THESI‘;

LIBRARIES
MlCHlGAN STATE UNIVERSIT:
EAST LANSlNG, MICH. 4882

This is to certify that the

dissertation entitled

Excited State Proton Transfer

presented by

Nahid Shabestary

has been accepted towards fulfillment
of the requirements for

Ph . D. _ degree in Chemistry

 
    

Major pr essor

Date 17/97 (of .1?ij

MS U i: an Affirmative Action/Equal Opportunity Institution 0- 12771

 

 

MSU

LlBRARlES
_:_.

 

 

RETURNING MATERIALS:
Place in book drop to
remove this checkout from
your record. FINES will
be charged if book is
returned after the date
stamped below;

 

 

 

 

 

 

 

 

 

 

 

fiEEG~

~_. . i-_ - ._Wl

 

 

EXCITED STATE PROTON TRANSFER

By

Nahid Shabestary

A DISSERTATION

Submitted to
Michigan State University
in partial fulfillment of the requirements

for the degree of

DOCTOR OF PHILOSOPHY

Department of Chemistry

1982

 

 

]—_i ,

ABSTRACT
EXCITED STATE PROTON TRANSFER
By
Nahid Shabestary

Excited state proton transfer represents one of the

 

‘ Simplest photoreactions that is expected to occur when the
acidity or the basicity of a functional group is enhanced

upon excitation. Phenomenologically, excited state proton

fluorescence band which has a large Stokes shift. For ex—
ample, methyl salicylate in hydrocarbon solvents exhibits

l
l
l transfer manifests itself usually by the appearance of a new
I
l
: two fluorescence bands, one at 3A0 nm and the other with a

l

large Stokes shift of 10,000 cm"1

at A50 nm. The latter
fluorescence band results from a tautomer produced in the
excited singlet state as a result of an intramolecular pro—
ton transfer. The methoxy derivatives which cannot under—
go excited state proton transfer exhibit a normal fluores—
cence, i.e., a band at N340 nm. The first chapter deals

with examples of excited state proton transfer in solution, the

solid phase and the vapor phase. Several aspects related to

 

Nahid Shabestary

the mechanism of excited state proton transfer are sum—
marized. The question of whether the emission results are
explained in terms of an excited state equilibrium or in
terms of equilibrium between different ground state con—
formers is discussed.

We have studied the absorption and luminescence prop-
erties of p-amino salicylic acid (PAS) and N,N dimethyl p-
amino salicylic acid (DPAS) and their methyl esters (PASE,
DPASE), under various conditions. Effects of the medium,
pH and substituents helped in the interpretation of the ab-
sorption spectra which are dominated by "charge-transfer
transitions. In the "charge transfer" states of these
molecules the hydroxyl and amino groups act as electron
donors while the carboxyl group acts as an electron ac-
ceptor. The emission spectra clearly reveal the occurrence
of proton transfer from the hydroxyl group to the carbonyl
oxygen during the lifetime of the first excited singlet
state. In the case of DPAS the charge transfer state
where the hydroxyl group acts as an electron donor (O)CT,
is lower in energy than the (N)CT where the amino group
acts as an electron donor. The N,N dimethyl amino group
has a lower ionization potential than the NH2 group. Thus
in the case of DPAS an energy level reversal occurs which
has an important consequence on the emission spectrum of
DPAS. A unique excitation wavelength dependence has been

observed. In the (N)CT state, the basicity of the

Nahid Shabestary

oxygen of the carboxyl group is expected to be larger than
that of the ground state, but the acidity of the hydroxyl
group is not expected to change appreciably. In this case
no proton transfer is expected for the hydroxyl group as

a result of excitation of the (N)CT state. In contrast, for
the excitation of the (O)CT state, both the acidity of the
phenol group and the basicity of the oxygen of the car-
boxyl group are expected to be increased relative to their
values in the ground state. Experiments show that in the
case of DPAS a longer excitation wavelength (330 nm) leads
to emission at 354 nm while a shorter excitation wave-
length (300 nm) gives rise to an emission at mA6O nm

as well as at 35A nm. In the case of DPAS, we are dealing
with excitation of two different electronic states of the
same conformer. In contrast, the emission of MSA is
varied as a result of excitation of different conformers
which are at equilibrium in the ground state. All these
results have been discussed in detail in Chapter II. Our
studies Show that PAS also exhibits excited state proton
transfer in the solid phase. The crystal structure of PAS
shows intra- and intermolecular hydrogen bridges.

Studies of l-azacarbazole in the solid phase, also
shows that excited state double proton transfer accounts
for the observed emission. The crystal structure,as de-
termined by x-ray diffraction methods shows that the crystal
contains cyclic dimers only. A summary of these results

is given in Chapter III.

 

 

 

Nahid Shabestary

Chapter IV summarizes experimental procedures and puri—
fication of solvents and compounds and Chapter V gives a

brief conclusion with some suggestions for future work.

TO Yasha

 

ii

   

ACKNOWLEGMENTS

 

I wish to express my deep appreciation to my
advisor, Professor M. Ashraf El-Bayoumi for his
guidance, encouragement and friendship during the
course of this study.

I am grateful to professor James L. Dye. for
presiding as my second reader and for the editorial
assistance in improving my writing skill. I also thank
Professor Andrew Timnick and Professor Michael W. Rathke.
for serving on my committee.

To Dr. Paul Hunter goes my thanks for being such a good
friend and coworker throuhg my many terms of teaching.
I would also like to thank Dr. D.W. Ward for his great
help in defining crystal structure.

Many thanks go to Mr. Ron Hass and Scott K. Sanderson
for their expert assistance in electronics. Special thanks
to my good friends and colleagues, Khader Ahmad Al-Hassan

and Kamal Ismail. The author wishes to express her

thanks to the chemistry department of MSU for the
assistantships providing finacial support.
The author is particularly grateful to her husband,

Abbas, for his constant encouragement and understanding.

iii

Chapter

LIST OF
LIST OF
CHAPTER

I.

TABLE OF CONTENTS

TABLES.

FIGURES

I - EXCITED STATE PROTON-TRANSFER

Introduction.

Theoretical Models of Hydrogen—Bond

Proton Transfer in the Ground State

Hydrogen-Bonding in the Excited State

Excited State Acidity and Basicity.

Proton Transfer in the Excited State.

A.

Intramolecular Proton Transfer.
Salicylic Acid and Its Derivatives.
Salicylamide and Salicylanilide
Naphthoic Acid and its Derivatives.

Intermolecular Proton Transfer.
Naphthol-Triethylamine Systems.
3-Hydroxypyrene-TEA System.

7-Azaindole Dimers (Double Proton
Transfer) . . . . . . . .

7-Azaindole-Alcohol Complexes

Kinetic Isotope Effects in Proton
Transfer Reactions. . . . . .

IsotOpe Effect in Excited State
Proton-Transfer in 7-Azindole (7AI)
Dimers. . . . . . . . . . . . . . .

Proton Tunnelling

iv

Page

viii

l3
17
27

28
28

A7
55

6A
6A
71

72
78

82

85
88

Chapter

CHAPTER

Page
Experimental Evidence in 7A1 Dimers . . 9l
Excited State Proton Transfer in the
Vapor Phase, Methyl Salicylate. . . . . . . 92
Excited State Proton Transfer in the
Solid Phase, Acridine Derivatives . . . . . 97
II - EXCITED STATE PROTON TRANSFER IN
P—AMINO SALICYLIC ACID DERIVATIVES . . 103
Introduction. . . . . . . . . . . . . . . . 103
Excited State Proton Transfer in p-Amino
Salicylic Acid (PAS) and its Methyl
Ester (PASE). . . . . . . . . . . . . . . . 105
l. Solvent Effect on the Absorption
Spectra of PAS and PASE . . . . . . 11A
2. pH Effect on the Absorption Spectra
of PAS and PASE . . . . . . . I20
3. Emission Spectra of PAS and PASE as a
Function of pH at Room Temperature. . . 126
A. Fluorescence Spectra of PAS and
PASE in Different Solvents at
Room Temperature. . . . . . . . . . . . l3l
Excited State Proton Transfer in Solid
Phase of p-Amino Salicylic Acid (PAS) . . . 1A0
Excited State Proton Transfer in N,N-
Dimethyl p-Amino Salicylic Acid (DPAS)
and its Methyl Ester (DPASE). . . . . . . . 1A5
l. Solvent Effect on the Absorption
Spectra of DPAS and DPASE . . . . . . . 1A5
2. pH Effect on the Absorption Spectra
of DPAS and DPASE . . . . . . . . . . 1A8
3. Emission Spectra of DPAS and DPASE
As a Function of pH at Room Tempera-
ture. . . . . . . . . . . . . . . . . . 156

Chapter

E.

CHAPTER

CHAPTER

A.

A. Fluorescence Spectra of DPAS and
DPASE in Different Solvents at
Room Temperature.

A Unique Excitation Wavelength Dependence
of Excited—State Proton—Transfer. . .

III - EXCITED STATE DOUBLE PROTON TRANS-
FER IN l—AZACARBAZOLE HYDROGEN-
BONDED DIMERS, ITS ACETIC ACID
COMPLEX AND IN THE SOLID PHASE.

Introduction.
1. Excited State Double Proton Transfer
in 1-Azacarbazole Hydrogen-bonded

Dimers.

2. Phosphorescence Spectrum of l-
Azacarbazole.

3. Excited State Double Proton Transfer
in l-Azacarbazole-Acetic Acid Complex

A. Excited State Double Proton Transfer
in l—Azacarbazole in the Solid Phase.

Crystalographic Data.

Room Temperature UV Absorption and
Fluorescence Spectra of l-Aza-
carbazole in the Solid Phase.

IV - EXPERIMENTAL

Experimentally Studied Molecules.

l. Para Amino Salicylic Acid (PAS)

2. Methyl Para Amino Salicylic Acid
(PASE). . . . . . . . . .

3. N,N Dimethyl Para Amino Salicylic
Acid (DPAS) . . . . . . . . . . .

A. Methyl N,N Dimethyl Para Amino
Salicylic ACid (DPASE).

5. l-Azacarbazole (lAC).

vi

Page

159

163

.170

170

171

181

183

186
186

188
193
193
193

193

193

193
19A

Chapter

6.

l.

 

 

 

7IIIIIIIIIIIIIIIIII---E_________________

Nl—Methyl 1—Azacarbazole Tautomer

B. Solvents.

3—Methylpentane (3MP)
Ethanol

Ether

Water

Cyclohexane

Hexane.

C. Spectral Measurements

l.
2.
3.
A.
5.

Absorption Spectra.
Emission Spectra.
Infrared Spectra.
Mass Spectra.l

NMR Spectra

CHAPTER V - CONCLUSION AND FUTURE WORK.

REFERENCES.

V1].

Page

19A
195
195
196
196
196
196
196
197
197
197
198
198
198
199
20A

Table

LIST OF TABLES

*
pKa and pKa Values for Selected

Molecules

Room Temperature Absorption Maxima

(nm) and Spectral Shifts (cm-l) Relative
to 3MP for PAS in Different Solvents.
Room Temperature Absorption Maxima (nm)
and Spectral Shifts (cm-l) Relative to
3MP for PASE in Different Solvents.
Room Temperature Absorption and Emis-
sion Maxima for PAS at Different pH
Values, Ground and Excited States
Species

Room Temperature Absorption and Emis-
sion Maxima for PASE at Different pH
Value, Ground and Excited State Species
Emission Maxima in nm for PAS in Dif-
ferent Solvents

Emission Maxima in nm for PASE in Dif-

ferent Solvents

viii

Page

19

53

116

120

121

122

13A

13A

Table Page

9 Room Temperature Absorption Maxima

(nm) and Spectral Shifts (cm‘l) Rela-

tive to 3MP for DPAS Acid in Different

Solvents. . . . . . . . . . . . . . . . . . 1A7
10 Room Temperature Absorption Maxima

(nm) and Spectral Shifts (cm-1) Rela-

tive to 3MP for DPASE in Different

Solvents. . . . . . . . . . . . . . . . . . 1A7
11 Room Temperature Absorption and Emis-

sion Maxima for DPAS at Different pH

Value, Ground and Selected States

Species . . . . . . . . . . . . . . . . . . 152
12 Room Temperature Absorption and Emission

Maxima for DPASE at Different pH Value,

Ground and Excited States Species . . . . . 155
13 Room Temperature Emission Maxima in nm

for DPAS in Different Solvents. . . . . . . 159
1A Room Temperature Emission Maxima in nm

for DPASE in Different Solvents . . . . . . 161

ix

 

 

 

 

 

LIST OF FIGURES

Figure Page

1 Absorption spectra of Schiff base
(PLP 7 x 10’5 M, n butylamine 2.5 x
lo-Ll M) in various solvents; di-
oxane, AAA DMF, ———— H2O, the dotted

line represents decomposition of this

spectrum.

2 Potential energy diagrams for the
motion of the proton in a hydrogen
bond A—H...B.

3 Hypothetical potential energy curves
for the formation of a H—bond A-H...B

in Ground and excited state, (1) wg

16

>w (2) wg<we

e’
A Effect of pH on the fluorescence of

an aqueous solution of indole at 25°C . . . 25
5 Absorption and fluorescence spectra

of salicylic acid and 2—methoxyben—

zoic acid in methanol at 20°C . . . . . . . 29
6 Absorption and fluorescence spectra of

methyl salicylate in methylcyclohexane

(

) and fluorescence spectrum of

 

 

F igure

10

11

Page

methyl 2-methoxy benzoate in the same
solvent (---) . . . . . . . . . . . . . . . 31

Fluorescence spectra (uncorrected)

-A 1

of methyl salicy1ate, 10 mole A-
in methanol; at different excitation

wavelengths . . . . . . . . . . . . . . . . 35
Fluorescence excitation spectra of

-A 1

methyl salicy1ate, 10 mole A-

in methanol . . . . . . . . . . . . . . . . 36

Fluorescence spectra of 2 x 10—“

3

mole cm- methyl salicylate in cyclo—

hexane (————), ethanol (——-—), methanol

(-—--), and 2,2,2-trif1uoroethanol (....).

The spectra are shown in relative num-

bers of quanta vs. wavenumber in um-l

The excitation wavelength is 286 nm . . . . 38
Excitation spectra of methyl salicylate

in methylcyclohexane. Solid dots

represent the ASO-nm band, while open

circles represent the 350 nm fluorescence . A2
Temperature variation of the fluorescence
spectrum of methyl salicylate in cyclo-

hexane (1.A x 10-“

M). As temperature
increases the short wavelength emission

intensity raises and the long wavelength

xi

IPjsgure

12a

12b

13

1A

15

16

Page

intensity decreases. Temperatures (°C):
21.2, 29.6, 39.1, A8.3, 57.3, 67.5. . . . . A6

Fluorescence spectra of salicylamide:

(a) 2.A x 10—“ M in ethanol, Aex =

265 nm; (b) 1.3 x 10-5 M in cyclo-

A

hexane, A = 300 nm; (0) 1.0 x 10- M

ex
in water, Aex = 265 nm. . . . . . . . . . . A9
Effect of excitation wavelength on the
fluorescence spectrum of 2.5 x 10-5 M
salicylamide in ethanol: (a) Sex =

290 nm; (b) Aex = 310 nm. . . . . . . . . . 5O
Fluorescence decay curve of 10—3 M

aqueous salicylamide (all emission

above A00 nm) excited at 265 nm; sweep

speed 16 PS channel-l. The upper curve

is a semi—log plot. . . . . . . . . . . . . 52
Fluorescence of salicylamide in D20

(1.7 x 10'“

M). Excitation wave-

lengths (nm): (a) 330; (b) 330;

(c) 265 . . . . . . . . . . . . . . . . . . 5A
Fluorescence excitation spectra for

salicylamide in D20. Emission wave-

lengths (nm): (a) 350; (b) A30. . . . . . . 56
Fluorescence of methyl 3-hydroxy-2-

naphthoate in cyclohexane: (a) 1 x 10'5 M

xii

Figure

17

18

19

20

(b) 1 x 10‘“ M, 16X = 260 nm, (c)
u

1 x 10 M, Aex = 390

Fluorescence excitation spectra of
methyl 3-hydroxy-2—naphthoate in
cyclohexane (a) short wavelength

emission and (b) long wavelength

emission.

Change of the absorption spectrum of

B-naphthol in cyclohexane produced by

added TEA ( 15°C). Concentration

A 1

of B-naphthaol: 2 x 10' mole £—

Concentration of TEA: (l) 0 mole 2-

2

(u) 1.12 x 10' mole 2‘1 (6) 1.80

1 1

x 10- mole £-
Fluorescence spectra of B-naphthol
in methylcyclohexane at different
concentrations of triethylamine (1)
0.000 M; (2) 0.002 M; (3) 0.00A M,
(A) 0.008 M; (5) 0.020 M; (6) 0.10
and 0.15 M; (7) as solvent.
Fluorescence spectra of B-naphthol
in methylcyclohexane +0.1A M tri-
ethylamine at different tempera-

tures O O O O O C O 0 O O O O O O O

xiii

Page

58

59

65

67

69

 

Figure

21

22

23

 

Fluorescence Spectra of B-naphthol under
various conditions. The time—resolved
spectra of the B—naphthol—TEA system in
toluene at 25°C observed 0 ns (curve a)
and 10 ns (curve b) after the fluores—
cence of free B—naphthol reached the max—
imum intensity. The concentration of

B-naphthol and TEA are 1.2 x 10""!4

M and
0.705 x 10—2 M, respectively. Curve

C is the steady state fluorescence
spectrum of the toluene solution of

3—naphthol (0.705 x 10-“

M) . . . . . . . . 70
Room temperature absorption (left)
and fluorescence (right) of 7—azaindole

in 3MP at 1.0 x 10‘5 M (
2

) and 1.0

 

x 10- M (——-). Excitation wavelength
285 nm. Front surface excitation was
used for the concentrated sample.

The signals were arbitrarily adjusted
to give comparable intensity in the F1
region. . . . . . . . . . . . . . . . . . . 75
Corrected room temperature fluores—

cence spectra of 7A1 in EtOH (———-)

and 7A1 in EtOD (— --). . . . . . . . . . . 79

xiv

 

 

IPiTgure

2A

25

26

27

28a

Model barrier for proton tunnelling
calculation

Vapor phase fluorescence spectra (un-
corrected) of MSA at 50°C, equilibrium
pressure under N2, excitation wave-
lengths indicated. The experimental
conditions are identical for both
spectra

Vapor phase fluorescence of MSA,

under N2, as a function of temperature
(indicated in °C). Insert: semiloga—
rithmic plot of the intensity ratio of
short-to-long wavelength emission vs.
reciprocal temperature.

Fluorescence spectra of a solution of
9-(p-hydroxy) phenylaridine in ethanol
at different pH values. Curve l-pH

8; curve 2 - pH 2

Fluorescence spectra of a film of
9-(p-hydroxy) phenylacridine in vacuum
at -180°C. Curve 1 - before eXposure;
curve 2 - after exposure for 10 min to
the total radiation of an SVDSH-250

lamp; curve 3 - after regeneration at

room temperature; curve A - acetic acid

vapor; curve 5 - in pyradine vapor.

XV

Page

90

93

9A

99

99

Figure

28b

29

3O

31

32

33

Page

Potential energy curve scheme for

proton phototransfer in films at

-180°C. . . . . . . . . . . . . . . . . . . 99
Room temperature absorption spectra
of p—nitroaniline in different sol—
vents: (...) methylcyclohexane,

( ooooo ) ether, and (----) ethanol
(concentration in alcohol = 6.52 x
10-5 gm moles lit-l). 108
Room temperature absorption spectrum

of tri-l—naphthylboron in methyl—

cyclohexane (dashed line); and same

after bubbling in dry ammonia gas

(solid line). . . . . . . . . . . . . . . . 109
Qualitatively molecular orbital en-

ergy level for benzene, phenol and

lowest vacant orbitals of substi-

tuent . . . . . . . . . . . . . . . . . . . 111
Room temperature absorption spectra

of salicylic acid (SA) in different

media: (————) N/lO—HCl; ( ----- ) N/20-

NaOH; (-..-.._) 5-n NaOH; and methyl

salicylate in (----) N/20-NaOH. . . . . . . 113

Room temperature absorption spectra

of PAS in different solvents: (————)

xvi

 

Figure

3A

35

36

37

38

39

 

 

Page

3MP, (—-~-) ether, ( ..... ) ethanol, and

(....) H20 (pH = 2) . . . . . . . . . . . . 115
Room temperature absorption spectra

of PASE in different solvents: (————) 3MP,
(--—-) ether; ( ————— ) ethanol; and (-.._..)

H20 . . . . . . . . . . . . . . . . . . . . 119
Room temperature absorption spectra

of PAS at different pH values: (——-—)

0.85, (————) 2.0, ( ————— ) H20 (2.9-12.1);

and (-..-..-) 13.1. . . . . . . . . . . . . 123
Room temperature absorption spectra

of PASE at different pH values:

(----) 1.3; (———~) H20 (2.3-10.3)3

and (....) 11.0 . . . . . . . . . . . . . . 12A
Room temperature fluorescence spectra

of PAS at different pH values: (—-——)

0.85, (————) 2.0, ( ----- ) H20, and

(—~-—--—) 13.1. . . . . . . . . . . . . . . 127
Room temperature fluorescence spectra

of PASE at different pH values: (-——-)

1.3, (————) H20, and (....) 11.0. . . . . . 128
Room temperature fluorescence spectra

of PAS in different media: (————) 3MP;

(----) ether; ( ----- ) ethanol; and

(....) H2O (pH = 2) . . . . . . . . . . . . 132

xvii

 

 

Figure

no

A1

A2

A3

AA

A5

'A6

Page

Room temperature fluorescence spectra

of PASE in different media: (————)

3MP; (----) ether; (-.---) ethanol;

and (....) H20. . . . . . . . . . . . . . . 133
Ground-state conformers of methyl

salicylate: (a) "ether bonded", (b)

"open ring" (trans), and (c) "closed

ring" (cis) . . . . . . . . . . . . . . . . 138
Room temperature fluorescence spectrum

of PAS in KCl pellet. . . . . . . . . . . . 1A3
Room temperature absorption spectra

of DPAS in different solvents: (————)

3MP; (----) ether; ( ----- ) ethanol; and

(....) H20 (pH = 2.A) . . . . . . . . . . . 1A6
Room temperature absorption spectra

of (————) N,N-dimethyl p—amino salicylic

acid, (....) N,N-dimethyl p-amino

benzoic acid, and (----) salicylic

acid in their neutral form in H2O . . . . . 1A9
Room temperature absorption spectra

of DPASE in different solvents:

(————) 3MP; (----) ether; ( ----- )

ethanol; and (-..-..) H20 . . . . . . . . . 150
Room temperature absorption spectra of

DPAS at different pH values: (----) 0.8;

xviii

 

 

 

 

Figure Page
(————) 2.A, ( ————— ) H2O (A.1—12.5)
and (—--—-.-) 13.6. . . . . . . . . . . . . 153
A7 Room temperature absorption spectra

of DPASE at different pH values:

 

(—-—-) 1.1, (————) H20; and (....)

11.9. . . . . . . . . . . . . . . . . . . . 15A
A8 Room temperature fluorescence spectra

of DPAS at different pH values: (————)

0.8, (—~———) 2.A, ( ..... ) H2O, (-.._..__)

13.6. . . . . . . . . . . . . . . . . . . . 157
A9 Room temperature fluorescence spectra

of DPASE at different pH values: (———-)

1.1, (————) H20, and (....) 11.9. . . . . . 158
50 Room temperature fluorescence spectra

of DPAS in different solvents: (———~)

3MP; (————) ether; ( ----- ) ethanol,

(....) H2O (pH = 2.A) . . . . . . . . . . . 160
51 Room temperature fluorescence spectra of

DPASE in different solvents: (————) 3MP;

(-———) ether; ( ————— ) ethanol; and

(....) H2O. . . . . . . . . . . . . . . . . 162
52 Room temperature fluorescence spectra

of DPAS in 3MP at two different excita—

tion wavelengths: (————) Sex = 300;

(____);\ex=330..............1621

 

 

 

 

Figure

53

5A

55

56

57

(a) Energy level diagram for DPAS
(change in emission due to excitation
of different electronic state of the
same conformer)

(b) Energy level diagram for two con—
formers of MSA (change in emission

due to excitation of different con-
formers).

Room temperature excitation spectra

of DPAS in 3MP (uncorrected spectra)

(

while (---—) represents the A6A nm

) represents the 35A nm band,

 

fluorescence.

Room temperature fluorescence spectra
of 1—azacarbazole in 3MP at two dif—
ferent concentrations

Room temperature ultraviolet absorp—
tion spectra of 1—azacarbazole (2 x

10‘5 M) in different media: (———_)

 

cyclohexane; ( ) ether; and (....)
ethanol . . . . . . . . . . . . . .
Ultraviolet absorption spectra of
l—azacarbazole (2.5 x 10_5 M) in 3-
methylpentane at room temperature

( )5 770K (——-—) . . . . . . . .

 

XX

 

Page

166

166

 

169

172

173

. . 175

 

 

 

F7— ’

Figure Page

58 Room temperature fluorescence

spectra of 1—azacarbazole at dif—
feren pH values . . . . . . . . . . . . . 176
’ 59 Effect of pH on the fluorescence
spectra of l—azacarbazole at room
temperature . . . . . . . . . . . . . . . . 178
60 Room temperature absorption spectra
of l—azacarbazole in 3—methylpentane
at different concentration. . . . . . . . . 179
61 Qualitative double minimum potential
energy diagram illustrating excited—
State double proton transfer in 1—
| azacarbazole hydrogen-bonded dimers . . . . 182
i 62 Phosphorescence spectrum of 1—aza—
‘ carbazole (6 x 10_4 M) in ether at
. 77°K excited at 330 nm. . . . . . . . . . . 18A
I 53 Room temperature absorption spectra
of l—azacarbazole (2 x 10—5 M) in hexane,
in presence of variable amounts of
6

acetic acid, 0, 1.66 x 10— , 3.32 x 10'6,

A.98 x 10‘6, 6.6A x 10—6, 8.30 x 10—6
—6

3
9.96 x 10 , 1.16 x 10'5 M in 10 .cm cell. . 185

6A The emission of 5 x 10‘“ M of l—azacar-
bazole in 3—methy1pentane in the absence

(——*—), presence of excess acetic acid

; (—-—-).................,.187

Figure

65

66

67

Stereoscopic View of a hydrogen-
bonded dimer molecule . . . . . . .
Stereosc0pic views of the packing of
dimer molecules into the unit cel
Room temperature fluorescence spec—
trum of 1-azacarbazole in the solid

phase

xxii

Page

189

190

192

 

 

 

 

 

CHAPTER I

EXCITED STATE PROTON-TRANSFER

1. Introduction

 

Proton transfer (PT) representscnuaof the simplest chemi—
cal reactions. Unlike electron transfer which involves mainly
the exchange of charge between reactants, the transfer of
a proton also results in the transport of mass. Because
of their simplicity, reactions involving protons are
ubiquitous. They have been studied in a large number of
systems ranging from the base catalyzed decomposition of hi-

tramide to the dissociation of water.(l’2)

A large pro—
portion of our knowledge about equilibria, chemical kin-
etics, isotope effects, mass transport in solution, free
energy relationship, and chemical reactivity has developed
from the study of proton transfer reactions. Reactions
involving protons are not only very common and instructive
but are found in a number of processes important to our
everyday existence. In biology, protons are involved in a
host of enzymatic reactions, the most important of which

is the synthesis of ATP in mitochondria and chloroplasts.(3)

The rates of proton reactions have been shown to span

at least 15 orders of magnitude.<u) A wealth of kinetic

 

 

 

 

 

 

 

data has produced a good qualitative description of the pro—
ton transfer process in the form of reaction mechanisms.(5)
Detailed theories of proton transfer reactions are still
developing. However, these are often based on informa—
tion obtained from gas phase reactions (i.e., molecular
beam studies) which are free of solvent participation.(6)
NMR, rapid mixing and perturbation techniques have
been the most useful in the study of hydrogen ion reac—
tions and transport. Since stopped flow is limited to
the study of relatively slow reactions, most of our in—
formation about rapid proton transfer reactions such as
those between oxygen or nitrogen atoms is due either to

(7,8)

NMR or to perturbation techniques.

