<%BANNER%>
Proton transfer and operational pH in aqueous-organic solvent systems
CITATION THUMBNAILS PAGE IMAGE ZOOMABLE
Full Citation
STANDARD VIEW MARC VIEW
Permanent Link: http://ufdc.ufl.edu/AA00011847/00001
 Material Information
Title: Proton transfer and operational pH in aqueous-organic solvent systems
Physical Description: xvii, 230 leaves : ill. ; 29 cm.
Language: English
Creator: Townsend, Robert W., 1953-
Publication Date: 1992
 Subjects
Subjects / Keywords: Research   ( mesh )
Protons -- chemistry   ( mesh )
Hydrogen-Ion Concentration   ( mesh )
Solvents   ( mesh )
Biphenyl Compounds -- chemistry   ( mesh )
Quinolines -- chemistry   ( mesh )
Isomerism   ( mesh )
Spectrometry, Fluorescence -- methods   ( mesh )
Department of Pharmaceutics thesis Ph.D   ( mesh )
Dissertations, Academic -- College of Pharmacy -- Department of Pharmaceutics -- UF   ( mesh )
Genre: bibliography   ( marcgt )
non-fiction   ( marcgt )
 Notes
Thesis: Thesis (Ph.D.)--University of Florida, 1992.
Bibliography: Bibliography: leaves 218-229.
Statement of Responsibility: by Robert W. Townsend.
General Note: Typescript.
General Note: Vita.
 Record Information
Source Institution: University of Florida
Rights Management: All rights reserved by the source institution and holding location.
Resource Identifier: aleph - 002293326
oclc - 49673996
notis - ALP6525
System ID: AA00011847:00001

Full Text











PROTON TRANSFER AND OPERATIONAL PH
IN AQUEOUS-ORGANIC SOLVENT SYSTEMS




















By

ROBERT W. TOWNSEND


















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

UNIVERSITY OF FLORIDA

1992




































For my parents Leland and Frances Townsend, who dared me to

dream, and my daughter, Harmony, for the desire. This is for

you.














ACKNOWLEDGEMENTS


I would like to thank the members of my committee, Dr.

John Perrin, Dr. Hans Schreier and Dr. Ken Sloan, for their

friendship and support throughout this undertaking. I would

like to express my gratitude to Dr. Shangxian Chen for his

gentle guidance and valuable insight into the

physicochemical aspects of this project. Special thanks go

out to Dr. Kirk Schanze for his assistance and use of the

time-correlated single photon counter. I would also like to

recognize Anup Zutshi and Yi Zhang for being there through

it all. Finally, I would like to express my thanks and

deepest appreciation to my advisor, Dr. Stephen Schulman,

for his guidance and understanding and for the support

required to complete this dissertation.















TABLE OF CONTENTS



ACKNOWLEDGEMENTS . . . . . . . . ... iii

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

LIST OF FIGURES . . . . . . . . . . xi

ABSTRACT . . . . . . . . . . . .xvi

CHAPTERS

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

Physico-Chemical Properties of Water. . . ... 26
Solvation .... . . . . . . . 31
Hydrocarbons .. . . . . .. 33
Monohydroxylic Alcohols . . . . . 34
Ions . . . . . . . . . . 36
Relaxation Processes . . . . . .. 37
Proton Transfer. . . . . . . . . . 39
Proton Transfer in the Excited Singlet State . . 42
Electronic and Accompanying Spectral
Effects . . . . . . . ... . .43
Steady-State Kinetics . . . . . .. .44
Time-resolved Kinetics. . . . . . .. .52
Objectives . . . . . . . . .. . 58

2 EXPERIMENTAL . . . . . . . . ... .63

Materials. . . . . . . . . ... .63
Instrumental. . . . . . . . . 64
Methods . . . . . . . . ... . . 65
Steady-State Emission Spectroscopy. . . .. .65
Aqueous acid-base titrations . . .. .65
Aqueous-organic solvent
titrations . . . . . . . .. .67
Aqueous-organic acid-base
titrations . . . . . . . .. 68
Time-resolved Emission Spectroscopy . . .. .71
Aqueous acid-base titrations . . .. .74
Aqueous-organic solvent
titrations . . . . . . .. 75








Aqueous-organic acid-base
titrations . . . . . . . .. 76

3 RESULTS AND DISCUSSION . . . . . . . 77

Excited State Behavior in Aqueous Media. . . .. .77
6-methoxyquinoline. . . . . . . ... 77
Steady-state fluorescence . . . .. .77
Time-resolved fluorescence . . . .. .92
2-, 3- and 4-hydroxybiphenyl . . . .. .103
Steady-state fluorescence . . . .. .103
Time-resolved fluorescence. ..... . 118
Excited State Behavior in Mixed Aqueous-Organic
Media . . . . . . . . .. . . 124
Cosolvent Effects on the Forward Rate
Constant. . . . . . . . .. .124
6-methoxyquinoline. . . . . ... 124
2-hydroxybiphenyl .. . . . .. .129
Cosolvent Effects on the Reprotonation
Rate Constant. . . . . . . . ... 183

4 CONCLUSIONS . . . . . . . . ... .213

REFERENCES . . . . . . . . ... . .218

BIOGRAPHICAL SKETCH . . . . . . . ... 230














LIST OF TABLES

Table 1.1 Acid-base properties of 6-
methoxyquinoline in the lowest excited
singlet (S,) state . . . . . .. 61

Table 3.1 The Stern-Volmer constant, relative
quantum yield, lifetime estimates, and
hydroxide ion quenching rate constant of
neutral 6-methoxyquinoline . . . .88

Table 3.2 Experimentally determined constants,
lifetime and calculated hydrolysis and
reprotonation rate constants of 6-
methoxyquinoline in the lowest excited
singlet (Sj) state . . . . . .. .93

Table 3.3 Biexponential emission decay times and
chi-squared values of 6-methoxyquinoline
as a function of hydroxide ion
concentration. The excitation and emission
wavelengths are 280 and 350 nm,
respectively . . . . . . .. 95

Table 3.4 Maxima of the longest wavelength
absorption bands ('Lb and 'La) and
fluorescence bands of the neutral
molecules (N) and anions (A) derived from
the hydroxybiphenyls. Also presented are
their fluorescence decay times in the
absence of proton transfer. . . .. .104

Table 3.5 Rate constants (ka and kb) for prototropic
dissociation and reprotonation of the
hydroxybiphenyls (HBPs) in the S, state
and dissociation constants in ground
(pKso) and lowest excited singlet (pKs,)
states . . . . . . . . 116

Table 3.6 Results of the biexponential fitting of
the plots of (<'/c'o)/(P/po) vs the ethanol
mole fraction for 6-methoxyquinoline. . 127








Table 3.7 Results of the regression analysis of the
plots of log(/'/1'0)/(/90o) vs the ethanol
mole fraction for 6-methoxyquinoline. . 128

Table 3.8 Results of the regression analysis of the
plots of log(p'/p'0)/(q/Io) vs the mole
fraction of organic cosolvent for 6-
methoxyquinoline. . . . . . ... 131

Table 3.9 Results of the biexponential fitting of
the plots of (d'/!'o)/(t/40) vs the mole
fraction of organic cosolvent for 2-
hydroxybiphenyl . . . ... . .134

Table 3.10 A comparison of the inhibitory values; m,
obtained from the regression analysis of
the plots of log(7'/I'o)/(#/#o) for 6-
methoxyquinoline and m,, obtained from the
biexponential fitting of the plots of
()'/!'.)/(#/o.) for 2-hydroxybiphenyl as a
function of the organic cosolvent mole
fraction. . . . . . . . . 136

Table 3.11 The hydration 'sphere-of-action' radii of
6MQ and 2HBP in each of the mixed solvent
systems investigated . . . .. 144

Table 3.12 Results of the regression analysis of the
plots of log(q'/( 'o)/(9/po) vs -log(a,)
of 6-methoxyquinoline and 2-
hydroxybiphenyl . . . . . .. 154
Table 3.13 Results of the biexponential fitting of
the plots of Yi vs the mole fraction of
organic cosolvent for 2-hydroxybiphenyl 159

Table 3.14 The lifetime of neutral 2-hydroxybiphenyl
as a function of organic cosolvent mole
fraction. . . . . . . . .. 161

Table 3.15 Results of the biexponential fitting of
the plots of k, vs the mole fraction of
organic cosolvent for 2-hydroxybiphenyl 169

Table 3.16 The hydration sphere-of-action radii of
2HBP in mixed solvent systems . . .. .170

Table 3.17 Results of the linear regression analysis
of the plots of log(k) as a function of
-log(a,) of 2-hydroxybiphenyl . . .. .177








Table 3.18 Linear regression analysis of the
published water activity against the mole
fraction. . . . . . . . ... 182

Table 3.19 Debye-HUckel terms (A and B) at different
dielectric strengths (E) in methanol-water
solutions along with the products of the
rate constants and corresponding decay-
times kto, calculated from equation
(1.18) and the reciprocal of the intercept
of equation (1.17) and kb,', determined
from the product of kao and the slope of
equation (1.17) . . . . . ... .193

Table 3.20 Debye-Huckel terms (A and B) at different
dielectric strengths (E) in ethanol-water
solutions along with the products of the
rate constants and corresponding decay-
times krTo, calculated from equation
(1.18) and the reciprocal of the intercept
of equation (1.17) and kbt', determined
from the product of kt, and the slope of
equation (1.17) . . . . . . 194

Table 3.21 Debye-HUckel terms (A and B) at different
dielectric strengths (E) in isopropanol-
water solutions along with the products of
the rate constants and corresponding
decay-times kro, calculated from equation
(1.18) and the reciprocal of the intercept
of equation (1.17) and kbT', determined
from the product of ktg and the slope of
equation (1.17) . . . . . . 195

Table 3.22 Debye-HUckel terms (A and B) at different
dielectric strengths (E) in t-butanol-
water solutions along with the products of
the rate constants and corresponding
decay-times kTo, calculated from equation
(1.18) and the reciprocal of the intercept
of equation (1.17) and kbt', determined
from the product of kTo and the slope of
equation (1.17) . . . . . . 196

Table 3.23 Debye-HUckel terms (A and B) at different
dielectric strengths (E) in acetonitrile-
water solutions along with the products of
the rate constants and corresponding
decay-times kTo, calculated from equation








(1.18) and the reciprocal of the intercept
of equation (1.17) and kbTt, determined
from the product of k.Tr and the slope of
equation (1.17) . . . . . . 197

Table 3.24 Debye-HUckel terms (A and B) at different
dielectric strengths (E) in dioxane-water
solutions along with the products of the
rate constants and corresponding decay-
times kTo, calculated from equation
(1.18) and the reciprocal of the intercept
of equation (1.17) and kbTo, determined
from the product of kTo and the slope of
equation (1.17) . . . . . . 198

Table 3.25 pH of dilute sulfuric acid solutions in
aqueous methanol as a function of the
logarithm of the reciprocal hydrogen ion
formal concentration (-log[H*]) and the
mole fraction of methanol (X) . . .. .207

Table 3.26 pH of dilute sulfuric acid solutions in
aqueous ethanol as a function of the
logarithm of the reciprocal hydrogen ion
formal concentration (-log[H*]) and the
mole fraction of ethanol (X,) . . .. .208

Table 3.27 pH of dilute sulfuric acid solutions in
aqueous isopropanol as a function of the
logarithm of the reciprocal hydrogen ion
formal concentration (-log[H']) and the
mole fraction of isopropanol (X,) . . 209

Table 3.28 pH of dilute sulfuric acid solutions in
aqueous t-butanol as a function of the
logarithm of the reciprocal hydrogen ion
formal concentration (-log[H+]) and the
mole fraction of t-butanol (Xt) .... .210

Table 3.29 pH of dilute sulfuric acid solutions in
aqueous acetonitrile as a function of the
logarithm of the reciprocal hydrogen ion
formal concentration (-log[H']) and the
mole fraction of acetonitrile (X,) . . 211

Table 3.30 pH of dilute sulfuric acid solutions in
aqueous dioxane as a function of the
logarithm of the reciprocal hydrogen ion
formal concentration (-log[H']) and the








mole fraction of dioxane (Xd) . . .. .212














LIST OF FIGURES


Figure 1.1 The chemical structures of 6-
methoxyquinoline (6MQ), 2-hydroxybiphenyl
(2HBP), 3-hydroxybiphenyl and 4-
hydroxybiphenyl (4HBP) . . . .. 60

Figure 2.1 Time-correlated single photon counting
instrument schematic . . . . .. 71

Figure 3.1 The fluorescence spectra of the acidic (A)
and basic (B) species of 6-
methoxyquinoline. The excitation
wavelength was 289 nm. .. . . . .78

Figure 3.2 The RFI of the acidic (0) and basic (0)
species of 6-methoxyquinoline as a
function of pH . . . . . . .. 79

Figure 3.3 Extrapolation of fluorescence maxima for
6-methoxyquinoline . . . . . .. .81

Figure 3.4 The observed RQY of the acidic (0) and
basic (0) species of 6-methoxyquinoline as
a function of pH. The theoretical RQY
(-- ) for the acidic and basic species
were generated with equations (1.21) and
(3.2), respectively. . . . . ... 82

Figure 3.5 Variation of 9,/9 for the neutral species
of 6-methoxyquinoline as a function of
F[OH-] . . . . . . . ... 86

