AH° AS°
(Kcal) (Kcal) (Kcal) (e.u.)
2 98°K 298°K
surface binding
5.39 -7.35 -4.90 + 1.3 8.22 + 1.7
4.45 -6.07 -5.13 + 1.2 3.15 + 0.94
10.17
8.08
intercalative binding
-13.05 -13.7 + 5 -2.2 + 6
-11.02 -15.1 + 4 -14 + 13

APPENDIX III
COMPUTER PROGRAMS

These programs were designed for the calculations
of apparent association constants and related data
employing a Litton-Monroe model 326 Scientist computer.
Terms not defined here are either defined in the text
or are explained in the operator's manual for the
computer. Brackets signify molar concentrations, un¬
bracketed terms indicate moles. V denotes volume,
subscripts o, t, and 00 refer to values before any DNA
has been added, values at a given point in the titration,
and values after excess DNA has been added, respectively.
Program I
This program has been designed to compute the
following: [BH], [BHP], Kj, log([P / ]-2[BHP]), log([BHP]/
[BH]), K , and log([P /_]- [BHP]). For the program as
O L J
written, the total initial volume must be 10.00 ml with
a sample pathlength of 4.00 cm. Parameters which remain
constant throughout a given titration are stored in
registers 1, 2, 4, and 9 (their contents are listed in the
table below). All data is entered as written, using
scientific notation, and all results are displayed in
their final form, that is, no correction factors are
necessary. To run a program, enter total volume of DNA
282

283
solution (yl), press START, enter total absorbance of
the solution, press START. The program will be executed
automatically, pausing for 3 to 6 seconds to display each
of the above values in the sequence in which they are
listed. To initiate another computational cycle, repeat
the above procedure. At the completion of a computation,
the registers will have stored within them the following
information which may be recalled, if necessary.
Register
0
1
Information
2
3 [BHP]t
4
5
6
BH
o
[BH] t
[V3]t
Register Information
7 V in liters
8 Total 1 of DNA
added up to
point t
9 Pt//2 (m°les
DNA phosphate
per 1)
CHG K
SGN
log([Pt/2l-2[BHP]t
The individual steps of the program, along with their key¬
strokes and Munroe 326 code numbers are presented below.

1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
284
KEY
STROKES
CODE
STEP
KEY
STROKES
CODE
ST
300
51
)
027
8
010
52
T
024
STOP
033
53
RCL
310
ST
300
54
7
007
0
000
55
=
020
(
026
56
ST
300
RCL
310
57
5
005
8
010
58
fEXP
114
X
023
59
RCL
310
EXP
014
60
3
003
CHSN
013
61
T
024
6
006
62
RCL
310
)
027
63
7
007
+
021
64
=
020
•
012
65
ST
300
0
000
66
3
003
1
001
67
fEXP
114
=
020
68
RCL
310
ST
300
69
3
003
7
007
70
—
024
RCL
310
71
(
026
9
Oil
72
(
026
X
023
73
RCL
310
RCL
310
74
6
006
8
010
75
-
022
T
024
76
(
026
RCL
310
77
RCL
310
7
007
78
3
003
=
020
79
X
023
ST
300
80
2
002
6
006
81
)
027
(
026
82
)
027
RCL
310
83
Ax
025
1
001
84
2
002
-
022
85
X
023
RCL
310
86
RCL
310
0
000
87
5
005
)
027
88
)
027
X
023
39
= â– 
020
RCL
310
90
ST
300
2
002
91
CHSN
013
=
020
92
fEXP
114
ST
300
93
(
026
3
003
94
RCL
310
(
026
95
6
006
RCL
310
96
-
022
4
004
97
(
026
-
022
93
RCL
310
RCL
310
99
3
003
3
003
100
X
023