The perturbation methods of course utilize a change
in the equilibrium constant to observe the kinetics.
Equilibrium is often perturbed by a rapid change in tem—
perature or pressure. For a number of reactions, includ—
ing those involving protons, light can be used to quickly
change the equilibrium constant.(9) The perturbation
arises from the fact that the ground state properties of
a molecule are different than those of the excited state.
The excited state properties, such as the dipole moment,
ionic dissociation constant and hydrogen bonding ability
Amy undergo dramatic changes upon electronic excitation.
Anny aromatic compounds in their first excited singlet

states (Sl) display acidities very different from those

 

 

 

 

 

in the ground state due to changes in electronic distribu—
tion. In some cases the pH at which the absorption and
the fluorescence spectra undergo a sudden change differ

by six units or more. These studies indicate clearly that
P.T. takes place during the lifetime of the first excited
state (81).

Since molecules absorb a photon in 10-15 sec, in
principle it is possible to perturb the equilibrium on
this time scale. Thus, it is possible to rapidly ini—
tiate proton release or uptake with the absorption of a
photon.

The study of the energetics and dynamics of excited
state proton transfer is of great interest both experi-
mentally and theoretically. It serves as a model photo-
chemical reaction for which potential energy surfaces may
be calculated. It appears also that excited state proton‘
transfer occurs as a primary step in a number of photo-

(10) The occurrence

biological processes such as vision.
of excited state proton transfer in some enzymes may re-
veal interesting information regarding the nature and

(ll—13) Depending on

structure of their active sites.
the molecular system and the medium, intra— or intermole—
cular excited state proton transfer may occur.

In this review, a brief discussion of hydrogen-bonding

in the ground and excited state is given. Examples of

proton transfer in the ground state as well as in the

 

excited state are discussed. Kinetic data using nano-

and picosecond techniques provide rate constants for ex-
cited state proton transfer and the mechanism of excited
state proton transfer is discussed for intramolecular and
intermolecular proton transfer in derivatives of salicylic
acid and naphthaic acid, 7-azaindole dimers and hydrogen-
bonded complexes. Evidence for the possibility of proton
tunneling and values of energy barrier heights are reviewed.
Examples of solvent effects on the fluorescence spectra

and rate constants of systems that undergo excited state
proton transfer are provided. Finally, examples of excited
state proton transfer in the solid and in the vapor phase

are given.

2. Theoretical Models of the Hydrogen-Bond

 

Various models are used for H-bonding, the subject was

reviewed recently.(lu’15)
In the electrostatic model, the energy of the hydrogen

bond is calculated in terms of dipole interaction. Coul-

(l6)

son and Danielson used a valence-bond approach in
which the hydrogen-bond is treated as a four electron
problem. Other variations of this method were attempted,
nevertheless, the conclusions reached are similar. The
electrostatic forces, the short-range repulsion forces,

and the charge-transfer forces, are all of the same order

of magnitude but are of different signs. Often the second

 

 

 

 

and third effects cancel each other which explains the
success of electrostatic theories. The amount of charge

transferred from 0 to 0A is found to be nonnegligible

B
(but small) for short bonds and negligible for long bonds;
thus the long bonds are essentially electrostatic. The
charge—transfer approach of H-bonding is analogous(17)
t0 the charge—transfer theory of molecular complexes. The
contribution of the charge—transfer term increases as the
ionization potential of the lone pair becomes smaller.
Since a fraction of an electron is replaced on the anti—
AH orbital v, the OAH bond is thus weakened.

This is just what is observed experimentally. The large

bonding 0

migration increases the polarity of the Ofilt..OB systems
and intensifies the VOH IR bands. In the SCF—MO theory,
the complex AH...B is considered as a single large mole—
cule and is treated by the SCF—MO and CI techniques.
Although the electrostatic interaction appears to be the
principal interaction, the dipole—dipole term is also im—

portant.

3. Proton Transfer in the Ground State

A proton transfer reaction in which a proton involved
in an H-bond A—H...B is transferred from A to B, belongs
to the broad classification of acid—base reactions. The
transfer may occur along a potential energy curve similar

to the ones depicted in Figure 2 by two mechanisms, direct

 

transfer over the barrier and/or through quantum mechanical
tunneling through the potential barrier separating the two
minima. The phenomenon of tunnelling will be discussed
later.

The transition from an H-bonded complex A-H...B to a
proton transfer complex A-...H-B+ in a nonpolar solvent is
accompanied by significant changes. For example, in the
IR spectra, the disappearance of vOH(A-H) bands and the
appearance of vOH(B+-H) bands have been observed in the IR
spectra of mixtures of some carboxylic acids ranging from

acetic (the weakest, pKa = A.76) to trifluoroacetic (the)
(18)
3.

Evidence for ground state proton transfers has also

strongest, pKa = 0.23) with pyridine in CHCl

been found in ultraviolet absorption studies. Baba et
a1.(19) investigated the absorption spectra of the p-
nitrophenol-triethylamine system as a function of amine
concentration and found absorption bands corresponding to
an H-bonded complex and a proton transfer complex. The
appearance of two absorption bands of Schiff bases of
pyridoxal 5' phOSphate (PLP) in certain media have been
attributed to the presence of two tautomeric forms: an
enol-imine absorbing at 330 nm and a keto-enamine absorb-
ing at A10 nm. (See Figure 1). The following equilibrium
is solvent dependent, which explains the spectral changes

that occur due to a change of the dielectric constant of

the medium.

 

 

 

 

 

 

Wavelength n m

350 400
' l

 

250 270 300
Y I

Absovbnnc.

.0
A

 

 

 

 

Absorption spectra of Schiff base (PLP 7 x
10-5 M, n butylamine 2.5 x 10_ M) in various
solvents; dioxane, AAA DMF, H2O, the dotte2dO
line represents decomposition of this spectrum.

FiSure l.

 

 

 

 

R R
.2 1+
4? inf, tic \\fi
HC /
’ I o -—‘e-—K /| o
\g ‘\
330nm 4IO nm

The keto—enamine tautomer is formed due to a proton
transfer from the phenolic group to the azomethane nitro—
gen if the enol—imine is intramolecularly H—bonded. How-
ever, if intermolecular H—bonds are formed with the solvent,
a double proton transfer involving the solvent molecules

is required to form the keto-enamine tautomer.

The Double—Minimum Potential

Evidence for a double minimum unsymmetrical potential
energy curve (Figure 2b) where the lower of the two minima
corresponds to an H atom located near the donor atom was
provided by vibrational spectroscopy. In 1959 Bell and
Barrow(21) showed that the first and second overtone O-H
stretching vibrations of ethanol, phenol, and p—nitro—
phenol exhibit a pair of bands in the presence of bases.

The energy of the second potential minimum was shown to

depend on the strength of the proton donor and the acceptor

 

 

 

(a)

symmetrical double minimum

 

 

potential energy

 

v-2 ,‘i
I V=I
v:o .n---

potential barrier

A R(A-B)-—>

(b) (c)

symmetrical single minimum

 

 

 

unsymmetrical double minimum

 

 

 

 

 

 

 

 

 

 

 

 

x >~
a 0‘
b L
0 0
C C
0 O
I5 E
E E 1—‘ :1
2 . 3 :fi g
o V82 0 V=2C 7
° Val a V='
v.0 V30

 

 

 

 

 

 

' A R(A-B)- B A R(A-B)-> B

i Figure 2. Potential energy diagrams for the motion of
the proton in a hydrogen bond A—H...B.

 

 

 

 

 

strength of the base.(22) The same concept was used to

interpret the splitting of the first overtone of the NH2
symmetric stretching vibration band in a number of intra—
molecularly H-bonded ortho-anilines.(23) Some of the
basic criteria for the existence of a double-minimum po-
tential in an H-bonded system, with the second minimum at
about the u = 2 level, are (l) a splitting of the 002

but not the v band; (2) an increase in the magnitude of

01
the splitting with an increase in the strength of the H—

bond; (3) a reduction of the splitting upon deuteration; (A)
a temperature independent intensity ratio for the pair of
bands resulting from the splitting of v02, since both
transitions originate in the ground state. The band
splitting is concentration independent in an intramolecu—
larly H—bonded system, since the effect originates in

the monomer.

Some of the strongest evidence supporting a double-
minimum potential comes from neutron diffraction studies.
For example, neutron diffraction studies of KH2P0u show
the proton as an elongated shape between the two oxygen

A.(2u) This distribu—

atoms which are separated by 2.A9
tion could be interpreted in either of two different ways:
as a centrally located proton with a very anisotropic
motion or, more likely as a disordered distribution of

the protons between two possible positions, one on each

side of the mid—point between the oxygen atoms. In the

 

 

 

 

 

 

case of ice the 0-0 distance is much longer, being 2.76 A,
and the two possible hydrogen positions are much farther
apart (about 0.7A A). Because of this the disordered pro—
tons appear as separate well-defined peaks. Other indirect

evidence of double-minimum potentials in H-bonded systems

(25) from
(27,28)

comes from dielectric saturation experiments,

(26)

data on compressibility, thermal expansion,
the change in the bond energy,(29) change of intermolecu—
lar distance upon deuteration<30> and NMR studies.(3l)

At this point it should be mentioned that the inter—
pretation of the above experiments is not as decisive as
might be desired. For example even one of the best pieces
of evidence, i.e., the split of the IR bands, could be due
to an overtone of the AH bending mode, enhanced by Fermi
resonance with vStr°(32)

One of the early views of H—bonding was the concept
of "mesohydric tautomerism" according to which the H was
thought to resonate very rapidly between two equally prob—
able positions, one near the donor atom, and the other
near the acceptor atom. This concept would require that
the potential energy function for the H—atom either passes
two equal potential minima with a low barrier between the

tWO (Figure 2a) or be symmetrical, with one broad mini—

mum (Figure 2c).

(33)

Hadzi and collaborators have classified hydrogen

bonds of the O -H...O type into four groups. In group

A B

 

 

 

12

A are included the symmetric single minimum H—bonds; in
group B, the symmetric double minimum H—bonds with small
potential barrier, in group C, the symmetric double mini—
mum H—bonds with a potential barrier higher than in group
B; and in group D, the H—bonds with an asymmetric double
minimum potential curve. The H—bonds of type A and B

are uniformly very short, with R (0A...OB) varying from
2.A to 2.6 A. Some typical single minimum H—bonds are
observed in KH maleate, NaH diacetate, KH dibenzoate, KH
bisphenylacetate. The IR spectra are characterized by the

—1

J

absence of a v (O—H) band in the region above 1800 cm

01
the NMR signals are narrow and weak and remain unchanged
at low temperature. The H-bonded compounds belonging to
type B (KHEASOM, NH2H2POM, CaHPOu, NaHCOB) exhibit two
O—H bands in the 1900—3000 cm—1 region, separated by

300-500 cm—l°

5

the PMR Signals are strong and narrow at
room temperature and only slightly broader at —l80°C.(3”’35)
Group C includes various carboxylic acid dimers and group
D includes various phenols.

A very interesting investigation, through ab initio
techniques, of the double well picture of the hydrogen

(36)

bond was performed by Clementi et al. The study was

on the DNA base pair Guanine—cytosine and the motivation

(37) on the possible

was the theory put forward by wadin
importance of quantum mechanical tunnelling for the inter—

conversion between different tautomeric forms. Although

 

 

 

l

r'l

 

13

there was a noticeable shoulder where the second minimum
might have been expected no second minimum was found.
Clementi and coworkers gave various possible sources of
error in their prediction of the lack of a double minimum.
The most important is that no simultaneous motion of two
hydrogen bridges was considered. To test this possible
error source, Clementi and coworkers carried out calcula—
tions on the formic acid dimer<36> which has two hydrogen
bonds. First only the motion of a single hydrogen bridge
was considered. Using the Guanine—Cytosine basis set

only a single minimum was found. When a larger basis

set was used, a very pronounced shoulder was predicted but
no minimum. It was only when the coupled motion of the two
hydrogen bridges was considered and a double—zeta basis

set was used that the calculation yielded a double minimum.

A. Hydrogen-Bonding in the Excited State

A molecule in its first excited singlet state has a
different electronic distribution than the molecule in
its ground state. If the chromophoric portion of the
molecule is involved in H—bonding one would expect changes
in its strength. In this way, electronic absorption and
emission spectroscopy may provide experimental evidence
about H—bonding. Kasha<38> pointed out that absorption
bands corresponding to n + n* transitions are "blue shifted"

in H—bonding media. This is due to the decreases in charge

 

 

1A

density on the lone pair atom as a result of lone pair
promotion. Thus, the H-bond is always stronger in the

(39)

ground state. Bayliss and McRae considered H-bonding

interactions to be a special case of dipole-dipole inter-

(A0)

actions. However, Pimental pointed out that dipole
induced- dipole and dipole-dipole interactions produce
small solvent shifts compared with those due to H-bonding.
He discussed the importance of H—bonding effects compared
with other solvent effects and pointed out the role of the

Franck-Condon principle in H-bonding. Solvent shifts due

to H-bonding can be formulated as follows (See Figure 3):

Av = v — 0 W - W + w

Avf = vf - v = W - W - wg

where:

Av is the energy shift in the absorption maximum,

Av is the energy shift in the emission (fluorescence)

maximum,

v is the energy of the absorption maximum,

is the energy of the emission (fluorescence) maximum,

v is the energy of the absorption maximum in the gas
phase,

W is the enthalpy contribution in the ground state

from H-bonding interaction of the solvent with

 

 

15

the solute,

W is the enthalpy contribution in the excited state
from H—bonding interaction of the solvent with
solute upon excitation; and

w and Wg are the energies implied by the Franck—Condon

principle and are always positive.

Since we and WE are always positive, the shift in ab—
sorption and emission will depend on whether Wg is greater
or less than We.

When Wg > We, i.e., the H—bond is stronger in the
ground state than in the excited state as shown in Figure
3, a blue shift Ava > 0 which exceeds Wg—We by we will be
observed in the absorption spectrum. Similarly, in
emission a shift less than Wg—We by wg will be observed
(either a red or blue shift depending on the electronic
magnitudes of Wg—we and wg).

When We > Wg, i.e., the H—bond is stronger in the
excited state than in the ground state as shown in Figure
3, both absorption and emission spectra will Show a red
Shift, which is less than Wg-We by We for absorption
and is less than Wg—We by wg for emission.

The well characterized H—bonds have energies in the
range 1-7 Kcalmol‘l (350—2500 cm—l). According to the
above discussion, a blue shift in absorption may exceed

the ground state H—bonding energy. hence: the blue Shift

 

ENERGY

Figure 3.

 

16

 

 

 

 

 

 

 

 

 

 

 

 

 

R(A..B)

Hypothetical potential energy curves for the
formation of a H—bond A—H...B in ground and
excited state, (1) Wg > We, (2) Wg < We'

 

 

 

 

 

 

 

17

is expected in the range of 350-2500 cm-1 or even larger
than 2500 cm‘l. But a red shift in absorption should
never be as large as Wg, so should not be larger than
2500 cm‘l.

A red shift in the absorption spectrum indicates that
the H—bonding is stronger in the excited state. This
may indicate an increase in the acidity or basicity de—
pending on the functional group at the chromophore in—
volved in H-bonding. The spectra of nitrogen heterocyclic

bases in which the lowest transition is n + n* have been
found to undergo a red shift upon H—bonding. This is due
to an increase in the charge density of the H-bonded

nitrogen, which increases its basicity and in turn in-

creases the H-bond energy in the excited state.

5. Excited State Acidity and Basicity

In 1931, Weber<ul) noticed that the fluorescence of l—
naphthylamine-A—sulfonate changed color as the pH of the
solution was altered, although no corresponding change
in absorption was observed. Similar observations were
made<u2qu> which were interpreted by F6rster<u5-u7) in
terms of a change in the acid—base equilibrium upon elec—
tronic excitation, i.e., an excited state proton transfer
had occurred.

AS with other chemical prOperties the acid—base

pPOperties of molecules (pK values) depend on the electronic

 

 

 

 

 

 

18

charge distribution. Since charge densities are often
modified by electronic excitation (new species), one might
expect a change in the acidity or basicity of a molecule in
a different electronic statefu8’u9)

Compounds that change their pK values as a result of
excitation can be divided into two groups: one group in
which the pKa value decreases (increased acidity) upon
excitation, such as ArOH, ArNHR, ArN+HRR', ArSH<u2’SO’51)
where Ar is an aromatic group, the other group in which

the pKa value incgeases (decreased acidity), such as
B ' xR \\ <52 53)
ArCOH, ArC=O, Ar—C=N , ArN: ’

For hydroxy naphthalene derivatives, and hydroxy—
and aminopyrene sulphonates, the pH at which the absorption
and fluorescence spectra change is usually different by a
factor of 6 or more and thus the acidity constant in the
ground and first excited singlet state (81) differs by 106
or more (see Table 1). In these cases, a proton transfer
takes place during the first excited singlet state life—
time at a pH value where proton transfer does not occur
in the ground state.

Weller discussed methods of determining the excited-
Singlet state dissociation constants, pK*, from spectro-
scopic data.(9’u3’5u) Jackson and Porter(55) have applied
the same techniques to determine the pK of molecules in
the lowest excited triplet state. Their results showed

that the pK of the lowest triplet is usually close to that

 

 

 

 

l9

*
Table 1. pKa and pKa Values for Selected Molecules.

 

 

 

 

[ pK px* pK—pK*

Phenol 10 5.7 A.3
d—naphthol 9.23 2.0 7.23

‘ B—naphthol 9.A6 2.5 6.95

l B—hydroxy pyrene ‘

‘ 5,8,10-trlsulphonate 7.30 1.0 6.30

i A—ammoniumpyrene

1 (in ethanol) 3.7 -1.8 5.5

I l-naphthoic acid 3.7 7.7 —A.0
2—naphthoic acid A.2 6.6 —2 A

9-anthroic acid 3.0 6.5 -3 5

‘ Acridinium cation 5.A5 10.3 -A.85

__ *_

l~naphthoic acid cation —6.9 1.5 48.A

 

 

 

#
w—kfi‘

 

 

 

 

 

 

20

of the ground state rather than to that of the lowest
excited singlet state.

Various methods are used for the determinations of pK*.

The subject was reviewed recently.(56’57)

a. Fdrster Cycle Determinations

The Fdrster cycle method offers a simple method to
estimate pK* values, the method is based on a thermodynamic
cycle. This method requires a study of the emission and
absorption spectra as a function of pH. The change in
acidity as a result of excitation can be detected only
if the rate constant of excited state proton transfer is
comparable to or greater than the rate of fluorescence
decay. In such cases, the absorption spectrum exhibits a
pH dependence different from that of the fluorescence
spectrum, i.e., the ground and excited states have dif—
ferent pKa values. From the spectral shifts in absorp-
(A3,58)

tion and emission spectra, Weller estimated the

*
Change in the pK value (pKa - pKa) by using the equation:

 

ABS — EEH
*_ :
pK pK 2.303RT ’ Or

* _ _ ‘ _
pKa - pKa = 2.1 x 10 3 (vR - uRH) at 298°K.
In deriving this equation it was assumed that AS* 2 AS.

The equation does not take into account the fact that

 

 

 

 

21

equilibrium may not be reached before deactivation and
that the lifetime of S1 is too short to allow protolytic
reaction to reach equilibrium. It can, in principle, also
be applied to higher excited singlet states.

To obtain pK* it is necessary to know the ground state
equilibrium constant, pK (SO) for the reaction in question
and to have some measure of the energy difference between
the lowest vibrational level of the ground and the ex-
cited state in both the RH and R- forms, 143;, AEl and AE2.

Absorption spectrosc0py is commonly used to determine the

(59)

ground state pK. It is not always easy to obtain a

good value for the energy of O-O transitions except for
those cases in which the molecule shows vibrational fine
structure in its absorption and fluorescence spectra.

(60—62)

There are three approaches which can be used

to obtain a 0-0 energy approximation: (a) the use of
the absorption maxima,(b) the use of fluorescence maxima,
and (c) the averaging of the absorption and fluorescence

(A3,63)

maxima of each form according to:

_ _ abs. fluo.
AE1 - hvl - (hvl + hvl )/2 for RH
AE2 — hv2 — (hv2 + hv2 )/2 for R

One must realize that Av obtained from absorption and
fluorescence maxima are only approximate due to the Franck-

Condon factors in the excited and ground states because of

 

 

 

22

the effect of solvation on 0—0 transition. Since the
electronic charge distribution is significantly different
in the S1 and SO states, the equilibrium orientation of
the solvent cage necessary to keep a minimum of free en-
ergy<6u) must be different in the two states. Reorienta-
tion of the solvent shell does not occur during the absorp—
tion process and thus a Franck-Condom state of higher
energy than the equilibrium excited state is reached.

The energy of a 0:0 absorption band corresponds the alge—
braic sum of the 0—0 transition observed in the gas
phase, the stabilization of the ground state and the
Franck-Condon energy in the excited state. There is

no reason to believe that for two species electronically

as different as a dissociated and an undissociated acid

these solvent effects will be identical and cancel out in

the calculation of A5.

-3

A F

\_.7

Solution Gas phase
Energy of an electronic transition measured in
the gas phase and in solution.

So

 

 

 

In solution, the lifetime of the Franck-Condon state being

much smaller than that of the S1 state, reorientation of

 

 

 

 

 

 

 

23

the solvent cage will occur before fluorescence occurs.
After emission the molecule reaches a Franck—Condon ground
state, the overall result being that fluorescence is
shifted to the red compared to absorption. Therefore
A5 cannot be obtained from the absorption or fluorescence
spectra, only the average of these energies provides, a
"best value" of the energy gap between the ground and S1
states. In general a satisfactory agreement is obtained<63’
65) between methods (a) and (c), but method (b) leads to
larger pK*_(Sl) values that are relatively close to the pK
(SO) values.

The limitations of the Fdrster cycle have often been

(61,66)

discussed. To put the value of F6rster cycle in
perspective, one should state that "as with many other
rough rules in physical organic chemistry, it can serve
as a good guide even if it is not absolutely reliable

quantitatively and it can at least indicate the direction

of the pK—shift".

b. Fluorescence Titration

The change in molecular fluorescence with acidity gives
information about the protolytic behavior of the excited
Singlet state of a compound. A pK* (31) value calculated
from the Fbrster cycle gives an indication of the acidity
range in which a fluorescence change is expected, but,

since no account is taken of the relevant rates of the

 

 

 

 

 

 

 

2A

particular processes involved, the Fdrster cycle does not
indicate whether or not proton transfer in the excited
state is kinetically feasible. Improved instrumentation
has made fluorescence techniques more available and ex-
tended the range of compounds accessible to the method be—
cause of improved sensitivity. The techniques of fluores—
cence spectroscopy are well known.(67-69)

If equilibrium is established during the excited state

lifetime then the pK: and ng values can be determined in

a way analogous to the spectrophotometric determination of
the ground state pK's, by fluorometric titration. For

example, the fluorescence intensity of indole<70> as a

function of pH, shows that at both low and high pH
the indole fluorescence is quenced (Figure A). In the
high pH region the quenching is ascribed to the ionization

of the pyrrolic hydrogen according to the reaction:

@ . on (1’) H20

H

From the midpoint of the titration curve one gets pK*

b
12.3. Exactly the same number was obtained by E. vander

Donckt.<7l)

The excited state ng = 12.3 should be com—

pared to the ground state pr = 16.97(72) which shows that

indeed the pyrrolic hydrogen acidity has considerably

increased in the excited state.

 

 

 

 

 

 

 

25

 

 

 

‘

 

4
Avon: 3.03 E 0.03:

 

4|

 

 

 

. : masmwm

(mun Mmqu) Kigsualm aouaosaionH

 

 

 

 

 

 

 

 

26

In the low pH region, the quenching is considered to

be due to protonation of pyrrolic nitrogen according to

the reaction:

mlb : an“?
‘ H11

H

- 1.8, in very good agree-

In the same way one gets a pK:
- 1.7 obtained by Bridges and

 

ment with the value pK:

Williams.(73)

c. Photopotentiometry
Photopotentiometry<7u> is based on the measurement of
the potential developed between one illuminated electrode

and one dark electrode in a solution. This potential,

which is primarily a function of the various species in
solution, was used to produce data leading to pK: values.

The values found were —2.89 and 12.3 for 2-naphthy1amine

cation and anion, respectively. These values compare

reasonably well with the values obtained (—1.5 and 12.2)(u7)

by using the Fdrster cycle. Other values obtained using

this technique include, 13.5 for the 1-naphthylamine anion,

4.37 and 9.5 for 3—pyridinol anion and cation, respectively.

 

27

6. Proton Transfer in the Excited State

 

In the case of ground state proton transfer in ground
state "tautomerism", both the enol and the keto forms
may exist in equilibrium in solution. Excited state proton
transfer "phototautomerism", may also occur giving rise to
a tautomer with an emission shifted with respect to that
of the original species. The relative intensities of
these two emission bands depend on the efficiency of the
proton transfer process and the quantum yields of the
original molecule and its tautomer. For example, it has

(75)

been found that the shift of the fluorescence maximum
of B-naphthol in 06H6 brought about by the addition of tri-
ethylamine (TEA) is quite large (A000 cm-l) compared to the
shift of the 1Lb absorption band of B-naphthol caused by
hydrogen bonding with TEA which is onlyfifllcm-l. These
results suggested strongly that ion-pair formation, due
to a complete proton transfer rather than just stronger
H-bond formation, occurs in the fluorescent equilibrium
state.

Proton transfer may occur intramolecularly as in

(76-78)

salicylic acid or intermolecularly as in the 7-
azaindole hydrogen bonded dimer.(79’80) A detailed dis-
cussion of various examples of molecular systems that

undergo excited state proton transfer will be given.

 

 

 

 

 

a. Intramolecular Proton Transfer

Salicylic Acid (SA) and Its Derivatives — Studies of
excited state pK values reveal a general trend in which

aromatic alcohols become more acidic, while aromatic ke-
tones become more basic in their excited states.(8l)

This gives rise to an interesting effect when both groups
are present in the same molecule and are situated in

ortho positions. In such cases an intramolecular hydrogen

bond in the ground state is formed in non—hydrogen—bonding

solvents. The enthalpy of formation of such a hydrogen—

bond is estimated from infrared data to be in the range

of 7—8 Kcal mole—ZL indicating a strong hydrogen—bond and

a red shift of emission even larger than in aromatic mole—
cules with just one functional group is observed due to

intramolecular proton transfer. For example, in methyl

salicylate the fluorescence maximum is red shifted by
about 10,000 cm—1 from the absorption maximum, while in
phenol the corresponding energy difference is about A000

cm—l.(76’82) Weller had observed that although the maxima
of the long wavelength absorption bands of salicyclic acid

(I) and 2—methoxy—benzoic acid (11) differ in wavelength

by only about 1000 cm—1, the maximum of their fluorescence

bands differ by 5000 cm.1 (See Figure 5). The fluorescence

maximum occurs in the blue at A3A nm for I but that of the

ether II occurs in the ultra—violet at 357 nm. The large

St0kes shift observed in the case of salicylic acid

 

29

 

 

 

 

 

 

 

 

 

 

 

w l 8
f/fi' [3:0
S‘ [W ,
xr: Met/7000! 20 °C
l * 5
* -—-—l———-1a --~-——+ ;
E 7 "\ IQ
’3: / \ I
u“ / \
l \
05 +— / i
/ ;\
/ .
// I \\
I l x 0
w I l * 8
[(/\I}/[ g0 l
3"" A\0CH3
1r) Methanol 20°C
l l ‘ 5 4
(a. l ,1
13 .
e7 . a :“W’” 9
LP / \ Cu
/ \
l \
\
/’j \
14’ 7 0
77 25 35 q 45
jail 9770’3cm'9—-

Figure 5. Absorption and fluorescence spectra of salicylic
acid an 2-methoxybenzoic acid in methanol at
20°c.(7

 

 

3O
C|>H on
on ocu3
I 11

(76,83)

was interpreted as arising from a tautomer formed

in the excited state via an intramolecular proton trans—
fer. To further confirm these results Weller made similar
observations for methyl salicylate (MSA) in aprotic solvents
where two fluorescence bands occurred at room temperature.
The one at 3A0 nm is simply the mirror image of the lowest
near UV absorption band, while the second fluorescence is
strongly red shifted (m10,000 cm-l) with a maximum near

A50 nm and is much stronger than the other component at

3A0 nm (Figure 6). Proton transfer is driven by a pK

(8A,85)

change of about —6 for the phenolic oxygen and

a pK change which may be as large as +8 for the carbonyl

(65)

oxygen of the carboxyl group. Methyl 2-methoxy-ben_
zoate with no transferable proton emits a single band

at 320 nm with a normal Stokes shift. Furthermore, the
emission of methyl salicylate in 0.1 M NaOH/CH3OH exhibits
(86)

a maximum at A16 nm which corresponds to the anion
emission. Thus the A50 nm emission is assigned to a
zwitterion formed via an intramolecular proton transfer

in the excited state.