Figure 3.6 The variation of the relative quantum
yield ratio (q/po)/(1'/9j) of 6-
methoxyquinoline as a function of F[OH-] 91

Figure 3.7 The monoexponential emission decay
profile of the acidic (upper) and basic
(lower) species of 6-methoxyquinoline
at pH 7.0. . . . . . . .. .96





Figure 3.23


The slopes (m and m,) obtained for 6-
methoxyquinoline (0) and 2-hydroxybiphenyl


Figure 3.8 A biexponential emission decay profile of
the basic species of 6-methoxyquinoline,
taken at pH 11.6 . . . . . ... 97

Figure 3.9 A plot of (yi + Y2) vs YIy2 . . .. 100

Figure 3.10 A plot of (Yi + Y2) vs F[OH-] . . .. .102

Figure 3.11 Variation in the relative quantum yield
of 2-hydroxybiphenyl as a function of pH .105

Figure 3.12 Variation in the relative quantum yield
of 3-hydroxybiphenyl as a function of pH .106

Figure 3.13 Variation in the relative quantum yield
of 4-hydroxybiphenyl as a function of pH .107

Figure 3.14 Variation of ,o/+ of 4-hydroxybiphenyl as
a function of F[H] . . . . .. .110

Figure 3.15 Variation of 0,/+ of 4-hydroxybiphenyl as
a function of F[HSO;] . . . . .. 111

Figure 3.16 Variation of (io/9)/(4'/g'o) in 2-
hydroxybiphenyl as a function of F[H*] . 113

Figure 3.17 Variation of ( o/ )/(q'/I'o) in 3-
hydroxybiphenyl as a function of F[H'] . 114

Figure 3.18 Variation of ( o/ )/(4'/ 'o) in 4-
hydroxybiphenyl as a function of F[H*] . 115

Figure 3.19 Plot of log(7'/'O0)/(o0/0) vs mole fraction
of ethanol for 6-methoxyquinoline . .. 126

Figure 3.20 Plots of log(q'/p'o)/(9o/p) as a function
of the organic cosolvent mole fraction
(Xoc) for 6-methoxyquinoline. . . ... 130

Figure 3.21 Plots of log(9'/p'o)/(0p/q) as a function
of the organic cosolvent mole fraction
(XoR,) for 2-hydroxybiphenyl . . ... 132

Figure 3.22 Plots of log( '/'/0)/(o0/p) as a function
of the organic cosolvent mole fraction
(XoG) for 2-hydroxybiphenyl . . ... 133











Figure 3.24




Figure 3.25





Figure 3.26




Figure 3.27





Figure 3.28




Figure 3.29





Figure 3.30





Figure 3.31


The observed rate constant (yj) as a
function of the organic cosolvent mole
fraction for 2-hydroxybiphenyl in aqueous
methanol -0-, ethanol -0-, isopropanol
-a-, and tertiary butanol -o- . . .


157


The observed rate constant (y,) as a
function of the organic cosolvent mole
fraction for 2-hydroxybiphenyl in aqueous
acetonitrile -0-, dimethylsulfoxide -0-,
dioxane -a-, and formamide -o-. . . 158

The lifetime of 2-hydroxybiphenyl in the
absence of proton transfer against the
organic cosolvent mole fraction in aqueous
methanol -0- and aqueous ethanol -a-,
based upon equation (3.22). Solid circles
and triangles were obtained from equation
(3.23). . . . . . . . . . 165

xiii


(0) in aqueous alcoholas a function of the
molecular volume (VoRG) of alcohol . . 139

The sphere-of-action radii (R) of 6-
methoxyquinoline (D) and 2-hydroxybiphenyl
(o) in aqueous alcohol as a function of
the bulk dielectric constant. . . .. .146

Plots of log(/'/9'o)/(o0/9) as a function
of the -log(a,) for 6-methoxyquinoline in
aqueous methanol -0-, aqueous ethanol -0-,
aqueous isopropanol -a-, and aqueous
tertiary butanol -o-. . . . ... .150

Plots of log(7'/ ',)/(o,/ ) as a function
of the -log(a.) for 6-methoxyquinoline in
aqueous acetonitrile -0-, aqueous
formamide -0-, and aqueous dioxane -a-. 151

Plots of log(q'/q'o)/(Io/p) as a function
of the -log(a,) for 2-hydroxybiphenyl in
aqueous methanol -0-, aqueous ethanol -0-,
aqueous isopropanol -a-, and aqueous
tertiary butanol -o-. . . . ... .152

Plots of log(p'/p',)/(p ,/) as a function
of the -log(a,) for 2-hydroxybiphenyl in
aqueous acetonitrile -0-, aqueous
formamide -0-, and aqueous dioxane -a-. 153








Figure 3.32





Figure 3.33





Figure 3.34




Figure 3.35





Figure 3.36




Figure 3.37





Figure 3.38





Figure 3.39





Figure 3.40


Log(k) as a function of organic cosolvent
mole fraction (XORG) for 2-hydroxybiphenyl
in aqueous methanol -o-, aqueous ethanol
-0-, aqueous isopropanol -a- and aqueous
tertiary butanol -o- . . . . .. 167

Log(k) as a function of organic cosolvent
mole fraction (XoRC) for 2-hydroxybiphenyl
in aqueous acetonitrile -o-, aqueous
dimethylsulfoxide -0-, aqueous formamide
-A- and aqueous dioxane -o- . . .. 168

Plots of log(k) vs -log(a,) for 2-
hydroxybiphenyl in the aqueous alcohol
series; methanol -o-, ethanol -0-,
isopropanol -a- and tertiary butanol -o-. 175

Plots of log(k) vs -log(a,) for 2-
hydroxybiphenyl in aqueous acetonitrile
-0-, aqueous dimethylsulfoxide -0-,
aqueous formamide -A- and aqueous dioxane
-o-. . . . . . . . . . 176

Published water activity against the
organic cosolvent mole fraction, in -o-
methanol, -0- ethanol, -A- isopropanol and
-o- tertiary butanol. . . . . ... 181

Variation of the relative quantum yields
of fluorescence of the conjugate base
(('/I'o) species of 2-hydroxybiphenyl with
the formal hydrogen ion concentration [H']
in aqueous methanol .. .. . .. .. .184

Variation of the relative quantum yields
of fluorescence of the conjugate base
(9'/4'o) species of 2-hydroxybiphenyl with
the formal hydrogen ion concentration [H']
in aqueous ethanol. . . . . ... 185

Variation of the relative quantum yields
of fluorescence of the conjugate base
( '/e'o) species of 2-hydroxybiphenyl with
the formal hydrogen ion concentration [H']
in aqueous isopropanol. . . . ... 186

Variation of the relative quantum yields
of fluorescence of the conjugate base
(4'/('o) species of 2-hydroxybiphenyl with

xiv








the formal hydrogen ion concentration [H']
in aqueous t-butanol. . . . . ... 187

Variation of the relative quantum yields
of fluorescence of the conjugate base
(9'/4'o) species of 2-hydroxybiphenyl with
the formal hydrogen ion concentration [H*]
in aqueous acetonitrile . . . ... .188


Figure 3.41





Figure 3.42





Figure 3.43


Figure 3.44


Figure 3.45


Figure 3.46


Figure 3.47


Figure 3.47


Plots of (~0o/)/('/~' o) vs F[H*] for 2-
hydroxybiphenyl in aqueous methanol .

Plots of (Vo,/)/('/4'o) vs F[H'] for 2-
hydroxybiphenyl in aqueous ethanol. .


189


. 199


. 200


Plots of (4o,/)/()'/4'o) vs F[H*] for 2-
hydroxybiphenyl in aqueous isopropanol. 201


Plots of (q,/9)/(g'/9' ) vs F[H'] for 2-
hydroxybiphenyl in aqueous t-butanol.


202


Plots of (0o/p)/(q'/4'o) vs F[H'] for 2-
hydroxybiphenyl in aqueous acetonitrile 203


Plots of ( o/ )/('/'o) vs F[H'] for 2-
hydroxybiphenyl in aqueous dioxane. .


.204


Variation of the relative quantum yields
of fluorescence of the conjugate base
(q'//'o) species of 2-hydroxybiphenyl with
the formal hydrogen ion concentration [H']
in aqueous dioxane. . . . . . .














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

PROTON TRANSFER AND OPERATIONAL PH
IN AQUEOUS-ORGANIC SOLVENT SYSTEMS

By

Robert W. Townsend

December, 1992

Chairperson: S. G. Schulman
Major Department: Pharmaceutics

The lowest excited singlet state properties of 3

isomeric hydroxybiphenyls and 6-methoxyquinoline were

examined in pure water and mixed aqueous-organic cosolvent

systems with a combination of steady-state and time-resolved

fluorescence techniques. The results indicate that each of

these probes displays proton transfer in the lowest excited

singlet state in both the aqueous and mixed aqueous-organic

media. One finds that the pseudo-first-order acid

dissociation (or base hydrolysis) reaction is strongly

solvent dependent, yet independent of the acidity (or

basicity) of the solution. As the mole fraction of organic

cosolvent is increased, the rate constant decreases

biexponentially, eventually going to zero. The slope of the

logarithm of the dissociation (or hydrolysis) rate constant

against the mole fraction of organic cosolvent was found to








be proportional to the ratio of the diffusion volume of the

excited probe to the molecular volume of the organic

cosolvent. The logarithm of the dissociation (or hydrolysis)

rate constant against the logarithm of the water activity

can be used to characterize water structure and determine

the number of water molecules that participate in the

reaction at any given solvent composition. The second order

reprotonation reaction, on the other hand, depends

predominantly on solution acidity. In addition, one also

finds that the back reaction depends upon the continuum

properties of the solvent. Fortunately, steady-state and

time-resolved fluorescence methodologies allow reprotonation

to be treated independent of dissociation. The linear

relationship between the ratio of the relative fluorescence

efficiencies of the acid and conjugate base and the hydrogen

ion concentration is obtained if proper Brensted activity

factors are included in the relationship. These factors can

be calculated from classical electrostatics and are simple

powers of the activity coefficients necessary to convert the

formal hydrogen ion concentration to hydrogen ion activity.

Using this approach, it is possible to calculate pH from

hydrogen ion concentration in aqueous organic solutions.














CHAPTER 1
INTRODUCTION


Mixed aqueous-organic solvent systems are widely used

in the pharmaceutical sciences. They alter chemical

properties (stability and solubility) and therefore are

frequently used as vehicles to affect stabilization,

dissolution, dispersion and delivery of drugs. In addition,

mixed solvent systems are utilized as reaction media for

synthesis and as extraction media or mobile phase to aid in

the isolation and separation of complex mixtures of

compounds. However, much of the current knowledge regarding

mixed solvent systems is still subject to speculation.

The kinetics of excited state proton transfer are

affected by solvent composition in ways which reflect the

solvations of the reactants and products and the influences

of their structures on ionic or molecular diffusion.

If these relationships can be evaluated accurately, a

greater understanding of the nature of solute-solvent and

solvent-solvent relationships as well as a better

predictability of properties of solutes (drugs) in aqueous,

mixed aqueous-organic and more complex biological systems

may result.








2

In the past, excited state proton transfer has been

utilized as a simple means of probing solvation and solvent

structure, in aqueous and mixed aqueous-organic solvent

systems. For example, in 1963, using a steady-state kinetic

approach to excited state proton transfer, Urban and Weller

(1) examined the behavior of acidic probes in mixed

solvents. In their investigation of proton transfer from

ammonium pyrene derivatives in methanol, ethanol, n-propanol

and n-butanol, they found that the dissociation rate

constants decreased concurrently with temperature and that

this decrease was dependent upon the chemical nature of the

solvent.

The proton transfer process was depicted as follows in

scheme I:

k12 k23

AHz + L _= A'-1 HL* = AZ- + HL* Scheme I

k21 k32

where the acid (AH) with a valence (Z) donates a proton to

the accepting solvent (L). Resolvation of the probe was

considered to occur concurrently with the proton transfer

process because of a change in the charge of the reacting

species. It was also felt that the reduction in the

dissociation rate constant with temperature was due to a

retardation in the resolvation process based upon the

observed similarities that were found between the measured










activation energies for proton transfer and dielectric

relaxation in each of the solvents investigated.

In 1965, Trieff and Sundheim (2) determined the excited

state proton transfer rate and equilibrium constants for the

weakly acidic 2-naphthol (2N) in increasing amounts of

organic cosolvent using the steady-state technique. The rate

of deprotonation decreased with increasing amounts of

organic cosolvent yet no change was found in the

reprotonation process. The latter observation contradicted

the theoretical predictions of Weller (3) and Williamson and

La Mer (4) for diffusion controlled reactions.

The observed results were rationalized in terms of a

model which involved a preliminary ionization step, followed

by reaction of the ion with water. It was concluded that the

bulk dielectric constant must play a role in the ionization

process of 2N because in the solvents investigated, dioxane,

on a molar basis, exhibited a greater inhibitory effect on

the deprotonation process than did glycerol or methanol.