285
STEP
KEY CODE
STROKES
STEP
KEY
STROKES
CODE
101
2
002
129
3
003
102
)
027
130
024
103
)
027
131
(
026
104
ST
300
132
RCL
310
105
.
012
133
5
005
106
LOG
061
134
X
023
107
fEXP
114
135
(
026
108
fEXP
114
136
RCL
310
109
(
026
137
6
006
110
RCL
310
138
-
022
111
3
003
139
RCL
310
112
T
024
140
003
003
113
RCL
310
141
)
027
114
5
005
142
)
027
115
)
027
143
)
027
116
LOG
061
144
fEXP
114
117
fEXP
114
145
fEXP
114
118
RCL
310
146
(
026
119
6
006
147
RCL
310
120
f
024
148
6
006
121
1
001
149
-
022
122
.
012
150
RCL
310
123
5
005
151
3
003
124
=
020
152
)
027
125
ST
300
153
LOG
061
126
6
006
154
STOP
033
127
(
026
155
JUMP
350
128
RCL
310
156
START
033
Program 2
This commercially
available program (Litton-Monroe,
Orange, N.J.,
number 9042W)
solves
the simultaneous
equations (2-
8), (2-9),
and
(2-10)
for the
concentrations
of 7-aminoquinoline-containing species, [AHP], [AH], and
[A] .
Program
execution
is outlined in the
operator's
manual. The
following
table correlates the
terminology
of the program with the
quantities
used for
this specific
problem based on the general equations

286
A,X + BjY + C1Z = P
A2X + B2Y + C3Z = Q
A3X + B3Y + C3Z = R
Program
designation
Variables or
Coefficient
X
Y
Z
A1
B]
C]
P
A,
B„
[AHP]
[AH]
[A]
rAHp (5332 M_1cm"1)
£AH M lcm
e (347 M ^cm â– *")
v4
1
1
1
Q
A,
B
Ct = Co/Vt
Ka (2.239 x 10 7)
- [H+] (-1.259 x 10-6)
R
Program 3
This program calculates association constants for the
binding of 7-aminoquinolinium to DNA. The following are
computed and displayed sequentially, as written: Kp, log([AHP]/
[AH]), log([Pt/2]-2[AHP]), Kg, and log([P /3]-[AHP]). A total

287
initial volume of 8.03 ml and a solution pathlength of
4.00 cm is required. Concentrations of AHP and AH are
obtained from equations (2-8), (2-9), and (2-10), using
Program 2. Before running a program (moles of DNA
phosphate per yl)•1/2 is loaded into register 9. To
execute a program, enter yl of DNA solution, press START,
5 5
enter [AHP] x 10 , press START, enter [AH] x 10 , press
START. The concentrations of AHP and AH are converted to
the proper orders of magnitude in the program. At the
completion of a computation the registers will contain
the following information for retrieval, if necessary.
Register Information Register Information
0 [AHP]^ 7 V in liters
1 [AH] 8 Ml of DNA
2
9
Pt/2 (1/2 moles
DNA phosphate
per yl)
3
K
s
4
log ( [AHP] /[AH] t)
CHG
SGN
K
I
5
6
The individual steps of the program, along with their
keystrokes and Munroe 326 code numbers are presented below.

1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
288
KEY
STROKES
CODE
STEP
KEY
STROKES
CODE
ST
300
51
CHSN
013
8
010
52
5
005
STOP
033
53
=
020
ST
300
54
ST
300
0
000
55
1
001
STOP
033
56
RCL
310
ST
300
57
0
000
1
001
58
T
024
(
026
59
(
026
RCL
310
60
(
026
8
010
61
RCL
310
X
023
62
6
006
EXP
014
63
-
022
CHSN
013
64
(
026
6
006
65
RCL
310
)
027
66
0
000
+
021
67
X
023
•
012
68
2
002
0
000
69
)
027
0
000
70
)
027
8
010
71
Ax
025
=
020
72
2
002
ST
300
73
X
023
7
007
74
RCL
310
(
026
75
1
001
RCL
310
76
)
027
9
Oil
77
=
020
X
023
78
ST
300
RCL
310
79
CHSN
013
8
010
80
fEXP
114
)
027
81
(
026
024
82
RCL
310
RCL
310
83
0
000
7
007
84
7
024
=
020
85
RCL
310
ST
300
86
1
001
6
006
87
)
027
RCL
310
88
LOG
061
0
000
89
ST
300
X
023
90
4
004
EXP
014
91
fEXP
114
CHSN
013
92
(
026
5
005
93
RCL
310
=
020
94
6
006
ST
300
95
-
022
0
000
96
(
026
RCL
310
97
RCL
310
1
001
98
0
000
X
023
99
X
023
EXP
014
100
2
002