 

 

 

 

31

 

 

Rebbve quantuniintensfiytfi fluorescence

 

 

 

 

 

35 30 25 I 20 x 103 cm—1
VVavenunlber

FiSure 6. Absorption of fluorescence spectra of MSA in
methyl cyclohexane (solid lines) and fluorescence
spectrum of methyl 2—methoxy benzoate in the same
solvent dashed); fluorescence spectra after

Weller.< 3)
9C“: pan
c
\ C\
// 9 \\Oe

é—y o g
H

0” on:

9
A = 3A0 nm )em = 3A0 nm

em

 

 

 

 

 

 

 

 

 

32

Intramolecular proton transfers have also been re—

ported for the salicylanilid I,(81387’88) salicylamide

II,(88’89) phenyl salicylate 111,<90) and salicylaldehyde

IV(91>

¢ \k cpo new»;

H
t OH OH

I II

¢H\céo H\c,/O

Go“ on
III N

Several studies designed to investigate the mechanism
of excited state proton transfer led to many interesting
findings. Quenching experiments using 082 (carbon di-
SUlfide)<76) demonstrated that at room temperature the

3A0 and A50 nm emission of methylsalicylate were equally

This led to the speculation that the two forms
the

quenched.
of the excited molecule were in equilibrium, i.e.,

proton transfer reaction in the excited state reached

equilibrium in a much shorter time than the lifetime of

the excited molecules. This set a lower limit of 108 S—1

 

 

33

for the intramolecular proton transfer rate constant.

The relative contribution of the two components to the

total fluorescence was strongly dependent on temperature.

As the temperature was lowered, the intensity ratio of the

A50 nm to the 3A0 nm emission bands increased. From

the temperature dependence of this ratio in hydrocarbon

solvents, Weller(9) calculated a difference in enthalpy

of about AH* = -1.0iO.5 Kcal mole"l (—0.7 Kcal mole—l)(92)

between the neutral form and the zwitterion, and hence
I
)

= 5100 cm- he

from the fluorescence spectral shift (AU
deduced a ground state enthalpy change of 13.5i0.5 Kcal

The activation energy for proton transfer has
1

mole—1

(9A) to be E i 2 Kcal mole_ according to:

been estimated

k = v exp (-E/RT)

PT OH

1A

where vOH is the stretching frequency and is equal to 10

secIl, k > 1/TO (>108 8-1) and T = 77°K where proton trans-

fer is still observed. From A°K experiments an even smaller

limit of E i 0.1 Kcal mole-l is derived. A tunnelling

mechanism was suggested as the actual transfer mechan—
ism.(9u) In acetonitrile, the ratio of the A50-3A0 nm

emission is decreased.(92) This was interpreted in terms

of a shift in the excited state equilibrium to form the

neutral species. Because the neutral form has a large

dipole moment in the excited state compared to the

 

3A

zwitterion, it is stabilized to a large extent in aceto-
nitrile.* This interpretation was modified as will be
seen later.

(86)

Investigation by Klbpffer's group showed that in
hydrogen bonding solvents (methanol, ethanol, water, and
acetic acid) there are two strong emissions with maxima
at A50 and 355 nm, whose relative intensities depend on the
excitation wavelength (Figure 7). At excitation wave-
lengths greater than about 310 nm, the A50 nm band domin—
ates, while at excitation wavelengths below 310 nm, the
355 nm fluorescence is the stronger emission. This sug-
gests that different ground state species are being ex-
cited. The excitation spectra of the two bands are shown
in Figure 8. The excitation maximum of the blue fluores-
cence is at 310 nm, that of the UV component is at 300 nm.
The excitation maximum at 3A9 nm is due to traces of the
phenolate ion generated by ground state dissociation
of methyl salicylate.

The study of multiple fluorescence of 2,6-dihydroxy-

(93)

benzoate showed that the two emissions for methyl
salicylate arise from species which are not in equilibrium
in the excited state and which originate from different

ground state species. In the case of 2,6-dihydroxybenzoate

 

*
This strongly suggested considerable intramolecular charge
transfer in the excited state of salicylic ester.

fluorescence intensity

Figure 7.

f

 

35

300nm ' _ _'

313nm

 

350 1.00 1.50 nm 500
wavelength

Fluorescence spectra (uncorrected) of MSA, 10'“

mol 2’1 in methanol; excitation wavelengths and
emission slit width indicated; spectral excita-
tion slit width 20 nm.(86)

 

36

 

 

 

 

7 w IRE—‘7
a '
2
2
.E
s _, .
C
m
a
S
O
2
l
400

 

 

 

250 300 .
wavelength

Fluorescence excitation spectra of methyl

Figure 8.
10_u mole t—1 in methanol.(86)

salicylate,

 

37

which cannot undergo intramolecular H—bonding through ro-
tation around the ester-ring carbon-carbon bond, exhibits
only short wavelength emission in polar solvents, but in
non-polar solvents it emits only at longer wavelengths.
If there were an equilibrium between the neutral and the
zwitterion forms, one would have expected to observe both
forms, especially in polar solvents in which methyl

salicylate shows two emissions.

0 OH“ ° °OCH3
\ OCH3 \ O
0’" o/H
Methyl Salicylate 2.6-Dihydroxybenzoote

Sandros<9u> found that the ratio of the UV to blue
emission was very sensitive to the hydrogen—bonding ability
of the solvent and the excitation wavelength (Figure 9).
Contrary to the earlier findings,(92) this ratio is much
greater in alcohol solvents than in cyclohexane. The
DOSitioncxfthe short wavelength emission is seen to be
solvent dependent. The fluorescence maximum of methyl—
2—methoxylbenzoate shows similar solvent shifts. The
excitation wavelength dependence results from exciting

different proportions of the ground state conformer.

 

38

 

 

 

 

Figure 9.

 

 

Fluorescence spectra of 2 x 10’“ mole cm—3
methyl salicylate in cyclohexane ( ), ethanol
(— — —), methanol (————), and 2,2,2-trif1uoro—
ethanol (....). The spectra are shown in rela—
tive numbers of quanta vs wavenumber in um“ .
The excitation wavelength is 286 nm. 9”)

 

 

 

 

 

 

39

Sandros proposed that a ground state equilibrium between
cis and trans conformers rather than excited state equilib—
rium is responsible for the relative intensities of the

two fluorescence bands.

OR o
I I: ll
C§9 __f__x c‘OR
A *—
o’ "6 <3
cis trons H

The excitation of the cis form gives rise to a rapid and
virtually complete proton transfer, while the trans form
does not undergo proton transfer.

(95) performed a kinetic study at 25°C

Yasunaga gt a1.
by using ultrasonic absorption techniques and obtained the
rate constants for the interchange of the conformers:

1

kf = 9.5 x 105, kb = 2.6 x 107 s" , x = 0.0365
1

and AH = 2.5 Kcal mole-

Thus, the cis-trans interconversion may be slow on the time
scale of singlet state deactivation. The greater in—
tensity of the short wavelength fluorescence band in H—

bonding solvents can be attributed to stabilization of

 

 

 

 

 

 

 

AO

the trans ground state by hydrogen bonding with the sol—
vent. In non—H bonding solvents, the cis conformer is
the preferred species due to the strong intramolecular H—
bond.

An interpretation of the dual fluorescence of methyl
salicy1ate(87) in terms of an emission from an intra—
molecularly hydrogen-bonded species and from a species in
which the hydrogen-bond is ruptured in the excited state is
very unlikely. The energy difference between the UV and
blue emission bands is too great to be accounted for in
terms of hydrogen—bond breakage.

Kaufmann(96) studied the picosecond time resolved
fluorescence from MSA and SA in order to provide additional
information on the mechanism of proton transfer in the
excited state. Methylsalicylate in methylcyclohexane was

(96) (8 psec pulse) and the

excited with a 26A nm pic0pulse
appearance of A50 nm fluorescence was measured with a
streak camera. It was hoped that they could determine

the rate of intramolecular proton transfer. The re—
corded fluorescence signal was a convolution of the ap—
paratus function with the fluorescence from the sample.
Deconvolution of the data indicated that the transfer
rate constant must be greater than 1011 sec-l. The data
did not change even where the temperature was lowered

to A°K. Replacing the proton with a dueteron and cooling

(96)

to A°K also did not alter the data. Therefore, they

 

 

 

 

 

could not confirm or deny whether the proton transfer pro—
ceeds via a tunnelling mechanism.(97’98) The lifetime
measurement of methylsilicylate showed that the neutral
excited state Species had a fluorescence lifetime in
methylcyclohexane three times larger and in acetonitrile
10 times larger than that of the excited state zwitterion
at room temperature. The rapid formation of A50 nm emis—
sion and the very different lifetimes of the fluorescence
components would also indicate that two species are formed
immediately after the absorption of a photon and that they
originate from different ground state molecules, since the
excitation spectrum of the two emissions is different
(Figure 10). These results indicate that the ground
state equilibrium rather than the excited state equilibrium
is responsible for the distribution of the fluorescence
intensity between two wavelengths.

The case of salicylic acid is more complicated than
that of methyl salicylate since it can exist in several
ionic forms in the ground state and excited state. In
addition, it can form dimers.

Methyl—5—ethoxysalicylate I also shows two f1uores~
cences,<92) but in this case the neutral form is favored
in the excited state. The enthalpy has been determined
to be AH* = +0.9 Kcal mole—l. Thus, increasing the
temperature now favors the zwitterion. Time resolved

studies are needed to unravel the dynamics of such ethoxy

derivatives.

 

 

A2

 

 

Intensity (orbllrory Ul'llIS)

 

 

 

Figure 10. Excitation Spectra of methyl salicylate in
methylcyclohexane. Solid dots represent the
A50-hm band, while Open circles represent the

350 nm fluorescence.

A3

0430‘ c,0

OH

WW
I

Recently, a quenching study of methylsalicylate by
carbon tetrachloride<9o> showed that the quenching of the
short wavelength band is much more marked, the quenching
of the long wavelength band being barely detectable (con—
>176)

trary to the previous result Moreover the short
wavelength emission may contain contributions from two
"slowly" interconverting ground state conformers that are
not equally quenched. Preliminary fluorescence decay
measurements using single photon counting appear to verify
these results. A single exponential decay with a life-
time of 131:10 PS<9O) was observed for the long wave-
length emission band of methylsalicylate in methanol (in
methylcyclohexane the corresponding lifetime is 280
PS).(96) Therefore there exist at least three distinct
ground state conformers: I, II and III.

In non-H—bonding solvents such as cyclohexane, we

would expect 11 to be the major ground state precursor

of short wavelength emission, due to stabilization by

c
2’ "’ 9 i
I
cs9 ccqz c”: ego/C":
o’H 0’" <3
H

the intramolecular H-bond while I is expected to be the
precursor of the zwitterion emission at longer wave-
lengths. In methanol, however, there will be a stabiliza-
tion of structure 111 due to H-bonding with the solvent.
Structure III contributes a significant component to the
short wavelength emission in methanol: such a component

is more quenchable by CClu (lifii’ has a larger Stern—Volmer
constant). Structure 11 contributes also to the short
wavelength emission band. One should notice that the
relative intensity of the short wavelength emission in-
creases as the solvent proton's ability to form H-bonds in—
creases.

(99)

Mori gt g1. concluded from IR data that structure

(I) was the most stable; a fact confirmed in a later

(100)

study which used the CNDO/2 method and gave the cal-

culated stability order I > 11 > III. The calculated ener-

gy differences of (II) and (III) as compared with the

most stable ones, (III), are 1.8 and 12.7 Kcal mole-l

reSpectively- In another study using the INDO<101)

method a difference in energy of 18.3 Kcal mole"l between

 

 

45

I (cis) and III (trans) has been obtained which suggests
that a strong intramolecular H—bond must be present.

The effect of temperature on the fluorescence spectrum
of methylsalicylate in cyclohexane excited at 280 nm also
has been studied by Ford §E_§l-(9O) (Figure ll). The
essential feature is that a decline in the intensity of
the long wavelength band as the temperature is raised is
accompanied by an increase in the short wavelength band.
Weller<92> had earlier observed a similar but apparently
smaller effect. In contrast to the results of Ford et 3;.
Smith and Kaufmann reported that the intensity of the short
wavelength band remains substantially constant with
temperature in methylcyclohexane. Ford et al. have noted
one possible explanation for this discrepancy. The two
fluorescence bands overlap appreciably, so that if spectra
are recorded with a large emission monochromator slit
width then the bands will be less well resolved and an
apparent insensitivity to the relative intensities will
occur.

The detailed interpretation of temperature effects
is quite complex, because the nonradiative rates for each
of the excited state species and the ground state equilibria
between the three conformers are temperature dGPGNdent-
Moreover, the quantum yields of different emitting species
are different. It appears certain however, that the short

. . ' ' in
wavelength emission intenSity increases With increas g

 

‘

t I I'll! . .. ,
Illlll I . ..| ..! I ill-Ill]

>b.m2.w.»..2.-_.

 

 

Figure ll.

 

M6

INTENSHY

 

4bo ' sbo
x/NM

Temperature variation of the fluorescence
spectrum of methyl salicylate in cyclohexane
(1.4 x 10' M). As temperature increases the
short wavelength emission intensity raises and
the long wavelength intensity decreases. Tem—
peratures (°C): 21.2, 29.6, 39.1, “8.3, 57.3,

67.5.

 

 

47

temperature indicating a shift in the ground state
equilibria, since the rates of radiationless processes
seldom decrease with increasing temperature.

The pH dependence of the intramolecular proton trans—
fer in salicylic acid was examined,(85) the following
ionization scheme was postulated for the ground and ex—
cited singlet state. Very similar behavior was found

for salicylamide and salicylanilide.(102)

Salicylamide and Salicylanilide — In contrast to methyl—
salicylate in cyclohexane, salicylamide in cyclohexane
shows only a single fluorescence band at mMSO nm assigned
to the zwitterion species with a Stokes shift comparable
to the long wavelength band for methylsalicylate (ca.
10150 cm-l) (Figure 12a). This behavior is similar to
that of salicylanilide which also gives a single band at
“69 nm(8l) in non polar solvents. In ethanol the fluores—
cence of salicylanilide consists of a single band at #20 nm
attributed to the phenolate anion form. Salicylamide ex—
hibits two emission bands one at 342 nm due to the HGUtPal
form and the other at U26 nm which is a composite band due
to zwitterion and phenolate emission generated by exciting
different ground state conformers (Figure 12b)- In water
at pH 5.6 where salicylamide exists in its neutral form,
The

a Single fluorescence band at 420 nm is observed.

. - ' ra
fluorescence decay curve for salicylamide in wate

 

3'. and flair

148

Ground state

 

 

 

 

 

 

 

 

 

 

a" 0‘
I
C
0: So +H’ gcso +H’
- _H§ _ O
O ‘ OH H
pK(So) = 14 pK(So) = 3'0
(I)?! (IN-I
gcso +3. gcsén
on 41* OH
Lowest singlet state
a" 9“
C
[:1 §O +H’ gc§0 +H’
0. -H O- -H’
pK(S,) = 160
am a“
@Csan .u. aka.
0‘ 'H OH
pK(S,) = 7'0

Ionization Scheme for Salicylic Acid85

“9

INTENSITY

x3
X3

 

 

‘ f i

350 WAVE LENGTH I NM ‘50

Figure 12a. Fluorescence spectra of salicylamide: (a)

2.U x 10.“ M in ethanol, Aex = 265 nm; (b)
1.3 x 10"5 M in cyclohexane, Aex = 300 nm;
(0) 1.0 x 10'“ M in water, A = 265 nm.(89)

8X

 

 

..‘v4 uhv‘

 

 

50

 

).
t
a)
2
w
.-
g
b
a
350 WAVELENGTH I NM 450
Figure 12b. Effect of excitation wavelength on the fluor-

escence spectrum of 2.5 x 10'5 M salicylamide

in egganolz (a) Aex = 290 nm; (b) Aex - 310

nm

 

51

obtained by observing the whole fluorescence band from

400 nm upwards by picosecond spectroscopy, with a streak
camera, is shown in Figure 13. The semi—log plot of

the data (upper data points) clearly shows the nonexpon—
ential nature of the decay. The most obvious explana—
tion of these observations is that the observed decay is
the result of the simultaneous decay of the two species
with lifetimes of l.87i0.20 ns and 0.11:0.01 ns which have
overlapping fluorescence spectra. The fluorescence decay
curves have been recorded by observing different wave—
length sections of the emission, isolated by suitable
filters (Table 2).

The observation of the two lifetimes indicates two
separately emitting species, while the fluorescence and
excitation spectra in water give no indication that this
results from excitation of different ground state species.
Similar results were obtained for acidified solutions.

The possibility that one of the two decay components
observed with a lifetime of 6 ns is due to salicylate
anion(—COO') has been eliminated. However, the phenolate
anion(—0‘) of salicylamide gives a single exponential
with a lifetime of l.9i0.l2 ns which coincides with the
long lifetime component of the decay curve. This sug-
gests that the phenolate anion is a major contributor to
the long wavelength emission band of aqueous salicylamide.

Furthermore, the trend in the value of the fraction of

 

 

 

52

 

 

 

 

 

 

 

 

L
1202!-
23
5
‘5
C)
Q
E 8910»
U)
5 L
408r
r
Z_ 'i'f‘r-‘;_.-.._E.“::i,:__i‘-_"_;;_;, _ .
BD 6.4 8.21

TIME (NSEC)

Figure 13. Fluorescence decay curve of 10—3 M aqueous
salicylamide (all emission above #00 nm) ex-
cited at 265 nm; sweep speed 16 PS channel-1.
The upper curve is a semi—log plot.( 9)

 

 

 

 

 

 

 

53
Table 2.
Wavelength Fraction of
Band Fast Component
400 nm 0.3M i 0.09
400 nm 0.59 i 0.05
>095 nm 0.80 i 0.02

 

 

short-lived component with emission wavelength indicates
that the zwitterion fluorescence lies to the red of that
of phenolate anion. The most plausible explanation of
the results in water is that following excitation there

is an immediate loss of a proton from the phenolate group.
Some of these protons are transferred to the carbonyl
group to give the zwitterion species. The fast decaying
fluorescence can then be attributed to the zwitterion
species, while the slow component arises from the excited
phenolate anion. Whether the transfer of the proton

lost from the phenolate group to the carbonyl group depends
on the ground state conformation of salicylamide molecule
in water is not clear, although this appears reasonable.

The fluorescence Spectra for salicylamide in D O

2
(PD 0.5) at various excitation wavelengths are shown

in Figure 10. In changing from H2O to D20, a large

 

 

51I

INTENSITY

 

 

300 . 460 E03 ' 600
A / nm

Figure 1“. Fluorescence of salicylamide in D20 (1.7 x
10'“ M). Excitation wavelengths (nm): (a)
330; (b) 330; (C) 265.(88)

 

55

increase in the intensity of the 355 nm emission band
relative to that of the long wavelength band is observed.
The zwitterion concentration remains constant on deutera-
tion. The intramolecular proton transfer is expected

to be so rapid that even when slowed by deuteration, the
zwitterion yield is unaffected. The increased intensity
of the neutral molecule emission makes it possible to record
the excitation spectra of the neutral and zwitterion
species (Figure 15). As expected, the ground state

trans conformer of excited neutral salicylamide absorbs

at shorter wavelengths than the cis ground state conformer,

the latter can lead to excited state zwitterion formation.

Naphthoic Acid and Its Derivatives - 3—Hydroxy-2—

naphthoic acid (I) (87’103—105),and the corresponding

methoxy derivative (11) have been investigated.
COOH COOH

OH OCH3
I II

The infrared (and NMR) data suggested a relatively strong
(81)

intramolecular hydrogen bond. The ultraviolet absorp—

tion spectra of these derivatives show this bond to be

red shifted compared with the corresponding methoxy—

derivatives which lack the hydrogen bond.(105'107)

 

 

 

 

 

56
b
a
).
t
w
z
u:
f...
Z
I 1 I fi
260 300 340
K/nm
Figure 15. Fluorescence excitation spectra for salicyl—

amide in D20. Emission wavelengths (nm):
(a) 350; (b) H3O.(88>

 

 

57

The fluorescence spectrum of methyl-3—hydroxy-2—
naphthoate is the subject of some controversy. A fluores—
cence band at ca. H20 nm, corresponding to a normal Stokes
shift (3200 cm‘l) of emission is observed. Naboiken(105)
and co-workers have observed a weak, long wavelength emis-
sion (ca. 650 nm) in hydrocarbon solvents in concentrated
solutions where there is a noticeable quenching of the
short wavelength fluorescence band. The interpretation of
this band at 650 nm as being due to impurities, dimers or

excimers was refuted;(105) instead it was attributed to an

intramolecular proton transfer. The fluorescence decay
of this band<108) gave a single exponential lifetime of
60e6 PS. " i

The 650 nm band has not been observed by some
workers<77’lou) in either the methyl ester or in the free
acid. More recently, a band at 605 nm was observed<lO8>
(Figure 16). As we see in Figure 16 this band is more ap—
parent in concentrated solutions and at longer excitation
wavelengths.

The fluorescence excitation spectra (Figure 17) of the
short wavelength emission has two maxima and lies at shorter
wavelength than the featureless excitation spectrum of the
608 nm emission band. Neither excitation spectrum bears
a close relationship to the absorption spectrum, but this

is hardly surprising, since the absorption spectrum is a

superposition of the spectra of at least two distinct

 

 

 

 

 

58
a
?\
ll “
(\b \
t f) I
Q I I
If 1’ I
g I \
I \
I \
I “
I
\
I, X
'I \
I \\
I
I
I
l
I
I
’I
400 r 560 ' 6730 700
A / nm.
Figure 16. Fluorescence of methyl 3-hydroxy-2—naphthoate
U

(a) 1 x 10'5 M, (b) l x 10‘ M,
390.(108)

in cyclohexane:
Aex = 260 nm, (c) l x 10 M, Aex

INTENSITY

59

 

 

 

300

Figure 17.

350 360 aéo 460
X/nm

Fluorescence excitation spectra of methyl
3-hydroxy-2-naphthoate in cyclohexane (a)
short wavelength emission and (b) long wave-
length emission.(10

 

60

ground state species, as evidenced by the excitation spectra.
The long wavelength emission can be attributed to a
zwitterion species formed by intramolecular proton trans-
fer in the first excited singlet state. The ground state
precursor of the zwitterion is an intramolecularly H-

bonded conformer, III. The anion of the ester emits at

0-5 - by 0 +
(:40 ——'~ “OH
I I
III
0
OH
0/
\
OCH3

18,200 cm-1 (550 nm) in alcoholic NaOH, corresponding to an

absorption maximum near 25,000(105) cm_l.

The short wavelength emission band of methyl—3-hydroxy_
2-naphthoate in cyclohexane is characterized by a non—

exponential and concentration-dependent fluorescence

A

decay.(108) At a concentration of 10_ M, the lifetimes

that best fit a double exponential decay function are
23.3:0.2 ns and 5.1i0.9 ns. As the concentration is

increased, the longer of the two lifetimes is progressively

 

’/

61

reduced. This observation suggests that there are two
distinct contributions to this short wavelength band.
Quenching evidence for a dual contribution to the short
wavelength emission of methyl salicylate has been

(90)

given.
The interpretation is in terms of two interconvertible

ground-state conformers. Interconversion is sufficiently

slow that it does not occur during the excited state life-

time. The suggested conformers are ones in which the

hydroxylic proton is H-bonded to the "ether type" oxygen

of the carboxyl group (IV) and an "open-ring" form contain-

ing no intramolecular H-bonds (V).

'I'
O 0
\” OCH
C’OCHS c’ 3
II II
0 O

N Y

The 23.3 ns lifetime can be attributed to conformer IV

and the 5.1 ns lifetime to V. A calculation of the follow-
ing equilibrium constant from absorption data (AH = 5-
Kcal mole-l) yields a value of approximately 5 x 103.

The equilibrium should therefore be almost entirely in
favor of III. On this basis then, the absorption Spec-
trum would bear a close similarity to the excitation

spectrum of the long wavelength emission band which is

 

 

62

'I
O OxH
O ————s :
(3" (”*3 ‘:¢5CI
g l
ocu-I3

Y III

clearly not the case. McTigue(109) has shown that H—bonds

involving ether oxygen and carbonyl oxygens are of similar
strength for intermolecular H bonding between H20 and
various organic ketones and ethers in 001” solution.
Thus conformers IV and III will be of similar energy.
The absorption spectrum is virtually independent of tem—
perature between 293 and 3280K in cyclohexane and between
160 and 280°K in methylcyclohexane. This is consistent
with conformers III and IV being energetically similar,
with conformer V making up only a very small proportion
of the ground state mixture.

The study of 3-hydroxy—2—naphthoic acid(77) (free
acid) shows that an anomalous long wavelength fluorescence

appeared only in basic solvents or proton accepting sol—

vents such as acetone, acetonitrile, and in aprotic sol—

vents doped with small amounts of basic solvents, such

as a mixture of toluene and pyridine,(77’lou) whereas in

benzene only normal fluorescence with its maximum near

23700 cm—1 (420 nm) is observed. Two fluorescence bands

have been observed, one near 2M000cm”l (416 nm) and the

 

 

63

other at 19000 cm"1 (526 nm). These two bands have been

interpreted by Hirota in analogy to those of methyl salicyl—
ate. From the temperature dependence of the relative in—

tensity AH* = —3.1 Kcal mole"l has been calculated.(77)
Recently the hypothesis of true intramolecular proton trans-

fer in the excited state of 3—hydroxy—2-naphthoic acid

l.<lou>who employed fluores—

has been refuted by Ware gt

cence quenching and singlet—state lifetime measurement

techniques in a variety of solvents to show thatphotoauto—

merization in the lowest excited singlet state involves

intermolecular H-bonding with the solvent in the ground
(10A)

electronic state. According to Ware, pyridine

acts as static quencher by forming a ground state com-

plex. The proposed mechanism of the anomalous fluorescence

.H- ’\
C

0 Ce
: V— H?—
0’“ ' o’ o-

of this acid consists in the formation of the H—bonded
complex, followed by absorption of a photon and intra—
molecular proton transfer. Thus the enthalpy (AH = -3,1

Kcal mole—l) measured by Hirota should be attributed to
the enthalpy of formation of this complex. At high con-

centrations of this acid complex spectral behavior is

 

 

T——"—

64

observed due to association and to quenching by H+

(103)

ions.

 

The effect of pH on the absorption and fluorescence
spectra of 3-hydroxy—2-naphthoic acid has been investi—

gated.(110)

Although the ground-state behavior was similar
to that of salicylic acid, the molecules differed in the
prototropic reactions of the S1 state, phototautomerism
occurred only partially in the anion and not at all in the
neutral molecules. Similar experimental evidence has been
obtained for l—hydroxy—2-naphthoic acid and 2—hydroxy—l—

naphthoic acid.(lll)

In these cases the neutral molecules
undergo intermolecular phototautomerism to form the

zwitterion but, unlike 3—hydroxy—2—naphthoatethe anions

 

do not react in this way.

b. Intermolecular Excited State Proton Transfer

 

Naphthol—Triethylamine Systems — Nagakura and Gouter-
man<ll2>

 

have demonstrated that both u- and B—naphthol
form quite strong H—bonds with triethylamine (TEA). The

change of absorption caused by H—bonding in the case of the

 

B—naphthcfl-TEA system in cyclohexane is shown in Figure
18. The ground and excited state equilibrium constants
for this system in cyclohexane have been estimated to be

180 and 3300 respectively. The dipole moment<ll3>

of
naphthol, its hydrogen-bonding ability and its acidity

are greater in the first excited singlet state than in the

 

 

 

65

   
 

 

0.5- (I) (3)
04--
0.3-
0.2-
0.!-
m
I I f F f r
300 310. 320 330 340 350
Tn
Figure 18. Change of the absorption spectrum of B-

naphthol in cyclohexane produced by added

TEA (N150C). Concentration of B—naphthaol:

2 x 10‘“ mol 2-1. Concentration of TEA: (1)
0 mole 2'1, (2) 1.12 x 10‘2 mole 2‘1 (3) 1.80

x 10-1 mole £‘l.(75)

 

 

 

 

66

ground state (pKa= 9.1 and pK: = 2.8).(llu) These facts
have been ascribed to the increased migration of n-elec—
trons on the oxygen to the vacant w—MO's of the naphthalene

ring in the excited electronic state.<113’115’116)

In
View of the above facts, one might expect proton transfer
in the excited state of the naphthol—TEA system in non-
polar solvents.