In 1978, Huppert et al. (5), with the help of a

picosecond spectroscopic technique, time-resolved emission

spectroscopy (TRES) and complimentary steady-state

fluorescence spectroscopy, measured the rate of proton

dissociation from 8-hydroxy-1,3,6-pyrene trisulfonate (HPTS)

in mixed aqueous ethanol solutions. They found from

semilogarithmic plots of the rates of proton dissociation as








4
a function of the mole fraction of ethanol that proton

dissociation decreases sharply as the ethanolic mole

fraction is increased and speculated that the rate of proton

transfer was directly influenced by the ability of water to

form clusters. They cited the prototropic conductivity

experiments by Erdey-Gruz and Lengyel (6) which give

support to a cluster order in water and referred to the

theoretical quantum mechanical calculations of Kaemer and

Diereksen (7), along with the ion clustering experiment

conducted by Kebarle (8), for the existence of water dimers

as an energetically favored species of water relative to the

monomer species.

In 1981, Huppert et al. (9) extended studies in mixed

solvent systems by examining the excited state proton

transfer of 2N, 2-naphthol-6-sulfonate (2N6S) and HPTS in

aqueous alcohol systems. Semilogarithmic plots 'of the

dissociation rate constant versus the alcoholic mole

fraction were found to be linear with slopes, independent of

the probe, increasing in the following order: methanol,

ethanol and n-propanol. Since the probes in this study

(i.e., weak acids with pK' > 0) did not exhibit excited

state behavior in the pure organic solvents and since the

bulk dielectric constant could not account for the reduction

in proton transfer rate, they felt that the dependence of

proton transfer in mixed solvents was contingent solely upon








5

the properties of water and reasoned that a structural model

of water in which diffusion of protons is mediated through

rapid exchange of hydrogen bonds-covalent bonds throughout

the quasistatic continuous network of hydrogen bonds must be

responsible for proton transport. Support for a structural

model came from Stillinger (10) who had concurred that the

distinct structural features and properties of liquid water

are related to the shape of the individual water molecules

and their tendency to hydrogen bond with each other.

Further, the x-ray and neutron diffraction studies on water

conducted by Narten and Levy (11), the subpicosecond

experiments on the photoejection of electrons in aqueous

solutions by Wiesenfield and Ippen (12) and the theoretical

studies conducted by Lie and Clementi (13) as well as those

from Stillinger and Rehman (14) provided strong evidence in

support of a local tetrahedral ordering of watdr, primarily

through hydrogen bonds. In addition, the dynamic flow

studies of Searcy and Fenn (15) demonstrated the existence

of H'(H0O)n clusters which vary in size from one to more than

twenty water molecules. Finally, the quantum mechanical

calculations of Newton and Ehrenson (16) in addition to

those of Kraemer and Diereksen (7) along with the results of

Kebarle (17) provided thermochemical evidence in support of

the proton hydration process.

With this background, Huppert and coworkers considered








6

that in the final state of proton transfer, the excited

anionic species and a proton would exist hydrated in a

complex of n water molecules. Further, since the free energy

of the reaction is dependent on the enthalpy of proton

hydration and since a linear relationship between the

dissociation rate constant and the mole fraction of organic

cosolvent was observed, they concluded that the addition of

organic solvents to the water network must disrupt the

hydrogen bonding network of water, leading to a decrease in

water cluster size, and that such behavior directly affects

the rate of proton transfer, possibly by delaying the rate

of proton hydration.

In 1981, Schulman and Vogt (18) investigated the proton

transfer kinetics of 2- and 4-quinolone in moderately

concentrated acid solutions using steady-state fluorimetry.

By assuming that the rate of dissociation was limited by the

rate at which the dissociating acid was hydrated to form the

transition state (or the solvent-separated ion-pair which is

the encounter complex of the fully hydrated proton and the

fully hydrated conjugate base), they found that they could

quantitate excited state proton transfer. This model was

later extended to the excited state investigation of

molecules containing either neutral or monocationic

conjugate bases and monocationic or dicationic conjugate

acids, respectively (19,20). The approach employed a








7
relationship between the relative quantum yield ratios of

the conjugate acid and base with the Hammet acidity function

hand the hydration requirements for the reaction of

interest and the dissociation of the Hammet indicator. They

found that the hydration requirements in concentrated acid

solutions for proton transfer ranged from 1 to 8 water

molecules. To explain this observation, it was suggested

that one or more of the reactants may not be in its maximum

state of hydration because free water in concentrated acids

is scarce.

In 1984, Huppert et al. (21) investigated the rate of

proton transfer of 2N, 2N6S, 2-naphthol-3,6-disulfonate

(2N36DS), 2-naphthol-6,8-disulfonate (2N68DS) and HPTS in

concentrated electrolyte solutions at pH >> pK', where pK"

is the pK of the probe in the excited state. In this study,

they found that the dissociation rate constant'was linearly

related to the chemical potential of water and the following

relationship was found to hold


log(k) log(k.) rlog(a,) (1.1)


where k and k, are the rate constants for excited state

deprotonation in the salt solutions and in pure water,

respectively. The slope r is an empirically determined

constant, which varied from 6-10 depending on the probe, and

a, is the activity of water. It was observed that probes








8

containing hydrophilic moieties in proximity to the

dissociating proton yielded smaller values for r. Thus, r

was deemed the probability of proton transfer from the

excited probe to nearby water molecules recognized as being

present at different degrees of aggregation and ordered at a

proper distance from the dissociating proton. Further, it

was reasoned that if proton transfer was considered to occur

inside a coulombic cage, then the decrease in proton

transfer with increasing salt concentration must be related

to ion-pair formation (a process in which the O-H bond of

the excited probe forms a hydrogen bond with the nearest

molecule which itself is hydrogen bonded to nearby water

molecules). However, according to Eigen (22) a combination

of electrostatic interactions and reduced proton

diffusibility, effects which dominate ion-pair recombination

and ion-pair separation, would accelerate the rhte of proton

dissociation. Thus, to support their ion-pair concept, they

cited Conway (23), who attributed the enthalpy difference

between proton hydration and hydronium ion formation to a

further solvation of the hydronium ion by additional

molecules of water. Finally, with the arguments developed

above as well as those from their previous publications

(5,9,21) and the experimental evidence presented, they

claimed that the hydrated complex of a proton in dilute

electrolyte solutions varies in size, but that it is








9
composed of no more than 10 molecules of water.

In biological investigations, Huppert and coworkers

(24,25) attempted to characterize the properties of water

within poorly hydrated binding sites by making use of the

linear relationship that was found between the dissociation

rate constant and the activity of water in electrolyte

solutions (equation (1.1)). From these studies they reported

that the ability of water to hydrate dissociating protons

within the binding sites of bovine serum albumin and

apomyoglobin was equivalent to the ability of water to

hydrate protons in salt solutions where water had activities

of 0.8 and 0.68, respectively.

In 1984, Bardez et al. (26) tried to extend this

relationship to reverse micelles using phase resolved and

steady-state fluorimetry. Preliminary studies on the proton

transfer of pyranine suggested that the activity of

solubilized water in reverse liposomes was in good agreement

with both the activities yielded by measurement of vapor

pressures (27) and those calculated from the theoretical

electrostatic model of J6nsson and Wennerstrom (28).

However, in a subsequent investigation with 2N and its

sulfonated derivatives (2N6S and 2N68S), Bardez et al. (29)

found that both the efficiency and kinetics of proton

transfer were strongly dependent upon the localization of

the probe within the micelle. It was concluded that the








10

validity of Huppert's method for determining the activity of

water in reversed micelles should be used with some degree

of caution.

Recently, Politi and Chaimovich (30) examined the rate

of proton dissociation in aqueous urea solutions with HPTS

and found that the forward rate constant could not be

related unequivocally with the activity of water. Instead, a

urea-water cluster with properties similar to pure water was

proposed to account for the observed behavior. It was

concluded, that since the effects of the groups at the

interface on the properties of surrounding water are

unknown, that any information obtained by proton transfer or

related studies which probe the properties of an interface

must be regarded with great care.

In 1984, in an ongoing attempt to more fully comprehend

the proton transfer process, Huppert et al. (31) carried out

picosecond studies of HPTS in both water and deuterated

water at pressures up to the ice transition point. They

found that proton dissociation was dependent upon a specific

water structure, in support of previous investigations

(5,9,21) and proposed the following model. The first step in

proton transfer was considered to be the transfer of a

proton from an excited probe to an adjacent properly aligned

water molecule forming H,30. Such behavior would affect the

local solvent environment, which would then lead to further








11

structural changes and increases in the proton hydration

cluster size. They then suggested that successive proton

transfers between large labile hydration complexes would

increase the distance (r) between the excited anion and the

proton to the point where the Coulomb field interaction or

R, (the "effective range" of the electrostatic interaction

potential between the excited anion and the proton) is

smaller than the thermal energy, and that under these

conditions, the proton will escape. They concluded that the

rate-determining step in proton transfer for the proposed

model must be the rate at which the proton is transferred to

the adjacent properly aligned water molecule, and that this

step ultimately depends upon the donor-acceptor distance.

In 1985, Robinson and coworkers (32) examined the

proton transfer of 2N in the pH independent portion of the

titration curve (pH >> pK', where the dissociation process

dominates at differing temperatures) in methanol-water

mixtures using the picosecond spectroscopic technique, time

correlated single photon counting accompanied by steady-

state fluorescence spectroscopy. They observed an increase

in the proton transfer rate with increasing temperature

which could not be explained by the temperature dependence

of pH as described by Kolthoff and Rosenblum (33) because of

the initial experimental conditions. They also found that as

the alcoholic component of the reaction mixture was








12

increased, the effect of temperature on the proton transfer

rate decreased. This behavior coincided with observations

made earlier by Urban and Weller (1). They also cited that

the rate of proton transfer decreased as the methanol

concentration increased which served to validate previously

claims (5,9). From this study, Robinson and coworkers

concluded that although solvent structure plays a role in

the proton transfer, the water cluster model (H30,[H,0]n, n >

1) proposed by Huppert et al. (9,21) fails to provide

adequate details about the cluster size or its specific

relationship to proton kinetics. In an effort to provide

additional information about the proton transfer process,

they adapted a Markov random walk theory previously used in

the kinetic analysis of electron transfer (34-36). Results

indicated that a 4 + 1 member water cluster was the primary

proton acceptor. In support, they cited Eigen and coworkers

(37-39), who suggested that the condition for proton

migration is a disturbed structure of water represented by a

hydrogen complex which is comprised of exactly four water

molecules.

Finally, the observed activation energy for proton

transfer in pure water (3.45 Kcal/mole) was attributed to

the energy required to reorganize the hydrogen bonded water

network to form the proton accepting cluster. From this they

implied that a specific structure of the critical 4 + 1








13

molecule water cluster apparently determines the rate of

proton transfer in aqueous media. Any additional water or

cosolvent molecules must exist in secondary coordination

shells and therefore, they only play secondary roles in the

proton transfer process.

In 1986, Shizuka et al. (40) investigated the proton

dissociation reaction of 1- and 2-naphthol in the presence

of NaCl at 300 K by means of nanosecond and picosecond

spectroscopy with steady-state fluorimetry. The observed

proton dissociation rates of 1- and 2-naphthol decreased

considerably upon addition of NaCI, in agreement with

Huppert et al. (21). It was found that k,/k?, when plotted

as a function of NaCI concentration, decreased linearly and

that this behavior could be described by the following

equation


k, k(l a[NaCl]) (1.2)


where ki and k, represent the proton dissociation rate

constants with and without NaCl and a denotes a rate

parameter which is determined from the slope of the line.

This equation holds for a[NaCl] s 1, and a was found to be

0.2 for the naphthols.

Shizuka et al. (40) implied that proton transfer from

naphthols to water clusters is a second order process.








14

Consequently, the rate constant in pure water is given by


k k0[(H,0)n] (1.3)


where k, represents the second order rate constant in units

of M -s-1 and [(H,0)n] denotes the concentration of water

clusters. Evidence for a 4 member water cluster had been

provided (32,38,39). Further, Huppert et al. (21) had

suggested that protons produced in the dissociation process

become entrapped in clusters of up to 10 molecules of water.

They figured that it was reasonable to assume that the

destruction of proton-accepting water clusters by NaCl shuts

down proton transfer and that because of this, the proton

transfer rate can be expressed as

k, k[(H20)4](1 a[NaCl]) (1.4)


where a[NaCl] s 1.

Shizuka et al. (40) also speculated that the

concentration of water clusters could be estimated to be

13.8 M in pure water; and since a = 0.2, they infer that in

a 1.0 M NaCl solution 20% of the water clusters (2.7 M) were

destroyed, which leads to the production of 11 water

molecules that hydrate the Na' and Cl- ions. In their

investigation, they report that the average hydration

numbers for the Na' and Cl- ions were 4.4 and 6.6,








15

respectively. These values are comparable to literature

values obtained through x-ray diffraction studies (41,42).

Around the same time, Pines et al. (43) reported the

excited state behavior of the heterocyclic bases, acridine

and 6-methoxyquinoline in both methanol-water and

concentrated electrolyte solutions using steady-state and

time-resolved emission spectroscopy. The rate of

heterocyclic protonation was noted to decrease with

increasing methanol concentrations. However, the

investigators felt that the role of water on the rate of

proton abstraction could not be properly evaluated because

the probability of nonradiative decay S, -- So or triplet

formation SI--Ti is higher in the nonpolar organic solvents

than it is in water with heterocyclic compounds (44-46).

They turned instead to a series of concentrated electrolyte

solutions and found that the rates of protonation for the

heterocyclic bases increased with increasing salt

concentrations. From this observation, it was suggested that

the addition of strong electrolytes to water decreased the

concentration of free bulk water as a result of the ion

hydrating process. Therefore, they claimed that water

molecules that hydrate cations tend to be more acidic than

free bulk water because they increase the rate of proton

transfer.