289
STEP
KEY
STROKES
CODE
STEP
KEY
STROKES
CODE
101
)
027
124
RCL
310
102
)
027
125
0
000
103
LOG
061
126
)
027
104
STOP
033
127
X
023
105
RCL
310
128
RCL
310
106
6
006
129
1
001
107
X
023
130
)
027
108
#
012
131
=
020
109
6
006
132
ST
300
110
6
006
133
#
012
111
6
006
134
fEXP
114
112
6
006
135
(
026
113
=
020
136
RCL
310
114
ST
300
137
2
002
115
2
002
138
-
022
116
RCL
310
139
RCL
310
117
0
000
140
0
000
118
%
024
141
)
027
119
(
026
142
LOG
061
120
(
026
143
STOP
033
121
RCL
310
144
JUMP
350
122
2
022
145
START
033
123
-
022
Program 4
This program calculates the activity coefficients of
species at a given ionic strength based on equation (1-30).
The coefficients may be stored for subsequent determination
I
of K from K [equation (2-17)] . The program is designed for
aqueous systems at 25.0°C for which B is equal to 3.3 x 10^.
To execute a program, enter I, press ST 7, enter d for the
ith species, press ST 8, enter Z^, press ST 6, press JUMP,
CHG SGN, START. The value of will be displayed. Store
the values of aBHp/ “bh' anc^ ap or t*le resPecti-ve values
for AHP, AH, and A in registers 4, 2, and 5, respectively.
Upon storing these values, the correction term for

290
intercalative binding [n = 2 of equation (2-17) may be
computed by pressing JUMP, 9, START].. For the surface
binding correction factor press JUMP, ., START.
At the completion of a series of computations, the
registers will contain the following information which
may be retrieved, if necessary.
Register Information
Register Information
0
1
°1BHP//{aBH'aPt)
aBHP/(aBH*aPt2)
2
3
4
a
BH
a
BHP
6 Z^ (the last one
entered)
7 I
8 d^ (the last one
entered)
9
CHG
SGN
The individual steps of the program, along with their
keystrokes and Munroe 326 code numbers are presented below.
STEP
KEY
STROKES
CODE
1
f.LABEL
213
2
(
026
3
(
026
4
RCL
310
5
6
006
6
AX
025
7
2
002
3
X
023
9
(
026
10
RCL
310
STEP
KEY
STROKES
CODE
11
7
007
12
Ax
025
13
»
012
14
5
005
15
X
023
16
#
012
17
5
005
18
0
000
19
9
Oil
20
)
027

291
STEP
KEY
STROKES
CODE
STEP
KEY
STROKES
CODE
21
)
027
51
RCL
310
22
i
024
52
4
004
23
(
026
53
T
024
24
(
026
54
(
026
25
RCL
310
55
RCL
310
26
7
007
56
5
005
27
Ax
025
57
X
023
28
012
58
RCL
310
29
5
005
59
2
002
30
X
023
60
)
027
31
RCL
310
61
)
027
32
8
010
62
ST
300
33
X
023
63
0
000
34
3
003
64
STOP
033
35
012
65
JUMP
350
36
3
003
66
#
012
37
EXP
014
67
fLABEL9
211
38
7
007
68
(
026
39
)
027
69
RCL
310
40
+
021
70
0
000
41
1
001
71
T
024
42
)
027
72
RCL
310
43
)
027
73
5
005
44
CHSN
013
74
)
027
45
ALOG
161
75
ST
300
46
STOP
033
76
1
001
47
JUMP
350
77
STOP
033
48
CHSN
013
78
JUMP
350
49
fLABEL.
212
79
9
Oil
50
(
026
Program 5
This program calculates the ionic strength of a
solution resulting from the addition of a given volume of
DNA to 10.00 ml of solution having a known initial ionic
strength. The contribution of background electrolyte in
the DNA solution is included. An initial and final volume
of DNA solution may be entered to obtain an averaged ionic
strength over a range. Calculations for solutions containing
monovalent or divalent cations may be made depending on the

292
constants used for steps 51 and 67 of the program. Before
running the program, the normality of the 10.00 ml solution
containing the BH or AH is stored in register 0 and 1/2
(moles of DNA phosphate + equivalents of supporting elec¬
trolyte in the DNA solution) per pi is stored in register
2. To execute a program, enter of DNA solution (pi),
press START, enter V2 of DNA solution (pi), press START.
The averaged ionic strength will be displayed. If the
total final ionic strength of a solution containing a
given, single, volume of DNA is desired, execute the
program by entering that value (pi) and pressing START,
START. At the completion of a calculation the registers
will contain the following information which may be
retrieved, if necessary.
Register Information
0 Normality, N, of the 10.00 ml solution before
any DNA has been added.
1 Average volume of DNA added (pi) = (V,DNA +
V2DNA)/2. 1
2 1/2(moles of DNA phosphate + equivalents of
supporting electrolyte in the DNA solution).
3 Total final average volume of solution =
0.01000 + average volume of DNA added, in
liters.
2
4 Total final volume of solution x 10 .
5 Ionic strength, I, of the final solution if
register 0 contains 0.
6 I of the final solution if register 2 contains
0 (dilution effect only).