The change in the fluorescence spectrum of B-naphthol
caused by the addition of TEA in methylcyclohexane<92)
is shown in Figure 19. Up to about 0.1 M TEA, the change
undoubtedly is due to hydrogen-bond formation. However,
the broadness of the spectrum observed at and above 0.1 M
triethylamine, when practically all excited naphthol mole—
cules are complexed, indicates that this spectrum is due
to more than one emitting species. Comparison with the
fluorescence spectrum in pure TEA suggests that the con—
tact ion—pair (III) is the other component. The shift in
the fluorescence spectrum due to ion—pair formation is
about 5500 cmfll. The observed phenomenon is interpreted
as due to a strong charge—transfer from the nitrogen non—
bonding orbital to naphthol in the Franck—Condon excited
state, which is followed by almost complete proton trans—
fer from oxygen to nitrogen in the excited equilibrium
state. Fluorescence quenching with oxygen showed that

(92)

both components are equally quenched. This indicates

that the proton transfer equilibrium between II and III

 

 

 

ID I-

h'b\

 

67

 

Ind

 

Figure 19.

 

 

 

Fluorescence spectra of B-napthol in methyl-
cyclohexane at different concentrations of tri-
ethylamine (1) 0.000 M; (2) 0.002 M; (3) 0.00M
M, (A) 0.008 M; (5) 0.020 M; (6) 0.10 and 0.15
M; (7) as solvent.(9

 

 

I'——7—

68

is fully established in a time shorter than the lifetime

of the excited molecules.

 

1.5 +

ArOH + NB + Ar*OH...NR Ar*O-...HNR

3 3

“H 03*
W

(l)
I 11 III

Lowering of the temperature, however, changes the
fluorescence spectrum of the H-bond complex in favor
of the ion—pair III, as shown in Figure 20. From the
intensity ratio at separate frequencies (22,500 cm"1 and

28,000 cm-l) the equilibrium constant, K* has been found

 

23
to be 1.5 and from the temperature dependence of K53,
the enthalpy, AHE‘3 of the excited state equilibrium is

-0.9 Kcal mole—l.
The interaction of B-napthol in the excited state with
triethylamine was studied by measuring nanosecond time—
resolved fluorescence spectra and by analyzing the decay
fluorescence bands due to the free and complexed proton
doners. The time—resolved fluorescence spectrum has a

new band at 010 nm in addition to the band of the free

B—naphthol molecule at 355 nm. The former corresponds to

 

the band found in the steady—excitation fluorescence
spectrum of the B—naphthol-TEA-benzene system and ascribed
to ion-pair formation by Mataga and Kaifu.(75’ll7) This
band decays more rapidly than the band at 355 nm (Figure

21).(118) (119)

Using the Stern—Volmer equation, the

 

 

K

V

'5’"

 

69

 

 

 

 

 

I I l
Methylcyclohexane
I-0 -
‘7 0-5 —
0
20 25 30
v003mnfi)

FiSure 20. Fluorescence spectra of B—naphthol in methyl—
cyclohexane +0.1“ M triethylamine at dif—
ferent temperatures.(92

/

 

>«.m..Cqu«C_

 

QUCWUmuachD ~ 1

, x

 

7O

0
'a
°o
3
o
3
O
'0

Intensity

Fluorescence

 

 

 

'400' 450
7((nm)

Figure 21. Fluorescence Spectra of B-naphthol under various
conditions. The time-resolved spectra of the
B—naphthol—TEA system in toluene at 25°C ob—
served 0 ns curve (a) and 10 ns (curve b) after
the fluorescence of the B—naphthol reached the
maximum intensity. The concentrations of B-
naphthol and TEA are 1.2x10'u M and 0.705x10-2 M
respectively. Curve 0 is the steady state a
fluorescence spectrum of the tolgene solution
of B—naphthol (0.705x10‘” ) ll )

 

 

 

 

71

lifetime of B-naphthol and the rate constant for complex
formation in the excited state were determined at 30°C
to be 11.6 ns and (2.17i0.01) x 109 M-ls-l, respectively.
From the temperature dependence of the rate constant an
activation energy Ea and a frequency factor A were de-
termined as follows:

1.6u¢0.11 Kcal mo1e"l

U1
ll

10 —lS—l

M

11>
II

(3.05i0.12) X 10

This Ea value is comparable to the activation energy for

the diffusion of toluene as determined from the temperature

coefficient of its viscosity.(120) The rate constants ob-
tained from analysis of the fluorescence quantum<l2l’ll6d)
(122,123)

yield and by using Somoluchowski's equation

—1

are 2 m 3 x 109 IVI—ls—l and 3 m 8 x 109 M-ls , respectively,

which are in good agreement with the results of time-re—

solved fluorescence spectra.

3—Hydroxypyrene—TEA System - The change in the fluores—
cence spectrum in methylcyclohexaneI92) upon lowering the
temperature is much less significant in this case than

with B—naphthol and can be interpreted as a slight Shar—

pening of the structure. Thus AH§3 is either around
zero or strongly positive. The latter assumption implies

that the spectrum in methylcyclohexane is solely due to

 

 

 

72

the form (II). In toluene, evidently the ion—pair (III)
is favored over from (I) and the same is true in pure tri-
ethylamine. The relative values of the equilibrium con-

in methylcyclohexane, toluene and triethylamine,
l

stant K*

23

calculated from the intensity ratio at 21,000 cm- and

25,600 cm_l, are <1, 10 and 7.6, respectively.

7-Azaindole Dimers (Double Proton Transfer) — Bi—

protonic phototautomerism in 7—azaindole hydrogen—bonded
dimers was first discovered by Taylor, El—Bayoumi, and
Kasha<lgu> when they examined the fluorescence spectrum
of 7—azaindo1e (7AI) in 3-methylpentane (3MP) as a func—
tion of concentration. It was known from earlier work

by El—Bayoumi<l25) that in carbon tetrachloride an ap—
preciable fraction of the 7AI molecules is present in the
form of doubly hydrogen bonded dimers. The equilibrium
constant for this self—association was estimated to be

118 M_1 based on an infrared study in C01“, at 35°,(125)

1.8 x 103 and 2.5 x 102 M_1 in 3MP and CClu, respectively,

at 25°C, based on absorption spectra,<79) at 315 nm. These
correspond to AGO values of —A.5 and -3.3 Kcalmol“l at room
temperature. The lower association constant in CClu can
be rationalized on the basis of the high polarizability of
this solvent relative to 3MP. By studying the effect of
temperature on the association constant in 3MP, AH = —9.6

Kcalmol-l and AS = —18 e.u. have been estimated. Using

 

 

73

the value of 1.8 x 103 M"1 for the association constant in
3MP at 25°, only about H% of 7AI molecules are dimerized

5 M while 85% are dimerized at 10.2 M. Thus the

at 10-
solid curve in Figure 22 is due mainly to the monomer
while the dashed curve represents the dimer. Because of
the large negative value of AH, the extent of dimerization
will increase dramatically on cooling. This was confirmed
by measuring the absorption spectrum in 3MP at 77°K, where

u and 10'2 M.

the spectra were identical ath-
In Figure 22, the room temperature fluorescence spec-
trum of 7AI in 3MP is seen to consist of two fluorescence

bands. The first band, F has a maximum at about 325 nm

1,
and is attributed to the normal fluorescence of the mono-
mer with possibly a small contribution from the dimer.
The second band F2, has a maximum at about A80 nm and an
intensity relative to F1 which depends on concentration,
excitation wavelength and temperature. This second band
was attributed by the authors to fluorescence from a new
tautomeric species which forms in the excited state of the
dimer by the simultaneous intermolecular proton transfer
of the two pyrrolic hydrogens, as shown below. The authors
took great care to eliminate alternative explanations of
F2 such as impurities, excimers, and photodecomposition
products.

Further evidence supporting this mechanism was pro-

(126)

vided by comparing the fluorescence spectrum of 7AI

 

 

7M

/ / hv \ \
2 \~ I '\ :==3 \ I ‘\ I (e
N N .N N N
,1, (= l I =
H H '5 H
I I y) 1 I
N N\ N~ N
\ I / \ \ I

]I III

with that of its two N-methyl derivatives IV and V.

"'| \ |\‘ \

\ §

”- 'I' ”I N
CH3 CH3

N I

The fluorescence spectrum of IV exhibits only the normal

fluorescence analogous to F in the 7AI system. In par—

1
ticular the V compound, which has a formal electronic
structure similar to Species III, gives a fluorescence
spectrum very similar to F2.

An important question was the mechanism by which this
double proton transfer takes place (whether the reaction

is thermal or proceeds through quantum mechanical

 

 

 

z

'-

75

use manmpmdsoo o>fim op ompmznom zaflswspflnpw ope: mfimcwfiw one
pouwspcoocOo esp pom com: me: coapwpfioxm oommLSm pcopm

Im>m3 coflumpfloxm

.AIII-V z muoa x o.H use A

 

.conoL am one CH mpfimcmp
.mHQEwm
.s: mmm summed
v z muoe x o.H em AZm ea

oaoucfimmmus mo Apcwfipv oocmomoposam pew Apmoav cofipapomnmlmpzpmpodsop Eoom

:27: Ikozw-_w>§§

 

0mm. 000 On. v 00v Om m
a a d \\ . a
7/1 ex
/ \
.t x, x
C I x
mm /, ~ \
C I m x
m“ ,/ \
R / \
O z
m“ ,x
c. // \\
F. //l\\
N”
T
.A
.L
F.
R

 

 

 

 

 

 

 

 

 

.mm messes

(E- 0| " I NOLLDNLLXB HV'IOW

 

76

tunnelling). By studying the effect of temperature on
the relative intensities of the normal violet fluores—
cence (F1) and the green tautomer fluorescence (F2) in
3MP, a value of 1.4 Kcalmole"l (500 cm‘l) was obtained
as the potential energy barrier for double proton trans-
fer. It was found, however, that the relative integrated

intensities of the tautomer and dimer fluorescence are the

same at 77 and H.2°K,

If the double proton transfer is a thermal process in
which the two proton move over the energy barrier, then

by assuming an Arrhenius behavior with a pre~exponential

factor of 1012 _ 1013 sec—1, characteristic of unimolecular

reactions, an energy barrier of the order 20 - 30 cm-1

for the reaction at H.2°K is obtained.<80) Thus at this

temperature thermal energy is not sufficient to allow ap—
preciable transfer over the barrier. However, substantial
F2 is observed. Clearly these observations pointed strongly
to quantum mechanical tunnelling. At higher temperature
the proton transfer can take place by both tunnelling and
thermal activation in which the two protons move over the
energy barrier from high vibrational levels populated ac-

cording to a Boltzmann distribution, while at very low

temperature tunnelling may be the only reaction pathway.

 

 

 

 

77

In order to obtain further support for the occurrence
of proton tunnelling at low temperature, a sample of 7AI
in which the hydrogen atom at the N1 position was replaced
by deuterium was used.(79) At 77°K where proton transfer
over the barrier is presumably negligible, there is a
large effect on the F2/Fl ratio. It was noted that
deuterium substitution enhances Fl at the expense of F2
indicating that proton transfer is less efficient in the

deuterated sample. The ratio of proton tunnelling rates

for the deuterated and nondeuterated compounds was 2.9.

h e
kPT _ F2 h b1 d
—— — (—-) (-—)
kd F1 F2

PT

Recently Bulska and Chodkouska<lz7> have studied the
effect of temperature on the F2/Fl ratio by exciting the
dimer specifically and they found that the ratio F2/Fl
approaches zero at low temperatures where the phosphor—
escence was not yet observable. At low temperature the
tautomeric fluorescence overlaps the isoelectronic mono—
meric phosphorescence. From these results they concluded
that a tunnelling mechanism is not required to account
for the observed experimental effects. The role of proton
tunnelling in excited—state proton transfer is still an

open question that requires further study.

 

 

 

 

78

The dynamics of double proton transfer in the excited
state of the 7AI hydrogen—bonded dimer has been studied<80>
in our lab by nanosecond time resolved spectroscopy, the

rate constants for proton transfer in dueterated 7AI

(Nl-D-7AI) at 77°K are

k? = 1.9 x 108 sec-1
xi = 1.2 x 107 see‘1

A picosecond study of the 7AI dimer by Eisentha1(128’129)
shows dependence of the rate constant for excited state
proton transfer upon wavelength which offers an explana—
tion of the relatively slow risetime (A ns) for the tauto-

mer fluorescence observed in dueterated 7AI at 77°K.

7—Azaindole—Alcohol Complexes — The fluorescence
spectra of 7AI in alcohols are composed of two bands com—
pletely analogous to those in hydrocarbon solvents.(12u)
Besides the normal fluorescence band of 7AI, Fl’ another
broad (excimer like) fluorescence, F2, appears with a
maximum at A50 nm (Figure 23). In the case of alcoholic
solvents, the relative intensities of F1 and F2 depend
only on temperature and not on concentration and excita—
tion wavelength. The F2 fluorescence in alcohols is

not due to excimer formation since the F2/Fl intensity

ratio is independent of concentration.

 

 

 

 

>~.U:roo:.

to your. vast-sou... s

 

 

 

79

Fluorescence Intensity

 

K

 

I I I I I
3“) 350 400 450 500 550 600 650
A(nm)

Figure 23. Corrected room temperature fluorescence
spectra of 7AI in EtOH ( ) and 7AI in
EtOD ( ————— ).

 

 

 

 

80

Taylor gt gt. attributed the F2 band to a tautomer
resulting from a double proton transfer in the excited
state between 7A1 and an alcohol molecule in an analogous
way to the tautomerization of H-bonded 7—azaindole dimers

in hydrocarbon solvents. The tautomer is the same in both

cases and only the mechanism of its production is dif—

ferent.

k

/ I \ __f_. \ \
‘—— \

5‘?! 1‘ kb \Wq !!

: | 5

fix I“ H'- [H

Fl ('5 If IL"'2

Et fit

The effect of pH, temperature and solvent deutera-

tion on the fluorescence spectra and quantum yields of 7A1

and other model compounds in ethanol<l30> gives some evi-

dence supporting the above mechanism of excited state pro—
ton transfer between a 7AI molecule and an alcohol molecule.
One may visualize two kinds of complexes between 7AI
and alcohol: a 1:2 complex (I) or a cyclic 1:1 com—
plex (II).(131)
When small amounts of ethanol are added to a dilute

(1.0 x 10_u M) solution of 7AI in 3MP at room temperature,

 

 

 

81

’I \ ’I \

‘\ ‘\

~ 7 .. 1,1
o/H H (moi;
é ’0—C2H5 I
”52 H CH

the absorption spectra are perturbed in the same way as
during 7AI dimerization (red shifts). The absorption

at 310 nm is due to primarily to the complex and can be
used to determine the stoichiometry and the association
constant of the complex. The results confirm the assump-
tion of a 1:1 complex and an equilibrium constant of

49 Mn1 is obtained. The association constant of the al—
cohol complex in 3MP is much smaller than the self-associa—
tion constant of 7AI in 3MP which is 2.5 x 102 M‘1 at 25°C.
This has been attributed<79> to cooperative effects present

in the self-dimer and is absent in the case of the 7AI—

ethanol complex.

 

82

c. Kinetic Isotope Effects in Proton Transfer Re—

actions

Kinetic isotope effects are particularly common in the
case of hydrogen isotope effects, partly because hydrogen
is involved in so many reactions, and partly because such
effects are much larger for hydrogen atoms than for the
isotOpes of heavier atoms. It is interesting to note
that the rate differences between reactions of hydrogen
and deuterium compounds are sometimes so large that the
use of deuterium compounds has been proposed as a practical
means for slowing down harmful reactions.

The starting point for most discussions of isotope
effects is the implication of the Born—Oppenheimer approxi—
mation that nuclear substitution leads to no appreciable
change in electronic energy. Of the other factors lead—
ing to energy differences between molecules it is a feature
of primary hydrogen isotope effects that zero—point energy
changes are usually much more important than differences
in isotopic partition functions. The formalism of the
kinetic isotope effect has been given by many workers.(l32)
Within the framework of transition—state theory, kinetic
isotope effects may be expressed in terms of differences
and ratios of isotopic zero point energies and partition

functions between reactants and the transition state, i.e.,

kH/kD = MMI-EXC-TUN°epr(A€R — A€#)/2KT}

 

 

83

where MMI (masses and moments of inertia) and EXC (vi-
brational partition function) include reactant and transi—
tion—state partition functions, and A55 is the zero-point
energy difference between transition states containing
H and D. TUN is the 'correction' for quantum—mechanical
tunnelling through the energy barrier to reaction. There
are two qualitative manifestations of the importance of
zero—point energy changes in the primary isotope effect.
The first is the wide applicability of Swain's relation—
ship<l33) between deuterium and tritium isotope effects:
kH/kT = (kH/IzcrlJl‘LHIl2 (independent of the force constant

change). The second factor is the common observation of

an Arrhenius temperature-dependence,
k /k = (A /A ) expt-(EH - ED)/KTJ
H D H‘ D a a

with AH/AD close to unity.

These results should, of course, be put in perspective.
Swain's relationship is not very sensitive to the pres—
ence of factors other than zero-point energy,(l3u-l36) and
deviations from linear Arrhenius dependences and values of
AH/AD < l are found and normally are taken as indicating
(137,138)

the presence of tunnelling. Westheimer pointed

out that for the two cases of kH = kD and kH >> kD, the
preexponential factor is near unity. The factor that

determines the isotope effect is then just the difference

 

 

84

between the zero point energies.

The substitution of D for H is an elegant probe to use
in the elucidation of chemical dynamics. In particular
the solvent isotope effect is a powerful technique for
investigating the role of the solvent. When H20 is re-
placed by D2O it is inevitable that we are dealing with
the effects of substitution at many different sites.

The effects of these substitutions can be described by
fractionation factor theory(l39) and, because many sites
are involved, more information can be found if measure—
ments are made in equimolar mixtures of H20 and D20 as well

O and in pure D20.(140) The two forms of

(1A1—1A3)

as in pure H2

water differ in structure and differences exist

as evident from differences in thermodynamic prOperties
of the two liquids, in the heats of hydration and solu—

(144,145)

bilities of salts in the two forms of water.

Swain and Bader<lu6’lu7) prOposed a model which explains
the solvent isotOpe effect in terms of zero point energies
associated with hindered translations and hindered rota-
tions (librations) of H20 and D20 (as well as the internal
vibrations). The hindered translations are of very low
frequency (m170 cm_l) and are treated classically. The
librational frequencies of H20 and D2O are different; for
example, at 10°C 5H20 = 710 cm'1 and 5020 = 530 cm‘l.

The conclusion is that solvent isotope effects can be

calculated rather accurately by (l) determining the

 

 

85

internal vibrations of any molecules which have protons
which would exchange with water and become deuterons in
D2O solution, (2) determining the librational frequencies
of the water molecules solvating each molecule appearing
in the reaction. AS a crude generalization, it can be
expected that ion—forming reactions will have a solvation
isotope effect contribution to kHZO/kD20 > 1, corresponding
to a net increase in structure breaking (net decrease in
librational frequencies) on going from reactants to transi-
tion state, while effects <1 are expected for ion—destroy—
ing reactions, and effects close to unity are expected for
reactions which do not produce or destroy ions. Solvation
isotope effects will generally be small, probably no more
than 20—30% except for small ions which can have a con—

siderable effect on the water structure (librational fre-

quencies).

Isotgpe Effects in Excited State Proton Transfer in 7-
Azaindole Dimers - Let us consider the kinetic scheme of

double proton transfer:

kTD
.1 (2T)* k = k + k

«7 «:\ / \ .T

D + th 2T + th

 

 

86

The kinetic description of the system under conditions of

continuous illumination is as follows:

d[D*]/dT = IO - (kD + kTD)[D*] + kDTI(2T)*l
(1)
d[(2T)*]/dT = kTD[D*l - (kT + kDT)[(2T)*l

Steady state conditions give d[D*]/dT = d[(2T)*]/dT = 0.

By using the definition for quantum yields

¢FD = kFD[D*]/IO and ¢FT = kFT[(2T)*]/IO (2)

and since after correction for different instrumental

sensitivity in the two different spectral regions

F /F

¢FT/¢FD = 2 1 (3)

we obtain from Equations 1, 2 and 3

F2/Fl = (kFT[(2T)*]/(kFD[D*]) = (kFTkTD)/<kFD(kDT+kT>)
SO
(F2/F1)H/(F2/F1)D = kQD/k D (kgT+k$>/(kgT+k$) (u)

since isotopic substitution is not expected to change the

 

 

 

 

87

radiative lifetimes. Further, since kT + kDT z kT and
. _ h . d
Since kT — kFT + knr’ kT is expected to be equal to kT
except in the improbable case of a very large and specific
effect of one H/D atom on the non-radiative rates. Under

these assumptions equation 4 can be Simplified to
(F /F > /(F /F > = kh /kd (5)
2 l H 2 l D TD TD

The experimental isotope effect measured as in Equation 5,

as the ratio of F2/Fl for protonated and deuterated 7-
azaindole in 3MP matrix at 77°K is 2.9. So,

h . d _
(kTD/kTD>77OK ‘ 2'9

Using the result for k%D(8.6 x 107 s_% one can obtain

ng = 2.5 x 108 S-l.(80) Further, as we discussed earlier,
this kinetic isotope effect is connected to the zero—point

energy difference (ZPE). So,

kh /xd = exp(—ZPE/RT)

TD TD

where
h -d _
kTD/KTD — 2.9.

The zps thus obtained is around 60 cm’i. If the

 

 

88

activation energy for double proton transfer is 500 cm-1

then the activation energy for the double deuteron trans-
fer is around 560 cm_l. This is in good agreement with
the estimated activation energy of 500-600 cm-1 consider-
ing the double deuteron transfer at 77°K as a thermal
process. It appears that a mechanism that involves mo-
tion of the deuterons over the potential barrier can ex-
plain the experimental data at 77°K because of the very

small energy barrier.

d. Proton Tunnelling

 

One of the most unusual and interesting results of
quantum mechanics is the prediction of tunnelling, i.e.,
the ability of a particle to exist in, or pass through, a
region of space where its total energy is less than its
potential energy. According to classical mechanics, such
a phenomenon is impossible. As early as 1927(1UB)
Hund discussed the probability of intramolecular rearrange—

(149)

ments via tunnelling and Wigner in 1932 discussed
tunnelling with a view aimed at chemical kinetics. A
tunnelling mechanism was prOposed for the doubling of

(150) When proton

certain Spectral bands of ammonia.
transfer is very fast, such as in acid-base reactions in
aqueous solutions, there exists the possibility that the

proton in the incipient H-bonded complex A-A...B tunnels

through the potential barrier between the two minima.

 

 

 

 

89

This possibility was pointed out by Bell in 1935. Since
then, a great deal of indirect evidence has been accumu-
lated which indicates the occurrence of proton tunnelling

(151,152) (153)

in ultrafast proton transfer. Johnston

(154)

and Caldin have provided reviews of various aspects
of proton tunnelling in ordinary chemical reactions, and

the significance of tunnelling to an understanding of the
hydrogen-bond has been well described.<155_157) L6wdin<l§8>
has discussed tunnelling as a possible mechanism of causing
mutation and genetic error.

The JWKB method was used to obtain the transmission co—
efficient for an arbitrarily shaped potential. The JWKB
solution has an oscillatory behavior in the "permitted"
region and an exponential behavior in the forbidden region.
In ordertmiderive a formula for the transmission co-
efficient g (the probability of getting through the bar-
rier) in general, we will consider a single incident wave
hitting a barrier with two turning points (Figure 24). If
the barrier is parabolic at its top (Figure 24), one can
obtain for KO,

K = (1/4w2fi/ha )(2m(V —E ))1/2
0 o 2 o
The factor w/4 comes here from the shape of the barrier.

If one measures the energy from the top as a fraction

k of V0, so that E = V2 — kVO one obtains

 

 

9O

 

Figure 24. Model barrier for proton tunnelling calcula—
tion.

91

and

1/2

09
ll

exp(-kfi2/hao)(2m(V2-EO))

Experimental Evidence in 7—Azaindole Dimers — By
studying the effect of temperature on the fluorescence
Spectra of 7AI under conditions of continuous illumination
a value of 500 cm—1 was obtained<79) as the potential
energy barrier for double proton transfer. It was found,
however, that the relative integrated intensities of the

tautomer and dimer fluorescence are the same at 77°K and

4.2°K, i.e., (kTD)77°K= (kTD)u.2QKii'one makes the reason—

able assumption that the radiative and nonradiative

(k ) rates do not change from 77 to 4.2°K. If the

ID,kIT
double proton transfer is a thermal process where the

two protons move over the energy barrier then by assuming

an Arrhenius behavior with pre—exponential factors of

1012 — 10l3 sec—l, characteristic of unimolecular reac-

-l
tions we obtain energy barriers of the order of 20-30 cm

for the reaction at 4.2K. Clearly these observations

point strongly to quantum mechanical tunnelling. At

higher temperatures the proton transfer can take place

by both thermally activated and tunnelling mechanisms,

while at very low temperature tunnelling may be the only

92

reaction path. The effects of deuterium substitution of
the pyrrolic hydrogen on the magnitude of F2/Fl ratio at
77°K are consistent with current theories of proton tun-
nelling. However, recent studieél27) of the effect of
temperature on F2/Fl ratio appears to indicate that the
proton-tunnelling mechanism is not necessary to account
for the observed results. Further studies are required
to examine the role of tunnelling in excited state proton

transfer.

e. Excited State Proton Transfer in the Vapor Phase

 

Excited state proton transfer in the vapor phase was
observed for the case of methyl salicylate<l59) in 1924.
At 110°C (17 mbar) two broad emissions of comparable
intensity in the spectral range between 320 and 480 nm,
similar to those in solutions of methylsalicylate have
been observed.

The vapor—phase UV-absorption spectrum of methyl-
salicylate shows an absorption band with a maximum at 303 nm,

(81) and

probably an intramolecular charge transfer band,
a structured absorption with a maximum at 234 nm (similar
to the usual absorption of aromatic hydrocarbons).
Fluorescence spectra of MSA vapor in dry nitrogen at
50°C are shown in Figure 25 for two exciting light wave-
lengths: Methyl salicylate in the vapor phase clearly

shows two different fluorescence bands whose relative

 

93

 

 

fluorescence intensity

 

 

Figure 25.

 

0
wavelength

Vapor phase fluorescence spectra (uncorrected)
of MSA at 50°C, equilibrium pressure under N2,
excitation wavelengths indicated. The experi—
mental conditions are identical for both spectral.6O

I.
(I
l‘
«I.

O.
0080!.- ..llt. .Cto.

 

 

94

 

 

10

”333)

[(450) _ \\

 

 

 

 

 

— 82
3‘
'2’ 55 _
2 57
s
0 0.1 J 1
g 25 30 35 K”
3 — 1000/T
Q 48
8
z 1.0
30
I I I I I
300 350 400 450 500 nni
wavelength
Figure 26. Vapor phase fluorescence of MSA, under N , as

a function of temperature (indicated in C).
Insert: semilogarithmic plot of the intensity
ratio of short—to—long wavelength emission vs.
reciprocal temperature.160

 

 

 

 

 

95

intensity depends on the excitation wavelength and the
temperature. These are a blue fluorescence (emission
maximum 450 nm, excitation maximum 319 nm) and a UV fluores-
cence (emission maximum 335 nm, excitation maximum 303 nm).
The existence of distinctively different excitation max—
ima for the two emissions points to two ground state
species as the origin of the two fluorescence bands rather
than to an excited state equilibrium. The species respon-
sible for the blue fluorescence corresponds to the form
present in non-polar solvents, which is an intramolecularly
H-bonded species (closed ring structure). The rate of
excited state intramolecular proton transfer has been
measured directly<98> for N-heterocycles containing
intramolecular H—bonds, to be the order of 1010s.1 and by
a picosecond study for methyl salicylate in hydrocarbon sol—

lls_l.(96) Proton transfer

vent, to have a lower limit of 10
therefor effectively competes with other deactivation pro-
cesses. The UV fluorescent species emit at shorter wavelengths

than observed for the solvated form (355 nm), these species

9...
(:§§c> <:§()
I; :E;:__. (l)
O ’ ('3
H
IflSUKCL IISARD

 

 

96

are most probably the open structure without a hydrogen
bond. The equilibrium constant for this reaction oc-

curring in the electronic ground state is given by:

K = [MSAOJ/[MSAC]

Fluorescence obtained at different temperatures, using
290 nm excitation, are shown in Figure 26. Increasing
the temperature strongly increases the short wavelength
fluorescence intensity, whereas the blue fluorescence
turned out to be nearly temperature independent. Lowering
the temperature after heating again yields the spectrum
characteristic of the lower temperature, only after an
extended irradiation time do new peaks appear which can
probably be attributed to irreversible photochemical
reactions. The insert in Figure 26 shows that the intensity
ratio of short—to—long wavelength fluorescence decreases
exponentially with l/T. The slope of the semilogarithmic
plot in Figure 26 gives the enthalpy of the reaction (1).
This quantity is found to be 3.59 Kcalmole-l and corresponds
to the energy needed for breaking the intramolecular hy—
drogen—bond. It is substantially smaller than the value
estimated for the strength of the H—bond in MSA based on
Spectral shifts caused by C=O...H-O in stretching vibra_

-l>(161)

tion bands of the 0:0 group (6.7 Kcalmole and in

much better agreement with the results of a measurement

 

97

using ultrasonic relaxation in liquid MSA (2.5 Kcalmole— ,
K = 0.0365).<95) Since the entropy change, AS, of reac—
tion (1) is most probably very small, we can estimate

the equilibrium constant K to be 0.0025 at 25°C.