In 1986, Robinson and coworkers (47-50) presented a








16

series of papers in support of a new proton hydration model

for excited state proton transfer. In this model, the direct

kinetic product of the acid dissociation is considered to be

H,0,4 ion. Further, they argue that photokinetic experiments

on ultrafast timescales suggest that the integrity of the

quasi-tetrahedral oxygen structure of liquid water controls

the proton hydration process. Support for this belief stems

from the early works of Kamb (51) and Triolo and Narten (52)

who inferred that a key to understanding the dynamic

processes in water may be the stability, on characteristic

timescales, of the five-membered quasitetrahedral

arrangement of oxygen atoms. With the oxygen atoms forming a

more or less static background, rapid reorientational

motions in the liquid phase occur primarily through the

motions of the hydrogen atoms. This implies that proton

transfer and other properties of liquid water, 'such as

dielectric relaxation, (53) shear viscosity (54,55) and

spin-lattice relaxation processes (56,57) are dependent upon

these rapid reorientational motions.

To help comprehend the temperature dependence of these

dynamic processes in pure liquid water, Robinson et al. (47)

developed an analytical approximation to a temperature

dependent rotational rate parameter K(T) which has the

pseudo-Arrhenius form











K(T) KO(T)exp(-AH,/RT)
(1.5)


where K is the experimental prefactor, AH, is the kinetic

molar enthalpy of activation, R is the gas constant and T is

the absolute temperature. They felt that the temperature

dependent rate parameter which is related to the Debye

relaxation time (rT = K(T)-1) simplifies the multitude of

vibrations and structures of water to a single set of

temperature-sensitive frequencies while maintaining much of

the anomalous physics associated with liquid water. However,

since most dynamical and transport properties associated

with liquid water do not exhibit typical Arrhenius behavior,

they turned to the mathematical models developed for

supercooled liquid water by Lang and LUdeman (56-58) and

Angell and coworkers (59-61) for support because they

approximate the thermodynamic and kinetic data over a wider

range of temperatures. The heat capacity data generated by

Oquini and Angell (61) appear to be of central importance in

support of their model.

The differences observed between the calculated heat

capacities and those obtained by Oquini and Angell (61) were

speculated to arise from an assumption used in the model

development which fails to include contributions from five

intermolecular vibrations (water exists as a simple rotor).








18

However, the level of agreement, which existed between the

two sets of data seemed to indicate that the temperature

dependence observed in the data could account for the

anomalies associated with liquid water. Further,

investigations on the temperature dependence of proton

recombination and proton-induced quenching for 2-naptholate

(50) showed that the dielectric relaxation times Td listed

by Hasted (53) are essentially proportional to the newly

developed pseudo-Arrehnius rate parameter over the

temperature range of interest. Therefore, they concluded

that their simplistic model provides a basis through which

the temperature dependent properties of water can be

understood, and that cooperative hydrogen bonding produces

the anomalous behaviors associated with water. They also

indicate that this cooperativity leads to an effective

hindering potential which is generated through the hindered

rotational motions of hydrogen bonds and that this behavior,

in itself, provides the foundations from which the kinetic

and thermodynamic properties associated with water in the

pure and binary states arise.

In a subsequent paper, Robinson et al. (48) derive a

set of absolute rate expressions for proton transfer

processes which incorporate the Debye relaxation time, T,,

and which correlate the dissociation and recombination rate

constants of weak acids with thermodynamic ionization








19

enthalpies (AH?) and entropies (AS ).

The dissociation and recombination rate constants are

defined as follows:


kds QTDexp(ASR/R)exp(-AHl,/RT) (1.6)



krc QT exp[ -(AHi, AH)/RT] (1.7)


where Q is a mobility/steric factor introduced by Eigen and

Kustin (62); Tr is the Debye relaxation time (which is

related to the temperature dependent rotational rate

parameter K(T), Tr = K') ; AS, and AH? are the thermodynamic

entropy and enthalpy values for acid ionization; and AH;: is

the kinetic enthalpy for acid dissociation. Standard-state

conditions require unit molarities throughout.

Support for the new rate expressions was based upon a

review of weak electrolytes by Harned and Owen (63) which

revealed that many weak univalent acids at room temperature

have enthalpies of ionization near zero leaving only

entropic barriers of -20 cal/mole to stand in the way of

acid dissociation. This behavior was considered to be

normal, and the group turned to the picosecond

investigations on 1- and 2-naphthol conducted by Webb et al.

(64) along with their previous work on the hydration

dynamics of protons from the photon initiated acids, 1-and








20

2-naphthol in water/alcohol mixtures (49), to suggest that

naphthols fit the weak acid criteria.

When proton dissociation and recombination rates are

described in terms of the Arrhenius equation, where AH =

[AH, AHfiec, Robinson and coworkers (48) show that the

free energy of dissociation (AG w, 0) for the naphthols is

consistent with Harned and Owen's weak acid picture. They

also show that calculations from the Arrhenius equation

provide additional evidence for a hydration model in which a

solvent rotational bottleneck stands in the way of the

overall ion hydration reaction. For example, the inverse of

basic rates, i.e., rates that the dissociation process would

have if it were unencumbered by an entropy or enthalpy of

activation, are roughly equal to the Debye correlational

times rT for pure HO and D20. This is consistent with

results obtained by Eigen et al. (62,65), who found that

recombination rates are dependent upon frictional diffusion,

a process which can be described by the Debye equation.

Robinson and coworkers are quick to point out that

Eigen's indirect methods for determining the dissociation

rate fails to provide adequate detail about ion hydration.

They infer that the mixed alcohol-water solvent systems used

throughout their studies serve as effective probes in the

ion hydration process because the sterically bulky alcohol

molecules require more time than is available in the excited








21

state to disentangle and reorient to surround the small

charge. Further, with their new analysis technique (34-36)

they indicate that the water cluster H,4O is the direct

kinetic product of acid dissociation. Reviews by Conway (23)

and Newton and Ehrenson (16) lend support to Hg90 as a

specific entity for the hydrated proton. In addition, recent

gas phase studies by Kebarle (66) provide firm

thermochemical evidence for successive binding of water

molecules (H,0)n to H30'. Further, the theoretical

calculations by Pople (67-69) and Clementi (70) appear to

confirm that these energy relationships exist. Robinson and

coworkers (48) conclude that the the hydration model

currently provides the best means through which ion

solvation can be understood.

In 1989, Lee (71) conducted picosecond studies of 2N in

various salt solutions at different temperatures in an

attempt to elucidate the hydration structure of an ion in an

electrolyte solution. Proton transfer was found to decrease

with increasing salt concentrations, consistent with the

results obtained earlier by Huppert et al. (21) and Shizuka

et al. (40). The relationship that Huppert used to correlate

the dissociation rate constant with the activity of water

was redefined in terms of the osmotic coefficient h of the

solutions,











kd = ks, exp[-nhEm v//55.56]
(1.8)


where k8, and kdi, are the rates of proton dissociation in

water with and without salt, n is the proton hydration

number, and m, and v, are the molality and the

stoichiometric number of the ions, respectively. Plots of

log (k,,,) versus log (a,) of 2N in NaCI at differing

temperatures were found to coincide with Huppert's work,

however, Lee questioned the validity of this model because n

was found to be solvent dependent and because the modified

osmotic equation gave dissociation rates that were greater

than observed rates in LiCd, MgCl, and CaCl2 solutions.

To help explain the dynamics between the proton

dissociation-hydration and the ion hydration processes

during the lifetime of the excited state, kinetic solvation

numbers as proposed by Taube (72) were generated for each

electrolyte system. Lee (71) inferred that in ion solvation

only the most loosely attached water molecules can detach

from the ion and reorient to hydrate the proton during the

lifetime of the excited state, and that this would

constitute an intermediate region which bridges the gap

between the strongly bound primary solvation shell and the

bulk region where the solvent properties are retained.

Retardation of the dissociation rate of 2N in increasing








23

concentration of salt thus results from the loss of free

water, i.e., water used in ion hydration.

Lee (71), in support of Pines et al. (43), suggests

that cations inhibit dissociation to a greater extent than

anions because they compete with protons for the same

negative part of the water molecule. Experimental evidence

in NaC1 and NaC0I3 solutions support this notion because the

only noticeable difference in the dissociation process of 2N

could be attributed to ionic quenching parameters. Thus, it

was concluded that kinetic solvation numbers provide a more

effective means through which the various salt effects on

the deprotonation process can be explained.

In 1990, Robinson and coworkers (73) extended the

hydration model to a more complicated weak acid system

(i.e., one in which intramolecular hydrogen bonding

interferes with proton transfer). Using picosecond laser

technology, they investigated the excited state behavior of

l-napthol-2-sulfonic acid in ethanol/water solvent systems

at different temperatures. When these results were compared

with the results obtained with l-naphthol (49), it was found

that the intramolecular hydrogen bonding in the sulfonate

derivative sterically reduced the extramolecular proton

dissociation and recombination rates. However, the overall

rate was still controlled by the reorientational motions of

the adjacent water molecules; and, if dissociation was to








24

occur, a cluster of 4 + 1 water molecules was required to

solvate the transferring proton. Thus, this investigation

provided additional support to their proposed hydration

model for endothermically produced protons in aqueous

environments.

Recently, Schulman and Kelly (74,75) investigated the

excited state behavior of 2-naphthol and some sulfonated

derivatives (2-naphthol-6-sulfonate, 2-naphthol-8-sulfonate,

2-naphthol-3,6-disulfonate and 2-naphthol-6,8-disulfonate)

in dimethyl sulfoxide/water and ethanol/water mixtures using

steady-state fluorimetry. It was found that semilogarithmic

plots of the relative quantum yield ratios (9 /90)/(9/9)

versus the organic mole fraction decreased linearly and that

this behavior could be fit to the following empirical

relationship


log( ) mXor log(kr0) (1.9)




where #'/9 and 0/o are the relative quantum yields of the

conjugate base and the acid, respectively; m is the slope

and X is the mole fraction of the organic solvent. Schulman

and Kelly suggested that the separation of the initial

ionization products by frictional diffusion was too slow to

allow dissociation to compete with recombination and that








25

the much faster proton jump (Grotthus) mechanisms, which

imply an extended hydrogen bonded network in the solvent,

are required for proton transfer to occur on the nanosecond

time scale. In addition, it was felt that the slope

represented some energetic parameter and since equation

(1.9) was similar to equation (1.1), which relates the

activity of water to the dissociation rate constant in

concentrated electrolyte solutions (21), it seemed

reasonable to equate the two equations through the following

relationship:


rlog(a), -mXorg (1.10)


Initial results in an ethanol-water system suggest that

equation (1.10) may be valid. For example, a plot of

log(krT) vs. log(aw) was linear over a narrow range of the

ethanol mole fraction. The slope, r, which was -designated as

the total number of water molecules involved in the proton

transfer reaction, was found to be 12.4. One would expect a

much smaller value of r (3 or 4) if only waters of hydration

(i.e., water molecules in the primary hydration spheres of

the products and reactants) were involved. They proposed

that the hydration behavior of acids in mixed solvents is

"water structure making," which is quite different from the

"water structure breaking" behavior of concentrated

electrolyte solutions, although reasons for this claim were










not well defined.

From practical and fundamental points of view, water is

arguably the most important component in a mixed solvent

system. Therefore, a basic understanding of the fundamental

physical, chemical and structural properties associated with

liquid water (which include solvation phenomena and what is

currently known about proton transfer) is required to

develop a better understanding of how cosolvents affect

proton transfer. This is because the composition and

structure of the molecules contained within the liquid

system will ultimately influence its behavior. Thus, these

subjects are briefly reviewed, before the concepts of

excited state proton transfer are developed.


Physico-Chemical Properties of Water

The dynamic properties of liquid water ultimately

depend upon the structure and properties of its individual

molecules. Spectroscopic analysis of water in the gas phase

has shown that water molecules are not linear. An angle of

105'03' exists between the H-O-H bonds, while the

internuclear 0-H and H-H distances are 0.9568 A and 1.54 A,

respectively (54). In addition, dielectric investigations of

water in the gas phase indicate that the isolated water

molecule has a permanent electric dipole moment of 1.84

Debye (76). These observations do not yield a complete








27

picture of the isolated water molecule and its accompanying

electronic charge distribution. However, with the aid of

electrostatic models, molecular orbital theory, electron

cloud distribution, wave functions and the calculation of

physical properties, such as those described above, more

precise descriptions of liquid water can be obtained.

Perhaps the most reliable picture of water is Bjerrum's

quadrupole model (77) in which eight ordered electrons move

along four elliptical orbitals with two electrons in each

orbital. The lobes of positive charge are directed along the

O-H axis and are located primarily in the region of the two

protons. The lobes of negative charge, essentially the two

lone pair orbitals of the oxygen, are directed above and

below the plane which contains both the oxygen and hydrogen

nuclei. Conceptually, water takes the shape of a tetrahedron

with its positive and negative charges being directed

towards the apices of this tetrahedron (78).