293
Register
Information
7 V1DNA (yl).
8 V2DNA (yl).
9 I of final solution (averaged volume and
averaged ionic strength).
CHG
SGN
The individual steps
keystrokes and Munroe 326
of the program, along with their
code numbers are presented below.
STEP
KEY
STROKES
CODE
STEP
KEY
STROKES
CODE
1
ST
300
31
2
002
2
7
007
32
)
027
3
STOP
033
33
ST
300
4
ST
300
34
3
003
5
8
010
35
(
026
6
(
026
36
RCL
310
7
(
026
37
3
003
8
RCL
310
38
X
023
9
7
007
39
EXP
014
10
+
021
40
2
002
11
RCL
310
41
)
027
12
8
010
42
ST
300
13
)
027
43
4
004
14
T
024
44
(
026
15
2
002
45
(
026
16
)
027
46
RCL
310
17
ST
300
47
1
001
18
1
001
48
X
023
19
(
026
49
RCL
310
20
(
026
50
2
002
21
RCL
310
51
X
023
22
1
001
52
3*
003*
23
X
023
53
)
027
24
EXP
014
54
T
024
25
CHSN
013
55
RCL
310
26
6
006
56
3
003
27
)
027
57
)
027
28
+
021
58
ST
300
29
EXP
014
59
5
005
30
CHSN
013
60
(
026

61
62
63
64
65
66
67
68
69
70
71
72
73
74
75
76
77
78
294
KEY
STROKES
CODE
STEP
KEY
STROKES
CODE
(
026
79
RCL
310
RCL
310
80
6
006
0
000
81
)
027
X
023
82
ST
300
1
001
83
9
011
•
012
84
STOP
033
5* *
005**
85
JUMP
350
)
027
86
START
033
T
024
RCL
310
4
004
*
51
3
003
for
m2+
)
027
51
2
002
for
M+1
ST
300
6
(
006
026
**
67
5
005
for
„2 +
M ,
RCL
310
67
0
000
for
M1+
5
005
+
021

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300
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BIOGRAPHICAL SKETCH
Peter F. Eisenhardt was born on March 18, 1945, in
Cortland, New York. He graduated from Dryden Central High
School in June, 1963. In September, 1963, he entered SUNY
at Cortland and graduated in June, 1967, with a Bachelor
of Science degree in chemistry. He entered the Department
of Chemistry at the University of Florida in the Fall of
1968 and received a Master of Science degree in inorganic
chemistry in June, 1971. For the next three years he was
employed at the College of Pharmacy, University of Florida
where he developed analytical methods for an ophthalmic
ointment development program, before beginning studies
toward his doctorate in September, 1974. He is a member
of the American Pharmaceutical Association, Academy of
Pharmaceutical Sciences, American Chemical Society, Florida
Section of the American Chemical Society, Alpha Chi Sigma
Chemistry Fraternity, and Delta Kappa Pi.
Upon receiving his doctorate he will join Endo
Laboratories as an Analytical Chemist.
He married the former Joanne Rissmann in June, 1972.
301

I certify that I have read this study and that in
my opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
Stephen G. Schulman,/Chairman
Professor of Pharmaceutical
Chemistry
I certify that I have read this study and that in
my opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
j
U ár'X /
Andre seri
Brian D.
Assistant Professor of
Pharmaceutical Chemistry
I certify that I have read this study and that in
my opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
l A
Merle A. Battiste
Professor of Chemistry

I certify that I have read this study and that in
my opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
Professor of Pharmacy
I certify that I have read this study and that in
my opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
Richard H. Hammer
Professor of Pharmaceutical
Chemistry
This dissertation was submitted to the Graduate Faculty of
the College of Pharmacy and to the Graduate Council, and was
accepted as partial fulfillment of the requirements for the
degree of Doctor of Philosophy.
August, 1977
Dean, College of Pharmacy
Dean, Graduate School

UNIVERSITY OF FLORIDA
3 1262 08554 5225



UNIVERSITY OF FLORIDA
3 1262 08554 5225


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