F. Excited State Proton Transfer in Acridine Derivatives —
In 1946, a reversible color and fluorescence spectral

change were detected in acridine at —l80°C. The acridine
was present as a minor constituent in a vacuum-sublimed
crystalline organic acid such as oxalic, terphthalic, or
succinic. The transition from the spectrum of the acridine
cation to the spectrum of the neutral acridine molecule
under illumination indicates clearly that a transfer of a
proton occurred from the acridine cation back to the acid
anion as a result of electronic excitation. The study<l62>
of the fluorescence and absorption spectra of acridine and
its derivatives at different pH values in aqueous media
showed that acridine has a higher affinity for the proton
in the excited state than in the ground state. These
results were in contrast to the observation<l63> that the
phototransfer of a proton occurs from an acridine cation
to an acid anion. This may be attributed to the fact

that the conditions were different; the acridine was con—
tained in a crystalline acid at low temperature, and not
in aqueous solution.

Further investigation on other compounds and media,

 

 

98

condensed together at low temperatures in a solid matrix
gave similar results<l6u) (a film of acridine and an
organic acid). Another study used "amphotoric" derivatives
of acridine, which have acidic and basic groups in the same
molecules,(l6u) one of these compounds is 9—(p-hydroxy)—
phenylacridine, which forms chains by means of strong inter—
molecular H-bonds.<l65) In aqueous media at low pH values
the absorption maximum of the long wavelength absorption
band undergoes a red Shift into the visible region as in
the case of acridine.(l65) A similar effect was observed
in the fluorescence spectrum of an ethanol solution of 9-

(p-hydroxy)phenylacridine (Figure 28).

 

Xm=450nm xm .530...“

A film of 9—(p—hydroxy)phenylacridine, immediately
after sublimation under vacuum at ~180°C has a green
luminescence with a maximum near 530 nm (Figure 28a),
coinciding with the maximum observed for this compound

in acid solution in Figure 27. This undoubtedly is due

 

 

Figure 27.

 

 

99

 
 

 

 

:00 as} 550 530
A071“)

Fluorescence spectra of a solution of 9-(pr
hydroxy) phenylaridine in ethanol at differegt
pH values. Curve 1 — pH 8; curve 2 - pH 2.1 4

 

041 N 0’ INF

 

L._

 

 

 

NM

Figure 28

550 L 53:7 W35. )3:

A (mu)

dfve 101N701

(a) Fluorescence spectra of a film of 9-(p-
hydroxy) phenylacridine in vacuum at —l80°C.
Curve 1 - before exposure; curve 2 - after ex—
posure for 10 min to the total radiation of an
SVDSH-250 lamp; curve 3 - after regeneration
at room temperature; curve 4 — acetic acid
vapor; curve 5 - in pyradine vapor.

(b) Potential energy curve scheme for proton
phototransfer in films at —l80°C.l

 

 

100

to the combination of the proton of the phenol group with

the nitrogen of the acridine ring (Structure 1):

-0...,:NI_\ O
I

Upon irradiation at —180°C, the green luminescence is re-

+z’
0"“HN \

 

placed by a blue luminescence in about 10 min, with a
maximum near 450 nm. The latter luminescence is observed
in 9-(p—hydroxy)phenylacridine in alcoholic solution, ttgt,
under conditions in which the phenol group is unionized
(Figure 28). If the film is kept at room temperature for

2 to 5 hours, the green luminescence is restored, and

this experiment can be repeated reversibly several times.
Thus it appears that upon irradiation there is photo-
transfer of a proton from the nitrogen atom (Structure II)

to the negatively charged phenolic group.

 

Therefore, it follows that in condense films, 9—
(p—hydroxy)phenylacridine exists in the O— (I) and not in

the hydroxy form (II), as it does in neutral solvents.(l65)

 

 

 

101

Here the proton transfer process uses the energy absorbed
by the O_ form of the molecule of 9—(p—hydroxy)phenyl—
acridine in the short wavelength absorption band. The
return of the proton from the phenolic hydroxy group to
the heterocyclic nitrogen is a thermal process. A sim—
ilar phenomenon occurred in the two-component film (acri-
dine + acid), in which the phototransfer of proton pro—
ceeded from an acridine cation to an acid anion.(l63)
Structure II, which has the higher energy, can be suc-
cessfully stabilized at low temperatures.

The process described is apparently possible only in
systems in which there is intimate contact between mole—
cules and little probability of dissipation of the excita—
tion energy into the thermal energy of the solvent. It
was assumed that the transfer of a proton to the acid
anion does not occur during the lifetime of the excited
state of the fluorescing molecule, but after its transition
to the ground electronic state at the expense of its high
vibrational energy, which is acquired immediately after
the emission of a light quantum. In fact, from the magni—
tude of the Stokes shift of the luminescence spectra of
acridine, it follows that the emitting molecule in its
ground state has initially a supply of vibrational energy
approaching 11—12 KcalunoleJl The two minima in the po_
tential energy of the proton near the ionized oxygen 0-

of the phenol group on the one hand, and the nitrogen

 

102

atom of the acridine ring on the other (Figure are
separated by an energy barrier of about 5 Kcal, as is
characteristic of a strong hydrogen bond. In such a case,
the vibrational energy liberated in the process of emitting
a quantum of fluorescence should obviously be sufficient
to overcome the indicated potential barrier. For the
reverse process the potential barrier is lower, as follows
from the rapid transfer of the proton to the nitrogen,
even at room temperature. The same things occur in a 9-
(p—hydroxy)phenylacridine, the ionized phenolic group of
the neighboring molecule playing the role of the acid

anion.

 

CHAPTER II

EXCITED STATE PROTON TRANSFER IN P-AMINO

SALICYLIC ACID DERIVATIVES

A. Introduction

 

Excited state proton transfer reactions are among the
most common adiabatic photoreactions; that is reactions
in which the products are generated in an excited state.
This fact allows the study of these phenomena by fluores-
cence spectroscopy.

This chapter deals with excited state proton transfer
in hydrogen bonded systems both in solution and in the
solid phase. The hydrogen bonded systems we have studied
are p-amino salicylic acid (PAS) (I) and its N,N dimethyl

derivative (DPAS) (II).

HO\C,’O HO\C¢O

OH OH

103

104

Absorption and luminescence properties of p-amino salicylic '
acid (PAS) derivatives were investigated under various
conditions. The effect of medium, pH and substituents on
the absorption spectra helped in the interpretation of the
absorption bands which are dominated by "charge—transfer"
transitions. In the "charge—transfer" states of these
molecules the hydroxyl and amino groups may act as an
electron-donors whfljathe carboxyl group acts as an elec—
tron acceptor. The various species responsible for the
absorption and emission at different pH values and the
emission spectra in different solvents clearly reveal the
occurrence of proton transfer in the excited state. Ex-
cited state proton transfer occurs either intramolecularly
or intermolecularly depending on the medium. In order

to get more information about proton transfer and to
confirm the occurrence of proton transfer in PAS and DPAS
we also synthesized the methyl ester of these compounds
III, IV and we have studied their absorption and emis-

sion spectra under the same conditions.

CH30\ [/0 CH30\ //O
C C
CH
CH3 CH3

 

 

105

The crystal structure of PAS<166> intra and inter—
molecular hydrogen bridges. We measured the fluores-
cence spectrum of PAS in the solid phase. Information
from the solution spectra of PAS helped us to interpret
the emission band in the solid phase of PAS in terms of
an excited state intramolecular proton transfer. At
the beginning of the chapter the spectrum of PAS is related
to parent and related molecules, ttgt, benzene, phenol and

salicylic acid.

B. Excited State Proton Transfer in P—Amino Salicylic

 

Acid (PAS) and its Methyl Ester (PASE)

 

Electronic absorption spectra of substituted benzenes
are usually compared with the spectrum of benzene itself.
This suggests that the differences in Spectra of substi—
tuted benzenes reflect the perturbation effects of the
various substituents. Such a picture is valid if or-
bitals are localized either on benzene or on the sub-
stituent and if no charge migration occurs between the
substituent and benzene. However in order to interpret
certain intense transitions which occur in the spectra of
substituted benzenes but do not correspond to any benzene
transition (B2u (lLb), Blu (lLa), E2u (1B), charge—trans—
fer between the substituent and benzene orbitals must be

considered. One may classify an excited electronic state

 

 

106

of a substituted molecule as either locally excited (LE)
state or a charge—transfer (CT) state only as a zeroth—
order description of the electronic state, since mixing
between LE and CT states occur.

A locally excited state involves an excitation between

 

different orbitals of the hydrocarbon, or between dif-
ferent orbitals of the substituent. Transitions involving
LE states will appear in the substituted molecule at dif—
ferent energies than the corresponding transition of the
unsubstituted hydrocarbon. Their intensities are also
different due to the influence of the perturbing potential
of the substituent.

A charge—transfer state involves the transfer of an

 

electron between the hydrocarbon and the substituent. A
CT transition gives rise to a CT band which is inherent
to neither the hydrocarbon nor the substituent. Intense
absorption bands have been observed frequently in com—
plexes involving an electron-donating molecule and an

electron—accepting molecule. A quantum—mechanical theory

for these complexes has been developed by Mulliken<167).

(Excellent reviews of this subject have been given by

(168) (169) (170)

McGlynn, Briegleb, Murrel and Mulliken

(171)), which explains the observed band as an

and Pearson
intermolecular CT transition. One would expect similar
bands to be observed in the Spectra of poly substituted

benzene molecules, in which an electron—donor group and an

 

 

107

electron-acceptor group are parts of the same molecule.
Transitions which result from CT interaction between these
groups consequently are called intramolecular CT transi—
tions.

Many examples of intramolecular CT transitions are
known. For example, the long wavelength absorption band
of p-nitro—aniline which is solvent sensitive, is charac-
terized as an intramolecular CT transition involving charge
migration from the amino- to the nitro-group (Figure 29).
Arylboron compounds provide an interesting converse ex—
ample where intramolecular CT transitions are expected
to occur, thus the spectrum of tri-l-naphthylboron shows
an intense absorption band at 3525 A which was inter—

(172)

preted as an intramolecular CT band in which the
naphthyl group acts as the electron—donor and the boron
atom as the electron—acceptor. Comparison of a given
arylboron compound in methylcyclohexane as solvent before
and after bubbling dry NH3 gas through the solution

shows that the band assigned as an intramolecular CT band
disappears in the presence of ammonia. This was explained
in terms of the formation of an additive compound with

NH3 in which the lowest vacant orbital of boron,involved
in the intramolecular CT transition,is now filled with

the nitrOgen lone pair electrons (Figure 30).

The localized—orbital model is the most preferred model

for the treatment of substituted benzene. One of the

 

 

108

 

6 no"

 

\J/R\I ,1 ‘ . . .
5 \\ // / \ \\ _.

. —— —enm ALCOIOL
,o’ '- A- --------- METHYL cmorcx»:

I/ \ -* ,__. _ .—. amen
/'}:"- \- I
/ I": /)\~

. ' / '

\\
2/ \ \

 

/"-., '\\

2 \ss' ,./// i I." - \ :
I I“ I I - I I .I"MI_A

 

 

 

2000 2400 2000 3200 3600 4000 4400 4000

 

WAVE LENGTH IN ANGSTROMS

 

Figure 29.

Room temperature absorption spectra of p—
nitroaniline in different solvents: (...) methyl-
cyclohexane, ( ----- ) ether, and (————) ethanol
(concentration in alcohol = 6.52 x 10- gm

moles lit-1.).

 

109

Wovemmbers x I0"3
315 40 4s 50

30
L

 

 

‘\~

    

 

 

I
I
/
o 4/ 1 T ‘
‘000 3500 3000
A, Angstroms
Figure 30. Room temperature absorption spectrum of tri-l-
naphthylboron in methylcyclohexane (dashed
line); and same after bubbling in dry ammonia

gas (solid line).

 

110

 

reason is that it allows a correlation between the ab-
sorption bands of various substituted benzenes. The CT

state energies can be calculated by using the equation:

 

where ID is the ionization potential of the donor, EA is
the electron affinity of the acceptor and C is the Coulombic
attraction energy between the electron and the hole.

In monosubstituted benzenes symmetric and antisym-
metric CT states result, due to the degeneracy of the
highest filled and lowest vacant benzene orbitals, their

energies are different due to the difference in the electro—

 

static term. In disubstituted benzenes of the type
ID-<:)-A, one can use two modifications.(l73) The first

is the use of the ionization potential of D—O rather
than that of benzene. The second arises from the realiza—
tion that the degeneracy of the highest filled benzene
orbitals is removed as a result of substitution. The
highest filled MO of benzene is doubly degenerate, but

this is not the case in phenol or aniline (Figure 31).

In para disubstituted benzene the transition to the

 

symmetric state is allowed but it is forbidden for the
antisymmetric state, due to the zero coefficient of the
wave function on the carbon atom which is linked to the

substituent. The CT band observed for the para isomer

 

 

111

 

 

 

.__JLI__.

I
,0 I
‘Q’

Figure 31. Qualitatively molecular orbital energy level

for benzene, phenol and lowest vacant orbitals
of substituent.

 

 

112

corresponds to the symmetric CT state and it borrows most
of its intensity from lBlu band. For meta and ortho di—
substituted benzene both transitions are symmetry allowed.
Let us consider the absorption spectra of salicylic
acid (SA) in different media (Figure 32). For SA in its
neutral form there exist two CT bands, one at 302 nm and
the other at 237 nm. These bands are due to a transition
to a CT state in which the hydroxyl group acts as an
electron donor and the carboxyl group acts as an electron
acceptor. The n—n and B2u bands of benzene are hidden
under the intense CT band at 302 nm, but the higher energy
CT band at 237 nm of the ortho and meta isomers borrows
most of its intensity from the 1B LE band of benzene.

lu

The intensity of the lBlu band is approximately 10 fold

greater than that of the 1B211 band and therefore the band

at 302 nm has greater CT character. When we remove a
proton from the carboxyl group, the two CT bands are shifted
to the blue because of the lower electron affinity of the
carboxylate anion compared to that of the carboxyl group.
The removal of the phenolic proton causes a red shift for
both CT bands, because of the lower ionization potential

of the phenolate anion with respect to that of the phenol
group. The red shifts for the doubly—Charged anion for

the two CT bands are smaller than that of methylated
salicylic acid. This reflects the electron repulsion of

these two negative charges which forces the carboxyl group

 

 

 

 

 

4v.» 3.38:: 088059....“ \I // \x. 2

 

113

 

 

 

 

 

3
xIO
xI03 I5
>\
E 5'
:5... 4 n " IO
3-
8 2 ' \‘—' '5
‘8 I \
ES J -1 I \\
”L 250 300 350 400
Wavelength (nm)
-0 -0 H03 \ ceo
296 302 306 358

Figure 32. Room temperature absorption spectra of salicylic

acid (SA) in different media: N/lO—
HCl; ( ..... ) N/20-NaOH; (-..— .—) 5—n NaOH; and
methyl salicylate in (-—-—) N/20-NaOH.

 

114

to be out of plane. The spectrum in this case is very

similar to the spectrum of the phenolate anion.

l. Solvent Effect on the Absorption Spectra of PAS

 

and PASE

 

Environmental factors affecting the absorption spectra
can be conveniently divided into three types: intermolec-
ular hydrogen-bonding between solute molecules only; inter-
molecular hydrogen-bonding between solute and solvent mole-
cules; and environmental effects not involving the forma-
tion of hydrogen bonds such as dipole—dipole effects. All
of our experiments were performed for very low concentra-
tions of the compound in order to eliminate the possibility
of intermolecular hydrogen bonding between solute mole-
cules.

Absorption spectra of dilute solutions (5 x 10-5 M)
of PAS has been taken in different media and are shown
in Figure 33. The energies of the two near UV absorption
bands in different media and their spectral shifts with
respect to 3-methylpentane (3MP) are listed in Table 3.
Depending on the solvent, an intramolecular hydrogen bond
or intermolecular hydrogen bond can be formed. In 3MP,
one expects that the predominant species will have intra-
molecular hydrogen bonds involving the hydroxyl and car-

boxyl groups. We have assigned the 275.5 nm absorption

200

115

 

 

 

 

 

 

0.8 ‘ ' -
0.7 - .
0.6 - -
0.5 ' "
§04- -
B
.203- q
<[
0.2 - -
0.l - -
l l
200 250 300 350
Wavelength (nm)
Figure 33. Room temperature absorption spectra of PAS
in different solvents: ( ) 3MP, (----)
ether, ( ----- ) ethanol, and (....) H2O (pH

— 2).

 

116

Table 3. Room Temperature Absorption Maxima (nm) and
Spectral Shifts (om-l) Relative to 3MP for PAS

in Different Solvents.

 

 

 

 

 

Solvent 1(nm) A§(cm-l) 1(nm) A5(cm_l)
(N)CT (O)CT
3MP 275.5 0 302 0
Ether 278 —327 302 0
EtOH 280 -584 303.5 -164
H20(PH=2) 279 —456 299 +332
— red shift + blue shift

band to a transition to a charge transfer state in which
the amino group acts as an electron donor and the aromatic
ring as an electron acceptor, (N)CT band. The 302 nm
band is assigned as a charge transfer band involving the
hydroxyl as an electron donor and the carboxyl group as
an electron acceptor, (O)CT band. Of course, this is a
zeroth—order description of the electronic states Since
mixing between locally—excited and charge transfer states
occurs. If the transition involves mainly an electron
transfer from the highest filled orbital of the donor
group to the lowest vacant orbital of the acceptor, then
one would expect that the band would disappear if the
energy of the highest filled orbital of the donor group

becomes much deeper in energy as a result of protonations.

 

 

117

Charge transfer transitions involving promotion of an
electron from a w—orbital "lone pair"* which has n-sym—
metry and can therefore conjugate with the benzene molecule,
virtually disappear as a result of protonation. The ab—
sorption spectrum of PAS at a low pH value 0.85) is con—
sistent with our assignments since the 275.5 nm band at—
tributed to a charge-transfer band involving the amino
group, disappears when the amino group is protonated and
the spectrum becomes nearly the same as that of salicylic
acid (Figure 35). The fact that neutral salicylic acid
exhibits an absorption band at m302 nm and that p—amino
benzoic acid exhibits an absorption band at m270 nm sup—
ports our assignment.

In ether, where intermolecular hydrogen bonds in?
volving the hydroxyl and amino groups are formed, the
spectrum is very similar to that in 3MP where an intra—
molecular hydrogen band is formed. Dipole—dipole inter~
actions contribute to the observed transition energies
in ether.

In ethanol, in addition to the previously mentioned

interactions, there are intermolecular hydrogen bonds

 

* .
The 2p7rorbitals on the nitrogen atom of aniline and pyrrole
are w-orbitals while the 2p orbitals on the aza-nitrogen
of pyridine and on the oxygen atom of the carbonyl group
are n—orbitals.

 

 

 

 

118

involving the proton of ethanol. The net result is that
both absorption bands are shifted slightly to the red
compared with those in ether.

In water, at pH 2, the Spectrum shown in Figure 33
corresponds to the neutral form and is Similar to that
in alcohol. However, the relative intensities of the
two absorption bands are reversed and the bands occur at
slightly higher energies. Hydrogen bonding involving
the protons of water with the amino group lone-pair is
probably responsible for these effects. In water the
intensity of the (N)CT band decreases due to H—bonding with
the amino nitrogen lone—pair which if carried to the ex—
treme situation, would be the addition of H+ which would
eliminate the amino CT band. In water we expect a larger
red Shift, but the hydrogen bond between the amino group
lone—pair and the alcohol proton is weakened as a result
of CT excitation and will cause a blue shift which is
superimposed on the red shift due to dipole-dipole inter-
actions. This effect is more pronounced in water than in
alcohol because of the higher acidity of water. Hydrogen
bonds formed between the amino group hydrogen atoms and
the solvent molecules become more stable in the CT state
and lead to a red shift.

The absorption spectra of PASE (5 x 10‘5 M) have been
measured in different media and the results are shown in

Figure 34. Absorption maxima and spectral Shifts are

 

119

 

C)
e 8

Absorbmce

_o
o:

(D
A)

DJ

 

 

 

 

200 250 300 W 350
Wavelength (nm)

Figure 34. Room temperature absorption spectra of PASE
in different solvents: ) 3MP, (-——-)
ether; ( ----- ) ethanol; and (—--—--) H20.

 

 

120

summarized in Table 4. The results are very similar to

that of PAS.

Table 4. Room Temperature Abggrption Maxima (nm) and
Spectral Shifts (cm ) Relative to 3MP for PASE
in Different Solvents.

 

 

 

1(nm) _ —1. 1(nm) __ —l
Solvent (N)CT Av(cm ) (O)CT Av(cm )
3MP 273 0 302 0
Ether 281 —1043 303 -110
EtOH 285 -1542 306 —433
H20 276 —398 301 +110

 

 

2. pH Effect on the Absorption Spectra of PAS and

PASE

_..._——

We have studied the effect of pH on the absorption
spectra of PAS and its ester (PASE) in order to determine
the absorption characteristics of different species which
may exist in the ground state. The results are summarized
in Tables 5 and 6 and in Figures 35 and 36. For PAS
at pH 0.85, the cation exhibits only one charge transfer
band involving the hydroxyl group as an electron donor

since the amino group is protonated at this pH value.

 

 

 

 

121

 

mmz" mmzn
co: oom mom oa.mH
Io Io
owe/o- owe/o-
mmz" mm
00:2 mam 3mm om.m
0 mo
O“O/OI O“O/OI
mmz“ mmz”
3H: mam msm oo.m
Io mo
owe/om owe/om
mmz" mmz+
02:2 mam III mw.o
Io mo
+mosoxom ouarom
AEQV mofiooam opmpm popfloxm mofloodm mumpm pcsopw AECV ma
meflxmz wefixez coflpdpomn<
coemmfiem

 

 

.mofloodm mopmpm woufloxm one pcsopw smosfim>
ma pcosoeefim pm m<m soe meflxmz coemmflem one :ofipdeomp< obsessoQEoB Eoom .m mange

 

, _fi_ _w.....-a

122

 

 

(\l
("\J

 

m2" mz”
mo: 0 0 0mm SN 0.: _
Io Io
owe/Ommo omozommo
mmz. mmz” mmz
om: mm 0 Q Q Sm Em m.m
Io mo mo
Homo/ammo eve/ammo eve/0mg
mmz mmz mmz+
Hm: mm mm III m;
o no mo w
+mo“o/ommo eve/cmmo owe/cmmo
AEQV mmflommm mumpm cmpfloxm mmwommm mumpm UCSOLU AEQV ma
mEflxmz wEwaz QOflthomn<
COflmeEm

 

 

.mmflomam mmpmpm umpfloxm wcm 625096 am®5Hw>

mm pcmgm%%flm pm mm<m L0% mEwaz COHmmHEm ccm coapgpomn< mQSpmmeEmB Eoom .w mHQmB

 

 

 

08"—
1

07"

0.6 "

05F

»
4
0

hi! P
3 2
0 O
09.50550me

b
0

Lm

123

 

0.7 "

0.6 '

Absorbance
o
o:

 

 

 

 

 

0.2 _
O.I '
l 1 . '-
200 250 300 350
Wavelength (nm)
Figure 35. Room temperature absorption spectra of PAS
at different pH values: (—--—) 0.85, (————)

2.0, ( ----- ) H20 (2.9-12.1); and (- .-- -)
13.1.

 

 

 

 

 

 

|ll|
I I ‘ _

04"

”080.50% 94‘

 

124

 

0.7 . I , 1
0.6-

0.5 "

 

0.4 *'

 

 

O. 3 P \l 1

Absorbonce

|

‘ \
<)£2" ‘ /. \

I

()J" \ //

 

 

 

l l
0.0200 250 300
Wavelength (nm)

 

Figure 36. Room temperature absorption spectra of PASE
at different pH values: (—-——) 1.3; (————)
H20 (2.3-10.3); and (....) 11.0.

 

125

At pH 2, the PAS exists in the neutral form, and two ab—
sorption bands comparable in intensity are observed. At
pH range 2.9—12.1, PAS exists as an anion, the (N)CT ab-
sorption band occurs at higher energy (264 nm) compared
with its value for the neutral form. This is due to the
lower electron affinity of the carboxylate anion compared
with the carboxyl group. The intensity of the band is
increased relative to that of neutral PAS, probably due
to more charge migration from the amino group to the
carboxyl group and to the planar structure. At pH 13.1,
the doubly-charged anion spectrum shows a decreased in—
tensity of both bands, probably reflecting a partially
twisted carboxylate group due to electron repulsion in—
volving the 0‘ group.

For PASE, the results are very similar to those for
PAS. An exception is that in the case of the PASE species,
no carboxylate anions or doubly-charged anions are formed,
instead the phenolate anion is formed at higher pH values.
The absorption spectrum of the phenolate anion lies at
longer wavelengths than that of the doubly-charged anion

of PAS, supporting the notion that repulsion in the doubly—

charged anion of PAS forces the carboxyl group out of plane.

A

 

126

3. Emission Spectra of PAS and PASE as a Function of

pH at Room Temperature

 

The room temperature emission spectra of PAS and its
ester (PASE) at selected pH values are shown in Figures
37 and 38. The excitation wavelengths are 286 and 280 nm
for PAS and PASE, respectively, which correspond to the
first absorption bands. The ground and excited state
species at a given pH value are shown in Tables 5 and 6.

In water at pH 2, where PAS exists predominantly
in its neutral form in the ground state, the emission band
occurs at #1“ nm. This band is due to intermolecular
proton transfer from the hydroxyl group to the solvent
molecules, producing the phenolate anion as shown in
Table 5. The emission of PASE in ethanol in the presence
of sodium hydroxide shows a band at mulfl nm which provides
further support to this conclusion. At a pH value of 5.6
where the dominant species in the ground state is the
carboxylate anion, the emission maximum occurs at “00 nm.
Under such conditions the spectrum is identical to the
emission observed at pH 13.1 where the dominant species
in the ground state is the doubly—charged anion. This
indicates that the species produced in the excited state
under these conditions is the doubly—charged anion.
Further confirmation of this conclusion comes from the
observation that the emission spectrum Of PAS in ethanol

in the presence of sodium hydroxide and that of PAS in

 

“a
1

)u’rturs~ .5 1:9. y'nvill ‘

35C

3’)?

127

 

Fluorescence Intensity

 

 

 

 

I l ' 1
I ’g/ l 1 $T=wn
300 350 400 450 500 550
Wavelength (nm)

Figure 37. Room temperature fluorescence spectra of PAS
at different pH values: (-—--) 0.85, (———-)
2.0, ( ----- ) H20, and (---—--—) 13.1.

 

/

1!. z: .
sh a. v9! v"sa

P

is e

350

 

128

 

Fluorescence Intensify

 

 

J

 

 

'ul 1 l
350 400 450
Wavelength ( nm)

500 550

l
600

Figure 38. Room temperature fluorescence spectra of PASE

at different pH values:
H20, and (....) 11.0.

 

(~——-) 1.3, (

 

)

 

 

129

3MP in the presence of pipyridine are both nearly the
same as that observed in water at pH 5.6.

The emission of the phenolate anion occurs at longer
wavelengths than that of the doub1y~charged anion, probably
due to a smaller electron affinity of the carboxylate anion
and the non-planar configuration of the doubly—charged
Species. At pH 0.85, where the amino group is protonated
in the ground state, emission occurs at MAO nm which is
due to the zwitterion formed upon the excitation of PAS
at this pH value. This emission occurs at longer wave—
length compared to the singly—charged anion emission.