The unusual tetrahedral configuration of the isolated

water molecule (each oxygen atom sp3 hybridized) provides

liquid water with its unique ability to form four hydrogen

bonds. The two hydrogens of liquid water are capable of

donating hydrogen bonds, while each of the two lone pair

electrons of the oxygen atom are capable of accepting

hydrogen bonds (79). It has been suggested that the

anomalous properties of liquid water, such as the maximum








28

density at 4'C, the high dielectric constant, heat capacity,

viscosity, thermal conductivity and critical temperature,

may arise from the covalent characteristics of these

hydrogen bonds (80).

The structure of liquid water, that is, the relative

positions and motions of its individual molecules, has been

scrutinized for decades. The first major breakthrough came

in 1933, when Bernal and Fowler (78) inferred that a

tetrahedral grouping (oxygen located at the center of a

tetrahedron formed by four oxygen atoms, each 2.76 A away)

frequently occurs in the liquid as each water molecule is

hydrogen bonded to its four nearest neighbors. In support,

they cited that the structure of ice was partially retained

upon melting and that within small regions of the liquid,

water monomers remained in a lattice-like order.

Investigations of the radial distribution functions of water

with x-ray diffraction techniques have confirmed this

conclusion (76,81).

Since that time, various model systems have been

developed (for reviews see references 54,82-84). The most

popular are based upon the existence of structured water as

is readily seen in the "flickering cluster" model developed

by Frank and Wen (80). They reasoned that highly directional

cooperative covalent hydrogen bonds lead to the formation

and dissolution of 3-dimensional hydrogen-bonded water








29

clusters. Coexisting with the ice-like, short-lived clusters

which form and melt as a consequence of local energy

fluctuations is a dense fluid made up of nonhydrogen bonded

water molecules. According to Frank and Wen, the dissolution

and formation of clusters is the rate-limiting step in the

structural rearrangement of water. This is because

perturbations of the liquid water system, which increased

the population of hydrogen bonded molecules, would rigidify

the structure of water on the subnanosecond time scale. On

the other hand, any perturbation which would decrease

hydrogen bond formation would ultimately loosen water

structure by increasing the population of individual

nonhydrogen bonded water molecules. The dielectric

relaxation time (- 10'" sec) for liquid water (considered to

be the lifetime of a hydrogen bond) (53) was cited as

support for the short-lived clusters. Further, the uniform

activation energies (4.6 kcal/mole) for dielectric

relaxation (53), viscous flow (4.8 kcal/mole) and self-

diffusion (4.6 kcal/mole) (85,86), which correspond roughly

to the energy required to break and form hydrogen bonds,

provide additional support to the idea that clustering was

the rate limiting step in the structural rearrangement of

water. However, in a review of the relaxation processes of

water, Davis and Jarzynski (87) provide evidence which

clearly shows that the rate determining step in water's








30

structural rearrangement is not clustering, but the jumping

of individual water molecules from their temporary

equilibrium positions. Thus, Frank and Wen's "flickering

cluster" model and, for that matter any structural model of

liquid water, is at best an approximation of reality. For a

more accurate model, it becomes necessary to consider the

relative positions and motions of water's individual

molecules that arise in the process of thermal motion.

According to Eisenburg and Kauzmann (54), the characteristic

features of liquid water arise from two types of thermal

motions. The first is a rapid oscillation which occurs about

a temporal equilibrium position generally within a time-

frame of about 10-" sec. The second is a much slower

displacement of the equilibrium position itself which occurs

within a time frame of approximately 10-1 sec at O'C.

Unfortunately, current experimental techniques'are not

sensitive enough to provide the necessary details of water's

momentary structure. Instead, one obtains an averaged

picture of the immediate surroundings around an arbitrarily

chosen water molecule during a time period which includes

about a thousand molecular oscillations around a temporary

equilibrium position. To circumvent this limitation,

theoretical models like the Monte-Carlo (13,88,89) and

molecular dynamics approach (14,90,91) to liquid water have

been developed which tend to coincide with what has been








31

obtained experimentally. Consequently, water can be pictured

as a statistically ordered three dimensional network of

water molecules with different degrees of cooperative

hydrogen bonding which are ordered in a tetrahedral fashion,

complete with structural defects or vacancies.

Solvation

The process of solvation, where individually dissolved

solute molecules become encapsulated by a coordination shell

of more or less tightly bound solvent molecules, follows the

addition of a solute to water. The magnitude of structural

change that accompanies this process is not only dependent

upon solute size, but also upon the degree and types of

interactions which can occur between the solute and solvent

molecules. For example, interactions may be solute-solute,

solute-solvent or solvent-solvent in nature and may involve

the so-called directional, induction and dispersion forces

which are non-specific and unsaturable (i.e., ion-dipole,

dipole-dipole, dipole-induced dipole or instantaneous

dipole-induced dipole forces), or may involve the highly

directional, saturating forces which can lead to

stoichiometric molecular compounds (i.e., hydrogen bonds,

charge transfer and electron-pair donor acceptor forces). In

addition, solvophobic forces may play a role (92). Thus, the

process of solvation is extremely complicated.

Current knowledge of solvation stems primarily from








32

fundamental thermodynamics. The energy changes that

accompany this process are normally depicted by enthalpy

characteristics, while any structural changes that occur are

normally depicted by entropy characteristics. When combined,

these changes constitute the change in Gibbs free energy of

solvation, and therefore are representive of the chemical

nature of the solution. They are interrelated through the

following relationship:


AGsov = TH"asv TAS'so, (.11


Daudel (93) claims that the measure of a solvent's

solvation ability (i.e., aGsoiv) is directly related to four

distinct energy components that arise within the particular

solvent upon solute dissolution. These components are:

1) The cavitation energy or the energy required to make

a hole in the solvent system upon solute dissolution.

2) The orientation energy or the energy which arises

from the partial orientation of the dipolar solvent

molecules while in the presence of solvated solute

molecules.

3) The isotropic interaction energy which results from

such intermolecular forces as the electrostatic,

polarization and dispersion energies.

4) The anisotropic interaction energy or the energy

that accompanies specific interactions such as hydrogen








33

bonding or electron-pair donor acceptor bonds at well

localized points in the dissolved molecules.

Unfortunately, a measure of solvation energetic from the

Gibbs energy of solvation is not directly obtainable as it

involves the energy change associated with the transfer of a

solute from the gas phase into a solvent (93). However, it

can be calculated with theoretical models (for reviews, see

references 23,95-97). It is frequently replaced by the Gibbs

free energy of transfer (AG,), which is a measure of the

change in energy associated with the transfer of a solute

from water (the reference solvent) to another solvent system

(98,99).

Hydrocarbons

It has been suggested that the abnormal thermodynamic

properties that arise when hydrocarbons are added to water,

can be ascribed to an increase in the structuring of water

(100,101). The behavior of the tetraalkylammonium halides

and long chain organic salts in water, also appear to be

related to this phenomenon (80).

The large positive molar heat capacities and the large

negative heats and entropies of solution observed for

nonelectrolytes in water have been attributed to the

formation of water cages (100) or icebergs (101) around the

solute molecules. Frank and Wen (80) claim that nonpolar

solutes stabilize the ice-like clusters of water by








34

protecting them against the local energy fluctuations as

pictured in their "flickering cluster" model. N4methy and

Scheraga (102) suggest that nonpolar groups or molecules act

as a fifth coordination neighbor to the tetrahedrally

hydrogen-bonded water molecules. Thus, the weak Van der

Waals forces which accompany such processes lower the energy

of the associated water molecules, thereby stabilizing the

cluster.

Monohydroxylic Alcohols

When placed in water, the monohydroxylic alcohols, tend

to possess characteristics which are similar to hydrocarbons

because of their apolar groups. They show large negative

entropies and large positive molar heat capacities of

solution (103,104) and their negative partial molar volumes

in water increase with temperature and molecular size (105).

Further, the enthapies of transfer of alcohols'from H20 to

D,O, suggest that the apolar groups of the alcohols

structure water (106). Support for these concepts can be

found from the X-ray diffraction studies of ethanol (107)

and t-butyl alcohol (108) in water. Results indicate that

the ability of alcohols to structure water approaches a

maxima when the alcohol mole-fraction is approximately 10-12

%. Any additional alcohol, results in the disruption of

water structure; probably by breaking hydrogen bonds.

In addition to their nonpolar group, the monohydroxylic








35

alcohols possess a polar functional grouping (-OH) which

adds to their rather unique properties in water. Unlike

nonpolar molecules which tend to be forced out of solution

by virtue of their large negative entropies of hydration,

alcohols (at least the smaller ones) tend to remain in

solution (109). Further, their ability to form cooperative

hydrogen bonds leads to nonplanar structural arrangements of

alcohol molecules (110).

At one time it was felt that alcohol dimers dominated

the liquid structure of alcohols. However, Tucker and

Christian (111) have found that dimers are relatively

unimportant, even at low concentrations. Instead, evidence

now suggests that tetramers (112) and/or pentamers (113)

dominate their liquid structures. In fact, networks of

cooperative hydrogen-bonded alcohols may even exist (113-

115).

Recently, Krestov (116) has proposed that three

distinct structural regions exist in mixed aqueous-alcohol

solvent systems. In regions of high water concentration,

alcohol molecules can be seen as preferentially filling the

voids and interstitial spaces of water. Any displaced water

becomes incorporated into the ice-like network (117,118).

Thus, alcohols become incorporated into the tetrahedral

stucturing of water, at least up to the concentration of

alcohol which correlates with the maximum stabilization of








36

water structure (119). However, with continued addition of

alcohol to water, solvent structure in dynamic equilibrium

goes from predominately water in nature to predominately

alcohol in nature. Finally, at high concentrations of

alcohol, there appears to be interstitial insertion of water

molecules into the networks of alcohol structure (120),

which tends to preserve the stucturing of alcohols over a

wide range of concentration (120,121).

Ions

In the solvation of ions in water, where electrostatic

forces dominate, Frank and Wen (80), have considered that

three distinct regions of solvent exist. The first of these

regions consists of an ion surrounded by a coordination

sphere of strongly bound solvent molecules which tend to be

less mobile than the bulk solvent. At a distance well

removed from this centralized ion, bulk solvent exists in

its normal state which leaves an intermediate region of

highly mobile, disordered solvent molecules. In this way,

they were able to explain the "structure making" and

"structure breaking" properties of ions of different charge

and size in solution. Gurney (122), elaborated on the

concept of different regions around the dissolved ion when

he introduced the term cosphere for the zone surrounding a

spherical ion in which significant differences in structure

and properties are to be expected. Samoilov (123) found that








37

in some instances, the water molecules which surround a

dissolved ion are more mobile than those found in the bulk.

This is in contrast to the ordinary positive hydration of

small spherical ions, which possess structure making effects

on the solvent. This helped to explain the observation that

aqueous solutions of certain salts (i.e., KI) are more fluid

than pure water. It was suggested that when the exchange

frequency of water molecules in the solvent shell is greater

than that which occurs in the bulk, negative hydration

occurs. Such ideas support the structure breaking effect nf

large singly-charged spherical ions on solvent molecules

proposed earlier by Frank and Wen.

Relaxation Processes

The addition of solutes to water has been shown to

rigidify water structure on the nanosecond time scale (80).

Further, this phenomenon is reflected in the characteristic

relaxation times of water (87).

Dielectric studies show that water molecules are able

to respond to alternating electric fields up to very high

frequencies of about 1010 sec-', and in the literature two

types of dielectric relaxational times (rD and TL) have been

discussed (124). The Debye relaxation time, To, which is a

phenomenological time constant related to the

reorientational time constant of the solvent dipole, and the

longitudinal relaxational time, rL, which is the time scale








38

for excited state solvation of a Debye solvent as well as

the time scale for both ground and excited state solvation

of dipolar and ionic solutes (125).

In pure liquid water, Pottel (126) states that small

reversible shifts and rotations of atoms and molecules away

from temporary equilibrium positions around which

oscillations are possible relate to TL, while displacement

(or irreversible rotations) relate to TD. According to

Onsager (127), relaxational processes associated with

solvation start far from the probe (i.e., in the bulk

solvent and approximate the T timeframe) and proceed inward

towards the probe where the relaxation time approximates

To - the "inverted snowball" effect. However, he later

points out that a continuum of relaxation times between tr

and t, is possible (128,129).

Recently, Robinson and coworkers (48) claimed that to

accommodate the altered polarity that accompanies the

ionization of weak acids, water molecules must rearrange

themselves around the conjugate base in the TL time frame.

In their defense, they cite that that the solvation dynamics

of large polar molecules correlate well with the

longitudinal relaxational times of alcohol molecules (130).

Similar behavior has now been found in polar aprotic

solvents (131). On the other hand, to solvate protons,

Robinson and coworkers have envisioned that water molecules








39

must physically wrap themselves around protons, and that

such motion can only occur on the r, time scale, leading to

their backward sounding concept that "small elementary ions

hydrate more slowly than larger ions." In defense, they cite

their recent work in mixed solvents which spans the gap

between a diffusion-controlled regime (translation of water

molecules to satisfy the local concentration requirement)

and a hydration-controlled regime (rotational diffusion of

water molecules to satisfy the local structural requirement)

(48).

It is readily seen that relaxational processes provides

additional insight into the mechanism of solvation. However,

the structural role played by the surrounding solvent

remains open to debate.


Proton Transfer

The transfer of protons between molecules is one of the

most important reactions found in chemistry today (132). It

has been suggested that the kinetics associated with proton

transfer are directly influenced by both the molecular

dynamics and the molecular mechanics of the accompanying

proton solvation (23). However, as in solvation, the

molecular nature of these reactions remains unclear.

From Eigen's (37-39,62,66) theory of proton transfer,

the solvated proton has been described in terms of the










following structural species:

1) the hydronium ion, H30' (primary solvation)

2) the complex ion, H,04+ (secondary solvation)

3) the hydrated complex ion (tertiary solvation)

The existence of hydronium and the complex ions have

previously been demonstrated through mass spectroscopy

(133,134). Further, neutron scattering (135) has led to the

characterization of the hydronium ion (pyramidal in shape,

with its oxygen atom residing 0.32 A above the plane of the

three protons, its H-O-H angles were 110.4 and its O-H

internuclear distances were 1.01 A in length). This gives

strength to Eigen's claim that the excess positive charge on

the hydronium ion was equally distributed between the three

protons.

According to Eigen, the spatial characteristics of the

hydronium ion provide it with its unique ability to form

stable hydrogen bonds with neighboring water molecules which

in turn leads to the formation of stable H904* complex ions.

Further, it is Eigen's belief that the formation and

destruction of additional, relatively labile hydrogen bonds

along the border of the stable H,04' complex ion leads to its

"structural diffusion" in water. He claims that this is the

manner in which rapid proton transfer takes place, the rate

limiting being structural diffusion or the rate at which

molecules along the periphery of the hydrated complex ion








41

came into positions to allow hydrogen bond formation and

disruption. Unfortunately, experimental evidence does not

fully support Eigen's proton transfer mechanism. Moreover,

traditional theories of proton transfer which dated back to

1905, have relied on the Grotthus "chain" mechanism to

qualitatively explain prototropic behavior in solution.

In this classical mechanism, protons are considered to

be transferred along a linear chain of hydrogen-bonded water

molecules. For successive proton transfers to occur, each

"proton jump" has to be followed by a structural

rearrangement of the surrounding water molecules. Conway et

al. (23, 136) suggest that the transfer of a proton from a

hydronium ion to an adjacent water molecule leads to an

unfavorable orientation of neighboring water molecules. They

claimed that classical proton transfer mechanisms were too

slow and that proton tunneling across hydrogen bonds was to

fast to account for proton mobilities. Thus, the rate-

limiting step in proton transfer must be the rotation of

hydrogen-bonded water molecules near the H30' ion, i.e.,

rotation of molecules within the immediate proton hydration

sphere. They also found that this type of mechanism could

account for hydroxide's mobility in water.

According to Erdey-Gruz and Lengyel (138), different

mechanisms for proton transfer must be assumed because the

proton transfer process depends both on the nature and








42

composition of the solvent system. In support, Robinson et

al. (48), have recently shown that in aqueous-alcohol

solvent systems, the mechanism of proton transfer spans the

gap between a diffusion-controlled regime (Eigen's model)

and a hydration-controlled regime (Conway et al.'s model).

They also found that acid dissociation and its attendant

proton hydration, produced the H,O* ion as its direct

kinetic product. This tends to substantiate Eigen's claim

that rapid proton transfer occurs within a strongly hydrated

complex. Further, Castleman and coworkers (137) have now

provided direct experimental evidence of stable cluster ions

through vacuum-UV photoionization. However, additional

investigations, preferentially in solution, are still

required to substantiate Eigen's claim.


Proton Transfer in the Excited Singlet State

Excited state proton transfer, a special form of proton

transfer exhibited by certain aromatic acids and bases

following photon absorption, will be used as a tool to probe

solvation phenomenon and solvent structure in mixed aqueous

solvent systems. The kinetics associated with such

processes, which occur during the lifetime (typically 107

to 10-11 s) of the lowest excited singlet state (S,), can be

determined with steady-state and time-resolved fluorometric

techniques. The underlying principles associated with each








43

of these techniques will be preceded by a brief discussion

on the electronic and accompanying spectral effects of

proton transfer in electronically excited singlet states.

Electronic and Accompanying Spectral Effects

In 1949, F6rster (139) concluded that the spectral

dependence on pH of electronically excited aromatic acids

and bases, which was initially observed by Weber (140),

arose from proton transfer processes that took place during

the lifetime of the excited state. This behavior was

subsequently quantified by Weller (141) with his steady-

state kinetic approach to these dynamic processes. Several

advances have stemmed from these pioneering efforts. For

example, it is now recognized that aromatic molecules

possessing functional groups which donate electronic charge

to the aromatic system become stronger acids in the excited

state because the electrostatic attraction between the

proton and the functional group decreases, (pKa < pKa).

Hydroxyl, sulfhydryl, and amino groups along with pyrrolic

nitrogens all donate electronic charge. On the other hand,

aromatic molecules possessing functional groups with vacant

low lying n orbitals accept electronic charge from the

aromatic system upon excitation, making proton abstraction

more difficult, (pK, > pK,). Thus, aromatic compounds of

this type become stronger bases in the excited state.

Carbonyl, carboxyl, carboxylate, and amide groups are








44

capable of accepting electronic charge. In addition,

pyridine nitrogen atoms (such as those found in quinoline)

show increased basicity upon excitation even though they do

not posses vacant low lying i orbitals because electronic

charge becomes localized on the nitrogen atom. Further, the

absorption and fluorescence spectra of aromatic molecules

capable of transferring protons in the excited state will

reflect these changes. For example, upon excitation the

longest wavelength absorption and fluorescence bands of

aromatic molecules which possess electron donor groups shift

to shorter wavelengths upon protonation and to longer

wavelengths upon dissociation. Such behavior is indicative

of a decrease in basicity and an increase in acidity upon

going from the ground to excited states. Similarly, the

long-wavelength absorption and fluorescence bands of

aromatic molecules possessing electron acceptor groups shift

to longer wavelengths upon protonation and shorter

wavelengths upon dissociation. This results from an increase

in the basicity and a decrease in the acidity of such

molecules upon going from ground to excited state.

Steady-State Kinetics

The approach developed by Weller (141) to determine the

excited state proton transfer rate constants is based upon

the pH dependent fluorescent intensities of the molecular

species involved and the assumption that all photophysical








45

and photochemical processes deactivating Sz have achieved

steady-state. With this in mind, proton transfer from

excited aromatic acids and bases in dilute aqueous solutions

can be considered to occur by one of two simple reversible

two-state reaction schemes, shown below. Of course, more

complex behavior is often observed, but Weller's simple

model systems provide the conceptual framework for the more

complex processes. In the first case, which describes the

excited state dissociation of 2-,3- and 4-hydroxybiphenyl,

weak aromatic acids examined in this investigation, water

acts as the proton acceptor while the solvated

hydroxybiphenyl(s) act as the proton donor.


ka
A' + rHO H' + B
k,
Case 1 1/To 1

A B


The second case describes the excited state reaction of 6-

methoxyquinoline, a weak base that was examined in this

investigation. Here, water acts as the proton donor and the

solvated 6-methoxyquinoline as the proton acceptor.


kb
B' + rH,0 OH- + A'
ka
Case 2 1/To 1/To'

B A








46

In both cases, r represents the difference between the

number of water molecules solvating both reactants and

products. In turn, To and To' represent the lifetimes of the

reactants and products in S, in the absence of proton

transfer, which are equal to the reciprocals of the rate

constants for photophysical deactivation of S, are given as:


r = 1 and rT 1
(kT'f, k ^)

where kf and k. are the fluorescence rate constants and kd

and kd are the rate constants for radiationless

deactivation.

From Case I, the relative quantum yields of

fluorescence (9/qo and ('/9p) for the excited acid (A') and

conjugate base (B') respectively, are obtained by

simultaneously solving the following two equations for the

disappearance of A' and B' from S,:


-fd[A.] f(I +k)[A*]dt fkbf[H][B]dt (1.12)


-fd[B-] = ( +kb[H'])[B']dt fka[A ]dt (1.13)


In the equations above, [A'] and [B'] are the respective

relative probabilities of finding an A" and a B" molecule in

the excited state at any time (t) following excitation. The

rate constants (or probabilities) of dissociation and








47

protonation of A" and B' are ka and kb, respectively, and [H']

is the molar concentration of hydrogen ion. The lower limits

of integration, (A and (x,, are the fractions of excited A

and B or the fractions of total light energy absorbed by A

and B at the wavelength of excitation, respectively, and are

defined as follow:

c,[A] EB[B]
(U, and a [
E,[A] + CB[B] B E,[A] + :B[B]

where the molar absorptivities of A and B at the excitation

wavelengths are given by EA and EB. The ground state

concentrations of A and B at equilibrium and at known pH are

[A] and [B]. In turn, the relative quantum yields of

fluorescence are defined by the integrals (141):


A ]dt = T r and j '[B']dt = = IP


where the actual lifetimes of A' and B' are T and T' and are

proportional to relative fluorescent intensity measurements

when an isosbestic point (EA = E) is used for excitation.

Integration and rearrangement of equations (1.12) and (1.13)

gives:


aA + k[ H']
T0 1 + kT + kt'o[H] (1.14)











/ a + kato
1 + kr, + kbr[H']
(1.15)



based upon the relationship that the fractions of directly

excited A and B are equal to unity, (a. + (0, = 1). Equations

(1.14) and (1.15) can then be combined to give:


9~/~ 1 kTo[H-] / (1.16)
p'/', a, ,


where it is readily seen, that a plot of:


vs. [H']
q'/9P -/ B /qO t/9 ,

provides a means by which to calculate k, and k, (the rate

constants of proton exchange) provided the lifetimes of the

species in the excited state are known.

When protonation of the excited conjugate base occurs

at [H+] >> K,, where K, is the ground state dissociation

constant, aB = 0 and equation (1.16) reduces to

Weller's equation (141).


Io 1 kbt '
+ [H ] (1.17)


Under these conditions, the excited state reaction is

reversible. However, at pH > 4 ([H'] < IxlO-'M) the reaction








49
is virtually independent of pH and irreversible. This is

because k, is diffusion controlled (k, s 5 x 10 M-'s-1) and Tr

is approximately 1 x 10'-s, making kbr[H'] s 5 x 10-2. Under

these conditions, equation (1.17) to be reduced to:


S= kj, (1.18)

which provides a means to directly determine k., provided r,

is known.

From Case 2, the relative quantum yields of

fluorescence (0/0o and p'/)') for the excited base (B') and

conjugate acid (A') respectively, are obtained by

simultaneously solving the following steady-state equations

for the disappearance of B' and A' from S,:


Jd[B] ( + kb)[B']dt k,[O -][A-]dt (1.19)




fd[A-] ( + k,[OH-])[A']dt fk[B]dt (1.20)


As in Case 1, the excited state rate constants, ka and k,,

can be determined. Integration and rearrangement of

equations (1.19) and (1.20) gives the following equations:


(P ia + ka'o[OH-]
90 1 + kr, + kar o[OH- 21)












(i A + kbo
1 + kto + kt'oOH ] (1.22)

which when combined leads to the following complex

expression:


+ S k [OH- / ) (1.23)


which can be reduced, when [H+] << K,, tA = 0, to another one

of Weller's reversible two-state kinetic equations (142):


1P / 1o 1 kao[OH
qIp + k.O (1.24)


which provides the means to calculate the individual rate

constants, provided appropriate lifetimes are known. The

relative fluorescence intensities and their ratios become

irreversible and independent of pH when pH < 9, because

kaj[OH-] << 1. Under these conditions, equation (1.24)

reduces to:


kb (1.25)

which allows direct determination of kb, provided T, is

known.

It is important to recognize that the preceding steady-

state development of excited state proton transfer is based

upon the assumption that the presence of solute molecules or








51

ions in solution does not affect the properties of the

reactants or the activated complex. However, when ions are

present, deviations from ideality become apparent even at

low concentrations. Thus, the observed rate constants, as

defined in equations (1.14), (1.15), (1.21) and (1.22), are

not necessarily those which would be obtained at infinite

dilution. According to Bransted (143,144), the medium

effects on the rate constants in relatively dilute solutions

of acid and base within the pH range of 1-11, can be

compensated for by using the kinetic activity factor (F),

where F is of the form:


F = exp( ZZ e2 K (1.26)
EkTa 1 +K

where Zle and Z2e are the charges of the ionic reactants, r

is the dielectric constant, kT is the Boltzmann's factor, a

is the ionic encounter distance and K is the reciprocal

thickness of the ionic atmosphere. When compensation is

made, equations (1.16) and (1.23) take on the following

form:


/0, 1 kb, o[H] '/ ) ('
-77 F77 (1.27)
0'/o -a, + '/p 1-




/Io 1 k .T[OH-] /' )
('/9~ a T T ('/9' a








52

Perhaps the single most important advance in the study

of excited state proton transfer processes since Weller

developed the steady-state equations has been the

development of sensitive detection methods such as pulsed

and phase-shift demodulation fluorimetry (145). With these

techniques it is now possible to look directly into the

reaction process by monitoring the fluorescence decay of the

sample (146). The pulsed method that was employed in this

investigation was time-correlated single photon counting

(147). The theory associated with the determination of

kinetic rate constants with this method are elaborated upon

below. However, the principles behind time-resolved

measurements is saved for the experimental section.

Time-resolved Kinetics

According to Laws and Brand (148), excited state

proton transfer rate constants can be evaluated using pulsed

methods, provided the following assumptions are made. First,

the excited state molecule exists in two states, each with a

distinct emission spectrum and second, that the kinetic rate

constants of these states are independent of the emission

wavelength.