The dissociation of the hydroxyl group and the protonation

 

of the carboxyl group both stabilize the excited state.
PAS in 3MP where an intramolecular hydrogen bond is formed
in the ground state gives a similar emission band at 450 nm.
Here an intermolecular double proton transfer involving
the solvent molecules occurs in the excited state. At
this pH value the amino group loses the proton in the
excited CT state because of more charge migration to the
carboxyl group.

For PASE at pH 1.3 two emission bands are observed,
one at W365 nm which appears as a shoulder and one at
mHBO nm. The 365 nm band emission is assigned to the
neutral form of PASE, but the long wavelength emission
band (W450 nm) is assigned to the zwitterion formed during

the lifetime of the PASE excited state. The emission of

 

130

the zwitterion of PASE is red shifted compared to that of
PAS, the same is true for methyl salicylate (MSA) and
salicylic acid (SA). In the case of PAS at low pH values
we have observed only one emission band at NUAO nm. The
results for PAS and PASE at low pH are very similar to the
results of SA and its ester (MSA) atlow pH values.

The phototautomerization of methyl salicylate in

(85)

aqueous solutions shows an acidity dependence which
indicates that phototautomerism in aqueous solutions is
biprotonic involving diffusion-limited protonation of the
carboxyl group in the excited state, followed by dissocia-
tion of the hydroxyl group by the solvent, the former step
is slow and incomplete in dilute acid resulting in some of
the excited MSA fluorescing from the neutral molecular
form. The biprotonic process in methyl salicylate is
complete during the lifetime of the excited state in
moderately concentrated acid solutions.(85)

In the case of PASE, the observation of short wave-
length emission could be interpreted as in MSA.

Excitation of PASE in 3MP is very similar to that of
PAS giving rise to a A50 nm emission band of the zwitterion
which results from an intramolecular proton transfer
during the lifetime of its excited state.

When the neutral form of PASE is excited, two emission

bands are observed. One corresponds to the excited neutral

form (N365 nm) and the other corresponds to the zwitterion

nu

. .

 

 

 

131

(m450 nm). In the case of PAS, excitation of the neutral
form gives rise to one band at A14 nm corresponding to emis-
sion from the phenolate anion. In the PASE case, the
carboxyl oxygen is more basic and tends to be protonated
in the excited state. Therefore intermolecular proton
transfer occurs in the excited state giving rise to a
zwitterion emission band. At high pH values PASE exists
mainly in the phenolate anion form in the ground state and

exhibits an emission maximum at mAlO nm.

4. Fluorescence Spectra of PAS and PASE in Different

Solvents at Room Temperature

The room temperature fluorescence spectra of PAS
and PASE in different solvents are shown in Figures 39
and A0 and the emission band maxima in different solvents
are summarized in Tables 7 and 8.

In 3MP where there is an intramolecular hydrogen bond
involving the hydroxyl and the carboxyl groups of PAS,
we have observed two emission bands, one at m3MO nm and
the other with a large Stoke shift (10890 cm_l) at m450 nm.
The short wavelength band at m3AO nm corresponds to the
normal fluorescence of the molecule from its lowest singlet
excited state. But the long wavelength emission band at
1450 nm is due to intramolecular proton transfer from the
hydroxyl group to the carboxyl group WhiCh occurs during

the lifetime of the excited state giving rise to the

 

 

 

gum

 

 

 

)uagi-Ch UUC“U:§:.R

 

132

 

 

Fluorescence Intensity

 

 

 

 

n} '

. .' I l l h ‘ u- o.
250 300 50 400 450 500 550
Wavelength (nm)

 

Room temperature fluorescence spectra of PAS
in different media: ( ) 3MP; (—--—) ether;
(--—-—) ethanol; and (....) H2O (pH = 2).

Figure 39.

 

 

 

 

 

x..mC-I:C~ flutturtsCD. a

 

 

 

Fluorescence Intensity

 

 

 

 

 

 

”/ _/. J 1 J l 1
250 300 350 400 450 500 550
Wavelength (nm)
Figure MO. Room temperature fluorescence spectra of PASE

 

in different media: ( P; (_-__) ether;
( ----- ) ethanol; and (....) H2O.

 

 

13u

Table 7. Emission Maxima in nm for PAS in Different
Solvents.

 

 

 

Solvent Emissions Maxima in nm
3MP 3A0 A50
Ether 338 nus
EtOH 3UO A16 -—-
H20 (pH=2) --- ulu ---

 

 

Table 8. Emission Maxima in nm for PASE in Different

 

 

 

Solvents.

Solvent Emissions Maxima in nm
3MP 33A “50
Ether 3&2 U52
EtOH 350 “58
H O 365 “50

2

 

 

 

 

HO,

 

 

 

135

formation of a zwitterion.
The observation of these two bands is probably due
to an excited state equilibrium between the neutral form

and zwitterion according to:

+
HO\C,/O... r H O\C,/OH
O-
.J!Z_i. (I)
V‘—

Alternatively, they may have originated from a ground
state equilibrium between different conformers followed
by a rapid intramolecular proton transfer in the excited
state. It seems that the latter process is more probable
than the first. This is because picosecond time resolved

spectroscopy for MSA showed<96> that the rate of intra—

molecular proton transfer is greater than 1011 s"1 even

at A°K, whereas the interconversion between cis and trans

d(94>

conformers is slow. The rate constants foun in

neat MSA between cis and trans forms, at 25°C were k =

f
—l

9.5 x 105 and k = 2.6 x 107 s If excited MSA in

b
cyclohexane solution has a similar value of kb, the forma_
tion of the cis form, a prerequisite for the intramolecular

proton transfer, cannot efficiently take place during the

 

 

 

 

 

 

 

136

OR 0

' I:
: = (II)

b l

H

lifetime of the excited state trans form. In air saturated
solutions this lifetime is probably less than 10-8 s.

It might well be that the proton transfer in excited

cis form molecules is very efficient, i;g;, that equilib—
rium (I) (which is similar to that of MSA) is almost tot—
ally displaced to the right. The temperature effect on
equilibrium (II) should then be the main cause of the
fluorescence temperature dependence in methylcyclohexane
solution found by Weller. The lifetime measurement of
MSA showed<96> that the neutral excited species has a
fluorescence lifetime three times larger than the excited
zwitterion at room temperature. Rapid formation of 450

nm emission and the very different lifetime of the fluores—
cence components would also indicate that the two species
are formed immediately after absorption of a photon. Per—
haps they originate from different ground state molecules.
In the case of MSA this is suggested by the observation

of differences in the excitation spectrum of the two emis—
sions in methylcyclohexane. One would expect in hydro—

carbon solvents where there is an intramolecular hydrogen

 

 

 

 

 

137

bond (cis conformer) and not an intermolecular hydrogen
bond between solute and solvent molecules, to get emis—
sion at longer wavelengths due to the zwitterion, whereas
for both MSA and PAS we see the short wavelength emission
also. The interpretation is as follows: Fluores—

(90)

cence quenching measurements on the short wavelength
emission band of MSA indicated that the short wavelength
emission may contain contributions from two "slowly”
interconverting ground state conformers that are not
equally quenched and which upon excitation, give rise to
an emission in the short wavelength band. The suggested
conformers were ones in which the phenolic proton is H-
bonded to the "ether" oxygen atom of the ester group

(Ala), and a nonintramolecularly H—bonded, or "open—ring",
trans conformer (Alb). The latter is stabilized in hydroxyl—
ic solvents by solute—solvent H—bonding. Excited—state
prototrOpism depends in part on an increase in the basicity
upon excitation, of the carbonyl oxygen, no such change

in basicity is expected for the ether oxygen. Prelimin—
ary fluorescence decay measurements using single-photon
counting appears to verify the results of the quenching

(108)

experiments in MSA. In a third conformer (Alc),
cis form, in which the phenolic proton is H-bonded to the
carbonyl oxygen, intramolecular proton transfer occurs

upon excitation, giving rise to a zwitterion which emits

at A50 nm.

 

 

 

 

 

 

138

l‘
Cx\ti' (D‘s CL\+'
,O- ‘ ,o
5 ‘CH3 El CH3 ('2’
° 0 0.
CH3
0 b C
Figure U1. Ground—state conformers of methyl salicylate:

(b) ”open—ring" (trans),
"closed ring" (cis).

(a) "ether bonded",
and (c)

The observation of a nonexponential fluorescence decay and
two fluorescence lifetimes (T1 = 23.3:O.2 ns, and T2 =

5.1:O.9 ns) for the normally Stoke shifted emission of

(12)

methyl 3-hydroxy-2—naphthoate in cyclohexane also was

interpreted in terms of more
(originating from excitation
ers) emitting within the two

the observation of the short

than one excited state
of distinct ground state conform—
fluorescence band. Therefore,

wavelength emission in MSA

and PAS could be mainly due to the existence of the
ether bonded conformer in the ground state.
In ether the results are very similar to those

3MP. There are two emission bands, one at m338 nm

responding to the neutral form of PAS and the other

stable

in

COP-

at

 

muu6 nm, corresponding to the zwitterion. In ether, intra—

molecular hydrogen bonds are formed in PAS.

In ethanol there are mainly intermolecular hydrogen

 

 

 

 

 

 

 

 

139

bonds between the hydroxyl group and solvent molecules in
the ground state. Here we observed two emission bands,
one at m416 nm and the other at m3AM nm. The short wave-
length emission band at t3uu nm arises from the neutral
form of PAS after excitation. The long wavelength emis-
sion band can be attributed to a species produced by an
intermolecular proton transfer from the hydroxyl group to
the solvent ethanol molecule producing a singly-charged
anion (phenolate anion). In order to prove this assign-
ment we produced this species in the ground state by
using sodium hydroxide and found that the emission occurs
at the same wavelength. Therefore, in ethanol, the phenolate
anion is created after excitation of a neutral molecule
ground state rather than by direct excitation of a ground
state phenolate anion.

In water at pH 2, where PAS exists in its neutral form
in the ground state, there is only one emission band at
mulu nm, which again corresponds to a species produced via
an intermolecular proton transfer between the hydroxyl
group and solvent molecules. The shift in emission is
due to a solvent effect. In ethanol and in water the
ground state equilibria between different conformers play
an important role in determining the emission properties.

In the case of PASEZin all four solvents two emissions
are observed. One occurred at shorter wavelengths and cor—
responds to emission from the excited neutral form. The

wavelength of this emission is solvent dependent as with

 

 

 

 

 

an...

 

140

 

MSA.(94) The long wavelength emission occurs at nearly
the same wavelength (m452 nm) and corresponds to a zwit—
terion emission. The zwitterion results from intra-
molecular proton transfer in the case of 3MP and ether
and from intermolecular proton transfer in the case of
ethanol and water. Again, one could consider the ground
state equilibrium between conformers to be the origin of

the two emission bands in different solvents. Depending

 

on the solvent and which conformer is more stable in the

solvent we have inter— or intramolecular proton transfer.

C. Excited State Proton Transfer in the Solid Phase of
P—Amino Salicylic Acid (PAS)

The crystal structure of PAS has been investigated<lb6>

by x-ray diffraction. The packing of the molecules in
the unit cell, the intermolecular distances, the bond
distances and the positions of the hydrogen atoms were

all consistent with the following scheme for the hydrogen

bond system in PAS.

11: 2705
O-l-l r.“ fin

 

ou-H—o

 

 

 

 

 

1H1

H
The different band distances found for C-O and C-O”

1 II
and the positions of the hydrogens show clearly that these
groups are not equivalent. Ol forms two hydrogen bonds,
an intramolecular one with the phenolic OH and an inter—
molecular one with the carboxylic OH of a second molecule
related by a center of symmetry. The intramolecular H—
bonds are stronger than the intermolecular H—bonds
(dOI'Olil= 2.62 K and dOl_OII = 2.70 A). The inter—
molecular bridge is responsible for the relatively high
melting point and low solubility of PAS, while the intra—
molecular hydrogen bridge might well explain the fact that
the dipole moment of p—amino-benzoic acid does not change
appreciably when the OH group of PAS is introduced.<l7u)

In contrast to the results obtained for aliphatic
amino acids, a hydrogen bridge between the carboxyl groups
is preferred in this case to that between an amino group
and the carboxyl group. This can be explained by the
fact that the NH2 group is conjugated with the benzene
ring and therefore zwitterion cannot be formed. Thus,
the molecules are uncharged in the crystal.

Assuming that the four formulas shown on the following
page contribute to the ground state of PAS, the best agree—
ment between measured and calculated distances was ob—

tained<l66>

by assuming the following contributions:
I 58%, (III) 20%, (Iv) 13% and (II) 9%.

The Kekulé formulasaccount for only about 60% of the

 

 

 

 

1H2

H H
I _ '4.
l I
O 0
NHz NH2
I N
H H
| _ l
H O O H‘ O 0
9° 0
buiz +IVH2

11 III

total contribution, and this may explain several properties
of this substance such as the high rate of decarboxylation,
the low pK value, corresponding to the ionization of the
amino group, and the marked variation of the absorption
spectrum with pH.

The existence of intramolecular hydrogen bonds in
PAS in the solid phase encouraged us to look for intra—
molecular proton transfer in the excited state. For this
reason we measured the fluorescence spectrum of PAS in a
KCl pellet. We observed a broad emission band with a
maximum in the visible region (WANG nm) (Figure H2). Ex—
citation at 300 nm and 370 nm gives the same emission
spectrum. The observation of this emission band in the

solid phase may be rationalized as due to proton transfer

 

 

143

 

poaaoa Hum SH m<m go Ednpooam cocoomosOSHm osSmeoQEop Eoom

AES 265333
00m omv 00¢
. .

 

 

 

 

 

.m: opsmfim

Ausuelul eoueosalanld

 

 

 

 

 

 

14H

occurring in the excited state of the solid phase since
excitonic emission cannot account for such a large Stakes
shift. Our fluorescence spectral measurements of PAS in
solutions under various conditions such as the effect of
solvent and pH, support the suggestion that the emission
at mAMO nm originates from the zwitterion species which is
produced in the excited state Via an intramolecular proton
transfer from the hydroxyl group to the oxygen of the car—
boxyl group.

The study of fluorescence spectra of MSA in solution
made it clear that the ground state equilibrium between
different conformers is important in understanding the
observed emission. Thus, in non hydrogen bonding solvents
the closed ring conformer (cis form) and hence the zwitterion
emission was the most predominant one. However, in hydrogen
bonding solvents a longer excitation wavelength gives rise
to emission from the zwitterion while a shorter excitation
wavelength gives rise to a normal fluorescence emission
band reflecting the existence of at least two conformers.
Since the crystal structure shows that all of the molecules

exist in an intramolecular hydrogen—bonded form, i.e. that

 

the most predominant conformer in the solid phase is the
intramolecular hydrogen banded cis form, one would expect
only Zwitterionic species at «#40 nm. Excited state inter—
molecular proton transfer cannot occur in the solid phase
of PAS since the proton of the carboxyl groups are involved

in intermolecular hydrogen bonds and the acidity of this

 

 

 

 

 

 

 

 

 

1H5

proton does not increase substantially upon excitation.
The absence of the normal fluorescence band in the UV
region which is due to neutral PAS in Solution is due to
the type of inter- and intramolecular hydrogen bonds

present in the solid phase of PAS.

D. Excited State Proton Transfer in N,N-Dimethyl P-Amino

Solicylic Acid (DPAS) and Its Methyl Ester (DPASE)

l. Solvent Effect on The Absorption Spectra DPAS

and DPASE

 

The room temperature absorption spectra of DPAS (2.5
x lO_5 M) have been measured in different solvents and
are shown in Figure 43. The energies of the two near UV
absorption bands in different solvents are summarized in
Table 9. In water at pH 2.“ the spectrum shown in Figure
”3 corresponds to the neutral form. Depending on the
solvent intramolecular or intermolecular hydrogen bonds
are formed. The results are similar to that of PAS with
a very important exception; namely, an energy—level reversal
of the two charge—transfer states. The N,N—Dimethyl amino
group has a lower ionization potential than the amino group,‘
which lowers the energy level corresponding to the charge—
transfer state in which the N,N—dimethyl amino group acts
as an electron-donor. In fact, the electronic transition

to this state is lower in energy than C.T. transition in

 

 

 

l
l

hr

LO

 

0.6-

0.55,:

0.4 "

s
3 2
O O
OUCOOuOmD d.

»
AU.

2

 

 

1A6

 

0.6

o
w

A bsorbance

.0
m

O.|

 

 

 

 

200

250 300 350
Wavelength (n m)

Figure U3. Room temperature absorption spectra of DPAS

in different solvents: (
ether; ( ----- ) ethanol; and (....) H20 (pH
U).

 

) 3MP; (----)

1A7

 

 

 

 

 

 

 

 

Table 9. Room Temperature Absorption Maxima (nm) and
Spectral Shifts (cm-1) Relative to 3MP for
DPAS in Different Solvents.
A(nm) _ -l l(nm) l
Solvent (O)CT Av(cm ) (N)CT A§(cm' )
3MP 295 O 319 O
Ether 291 +U66 311 +806
EtOH 292 +3U9 313 +601
H20 295 O 313 +601
Table 10. Room Temperature Absorption Maxima (nm) and
Spectral Shifts (cm-l) Relative to 3MP for DPASE
in Different Solvents.
y<nm) _ ‘1 l(nm) _ _1.
Solvent (O)CT Av(cm ) (N)CT Av(cm )
3MP 29H 0 309 O
Ether 295 -116 313.5 -A6A
EtOH 295 —116 316 —7l6
H20 290 +u69 318 ~915

 

 

 

 

 

 

1A8

which the hydroxyl group acts as an electron—donor. The
room temperature absorption Spectra of P—amino benzoic acid
(PAB) and its N,N-dimethyl derivative (DPAB) have been
measured in water at a pH value of 2.A where the amino
group is not protonated. For these two compounds the only
charge—transfer transitions occur from the amino group to
the aromatic ring. This charge transfer band lies at ”277
nm for PAB and at “315 nm for the N,N-dimethyl derivative.
The comparison of the spectra of the neutral forms of DPAS
(m290, 313), DPAB (315), and SA (m302), (in SA there is only
one charge transfer band due to hydroxyl group) supports
our assignments of the charge transfer bands (Figure 44).
Another confirmation of our assignment has been obtained
from the study of the absorption spectra of DPAS in ethanol
and in water at different pH values. This energy level-
reversal has an important consequence to the emission
properties of DPAS which will be discussed later.

The absorption Spectra of DPASE (2.5xlO-5 M) have been
measured in different solvents and are Shown in Figure 45.
Absorption maxima and Spectral shifts are summarized in

Table 10. The results are very Similar to those of DPAS.

2. pH Effect on The Absorption Spectra of DPAS and

DPASE

The spectra of various species of DPAS have been

identified by studying the effect of pH on the absorption

 

 

0.4 ~

II
8033

8 h:
B 0.2

0.| ~

 

Absorbance

.0

1U9

 

 

 

l ’T'
3l3
: l n
= l- 'I’ ‘ 35
2 V 1 \ ~v2§9<3 .‘O'X -
' 3 I: ‘ I 1
.| I: \ . '
l
I l
- ‘| I l a
l l
‘ 302
f‘\
l // \
r l / \ -
| / \
/ \
l / \
|\._.—/ I l ‘\._

 

Figure 44.

200 250 300 350

Wavelength (nm)

Room temperature absorption spectra of (————)
N, N— —dimethyl p— —amino salicylic acid, (. ,)

N ,N— —dimethyl p- amino benzoic acid, and (————)
salicylic acid in their neutral form in H20

 

 

 

 

150

 

N
.b

20

Molar Absorbabllity x U3

 

 

 

I l J l \n. I.

200 J ' .256 300 350
Wavelength (nm)

1 j

 

 

Figure 45. Room temperature absorption spectra of DPASE
in different solvents: ( ) 3MP; (____)
ether; ( ————— ) ethanol; and (-.._..) H2O.

 

 

 

 

 

 

 

 

151

spectra and the results at selected pH values are shown in
Table 11 and Figure A6.

At pH 0.8, where the amino group is protonated, there
is only one charge—transfer band at N300 nm involving the
hydroxyl group. The spectrum in this case is very similar
to that of salicylic acid at the same pH value. At pH 2.A
in which DPAS exists in its neutral form, the charge-
transfer band due to the amino group appears at 313 nm
while a shoulder appears at 290 nm. In the pH range A.l—
12.5 DPAS exists as an anion and the whole spectrum is
shifted to higher energies compared to the neutral form.
This probably reflects the lower electron affinity of the
carboxylate anion compared with the carboxyl group giving
rise to an increase in the energies of both charge transfer
states. At pH 13.6 DPAS exists as a doubly-charged anion
and the charge-transfer band due to the carboxylate anion
is shifted to higher energies compared to the same form of
PAS or SA. The repulsion of the two negative charges is
probably greater in this molecule, reflecting the higher
electron density of the carboxylate group in DPAS.

The effect of pH on the absorption Spectra of DPASE
are shown in Figure A7 and the results are summarized in
Table 12. The spectra of the cation as well as the neutral
form are very similar to those observed for DPAS. In this
case, neither carboxylate nor doubly-charged anions are
formed. Instead we have the hydroxylate anion formed at

high pH values. The absorption Spectrum of the phenolate

 

 

 

 

152

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

mflmmovz" mflmmovz”
we: . som 0mm e.ma
I0 I0
O“U/Ol O“O/Ol
mammovz” mfimmovzn
OH: moxflmu mom 2mm m.m
Io
mflmmevz” mgmmevz”
3H: MHM O N .
o moIhflwu : m
emu/om one/om
m m +
mflmmovz. A movzm
am: e as I.. com s.s
mocoIom one/om
+
AEQV moflooam opmum oopfloxm woflooam oompm US$090 AEQV ma
cause: cause: soapstonea
soawmflsm
.mofloomm mopmpm oopfioxm use ocsonw «mosam>
mm psosommflm pm m<mo Log mEflme COHmmfiEm ocm COHpaLOmQ< chapmpoQEoB Eoom .HH oHomB

 

 

 

 

v!‘-.~- r- ‘Cl .—

 

08 -

07 r

0.6 -

m to
O O
QUSOLOwDQ

 

 

OB

 

 

 

 

I I | ' ' '
' l
I . l /’\.
(3:7’— ‘ ! 1 xx / \ ‘-
I 1 I \z' \
I ‘ 1 .I l
0.6” l - \ .I l "
I! ! l
l I \ [I l
05 - l '\ ‘x \ . ,. .\ . "l
. . I \ / -. I
\\ \.-/ \ '- I'./ V \ I
g; (1;: _ \ \'\/J l -
c3 \ \J! _
.D \ '
8 l l)
3 0-3‘ \ 'Il ‘
4 l :l
‘l
- l /‘ " -
0.2 ‘ / \\ I
l / \ I
l / \ I
01‘ \ // \ '
\_,/ \ \
\ \
\ 3
00 I I ‘\
200 250 300 350
Wavelength (nm)
Figure A6. Room temperature absorption spectra of DPAS
at different pH values: (--—-) O. 8; (——-—)
2 A, (- - —) H2O (A. 1-12. 5) and (_. _ .-)

13.6.

15A

 

.0
on

 

£3
J5
I

 

.0
04
l

Absorbance

 

.0
DD
1

Q
l

 

 

 

 

 

200 250 300 350'”
Wavelength (nm)

 

Figure A7. Room temperature absorption spectra of DPASE
at different pH values: (-——-) 1.1, (————)
H20; and (....) 11.9.

 

 

 

 

 

 

 

 

155

 

 

 

mAmmovz”

 

 

 

mgmmevz”
am: 0 mmm 0mm 9:
IO IO
one/Ommo owe/cmmo
mammovz" mAmmovz“
cam O Q a saw 11m
$0 $0
owe/cmmo one/cmmo
+
mAmmovz" mflmmovzm
om III oom H.H
: I0 so
+30¢oxommo one/cmmo
AEQV ma
woflooa o m o flex moflooa o mam oesonw AEQV
causes m p cm s e m m c cease: scaeascaca
COHmmHEm
.moflooam mopmpm oopfioxm ocm ©250h©.mo3Hm> . o nob
ma unapomMHQ pm mm<mm pom mEmez scammflsm one cospdnomn< asapmsoasoe Eoom ma H

 

 

 

 

 

 

 

156

anion lies at longer wavelengths than that of the doubly—

charged anion, reflecting a larger electron affinity of

the carboxyl group.

3. Emission Spectra of DPAS and DPASE As A Function

of pH at Room Temperature

The room temperature emission spectra of DPAS in water
at different pH values are Shown in Figure A8. The
excitation wavelength is 290 nm which corresponds to the
first absorption band. The ground state and excited state
Species at given pH values are shown in Table 11. The
results are very Similar to those of PAS.

The results of the pH effect on the emission spectra
of DPASE are shown in Figure A9 and Table 12. At pH 1.1,
there is only one emission band at A50 nm which is attributed
to the zwitterion formed during the lifetime of the DPASE
excited-state via a double—Proton Transfer process involving
two solvent molecules. This is in contrast to PASE at about
the same pH region, where a very weak, short wavelength
Shoulder is seen. When the neutral form is excited,there
is only one emission band at m380 nm which is probably due
to the neutral form, but in the case of PASE the excitation
of the neutral form results in emission at two wavelengths,
one at m367 nm which is due to the neutral form, and the
other at A50 nm which is due to the zwitterion. At pH 11.9

where DPASE exists in the anion form emission occurs at

 

 

 

 

 

 

 

157

 

 

 

 

 

 

A?
(D
C
0
E
0—0
G)
U
C
Q)
U
(D
0)
n.
O
2
LL
1 ‘-
350 400 450 500 550
Wavelength (nm)
Figure A8. Room temperature fluorescence spectra of DPAS
at different pH values: (----) 0.8, ( )

2,A, ( _____ ) H2O, (---—---)13.6

 

 

a

 

 

xtntmzpfi 02500.3.0503‘0

 

 

 

 

158

 

Fluorescence Intensity

 

 

 

 

Figure A9.

-1 l ‘

350 400 450 560 550
Wavelength (nm)

Room temperature fluorescence spectra of DPASE

at different pH values: (———-) 1.1, ( )
H20, and (....) 11.9.

 

 

 

 

 

159

A20 nm which is at longer wavelengths than that of the

doubly-charged anion of DPAS (mAlO nm).

A. Fluorescence Spectra of DPAS and DPASE in

Different Solvents at Room Temperature

The room temperature fluorescence spectra of DPAS in
different media are Shown in Figure 50, and the emission
maxima in Table 13. There are two emission bands in 3MP
and ether, one at mA60 nm which is attributed to a zwitter-
ion produced via an intramolecular Proton Transfer from the
hydroxyl group to the carboxyl group, which occurs during
the lifetime of the excited state. The other emission band

at 35A nm corresponds to the neutral form of DPAS.

Table 13. Room Temperature Emission Maxima in nm for
DPAS in Different Solvents.

 

 

 

Solvent Emission Maxima in nm
3MP 35A A60
Ether 35A “58
EToH - A00 A60
H2O (2.A) AlA

 

In ethanol there are two emission bands, one at mAOO nm

due to a species produced via an intermolecular Proton

 

 

III]

. >.3C0.CH Ouguntg. a

 

 

160

 

Fluorescence Intensity

 

 

 

 

250

Figure 50.

4 " "nu-"1.4 1 1...”...
300 350 400 450 500' 550
Wavelength (nm)

Room temperature fluorescence spectra of DPAS
in different solvents: ( ) 3MP; (----)
ether; ( ----- ) ethanol, (....) H20 (pH = 2.A).

 

 

 

 

161

Transfer from the hydroxyl group to a solvent ethanol
molecule producing a singly-charged anion, and the long
wavelength emission band at WA6O nm which is a shoulder
and is attributed to the zwitterion produce via an inter-
molecular double proton transfer between the solute and
the solvent. One important difference between PAS and its
N,N-dimethyl derivative is that in DPAS the carboxyl oxygen
is more basic than in the case of PAS and tends to be
protonated in the excited state. In the case of PAS in
ethanol we have two emission bands, one due to the neutral
form (“3AA nm) and the other due to the phenolate anion
(“A16 nm). In water at pH 2.A, where DPAS exists in its
neutral form, there is only one emission band at AlA nm
which is due to a singly-charged anion as was observed in
PAS in the same pH region.

The fluorescence spectra and emission maxima of DPASE

in different solvents are shown in Figure 51 and Table 1A.

Table 1A. Room Temperature Emission Maxima in nm for
DPASE in Different Solvents

 

 

Solvent Emission Maxima in nm
3MP 350 A68
Ether 360 A5A
EtOH 366 A92
H O 380

2

 

 

162

 

Fluorescence Intensity

 

Figure 51.

 

300 350 400 450 500 550 600

 

 

Wavelength (nm)

Room temperature fluorescence spectra of DPASE
in different solvents: ( ) 3MP; (-——-)
ether; ( ----- ) ethanol; and (....) H2O.

 

 

 

 

163

In 3MP, ether and ethanol there are two emission bands,

one at short wavelengths with a normal Stokes' shift and

is due to the neutral form. The other occurs at long wave-
lengths with a large Stokes' shift due to the zwitterionic
species. In 3MP and ether, intramolecular proton transfer
occurs whereas in ethanol intermolecular proton transfer
occurs. In water, as we mentioned earlier, there is only

one emission band due to the neutral form of the molecule.

E. A Unique Excitation Wavelength Dependence of Excited—

State Proton—Transfer

 

A comparison of the relative intensities of the Short
(35A nm) and long (m A60 nm) wavelength emission bands of
DPAS in 3MP indicates that the intensity of the 35A nm
emission band is higher than that of the A60 nm emission
band which is the reverse of the Situation with PASE and
MSA. The high intensity of the short—wavelength emission
band relative to that of long wavelength emission band
prompted us to investigate whether the energy—level reversal
Of the two charge transfer states in DPAS could be respons-
ible for this observation.

To verify this suggestion we studied the fluorescence
spectra of DPAS in 3MP at two different excitation wave—
lengths (Figure 52). Excitation at the 330 nm band gave
mainly the emission at 35A nm, while excitation at 300 nm

gave two emission bands, one at 35A, the other at mA60 nm.

 

 

 

 

16A

 

 

 

 

T A I I l
I
l
l
l
\
l
>‘ I
3:
m l
#5. I /’ \\
5 I / \
(D / \
2 I I / \
/
A ‘ / \
g ‘I [x \
g \\ // \\
u- "" \
\
__;____““~___
1 J 1 J __
350 400 450 500
Wavelength (nm)
Figure 52. Room temperature fluorescence spectra of DPAS

in 3MP at two different excitation wavelengths:

(---~) Aex = 300; < ) *ex = 330.

 

 

 

 

 

 

165

It is most interesting that at a longer excitation wave—
length (330 nm) the emission occurs at 35A nm while a
shorter excitation wavelength (300 nm) gives rise to an
emission at A60 nm as well as at 35A nm. It is important
to note that in the case of MSA, shorter excitation wave—
lengths gave mostly an emission at 3A0 nm while longer
excitation wavelengths gave mainly emission at A50 nm.

To interpret these data, we have constructed an energy
level diagram for DPAS (Figure 53a). From the absorption
maxima we calculated that the energy of the (0) charge
transfer state lies at roughly 2600 cm_1 above that of the
(N) charge transfer state.

In the (N) C.T state where the amino group acts as the
electron donor, the basicity of the oxygen of the carboxyl
group is expected to be larger than that of the ground state,
but the acidity of the hydroxyl group is not expected to
Change appreciably. In this case, no proton transfer is
expected for the hydroxyl group as a result of excitation
of the (N) C.T state. In contrast, an excitation of the
(O) C.T state where the hydroxyl group acts as an electron
donor, both the acidity of the phenol group and the basicity
of the oxygen of the carboxyl group are expected to be
increased relative to their values in the ground state.
Thus excited state proton transfer is expected to occur as
a result of excitation to the (O) C.T state unless other
deactivation processes are much faster than the rate of

proton transfer. One can therefore expect two possible

 

 

 

——_—-————

166

 

 

 

““92 l
(O)CT 1"- 12. M -" --
: '0 2600cn‘l'
(N)CT 11:1 —T—-
. —- -
' l
l I
300 '
I 4§O
'330 l
l
I ___L_
I 359
I I +
_J_ 01......Hogc
0-Hmmno.c

Figure 53(a). Energy level diagram for DPAS (change in em-
ission due to excitation of different elec-
tronic state of the same conformer).

 

 

*I‘
O\\c/O\r
6°
‘\-- --
HO\ ‘ ..'H/-“ | :
C I II H
0 I I l
: I axe/10H
3'0: : 450 :9"
3?O| l
'390
I l
g I
=E'1—J' - 1...... 600cm"

‘T"

Figure 53(b). Energy level diagram for two conformers of
MSA (change in emission due to excitation of
different conformers).

167

pathways depending on the excitation wavelength. If the

(0) CT state is excited radiationless transition will occur
with a rate constant of 1012 sec.1 to the lowest excited
singlet state, (N)CT; giving rise to emission at 35A nm.
However, excited state proton transfer may occur during the
lifetime of the (0) CT state giving rise to the zwitterion
Species in the excited state and leading to an emission at
A60 nm. The intensity of the A60 nm band indicates that the
rate constant for proton transfer must be of the order of
lo12 sec—l. Excitation of the (N)CT band will give rise

to 35A nm emission only.

Figure 53b also shows the energy level diagram of MSA.

-1
The diagram shows an energy gap of 600 cm for the two
ground state conformers. This value is obtained from
(100)

theoretical calculations In this case excitation to

the lower level corresponding to the cis conformer leads
to proton transfer with a rate constant of 1011 sec—1
resulting in an emission band at A50 nm. However, excita-
tion to the higher energy level, i.e., excitation of the
ether bonded conformer where no proton transfer may occur
gives rise to an emission band at 3A0 nm.

A comparison of the two energy level diagrams clearly
shows that in the case of DPAS, we are dealing with excita-
tion of two different electronic states of the same conformer.

In contrast, the emission of MSA is varied as a result of

excitation of different conformers which are at equilibrium

in the ground state.

168

Consistent with our suggestion, we expected different
excitation spectra for these two emission bands. The excita—
tion spectra monitored at two different emission wavelengths
are shown in Figure 5A. The excitation maximum corresponding
to the 35A nm emission occurs at 310 nm while the excitation
maximum corresponding to the A6A nm emission occurs at 300
nm. In the case of MSA<96), emission at 3A0 nm gives an
excitation maximum at 308 nm and emission at A50 nm gives
an excitation maximum at 311 nm. Our experiments also sup-
port our earlier assignments of charge transfer states, and
to our knowledge, this result is the first example of this

unique excitation wavelength dependence.

169

 

 

 

 

 

 

T 1 ‘57
3IO
>5
4-
"I5
C
.9.’
S
Q)
U
C
8
U5
8
O
9..
LL
300 350
Wavelength (nm)
Figure 5A. Room temperature excitation Spectra of DPAS

 

in 3MP (uncorrected spectra) ( ) represents
the 35A nm band, while (--—-) represents the
A6A nm fluorescence.

CHAPTER III

EXCITED STATE DOUBLE PROTON TRANSFER IN
1-AZACARBAZOLE HYDROGEN BONDED DIMERS
ITS ACETIC ACID COMPLEX AND

IN THE SOLID PHASE

A. Introduction

 

Both l-azacarbazole hydrogen bonded dimers and the
1-azacarbazole-acetic acid complex were shown(175) to
undergo photoinduced double proton transfer reactions in
their lowest excited singlet state. An emission band with

a maximum at 510 nm arises from a tautomer formed in the
excited state as a result of the double proton transfer
process.

Studies of l-azacarbazole in the solid phase also Show
that excited state double proton transfer accounts for the
observed emission. The crystal structure, determined by
x-ray diffraction methods(l76), shows that the crystal
contains cyclic dimers only. Spectral measurements of
l—azacarbazole in solution under various conditions provide
further details regarding the mechanism of excited state

proton transfer. The absorption and emission properties of

l-azacarbazole- were reexamined under various conditions.

170

171

l. Excited State Double Proton Transfer in l—AZA

Carbazole Hydrogen—Bonded Dimers

Dilute solutions of l—azacarbazole in 3-methylpentane
(3MP) exhibit a fluorescence band with a maximum near 367 nm.
Concentrated solutions of l-azacarbazole in 3MP, in addition
to the normal fluorescence (Fl) at 374 nm with a small
contribution from the dimer, exhibit a second fluorescence
band (F2) with a maximum at about 510 nm (Figure 55). The
relative intensities of Fl and F2 depend on the concentration,
excitation wavelength and temperature.

The room temperature absorption spectrum of l-aza
carbazole in 3MP (Figure 56) exhibits an absorption band
covering the range from 300 to 350 nm. This band corresponds

to the lAlI'lAl (lLa) transition of carbazole which has a

transition moment along the short axis of the molecule.
Such an assignment was obtained from polarized absorption
of carbazole in a single—crystal matrix of fluorene at about

1 8
l5 K<l77). An emission polarization study( 7 ) of stretched

polyvinylalcohol films confirmed this result. The same

Study showed that the second absorption band with a maximum

1 l l . . f
at 296 nm corresponds to the Al+ B2 ( Lb) tranSition o

carbazole and is polarized along the long axis of the mole-

cule. These results are consistent with relative polariza—

tions obtained from luminescence polarization studies of

(179).

frozen solutions

172

 

Relative Intensity

 

300

Figure 55.

 
 
 

 

 

. 2XIO‘5 M
——-—6X|O'4 M

 

I \

 

 

350 400 450 500 550 600
Wavelength (nm)

Room temperature
azacarbazole in 3
trations.

fluorescence spectra of 1-
MP at two different concen-

173

 

 

 

 

1
\
‘ I}
l I ‘
ao— \ / ‘
\ I ‘
z \ I
9 25-- 2;\ /
" '. \ I
5 '1, \ I
§ 20... 1. \\ I,
g . \v“ I '
f L... A, l’
2 -'l .
g ‘ / I
I.O- ‘\ I \
\ l: \\A’,
\ .' .
05*- V. ' ....... \.
'\- .
l J Li J J l\\"
240 260 280 300 320 340

Wavelength nm

Figure 56. Room temperature ultraviolet absorption spectra
of l-azacarbazole (2 x 10-5 M) in different
media: (---—) cyclohexane; (————) ether; and
(....) ethanol.

17M

Solvent effects on the absorption spectrum of l—aza

carbazole indicate that the hydrogen bond energy has the

l 1

following order:1L > Lb > A. Analysis of the spectral

a

shifts due to change in medium leads to the qualitative
conclusion that the acidity of the pyrrolic proton and the
basicity of the aza—nitrogen of l—azecarbazole are both
enhanced, particularly in the 1La state, the lowest excited

singlet state. Figure 57 clearly shows the large red shift

of 1La band of dilute solution of l—azacarbazole in 3MP at

77 K. Solvent effects on the emission spectrum of l-aza

carbazole also are consistent with a larger dipole moment
in the 1La state compared with that of the ground state.

The ground and excited state dipole moments of carbazole

were found(180) to be 1.7 and 3.6 D respectively.

If equilibrium is established during the excited state

* a a
lifetime, then the pKa and pK; can be determined in a way

analogous to the spectraphotometeric determination of the
ground state pK's, by fluorometric titration. For this

reason we have repeated earlier measurements of the fluores—

cence and absorption intensity of l-azacarbazole in water

as a function of pH. In the range from 7 to 13, neither

the absorption spectra nor the emission spectra show any

change. Above a pH value of l3, the absorption spectra

remain the same; however, the emission spectrum exhibits

the A00 nm band as well as a shoulder around the 500 nm

region (See Figure 58). The new emission band can be

clearly identified as the emission of the anion of l—aza

 

175

 

 

Absorbance

 

 

 

 

O 1
250 300 350

Figure 57. Ultraviolet absorgtion spectra of l-azacar—
bazole (2.5 x 10‘ M) in 3—methylpentane at

room temperature (~—-)3 77°K (————).

 

 

176

 

Fluorescence Intensity

 

 

 

 

l 1 1 L l
350 400 450 500 550 600 650
Wavelength (nm)

Figure 58. Room temperature fluorescence spectra of l-
azacarbazole at different pH values: ( )

“'13), (....) 13-59 ( """ ) 1307, ("’")>13-7-

 

177

carbazole which exists in pH Z 14 solution. From the above
discussion, we can conclude that the acidity increases upon
excitation to the La state. In other words, the proton
will dissociate more easily in the lowest excited singlet
state than the ground state. Other evidence that supports
our idea is the dramatic decrease of the fluorescence
quantum yield around pH 13.5 (see Figure 59) which is due
to the proton dissociation of the pyrrolic nitrogen and
production of the weakly emitting anion. From the midpoint

of the fluorescence and absorption titration curves we
*
b b

ground states respectively which show that indeed the acidity

obtain a pK = 12.12 and pK = 13.A2 for the excited and
of pyrrolic hydrogen increases by 1.3 pK unit. In the same
way we get pK: = A.2 and pKa = 2.A at low pH due to protona—
tion of the aza nitrogen which indicates that the basicity

of the aza nitrogen increases by 1.8 pK unit. The actual
magnitude of the pK value change due to excitation has been
estimated (by using the Forster cycle) to be Abe = 1.6 and
ApKa = “.9 for the pyrrolic hydrogen and aza nitrogen respec-
tively.

Hydrogen bonded dimer formation of 1—azacarbazole has
been studied by comparing the ultraviolet absorption Spectra
at different concentrations in 3MP at room temperature
(Figure 60). A shoulder appears at 353 nm in the more
concentrated solution which corresponds to the dimer. The

association constant of dimer formation as obtained from

absorption spectra at different concentrations is

178

.mpzpmmeEmu
Eoop pm maowmnhmomwmla mo wLuomdm mocoommzosfim one no mo mo powwow .mm mpzmfim

V. m_ N. : 9m m N m n v m N _

d d

 

f

 

Ausualul allumaa

 

 

 

 

 

 

 

 

 

179

 

 

 

 

 

 

F'TTT‘1 I l I I I 1 ~1- 1
“5
4 l 6xlO
2 8xl0‘5
3 l x IO”4
\N 4 2xl0‘4
8
c:
8
§ 3
.0 2
<1 I
l I ii I L_ 1 i4. .L

 

 

 

 

300 350
Wavelength (nm)

Figure 60. Room temperature absorption spectra of l—aza-
carbazole in 3—methylpentane at different con_

centrations.

 

180

-l

(6.8:0.7)x102 M at room temperature (21°C). This

corresponds to a AG value of 3.8 Kcal mole_l at room

temperature (8.1i0.8)x103 M-1 at 0°C and W108 M-1 at 77°K.

The association constant in 3MP as a function of temperature
resulted in an estimate of -l8.25 Kcal mole_l for AH and
-A9 eu for AS. The large negative value of the enthalpy
change means the system is energetically favorable towards
dimerization at low temperature.

The excitation spectrum (mlO—uM) for the F band is

2

shifted to longer wavelengths compared with that for the F1

band indicating that the F2 fluorescence band arises from

excitation of the hydrogen bonded dimer which absorbs at
longer wavelength compared to the monomer. The F1 fluores—
cence band arises mainly from monomer excitation.

The assignment of the F band as being due to a tautomer

2
formed as a result of double proton transfer in the excited
state of the '1—azacarbazole hydrogen-bonded dimer has been
further confirmed by studying the emission spectrum of the

N -methyl tautomer of l-azacarbazole (I) which we have

1

synthesized.

CH3
I

This molecule has the same w-electron system as the prOposed

tautomer. Its emission spectrum in ethanol occurs at the

 

181

same energy and exhibits similar fine structure to the F2
band observed by excitation of the 1-azacarbazole hydrogen—
bonded dimer.

By using nanosecond time resolved spectroscopy(l8l),

the rate constants of double deuteron transfer has been

estimated at 77°K to be k? = 6.1 x 107 and k3 = 2.1 x 107
sec—l. The large rate of the double deuteron transfer points

to a very low barrier energy for the process. The energy
barrier for excited—state proton transfer in the l—aza
carbazole hydrogen—bonded dimer can be estimated by studying
the temperature effect on the F2/Fl intensity ratio. An

1

estimated energy barrier is about 3 Kcal mole— A potential

energy diagram which qualitatively describes the energetics
of the double proton transfer process is shown in Figure 61.
Potential energy is plotted versus a reaction coordinate
associated with the simultaneous motion of the two protons

in the dimer. It is assumed that the tautomer is unstable

in the ground state.
-1
By using an activation energy of 3 Kcal mole and

l
k = 6.1 x 107 sec—1 one may obtain a value of 10 3 for the

pt
frequency factor A in the equation for the absolute rate

constant, which is similar to the expected theoretical value

12

10 m 1013.

2. Phosphorescence of l—Azacarbazole

In the case of 7—azaindole where photo-induced double

(79)
3

proton transfer occurs in concentrated solution in 3MP

182

 

 

 

 

F?h+_++-———a-

c 03"" W

I?

33..
._;J: :

Dimer Tautomer

 

 

 

Figure 61. Qualitative double minimum potential energy
diagram illustrating excited-state double
proton transfer in l-azacarbazole hydrogen-

bonded dimers.

 

183

the low temperature fluorescence spectrum overlaps with the

isoenergetic monomeric phosphorescence(127). This coinci-

dence makes it difficult to remove ambiguities regarding

the temperature dependence of the two fluorescence bands

and makes the mechanism of proton transfer less certain.

The phosphorescence Spectrum of 1—azacarbazole was

measured at 77°K and is shown in Figure 62. The maximum of

the phosphorescence band occurs at MAO nm while the fluores—

cence band due to the excited state double proton transfer

of hydrogen—bonded dimers occurs at 510 nm. Thus, in the

case of l-azacarbazole no ambiguity is present.

3. Excited State Double Proton Transfer in the

l—Azacarbazole-Acetic Acid Complex

1-Azacarbazole forms a complex with acetic acid. The

room temperature absorption spectra of a dilute solution

in hexane, 2 x 10"5 M of l—azacarbazole in the presence of

various amounts of acetic acid are shown
absorption at 353 nm is due primarily to
can be used to determine the association
complex. A value of A X 105M—1 has been

association constant of the complex. It

in Figure 63. The
the complex and
constant of the
estimated for the

should be noted that

the association constant for the l—azacarbazole acetic acid

complex is much larger than that for the

mer (6.8 x 102M_1).

l—azacarbazole di—

The room temperature emission spectra of l—azacarbazole

in the presence of excess acetic acid exhibits only the F2

184

 

Intensity

 

 

 

J

 

l n LJJI I 1 1

300 40 500 GOO
Wavelength (nm)

Figure 62. Phosphorescence spectrum of 1—azacarbazole
(6 x 10‘ M) in ether at 77°K excited at 330

nm .

 

 

     
 

 

 

l L \

 

413—-
3£5-
,0 3.0 —
O
S?
)<
> 2.5—
3
l.
o
m 2x)-—
.0
O
L
3
0 L5 '—
:5
LO--
015——
Figure 63.

320 340 360 380
Wavelength

Room temperature absorption spectra of l—
azacarbazole (2 x 10_5 M) in hexane, in the
presence of variable amounts of acetic acid,
0, 1.66 x 10‘6, 3.32 x 10'6, 4.98 x 10‘6,
6.6A x 10‘6, 8.30 x 10-6, 9.96 x 10‘6, 1.16
x 10‘5 M in 10 cm cell.

186

band as shown in Figure 64. Similarly, at 77°K only F2

emission is observed. This means that the proton transfer

is much faster than the fluorescence rate of 1—azacarbazole

The estimated value for kpt is 1.“ x 109 sec—l, so that the

activation energy in the excited state must be less than

I

0‘16 Kcal mole_ and consequently, the thermal energy at

77°K is enough to overcome the barrier.

A- Excited State Double Proton Transfer in l—Aza

 

Carbazole in the Solid Phase

Crystallographic data

The crystal structure of l-azacarbazole has been deter-
mined(l76) by x-ray diffraction. Crystals of l-azacarbazole

are monoclinic and belong to the space group P2l/n; the
dimensions of the unit cell are: a = 13.516(A), b = 5.526
(2), c = 11.336(5)Z, B = 95.77(3)O; z = u, M = 168.20, pC =
1.326 g cm_3. Lattice dimensions were determined using a
Picker FACS-I diffractometer and IVIOKOLl (A = 0.70926 A

radiation. Intensity data were measured by using Mqu

radiation (26 ax = 55°) yielding 1940 total unique data
m

points and, based on I>30(I), 1236 observed data points.

The data were reduced(l82), the structures were solved by

direct methods(183), and the refinement was by full-matrix

least squares techniques<18u). The finalll.value was 0.035.

The final difference Fourier map showed densities ranging

187

 

Relative Intensity

 

 

 

/”\\
/ \ \
/
/ \
l ’ J \.\
300 400 500 500

Wavelength (nm)

Figure 6A. The emission of 5 x 10‘ M of l-azacarbazole
in 3—methy1pentane in the absence ( ,
presence of excess acetic acid (—-——).

 

 

188

from +.19 to -.28 with no indication of missing or
incorrectly placed atoms.

The planes of the three rings individually are flat to
within $0.01A. The plane of the three rings together is
parallel to the plane of the equivalent molecule joined by
hydrogen-bonds and related by a crystalographic center of
symmetry, but these two planes are not coincident. The
distance between the two planes is’V0.5lA and the angle
between either plane and the plane of A nitrogen atoms
involved in the H-bond is 10.250. Figure 65 shows stereo-

(185) D

scopic View 01 a hydrogen-bonded dimer molecule. The
stereoscopic views of the packing of the dimer molecules

into the unit all are shown in Figure 66. The intermolecular
distances, the bond distances and the positions of the

hydrogen atoms were all consistent with doubly hydrogen—

bonded '1—azacarbazole dimers in the crystal.

Room Temperature Absorption and Fluorescence Spectra

of l—AzaCarbazole in the Solid Phase

The absorption spectrum of l-azacarbazole in a nujol
mull at room temperature is shifted to the red similar to
that observed for concentrated solutions in 3MP, which

indicates the existence of hydrogen-bonded dimers in the

solid phase of 1—azacarbazole..

The infrared spectrum was measured in the region of A000

Cm—l to 2000 cm—1 in two different media, a nujOl mull and

 

 

 

189

 

 

 

 

Stereoscopic View of a hydrogen-bonded dimer

molecule.

Figure 65.

Figure 66.

 

 

./ j x”,
’§"
/
\ I
’ \ l
I I- I-\ \cl 1
|'\ - ’7 'I \' \r‘
’\1\/ ‘ l—\ \'| -\ l-
‘ L \ \" '\/\l \ ‘\ ‘-\
\;-\ "' ‘ .. ‘ ‘ \o\ 'V\I .\ ’\I\’
I." \ Q ; ’\/\ \’\ \\. 1‘ ‘1‘ , l \
- \ l “A sax \ I.
' \ f\ “ ’ l l .\,\ "\
.- t.- -: . u \r n l
l.\ \\ \ul " ~ I-‘
\ it“ 1 \ \w ‘l " -‘J-
. \ , I \
I :\ est}; ”‘0?\ \J \ )1“ \ g
‘ n J \ z‘ .
\IJ -‘ . II _I:\ “V“ 7 I“
.\\ ”\a l ~\ \t- .i\e- -‘\-
\ ‘ {\. \ \{\ \’\l \\ f\ \yf: ‘ rt
. ‘ \ . l " '\'\. l \‘ x. 4 \
\ r '\ \l\' \J \. ‘ " ‘\ \’\’
- \ \ \J ‘\ \\, \.\ "\}\‘ ‘\\
J a a
\J- \ \~ JJ \ \-‘ ‘\
l .1 \\ J
' \ or \ s
u .1
' \

 

Stereoscopic Views of the packing of dimer

molecules into the unit cell.

a. Viewed down
towards the
b. Viewed down
towards the
0. Viewed down

towards the

the
top
the
top
the
top

a* axis with +b to the right and +9

of the page.
b axis with +9 to the right and +3

of the page.
3* axis with +a to the right and +2

of the page.

191

a KBr’pellet. In both cases the free N-H stretching vibration
(m3A80 cm-l) is absent.

The room temperature fluorescence Spectrum of a single
crystal of 1-azacarbazole which has been grown by slowly
cooling the liquid is shown in Figure 67. The spectrum shows
clearly that an emission band occurs at W500 nm which is
similar to the band observed in concentrated solution of
l-azacarbazole. The emission is therefore interpreted as
being due to a tautomer formed via double proton transfer in
the excited state. This result is what is expected from the

fact that the crystal consists of hydrogen bonded dimers.

192

 

.mmmcq UHHom
as» CH mHonopmomNth mo ESLpoodm oocmommposam apzpmpmoEop Eoom

AES £32903
0mm 00m 09v
1 _

q u 4 d —

.So opswwm

Mgsualul aoueosetonl :1

 

 

CHAPTER IV

EXPERIMENTAL

A. Experimentally Studied Molecules

 

1. Para Amino Salicylic Acid (PAS)

 

PAS was obtained from Aldrich Chemical Company, and was
further purified by crystallization from ethanol M.W. =

153.1u, n = lAA-1A5 (dec.)

2. Methyl Para Amino Salicylic Acid (PASE)

 

PASE was synthesized by using the method of selective es-
terification of aromatic carboxylic acids which have other
reactive groups by using a Boron Trifluoride EtherateeAlcohol

(186) and was purified by crystallization from ethanol.

reagent
Confirmation of the identity of the compound was done by

measuring the melting point (120°C), the mass spectrum (MW=
167.1“), and NMR and infrared spectra (in a KBr disk). The

UV absorption spectrum is almost identical to that of PAS.

3. N,N Dimethyl Para Amino Salicylic Acid (DPAS)

 

DPAS was purchased from Aldrich Chemical Company, and

was crystallized from ethanol, MW = 181.19, MP = 1A0 (dec.)

A. Methyl NyN Dimethyl Para Amino Salicylic Acid (DPASE)

 

193

 

 

19A

DPASE was synthesized in the same way as PASE. Charac-
terization of the compound was done by melting point 85°C,
and by obtaining thermassspectrum (M.W. = 195.19), NMR
and infrared (in KBr) spectra. The u.v. absorption spectrum

is almost identical to that of DPAS.

5. l—Azacarbazole (lAC)

lAC was synthesized by using the method of L. Stephenson
and W. K. Warburton<187) with more attention to the effect
of light, and purification of starting materials. It was
recrystalized from ethanol several times. Confirmation of
the identity of the compound was done by measuring the melting
point (217°C) and from the mass spectrum. The pyrrolic
hydrogen was identified by N.M.R. with .[= -l.79. The UV
absorption spectrum is almost identical to that of carbazole.
The expected rrflfli transition must be completely hidden by
the more intense nem*' absorption band.

6. N -Methy1 1-Azacarbazole Tautomer

1

After dissolving 1 gm of lAC in toluene, we added 1.5
ml. dimethylsulfate. The solution was then refluxed for 10
hours, and cooled down. The solid was filtered and recryg
stalized from acetone and methanol. The white crystals have
been identified as d—carbolium methylsulphate. After dis—
solving in acetone, it was neutralized by adding NaoH

solution. The yellow precipitate was separated. The final

 

195

bright yellow crystals with m.p. 130 to 132°C have been
identified as the Nl-methyl 1—azacarbazole tautomer by Mass,

NMR, and UV, spectra.

B. Solvents

 

l. 3-Methylpentane (3MP)

 

A modified version of the purification method of Potts
(188) was used. Phillips pure grade 3-methylpentane (3MP)
was mixed with a 50:50 mixture of concentrated sulfuric acid
and concentrated nitric acid, then stirred for two days. Then
it was separated and stirred with concentrated sulfuric acid
for several hours until the dark red color disappeared and
the acid layer became yellowish brown. After separation, it
was stirred with a dilute solution of sodium carbonate for
one hour until CO2 production ceased. The 3MP was then
stirred several times with distilled water until the water
remained clear compared to the initial yellow color it
attained after the first stirring. After drying the 3MP
overnight over anhydrous sodium sulfate, the solution was
refluxed over sodium wires for two days and distilled. The
vapor was passed through a four foot vacuum-jacketed column
and condensed at a speed of two drops per minute. The purity
was checked by obtaining the absorption Spectrum. Passing
the distillate through a 1m column of activated silica gel

did not alter its absorption characteristics so this was not

required in the purification process.

196

2. Ethanol

Absolute ethanol was fractionally distilled through a
1 meter vacuum jacketed column. The distillation rate was
adjusted so that a very slow rate (about 5 drops per
minute) was maintained. Distillation continued until the
benzene—alcohol azeotrope was no longer present as deter-
mined by an absorption spectrum of the distilled alcohol
in a 10 cm cell. That is, the characteristic benzene UV
absorption was no longer apparent. Ethanol was then

distilled and used as needed.

3. Ether

Ether was distilled over sodium hydroxide.

A. Water

Only doubly-distilled water was used.

5. Cyclohexane

 

Aldrich Chemical Company (spectra grade).

6. Hexane

This solvent was slowly refluxed over benzophenone and
sodium wire until the blue color turned to white. Dry hexane

was used after distillation.

197

C. Spectral Measurements

 

1. Absorption Spectra

 

All reported absorption spectra were run on a Cary Model

17 spectrophotometer.

2. Emission Spectra

 

Most of the fluorescence spectra were obtained with an
Aminco—Bowman spectrofluorometer equipped with a high
pressure Xenon arc lamp and an EMT 9781 R photomultiplier
tube.

To obtain better resolved emission spectra, a multi-
component system was used. A high intensity Xenon lamp
(500W) was used as the light source, the excitation wave-
lengths were reflected by a Bausch & Lomb 10 cm grating
glazed at 300 A in a B & L 500 mm monochromator (which pro-
vides a narrow excitation band width). The excitation
illumination was focused on the sample by using quartz
lenses, and emission was detected at a right angle relative
to excitation. The emission monochromator was a Spex 1700-11
which utilized a 10 cm B & L grating glazed at 5000A. The
emission spectrum was detected with an EMI 9558 QA photo-
multiplier tube. The tube's input voltage was maintained by
a Fluke A12B power supply which was normally Operated at
1100V. Signals from the detected emission were fed to a
Princeton Applied Research HR-8 Lock-in Amplifier whose

reference was provided by a light chopper. Finally, the

 

198

amplified emission signal was displayed on a strip-chart

recorder. Phosphorescence spectra were obtained with Amino

SPF 500 equipped with a rotating can.(SLM Instruments Inc)
chopper.

3. Infrared Spectra

Infrared spectra were obtained by using a Perkin-Elmer
A57 Grating Infrared spectrophotometer. Calibration of

frequency reading was made with a polystyrene film. Samples

were examined as KBr disks.

A. Mass Spectra

Mass spectra were taken with a Finnigan EI.CI gas chromato-

graph—mass spectrometer.

5. NMR Spectra

NMR Spectra were run on a varian T—60 NMR Spectrometer

system.

 

 

 

 

 

r—'

 

 

CHAPTER V

CONCLUSION AND FUTURE WORK

Absorption spectra of dilute solutions of PAS have been
measured in different media. Depending on the solvent, an
intramolecular hydrogen bond or intermolecular hydrogen
bonds are formed. The interpretation of the absorption
spectra which are dominated by "charge—transfer” transitions
has been obtained by studying the effect of medium, pH and
substituents effect and comparison with the related compounds.
The absorption and emission characteristics of different
species which may exist in the ground and excited states have
been obtained by studying the effect of pH on the absorption
and emission spectra of PAS and its methyl ester (PASE).

This study has revealed the occurrence of excited state
proton transfer in both PAS and PASE. The fluorescence
spectra of PAS in 3MP and ether consist of two bands, one

at 3A0 nm which is due to the neutral form, and the other

at A50 nm and is due to zwitterion species produced via an
intramolecular PT in the excited state (81). In ethanol and
water we have intermolecular proton transfer from hydroxyl
group to the solvents molecule. In all solvents PASE exhibit

two emission bands, the long wavelength emission band in

 

 

 

200

3MP and ether arises from intramolecular PT. The short
wavelength emission band is solvent dependent.

The observation of these two emission bands is probably
due to an equilibrium between the ground state conformers.
Further study is needed specifically for these compounds of
three functional groups to account for ground state conformers.
Excitation wavelength dependence in both hydrocarbon and
hydrogen bonding solvents could confirm such possibility.

To identify different species in the excited state and to
measure the rate of excited state proton transfer, pico—
second time—resolved studies are required.

The crystal structure of PAS shows intra— and intermole—
cular hydrogen bridges. We measured the fluorescence spec—
trum of PAS in the solid phase. Information from the solu—
tion spectra of PAS helped to interpret the emission band
in the solid phase of PAS in terms of an excited state intra—
molecular proton transfer. Further studies including the
effect of temperature and time resolved technique may reveal
interesting information about excited state proton transfer
(EPT), in the solid phase.

Absorption and luminescence properties of DPAS and its
methyl ester DPASE also have been investigated under various
conditions. The results of absorption Spectra of DPAS in
different solvents are similar to that of PAS with a very
important exception, namely, an energy-level reversal of the
two charge—transfer states. The emission spectra clearly
reveal the occurrence of PT in the excited state (81).

Excited state PT occurs either intramolecularly or

 

201

intermolecularly depending on the medium. For DPAS in 3MP,
the intensity of the short wavelength (35A nm) emission band
is higher than that of the A60 nm emission band which is the
reverse situationwith.PASE and MSA. Excitation wavelength
dependence clearly demonstrates that at a longer excitation
wavelength (330 nm) the emission occurs at 35A nm while a
shorter excitation wavelength (300 nm) gives rise to an
emission at mA60 nm as well as at 35A nm. In the case of
DPAS, depending on the electronic state excited two pathways
are possible. Excitation of the (0) C.T. state, may lead to
radiationless transition to the lowest excited singlet state,
(N) C.T., giving rise to an emission at 35A nm or an excited
state proton transfer may occur giving rise to the zwitterion
species and leading to an emission at A60 nm. Excitation of
the (N) CT band on the other hand, will give rise to 35A nm
emission only. In the case of DPAS, we are dealing with
excitation of two different electronic states of the same
conformer. In contrast, the emission of MSA is varied as a
result of excitation of different conformers which are at
equilibrium in the ground state. Further studies such as
temperature effect and time-resolved Spectroscopy will be
very helpful in revealing more information about the mechanism
and the rate of P.T. in this molecular system.

Solutions of l—azacarbazole in 3MP at concentrations
where hydrogen-bonded dimers are formed (6 x lo'um) exhibit

both F and F fluorescence. The relative intensities of Fl

1 2

and F depend on the concentration, excitation wavelength and

2

 

 

202

temperature. All these studies as well as the study of model
compounds indicates that the record fluorescence band origin—
ates from a tautomer formed by excited state double P.T.
within a hydrogen bonded dimer. The enhanced acidity of the
pyrrolic proton and basicity of theazarnitrogen in the lowest
excited state of l—azacarbazole (lLa state) makes the proton
transfer process favorable. A preliminary study for the rate
of P.T. set 6.1 x 107 sec_1 for the rate of P.T. and the energy
barrier is estimated to be 2.95 Kcal mole—l. The l-aza-
carbazole—acetic acid complex undergo also an excited state
proton transfer that is much more efficient than in l-aza—
carbazole dimer. The phosphorescence spectrum of this compound
occurs at NAAO nm which is far from F2 emission band. This
system is ideal to study the temperature and deuterium—isotope
effect in order to determine whether a tunneling proton trans—
fer mechanism is operative at low temperature. A study of
the dynamics of excited—state proton transfer in 1-azar
carbazole using picosecond time resolves spectroscopy is
also needed.

The crystal structure as determined by x-ray diffraction
methods shows that the crystal contains cyclic dimers only.
We have also measured the UV absorption and infrared
spectra of this molecule in nujol mull and in KBr pellets;
both are consistent with the existence of hydrogen bonded
dimer in this system. A preliminary study of the fluores—

cence Spectrum of a single crystal shows an emission band at

“500 nm which is similar to the band observed for concentrated

 

 

 

 

 

 

 

203

solution of l—azacarbazole. This indicates that excited

state proton transfer occurs in the crystal.

 

 

 

 

 

REFERENCES

10.

11.

12.

13.

1A.

15.
l6.

 

 

 

REFERENCES

J. N. Brbnsted and K. J. Pederson, Z. Phys. Chem. 108,
185 (1928).

 

M. Eigen and L. deMaeyer, Z. Electrochem. 59, 986
1955). ‘—

For this theory, P. Mitchell has been awarded the Nobel
Prize in Chemistry for 1978; see A. L. Lehninger,
Biochemistry (Worth, NY, 1970).

A. J. Kresge, Accts. Chem. Res. 8, 35A (1975).

Isotopes in Organic Chemistry, E. Buncel and C. C.
Lee, ed. (Elsevier, Amsterdam 1976).

R. A. Marcus, Faraday Symp. 10, 60 (1975).

M. Eigen, Angew. (Chem. Int. Ed. Eng. 3, 1 (196A))
and references therein.

E. Grunwald and E. K. Ralph, Accts. Chem. Res. A,
107 (1971).

A. Weller, Progress in Reaction Kinetics 1, 189 (1961).

M. R. Loken, J. W. Hayes, J. R. Gohlke, and L. Brand,
Biochem. J., 11, A779 (1972).

G. Weber and K. Rosentieck, Biochem. Symp. 1, 333
(196A).

G. Weber, D. J. R. Laurence, Proc. Biochem. J., 56,
(195A).

M. Deluca, L. Brand, T. A. Cebula, H. H. Seliger, and
A. F. Makula, Time—resolved Proton—Transfer, 2A6,

6702 (1971).

P. A. Kollman, and L. C. Allen, Chem. Rev., 72, 283
(1972).

J. E. DelBene, J. Chem. Phys., §§, 3139 (1973).

C. A. Coulson, and V. Danielson, Arkiv, Kysik. 8, 205
(1955).

20A

 

205

17. S. Bratoz, Symp. Forces Intermoleculaires, Bordeaux,

1955.
18. G. M. Barrow, J. Am. Soc., 8, 5802 (1956).

19. H. Baba, A. Matsuyama, and H. Kokubum, J. Chem. Phys.,
.51. 895 (196A).

20. M. Arrio—DuPont, Biochem. Biophys. Research Communica—

tions 33, 653 (1971).

21. G. L. Bell and G. M. Barrow, J. Chem. Phys., 31,
1158 (1959).

22. G. M. Barrow, Spectrochim. Acta 18, 799 (1960).
23. Kreuger, Can. J. Chem., 83, 201 (196A).

2A. G. E. Bacon, in D. Hadzi (editor), Hydrogen Bonding,
Pergamon London, 1959, p. 23.

25. A. Piekara, J. Chem. Phys., 6, 21A5 (1962).

——

26. A. R. Ubbelohde, and K. J. Gallagher, Acta Cryst., 8,
71 (1955).

27. Y. Imry, I. Pelah, and E. Wiener, J. Chem. Phys.,
5;. 2332 (1965).

28. S. Ganesan, Nature, 188, 757 (1960).
29. A. Grimson, J. Phys. Chem., 81, 962 (1963).

30. C. C. Costain, and G. P. Srivastava, J. Chem. Phys.,
31, 1620 (196A).

31. R. Blinc, and M. Pintar, J. Chem. Phys., 38, llAO
(1961).

32. J. L. Wood, J. Mol. Structure, 18, 1A1 (1972).

33. R. Blinc, D. Hadzi, and A. Novak, Z. Elektrochem., 8A
567 (1960).

3A. R. Blinc, and D. Hadzi, Spectrochim. Acta, 18, 852
(1960).

35. D. Hadzi, Pure and Appl. Chem., 11, A35 (1965).

36. E. Clementi, J. Mehl, and W. von Niesen, J. Chem. Phys.,
53. 508 (1971).

 

 

 

37.
38.

39.

A0.
A1.
A2.
A3.
AA.
A5.
A6.
A7.
A8.
A9.
50.
51.
52.
53.
5A.
55.

56.

58.
59.

60.

206

P. 0. wadin, Rev. Mod. Phys., 5, 72A (1963).

M. Kasha, Disc. Faraday Soc., 8, 1A (1950).

N. S. Bayliss and E. G. McRae, J. Phys. Chem., 88,
1002 (195A).

G. C. Pimentel, J. Am. Chem. Soc., 12, 3323 (1957).
K. Weber, Z. Phys. Chem., 818, 18 (1931).

A. Weller, z. Physik. Chem., 3, 238 (1955).

A. Weller, Z. Elektrochem., 88, 662 (1952).

. Weller, Z. Physik. Chem., 18, A38 (1958).

. Férster, Naturwiss. 88, 186 (19A9).

. F6rster, z. Elektrochem. 83, A2 (1950).

. Férster, Z. Elektrochem., 88, 531 (1950).

Z. Lippert, Naturforsch, 188, 5A1 (1955).

Z. Lippert, Elektrochem., 81, 962 (1957).

:Dtill'flt—Eit-Ht-Jib

. Weller and W. Urban, Angew Chem. 88, 336 (195A).
Kokubun, Naturwiss, 33: 233 (1957).

Weller, Z. Elektrochem., 88, 8A9 (195A).

. Weller, Z. Elektrochem., 81, 956 (1957).

Weller, Z. Physik. Chem., 18, 163 (1958).

Q3>3>3>III

. Jackson, and G. Porter, Proc. Roy. Soc., (London),
A2 , 13 (1961).

E. Vander Donckt, Progr. Reaction Kinetics, 5, 273
(1970). I

J. F. Ireland, and P. A. H. Wyatt, Adv. Phys. Org.
Chem., 13, 131 (1976).

C. Parker, "Luminescence of Solution", 1970.

Albert and E. P. Sargent, "Ionization Constants of
Acids and Bases", Methuen, London, 1962.

S. G. Schulman, P. T. Tidwell, J. J. Cetroelli, and
J. D. Winefordner, J. Am. Chem. Soc., 88, 3179 (1971).

 

 

 

 

 

 

 

 

 

 

61.

62.

63.

6A.

65.

66.

67.

68.

69.

70.

71.

72.
73.

7A.

75.
76.
77.
78.

207

S. G. Schulman, and 1. Pace, J. Phys. Chem., 6,
1966 (1972). "—

T. C. Werner, and D. M. Hercules, T. Phys. Chem.,
15. 1030 (1970).

E. L. Wehry, and L. B. Rogers, Spectrochim. Acta, 81,
1976 (1965).

E. Lippert, W. Lfider and F. M611, Spectrochim. Acta,
.15. 858 (1959).

E. Vander Donckt, and G. Porter, Trans. Faraday Soc.,
83, 3215 (1968).

H. H. Jaffe and H. Lloyd Jones, J. Org. Chem., 88,
96A (1956).

S. Udenfriend, "Fluorescence Assay in Biology and
Medicine", S. Schulman, and H. Gershon, J. Phys. Chem.,

13. 3297 (1968).

 

C. A. Parker, "Photoluminescence of Solutions", Else-
vier, London (1968).

 

R. Argauer, and C. E. White, "Fluorescence Analysis",
Dekker, New York (1970), Academic Press, NY (1962).

Li. L. Yang, M. S. Thesis, Michigan State University,
197A.

E. Vander Donckt, Bull. Soc. Chim. Belges, Z8, 69
(1969).

G. Yagil, J. Phys. Chem., 11, 103A (1967).

J. W. Bridges, and R. T. Williams, Biochem. J.
107, 225 (1968).

D. D. Rosebrook, and W. W. Brandt, J. Phys. Chem.,
70. 3857 (1966).

N. Mataga and Y. Kaifu, Mol. Phys., 1, 137 (1963).
A. Weller, Z. Elektrochem., 88, llAA (1956).
K. Hirota, z. Phys. Chem. N. F., 38, 222 (1962).

D. L. Williams and A. Weller, J. Phys. Chem., 15,
AA73 (1970).

 

 

 

 

79.

80.

81.
82.

83.
8A.

85.

86.

87.

88.

89.

90.

91.

92.

93.

9A.
95.

96.

208

K. C. Ingham and M. A. El-Bayoumi, J. Amer. Chem. Soc.
88, 167A (197A).

M. A. El-Bayoumi and P. Avouris and W. R. Ware, J.
Chem. Phys., 23: 2499 (1975).

W. Klépffer, Adv. Photochem., 18, 311 (1977).

W. Bartok, P. J. Lucchesi, and N. U. Snider, J. Am.
Chem. Soc., 88, 18A2 (1962).

A. Weller, Naturwissenshaften, 88, 175 (1955).

G. M. Gantz, and W. G. Summer, Textile Rev. J.,
31. 299 (1957).

P. J. Kovi, C. L. Miller, and S. G. Schulman, Anal.
Chim. Acta, 81, 7 (1972).

W. Klépffer and G. Naundorf, J. Luminescence, 8, A57
(1979). '

IU. V. Naboikin, E. N. Pavolva, and B. A. Zaboro-
zhnyi, Opt. Spectrosc. (Engl. Transl.), 6, 231 (1959).

G. J. Woolfe and P. J. Thistlethwaite, J. Am. Chem.
Soc., 102, 6917 (1980).

P. J. Thistlethwaite and G. J. Woolfe, J. Chem. Phys.
Lett., 93. 401 (1979).

D. Ford, P. J. Thistlethwaite, and G. J. Woolfe, J.
Chem. Phys. Lett., 88, 2A6 (1980).

K.BZ. Ismail, Ph.D. Thesis, Michigan State University
19 2.

H. Beens, K. H. Grellman, M. Gurr and A. H. Weller,
Disc. Faraday Soc. 88, 183 (1965).

E. M. Kosower and Hanna Dodiuk, J. Luminescence, 11,

2“9 {1975/76).
K. Sandros, Acta. Chem. Scand., 188, 761 (1976).

T. Yasunaga, N. Tatsumoto, H. Inoue and M. Miura, J.
Phys. Chem., 18, A77 (1969).

K. K. Smith and K. J. Kaufmann, J. Phys. Chem., 83,
2286 (1978).

97.

98.

99.

100.

101.

102.

103.

10A.

105.

106.

107.

108.

 

 

109.

110.

111.

112.

113.

H. Shizuka, K. Matsui, Y. Hirata, and 1. Tanaka, J.
Phys. Chem., 88, 2070 (1976).

H. Shizuka, K. Matsui, Y. Hirata, and 1. Tanaka, J.
Phys. Chem., 81, 22A3 (1977).

N. Mori, Y. Asano, T. Irie and Y. Tsuzuki, Bull. Chem.
Soc., Japan, 88, A82 (1969).

J. Catalan and J. I. Fernandez—Alonso, Chem. Phys.
Lett . 19. 37 (1973).

E. Canadell, J. Catalan and J. I. Fernandez—Alonso,
Advance in Molecular Relaxation and Interaction Pro—
cesses, 1g, 265 (1978).

S. G. Schulman, P. J. Kovi and J. F. Young, J. Pharm.,
Sci . fig. 1197 (1973).

Tu. V. Naboikin, and B. A. Zadorozhnyi, Opt. Spec-
trosc. (Engl. Transl.), 8, 3A7 (1960).

W. R. Ware, P. R. Shukla, P. J. Sullivan, and R. V.
Bremphis, J. Chem. Phys., 88, AOA8 (1971).

Iu. V. Naboikin, B. A. Zadorozhnyi, and E. N. Pavlova,
Opt. Spectrosc. (Engl. Transl.), 6, 312 (1959).

M. Oki, M. Hirota, and S. Hirofuji, Spectrochim.,
Acta. 88, 1537 (1966).

E. D. Bergmann, Y. Hirshberg, and S. Pinchas, J.
Chem. Soc., 2351 (1950).

G. J. Woolfe, and P. J. Thistlethwaite, J. Am. Chem.
Soc., 103, 38A9 (1981).

P. T. McTigue, and P. T. Renowden, J. Chem. Soc.,
Faraday Trans. 1., 11, 178A (1975).

P. J. Kovi, and S. G. Schulman, Anal. Chem., 88,
989 (1973)-

S. G. Schulman, P. J. Kovi, Anal. Chim. Acta, 81,
259 (1973).

S. Nagakura, and M. Gouterman, J. Chem. Phys., 88,
881 (1957).

(a) E. Lippert, Z. Naturf., 188, 5A1 (1955).

(b) E. Lippert, z. Elektrochem., 81, 962 (1957).

 

 

 

113.

11A.

115.

117.

118.

119.

120.

121.

123.

12A.

126.

210

(c) N. Mataga, Y. Kaifu, and M. Koizumi, Bull. Chem.
Soc. Japan, 38, 690 (1955)-

(d) N Mataga, Y. Kaifu, and M. Koizumi, Bull. Chem.
Soc. Japan, 88, A65 (1956)~

(e) J. Czekalla, Z. Elektrochem., 88, 1221 (1960).

A. Weller, Z. Elektrochem. 88, 662 (1952).

(a) N. Mataga, Y. Kaifu, and M. Koizumi, Bull. Chem.
Soc. Japan, 33, 115 (1956).

(b) N. Mataga, Bull. Chem. Soc. Japan, 31, A81 (1958).

(a)

Th. Farster, Naturwissenschaften, 38, 108 (19A9).
(b) Th. Fbrster, Z. Elektrochem., 88, A2 (1950).

(0) Th. FCrster, Z. Elektrochem., 88, 531 (1950).
(d) A. Weller, Z. Elektrochem., 88, 55 (1960).

N. Mataga and Y. Kaifu, J. Chem. Phys., 38, 280A
(1962).

A. Matsuzaki, Nagakura, and K. Yoshihara, Bull.
Chem. Soc

. Japan, 81, 1152 (197A).

0. Stern, and M. Volmer, Phys.

2., 39, 183 (1919).

"Handbook of Chemistry and Physics“, ed by C. Hodg—
mon, Chemical Rubber, Cleveland, OH (1962), pp. 2258—
226A.

A. Weller, z. Phys. Chem. N. F., 11, 22A (1958).

M. V. Smoluchowski, Z Phys

. . Chem., (Leipzig), 93,
129 (1917)-

. R. Ware and J. S. Novros, J. Phys. Chem., 70
32A6 (1966)

C. A. Taylor, M. A. El—Bayoumi, and M. Kasha, Proc.
Nat. Acad. Sci. 0.8., 83, 253 (1969).

M. A El Bayoumi, Ph.D. Thesis, Florida State Uni—
versity, 1961.

K C Ingham, M. Abu—Elgheit and M. A. El—Bayoumi,
J. Amer. Chem. Soc., 83, 5023 (1971).

127.

128.

129.

130.

131.

132.

138.
139.
1A0.

1A1.

1A2.

1A3.

1AA.

211

H Bulska, and A. Chodkouska, J. Am. Chem. Soc., 102,
3259 (1980)-

K. B. Eisenthal,

K. Gnaedig, W. M. Hetherington, M
Crawford, and R. H. Micheels, Springer Ser Chem.
Phys. 8 (Picosecond Phenom.)

3A—9 (1978).
W M. Hetherington III, R. H. Micheels and K. B
Eisenthal, Chem. Phys. Lett. 66 230 (1979).

a_)
P Avouris, L. L. Yang, and M. A. El—Bayoumi, Photo-
chemistry and Photobiology, 38, 211 (1976).

C. A. Taylor, Masters Thesis, Florida State Uni—
versity, 1969.

W A. Van Hook in Isotope Effects in Chemical Re—
actions, eds., J. C. Collins and N. S. Bowman, ACS
Monograph 167, Van Nostrand Reinhold, New York (1971)
Earlier reviews are cited in this reference

C. G. Swain, E. C.

Stivers, J. F. Reuwer, Jr
J. Schaad, J. Amer.

, and L.
Chem. Soc.

. 99. 5885 (1958).
. Lewis and J. K. Robinson, J. Amer.

a SA337 (1968)

90

M. J. Stern and P. C. Vogel, J. Amer. Chem. Soc
33, A66A (1971); M.

J.

' 3
Stern and R. E. Weston, Jr
Chem. Phys., 88, 2815 (197A).

Chem. Soc.,

J.

Bigeleisen, Tritium in the Physical and Biological
Sciences, IAEA, Vienna, Vol. 1 (1962), p. 161.

R. P Bell The Proton in Chemistry, 2nd Ed., Chapman
and Hall (1973).

E. F. Caldin, Chem. Revs., 83, 135 (1969)

A. J. Kresge, Pure Appl. Chem., 8, 2A3 (196A).
V. Gold, Adv. Phys. Org. Chem., 1, 259 (1969).
P. M. Laughton, and R. E. Robertson, Can. J. Chem
35, 171A (1956).

R. E. Robertson, Can. J. Chem., 38, 613 (1957).

R. E Robertson, and P. M. Laughton, Can. J. Chem
35. 1319 (1957).

F D Ross1ni, J. W. Knowlton, and H. L. Johnston,
J. Res. Natl. Bur. Stand. 2A, 369 (19A0).

3

1A5.
1A6.

1A7.

1A8.
1A9.

151.

152.

162.

163.

16A.

212

W. M. Jones, J. Chem. Phys., 88, 207 (1968).

. G. Swain, and R. F. W. Bader, Tetrahedron, 10,
182 (1960)
C. G.

Swain, A. D. Ketley, and R. F. W. Bader, J
Am. Chem. Soc., 18, 2353 (1959).

F. Hund, z. Physik, 53, 805 (1927).

E. P. Wigner, Z. Phys. Chem. (Leipzig), B19, 203
(1932).

D. M. Dennison, and G. E. Uhlenbeck, Phys. Rev., 81,
313 (1932)-

E. Grunwald, in "Progress in Physical Organic Chem—
istry".

Cohen, Streitwieser, and Taft, eds. Wiley, NY 1965
Vol. 3.

H. S. Johnston, Adv.

Chem. Phys., 3, 131 (1961).

E. F. Caldin, Chem. Rev., 88

. 135 (1969).
C. Haas, and D. F. Horning, J. Chem. Phys.
(1960).

G.

. 32. 1763

M. Barrow, Spectrochim. Acta.

19. 799 (1960).
D. Hadzi, J. Chem. Phys., 33, 1AA5 (1961).

P. 0. wadin, Adv. Quantum Chem.

9 a. 213 (1965)-
J. Soc., 188, A18 (192A).

W Klopffer and G. Kaufmann, J. Luminescence, 20,
283 (1979).

K. Marsh, J. Chem.

B. A Zadorozhnyi and I. K. Ishchenko, Opt Spectry
(English. Transl.) 18, 306 (1965).

H. Kokubun, Z. Physik. Chem., 13, 386 (1957); Natur—
wissenschaften AA,

233 (1957);_Z. Elektrochem. 88,
599 (1958)

A N Terenin and A. V. Karyakin, Doklady Akad
Nauk SSSR 8, A25 (19A7); Nature 159, 115 (19A7).

A. V. Sheblya and A N. Terenin, Optika. Spektr. 18,
617 (1961).

165.

166.

167.
168.
169.

170.

171.

172.

173.

17A.

175.

176.

177.

178.

179.

180.

181.

182.

213

Yu. N. Sheinker and 1. Ya. Partovskii, Zhur. Fiz.

Khim. 33. 39A (1958).

F. Bertinotti, G. Giacomello and A. M. Liquori,
Acta Cryst. Z, 808 (195A).

R. S. Mulliken, J. Am. Chem. Soc., :8, 811 (1952).
S. P. McGlynn, Chem. Rev., 88, 1113 (1958).

G. Briegleb, Elektronen-Donator—Acceptor Komplexe
Springer—Verlag, Berlin-Gottingen—Heidelberg, 1961.

J. N. Murrel, Quart. Rev., 18, 191 (1961).

R. S. Mulliken and W. B. Person, Ann. Rev. Phys. Chem.
13, 107 (1962).

B. Ramsey, M. A. El—Bayoumi, and M. Kasha, J. Chem.
Phys., 38, 1502 (1961).

J. N. Murrell, The Theory of Electronic Spectra of
Organic Molecules, Wiley (1963).

L. Lannelli, P. Giordano Orsini and G. Daniele, R. C.
Accad. Napoli, 38, 1 (1953).

C. Chang, N. Sharbestury and M. A. El—Bayoumi, J.
Chem. Phys. Lett. Z8, 107 (1980).

X—ray structure was determined by D. L. Ward, Michigan
State University in 1978.

A. Bree and R. Zwarick, J. Chem. Phys., 88, 3355
(1968)-

K. R. Popov, L. V. Smirno, L. A. Nakhimovskaya and
V. L. Grebneva, Opt. i Spektroskopiya 31, 363 (1971).

H. U. Schuett and H. Zimmermann, Ber. Bunsenges.
Physik. Chem. 81, 5A (1963).

A. Marty and P. Viallet, J. Photochem., 1 (6), AA3
(1973).

C. Chang, Ph.D. Thesis, Michigan State University,
1978.

K. T. Wei and D. L. Ward, Acta Crystallographica,
B32, 2768 (1976).

183.

18A.
185.

186.
187.

188.

21A

G. G. Multan, P. Main and M. M. Woolfson, Acta
Crystallographica, A27, 368 (1971).

A. Zalkin, private communication (197A).

C. K. Johnson (1965). "ORTEP. A FORTRAN Thermal-
Ellipsoid Plot Program for Crystal Structure 11-
lustrations", Report 0RNL—379A, Oak Ridge National
Laboratory, Oak Ridge, Tennessee.

P. K. Kadaba, J. Pharm. Sci., 83, 1333 (197A).

L. Stephenson, and W. K. Warburton, J. Chem. Soc.
(C). 1355 (1970).

N. S. Potts, J. Chem. Phys., 28, 809 (1952).