In the first case, which describes excited state

dissociation of 2-,3- and 4-hydroxybiphenyl, the following

differential equations apply:











d[A'] 1
d ( + k )[A'] k b[B'I[H']

(1.29)




d[B'] 1
S( + k [H'])[B*] ka[A'] (1.30)



Boundary conditions are chosen such that [A'] = [Ao] and [B']

= 0 at t = 0 (148,149). Following 6-pulse excitation, the

time-dependent intensities of A' and B' at any wave number v

are given by:


I,(V,t) a (V)exp-t" + 2(V)exp- (1.31)


I,(V,t) fl(V)exp t'/ + f,(V)exp (1.32)


The decay times of A' and B' will be given by:


S 2 1[(Ya + Y.) + ((Y. YA)2 + 4kk [ ])1/2] 1.33



when y, and y. are defined as follows:


YA + k and y, 1 + k[H]
to


The pre-exponential factors in equations (1.31) and (1.32)

are given by:











(X(V) CA(V)[Ao] YA Y2kA
Y1 Y
(1.34)




U,(V) = CA(V)[A0] k, (1.35)
y1 AY2




-f1(v) P2(V) c(V )k [A]B (1.36)
Y Y2


In the equations above, Yi = r1i-, Y2 2-', kA, and kB are

the probabilities of radiative deactivation, while C,(,) and

CB,() are the emission spectra of species A" and B',

respectively, normalized to unit area:


C,(v) ()
a ,(V)dv (1.37)





f,(V)
C,(V) e((v)
SB,(V)dv (1.38)


where OA(v) and B,(V) are the number of quanta emitted at any

wavenumber by the A" and B' species, respectively.

At any given wavenumber the time-resolved decay of










fluorescence is given by:

(1.39)
I(V,t) IAV) + (,t)


where the individual decays of A' and B' are given by

equations (1.31) and (1.32), respectively, such that their

substitution into the above equation leads to

(1.40)
I(V,t) (at + l)exp-"'' + (0( + P2)exp- (1.40)


In case 2, the equation for time-resolved decay is identical

to the ones given above, provided one makes the appropriate

corrections for the respective decay times and pre-

exponential factors. These components are given as follows:


Ti1,2 [(YA + Y2)+((Y, Y )2 + 4kk,[OH-]) 2] (1.41)




-a(V) a,(V) CA(V) kbB]ktA (1.42)
Y1 Y2




P(V) C,(V)[B] YB k,, (1.43)
( C k (1.44)



P2(V) C(V)[B ] Y1 Y/kkB (1.44)
Y1 Y2


These results are readily understood when the








56

differential kinetic equations of case 2, seen below, are

compared with those of case 1.


-d ] ( + kb)[B'] k[A'][OH-] (1.45)





d[ ] + k,[OH-])[A'] k[B"] (1.46)



Returning to equation (1.40), one can see that at any

given wavelength the fluorescence decay appears to be

biexponential. It should be pointed out, however, that both

the decay times and pre-exponential factors are actually

complex functions of the kinetic constants of the system.

Further, the pre-exponential factors are also functions of

the spectral distributions of the species present.

Therefore, one must not make the assumption that the decay

times obtained from this expression represent unique

components of the sample system unless other indications are

available to support this interpretation.

There are several techniques by which rate parameters

can be extracted from fluorescent decay measurements

(145,150). In this investigation, the "difference of

lifetimes" method was used to extract the protonation and

dissociation rate constants from 6MQ and 2HBP, respectively.

The principles behind this technique are as follows. In case








57

1, if pH is adjusted such that [B'] = 0 and k, -* 0,

fluorescence decay becomes monoexponential:

(1.47)
I(, t) IA( ,t 0)exp (1.47)


where y = I/T,. Similarly, at neutral pH where [H'] and

kb[H'] 0, decay is also found to be monoexponential.

However, under these conditions, yi = (1/to + k,). Thus, if

t, can be determined and it is independent of the

concentrations of A and H*, the difference between the

reciprocals of the two lifetimes yields the forward rate

constant ka. Finally, in alkali boundary conditions can be

chosen such that [B'] = [B;] and [A'] = 0 at t = 0. Under

these conditions, the decay profile of the conjugate base is

monoexponential, and the following equation applies:

(1.48)
l(V,t), I(v,t = 0)exp-'' (


where y, = 1/Tr, provided T- is independent of B and OH-.

The use of this method provides a means by which the forward

rate constant can be determined. It also serves as a means

by which to obtain the lifetime of the individual species in

the absence of proton transfer, thus allowing for the

solution of the steady-state equations. The difference of

lifetime method can be developed for Case II in a similar

fashion. However, in the case of 6MQ, quenching of

fluorescence prevented the direct measurement of T,. An








58

additional set of experiments were required to estimate To

before the difference method could be used to calculate the

protonation constant for 6MQ. The rationale behind these

experiments will be better understood if it is left for the

results and discussion section.


Objectives

Recently Kelly (151), using traditional steady-state

spectroscopic technique, claimed that the activity of water

provides the best means through which the reduction in

excited state proton transfer in mixed aqueous-organic

solvent systems can be described. However, since his results

were not conclusive, it was suggested that excited state

proton transfer must be dependent upon some yet unidentified

physical chemical property of the mixed solvent system, such

as changes in the secondary or tertiary solvation of the

functional group bearing the reactive proton. Ultimately, he

could only conclude that additional investigations in mixed

aqueous organic solvent systems would be required before a

proton transfer model acceptable in all respects could be

proposed.

The present study was, therefore, undertaken to further

investigate and attempt to correlate the medium dependence

of excited state proton transfer in mixed aqueous organic

solvent systems with the activity of water in a manner








59

similar to that adopted by Schulman and Kelly (74,75).

However,in this investigation, time-resolved spectroscopic

technique will be combined with traditional steady-state

techniques because the union of these techniques allows the

rate constants of these elementary proton transfer reactions

to be determined with greater accuracy than by either

technique alone. Further, since traditional steady-state

methods can only focus on the role of solute structure and

reactivity in excited state proton transfer reactions, the

direct observation of ionic dissociation processes in water

with time-resolved spectroscopic methods should provide

additional insight into the effects of both solute and

solvent properites on the excited state proton transfer

process.

The weak base, 6-methoxyquinoline (6MQ) and the weak

acids, 2-, 3- and 4-hydroxybiphenyl (2HBP, 3HBP and 4HBP)

depicted in figure 1.1, were used as excited state probes in

the current mixed solvent investigations. It was felt that

an examination of 6MQ in mixed solvents would help to

determine if the steady-state approach developed by Schulman

and Kelly with the weakly acidic, sulfonated naphthols

(74,75) could be extended to weak bases. However, before

these studies could be conducted, it was imperative that the

behavior of 6MQ in aqueous media be properly identified to

provide a much needed reference point from which to base the














H3CO




6MQ


N


3HBP


OH
OH
4HBP


Figue 1.1. The chemical structures of 6-methoxyquinoline
(6MQ), 2-hydroxybiphenyl, 3-hydroxybiphenyl
and 4-hydroxybiphenyl (2, 3, and 4HBP).


2HBP








61

prototropic behaviors of 6MQ in mixed aqueous-organic

solvent systems. In addition, aqueous studies will hopefully

clear up some conflicting reports in the literature on 6MQ.

The major discrepancies have been listed in table 1.1,

below.


Table 1.1. Acid-base properties of 6-methoxyquinoline
in the lowest excited singlet (Si) state.

kb k, pK: Reference
(x 10-e s-') (x 10-1" M-1 s-1)

0.64 2.6 10.4 152
2.0 + 0.6 3.1 + 0.6 10.8 + 0.1 43
1.2 N.R. N.R. 153



It was felt that an investigation of the

hydroxybiphenyls, first, would provide a means to determine

if the steady state approach developed by Schulman and Kelly

could be extended to weak acids systems, which are composed

of two chromophores that are coupled by a single carbon-

carbon bond. Second, since a 42 + 2 dihedral angle exists

between the two rings of the biphenyl molecule in the ground

state (154) since the molecule becomes planar upon

excitation (155), it was felt that fluorometric

investigations of these simple nonrigid acids may provide

some additional insight into the behavior of larger nonrigid

molecular systems. Finally, before an investigation can be

conducted in mixed solvents, the aqueous properties of these









62

probes in the lowest excited state must be determined,

because even though Bridges et al. (156) have found that the

hydroxybiphenyls exhibit prototropic activity, they did not

characterize this activity.














CHAPTER 2
EXPERIMENTAL



Materials


The probes, 2-, 3-, and 4-hydroxybiphenyl, were

purchased from Aldrich Chemical Company, Inc., Milwaukee,

Wis. The 2-isomer, 99+ % purity, was used as supplied.

However, a triple recrystallization procedure from 10%

ethanol was required to clean up the 3- and 4- isomers prior

to use. The probe, 6-methoxyquinoline, was purchased from

Pfaltz and Bauer, Inc., Flushing, N.Y. To remove

contaminants, the compound was converted to its perchlorate

salt in 90 % ethanol with the addition of 70% perchloric

acid. The perchlorate salt was then recrystallized twice

from 90 % ethanol solutions prior to use. The purities of

these probes were subsequently confirmed when

absorptiometric pKa determinations were found to be in

agreement with the observed literature values (152,157).

Acetonitrile, dimethyl sulfoxide, formamide, dioxane,

methanol, isopropanol, tertiary butanol, perchloric acid

(70%), sulfuric acid (all reagent grade), sodium hydroxide

(1.0 M and 5.0 M) (analytical grade), certified pH 4.0, 7.0

63








64

and 10.0 buffer solutions and Davison 4 A molecular sieves

were purchased from Fisher Scientific Co., Fairlawn N.J.

Absolute ethanol was purchased from AAPER Alcohol and

Chemical Co., Shelbyville, Ky. Boiled deionized water was

used in the preparation of all aqueous solutions.


Instrumental


Absorption measurements were made on a Perkin Elmer

Lambda-3 spectrophotometer. Absorption spectra on all probes

were monitored from 500 nm to 210 nm.

Fluorescence measurements were made on a Perkin Elmer

LS-5 fluorescence spectrophotometer whose monochrometers

were calibrated against the xenon line emission spectrum.

The wavelength variable spectral output was uncorrected for

instrumental response. Excitation of the probes was effected

at an isosbestic point found in the absorption'spectra. For

the 2-, 3- and 4-hydroxybiphenyl isomers, the excitation

wavelengths were 290 nm, 284 nm, and 270 nm, respectively.

For 6-methoxyquinoline, the excitation wavelength was 289

nm. In all cases, emission was monitored from 320 nm to 540

nm. Excitation and emission slits were set at 3 nm.

Emission decay measurements were made on a

Photochemical Research Associates (PRA) model 1551 time-

correlated, single photon counter. The excitation source was

an H2-filled spark gap (PRA model 510 B). The instrument








65

response was 2.0 ns fwhm. In all cases, excitation light

was filtered with a Corion near-UV 10 nm bandpass filter

that had a maximum transmittance at 280 nm. Emission light

was filtered with a Schott UG-11 near-UV 10 nm bandpass

filter that had a maximum transmittance at 350 nm or with

Schoeffel UV-cutoff filters at 418 nm or 450 nm. The

excitation repetition rate was 30 KHz and the sampling rate

was < 500 Hz. Analysis of emission decay data was carried

out on an IBM-PC using a deconvolution program from PRA. A

Fisher Accumet 950 pH meter which employed a Fisher pencil

gel-filled silver/silver chloride combination electrode

calibrated against pH certified buffer solutions was used

for pH measurements. Gilmont pipets (0.02, 0.2 and 1.0 ml)

were calibrated and employed to deliver volumetric amounts

of water and cosolvent mixtures.


Methods


Steady-State Emission Spectroscopy

Aqueous acid-base titrations

A known volume of a 10-2 M ethanolic stock solution,

prepared prior to investigation, was micropipetted into a

series of 10 ml volumetric flasks. The aliquot was then

evaporated under a stream of nitrogen gas before the residue

was brought to volume by dilution to the mark with water,

0.01 M, 0.1 M, 1.0 M, 3.0 M, or 5.0 M standardized HSO, or








66

NaOH. The final probe concentration was 105 M, chosen to

keep the absorbance at each of the excitation wavelengths to

less than 0.02 absorption units, thereby reducing the

probability of nonlinear fluorescence.

The aqueous acid-base titrations, preformed at room

temperature, were conducted as follows:

1) Two milliliters of the aqueous probe solution was

pipetted into a 1 cm2 cuvette having a 4 ml volume, and its

emission spectrum and pH were obtained.

2) The above solution was then titrated with an acidic probe

solution. Fluorescence spectra were scanned and pH was

measured after each increment of titrant was added. The

titration was carried out until the excited state reaction

could no longer be detected, (i.e., the fluorescence species

either reached a maxima or went to zero, respectively).

Calculation of ionic strength at each point in'the titration

curve was carried out using the formal concentration of the

acid.

3) Steps 1-2 were repeated with alkaline probe solutions.

4) The relative fluorescence intensities for both the acidic

and basic species of excited probe, as a function of pH,

were converted to relative quantum yields for analysis. The

relative quantum yields for the neutral species of the

hydroxybiphenyl isomers or 6-methoxyquinoline were obtained

with the following equation, g/T0 = [(F, F,)/ (F. F,)],,








67
where F,, F,, are the relative fluorescence intensities of

the neutral species at intermediate and maximal pH values

and Fc is the residual relative fluorescence intensity of

the charged species, all at the analytical wavelength, .,

which is generally taken to be the wavelength at which the

fluorescence of the individual species is maximum.

Aqueous-organic solvent titrations

Working solutions for investigation of the forward

reaction were prepared from stock solution as above, with

one exception. In this case, they were brought to final

volume by either the addition of water or a dried organic

solvent, which had been stored over 4 A molecular sieves for

a minimum of 24 hrs.

The mixed aqueous-organic solvent titrations for

investigation of the forward reaction, performed at room

temperature were conducted as follows:

1) Two milliliters of the aqueous probe solution was

pipetted into a 1 cm2cuvette having a 4 ml volume and its

fluorescence spectrum was recorded.

2) The above solution was then titrated with the organic

probe solution. Fluorescence spectra were scanned after each

increment of titrant was added until the mole fraction of

the organic component approached unity, provided emission

from the excited state reaction could still be detected.

3) The fluorescence spectrum of the non-aqueous solution was










also recorded.

4) The relative fluorescence intensity obtained for each

species of probe as a function of the organic cosolvent mole

fraction was converted to relative quantum yield in a

fashion similar to above. However, rather than isolating the

individual fluorescence maxima for each species of probe at

each point in the mixed solvent titration, fluorescence

maxima were fixed to the initial aqueous values. This

normalization procedure provided a consistent means by which

the effects of the individual solvents on the forward rate

constants could be compared.

Aqueous-organic acid-base titrations

Solutions of the probe, 2-hydroxybiphenyl, were

prepared by pipetting a known volume of dried organic

cosolvent to three 10 ml volumetric flask, each of which

contained probe such that the final concentration was 10-5

M. One of the 10 ml flasks was brought to volume with water.

The remaining flasks were brought to volume with water after

addition of a known aliquot of standardized H2SO, or NaOH,

respectively. Thus, each flasks contained the same

concentration of organic cosolvent.

The mixed solvent acid-base titrations for the

investigation of the back reaction, preformed at room

temperature were conducted as follows:

1) Two milliliters of the aqueous-cosolvent mixture was








69

pipetted into a 4.0 ml cuvette and its fluorescence spectrum

was recorded.

2) The above solution was then titrated with the acidic

mixed solvent mixture. Fluorescence spectra were scanned

after each increment of titrant was added, and the titration

was carried out until the fluorescence of the conjugate base

could no longer be observed and the fluorescence of the acid

was maximal and constant. Calculation of ionic strength at

each point of the titration curve was carried out using the

formal concentration of the acid.

3) The fluorescence spectrum of the basic probe solution was

subsequently obtained to allow for the generation of

relative quantum yield.

4) The relative fluorescence intensity obtained for each

species of probe, as a function of the formal hydrogen ion

concentration in each of the mixed solvent systems

investigated, was converted to relative quantum yield after

the system was normalized. In the normalization procedure,

the relative quantum yield of the conjugate base at the

specific organic cosolvent mole fraction investigated was

fixed to the relative quantum yield obtained in the previous

titration. However, the relative quantum yield for the acid

under identical conditions was obtained through the

following relationship (141), 1 p'/(p = p/ ,. The

rationale for these procedures are given below.








70

Upon analysis, the sum of the observed relative quantum

yield for both excited state species of probe did not add up

to one. In fact, as organic concentrations were increased,

the sum of the observed relative quantum yield from both

species was found to deviate from 1 by as much as 40 %. Any

calculation and comparison of kinetic rate constants under

these conditions would have be meaningless. With further

investigation, it was found that the neutral species

undergoes extensive quenching in the acidic range as the

concentration of the organic solvent is increased, due in

part to the organic concentration and in part to the acid

concentration. The conjugate base on the other hand did not

exhibit any quenching processes. In fact, the relative

quantum yield values obtained from both mixed solvent

studies for the conjugate base were generally found to be

within 3 % of each other. The effect of quenching by organic

cosolvents was minimized when the relative quantum yield of

the neutral species at each mole fraction investigated was

fixed to the relative quantum yield of the conjugate base as

noted above. When this is done, the sum of relative quantum

yield as a function of the formal hydrogen ion concentration

at each of the organic cosolvent mole fractions investigated

was generally found to be within 5 % of 1, sometimes up

though 0.1 M [H']. This observation provides support to the

validity of this method. Further, normalization provided a








71

consistent means by which the effects of the individual

solvents on the reprotonation rate constant could be

compared.


Time-resolved Emission Spectroscopy

Emission decay measurements were carried out with the

time-correlated single photon counting technique (147). A

generalized schematic diagram of the instrumentation is

shown in Figure 2.1, below. The important components are the

start photomultiplier tube (PMT), the stop PMT, the flash

lamp (L), the time-to-amplitude converter (TAC), the

multichannel pulse height analyzer (MCPHA) and of course,

the sample (S).



Stop PMT


t output

TAC

Start PMT -couI


timing pulse


Figure 2.1. Time-correlated single photon counting
instrument schematic.








72

The basic principles behind emission decay measurements

are as follows. A pulse of light (L), simultaneously excites

the sample (S) and initiates the start PMT which signals the

TAC, whose function it is to charge a capacitor linearly

with time. After some period of time, t, the sample may or

may not emit a photon. If it does, the photon is detected by

the stop PMT which in turn signals the TAC to stop charging.

Thus, the voltage stored by the TAC is proportional to the

time period between start and stop pulses. If no photon is

detected after a predefined period of time, then the TAC

resets itself and prepares for a new start pulse. However,

if a signal is collected, the voltage is sent to the MCPHA,

whose function is to digitalize and store the analog signal

as a single count in a memory location or channel that

corresponds to the pulse height. Thus, the pulse height is

also proportional to the time that the photon is detected

after initiation of the start pulse.

Quantitatively, the observed decay curve generated by

the MCPHA is a histogram that describes the probability that

a sample will emit a photon as a function of time following

excitation. If the distortions from the instrument response

time which arise because the excitation pulse is not

infinitely narrow in width (a typical pulse is 2 ns in

width) are accounted for, the histogram collected on the

MCHPA would correspond exactly to the excited state profile








73

of the emitting species. However, since the timescale of

most emissive decays is generally in the nanosecond time

regime, the observed emissive decay is a convoluted function

of the time profile of the excitation pulse and the decay of

the excited probe.

The observed decay function R(t) can be expressed as a

function of the time profile of the excitation pulse L(t)

and the impulse response of the sample as (147):
(2.1)
R(t) L(t')F(t-t')dt'

where F(t-t') is the response function of the probe and

L(t') is the time distribution of the lamp. The derivation

for this expression has been given elsewhere (158,159).

To obtain the undistorted sample decay function, F(t),

from the experimentally determined functions R(t) and L(t),

the method of iterative convolution is employed (145,147).

This procedure convolutes an assumed decay law F,(t) of the

form given by equation (1.40) and generates a calculated

decay function Re(t) while minimizing the deviation between

the experimentally observed decay, R(t), and the calculated

decay, Rc(t). The goodness of fit is statistically

determined from the reduced chi-square x!:


X W,[R(t)- Re(t)]2 (2.2)
i-n,













X -, X (2.3)
n, n, + 1 p



where X2 is the statistical chi-square, Wi = 1/R(t) is a

weighting factor for the ith data point, n, + n, are the

starting and ending channels, respectively, of the MCPHA

that are being fit and p are the number of variables in the

fitting function (160).

Fits for convolution are considered to be good if X!

are s 1.2. A poor fit is distinguished by a X? > 2

(145,147,160).

Aqueous acid-base titrations

The probes examined had final concentrations of 10-5 M.

They were prepared in water and known concentrations of

HSO, and NaOH as noted in the steady-state section above.

The emission wavelength for the neutral species of all

probes examined was 350 nm. The emission wavelengths for the

anionic species of the hydroxybiphenyls and cationic species

of 6-methoxyquinoline were 418 nm and 450 nm, respectively.

Titrations, preformed at room temperature, were

conducted as follows:

1) Two milliliters of the aqueous probe solution was

pipetted into a 4 ml cuvette and placed into a photon

counter equipped with a multi-channel analyzer. At least








75

5000 counts were collected in the channel of maximum

intensity at each emission wavelength examined, provided a

signal could be measured.

2) Two milliliters of a scatter solution (see below) was

pipetted into a 4 ml cuvette and placed into the photon

counter with the emission filter removed. The amount of

photons collected in this step were the same as that

obtained in step 1. The scatter solution was prepared by

diluting a small amount of Kodak white reflectance coating

in a solution that was identical to, but not containing the

sample.

3) The counts obtained in steps 1 and 2 were subsequently

downloaded to a computer fitted with a convolution program

by PRA which allowed for the calculation of the emissive

decay (i.e., lifetime).

4) The above solution was then titrated with either an

acidic or basic probe solution and lifetime measurements

were obtained after each increment of titrant was added,

provided counts could be obtained. Calculation of the ionic

strength at each point in the titration curve was carried

out using the formal concentration of either acid or base.

Aqueous-organic solvent titrations

Working solutions of 2-hydroxybiphenyl were prepared as

noted in the steady state section above. The lifetime (r) of

the neutral species in water was obtained as described








76

above. The aqueous solution was subsequently titrated with

the organic probe solution. Lifetime were collected after

each increment of titrant was added. The titration was

carried out throughout the entire organic cosolvent mole

fraction range.

Aqueous-organic acid-base titrations

Working solutions of 2-hydroxybiphenyl in 0.01 N NaOH

at various mole fractions of organic cosolvent were prepared

as noted in the steady-state methods section above, such

that lifetime could be obtained for the 2-hydroxybiphenylate

anion (rt).














CHAPTER 3
RESULTS AND DISCUSSION


Excited State Behavior in Aqueous Media


6-methoxyquinoline

Steady-state fluorescence

The fluorescence spectra of the isolated species of 6-

methoxyquinoline (6MQ) is shown in figure 3.1. Notice that

the emission maximum of the neutral or basic form of 6MQ,

which is isolated in alkaline solutions when pH 2 13 is 368

nm and that the emission maximum of the charged or acidic

species of 6MQ, which is isolated in acidic media when pH s

2 is 440 nm. When the relative fluorescence intensity (RFI)

of 6MQ is plotted as a function of pH, as depicted in figure

3.2, one finds that emission from the neutral species is

significantly reduced in the alkaline range. It has been

suggested by Schulman et al. (152), that the observed

decrease probably results from quenching by hydroxide ion.

However, this behavior has never been quantified.

In steady state investigations of excited state proton

transfer, the RFI of each species must be converted to their

respective relative quantum yield (RQY) before the kinetic








100




80




60




40




20




0
500 450 400 350 300

Wavelength (nm)



Figure 3.1. The fluorescence spectra of the acidic (A) and
basic (B) species of 6-methoxyquinoline. The
excitation wavelength was 289 nm.
0 !^____ ___













excitation wavelength was 289 nm.








































T. I'.








80

rate constants can be determined. The procedure used for

this conversion is generally straightforward and can be

found in the methods section of Chapter 2, p.67. However,

when the true fluorescence of an emitting species cannot be

directly obtained, e.g., when quenching occurs, the process

of generating accurate and reliable RQY is often quite

difficult.

In an investigation of 6MQ by Schulman et al. (152),

the RQY of the neutral or quenched species of 6MQ (9/0)

were estimated by plotting 1 (<'/g'o) against pH, where

(p'/i' is the RQY of fluorescence of the 6MQ cation, with

the assumption that the RQY of the cationic species were

true values. Unfortunately, this was not the case. As a

result, the rate constants generated were in error. A more

effective means of estimating the RQY of neutral 6MQ was

achieved by extrapolating the RFI of the neutral species as

a function of pH in the range where quenching plays a major

role as depicted in figure 3.3. The RFI maxima obtained with

this procedure was then used to generate reliable RQY. The

fact that the sum of RQY for 6MQ approaches unity (9/9p +

'/fO = 1) in the pH range where quenching does not play a

significant role as shown in figure 3.4, provides support

for this technique.

The quenching of fluorescence with a consequent

decrease in the observed fluorescence lifetime of an excited










































.'J'















0 CN
*H -
cu

.IO ---

0 0







0-4

r H
-0000-"-- ^ ^









0



*HO 0
In4 .0
1u ,





0-
SI x-,I






m
C) *

0 .






0 04
z0 0








(o,/,,) pUD (o/)
C a) 1 1









-,
Sf5








83

species is a common photochemical event which must be

accounted for in the kinetic expressions (161). One of the

major prerequisites of traditional steady-state kinetic

methods is that there should be no pH dependent quenching

(141). However, in cases where the kinetic rate constants

are much larger than the quenching rate constant, steady-

state methods can still be used to provide reasonable

estimates of the individual kinetic rate constants. For

example, it was assumed in the development of Case 2 that

quenching does not occur. However, since quenching has been

shown to influence the emission of the neutral species, its

behavior must be accounted for. This is achieved through a

slight modification of Case 2 as shown below.

kb
B' + rH0 OH- + A'
k.
Case 2
(modified) (l/To + k,[OH-]) 1/T'

B A

The only noticeable difference is the inclusion of ko,[OH-]

in the deactivation pathway of B', where kon is the hydroxide

ion quenching rate constant and [OH-] is the molar

concentration of hydroxide ion.

To determine the rate constants of protonation and

deprotonation in Case 3, the following equations are

integrated: