THESIS

 

This is to certify that the

dissertation entitled

VITAMIN B6 MODEL COMPOUNDS AND THEIR

INTERACTIONS WITH METAL IONS
presented by

Kamal Zaki Ismail

has been accepted towards fulfillment
of the requirements for

Ph . D . degree in Chemistry

 

2y. War/{W

Majoi professo/ v
M. Ashraf El-Bayoumi

Date April 23, 1982

 

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VITAMIN B6 MODEL COMPOUNDS AND THEIR

INTERACTIONS WITH METAL IONS

By

Kamal Zaki Ismail

A DISSERTATION

Submitted to
Michigan State University
in partial fulfillment of the requirements
for the degree of

DOCTOR OF PHILOSOPHY
Department of Chemistry

1982

ABSTRACT

VITAMIN B6 MODEL COMPOUNDS AND THEIR

INTERACTIONS WITH METAL IONS

By

Kamal Zaki Ismail

Nearly all reactions concerned with amino acid trans-
formations are catalyzed by pyridoxal phosphate (PLP).
Such reactions involve racemization, transamination, de-
carboxylation, a,B-elimination, B, y-elimination and de-
aldolization. Pyridoxal phosphate, PLP, (vitamin B6) is
a substituted pyridine derivative with a hydroxyl group in
position 3- and an aldehyde group in position H-. The bio-
logically active form of pyridoxal phosphate, PLP, is its
Schiff base with an e-amino group of a lysine residue in the
active site of the enzyme. Salicylaldehyde derivatives
have been used as model compounds of PLP since they contain
its essential chromophore. We have studied both absorption
and emission properties of four model compounds namely:
salicylaldehyde, 2-methoxybenzaldehyde, 2-hydroxyacetophenone,
and N,6-dimethyl-2-hydroxyacetophenone and their corresponding
Schiff bases with an aliphatic amine, n-butylamine, and an

amino acid, DL-valine,in different media and at both room

Kamal Zaki Ismail

temperature and 77°K. Intermolecular and intramolecular
excited-state proton-transfer have been observed. The
steric hindrance in u,6-dimethyl-2-hydroxyacetophenone
model compound and its Schiff base clearly manifested
itself in their spectral properties.

The presence of monovalent cation is very essential
for activity in some PLP-enzymatic reactions. In trypto-
phanase both NH”+ and K+ are good activators whereas, Na+
and Li+ are not. Different techniques have been used ex-
tensively to study monovalent cation interactions with dif-
ferent enzymes. We have studied both absorption and emis-
sion spectra of some monovalent cation (e.g., Li+, Na+,

K+ and Tl+) salts of salicylaldehyde and three other para-
substituted phenols in THF and 111 other solvents. Stronger
cation-anion interaction (tighter ion-pair formation) was
observed in Salicylaldehyde salts. Even in polar solvents,
e.g., ethanol and DMSO an equilibrium between ion pairs and
free ions was observed.

Because of their intense and narrow emission, lanthanide
ions, Ln(III), have been extensively used to probe the
microenvironment at the active site in proteins. The
number of coordinated water molecules in the first shell
of Eu(III) are determined by following the fluorescence
decay of the metal ion in HZO/DZO mixtures. Energy transfer

2+ 2+

in Ln(III) bound protein, substituting Ca or Mg , is used

to measure the distances between different sites in proteins.

Kamal Zaki Ismail

To explore the potential of use of europium (III) in prob-
ing PLP-enzymes active sites we have prepared the europium
chelate of salicylidenevalinate Schiff base and studied its
magnetic and optical properties. The complex seems to be
in an octahedral configuration. Energy transfer from ligand
to metal was observed and an energy level diagram revealing
both absorption and emission properties of the complex was

constructed.

To my Mothejt,
My Wondenfiul 011.69.,

and my Son

ii

ACKNOWLEDGMENT

I would like to express my deep appreciation to my
advisor, Dr. M. Ashraf El-Bayoumi for his guidance, en-
couragement, and friendship during the course of this
study. I also thank the members of my committee, Dr.

Carl Brubaker, Dr. Thomas Pinnavaia, and Dr. Andrew Timnick.

My thanks also go to Mr. Ernest Oliver for the mass
spectrosc0pic measurements, Ms. Jo Kotarski for patiently
drawing all the figures, and Ms. Peri-Anne Warstler for
doing an excellent Job in typing this Dissertation.

I thank my former colleagues, Dr. Khader Al-Hassan
and Ms. Nahid Shabestary, for their friendship and support.

My thanks are extended to my beloved country, Egypt,
and to Alexandria University. I also extend my gratitude
to all my friends in the Egyptian Club in East Lansing for
making me feel at home with their friendship and brotherhood.

Last, but not least, I would like to thank my wife,
Eman, for her continuous encouragement and endless patience
throughout this work. To my son, Kareem, I give my ever-
lasting gratitude for waiting one entire week after my oral
defense to be born, on Friday, April 30, 1982, thus allowing

me a recovery period.

iii

Chapter

LIST OF
LIST OF
CHAPTER
CHAPTER

I.

II.

TABLES.

FIGURES

TABLE OF CONTENTS

I - INTRODUCTION.

II - EXPERIMENTAL

Systems Studied

A.

Model Compounds

a.
b.

Aldehydes .
Schiff bases.

Preparation and Purification of
Compounds . . . . . .

Monovalent-cation Salts

Potassium bis(salicylidene-DL—
valinato) Europium(III) Chelate

Materials . . . . . . .

A.

Purification of Solvents.

O\U'l 12'me

Ethanol . . . . . . .
3-Methylpentane (3MP)

Water . . . . . . . . . . .
Acetonitrile, P-Dioxane and
N,N-Dimethylformamide (DMF)
Dimethylsulfoxide (DMSO).
Tetrahydrofuran (THF)

Spectral Measurements

U'l-DUOI'UH

Absorption Spectra.
Emission Spectra. .
77°K Emission Spectra
Infrared Spectra.
Mass Spectra. .

iv

Page

vii

. viii

0rd «4 -4 «a l4

(I)

10
ll
11

ll
11
12

12
12
l2

l3

l3
l3
l3
I“
IN

Chapter

CHAPTER

II.

III.

CHAPTER

II.

III.

CHAPTER

III - OPTICAL SPECTRA OF SALICYLALDEHYDE ‘

MODEL COMPOUNDS . . .
Introduction. . . . . . . .

Room Temperature Absorption
Spectra. . . . . . . . . .

Emission Spectra and Excited—State
Proton—Transfer in Salicylaldehyde
Model Compounds . . . . . . . . .

IV - ABSORPTION AND EMISSION STUDY OF
THE ASSOCIATION OF MONOVALENT CAT-
IONS WITH PHENOXIDE AND SALICYL—
ALDEHYDE ANIONS. . . . . . . .

Introduction.

Monovalent Cation Interactions.

Spectroscopic Studies of Ion-Pair
Formation. . . . . . . . . . .

A. Ultraviolet-Visible .
B. Infrared and Raman.

C. Electron Spin Resonance
(ESR) . . . . . . .

D. Nuclear Magnetic Resonance (NMR).
E. Luminescence. . . . . . . . . . . .
Results and Discussion. . . . . . .

v - EUROPIUM (III) CHELATE WITH
SALICYLIDENE-VALINATE SCHIFF
BASE. . . . . . . . . . . .
Lanthanide Ions as Luminescence
Probes of the Structure of Biological
Macromolecules. . . . . . . .
A. Decay Lifetimes as a Measure

of the Number of Metal Co-
ordinated Water Molecules

Page
15
15

16

38

67
67
68

71
73
89

93
95
97
98
125

125

132

Chapter Page
B. Characterization of Individual -
Binding Sites . . . . . . . . . . . 136

C. Inter Metal Ion Energy Transfer
Distance Measurements . . . . . . . 138

II. Energy Transfer in Rare-Earth
Metal Chelates . . . . . . . . . . . . lHO

A. Paths of Energy Migration . . . . . 1H1
B. Rate of Energy Transfer . . . . . . 1&5
C. Mechanism of Energy Transfer. . . . 1H6

III. Europium Complex with Salicylidene-
DL-Valinate Schiff Base . . . . . . . 147

A. Introduction. . . . . . . . . . . . I“?
B. Results and Discussion. . . . . . . 151
CHAPTER VI - CONCLUSION AND FUTURE WORK . . . . . 160

REFERENCES. . . . . . . . . . . . . . . . . . . . 163

vi

Table

LIST OF TABLES

Solvent Effect on the Absorption
Spectra of Compounds I-IV .

Absorption Spectra of Schiff base
Model Compounds in Different Solvents
Emission Spectra of Compound I

in Different Media.

Emission Spectra of Compounds

(I-IV) in Ethanol and 3MP

Emission Spectra of Compounds

(I-IV) in Water . l . . . . . . .

Log K, AH, and TAS for the Formation
of 18-Crown-6 Complexes with Metal
Ions in H2O at 25°C

Rate Constants for Reaction a; 25°C in
H2O, u=0.3: M+ + 18-Crown-6 EdM-Crown+
Absorption Bands of Alkali Metal

Salts in Dimethylsulfoxide.

Room Temperature Absorption and Emission

Maxima of Para-Substituted Phenols in

THF . . .

vii

Page

23

33

“5

53

53

87

87

91

123

Figure

LIST OF FIGURES

Pyridoxal phosphate (vitamin B6)
related compounds

The effect of pH on the spectrum

of holotryptophanase.

Room temperature absorption spectra
of dilute solutions (5 x 10-5 M)

of compound I (-———), II (-..-.--),
III ( ------ ) and IV (-——-)

all in ethanol.

Room temperature absorption spectra
of dilute solutions (5 x 10'5 M) of

compound I ( ), II (------), III

 

(—-----) and IV (-—--) all in 3MP

Room temperature absorption spectra

of dilute alkaline solutions (5 x 10"5
M) of compound I (————), II (-------),
III (- ----- ) and IV (----) in ethanol
Room temperature absorption spectra

of dilute solutions (5 x 10-5 M) of
compound I-b (————), II-b ( ------ )

and IV-b (----) all in 3MP.

viii

Page

17

22

25

28

Figure

10

11

Room temperature absorption spec-
tra of dilute solutions (5 x 10-5 M)
of compound I-b in different sol-
vents: ethanol (————), dioxane
(----) and DMSO ( ----- )

Room temperature absorption spec-

tra of dilute solutions (5 x 10'5 M)

of compound II-b in ethanol ( ----- ),
acidic ethanol (---—), dioxane (....)
and 3MP (————).

Room temperature absorption
spectra of dilute solutions
(5 x 10"5 M) of compound I-v in

different solvents: ethanol

 

(----), dioxane ( ) and DMSO

(- ..... ). . . . . . . . .

Room temperature absorption

spectra of dilute solutions (5 x

10‘5 M) of compound I-b in: ethanol
(————), acidic ethanol (----), and
alkaline ethanol (- ----- ) . . . . . .
Room temperature absorption

spectra of dilute solutions (5 x

10-5 M) of compound I-v in

ethanol (---—), acidic ethanol

ix

Page

29

3O

32

3M

Figure

ll

l2

l3

1”

15

) and alkaline

 

solution (

ethanol solution (- ----- )

Room temperature absorption spectra

of dilute solutions (5 x 10"5 M)

of compound II-v in ethanol (----),

and in acidic ethanol solution

(-—-—0.

Room temperature absorption spectra

of dilute solutions (5 x 10"5 M)

of compound IV-b in: ethanol (----),

acidic ethanol solution (

alkaline ethanol solution ( ----- )

 

) and

Room temperature emission spec—

tra of a dilute solution (5 x 10-5 M)

of compound I in ethanol at dif-

ferent excitation wavelength

 

Aexc = 330 nm (

380 nm (----)

) and xex

Emission spectra of a dilute

5:

C

alkaline solution (5 x 10—5 M)

of compound I in ethanol, Xe
380, at room temperature (

and at 77°K (---—).

XC

 

),

Page

35

36

37

39

HO

Figure

l6

l7

l8

19

20

Room temperature emission spec-

tra of a dilute solution (5 x 10‘5 M)
of compound I in water at different
excitation wavelengths, xexc =

320 nm (
<---->.

Room temperature emission spec-

 

) and Aexc = 380 nm

tra of dilute solution (5 x 10‘5 M)
of compound I in 3MP xexc = 320 nm.
Room temperature emission spec-

tra of dilute solutions (5 x 10‘5 M)

of compound II, A 320 nm in

exc
3MP (------), ethanol (----)

).

Emission spectra of dilute solu-
u

 

and water (
tions (5 x 10- M) of compound

III in ethanol; neutral solution
(Aexc = 320 nm) at room tempera-

ture (---—---) and 77°K (----);

alkaline solution (Aexc = 380 nm)

 

at room temperature ( ) and

77°K (------) . . . .

Room temperature emission spectra
of dilute solutions (5 x 10"5 M)

of compound III, in aqueous solution

xi

Page

H2

an

M6

“7

Figure Page

Aexc = 320 nm (————), and in

alkaline aqueous solution, Aexc =

3&0 nm (-----). . . . . . . . . . . . . . H9
21 Room temperature emission spectra

of a dilute solution of compound

IV in 3MP at xexc = 320 nm. . . . . . . 51
22 Room temperature emission spectra

4

of dilute solutions (5 x 10- M)

of compound IV, Xe = 330 nm, in

xc

ethanol (—-—-) and in alkaline

ethanol (------). . . . . . . . . . . . 52
23 Absorbing and emitting species

in PLP-Schiff base systems. . . . . . . 56
2“ Room temperature emission spectra

of a dilute solution (5 x 10-5 M)

of compound I-b in ethanol at dif-

ferent excitation wavelengths:

Aexc = 320 nm (————), and Aexc =
380 nm (----) . . . . . . . . . . . . . 58
25 Emission spectra of dilute solu-

tions (5 x 10-5 M) of compound

I-b in ethanol at both room tem-
perature and 77°K: Alkaline solu-
tion Aexc = 380 nm at room tem-

perature (----) and at 77°K (-..-.._)

xii

Figure

26

27

28

29

acidic solution Aexc = 330-380 nm
at room temperature ( ) and at

77° (-'-°--). . . .

 

Room temperature emission spectra
of dilute solutions (5 x 10-5 M)
of compound I-v in ethanol: in
neutral solution Xe

(

in acidic solution Aexc = 365 nm

xc = 320 nm

 

) and Aexc = 400 nm (—..—..),

(----), and in alkaline solution
Aexc = 360 nm ( ----- ) . .
Room temperature emission spectra

of dilute solution (5 x 10-5 M) of

compound II-b in ethanol Aexc = 330

 

nm ( ) and in acidic ethanol

Aexc = 350 nm (----). . .

Room temperature emission spectra
of dilute solutions (5 x 10-5 M)
of compound II-v in dioxane

( ----- ) and in ethanol (————)

A c = 320 nm . . . . . . .

ex
Room temperature emission spectra

of dilute solutions (5 x lO'Ll M)
of compound IV-b in ethanol:
neutral solution, A = 320 nm

GXC

xiii

Page

59

61

62

6M

Figure

30

31

32

33

0-———), acidic solution Xe =

320 nm (----) and alkaline solu-

XC

tion Aexc = 360 nm (------)

Room temperature emission spectra
of dilute solutions (5 x 10-

of compound IV-v, Aex

C

u M)

= 320 nm:

in ethanol (-------), alkaline

ethanol (—-----), acidic ethanol

 

(

) and in dioxane (----).

Comparative effects of mono-

valent cations on the spectrum

of holotryptophanase at 21°.

Samples contained 2.0 mg of

enzyme per 1.0 ml of 0.02 M

gimidazole-HCI buffer (pH 8.0),

2 mM mercaptoethanol, and ether

0.1 M NaCl, 0.1 M KCl, or 0.1 M

imidazole-HCI (Im), as indicated.

Correlation between wave number

and the inverse of the cationic

radius for contact fluorenyl ion

pairs in THF at 25°C.

Absorption spectra of fluorenyl

salts in THF at -30°C for various

counterions

xiv

Page

65

66

7O

76

76

Figure

34

35

36

37

38

39

Page

Optical spectrum of contact and

separated ion pairs of fluorenyl-

lithium at 25°C in 3,“ dihydropyran,
.; 3-methyltetrahydrofuran,

-—--; 2,5 dihydrofuran, ; and

 

hexamethylphosphoramide, -----. . . . . 77
Optical absorption spectrum of
fluorenylsodium in THF at 25,

-30, and -50°C. . . . . . . . . . . . . 80
Temperature dependence of the
contact-solvent—separated ion-

pair equilibrium of 9-(2-hexyl)-
fluorenyllithium in 2,5-dimethyl-
tetrahydrofuran.t . . . . . . . . . . . 80
Pressure dependence of the con-
tact-solvent-separated ion-pair

equilibrium of fluorenylsodium

in THF at 25°C. . . . . . . . . . . . . 81
(I) Dibenzo-l8-crown-6. (II)
Dicylohexyl-lu-crown-u; (III)
monobenzo-lS-crown-S. . . . . . . . . . 83
Selectivity of l8-crown-6: log

K values for reaction of lB-crown-6

with metal cations in H 0 vs.

2
ratio of cation diameter to l8—crown-6

XV

Figure

40

Al

42

“3

an

cavity diameter. Value for

Ca2+ reported <o.5.

Optical absorption spectrum of
barium fluorenyl and its l:l
complex with dimethyldibenzo-l8-
crown-6 between 320 and “00 nm

in THF at 25°C. . . . . . .
Possible structure of the barium
fluorenyl-dimethyldibenzo-l8-
crown-6 complex in THF or pyridine.
Room temperature absorption spec-
tra of dilute solutions (5 x 10"5 M)
of Na+ salt of salicylaldehyde in
DMSO (----) and in DMSO in the

presence of 2 x 10"2 M NaBFu

(——-—). . . . . . . . . . . . . .
Room temperature absorption
spectra of dilute solutions (5 x
10"”5 M) of T1(I) salt of salicyl-
aldehyde in DMSO (----) and in
DMSO in the presence of 5 x 10‘3 M
TlNO3 (-——-).

Potential energy of ion-pair
formation as a function of the

interionic distance

xvi

Page

85

88

88

99

101

102

Figure

45

M6

“7

H8

49

Room temperature absorption
spectra of dilute solution of Na+
salt of salicylaldehyde: (u.8 x
10'5 M) in water (----),(U.5 x

10‘5 M) in ethanol (----—-) and
(u.5 x 10‘5 M) in THF (————).

Room temperature absorption
spectra of dilute solutions of
Tl(I) salt of salicylaldehyde:
(4.5 x 10"5 M) in water (----),
(5.2 x 10"5 M) in ethanol (-------)
and (u.5 x 10‘5 M) in THF (———)
Room temperature absorption spectra
of dilute solutions (4.5 x 10'5 M)
in THF of Ns'(----) and T1+(-——)
salts of salicylaldehyde. .

Room temperature emission spectra
of dilute solutions of Na salt

of salicylaldehyde (5 x lo"5 M)

in DMSO: A c = 320 nm (----),

8X

Xexc = 380 nm (

 

) and in the

2

presence of NaBFu (2 x 10' M),

Aexc = 320 (-°-°-')
Room temperature emission spectra

of dilute solutions (5 x 10.5 M)

xvii

Page

105

106

107

108

Figure

50

51

Page

in DMSO of: Tl(I) salt of salicyl-

aldehyde, Aexc = 320 nm (-——_)

and 420 nm (

 

); Tl(I) salt

of Salicylaldehyde in presence

3’ Aexc =

320 nm (-----). . . . . . . . . . . . . 110

of ~5 x lo"3 M TlNO

Room temperature emission spectra
of dilute solutions (5 x 10'5 M)
of Na+and TI’salts of salicyl-

+ -
aldehyde in THF. Na salt, Xexc —

320 and 390 nm (-——-); Nafsalt

in presence of N2 x 10"2 M

NaBFu, Aexc = 320 nm (----);
Tl(I) salt, Aexc = 320 and 390 nm
(----°--); and Tl(I) salt in

presence of =5 x 10'3 M TlNO3,
Aexc = 320 nm (-------) . . . . . . . . 112

Room temperature absorption spec-
tra of dilute solutions (5 x 10.5 M)
of alkali-metal salts of p-hydroxy-
benzaldehyde in THF. P-hydroxy-

benzaldehyde ( ); Li+salt (....);

 

NaIsalt (------) and K+salt

(----~--) . . . . . . . . . . . . . . . llu

xviii

Figure

52

53

54

Page

Room temperature emission
spectra of dilute solutions
(5.x 10‘‘5 M) of alkali-metal salts
of p-hydroxybenzaldehyde in THF.

+ _
Li salt, Aexc - 335 nm (--_-),

exc = 3145 nm (

KIsalt in presence of 10"3 M

K+salt, A

 

) and

1806, Aexc = 345 nm (—-----). . . . . . 116

Room temperature absorption

spectra of dilute solutions (1.8

x 10"L1 M) of alkali-metal salts

of p-methoxyphenol in THF. P-

methoxyphenol (————), Naisalt

(------) and K+salt (----). . . . . . . 118
Room temperature emission spec-

tra of dilute solutions (1.8 x

10‘“ M) of alkali-metal salts

of p-methoxyphenol in THF. P-

methoxyphenol, Aexc = 290 nm

. + =
(————), Na salt, Aexc 320 nm

(-°-°-); Na+salt in presence of

lo‘3 M 18c6, A c = 320 nm (....);

8X

K+salt, A x0 = 335 nm (-..-..-)

e
and K+salt in presence of 10"3 M

1806, Aexc = 335 nm (---—). . . . . . . 119

xix

Figure

55

56

57

Page

Room temperature absorption spec-
tra of dilute solutions (2.5 x

0-4 M) of alkali-metal salts

1
of p—methylphenol in THF. P—

methylphenol (————), Naisalt

(-.--_), K*sa1t (------) and K+

salt in presence of 10"3 M

18C6 (----) . . . . . . . . . . . . . . 121
Room temperature emission spec-

tra of dilute solutions (2.5 x

lo‘Ll M) of alkali-metal salts of

p—methylphenol in THF. P-methyl-

phenol, Aexc = 290 nm (——-—),

+ - -.—O-
Na salt, Aexc — 320 nm ( )
and Ktsalt, Aexc = 330 nm

(-------) . . . . . . . . . . . . . . .‘ 122
Electronic energy levels for

eurOpium(III) and terbium(III).

The two upward-pointing arrows

show the transitions which occur

upon laser excitation at the wave-

lengths indicated. The two down-
ward-pointing arrows label the

most intense emissive transitions

of the two ions. Radiationless

XX

Figure

58

Page

energy transfer competes with
the radiative processes through
coupling of the emissive states
to the O-H vibrational overtones
of coordinated H2O molecules.
(The energy level diagram of
Tb(III) has been displaced
slightly to position the highest
electronic acceptor state of the
ground 7F term to coincide with
the zero-point energies of the
vibrational overtone ladders.). . . . . 131
Plot of the observed recip-

rocal luminescence lifetimes

-l
Tobsd’ vs. the mole fraction of
H20, tzo, in DZO-HZO mixtures

of Eu(III) solutions. The ef-
fect of the chelating ligands NTA
and EDTA is to displace water
molecules from the first coor-
dination spheres of the aqua-
metal ions to yield the approxi-
mate values indicated along the

right-hand ordinate . . . . . . . . . 135

xxi

Figure

59

6O

61

Excitation spectral profiles of

7F0 + 5D0 transition ob-

tained during the course of the

the

titration of an aqueous EuCl3
solution with the sodium salt of
the dipicolinate anion (DPA).
Eu(III) to ligand ratios are
indicated to the left of each
trace

Schematic energy level diagram
for a rare earth chelate pos-
sessing low-lying 4f electronic
states:——r, radiative transi-
tions;vvq, radiationless transi-
tions . . . .

Energy level diagram for rare-
earth chelates: (----) lowest
triplet level for the complexes;
(

(—~r—) rare-earth ion resonance

 

) rare-earth ion level;

level in chelates. Excited singlet

states of the complexes have
much higher energies and are not

shown. Likewise only the lowest

triplet-state energy for a complex

is given.

xxii

Page

137

142

144

Figure Page

62 Diagram of Cu(II) complex with:

(A) salicylaldehyde-valine Schiff

base; (B) PLP-valine Schiff base. . . . 150
63 KEEu(C12Hl3NO3)2] . . . . . . . . . . . 153
64 Magnetic behavior of KEEu(C12Hl3N03)2]

complex . . . . . . . . . . . . . . . . 154
65 Absorption spectra of dilute solution

-5

(3 x 10 M) of K[Eu(C12Hl3NO3)2]

in ethanol (----) and DMSO (-———) . . . 156
66 Emission spectra of dilute solutions

(3 x 10"5 M) of KEEu(C12Hl3NO3)2] in

 

ethanol at room temperature (----),

ethanol at 77°C (— - -), and DMSO

at room temperature ( ). . . . . . . 157
67 Energy level diagram for KfEu-

(C12H13N03)2] . . . . . . . . . . . . . 159

xxiii

CHAPTER I

INTRODUCTION

Pyridoxal phosphate (PLP) is the most striking of
the known coenzymes in terms of the multiplicity of dif—
ferent enzymatic reactions that are dependent on its
presence. Nearly all these reactions are concerned with
transformations of amino acids. It is not surprising, then,
to learn that numerous nutritional studies in animals and
microorganisms have established that a deficiency of pyri-
doxal (pyridoxine), a biochemical precursor of PLP, results

(1) Vitamin

in many irregularities in protein metabolism.
B6 is a name presently used to refer to the family of sub-
stances of biochemical interest that are structurally re-
lated to pyridoxal phosphate, Figure (1).

Pyridoxal phosphate-dependent enzymes catalyze a strik-
ing variety of distinct reactions such as racemization,
transamination, decarboxylation, a,B-elimination, B,y-
elimination and dealdolization. The biologically active
form of PLP is itsmSchiffbase with an e—amino group of a
lysine residue in the active site of the enzyme. The pres-

ence of metal ion as an essential constituent of some

enzymes provides an obvious link between enzymatic reactions

' o

H‘c4p H\C49
II ’C o
"°-'.’-° H2 l\ H HO’CHZWOH
OH ’ I
N CH3 N CH3
Pyridoxal phosphate Pyridoxal

Ho’

Figure l.

(PLP)

,0H
CH2 cH’

2
CH OH C OH
2’ "x /’ \\
N

I
CH3 N CH3
Pyridoxol Pyridoxamine
Lyrflne

o “VIN

ll ,9 OH
Ho-I'D-o H2 l \

OH 51’

PLP Schiff Base 3

Pyridoxal phosphate (vitamin B6) related
compounds.

and coordination chemistry. A good example for the metal
ion requirement for activity is that of tryptophanase. All
tryptophanasessnlfar examined require NH: or K+, and the
concentration required for half-maximum activity is 2-6

+(2). Na+ and Li+ are essen-

mM for NH: and 8-20 mM for K
tially without activating effects in each case, but they
inhibit the E.coli enzyme when added together with NH:
and K+(3). At optimal concentrations, NH: permits a level
of activity at least equal to and usually slightly greater
than K+, Rb+ also activates but less effectively than
K+(u). Tl+ also replaces K+ for E.coli enzyme and has a
greater affinity (0.35 mM gives approximately half-maximum
activity). In the presence of K+, holoenzyme (tryptophan-
ase-PLP complex) undergoes a pH-dependent change (pK = 7.2)
from the protonated form with Amax at 420 nm to the de-
protonated form with Amax at 337 nm, Figure (2). Initially,
the first band was assigned to the inactive form and the

(5). Recent studiesf6’7)

second band to the active form
however, using rapid-scanning stopped-flow absorption study
of the interconversion between the 420 and 337 absorbing
species showed that the form absorbing at 337 nm is the in-
active form and that absorbing at 420 nm is the active form.
In order to understand the role of monovalent cations
and to possibly help in the interpretation of the kinetics

data, it became clear that careful spectral assignments of

different absorbing species (or conformers) as well as a

 

 

   

ABSORBANCE

 

 

 

ABSORBANCE

 

 

 

WAVE LENGTH (111»)

Figure 2. The effect of pH on the spectrum of holotrypto-
phanase. 5

study of the spectral effects of interaction with mono-
valent cations is important.

Our approach is to study the optical properties (ab—
sorption and emission) in different media at both room
temperature and 77°K of simple model compounds which maintain
the basic electronic structural features of PLP. Substitu-

tion at the 2- and 5-positions in PLP (-CH and -CH

(8)

2-OPO3H3)

even though

3
are expected to cause minimal spectral shifts

they are important for the PLP activity. Similarly the
pyridine nitrogen is an important structural feature for
the role of PLP as a cofactor but does not contribute sig-
nificantly to the spectral properties of PLP. Thus, the
study of salicylaldehyde derivatives and their corresponding
Schiff bases seem Justified. Since the phenolic group may
undergo a proton transfer either in the ground or excited
electronic states, the corresponding methoxy derivative

is also studied as a reference compound. Conformational
changes involving the ring-carbonyl group C-C bond seem to
play an important role in the enzymatic activity(9). Thus,
two compounds, 2-hydroxyacetophenone and 2-hydroxy-4,6-di-
methylacetophenone, were chosen to study the steric effects
which remove the coplanarity of the carbonyl group and the
benzene ring. The study of the absorption and emission
spectra of salicylaldehyde model compounds under different
conditions are discussed in Chapter 3. The spectra of the

Schiff bases of these model compounds with both n-butylamine

and potassium-DL-valinate were also studied, the results

are summarized and discussed in the latter part of Chapter 3.
To investigate the role of monovalent cations, we studied

the interaction of salicylaldehyde as well as some simple

para substituted phenols, e.g., p-methyl, p-methoxy, and p-

formyl phenols, with monovalent cations, such as Li+, Na+,

K+, and Tl+, in conditions favoring ion-pair formation.
Absbrption and emission properties of these ion pairs will
demonstrate the effect of the approach of a metal ion to the
organic moeity in both ground and excited states. The re-
sults of our ion-pair study is discussed in Chapter 4, this
is preceded by a summary of the use of several spectroscopic
techniques in studying ion-pair formation.

Lanthanide ions are very effectively used as probes in
biological systems(10) to report microenvironmental changes
at the active sites. Eu(III) in particular has been used

(11). Accordingly, we planned

extensively for this purpose
to prepare and study Eu(III) complex formation with potas-
sium salicylidene-DL-valinate Schiff base to explore its ef-
fectiveness as a probe to report the degree of hydrophobicity
or hydrophilicity at the active site of PLP-enzyme systems.
The results of our study for Eu(III) complex with salicyl-
aldehyde Schiff base is discussed in Chapter 5, preceded

by a short review on the use of lanthanide ions as probes in

biological systems. Finally, a conclusion and suggested

future experiments are given in Chapter 6.

I.

CHAPTER II

EXPERIMENTAL

Systems Studied

 

A. Model Compounds

 

a) Aldehydes

 

H\C40

O

2-hydroxybenzaldehyde
(salicylaldehyde)

I

(Hi

0196/0

0

methyl-Z-hydroxyphenyl ketone
(2-hydroxyacet0phenone)
III ‘

(”i

*L13453

l

2-methoxybenzaldehyde
(anisaldehyde)

II

(”13

methyl—2-hydroxy-4,6-
dimethyl phenyl ketone
(4,6-dimethyl-2-hydroxy-
acetophenone)
IV

b) Schiff bases

n-butylamine S.B. » potassium valinate S.B.
(it-If- (CH2): CH3 CH3\ I; +
CH— Cl-l-C . - K
'3;\ N / \‘

c’ CH334‘C4'!‘ 5

R3 OR'
EE‘ Ii (05“

Compound R R2 R R4 Compound

2 1 3
I-b H H H H I-v R
II-b CH3 H H H II-v 2
IV-b H CH3 CH3 CH3 IV-v

Preparation and Purification of Compounds

1. Compound I, II and III were obtained from Fisher
Scientific and were purified by vacuum distillation.

2. Compound IV was obtained from Aldrich Chemical and

_was purified by repeated crystallization from petroleum
ether 60-110°C.

3. Compounds I-b and II-b were prepared by refluxing 2
mmoles of the aldehyde with 2 mmoles of n-butylamine in
methanol for 30 minutes and the product was obtained and
purified by vacuum distillation. Compound I-b was obtained
as a yellow liquid and compound II-b as a colorless liquid.
Both compounds gave expected mass spectral results. The

purity was checked by using gas chromatography.

4. Compound IV—b was prepared similarly and was ob—
tained as a creamy white solid which was purified by re-
peated recrystallization from methanol and ether, m.p.
157-158°C.

5. CompoundsILAI,II-v and IV-v were prepared by the

method of Heinert and Martell.(l2)

Two mmoles of pure DL-
valine were added to a solution containing 2 moles carbonate-
free KOH in dry methanol. After the solid was dissolved
completely, the mixture was cooled to 0°C and 2 mmoles of the
aldehyde were added. The solution was stirred for 30 min-
utes between 0 and -10°C. The solid compounds were ob-
tained by evaporating the solvent at low temperatures. Then,
the compounds were recrystallized from methanol and ether.
Compound I-v was yellow; m.p. 150-152°C, compound II-v was
white; m.p. 260—262, and compound IV-v was pale yellow,

m.p. l98-200°C. Mass spectra of the prepared Schiff bases

confirmed their structure.

B. Monovalent-cation Salts

 

1. Preparation of Na+and Tl(I) Salts of Salicyl-

 

aldehyde - A modified version of the preparation method of
(13)

Maggio was used. Two mmoles of ethanolic Nafor Tl(I)
ethoxide and 2 mmoles salicylaldehyde were mixed under good
stirring conditions. On addition of chloroform and by

cooling to 0°C, a yellow solid was obtained in good yield.

The solid was recrystallized from boiling ethanol and

10

washed with ethanol-ether mixture and dried under argon.

o M+
L0.

ii. Preparation of Li: Na+ and K+ Salts of P-substituted
4

 

 

Phenols<1u> - To a dilute solution (10' to 10'5»M) of p-

 

methyl, p-methoxy, or p-formyl phenol in THF was added an
excess amount of the metal hydride. The solution was
stirred under dry argon, then,the spectra were measured for

a clear solution sample of this mixture.

 

 

r.
0' M"
.01

C. Potassium bis(salicylidene-DL-valinato) Europium(III)

Complex

To a clear solution of 2.2 mmoles of the Schiff base
was added 2.2 mmoles of carbonate—free KOH in methanol.

To this mixture was added 1 mmole of Eu(NO -5H20 solution

3’3
in methanol and the solution was refluxed for 30 minutes.
The complex was obtained by evaporating the solvent under

vacuum. On recrystallization of the crude complex from

11

methanol and ether, yellow crystals were isolated, de

composition temperature >250°C K[(C N03)2Eu]°2H O,

12Hl3 2
analysis, Calcd: C, 43.31; H, 4.54; N, 4.21; Eu, 22.83%.

Found: C, 43.17; H, 4.54; N, 4.36; Eu; 22.60%.

II. Materials

 

A. Purification of Solvents
l. Ethanol

Absolute ethanol was fractionally distilled through a
1 meter vacuum Jacket column, at a slow rate (about 5 drops
per minute). Portions of about 50 ml were collected, and
the absorption spectrum was taken in a 10 ml cell to check
for benzene. Distillation was continuediuuflj.the charac-
teristic benzene UV absorption was no longer apparent.

Ethanol was then distilled and used as needed.

2. 3-Methylpentane (3MP)

A modified version of the purification method of Potts(15)
was used. 3MP (Phillips Pure Grade) was shaken for 30
minutes with a 50:50 mixture of concentrated sulfuric acid
and concentrated nitric acid. It was then shaken 3 times
for 30 minutes each with concentrated sulfuric acid. The
solution was then treated with sodium carbonate solution

until CO2 production ceased. The 3MP was then washed

12

several times with water until the water remained clear
relative to the yellow color of the first wash. The 3MP
was stored over sodium ribbon in a flask overnight. It

was refluxed through a vacuum jacketed 1 meter column and
distilled for use as needed. Purity was checked by obtain-

ing the absorption spectrum.

3. Water

Was doubly distilled in this laboratory.

4. Acetonitrile, P-Dioxane and N,N-Dimethylforma-

mide (DMF)

 

Spectroquality solvents from Metheson Coleman and Bell

(MC/B) were used without further purification.

5. Dimethylsulfoxide (DMSO)

It was stored overnight over NaOH and distilled at
reduced pressure (m2-3 mm Hg, b.p. 50°) from NaOH pellets

and stored over molecular sieve 4A.

6. Tetrahydrofuran (THF)

Traces of peroxide were removed by refluxing a 0.5%
suspension of cuprous chloride for one-half hour, followed

by distillation before proceeding. Predried over KOH pellets

l3

and finally refluxed over and distilled from lithium aluminum

hydride as needed.

B. Spectral Measurements

1. Absorption Spectra

 

All reported absorption spectra were recorded by use of a

Cary 17 spectrophotometer.

2. Emission Spectra

 

Fluorescence spectra were recorded by use of a multi-
component system consisting of a 500 w xenon light source.
500 mm Bausch & Lomb excitation monochromator (which pro-
vides a narrow excitation band width), Spex 1700-11 emission
Monochromator and EMI Model 9558QA Phototube. Noise reduc-
tion and amplification of the PMT signal is achieved by
using a Princeton Applied Research Model HR-8-Lock In
Amplifier and an appropriate chopping apparatus. Some of
the room temperature emission spectra were recorded by using
an Aminco-Bowman spectrofluorometer. The phosphorescence
spectra were obtained with these instruments equipped with

a rotating can phosphorosc0pe.

3. 77°K Emission spectra

 

A quartz dewar with flat quartz excitation and emission

windows filled with liquid nitrogen, and a narrow glass

l4

tube for the sample were used.

4. Infrared Spectra

 

Infrared spectra were obtained by using a Perkin-Elmer
457 Grating Infrared Spectrophotometer. Calibration of
frequency reading was.made by polystyrene film. Samples

were examined as KBr disks.

5. Mass Spectra

 

Mass spectra were obtained with a Finnigan EI-CI gas

chromatograph-mass Spectrometer.

CHAPTER III

OPTICAL SPECTRA OF SALICYLALDEHYDE MODEL COMPOUNDS

I. Introduction

The purpose of this study is to investigate the ab-
sorption and emission properties of four salicylaldehyde
model compounds (I, II, III and IV) and their correspond-
ing Schiff bases with n-butylamine (I-b, II-b and IV—b)
and with potassium-DL-valinate (I-v, II-v and IV—v) under

various conditions. More specifically, we studied:

(1) The absorption and emission prOperties of pos-
sible cations and anions.

(ii) Solvent effects on the absorption and emission
spectra of the various possible species.

(iii) The possible occurrence of excited-state proton
transfer either intramolecularly (in hydrocarbon solvents,
e.g. 3MP) or intermolecularly (in hydrogen bonding solvents
e.g., water, alcohol,and dioxane); in the latter case sol-
vent molecules would be involved.

(iv) The steric effects on the spectra, since in com-
pound III and/or IV the carbonyl group may be forced out-

of-plane with respect to the benzene ring.

15

16

This study should provide relevant information regarding
the spectra of enzymes containing similar chromophores and
possible transient intermediates occurring as a result of
pH changes or during the course of enzymatic reactions.

To achieve these goals we have investigated the absorption
and emission spectra of the previously mentioned compounds
at room temperature as well as at liquid nitrogen tempera-
ture (77°K). The spectra were measured in water, dioxane,
N,N-dimethylformamide (DMF), dimethylsulfoxide (DMSO),
acetonitrile and 3-methy1pentane (3MP). In water and
ethanol the spectra were measured in both acidic and alka-
line solutions. The figures and tables shown in the sub-

sequent pages provide a summary of our results.

II. Room Temperature Absorption Spectra

 

In Figure (3), room temperature absorption Spectra of
dilute solutions (5 x 10-5 M) in ethanol are shown for
compound I-IV, the electronic absorption spectrum of com-
pound I, salicylaldehyde, exhibits two absorption bands in
the near ultraviolet, one at 323 nm with an absorptivity
(e = 4,000) and the second at 253 nm (e = 12,000). Con-
sidering salicyaldehyde as a disubstituted benzene, its
absorption spectrum may be interpreted, in the zeroth-
order approximation, as resulting from transitions to
either charge-transfer states or to locally-excited states.

1 l

The locally excited states of benzene are B2u’ Blu and

Absorbance

17

 

0.6 -

Q
A
I

Q
N
l

 

0.0

 

 

 

Figure 3.

l l l j
250 . 300 350 400

Wavelength (nm)

Room temperature absorption spectra of'dilute
solutions (5 x 10‘5 M) of compound I ( ),

II (-------), III (- ----- ) and IV (----) all

in ethanol.

 

18

1 (16)
E111 .

responding to the

In salicylaldehyde the absorption bands cor-
lB2u benzene transition and carbonyl
locally-excited n + n* transitions are probably buried
under the more intense absorption at 253 nm. Accordingly,
the first absorption band at 323 nm corresponds to an elec-
tronic transition to an essentially charge-transfer state,
where the phenolic group acts as an electron-donor and the
carbonyl group acts as an electron-acceptor. Since the

two highest filled orbitals in phenol do not differ greatly

1).

in energy (As is about 4000 cm' One may also expect

another charge transfer band at higher energies but in the
same region. Thus, the second absorption band at 253 nm
corresponds also to a transition to a charge-transfer state.

In the first order approximation, this charge-transfer state

lBlu locally-excited state of benzene. Such

assignment is consistent with electronic-energy calcula-

(l7) (18).

is mixed with the

tions as well as solvent and steric effects The
charge-transfer energies can be calculated using the equa-

tion(17),

where ID is the ionization potential of the donor, EA the

electron affinity of the acceptor,and C the coulombic

attraction energy between the transferred electron and the

hole it left behind.

19

(19) as a basis for

The use of localized-orbital model
a perturbation treatment of substituted benzene to cal-
culate the energy of charge-transfer states is very useful
because it is easily applied to polysubstituted benzenes,
it allows a correlation with the absorption bands of various
substituted benzenes and the model can account successfully
for steric effects. Although the characterization of ab-
sorption band as being charge-transfer (C.T.) or locally
excited (L.E.) bands is correct only in the zeroth-order
approximation, yet it is very useful in predicting changes
in basicities, acidities and dipole moment as a result of

(20) that steric effects forcing

excitation. It is known
the carbonyl acceptor group out of plane causes a decrease
in the intensity of the charge-transfer band. The inten-
sity of the charge-transfer band depends on the resonance
integral across the substituent-hydrocarbon bond which

is proportional to coso, ¢ being the angle of twist of the
substituent. The intensity of a charge-transfer band
diminishes gradually as o increases until it disappears
when 9 = 90°. For the 90° twisted substituted-benzene
molecule the spectrum approximated that of benzene together
with low lying absorption bands involving local excitation
of the substituent. The substituent in this case may
influence the benzene transitions only by inductive pertur-

bation. A charge-transfer state may change only very

slightly in energy as a result of twisting the substituent

20

involved in the charge-transfer transition in question.
Such small change is due to the possible variation of the
inductive parameter of the substituent with the angle

of twist. For example, the spectrum of 4-nitro-N,N-di-
methyl aniline approaches that of nitrobenzene when the
angle of twist of N,N-dimethylamino group approaches 90°.
N,N-dimethylamino group which has (+Ifl) effect due to its
H-lone pair electrons becomes less n—repelling as a result
of twist. The inductive parameter of the nitro group which
has (-Ifl) effect do not vary much with ¢~ The angle of
twist of the substituent may be calculated from the ex-
tinction coefficient of the charge-transfer band using the

(18,20)

prOportionality ea cos2¢. Steric effects show that

the following absorption bands of substituted-benzene
spectra are charge-transfer bands: the 322 nm band of p-
nitroaniline, the 239 nm of acetophenone, the 233 nm of
aniline and the 252 nm of nitrobenzene.

In ethanol two different intermolecular hydrogen
bondings may occur between the solvent and solute mole-
cules depending on whether the ethanolic oxygen atom or
proton is involved. In compound IV, the first absorp-
tion band is blue shifted (shifted to shorter wavelength)
and decreased in its intensity showing an absorption at
305 nm (e = 2,600). This blue shift as well as the reduc-
tion in absorptivity demonstrate the steric effects due
to the interaction between the acetyl methyl group and

the ring methyl group in position 6 in compound IV. If

21

the carbonyl group was completely out of plane with respect
to the benzene ring, one should expect compound IV to
absorb near 280 nm like methylated phenol (p-methylphenol
absorbs at 280 nm).

In 3MP in dilute solutions, only intramolecular hydrogen
bonding is expected between the phenolic and the carbonyl
oxygen through the phenolic proton. Compound II without
phenolic proton (methylated phenolic group) shows a blue
shifted absorption band at 307 nm, Figure (4), compared to
the rest of the compounds. The intramolecular hydrogen
bond in compound I, III and IV caused the red shift in
the first absorption band in 3MP. Contrary to its behavior
in ethanol, compound IV, shows an intense absorption band
at 335 nm reflecting the coplanarity of the carbonyl group
with the benzene ring in 3MP. This implies that the energy
gained through the formation of the hydrogen bond outweighs
the repulsion energy between the sterically hindered methyl
groups. The small red shift of the absorption band in
compound IV relative to that of compound III is due to the
methyl substitution in position 4 and 6.

In Table (I), the solvent effect on the first absorption
band of the four model compounds is summarized. The band
maximum for compound IV in dioxane lies at a much longer
wavelength compared to the band maximum in ethanol and in
water; one may conclude that in dioxane the carbonyl group

is relatively more coplanar compared to the situation

Absorbonce

22

 

0.6 -

0.0 -

 

0.4 -

 

 

 

 

Figure 4.

1 . 1 n I "
250 300 350 400

Wavelength (nm)

Room temperature absorption spectra of dilute
solutions (5 x 10-5 M) of compound I (———),

II (-~----), III (- ----- ) and IV (---) all in
3MP. .

23

Table l. Solvent Effect on the Absorption Spectra of
Compounds I-IV.

 

 

 

 

 

 

 

Compd. (I) Compd. (II)
SOlvent v-lo"3 *Av-lo‘3 v-lo‘3 *Av-10’3
-1 -1 -1 -l
A(nm) (cm ) (cm ) A(nm) (cm ) (cm )
3MP 328 30.5 0.0 307 32.06 0.0
Dioxane 323 30.9 +0.4 314 31.8 -0.8
Ethanol 323 30.9 +0.4 318 31.4 —l.2
Water 323 30.9 +0.4 322 31.1 -l.5
Compd. (III) Compd. (IV)
Solvent
v-lo‘3 *Av-lo‘3 v-lo'3 *Av-lo'3
A‘ -1 -l A -l -1
(nm) (em ) (cm ) (mm) (cm ) (cm )
3MP 325 30.8 .0.0 335 29.8 0.0—
Dioxane 321 31.2 +0.4 325 30.8 +1.0
Ethanol 324 30.9 +0.1 307 32.6 +2.8
Water 323 30.9 +0.1 3-5 32.8 +3.0

 

 

Av = vsolv. — v3MP

24

in protic solvents. Dioxane can form only an intermolecu-
lar hydrogen bond with the phenolic proton.

In an alkaline aqueous and ethanolic solution the ab-
sorption spectra of compound I, III and IV are shown in
Figure (5). The first absorption band of the anion is
red shifted compared to the corresponding first absorption
band in neutral solutions. In the case of compound II
no change was observed as expected. Due to a smaller ion—
ization potential energy of the phenolate compared with
phenol one should expect the anion to absorb at longer wave-
lengths. The increase in the intensity of this band in
the anion is a result of a relatively greater charge migration
in the charge-transfer state of the anion. The relative
intensities of this band in compounds I, III and IV reflects
the steric effect which is‘more dominant in the case of
compound IV.

In pyridoxal phosphate enzymes so far studied, the
carbonyl group of the coenzyme is combined with the c-
amino group of lysine residue to form a Schiff base<2l).
A number of PLP-enzymes have an absorption band at 410-
430 nm. In addition to this band, most PLP-enzymes have a
peak at 325-340 nm. Some enzymes display an absorption at

360—364 nm at high or at all pH values(21).

These absorp-
tion bands can be related to specific species in the acid-
base equilibria of the Schiff base of PLP. The 410-430 nm

band is assigned to the Schiff base with a protonated

25

 

 

 

 

I. I I I I
-."\, ..
’ l
‘. l ;
0.4 J- ‘ -
.09
Il
_ \' . _
8 It. ./ ‘\
s v. . -/
e \‘. x ="\ ,-’ ‘-
8 0'2 - 1‘ / .' '/ \ \ \- d
a \I \ " y} \ ‘
q ,‘I. \ ./ /./ \\ \
'- \\ \ . //./ . \ \ -
0°. «.1 / ‘ \ \
,__./ \ \ .
*\\. \ \.
' -. \\\'\. H, _
(DC) I I 1 F" I ‘I -' I
250 300 350 400 450
Wavelength (nm)
Figure 5. Room temperature absorption spectra of dilute

alkaline solutions (5 x 10-5 M) of compound
I (___), II (-..-..-), III ( ...... ) and IV
(—---) in ethanol.

26

azomethine nitrogen and phenolate group (ketoenamine
species), the 325-340 nm band to that of a phenol group
without a proton on the azomethine nitrogen (enolimine
species), and the 360-364 nm band to that with phenolate
group and unprotonated azomethine nitrogen (anionic Species).
The basis of this currently accepted assignment was provided

by Heinert and Martell<22)

who studied the absorption spectra
of the Schiff bases of 3-hydroxy-4—formylpyridine and re-

lated compounds.

i 5+ 5 3
NH H NH H N H N
‘c’ \c¢ ‘c’ \c”
O / O / OH / 0‘
I , «—» \ I \ I \ I
N N N N
(ketoenamine) (enolimine) (anion)
410-430 nm 325-340 nm 360-364 nm

It is important here to note that there were no profound
effects on the spectra of those model compounds when com-
pared with the spectra of PLP Schiff bases as a result

of the 2- and 55position substituents in PLP. A number
of arguments for this conclusion have been presented from
the spectral studies on solution equilibria of Schiff
bases of pyridoxal and related compounds and on PLP-

enzymes and from other physicochemical studies<21’23).

27

Figure (6) shows the room temperature absorption of
compound I-b, II-b and IV-b in 3MP. In this solvent, intra-
molecular hydrogen bond between the phenolic group and the
azomethine nitrogen gives rise to an absorption at 317 nm
in the case of Schiff base I-b. This absorption band can
be assigned to the neutral nonpolar form of the Schiff
base namely, enolimine form. Again if we treat these model
compound Schiff bases as substituted benzene derivatives
one should consider the smaller electron affinity of the
azomethine group when compared with the carbonyl group in
the model compounds. This explains the blue shift (1,060
cm'l) of the first absorption band of these Schiff bases
in 3MP compared with the corresponding model compounds in
the same solvent.

In polar solvents, Figure (7), a new absorption band is
develOped at 410 nm, the intensity of this band depends
upon the polarity of the solvent. In DMSO, for example,
the absorption band at 410 nm is assigned to an absorption

by the neutral bipolar form of the Schiff base, keto-
enamine form, this band does not exist in the absorption
spectrum of the methoxy compound II-b in DMSO, ethanol or
dioxane, Figure (8), because of the absence of the phenolic
proton and accordingly no ketoenamine formation. The ob-
servation that salicylaldehyde, compound I, does not absorb
at N410 nm in DMSO demonstrates the lack of intramolecular

proton transfer in the ground state in polar solvents,

28

 

Absorbance

 

 

 

 

 

250 300 350
Wavelength (nm)

Figure 6. Room temperature absorption spectra of dilute

solutions (5 x 10’5 M) of compound I-b ( ),

II-b (-°-°-°) and IV-b (----) all in 3MP.

 

29

 

Absorbance

 

 

 

 

()‘D" 1 1 I
300 350 400
Wavelength ( nm)
Figure 7. Room temperature absorption spectra of dilute

solutions (5 x 10'5 M) of compound I-b in dif-
ferent solvents, ethanol (-——), dioxane (---)
and DMSO (-----).

30

 

 

 

 

 

 

8
§ 0.2 - -
’5
U3
1:
'<I _ ._
L 31
0'0 300 350 400

Wavelength (nm)

Figure 8. Room temperature absorption spectra of dilute
solutions (5 x 10'5 M) of compound II-b in
ethanol (-----), acidic ethanol (--—-), dioxane
(....) and 3MP ( ).

 

31

i.e., the equilibrium between the tautomers in the ground
state is overwhelmingly shifted towards the enolimine
tautomer. This observation reflects the high basicity

of the azomethine nitrogen compared with the carbonyl
oxygen. The spectra of the Schiff bases with potassium
valinate in different solvents are shown in Figure (9).
Table (2) summarizes these results. In the case of
valinate Schiff bases the keto-enol equilibrium is shifted
more towards the more stable ketoenamine tautomer.

Figures (10) and (11) show the absorption spectra of
salicylaldehyde Schiff bases I-b and I-v respectively.
Each figure shows the spectrum of the neutral form as well
as that of the cation and anion. In acidic ethanolic solu-
tion a band at 350 nm appears, Figure (10), (11) and (12),
this band corresponds to the absorption of the cation form
of the Schiff base. The red shift in this band relative
to the absorption in pure ethanol is due to the increase
in the electron affinity of the azomethine nitrogen due to
protonation. This results in the lowering of the energy
of this charge-transfer absorption band. In the case of
compound IV-b, the absorption band of the cationic form
is weak and appears at 315 nm, Figure (13), the large
blue shift compared for example with the absorption of the
cationic form of II-b which appears at 350 nm is due to a

large steric effect.

32

 

 

 

 

 

0.4-
. m
C)
c -
I!
e
f»; 0.2 .
<1
0.0 1 . . T‘- ..
300 350 400 450

Wavelength (nm)

Figure 9. Room temperature absorption spectra of dilute
solutions (5 x 10'5 M) of compound I-v in dif-
ferent solvents, ethanol (----), dioxane (
and DMSO (- ----- ).

)

 

33

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ommm mafizom ho oppooam soathomn< .m manna

34

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.A ..... IV Hosanna
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35

 

 

 

 

 

 

 

0.6 b "
\
m0.4-1; -
Q, .
C:
C)
13 _
8
U)
8
0.2 ‘
(DE) I I I ll. -
300 350 400 450

Wavelength (nm)

Figure 11. Room temperature absorption spectra of dilute
solutions (5 x 10"5 M) of compound I-v in
ethanol (----), acidic ethanol solution (———)
and alkaline ethanol solution (------).

36

 

 

Absorbance

 

 

 

 

l l l l
250 300 350 400
' Wavelength (nm)

Figure 12. Room temperature absorption spectra of dilute
solutions (5 x 10-5 M) of compound II-v in

ethanol (----), and in acidic ethanol solu-
tion (

 

37

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.A v soapsfiom Hocmnuo ofipfiom .AIIIIV Hosanna ”ca nI>H ccsoasoo no
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38

III. Emission Spectra and Excited-State Proton Transfer

 

in Salicylaldehyde Model Compounds

 

Salicylaldehyde (Compound I) shows a medium as well as
an excitation wavelength dependent emission spectrum. In a
dilute ethanol solution excitation of the neutral molecule
at 330 nm gives two emission bands as shown in Figure (14).
The first band at 426 nm was assigned to an emission from
the neutral nonpolar form. The longer-wavelength emission
at 500 nm was assigned to an emission from a bipolar form
resulting from a double proton transfer involving solvent
molecules. One should point out that the phenolic and
carbonyl oxygens of the solute in ethanol are hydrogen
bonded intermolecularly to solvent molecules} Upon excita-
tion, charge transfer from the phenolic to the carbonyl
group results in an increaSe in the acidity of the phenolic

(24’25) and an increase in the basicity of the car-

proton
bonyl oxygen. Accordingly, in the excited state, the
phenolic group may loose a proton and the carbonyl oxygen
may pick up a proton giving rise to a bipolar species.
Excitation of the neutral solution at 380 nm,where the
phenolate ion absorbs,resulted in an emission at 490 nm.
Similar emission, Figure (15), was observed at 490 nm when
an alkaline ethanolic solution was excited at 380 nm.
These results suggest that a ground-state equilibrium

between the neutral and anion forms of compound I exists

in ethanol solutions. In a glass solution, at 77°K, the

39

 

Fluorescence Intensity

 

 

 

Figure 14.

400 450 500 550
Wavelength (nm)

Room temperature emission spectra of a dilute
M) of compound I in ethanol

solution (5 x 10'5
at different excitation wavelengths, A =

_ . exc
330 nm ( ) and Aexc - 380 nm (----).

 

40

 

Fluorescence Intensity

 

 

 

 

l I l J
400 450 500 550
Wavelength (nm)

Figure 15. Emission spectra of a dilute alkaline solution
(5 x 10-5 M) of compound I in ethanol, Aexc =
380, at room temperature ( ), and at

77°K (----).

 

41

emission of the anion is blue shifted to 434 nm. In such
a rigid medium, solvent relaxation does not occur during
the lifetime of the excited singlet state causing the
fluorescence maximum to be blue shifted compared to the

fluorescence maximum in a fluid polar medium.

0,5: H
I .1. ' 3.. i‘ A
_ ..- = C:O...H
°’°" 9"" <—"" O t
O-H: ’51 O' O-HD‘E'
“O 3
l H
H I H

(:‘PCV’
go

In an aqueous solution, the room temperature emission
spectrum shown in Figure (16) exhibits two bands upon ex-
citation at 330 nm, a shorter-wavelength emisSion band at
428 due to the neutral form and a longer-wavelength emis-
sion at 505 due to the bipolar form. Excitation of an
alkaline solution showed an emission at 500 nm and was
assigned to the anion form.

In 3MP, the phenolic proton and the carbonyl oxygen

are intramolecularly hydrogen bonded. Upon excitation,

42

 

Fluorescence Intensity

 

 

 

 

I
350

Figure 16.

l . l l l
400 450 500 550 600
Wavelength (nm)

Room temperature emission spectra of a dilute
solution (5 x 10"'5 M) of compound I in water
at different excitation wavelengths, Aexc =
320 nm ( ) and Aexc= 380 nm (----).

 

43

intramolecular excited—state proton transfer will occur
resulting in the formation of the bipolar form which emits
at 506 nm in this solvent, Figure (17). In Table (3), emis-
sion band maxima observed in different media are summarized
together with the corresponding emitting species.

In compound II, 2-methoxybenzaldehyde, the phenolic
proton is replaced by a methyl group and hence it cannot
undergo proton transfer. Figure (18) shows its room
temperature emission spectrum in 3MP, ethanol,and water.

In 3MP an emission band at 350 nm was observed, this cor-
responds to the neutral form. In ethanol, a single emission
band appears at 376 nm. This band is blue shifted com-
pared to the emission band of the neutral form of compound
I observed at 426 nm. Such a large shift is attributed to
the fact that in compound I the phenolic proton is hydrogen
bonded to ethanol and such hydrogen bond becomes stronger
in the excited state, in compound II this hydrogen bond is
absent. In water, the larger dipole moment of water com-
pared to that of ethanol led to further stabilization of
the excited state resulting in a red shifted emission band
in water at 440 nm. It is possible also that the 440 nm
band in H20 is due to a cation formed by an excited-state
proton transfer to the carbonyl oxygen.

'The emission spectra of ethanolic solutions of compound
III,2-hydroxyacet0phenone,at room temperature and 77°K

are shown in Figure (19). Intermolecular excited-state

44

 

Fluorescence Intensity

 

 

 

I I l l
450 500 550 600
Wavelength (nm)

Figure 17. Room temperature emission spectra of dilute

solution (5 x 10'5 M) of compound I in 3MP

Aexc = 320 nm.

45

Table 3. Emission Spectra of Compound I in Different Media.

 

 

 

Ground State Excited State Fluorescence
Medium Species Species max. (nm)
)5? le..éo
,.H’° C
H\ [p OH 426
Ethanol C IE,
°~H---°\ ,
H I-I\C,OH
GO.- 500
H 0 H 0
Ethanol/ \Cé ‘c"
NaOH -
(3 -
Cr 6°
+
«a»?
Mic/,0.)H 0"
3MP C)
506
H\c,0H
0 °
H/ O—H “\Céo

.0° H28
H\c,o 60H
H20 0‘ H... o’H +
\ H\ [’0
H C

O”
500

J:

 

 

46

 

Fluorescence Intensity

 

 

 

I.

 

Figure 18.

300

I l l l l
350 400 450 500 550
Wavelength (nm)

Room temperature emission spectra of dilute
solutions (5 x 10'5 M) of compound II, Aexc~=

320 nm in 3MP (---'-'), ethanol (----) and
water ( ).

 

Fluorescence Intensity

47

 

 

 

 

 

 

l I l l I
350 400 450 500 550 600

Figure 19.

Wavelength (nm)

Emission spectra of dilute solutions (5 x 10'“ M)
of compound III in ethanol; neutral solution

= 320 nm) at room temperature (-H - -)
andx $7

°K (--—-); alkaline solution (Aexc = 380 nm)
at room temperature (

 

) and 77°K (- °°°°° )

48

proton transfer in the neutral solution of compound III
in ethanol gives an emission band at 512 nm. The small
red shift of this emission band when compared with that of
compound I, which occurs at 500 nm, may be attributed to a
higher basicity of the carbonyl oxygen in compound III.
In an alkaline solution, the 478 nm band is assigned to the
anion emission. This band is blue shifted when compared
with the emission of the anion of compound I. This is
probably due to the lack of planarity of the carbonyl group
in the case of compound III anion. At 77°K, the emission
of the bipolar form occurs at 475 nm and that of the anion
at 414 nm, these blue shifts compared with the correspond-
ing emission maxima in ethanol at room temperature is due
to the freezing of solvent relaxation in rigid ethanol at
77°K. The emission of compound III in water is shown in
Figure (20), the bipolar form emits at 512 nm and the anion
emission occurs at 490 nm. The results in water are
similar to those observed in ethanol, small red shifts
in water are attributed to its larger dipole moment. In
3MP where intramolecular hydrogen bonding occurs, com-
pound III gives rise to an emission band at 500 nm cor-
responding to the bipolar form resulting from an intra-
molecular excited-state proton-transfer.

In compound IV, 4,6-dimethyl-2-hydroxyacetophenone,
the repulsion between the acetyl methyl group and the

substituted methyl group in position 6- forces the acetyl

49

 

Fluorescence Intensity

 

/'l \-
/' \.
/' \

'\

 

Figure 20.

l I L I
400 450 500 550 600

Wavelength (nm) .

Room temperature emission spectra of dilute
solutions (5 x 10'5 M) of compound III, in
aqueous solution, Aexc = 320 nm ( ), and
in alkaline aqueous solution, Aexc = 340 nm

<-----).

 

 

50

group to be out of plane with respect to the benzene ring.
These steric effects manifest themselves in the absorption
spectra as discussed before. The steric effect is expected
to be more pronounced in protic solvents where intermolecu-
lar hydrogen bonds with solvent molecules are found. In
nonpolar solvents like 3MP absorption spectra indicate

that the molecule assumes a nearly planar structure through
an intramolecular hydrogen bonding between the phenolic

and carbonyl groups. In 3MP compound IV exhibits two
emission bands, one at 503 nm corresponding to the bipolar
form and a shoulder at 440 nm corresponding to the neutral
form as shown in Figure (21). The appearance of the neutral
form emission is probably due to the existence of a ground-
state equilibrium between intramolecularly hydrogen bonded
species and species that are not hydrogen bonded, the latter
cannot undergo excited state proton transfer. Figure (22)
shows the room temperature emission spectra of compound IV
in ethanol, an emission band observed at 422 nm is attributed
to a non planar neutral molecule; no proton transfer occurs
in the excited state in neutral ethanol. Band maxima of

the emission spectra of compound I-IV in ethanol, water

and 3MP are summarized in Tables (4) and (5).

Pyridoxal phosphate dependent enzymes exhibit absorp-
tion bands around 415 nm and at 335 nm. The studies made
with simple model compounds have shown that the 415 nm band
is due to a Schiff base with structure [1], bipolar form,

which is stable in polar media (hydr0philic media). The

51

 

Fluorescence Intensity

 

 

 

Figure 21.

J, I ‘ l l 1
3150 400 450 500 550 600

Wavelength (nm)

Room temperature emission spectra of a dilute
solution of compound IV in 3MP at Aexc = 320 nm.

Fluorescence Intensity

52

 

 

/ \\ ./ \
/ \./ \
I .ll \
I I ‘ \
I . \
I ’ \ \
l I \ \
I] ’ \ \
I f \ \
I j \ .
I - \ . \
I .l .\ -
[I / \~ \\
\ \

 

 

l
4350 400 4 5 0 550 600

Wavelength (nm)

Figure 22. Room temperature emission spectra of dilute

solutions (5 x 10' M) of compound IV, A

C
-330 nm, in ethanol (-——-) and in alkalIne
ethanol (- ----- )..

53

Table 4. Emission Spectra of Compounds (I-IV) in Ethanol

 

 

 

and 3MP.
Ethanol Ethanol/NaOH 3MP
RT 77°K RT 77°K RT -, 77°K

Compound A(nm) A(nm) A(nm) A(nm) 1(nm) 4(nm) 1(nm)

 

neutral bipolar

I 426 500 431 490 434 506 506
I
II 376 450 350 450
III 512 475 478 414 500 485
IV 422 430 468 505 503

 

 

Table 5. Emission Spectra of Compounds (I—IV) in Water.

 

 

 

H20 H O/NaOH
Compound A(nm) 5(nm)
I 428 505 500
II 440 440
III 512 490

IV 437 422

 

 

54
335 nm band was assigned to structure [2], nonpolar form(26),
which is stable in nonpolar media (hydrophobic media). Ac—

cordingly, the ratio between the absorbance of these two

N" .,
\ O \ 0
l , I
IV
1 2

bands can be used as a measure of the polarity experienced
at the pyridoxal chromophore's site(27).
Solvent effects on the emission spectra of PLP-model
compounds should also be helpful in interpreting enzyme
emission data and predicting the microenvironment that the
pyridoxal moeity experiences at the active site of the
enzyme i.e., whether the site is a hydrophilic or hydro-
phobic in character. For example, glycogen phosphorylase
contains one PLP molecule per protein subunit, PLP is very
tightly attached to the protein through a Schiff base
bond with an e-amino group of a lysine residue of the

(28)

protein In the pH range of 5.0 to 9.5, at which the
enzyme is stable, two absorption bands due to PLP are
observed with peaks at 333 and 425 nm(29). These two ab-
sorption bands indicate that an equilibrium exists between
the enolimine and ketoenamine forms and the active site

is not hydrophilic. The enzyme shows two emission bands:

one occurs at 335 nm and corresponds to the emission of

55

tryptophane residues in the protein; the other band at

535 nm corresponds to the emission of the ketoenamine

form of PLP<30’31). In the excited state the keto-enol
equilibrium is completely shifted towards the ketoenamine
form. For model systems, n-butylamine, valine, or n-
hexylamine-PLP in various solvents, the observed properties
of Schiff bases are very different. The Schiff base of
PLP-n-butylamine in aqueous solution and the Schiff base
PLP-valine in DMF exhibit two emission bands: one oc-
curs at 510 nm from the ketoenamine form absorbing at

410 nm; the second emission at 430 nm occurs from the enol-
imine form which absorbs at 325 or 333 nm(32). In polar
solvents, excitation of the enolimine form at 325 nm re-
sulted in an emission from the ketoenamine form indicat-
ing that proton transfer occurred in the excited state.

In hydrogen-accepting solvents like dioxane, the two
emissions were observed from the n-butylamine Schiff base.
These results can be summarized in Figure (23). In non
polar solvents the ground state equilibrium is shifted
towards form [1], whereas, in the excited state it is
shifted towards form [2]*(32).

In contrast to the nitrogen hetero-atom, the phenolic
group plays a major role in the spectral behavior of
Schiff bases of aromatic ortho-hydroxy aldehydes. Trans-
fer of the phenolic proton causes tautomerism in the

ground state and phototautomerism in the excited state.

56

 

 

 

 

 

R R
I ' .. I a.
c¢ "y \c// '5’.
\ 0 4VI 1, \ o
l , 330nm l ,
N L N . w
[I] enolimine ' [|]*
(330nm) (430nm)
-x-
llK ll K
I ' I ‘*
+ +
\ 0' I‘ll/2 I \ 0"
I"! 4IO nm N/
[2] ketoenamine [2]*
(420nm) . (5lOnm)

Figure 23. Absorbing and emitting species in PLP- Schiff
base systems.

57

Blocking of the hydroxyl group thus dramatically modi-
fies the absorption and emission properties characteristic
of these Schiff bases(33). Thus Schiff bases of o—
methoxybenzaldehyde is expected to show little resem-
blance to the optical and luminescence properties of
the Schiff bases of the corresponding o-hydroxybenz-
aldehyde. The spectral properties of PLP-model com-
pounds depend on whether the phenolic group forms a hydro—
gen bond with the azomethine nitrogen of the Schiff
base or with solvent molecules. These considerations
lead to our choice of the model compounds that we have
studied.

Room temperature emission spectrum of compound I-b
in ethanol is shown in Figure (24). Two emissions were
observed when the neutral solution was excited at 320 nm
where the enolimine form absorbs. The shorter-wavelength
emission with a peak at 425 nm corresponds to an emission
from the enolimine form and the longer-wavelength emission
at 500 nm corresponds to an emission from a ketoenamine
form. Excitation at 420 nm where the ketoenamine form
absorbs gave only the longer-wavelength emission at 500
nm. In acidic ethanolic solution, Figure (25), an emis-
sion from the cationic form was observed at 472 nm. An
alkaline solution of compound I-b in ethanol gave an emis-
sion at 510 nm when excited at 380 nm where the anion

absorbs, this emission belongs to the anion form.

58

 

 

 

Fluorescence Intensity

 

 

I II I I

400 450 500 55
Wavelength (nm)

 

Figure 24. Room temperature emission spectra of a dilute
solution (5 x 10'5 M) of compound I-b in
ethanol at different excitation wavelengths:
Aexc = 320 nm ( ), and Aexc = 380 nm (----).

 

Fluorescence Intensity

59

 

 

 

 

 

I I I I I l l
350 400 450 500 550 600
Wavelength ( nm)

Figure 25. Emission spectra of dilute solutions (5 x 10'5 M)

of compound I-b in ethanol at both room tem-
perature and 77°K: alkaline solution Aexc =

380 nm at room temperature (----) and at 77°K
(-H - --), acidic solution Ae xc = 330- 380 nm at
room temperature ( ) and at 77° (-----).

 

60

Emission in both acidic and alkaline solutions were blue
shifted to 415 and 408 nm respectively at 77°K because of
the lack of solvent relaxation. Room temperature emis-
sion spectra of compound I-v in ethanol are shown in
Figure (26). Excitation at 320 nm led to an emission
from the enolimine form at 450 nm, whereas, excitation
of the ketoenamine form at 400 nm gave an emission at
495 nm. In acidic solution, an emission at 480 nm was
observed as a result of the cation form excitation at
365 nm. In alkaline solution, excitation of the anion
form gave an emission peak at 500 nm.

In compounds II-b and II-v, there is no phenolic pro-
ton. The emission spectra of compound II-b in ethanol are
shown in Figure (27). Excitation of the neutral ethanolic
solution at 300 nm resulted in two emission peaks. The
first peak at 366 nm corresponds to the neutral form, and
the second at 450 nm was assigned to an emission from a
cationic form where the azomethine nitrogen accepts a
proton from the solvent. This assignment was confirmed
by studying the emission in acidic ethanol shown also in
Figure (27). These results indicate that the basicity of
the azomethine nitrogen is apparently greatly enhanced in
the excited state. Figure (28) shows the emission spectra
of a dilute solution of compound II-v in dioxane and
ethanol. Emission from the neutral form was observed in

dioxane at 352 nm. The 450 nm emission in neutral ethanol

61

 

O. ‘.
\\,
\ ~

I l l l l
400 450 500 550 600
Wavelength (nm)

 

Fluorescence Intensity

 

 

 

Figure 26. Room temperature emission spectra of dilute
solutions (51:10'5 M) of compound I-v in ethanol:

 

in neutral solution A = 320 nm ( ) and
Aexc = 400 nm (----°- , in acidic solution
Aexc = 365 nm (----), and in alkaline solu-

tion Aexc = 360 nm (-----)._

62

 

Fluorescence lntensity

 

 

/ l I l 1 1‘-
350 400 450 500 550
Wavelength (n m)

 

Figure 27. Room temperature emission spectra of dilute
solution (5 x 10‘5 M) of compound II-b in
ethanol Aexc = 300 nm ( ) and in acidic
ethanol Aexc = 350 nm (----).

 

63

 

 

 

 

Fluorescence lntensity

 

' 1 l I l I I
350 400 450 500 550 600
Wavelength (nm)

Figure 28. Room temperature emission spectra of dilute
solutions (5 x 10'5 M) of compound II-v in
dioxane ( ----- ) and in ethanol ( . ) Aexc
= 320 nm.

 

64

which results from excitation of the neutral form at 320 nm
can be interpreted as resulting from a cation form ob-
tained by proton transfer from solvent to the azomethine
nitrogen.

Room temperature emission spectra of compound IV-b
in neutral, acidic and alkaline solutions in ethanol are
displayed in Figure (29). Excitation of the neutral form
at 320 nm produced a broad emission with a peak at 415 nm,
this band is probably a composite band consisting of emis-
sion of the neutral form at 415 nm and an emission at
N460 nm due to the anion emission. In acidic solution,
the cation emits at 450 nm whereas, in alkaline solution
the anion emits at 465 nm. In the corresponding valine
Schiff base, compound IV-v, excitation of the neutral solu-
tion at 320 nm results, Figure (30), in a broad emission
which can also be considered as a combination of two
bands: the first around N420 nm which corresponds to the
neutral form emission and the second at N465 nm which cor-
responds to the anion form. In acidic solution an emission
peak was observed at 435 nm. In alkaline solution, the
emission of the anion form was observed at 463 nm. The
lack of excited-state proton transfer in the case of
compound IV and its Schiff bases IV-b and IV-v demonstrates
the lack of planarity in these molecules because of steric

effects.

Fluorescence lntensit y

65

 

 

 

 

 

l J l l
400 450 500 550
Wavelength (nm)

Figure 29. Room temperature Emission spectra of dilute

solutions (53:10' M) of compound IV-b in
ethanol: neutral solution, Aex = 320 nm (
acidic solution Ae xc = 320 nm (-
line solution Aexc = 360 nm (- - - ).

 

---) and alka-

66

 

Fluorescence lntensitL,

 

 

 

 

400 480 500 550
Wavelength (nm)

Figure 30. Room temperature emission spectra of dilute

solutions (5 x 10' M) of compound IV-v,
Aexc = 320 nm in:ethanol (-'-----), alkaline
ethanol (-. ----- ), acidic ethanol ( ) and
in dioxane (----).

 

CHAPTER IV

ABSORPTION AND EMISSION STUDY OF THE ASSOCIATION OF

MONOVALENT CATIONS WITH PHENOXIDE AND

SALICYLALDEHYDE ANIONS

I. Introduction

 

The presence of a metal atom as an essential consti-
tuent of some enzymes, and the metal ion requirement of
others for maximum activity, provide an obvious link between
enzymatic reactions and coordination chemistry. The pos-
sible functions of metal ions in enzyme substrate systems

(34)'

are easily visualized as follows:

(i) The model may form a complex with donor atoms of
either the enzyme or substrate and thereby enhance their
tendency towards reaction;

(ii) It may serve merely as a bridge through common
coordination to bring the enzyme and substrate into
proximity;

(iii) While serving function (ii) it may provide as
well a chemical activating influence; and

(iv) While coordinated to either the enzyme or substrate

it may appropriately orient groups undergoing reaction.

67

68

The most obvious property of a metal ion is its positive
charge which makes it effectively a Lewis acid, and it

will have a tendency to withdraw electrons from atoms and
groups to which it is attached. The catalytic efficiency

will depend on the effective charge of the metal ion.

Monovalent Cation Interactions

 

All tryptophanases so far examined require NHu+ or
K+, and the concentration required for half-maximum activity
(NHu+, 2—6 mM; K+, 8-20 mM) is similar for each enzyme(2).
Na+ and Li+ are essentially without activating effects in
each case, but they inhibit the E.coli enzyme when added
together with NH“+ or K+(3). At optimal concentrations,
NH“+ permits a level of activity at least equal to and
usually slightly greater than K+, Rb+ also activates but
less effectively than K+(u). T1+ also replaces K+ for the E.
coli enzyme and has a greater affinity (0.35 mM gives ap-
proximately half-maximum activity<35)).

Equilibrium dialysis studies showed that apotrypto-
phanase binds one mole of pyridoxal phosphate per 57,500

grams of protein<36>

, that is four pyridoxal phosphate
molecules per molecule of native tetrameric enzyme (mol-
ecular weight 220,000). K+ and Na+ are equally effective
in permitting the dissociation of tetrameric to dimeric
apotryptophanase at low temperature<37), and it is unlikely

therefore, that this effect serves as a basis for their

69

markedly different catalytic effects. In addition, distinct
spectral differences are observed in the presence of K+
versus Na+. Holotryptophanase shows characteristic pH-
dependent absorption maxima at 420 and 337 nm in the pres-
ence of the catalytically essential K+ or NH: ions. At equiv-
alent concentrations of Na+ or imidazole, only the 420 nm
band is formed, Figure (31), at pH 8.0(5). In the presence
of K+, at the same pH, the 337 nm band is formed. In the
enzymatically inactive form with Amax 420 nm, intramolecular
hydrogen-bond between the phenolic group and the azomethine
nitrogen was proposed, whereas, in the active form with

A x = 337 nm, the phenolic group lost its proton(5).

ma
(6,7)

Recent studies ,using a rapid-scanning stopped-flow
absorption technique, of the interconversion between the
420 and 337 nm absorbing species showed that the form
absorbing at 337 nm is the inactive form but acts as a
reservoir for the active form which absorbs at 420 nm.

The activity as well as the spectroscopic behavior in
the presence of K+ or Na+ ions demonstrate a difference in
the tertiary structure of tryptophanase and in the mode of
binding of pyridoxal-p in the two environments that is
consistent with a more compact structure and tighter bind-
ing of pyridoxal-p in the K+ environment. This difference
is sufficient to explain the requirement for K+ (or NHu+)

for catalysis, but does not elucidate the molecular basis

for this behavior. In this regard, we have studied the

70

 

 

 

 

 

3 , , '
pH a0 w I"
f
i
In .2 °:
0
Z
(
m
K
8
a:
‘.1
300 350 . 400 450

WAVELENGTH (MIL)

Figure 31. Comparative effects of monovalent cations on
the spectrum of holotryptophanase at 21°.
Samples contained 2.0 mg of enzyme per 1.0 m1
of 0.02 M imidazole-HCI buffer (pH 8.0), 2 mM
mercaptoethanol, and either 0.1 M NaCl, 0.1 M 5
KCl, or 0.1 M imidazole-HCl (Im), as indicated.

71

the monovalent cation interaction with some pyridoxal
phosphate model compounds. In media favoring ion-associa—
tion, e.g., tetrahydrofurane, we have studied both absorp-
tion and emission properties in order to examine the affinity
of ion-pair formation both in the ground and excited states.
Before discussing our results, a short review of the
spectroscopic methods of studying ion association is

given.

II. Spectroscopic Studies of Ion-Pair Formation

The concept of ion pairs was introduced in 1926 by
(38)

Bjerrum It was known in those days that ionophores-
compounds built up of ions and not neutral molecules — are
completely dissociated in aqueous solution, and it was ex-
pected that they should behave in the same way in other
solvents. It therefore came as a surprise when Krauss<39)
reported that sodium chloride, a typical ion0phore, behaves
like a weak electrolyte, an ionogene, when dissolved in
liquid ammonia. The electric conductance of such a solu-
tion is given by the law governing the conductance of
aqueous solution of acetic acid indicating that only a
small fraction of the dissolved salt is dissociated into
free ions. To account for these observations, BJerrum
proposed that in liquid ammonia and in other nonaqueous

solvents the oppositely charged ions are associated into

neutral ion pairs which do not contribute to the electric

72

conductance.

Solute-solute interactions result in the formation of
ion pairs or of higher aggregates. The extent of ionic
aggregation will depend not only on the dielectric constant
of the solvent but also on its solvating ability (donicity)
as well as on the nature of the ions. Ion pairing can
result in contact pairs, solvent shared pairs or separated

pairs(40).

At higher concentrations and/or in solvents of
low dielectric constant and low donicity ionic triplets,
quadruplets, etc. may also form. It should be noted also
that the introduction of a salt into a solvent may affect
the solvent-solvent interaction especially in highly struc-
tured solvents such as water or dimethylsulphoxide (DMSO).
To solve this complex puzzle, we must have the means of
identifying the various chemical species present in a given
solution. Once all species are identified, we may then
study their interactions and the equilibria that exist
amongst them. Such data are the key for understanding
chemical processes in solutions. For many years studies
of electrolyte solutions were limited to electrochemical
measurements or measurements of colligative properties of
solutions. It was only since 1960 that the behavior and
structure of ion pairs and ion-pair complexes in solution
have been extensively investigated by using a variety of

spectroscopic techniques such as electron paramagnetic and

nuclear magnetic resonance, ultraviolet, visible, infrared,

73
and Raman spectroscopy. A brief summary of the type of

results obtained using these various techniques is now

given.

A. Ultraviolet-Visible,

 

One drawback of this technique is the inability to
decide whether an observed solvent effect is due to a
ground state interaction, an excited state interaction or
both. The large band width of absorption bands observed
in solution requires large spectral shifts if ion-pair

(5)

association are to be detected by this technique. In
spite of these shortcomings, ultraviolet-visible absorption
spectrOSCOpy is an attractive technique for studying ion-
pair interactions due to the simplicity of its excecution
and the straightforwardness of data analysis. Moreover, it
is not limited to paramagnetic species like electron spin
resonance methods, and does not require high concentrations
of reagents which is often imperative in the nuclear mag-
netic resonance studies.

Warhust ep_al,(ul) were perhaps the first to observe
shifts in optical spectra arising from ion pairing. They
reported that the absorption peaks of ion pairs of alkali
and alkaline earth salts of ketyls of benzophenone, fluor-
enone and other ketones shift towards longer wavelengths

when the radius of the cation increases. Similar effects

were later observed by Hogen(42) for fluorenyl salts, Waak<u3>

74

for organolithium salts and by Zaugg and Schaefer<lu),

for alkali salts of phenols and enols. Warhust gp_al.(ul)
pointed out that the perturbation of the molecular orbital
levels of an organic anion due to the field of the cation
associated with it depends on the electrostatic interaction
energy. Therefore its magnitude is approximately propor-
tional to the reciprocal of the interionic distance given
by the sum of the cation radius r and a constant repre-
senting the corresponding contribution of the anion. The
observed bathochromic shift as the radius of the cation

)

increases was interpreted<uu in terms of a greater de-
stabilization of the ion pair in the ground state. In

the ground state of an ion pair the cation is located

close to the atom possessing the highest electron density.
As the size of the cation increases the distance between
the cation and anion forming the ion pair increases leading
to a destabilization of the ion pair. Since redistribution
of charge occurs upon excitation, such delocalization
effect is smaller in the excited state and a progressive
blue shift is observed as the cation becomes smaller. Con-
sequently, the excited state is less stabilized by the
counter ion than the ground state, and the smaller the
cation the larger becomes the increase in the transition
frequency. Moreover, as expected from the Frank-Condon

principle, the cation has no time to readjust its position

during the excitation causing a further increase in the

75

transition frequency due to this Frank-Condom effect.
(45)

Kosower investigated the spectra of pyridinium
iodides, which are extensively sensitive to the variation
in solvent polarities. The relevant optical transition

is associated with the transfer of a negative charge from
the iodide anion to the pyridinium cation and indeed the
dissociation of these ion pairs eliminates the charge
transfer band. The interaction of an ion pair with the
oriented molecules of solvent, which stabilizes the ground
state, is therefore lost upon excitation.

The absorption maxima of a variety of anionic species
show hypsochromic shifts when they become paired with
cations. The absorption bands of the fluorenyl contact
ion pair shifts to lower wavelength as the radius of the
cation decreases. This is~shown in Figures (32) and (33)(46).
For example, the observed blue shift due to the forma-
tion of a contact ion pair between 9-fluoreny1 anion and

1

Li+ is 1,900 cm“ while with Na+ the blue shift is 1,300

cm-l, the radius of Li+ and Na+ are 0.60 and 0.96 A,
respectively. The formation of solvent-separated ion-
pairs is greatly facilitated by increasing solvent polarity.
Often, small changes in the solvent structure can dras-
tically affect the equilibrium between contact and solvent
separated ion pairs. Typical examples of the effects

exerted by solvents on the absorption spectrum of fluorenyl-

lithium are depicted in Figure (3”)(47). The spectral

76

 

3
U

 

 

 

 

 

 

 

 

 

Figure 32. Correlation between wave number and the inverse
of the cationic radius fog contact fluorenyl
ion pairs in THF at 25°C. 6
w

_. fur

---- r: u.’ in THF oI -3o'c

--- F', (:31

—0 10M”
Figure 33.

Absorption spectra of fluorenyl s

Egts in THF
at -30°C for various counterions.

Figure 34.

 

77

 

l— Spectrum at F: Li' in:

--l- 3-». THF ---- HMPA

 

 

 

1
340 350 380 '400
-> Almpl

Optical spectrum of contact and separated ion
pairs of fluorenyllithium at 25°C in 3,4 di-
hydropyran, ...; 3-methyltetrahydrofuran,
----; 2,5 dihydrofuran ; and hexamethyl-
phosphoramide, ------.fi7

 

78

shifts are the result of a decreased influence of the
cation on the anion, caused by specific cation-solvent
interactions which led to a partial separation of ions.

A pronounced effect was observed with solvents containing
small amounts of reagents capable of coordinating strongly
with the cation, for example, DMSO and hexamethylphos-

phoramide(42’u7).

The lithium salt was singled out for

these studies because of the strong specific interaction

of the small Li+ ion with solvent molecules permits

studies of low polarity media in which the Na+salt exists
only as a contact ion pair. Moreover, the overlap of the
absorption peaks corresponding to the two types of ion

pair (contact and solvent separated) is the smallest for

Li+ salt, and this facilitates the calculation of equilibrium
constants between contact and solventdseparated ion-pairs.
The enthalpy change, AH, governing the equilibrium between
the two kinds of ion pairs is a significant thermodynamic
quantity, knowledge of which permits the discussion of
energetics of ion-pair separation. It also plays an
important role in determining the temperature dependence

of the reactions in which the two kinds of ion pairs simul-
taneously participate, each with its own rate constant<u8’49).
The value of AH for a given equilibrium between contact and
solvent-separated ion-pairs can be obtained from a plot

of K (the ratio between both types of ion pairs) versus

1
l/T, which gives a straight line with slope = AH. A AH Value

79

of —7.5 Kcal/mole was obtained for the formation of solvent-
separated ion—pairs from contact ion pairs for fluorenyl-
lithium in THF. The first evidence for the existence of
contact and solvent-separated ion-pairs in solutions of
carbanion salts came from the study of the temperature—de-
pendent spectrum of fluorenylsodium in THF, Figure (35),
and even more dramatic spectral changes are depicted in
Figure (36)(39) which shows that 9(2-hexyl)f1uorenyllithium
in 2,5-dimethyltetrahydrofuran is transformed from a pre-
dominantly contact ion pair state at -20° to virtually

100% solvent separated ion pair at -40°.

The effect of pressure upon the equilibrium between
contact and solvent separated ion pairs of fluorenyl-
lithium and sodium in THF was investigated by Szwarc
£2 al.(50) and by LeNoble and Das<51). Spectroscopic ob-
servations, extended to pressure as high as 5000 atm.,
show that the equilibrium shifts to the separated ion
pairs as the pressure increases, Figure (37)(51). This ob-
servation was explained in terms of electrostriction (the
observed volume contraction resulting from a light binding
of solvent molecules around the cations of the separated
ion pairs). The binding of one mole of THF to separated
ion pairs is calculated to cause a volume contraction of
about 8 m1., approximately 10% of the solvent molar volume.
Such contraction of the volume of the pure solvent requires

pressures of about 2000 atm. It therefore seems that the

80

 

L5, 373
.0203 cm. CELLl 0.200 cm. CELL

 

ABSORPTION SPECYRUH OF FLUORENYL ’ WI

 

 

 

 

 

 

 

 

l l
440 460 400 500 520 540

Figure 35. Optical absorption spectrum of fluorenyl-
sodium in THF at 25, —30, and -50°c.39

 

 

 

 

 

340 360 3” 400

Figure 36. Temperature dependence of the contact-solvent-
separated ion-pair equilibrium of 9(2-hexy1)-.

fluoren llithium in 2,5-dimethyltetrahydro-
furan.3

81

 

 

 

 

Figure 37. Pressure dependence of the contact-solvent-
separated ion-pair equilibrium of fluorenyl-
sodium in THF at 25°c.51

82

forces contracting the solvent around the ion pairs are
equivalent to those produced by pressure of 2000 atm.

In powerful cation solvating media such as ethylene-
diamine, DMSO, or polyglycoldimethylethers, referred to as
glymes with general formula CH3O[CH2-CH2O]XCH3, most of
the fluorenyl salts form only separated ionapairs. Studies
on ion—pair solvation by such reagents may be carried out
in a mixture of the solvating agent with a relatively low-
polarity solvent in which the fluorenyl salt predominantly
exists as a contact ion—pair. For example, small quantities
of DMSO added to a fluorenyllithium solution in dioxane
converts the contact ion-pairs to DMSO-separated ion-pairs(u2).

Shinohara pp al.(52a’b) have shown that addition of
small quantities of glymes drastically increases the rate
of anionic polymerization of styrene in THF. For example,
at a polystyrylsodium concentration of 5 x 10_5 M, the
polymerization rate increases by nearly a factor of 200
when glyme-5 is present at a concentration of 10'3 M.

The formation of reactive glyme-separated ion-pair and an
increase in the fraction of reactive free ions was respon-
sible for this effect.

Pedersen<53a’b)

developed a class of very strong cation-
binding complexing agents referred to as macrocyclic poly—
ethers or crown ethers. Some of these compounds are dis-
played in Figure (38). The first number refers to the

total number of atoms in the ring, the second number cor-

responds to the number of the ring oxygen atoms. These

83

7 l
6‘0’0 ‘0’0

(I) Dibenzo-lB-crown-6.

c/C\

68

cc. to

is»

(II) Dicyclohexyl-l4-crown-4.

(3‘0\

0 0/
: \C
O
I 0‘0,
C I

‘C

(III) Monobenzo-lS-crown-S.

Figure 38.

84

crown ethers considerably increase the solubility of in-
organic salts in non-polar media and form crystalline
complexes with many salts-

The stability of these complexes depends on the size
of the cation relative to that of the hole of the macro-
cyclic polyether, the charge on the ion, the number of the
ring oxygen atoms, their basicity, coplanarity and sym-
metrical placement, steric-hindrance in the polyether ring,
and the extent of ion association with the solvent<53a).
For example, the small Li+ ion fits the hole of a l4-crown-
4 polyether, but the larger Na+ and K+ ions do not. How-
ever, the Na+ and K+ ions can be accommodated by the cavity
of an 18-crown-6 ether. The stability of the complex is
increased as more oxygen atoms are available for coordina-
tion, provided they are faVOrably located in the poly-
ether ring. In this respect, the dibenzo-l8-crown-6,
Figure (38), is one of the best complexing agents for Na+
and K+ ions.

Similar to the effect of glymes, addition of dibenzo-
or dicyclohexyl-l8-crown—6 to fluorenylsodium in THF yields
1:1 complex absorbing at 372 nm, the optical spectra are
identical to that of the solvent-separated ion-pairs(54).

The size of metal ion plays an important role in de-
termining the stability of its complex with the complexing

agent (crown ether or cryptand). It is evident from

Figure (39)(55) that maximum stability for complexes of

85

  
  

 

 

4.0 -
Ba”
3.5-3
3.0-4
r\ O 1
log K 2.5- (° 3
o o
k/°\/'
10‘ Cavity Diameter 2809M
(from Table 1)
1.5--l
1.0- Cs‘
0.5 If I I j I I

1
0.7 1.0 1.3
Din. Mm/ Ola. Cavity

Figure 39. Selectivity of 18-crown-6: log K values for re-
action of 18-crown-6 with metal cations in H20
vs.ratio of cation diameter to l8-crown-6 cavity
diameter. Value for Ca2+ reported <0.5.55

86

l8-crown-6 with both the alkali and alkaline earth cations
occur at metal ion to cavity diameter ratio of unity (K+
and Ba2+) i.e., greater stability for complexes with
cations diameter more closely matching that of the ligand
cavity<56). Table (6)(56) illustrates that in the case

of 18-crown-6, increased stability of the complexes of K+
and Ba2+ over those of other ions in the series is largely
due to the enthalpy term. This undoubtedly corresponds to
the greater electrostatic bond energy for those ions that
better fit the ligand cavity.

Erying and co-workers(57a’b) concluded that the high
selectivity of 18-crown-6 for K+ over other alkali cations,
Table (7), arises from the slowness of the dissociation
step.

The structure and properties of ion-pairs containing
divalent cations are less known than those of the alkali
salts. The barium salt of the fluorenyl carbanion was

(58) (59)

prepared by Hogen Esch and Smid and by Pascault

as orange crystals. A solution of this salt in THF showed
a sharp absorption maximum at 347 nm, Figure (40)(58).
The fraction of separated ion-pairs is low even at -70°C,
whereas for the sodium salt it exceeds 0.9 under these
conditions. In 1,2-dimethoxyethane (DME) at 25°C the
barium salt is also essentially a contact ion-pair. The

strontium salt yields higher fractions of separated ion—

pair. The spectrum of this salt in THF at -25°C reveals

87

Table 6. Log K, AH, and TAS for the Formation of 18—

 

 

 

Crown-6 Complexes with Metal Ions in H2O at
25°c.(5°)
AH _1 TAS -1

Ion Log K (Kcal mol ) (Kcal mol )
Na+ 0.8 -2.25 -1.16
K+ 2.03 -6.21 —3.40
Rb+ 1.56 -3.82 -1.7
Cs+ 0.99 -3.97 -2.6
Sr2+ 2.72 -3.61 0.1
Ba2+ 3.87 -7.58 -2.3

 

 

Table 7. Rate Constants for Reaction at 25°C in H
u = O.3(‘37a.b):

20’

kt
M+ + 18-Crown-6 E; M-crown+

 

 

 

kf - kd
Cation (M.l s-l) (3-1)
Na+ ~2.2 x 108 3.4 x 107
K+ ~4.3 x 108 3.7 x 106
Rb+ m4.4 x 108 1.2 x 107
Cs+ m4.3 x 108 4.4 x 107
Ag+ m11.3 x 108 3.6 x 107
T1+ 9.9 x 108 5.3 x 106
NHu+ ~5.6 x 108 4.4 x 107

 

 

88

 

0.5 *-

 

OO-

...ao“F5-cee in THF

 

 

k

300 350 400

 

Figure 40. Optical absorption spectrum of barium fluorenyl
and its 1:1 complex with dimethyldibenzo-l8- 8
crown-6 between 320 and 400 nm in THF at 25°C.5

 

Figure 41. Possible structure of the barium fluorenyl-

dimethyldigenzo-lB-crown-6 complex in THF or
pyridine.5

89

two absorption band maxima of approximately equal in-
tensities at 347 and 371 nm, indicating an appreciable
proportion of separated ion-pairs. A mixture of equi-
molar quantities of dibenzo-l8-crown-6 and barium difluorenyl
(or the corresponding strontium salt) in THF produces a

1:1 complex, the spectrum of which reveals approximately
equal fractions of tight and loose ion-pairs. Their pro-
portion is not changed on further addition of macrocyclic
ether. A similar behavior is observed with glyme-6.

A possible structure of the complex is shown in Figure
(Hl)(58) (page 88) in which the Ba2+-crown complex is sand-

wiched in between the two fluorenyl moieties.

B. Infrared and Raman

 

(60a,b) that alkali ions vibrate in solu-

The discovery
tion offers a new fundamental way of obtaining information
about electrolytic solutions. The infrared bands at long
wavelength associated with this phenomenon arise from the
excitation of quantum states of the solution connected
with motion of solution elements adjacent to the alkali
ion. Infrared spectroscopy can be used to obtain informa-
tion about the forces acting on the ions and the structure
of solvent molecules surrounding the ion.

In a study of far-infrared spectra of tetraalkyl-

(61)

ammonium salts in benzene,Evans and Lo found a band

which could not be assigned either to the solvent or to

90

the salt. The authors assumed that it was due to a
certain cation-anion vibration. Calculations based on a
simple diatomic model supported their assumption.
Far-infrared bands arising from the motion of alkali
cation in THF solutions of lithium, sodium and potassium
tetracarbonyl cobaltate, MECo(CO)u], and pentacarbonyl

(61). The

manganate were observed by Edgell and co-workers
large dependence of the frequency of the band on the nature
of the cation, Table (8), and to some extent on the anions,
suggests that the vibrating species are ion pairs or higher
aggregates. Far-infrared spectra were also obtained for

large numbers of alkali salts in a polar and highly solvat-

ing solvent, DMSO<62a’b).

The observed bands were strongly
dependent on the nature of the cation but completely in-
dependent of the anion. The bands were very broad and

their integrated intensities are directly proportional to
the concentration of the salt. The relative intensity,
however, decreases with the increasing mass of the cations.
These bands were assigned to the vibration of the cations

in a solvent cage. This was confirmed by studying the

(62a)

isotopic substitution effect by using 7Li+ and ND“+

instead of 6

L1+ and NHu+ in d6-DMSO which showed that both
the cation and the solvent participate in the observed
vibration. In solvents of low donicity and/or low di-
electric constant e.g., l-methyl-2-pyrolidone, solutions

of salts with polyatomic anions show a constant frequency

for the absorption band, while with some halide anions a

91

Table 8. Absorption Bands of Alkali Metal Salts in Di-
methylsulphoxide.(62a,b

 

 

 

Salt v(cm-l) Salt v(cm-l)
L101 ”29 KBr 153
LiBr H29 KI 153
L11 U29 KNO3 154
LiNO3 ‘ H28 KSCN 153
LiClOu 429

RbBr 125
NHuCl 21H RbI 123
NHuBr 21“ RbNO3 125
NHuI 21H
NHuNO3 21H
NHuClou 21” V C81 110
NHuSCN 214 03010“ 109
NaCl 199
NaBr 199
NaI 198
NaNOu 206
NaClOu 200
NaSCN 200

NaBPhu 198

 

 

92
10-15 crn-l shift to lower frequency was observed. It
seems reasonable to assume that the change in the frequency
of the solvation band is due to a change in the nature of
the solvent cage around the cation. A simple explanation
of this change would be that a small counter—ion, such as
a halide ion, replaces a solvent molecule in the inner
solvation shell forming a solvated contact ion pair. The
cation in such cases vibrates in a cage composed of sol-
vent molecules and a counter-ion.

(63a)

Tsatsas and Risen studied the far-infrared spectrum

of sodium tetrabutylaluminate (NaAlBuu). In cyclohexane

solutions two solvation bands at 195 and 160 cm"1 were ob-

served. In THF solution, however, only one band at 195

cm'1 was observed. Addition of THF to a sodium solution

of tetrabutylaluminate in cyclohexane resulted in the dis-

1

appearance of the 160 cm' band as the THF/NaAlBuu molar

ratio exceeded 1:1. It should be noted that Edgell gt
al.(62a’b)

showed that, in general, the far-infrared solva-
tion bands are Raman-inactive, which is indicative of

the largely electrostatic nature of the cation-solvent or
cationeanion interaction. Tsatsas and Risen, however,
found a 202 cm"1 Raman band in cyclohexane solutions of
NaAlBuu. This band, however, was not observed in THF
solution, which indicates that in cyclohexane the ion-

1

solvent interaction responsible for the 200 cm" band has

a significant degree of covalent character.

93

The addition of a complexing agent, dibenzo-l8-crown-6
to sodium and potassium salt solutions in THF and in pyridine
produced a marked change in the spectra<63b). The absorption
bands were shifted to higher frequency and became inde-
pendent of the solvent as well as of the anion. It is
clear that in this case the cation is accommodated in the
ligand and vibrates within the ring formed by the ether

oxygens.

C. Electron Spin Resonance (ESR)

Application of ESR spectroscopy to study ion-pair systems
has been recently reviewed by Simon<6u). Although it is
limited to the study of dilute paramagnetic solutes, the
technique of ESR spectroscopy has taken its place amongst
other forms of spectroscopy as a useful method for obtain—
ing structural and dynamic information about solvation.

In the field of ion-pair formation, ESR stands out as the
most powerful technique for learning about their presence,

structure and energies<65a-c). Weismann and his oo-

(66)

workers discovered that the ESR spectra of certain
radical-anions in low polarity solvents contained hyper-
fine features characteristic of the alkali-metal ions. The
reason alkali metal nuclei frequently give rise to detect-
able hyperfine coupling is that a slight electron-transfer

from the paramagnetic anions which are, of course, powerful

electron donors takes place into the outer s-orbital of

9“

the cation; and hence even a small transfer gives an
easily detectable coupling.

In the ESR spectra of the benzophenone ketyl anion and
the naphthalene radical anion, each line splits into four
lines by coupling to 23Na as a result of close association

(67’68). The effect of

of the anion with the sodium cation
ion-pairing on the g-value has been studied by Williams

at al.(69). Evidence that different types of ion pairs

are sometimes formed in a solution are strong. Hirota<7o>
found that lithium anthracenide in diethylether gave spec-
tral features for the normal ion pair showing cation hyper-
fine coupling together with a species showing no cation
coupling, but having proton splittings that were different
from the normal solvated ion values. A more definitive

(71) for sodium naphthalenide

study is that of Hofelmann et al.
in THF. Contact ion-pairs [A(23Na) = 1.23 G] together with
another type of ion pair, solvent separated, [A(23Na) =
0.38 G] were detected when tetraglyme was added to the
solution. Excess glyme gave only the latter species.
Triplet-ions formation has also been detected by ESR. The
addition of sodium tetraphenylborate to sodium-2,5-di-t—

butyl-pebenzoquinone in THF resulted in a spectrum charac-

teristic of the triplet ion(72).

95

D. Nuclear Magnetic Resonance (NMR)

 

In recent years the use of NMR in the study of electro-
lyte solutions has become quite popular. Proton and 130
NMR have been widely used for the studies of ionic inter-
actions in solutions. At low temperatures,exchange between
bulk water molecules and those in the inner solvation shell
ofzacation is slower than the NMR time scale and, therefore,
two separate resonances are observed for the free and

bound water. Hydration numbers for a number of transition

(73)

metal cations were obtained by using acetone-water sol-
vent mixture, acetone was assumed to act only as a com-
pletely inert diluent.

Alkali metal NMR as well as halogen NMR has been par-
ticularly fruitful in the elucidation of the structure of
alkali salt solutions in nonaqueous solvents. A compre-
hensive study of sodium iodide solution in nonaqueous

<7u>_

solvents has been reported by Bloor and Kidd Twelve
solvents were used and chemical shifts varied from +9.9 ppm
(upfield) shift for acetic anhydride to -l3.l ppm (down-
field) for ethylenediamine, the chemical shifts were re-
lated to the aqueous pKa of these solvents. More detailed
studies by Popov gt al.(75’76) of different salts indicated
that the counter-ions also play an important role in
determining the behavior of the 23Na chemical shifts.

The concentration dependence of 23Na chemical shift in

the sodium halides and thiocyanate solutions in different

96

solvents can be ascribed to the formation of contact ion-

pairs in these solutions<77).
The 7Li nucleus has a spin of 3/2 and thus is amenable

to NMR studies. Although the nucleus has a quadrupole

7

moment, the Li resonance lines for Li+ ion are exception-
ally narrow and the chemical shifts can be measured with

considerable accuracy.

(78) (79)

Studies by Maciel E; El- and by AkittanuiDowns
in several nonaqueous solvents showed that the frequency
of the 7Li resonance is very sensitive to the environment.

While a large number of 7Li, 23Na and 13303 NMR
studies have appeared, 39K and especially 89Rb NMR investi-
gations are scarce due to the low sensitivity in 39K and
the broad resonance lines obtained with 89Rb.

In dilute solutions of'lithium perchlorate in nitro-
methane and THF the width at half-height of the 3501
resonance was 10-20 Hz. Considerable broadening was ob-
tained with increasing the salt concentration in both sol-
vents whereas, very little dependence was observed in
acetone, methanol and acetonitrile solutions. This broad-
ening was due to the change in the electric field gradient
at the chlorine nucleus as the perchlorate ion interacts

with Li+ to form a contact ion pair<80).

97

E. Luminescence

 

In comparison with the other spectrOSCOpic methods,
luminescence spectroscopy was little used in studying ion-
pairs. Much information can be obtained using luminescence
techniques. Due to charge redistribution in the excited
state, ion-pair characteristics will be different in the
excited state compared to the situation in the ground state.
Excitation may cause a shift in the equilibrium between
free ions and ion pairs i.e., a change in the dissociation
constants.

(81) used the phosphorescence lifetimescn?phenyl-

McGlynn
carboxylic acids and their saltstx>show that the lowest
triplet state of benzoic acid is not an excitation localized
on the benzene ring but involves considerable delocalization
onto the carboxyl group. The large spin-orbit coupling
effect caused by the heavy lead ion in the lead salt of
benzoic acid was completely lost in the lead salt of A-
phenyl-butyric acid, in the latter case the benzene ring
is separated from the carboxylic group by three methylene
groups. These observations also show that the observed
effect of lead on the lifetime of the benzoic acid salt is
not an external heavy—atom effect on the benzene ring.

Recently, Carter and Gillispie<82)

studied the luminescence
of ion pairs of alkali-metal cations with dichromate anion.
The influence of other alkali-metal cations on the emission

of sodium dichromate was probed by adding their chloride

98

salts to the sodium dichromate solution. The wavelengths
of the emission maximum in aqueous solution at 77°K de-
creased continuously in going from Na+ to Cs+. The emis-
sion was assigned as phosphorescence. The greatly en-
hanced phosphorescence intensity when K+, Rb+ or Cs+
cations are present reflects the heavy-atom effect on the
spin-forbidden process (triplet + singlet). The magni-
tude of any heavy-atom effect is presumably sensitive to

(46). If the ion

whether the ion pairs are loose or tight
pairs are loose the heavy-atom effect may be an external
one whereas, in the tight ion pairs case it is better
described in terms of an internal heavy-atom effect, i.e.
because of orbital overlap between the already bended
emitting species and heavy atom not through collision
mechanism with the heavy atom. Accordingly,the emission

at 77°K was assigned to ion-pair formation between the

alkali metal cation and the dichromate anion.

III. Results and Discussion

 

The room temperature absorption spectrum of a dilute
solution of sodium salt of salicylaldehyde in DMSO exhibits
two absorption bands, Figure (“2). The band absorbing at
417 nm corresponds to the absorption of the free anion.

A solution of the same concentration of tetrabutylammonium
salt of salicylaldehyde in DMSO absorbs at A17 nm, the

bulkiness of the tetrabutylammonium cation prevents

99

 

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100

ion—pair formation. The other absorption at 323 nm is
interpreted as corresponding to the absorption of the as-
sociated metal cation with the organic anion. This inter-
pretation is consistent with the increase in absorptivity
of the 323 nm band upon addition of Na+ ions while that of
the ul7 nm decreases as shown in Figure (”2). Similar
behavior was observed for the thallous, Tl+, salt of
salicylaldehyde, Figure (“3). In the presence of excess
T1+ ions, the maximum of the A17 nm absorption band was blue
shifted, this probably reflects a change in the solvent
structure while the absorptivity of the 323 nm absorption
band is enhanced, Figure (A3). In water the 323 nm ab-
sorption band is weak and appears only as a shoulder.
Figure (AU) shows the potential energy of an ion pair as

a function of the interionic distance R. The absorption

of the free anion is represented by AE. For the formation
of an ion pair, the positively charged cation gets closer
to the anion, the smaller the ionic radius of the cation,
the closer the distance between anion and cation, resulting
in a greater stabilization of the ground state of the ion
pair. Because of the covalency in O-H bond, maximum
stabilization corresponds to the formation of the phenolic
derivative. In the excited state, the negative charge on
the oxygen will have partially migrated to the ring, and will
decrease the negative charge on the phenolic oxygen.

Accordingly, less interaction between the cation and anion

101

 

 

 

 

 

8

,C:
3 _

.8

.8

<1
0.0 b l 1 l 1 9

300 350 400 450
Wavelength (nm)

Figure H3. Room temperature absorption spectra of dilute
solutions (5 x 10'5 M) of Tl(I) salt of salicyl-
aldehyde in DMSO (----) and in DMSO in presence
of 5 x 10-3 M T1N03 ( ).

 

102

AB, ©—0'--~ M“

 

 

2§J59<C:>'()"fi

 

 

 

chation-anion interionic distance

Figure “A. Potential energy of ion-pair formation as a
function of the interionic distance.

103

will occur, resulting in an increase of R and less stabi-
lization of the ion pair in the ground state. During the
absorption,the interionic distance does not change and
therefore a Franck-Condom energy contributes to the transi-
tion energy. The result is a blue shift in the absorption
spectra of the ion pair, with respect to that of the free
ion. This blue shift increases with the decreasing cationic
radius. The blue shift of the “17 nm band, Figure (A3),
to MOO nm in presence of excess (5 x 10'3 M)Tl+ indicates
that the solvated cation interacts with the solvated anion,
i.e., solvent separated ion pairs absorbing at A00 nm

are formed. Thus the following equilibria occur in solu-
tion

M+SSL- : M+L'

M+S + L‘s

—> .
+.

solvated free ions solvent separated ion contact ion

pairs pair
(417 nm) (MOO nm) .323 or 326
S: solvent M+z Na+ or Tl+ L‘: salicylaldehyde
anion

In ethanol, both sodium and thallium salts of salicyl-
aldehyde show an equilibrium between the free ions which
absorb at 380 nm and contact ion-pairs which absorb at
325 nm. Thus adding an excess of Tl+ shifts the equilibrium
to the right; the relative concentration of the contact

ion-pairs increased. In addition, excess Tl+ induces

10A

structural changes in the liquid structure of the solvent.
The enhanced order in the solvent structure favors the
formation of solvent-separated ion-pairs at the expense

of solvated free ions. In THF solution of the sodium

salt of salicylaldehyde, the longer-wavelength absorption
band dominates the spectrum as shown in Figure (H5).

This indicates that in THF there is an equilibrium between
the free ions and contact ion-pairs with the equilibrium
shifted towards the free ion formation. This can be ex-
plained in view of the high solvation of Na+. For a thal-
lium salt in THF, Figure (N6), both absorptions of the ion
pair and free ions are of comparable absorptivity. This
reflects the lower solvation of the Tl+ ion with THF com—
pared with Na+ because of the larger radius of Tl+.

Figure (N7), shows a comparison between the absorption
spectra of Na*and Tl+salts in THF.

When the anion of salicylaldehyde is excited, charge
transfer occurs from the phenolate oxygen to the carbonyl
oxygen. As a result,the negative charge density on the
phenolate oxygen will be smaller in the excited state than
in the ground state. This weakens the interaction between
the positive and negative ions in the excited state‘ i.e.,
the ion pair becomes less stable, and may lead to dis-
sociation of the ion pair in the excited state. Figure (nu)
shows the energetics of both excited and ground states of

ion pairs. Figure (H8) displays the room temperature

Absorbcnce

105

 

 

 

 

 

l l l l
250 300 350 400
Wavelength (nm)

Figure #5. Room temperature absorption spectra of dilute
solution of Na+ salt of salicylaldehyde: (H. 8
x 10-5 M) in water (----) (u. 5 x 10-5 M) in
ethanol ( ----- ) and (u. 5 x 10-5 M) in THF

 

106

 

04—/I

 

 

 

 

Q)
U
C
£302.
3
(D
13
<1 -
0.0L 1 l l l
250 300 350 400
Wavelength (nm)

Figure H6. Room temperature absorption spectra of dilute
solutions of Tl(I) salt of salicylaldehyde:
(u.5 x 10-5 M) in water (--——), ( .2 x 10-5 M) .
inethanol ( ----- ) and (H.5 x 10- M) in THF

 

107

 

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Fluorescence Intensity

 

Figure H8.

 

108

 

 

 

 

l___l I 4 I
400 450 500 550 600

Wavelength(nm)

Room temperature emission spectra of dilute
solutions of Na+ salt of salicylaldehyde

(5 x 10-5 M) in DMSO: lexc = 320 nm (--—-),
Hexc = 380 nm ( ) and in DMSO in the pres-
ence of NaBFu (2 x 10'2 M),A exc 8320 (- ----- ).

 

109

emission spectra of the sodium salt of salicylaldehyde in
DMSO. Exciting at the absorption maximum of the free anion
produced an emission from the free anion at H80 nm. Ex-
citation at 323 nm, the absorption maximum of the contact
ion-pair gave two emissions. The shorter-wavelength
emission at H05 nm belongs to the associated pair and the
longer-wavelength emission at H80 nm to the free anion.

2 M) Na+ ions, excitation at

In the presence of (2 x 10-
323 nm gave only the H05 nm emission whereas, excitation

of the same solution at 380 nm produced only the free

anion emission. Thus, the production of two emissions

when exciting at the absorption band of the contact ion
pair indicates that the contact ion pair dissociates as

a result of excitation i.e., the equilibrium shifts towards
free ion formation in the excited state. Similar results
were obtained in the corresponding thallium (I) salt. In

a dilute solution of the thallium salt (5 x lo“'5 M) in

DMSO in the presence of (10'3 M) T1 N0 Figure (H9), the

3,
anion emission was observed upon excitation at the longer-

wavelength whereas, fluorescence quenching occurred when

at the absorption band of the contact ion-pair was excited,
this may reflect the high degree of covalency in the metal
cation-salicylaldehyde anion ion-pairs. Since in the

same solution the observation of the intense anion emis-

sion eliminates the quenching through collisional mechanisms

by the excess T1+ ion present in solution, i.e., the

110

 

Fluorescence Intensity

 

 

 

jf, l. 1 l 1 ° 'T' I
350 .400 450 500 550 600 650
Wavelength (nm)

 

Figure H9. Room temperature emission spectra of dilute solu-
tions (5 x 10-5 M) in DMSO of: Tl(I) salt of
salicylaldehyde, Aexc = 320 nm (----) and H20 nm
( ); Tl(I) salt of salicylaldehyde in pres-
ence of ~5 x 10-3 M T1N03, lexc = 320 nm (-----).

 

111

(“6). In

quenching is via internal heavy atom effect
solutions of Na+ and Tl+ salts of salicylaldehyde in THF,
Figure (50), both gave the free anion emission at 500 nm
when excited at 320 nm where the ion pair absorbs, or at
380 nmwhere the free anion absorbs. The red shift of
the emission of the anion in THF compared with that in
DMSO at H80 nm can be interpreted as a result of more
stabilization of the ground state in DMSO than in THF.
Excitation of the anion leaves a smaller negative charge
on the phenolate oxygen in the excited state compared with
the ground state charge distribution. In the presence of
excess metal ions the emission spectra started to show
ion pair emission peaks at ~H20 nm as well as the anion
emission. The intensity of this emission was higher for
the Na+salt than for the K*salt i.e., the shift in the
equilibrium between the ion pairs and free ions is more
shifted towards ion pair formation in the case of Tl+than
Na salts. In aqueous solutions,complete dissociation,at
least in the excited state,gave only the anion emission.
In this section, the results of our study of the
interaction of monovalent alkali metal ions with three
phenols,name1y: p-hydroxybenzaldehyde, p-methoxybenz-
aldehyde,and p-methylphenol are summarized. Para-sub-
stituted phenols were chosen to contrast possible inter-
action unique to ortho-substituted derivatives. It should

(1“)

be pointed out that Zaugg and Schafer studied the

112

 

Fluorescence lntenslty_

 

 

 

 

Figure 50.

400 500 660
Wavelength (nm)

Room temperature emission spectra of dilute
solutions (5 x 10'5 M) of Na+ and T1+ salts of
salicylaldehyde in THF. Na+ salt, Hexc = 320
and 390 nm ( ); Na salt in presence of

~2 x lo-2 M NaBFu,A e . 320 nm (----); Tl(I)
salt, Aexc = 320 and 390m W§ -); and Tl(I)
salt in presence of ~5 x 10- M TlNO3, Hexc a
320 nm (-------).

 

113

cation and solvent effects on the ultraviolet absorption
spectra of alkali metal salts of some phenols and enols.
They reported a red shift in the absorption spectrum of
alkali salts of a phenol in dimethoxyethane as the radius
of the cation is increased. These results were inter-
preted in terms of an increased interionic distance, and
hence smaller interaction between the anion and cation
forming the contact ion pairs.

The absorption spectra of dilute solutions (5 x 10'5 M)
of p-hydroxybenzaldehyde and its Li+,Na*and K*sa1ts in THF
are shown in Figure (51). The phenol absorbs at 272 nm.
When LiH was added to the phenolic solution, the obtained
absorption spectrum showed two bands: one identical to
that of the phenol and another band with maximum at 328 nm.
The shorter-wavelength absorption band is due to unreacted
phenol; while the absorption band at 328 nm is due to ion-
pair formation. Addition of NaH or KH to the phenolic
solution produced the sodium or potassium salt respectively,
and the reactions were complete. The absorption maxima
occurred at 33H nm for the Né’salt and at 3H7 nm for the K+
salt. The increase in the radius of the cation leads to
progressive red shift of the absorption of the ion pair
from 328 nm for Li" salt to 3H7 nm for the K"sa1t. The
free anion of p-hydroxybenzaldehyde absorbs at 360 nm in
THF as obtained from the absorption spectrum in the pres-

ence of tetrabutylammonium hydroxide. When 18-crown-6

11H

 

 

 

 

 

 

I f)? F?
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U.
U.
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.o l‘
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. l‘=
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\‘-.
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250 300 350
Wavelength (nm)

Figure 51.

Room temperature absorption spectra of dilute
solutions (5 x 10'5 M) of alkali-metal salts
of p-hydrOxybenzaldehyde in THF. P-hydroxy-
benzaldehyde ( ); Li+salt (....); Na+salt
(- ----- ) and K+salt (----'--).

 

115

was added to metal salt solution, no effect was observed
for the Li+sa1t. Li+ does not form stable complexes with 18-
crown-6. However, in the case of the Na*and KIsalts, red

1 and 900 cm"1

shifts of 500 cm' respectively,were ob-
served. The stability constant of the K+ complex with
18-crown-6 is larger than that of Na+. These red shifts
are interpreted in terms of the crown ether forming a
complex with metal cation in the ion pair. Such inter-
action will decrease the interionic separation,decreasing
the perturbation influence of the positive charge of the
cation,causing a red shift, i.e., the spectrum approaches
that of the free anion. The room temperature emission
spectra of p-hydroxybenzaldehyde salts in THF are shown
in Figure (52). The emission maxima of the Lit Na+and K+
salts occur at 357, 361,and 365 nm respectively. It is
clear that the emission maximum undergoes a red shift
with the increasing radius of the metal ion. In the
presence of l8-crown-6 both Na+and K+sa1ts produced the
same emission maximum at 372 nm which is the same emis-
sion of the free anion in THF. These results indicate
that the ion-pair, weakened by cation interaction with
18-crown-6, dissociates upon electronic excitation giving
rise to the excited free anion.

Comparing the absorption and emission spectra of the
salts of ortho and para hydroxybenzaldehyde leads to the
conclusion that a tighter ion pair is formed with salicyl—

aldehyde and that a metal chelate is formed with the former

116

 

Fluorescence Intensity

 

 

 

Figure 52.

 

Wavelength ( nm)

Room temperature emission spectra of dilute
solutions (5 x 10'5 M) of alkali-metal salts of
p-hydroxybenzaldehyde in THF. Li*salt, 1e c =
335 nm (--—-), K+salt, He c = 3H5 nm ( 1 and
K+salt in presence of 10-3 M 18C6, Ae c = 3H5 nm
(--—----), Na+salt l = 3H0 nm (....)‘, Na+ salt

 

in presence of 10‘3eMc18C6, Aexc = 3H0 nm (------

117

where the ortho-substituted anion acts as a bidentate ligand
with both phenolate and carbonyl oxygen atoms.

Similar results were obtained for p-methoxyphenol,
the absorption spectra of its metal salts are shown in
Figure (53). Addition of LiH did not change the absorp-
tion spectrum of the phenol, it appears that no salt for-
mation occurred. The absorption maxima of the Na+and K+
salts indicate that ion pairs are formed in THF solu-
tions. The corresponding emission spectra of the two
salts are displayed in Figure (5H). The emission spectra
of both Na+and K+ salts are blue shifted compared to that
of the parent phenol. When l8-crown-6 was added to the

Na+

salt in THF, a new emission was observed at HHO nm, but
the original emission at 36H nm of the salt while dimin-
ished in intensity still occurred. In the case of the K+
salt, addition of l8-crown-6 led to the disappearance of
the 367 nm emission of the ion pair and only one emission
band was observed at HHO nm. The latter emission cor-
responds to the free anion emission obtained in THF in the
presence of tetrabutylammonium hydroxide. These results
are interpreted in terms of the change in the stability of
the ion pair upon electronic excitation. In the excited
state, as mentioned before, the charge density at the
phenoxide oxygen is diminished due to the charge migration

to the benzene ring. This weakens the interaction with

the metal ion. However, this effect does not manifest

Absorbance

118

 

0.8 -

0.6 -

I

0.2

 

 

 

 

 

 

 

Figure 53.

2:50 360 350
Wavelength (nm)

Room temperature absgrption spectra of dilute
solutions (1.8 x 10' M) of alkali-metal salts
of p-methoxyphenol in .THF. P—methoxyphenol
(-———), Na+salt (~-- - -) and K+salt (-- - ).

119

 

 

 

 

 

l I l
)5
3: 1
(n a.
g [I \\
*- .’ / ‘ 3
.EE 3 I \\.H
0) \ ,3 I \ 1
o \ :' ’ ‘
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§ \ / \
\f\/ ‘{
e ' \ X \
O . / °\ \ .
3 )‘ '\ \\'"-.
LL / .\. .\ \.\'
/ .\,\\.
L l l N l
300 _ 350 400 450 500
Wavelength (nm)
Figure 5H.

Room temperature emifision spectra of dilute
solutions (1.8 x 10' M) of alkali-metal salts
of p-methoxyphenol in THF. P-methoxyphenol,
lexc =- 290 nm ( ); Na+salt, xexc -- 320 nm
(-'---); Natsalt in presence of 10-3 M 18C6,
lexc - 320 nm (....); xtsalt, lexc = 33; nm
(-°°-°--) and K*salt in presence of 10' M
18C6, Aexc = 335 nm (----).

 

120

itself except in the presence of the crown ether. Thus,
in the excited state, the crown ether competes effectively
with the phenoxide anion for the metal cation and is
particularly true in the case of K+ which has a larger
radius and forms a weaker bond with the anion and also
because it forms a more stable complex with the crown
ether compared with Na+. Thus, the two emission bands
observed in the case of the Na+salt in the presence of 18-
crown-6 are due to partial dissociation of the ion pair

in the excited state. In the case of the K+sa1t in the
presence of 18-crown-6, complete dissociation of the ion
pair occurs.

The absorption and emission spectra of alkali metal
salts with p-methylphenol in THF are shown in Figure (55)
and (56), respectively. Ion-pair formation occurs in
THF and their absorption maxima occur at 310 nm for Na+
and 317 nm for K1 In the presence of l8-crown-6 a further
red shift occurs for K+giving an absorption at 326 nm.

The emission spectra gives rise to a maximum at 3HH nm and
355 nm for Na+and Kt respectively. This should be com-
pared with 360 nm for the free anion emission in THF.
Table (9) summarizes the absorption and emission spectral
results of the three phenols in THF.

From the previous discussion, it is clear that mono-
valent cations form tight ion pairs in THF. The interionic

separation increases with the size of the cation giving

121

 

 

 

 

 

 

k u
0.8“ ‘
06* "
a)
o
8 F l -
.{3 l
o ‘ -
(n 0.4 F l
5 : l
<1 .- . \ q
° l
= l
02" ' = ‘ '1
. ( \
- '\ \\ .-
'\ '\\\
0.0 L l \l‘ O. "
250 300 350

Figure 55.

Wavelength (nm)

Room temperature aberption spectra of dilute
solutions (2.5 x 10' M) of alkali-metal salts
of p-methylphenol in THF. P-methylphenol (
Na*salt (-----), K+salt (—~----) and K’salt
in presence of 10'3 M 18C6 (----).

 

),

122

 

 

 

 

 

I
Z.‘
"(5
c
m
15
a)
o
c
m
0
U)
9
8
LT. ..
l l ." ..T
300 350 400
Wavelength (n m)
Figure 56.

Room temperature emission spectra of dilute
solutions (2.5 x 10' M) of alkali-metal salts
of p-methylphenol in THF p-methylphenol, Aexc
= 290 nm ( ); Natsalt, Aexc = 320 nm ( ----- )
and Kisalt, Aexc = 330 nm (-------).

 

123

 

 

 

 

 

 

 

 

 

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rise to a red shift in going from Na+ to K+. In the pres-
ence of 18-crown-6 it interacts with the cation resulting
in a looser ion-pair formation. In the excited state the
ion pair becomes weaker due to a decrease in the negative
charge on the phenoxide oxygen and in the case of bigger
cation (K+) in the presence of l8-crown-6 complete dis-

sociation of the ion pair may occur.

CHAPTER V

EUROPIUM (III) CHELATE WITH SALICYLIDENE-VALINATE

SCHIFF BASE

I. Lanthanide Ions as Luminescence Probes of the Structure

 

of Biological Macromolecules

 

Calcium and magnesium occur as essential inorganic
ions in living systems. A Ca2+ concentration gradient is
maintained across many cell membranes with the concentra-
tion within the cell lower than that outside. Ca2+ is
involved in muscle contraction, blood clotting, K+ trans-
port, neurotransmitter release, cellular adhesion, inter-
cellular communication, microtubule formation, and hor-
monal responses. Many of these Ca2+ related activities

occur by means of interaction with proteins, which Ca2+

may stabilize, activate, and modulate. In this way Ca2+
plays a significant regulatory role in many biological
processes. Because of the lack of suitable physical
techniques for studying Ca2+, it has not received the
attention it deserves. The electronic transitions of
Ca2+ cannot be studied by conventional optical absorp-

tion and emission spectroscopy, and the absence of un-

paired electrons precludes the use of magnetic resonance

125

126

techniques in probing the chemical and structural nature

of Ca2+ binding sites.

Fortunately, about 12 lanthanide ions (Ln3+) possess

properties which make them excellent probes for Ca2+(83).

2+ and Ln3+

In forming complexes, both Ca prefer charged

or uncharged oxygen donor groups to nitrogen donor atoms.
In aqueous solution, except for some multidentate ligands,
hydroxo complex formation almost always occurs before amine

(8“). Both Ca2+ and Ln3+

nitrogen coordination takes place
display a variable coordination number and a lack of strong
directionality in binding donor groups. Trivalent lan-
thanide ions Ln3+ exhibit effective ionic radii that show

a gradual contraction from one end of the series to the

3+ (atomic number 57)

other, for example, from 1.16 K for La
to 0.98 K for Lu3+ (atomic.number 71) in 8—fold coordina-
tion(85). For a given coordination number, the decrease

in the ionic radii in going from one end of the lanthanide
series to the other (0.17 - 0.19 K) is about equal to the
increase in ionic radii on going from 6- to 9-fold coor-
dination. Thus, there are Opportunities for fine adjust-
ment upon substitution of heavier Ln3+ for Ca2+ ions either
by small decreases in ionic radii or by an increase in co-
ordination number. As an example, the ionic radius of 7-
coordinate Ca2+ is nearly equal to the ionic radius of 8-
coordinate Eu3+ ion. The replacement of Ca2+ by Ln3+

has been demonstrated by x-ray diffraction to occur in two

(86a) (86b).

proteins, carp parvalbumin and thermolysin

127

In the latter case, the results have been reported with
sufficient precision to allow some comparisons of bond
distances and coordination numbers between Ca2+ and Ln3+
containing proteins. Accordingly, most Ln3+ should be
able to substitute Ca2+ in proteins without causing serious
structural modifications in the active sites. The unit
charge difference between Ca2+ and Ln3+ is apt to be of

3+ for Ca2+.

secondary importance for substitution of Ln
In biological systems Na+ and Ca2+ of comparable ionic
radii have been found to be competitive for sites as have

K+ and Ba2+<87’88).

Charge differences may assume more
importance in rate phenomena which may be sensitive to
net charge at a binding site.

Except for La3+ and Lu3+, all 12 other available Ln3+
ions contain unpaired f-electrons, and both nuclear mag-
netic resonance (NMR) and electron spin resonance (ESR)

methods may be used to investigate their environment<89a’b).

3+ may be used as chemical shift probes in NMR.

Several Ln
cd3+ with half-filled Hf subshell is used as a broadening
probe in NMR and also is employed in ESR studies. These
12 Ln3+ ions also exhibit absorption spectra due to inter-
configurational f-f transitions which are easily acces-
sible to study by conventional optical absorption tech-
niques. However, the molar absorptivity associated with
these transitions are generally so low (of the order of

unity) that the study of Ln3+-protein systems in which Ln3+

128

concentrations are smaller than 10'2 M, optical absorption

spectroscopy can be performed only near the limits of
instrumental sensitivity. 0n the other hand, luminescence
due to interconfigurational f-f transitions in Ln3+ ions bound

to protein systems remains well above detection limits

3+

even when Eu and Tb3+ are present at concentrations as

low as 10.6 M. Thus a factor of at least 10Ll in sensitivity

favors luminescence over absorption spectroscopy.
In order for luminescence to occur, appropriate excited
states of the emitting (luminescent) species must be

populated. The use of Ln3+

3+

ion emission spectra as a
probe of Ln complexes requires, then, excitation of
the Ln3+ f-f emitting states. Excitation may be accom-

plished either by direct excitation of the metal ion

 

 

chromOphore (the exciting light being absorbed directly

by the Ln3+ ion) or by an indirect mechanism involving

 

optical excitation of ligand chromophores follows by

sensitization of Ln3+

emission via radiationless energy
transfer (ligand to metal ion). Because of the weak
absorptivities of Ln3+ transitions in the visible and
near-ultraviolet spectral region, direct excitation of Ln3+
emission requires either a powerful excitation source (such
as a laser) or relatively high concentration of the metal
ion (>lO"2 M). 0n the other hand, indirect excitation of
Ln3+ emission may be accomplished using somewhat less

intense exciting light and at lower metal ion concentra-

tions if the ligand environment includes a highly absorbing

129

chromophore capable of acting as an efficient energy donor
to the Ln3+ acceptor-emitter species. Both the direct and
indirect excitation methods have been used in emission
studies of Ln3+-protein complexes.

In the indirect excitation method, in Ln3+-protein
3+

systems, the Ln ion may be excited by energy transfer
from nearby aromatic side chain chromophores which are
excited directly by near-ultraviolet radiation. The over-
all process is thus started by an absorption of an aro-
matic side chain of phenylalanine, tyrosine, or trypto-
phan in the 250-300 nm region; followed by a non-radiative
energy transfer from an aromatic side chain to a nearby
Ln3+; giving rise to a strongly enhanced (up to 105)
Ln3+ emission in the visible region. The energy transfer
step probably occurs by‘a Forster dipole-dipole reson-
ance transfer mechanism<90> with a r.6 dependence on the
distance between the donor and acceptor sites. A semi-
quantitative analysis suggests that 50% probability of
energy transfer occurs at a donor to acceptor distance
of 5-10 3 for tyrosine-Tb3+ and tryptophane-Tb3+ pairs<9l).
Using Tb(III) or Eu(III) as an acceptor, Horrocks gt
gl,(9l) measured the distance between tryptophane as an
energy donor and the calcium-binding sites in parvalbumin
from codfish. The 11.8 and 10.2 K distances obtained using
Tb3+ and Eu3+ as acceptor respectively, are in a good

agreement with the 11.6 A distance estimated from x-ray

structural results.

130

Direct excitation of Eu3+ and Tb3+ emission using a
pulsed laser source was employed in the study of lumin-
escence lifetimes for a variety of complexes in aqueous

(92).

solution The lifetime and emission decay charac-
teristics of Eu3+ and Tb3+ with a variety of ligands (in-
cluding the protein thermolysin) were measured as a func-
tion of D20 versus H20 content of the water solvent.

Such studies gave estimates of the number of water mole-

3+

cules bound to the Ln in each complex<92). This follows

from the somewhat different perturbative influences of D20
versus H20 upon the rate of radiative and radiationless pro-
cesses of Eu3+ and Tb3+ excited states(93).

Eu3+ and Tb3+ are the most strongly emitting members

3+

of the Ln3+ series. Energy level diagram for Eu and

(11).

Tb3+ are shown in Figure (57) The 5D and 5D“ emis-

0

sion levels of Eu3+ and Tb3+ can be excited by 579 and
H88 nm light, respectively. The use of visible light
for the excitation eliminates problems of protein-ultra-
violet photosensitivity. The excited state lifetimes of
these ions are environmentally sensitive and lies in the
conveniently long 100-300 us range. While ligand field
splittings of f-electron levels are much smaller than for
the d-electron levels of transition metal complexes, the
absorption bands are generally much narrower, and small

splittings can be readily detected. Individual emission

bands can be examined under high resolution to yield

Figure 57.

Energy (cfi' X l0'3)

131

z:-
Europium(III) “0*" "04” Tctbium (m)

'05 _. v07

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

’0:
"b ., W l ’°‘
n— ..4
’°° l \- ""
Is- "3 __ v“
- '03
'02-
: E 5 5
C C o
It)h . 2 _ I v
5 ‘.’ " : e
2. e «l- 1 e
2 .2 - vol .3 2
5» "‘ 0’0 - — '0 "o
‘ 7
7'2 '5
o- 'r.

 

__ 7"

Electronic energy levels for europium(III) and
terbium(III). The two upward-pointing arrows
show the transitions which occur upon laser ex-
citation at the wavelengths indicated. The

two downward-pointing arrows label the most
intense emissive transitions of the two ions.
Radiationless energy transfer competes with

the radiative processes through coupling of

the emissive states to the O-H vibrational over-
tones of coordinated H20 molecules. (The energy
level diagram of Tb(III) has been displaced
slightly to position the highest electronic
acceptor state of the ground 7F term to coin-
cide with the zero-point energies of the vibra-
tional overtone ladders.)11

132

information about the splitting of ground and excited

states. Excitation spectroscopy using tunable lasers also

yields information of this type.

A. Decay Lifetimes as a Measure of the Number of

 

Metal Coordinated Water Molecules

 

The early observation that hydrated crystals show a
much reduced fluorescence intensity in the middle of the
series of rare earth ions (Sm, Eu, Gd, Tb, Dy) and prac-
tically no emission in those at the beginning and end of
the series, led to the conjecture(9u) that the highly
energetic O-H vibrations in the crystal lattice act as
energy sinks. The role of OH vibrations in the quenching
of electronic f-f transitions in the rare earth transitions

' (95)

was demonstrated in crystals by measuring fluorescence
lifetimes of deuterated and protonated hydrous crystals

of EuCl and TbCl . Quenching occurs if the energy gap

3 3

between the lowest fluorescent and highest non-fluores-
cent level is matched by a single high energy (OH or OD)
vibration of one solvent molecule. The decisive parameter
is R, the ratio between the electronic energy gap (E)
)(96).

and the vibrational quantum energy (hm
R = E/hw

Of all the lanthanides Gd is in a class by itself with

133

R m 10 for H O and m15 for D20.

2
There is a very low probability of radiationless

transition to the very high vibrational quanta thus re-

(16)

quired. Accordingly addition of a single high-energy

vibrational quantum due to deuterium substitution make
no significant contribution to the radiationless dissipa-

tion of electronic energy. Hence, no isotOpic effect is

observed in the case of Gd. For Eu3+ (E m 12,200 cm’l)

3

the isotope effect is the greatest of all Ln(III) series
and is due to the requirement of v = H for OH and v = 5
for OD. Radiationless transition probability is not small

and will differ greatly depending on whether H O or D O

2 2

is interacting with Eu(III). For Tb3+, (E m 1H,800 cm‘l)

,

radiationless transition occurs to v = 5 for OH and to

v = 6 or 7 for OD resulting in smaller probability for
3+

the transition compared to the case of Eu and hence

longer lifetime for Tb3+ is observed, and the isotopic

effect is smaller.

The experimental reciprocal excited-state lifetime

(exponential decay constant), Toisd’ consists of several

terms

T-1 = T-l + T--l + -1
obsd nat nonrad T0H

where T-1 is the natural rate constant for the emission

nat

nonrad represents the rate constant for

of photons, T

13H

nonradiative deexcitation which does not involve OH oscil-

lators, and 16% is the rate constant for nonradiative

energy transfer to the OH vibrational manifold of OH

oscillators in the first coordination sphere (e.g., oo-

16% is very significant. For
instance, for Eu3+ (aq), 16% = 9.5 ms-l, whereas, 1-1

nat
= 0.19 ms‘l, Thinrad = 0.25 ms’1(93’97a).

ordinated water molecules);

Replacement

of OH oscillators by the OD variety causes the vibronic
deexcitation pathway to become exceedingly inefficient(97a'c)
because of the smaller orbital overlap as the vibronic
quantum increases, and enables one to determine the number
of OH oscillators in the first coordination sphere by

carrying out experiments independently on H20 and D20

solutions. In D20 solutions TOH vanishes and Tobsd —
-1 -1
Tnat + Tnonrad’ where the latter term includes any small

deexcitation via DO oscillators. 13%8d varies linearly

with the mole fraction of H20, XHZO’ in H20-D20 mix-

tures(92). Figure (58)(11), shows a plot of Tobsd vs

ngO for EuCl3. In pure H20, nine water molecules are co-
ordinated to Eu3+. When 1 mole of EDTA (ethylenediamine-
tetracetate) was added the number of coordinated water
molecules dropped to six. This is in agreement with the
tridentate nature of EDTA which substitutes water mole-
cules in the first coordination sphere. The same tech-
nique was used to estimate the number of coordinated water

molecules in crystalline solids.(98a)

Figure 58.

135

 

" Europium(III)

 

 

Plot of the observed reciprocal luminescence
lifetimes, ragga d, vs. the mole fraction of H20,
XHZO: in D2 0- H 0 mixtures of Eu(III) solutions.
The effect 2of ghe chelating ligands NTA and

EDTA is to displace water molecules from the
first coordination spheres of the aquametal ions
to yield the approximate values indicated along
the right-hand ordinate.1l

136

B. Characterization of Individual Binding Sites

Eu(III) ion is unique in that both the ground (7FO)
and the emissive excited (5DO) states are nondegenerate.
Since neither of these levels can be split by a ligand
field, the absorption band corresponding to a transition
between these levels must consist of a single, unsplit
line for a given Eu(III) ion environment. Different ionic

environments can, in principle, yield transitionszuzslightly

different frequencies. Because 5D + 7F transition is

O 0
highly forbidden, study of this band by ordinary absorption

spectroscopy on dilute solutionsijsnot feasible. However,

since 5D0 is emissive, this transition can be studied via

excitation spectroscopy by monitoring emitted photons

5 7
( D0 I Fz’

scanned through the 5D0 + 7F0

The position and shape of the excitation band will depend

612 nm) while a tunable laser is continuously

transition region (578-580 nm).

on the microenvironment around the Eu3+ ion.

Figure (59) shows the excitation spectra resulting from
the titration of the Eu(III) aqua-ion with tridentate
ligand, dipicolinate (DPA)(98b). Individual narrow peaks
characteristic of Eu3+ (aq), [Eu(DPA)]+, [Eu(DPA)2]’,
and [Eu(DPA)3]3' are apparent at the following respective
frequencies: 17272, 17263, 172u8, and 17231 cm'l. Eu(III)
excitation spectroscopy was used to characterize the in-

dividual metal ion binding sites in calcium-binding pro-

teins such as thermolysin. These various Eu(III) coordination

Figure 59.

137

C02

DRQLW=<:E:

co;

   
 

[EU (CPA 1313-

J

"31"

     

\

‘llEu‘DPAlzl-

"71‘.

l

Excitation spectral profiles of the 7F0 + 500
transition obtained during the course of the
titration of an aqueous EuCl3 solution with

the sodium salt of the dipicolinate anion (DPA).

Eu(III) to ligand ratios are indicated to the
left of each trace.9 9

 

138

environments were further characterized by their excited-
state reciprocal lifetime which can be used to determine

the number of coordinated water molecules in each site.

C. Inter Metal-Ion Energy-Transfer Distance Measurements

If there is a significant overlap between the emission
spectrum of an energy donor D and the absorption spectrum
of an energy acceptor A, energy transfer of a nonradiative

(99)

kind can take place between D and A. F6rster showed
that the efficiency of dipole-dipole energy transfer, E,
is inversely proportional to the sixth power of the D-A

separation, r, and is given by

T and T0 are the excited-state lifetime of the donor in
presence and absence of energy transfer.
r, is the actual donor-acceptor distance, R0, the

critical distance for 50% energy transfer

6 25 K2n-H 6

(R = 8.78 x 10' @J cm

0)
8.78 x 10-25 is the product of fundamental constants
K2, the orientation factor

n, the refractive index of the medium between the

interacting ions

 

139

0, the quantum yield of the donor in the absence of
the acceptor

J, the spectral overlap integral

J = ;E(v)e(v)s‘“dv
fF(v)dv

 

F(v) is the luminescence intensity of the donor at
frequency (cm-l)
8(v) is the molar extinction coefficient of the

acceptor (M'l cm'l)

Ln(III)-transition metal ion and Ln(IIIyLn(III)
energy transfer were applied in distance measurements.
Recently, energy transfer involving Tb(III) to Fe(III) in
transferin(loo) (R0 = 27.1 A, r = 2512 H) and Eu(III) to
Co(II) in galactosyltransferase(101) (R0 = 20 l, r = 18
i 3 K) were reported. In neither case are confirmatory
crystallographic studies available. Horrocks gt_gt.(ll)
used the inter Ln(III) ion energy transfer to measure the
distances between the calcium-binding sites in thermolysin.
Thermolysin has four Ca2+ binding sites, S(l) to S(H),
and one Zn2+ binding active site. The exchange-inert
nature of Ln(III) ions occupying site 8(1) of thermolysin
makes it possible to substitute a luminescent ion (e.g.,
Tb(III) or Eu(III)) at this site and a different absorbing
ion (e.g., Nd(III), Pr(III), Ho(III), or Er(III)) at sites

8(2) and S(H). Nonradiative energy transfer can be

1H0

monitored by measuring the effect of the presence of energy
acceptor ions on the excited-state reciprocal lifetime of
the donor ion. The r values estimated from the energy
transfer efficiencies are in good agreement with the X-

ray distance of 11.7 K between calcium sites 8(1) and S(H),
except for the Tb3+ - Nd3+ results due to some contribution

from a dipole-quadrupole mechanism of energy transfer.

11. Energy Transfer in Rare-Earth Metal Chelates

(102) in 19u2 that B—

It was discovered by Weissman
diketone coordination compounds of tervalent europium,
terbium, and Samarium exhibit unusual luminescence proper-
ties when excited by near ultraviolet light. The compounds
emit visible radiation characteristic of the rare-earth
ions. Thus, intra-Hf electronic transitions, which are
known to originate from levels derived from the electro-
static and spin-orbit interactions among the Hf electrons
within the rare-earth ions, occur whenever the excitation
is carried out in the intense ligand absorption bands.
These ligand bands (molar extinction coefficient ~10“ to
105) are n-n* in nature and are characteristic of the co-
ordinated ligands surrounding the central chelated ion.
Weissman realized that the energy was being pumped into
the electronic system of the complex characterized by the

w-electronic levels of the ligands and was subsequently

migrating to the control chelated rare-earth ion, from

1H1
whence characteristic luminescence (line emission) of the
ion occurred. He designated this process intramolecular

energy transfer.

A. Paths of Energy Migration

 

A schematic energy level diagram for rare earth chelate
possessing low-lying Hf electronic state is given in
Figure (6O)(103). After excitation of a chelate to a
vibrational level of the first excited singlet state
(S0 + Sl)’ the molecule undergoes rapid internal conver-
sion to the lower vibrational levels through interaction
with the solvent matrix. The excited singlet state may
be deactivated by combining radiatively with the ground
state (80 + 81), resulting in molecular fluorescence, or
the molecule may undergo radiationless intersystem cross-
ing from the singlet to the triplet system. Again by
internal conversion the molecule may reach the lowest
triplet state, Tl' From this state, it can then combine
radiatively with the ground state by means of a spin-
forbidden transition (SO + T1) giving rise to a typical
long-lived molecular phosphorescence. Alternatively, the
molecule may undergo a radiationless transition from the
triplet system to a low-lying rare-earth ion state(loua).
The latter states are derived from the Hf electronic con—
figuration of the coordinated trivalent rare—earth ion.

After this indirect excitation by energy transfer, the

Figure 60.

1H2

         
 

Rare Earth

Sunglat Triplet Ion States

511%

ENERGY

Schematic energy level diagram for a rare earth
chelate possessing low-lying Hf electronic states:

._,, radiative transitions;.~qh radiationless
transitions.lo

 

1H3

metal ion may undergo radiative transition to a lower ion
state resulting in a characteristic line emission, or it may
be deactivated via radiationless processes. Direct trans-
fer of energy from the excited singlet state to the low-
lying rare-earth ion states has been shown to be unim-
portant<loua).
Using emission spectrOSCOpy, Crosby gt gt (lOHa-c)
obtained a quantitative measurement of the positions of
the excited singlet, triplet, and rare-earthixnllevels
present in rare-earth chelates. By coorelating the rela—
tive positions of these levels with the observed emission,

an empirical rule for the occurrance of line-like emission

from these chelates was established. "The lowest triplet-

 

state energy level of a complex must be nearly equal to or
must lie above the resonance energy level 'the emitting_level'
of the rare-earth ion". Whenever a triplet state lies
decidedly below a resonance level of a given ion, no emis-
sionflmmlthat ion level is observed. Figure (61)(10Hd)

shows the energy levels of several rare-earth ions, along
with the measured triplet-state levels of various com-

plexes. Eu(III) has two resonance lines, one at 17,257

1 1

cm' and a second at 19,020 cm_ Excitation of Eu(8HQ)3,

tris(8-hydroxyquino1ate)europium(III), with ultraviolet

light results in an emission only from 5DO level and not

from the upper 5Dl level. Crosby gt g; (10Ha,e) inter-

preted this observation as a proof that the energy

 

lHH

 

 

 

 

 

 

 

 

 

25,-3;; ------ ‘B: ------- _-_=.-MA,acety1acetonate chelate
-_--_::—---_-_—-,- -------- “MB,Benzoylacetonate chelate
zo:-=——; ......... m’o..--r.wo,Dibenzoylmethide chelate
_— -'-| .
n __==r:gfu.-_F, ........... mmfih 8-Hydroxyquinolate chelate
e
w- a,
x :fi:::
2: :2—
“ ‘F a .
:2: I0- "7“. . ,H
z 91— I! I!
g g— "I. '1.
3 3:3"Wt *5 -—-W
5,. :B’: TF‘ ———é
—lh —-—-s 4 .__'31
—’I: —; ‘
._—7/2 _f
°‘ 17:; 7F,” 7? mg:
Sm“ Eu” Tb“ D!“
Figure 61. Energy level diagram for rare-earth chelates

(----) lowest triplet level for the complexes:

( ) rare-earth ion level; (T) rare earth
ion resonance level in chelates. Excited
singlet states of the complexes have much higher
energies and are not shown. Likewise only the
lowest tfiiplet-state energy for a complex is
given'.10-d

 

1H5

traveled through the triplet manifold and then down to

the 5D and could not travel uphill to the 5D

O 1'

Three paths have been prOposed for the energy trans-
fer process by which the excitation energy absorbed in the
n-n* electronic bands of the ligands migrates to the reson-

ance Hf electronic levels of the ions. They are:

I. [singlet(Sl)]«-[Triplet(Tl) or a triplet slightly
higher than Tl]«~v[Rare earth ion (RE) level]——p
[Emission](loua)

 

II. [Singlet(Sl)]~«»[Rare earth ion (RE) level]-——r
[Emission](105)
III. [Singlet(Sl)]-v[Rare earth ion (RE) levels higher

than (T1)]~“'[Trip1et (Tl)]**'[Rare earth ion (RE)

level lower than (Tl)]-—v[Emission](lO6a).

The weight of evidence in the literature is in favor of
paths I and III and against path II.

B. Rate of Energy Transfer

Based upon experiments involving intermolecular energy

transfer between benzophenone as a donor and rare-earth

chelates as acceptors,a lower limit of 105 sec"1 for

(106b).

the rate was deduced by El-Sayed and Bhaumik A

maximum upper limit on the transfer would, of course, be

12 13 -1

10 to 10 sec which is the rate of vibrational

 

1H6

relaxation. If the energy transfer occurs from the
singlet state, the transfer has to compete with the singlet-

triplet intersystem crossing which is estimated to be 0,1011

l (1060).

sec- in rare-earth complexes If, as appears more

probable, the mechanism is via the manifold of the triplet

-1

levels, a rate faster than 109 sec is expected.

C. Mechanism of Transfer

The possibility of a pure F6rster type mechanism seems
doubtful because the necessary condition of spectral over-
lap of the emission of the complexed ligand and the absorp-
tion of the metal ion is not always fulfilled. The study
of Eu(III) chelates, shows the efficiency of the energy
transfer to be dependent on the nature of the ligand-
metal bonds. A mechanism was suggested by Crosby gt gt (lOHa)
in which a vibronic coupling of the electronic states of
the rare-earth ion with those of the ligand passed through
an exchange mechanism while preserving the total spin of
the molecule. In the case of europium, as an example, 5D +
3T = 7F + 18.

Whenever the luminescence spectra observed from a series
of lanthanide chelates are too weak or too complex to ensure
accurate measurement of the phosphorescing (triplet) states,

one can employ energy transfer as an aid in bracketing the

energy of these states.

1H7

III. Europium Complex with Salicylidene-DL-Valinate

Schiff Base

 

A. Introduction

 

In many enzymes a proton may assume a critical role in
a chemical transformation at the active site, and a metal
ion or metal complex may serve as a model enzyme by pro-
viding a positive center which promotes a similar function.

Martell(107)

called them model enzymes or "Artificial
Enzymes". In the metal complex the metal ion may provide
both the positive center and through its coordination
requirements considerable stereospecificity. While the
steric arrangements of donor groups around the metal ion
as well as the intensity of metal-ligand interaction can
be varied widely by changing the metal, one must expect
the stereospecificity and adjacent group effects in the
metal coordination sphere generally to differ widely from
the stereospecificity of the enzymatic active site. Model
studies have the advantage of allowing more complete

study of structure-reactivity relationships and of achiev-
ing a better understanding of enzyme function than can be
accomplished by studies of enzymes alone. The structure
and reactivity of a model may be varied by the substitu-
tion of functional groups, charge of metal ion, and change

in the structure of the substrates. Such variations are

either not possible or limited when working with enzymes

1H8

themselves. The use of model systems in the study of
complex enzymatic reactions is well known and particular
interest has been shown in systems set up to explain the
role of vitamin B6 as a catalyst for some reactions
involving amino-acid transformations. The reaction
mechanisms of pyridoxal phosphate coenzymes have been

(108’109) and are

the subject of intensive investigations
now well documented as occurring through a Schiff base
intermediate formed between the vitamin and the amine-
containing substrate (normally an amino acid). It is also
established that a metal is Sometimes an additional co-

(110) and that

factor in enzymatic catalysis by pyridoxal
metals generally enhance the rate of some reactions even
in the absence of enzyme(107). In this instance it has
been proposed that the intermediate involves a coordina-
tion complex between the metal and the Schiff base. A
general theory on the mechanism of these nonenzymatic

(111,112) (1) Formation of

reactions involves four stages:
a Schiff base with subsequent electron displacement from
the a-carbon atom to the electronegative ring nitrogen
through a conjugated system of double bonds, weakening of
the bond to the a-carbon thus results. (2) Release of H+,
COOH+ or R+ from the a-carbon to produce a transitional
Schiff base. (3) Localization of the ring nitrogen lone
pair of electrons onto the a-carbon for racemization and

decarboxylation or onto the formyl carbon for transamina-

tion followed by protonation. (H) Hydrolysis of the

1H9

carbon-nitrogen bond to give products.

(113)

Pullman gt gt. showed that variation in resonance
energy and electronic distribution among the possible
transition states plays an important role in these
catalytic reactions. Change in bond geometries is critical
as well. x-ray-structure determination of two copper com-
plexes with two Schiff bases (pyridoxal—valine and

(“1), Figure (62).

salicylaldehyde-valine) was studied
In (pyridoxylidene-DL-valinato) copper(II) chelate, the
molecule is centered on five—coordinate copper in square-
pyramidal stereochemistry. The phenolic oxygen, imine
nitrogen and carboxylate oxygen atoms of the tridentate
Schiff base occupy three corners of the coordination
square; the remaining in-plane positions being filled by
the heterocyclic nitrogen of a neighboring molecule. Bind-
ing by the hydroxymethyl oxygen of yet another moiety in
the fifth apical site completes the square-pyramid. In
(salicylidene-DL-valinato) COpper(II) chelate, the metal
is four coordinated with planar geometry, one donor posi-
tion being filled by a water molecule. The two chelate rings
are somewhat more planar than in the pyridoxal complex
although a tetrahedral distortion of about the same magni-
tude is still apparent in the arrangement of donors about
the metal.

In aqueous media, Ln(III) ions form (1:2) complexes
with N—hydroxyacetOphenonevaline Schiff base<115). The

stability of these chelates follows the order: La(III) <

 

150

(A)

 

(B)

 

Figure 62. Diagram of Cu(II) complex with: (A) salicyl—
aldehyde-valine Schiff base; (B) PLP- valine
Schiff base.

 

151

Ce(III) < Pr(III) < Nd(III) < Sm(III) < Gd(III) < Tb(III)
< Dy(III) < Ho(III), which is in agreement with lanthanide
contraction.

We here prepared a solid europium chelate with salicyl-
aldehyde Schiff base of DL-valine. The absorption and
emission properties will be discussed in the following

section.

B. Results and Discussion

 

Europium chelate with salicylidene-DL-valine Schiff
base was characterized by elemental analysis, infrared and
optical spectroscopy and its magnetic properties. Com-
parison of the infrared spectra of the complex with that

of the pure ligand shows the following. The imine C=N

1

(amide I) absorption which appears at 1635 cm“ in the

(12) l

was red shifted to 1600 cm' in the metal

1

ligand

complex. The 1375 cm' symmetric stretching absorption

band of the -COO' group in the ligand was red shifted to

l 1

1350 cm' in the complex. The 1225 cm- phenolic C-0

1

stretching of the ligand appears at 1215 cm" in the complex.

The presence of broad water absorption bands prevented ob-
servation of bands in the neighborhood of 3000 cm'l.

In a recent study(11u)

of the x-ray structure of a
(1:1) complex of salicylene—DL-valine with Cu(II), it was
shown that the Schiff base acts as a tridentate ligand

through its phenolate and carboxylate oxygen atoms and

 

 

152

its azomethine nitrogen. The planar geometry of the complex
with one water molecule filling the fourth position shows
that the two chelating rings in the tridentate ligand are
somewhat planar. Someone has to keep in mind that Cu(II)

may force the planar configuration of the ligand through

its complexation. We wererufissuccessful in preparing good
crystals for x-ray measurements cfi' the prepared chelate;
however, according to our infrared and elemental analysis
results as well as considering the chelating nature of our
ligand,we can safely propose an octahedral structure, Figure
(63) (meridian isomer) for our complex. The complex contains
two water molecules. The magnetic behavior of the complex

at low temperatures showed temperature-independent para-
magnetism characteristic of Eu(III) rather than Eu(II)(ll6).
Figure (6H) displays a plot of the reciprocal of the molar
magnetic susceptibility (l/xM) vs T°K. At temperatures
higher than 100°K, XM follows Curie-Weiss law (XM = CM

(T +6)-l); CM and 6 are constants. The slope of the ob-
tained straight line equals l/CM, accordingly CM = 2.375.

Since ”eff = 2.837/3; “eff for our complex equal to H.37

(117)
203 .

, temperature-

B.M., a value of 3.53 B.M. is reported for Eu

Below 100°K, Van Vleck paramagnetism<ll6)

independent paramagnetism, was observed. Eu(III) has Hf6

configuration with 7FO ground state. Although the 7F

multiplet width for Eu(III) is 5360 cm'l; the value of

the interval between the two lowest components 7F0 and 7F1
is 5360/21 = 255 cm-1. At room temperature kT N 200 cm'l,

 

153

 

 

IT
K
O D
O u
xxcx 4%.
. EIO
wlwx
C C \C
/H3H
C

 

 

KEEu(012Hl3NO3)2].

Figure 63.

15H

 

0.24

 

 

 

l l ' l l
0'8 50 l00 I50 200
Termerature PK)

 

Figure 6H. Magnetic behavior of KEEu(C12H13NO3)2]
complex.

 

155

i.e., the transition (7FO + 7Fl)hv = kT, resulting in an
appreciable population of low-lying excited states which
contribute to the paramagnetism at low temperatures.
Room temperature absorption spectra of the Eu-complex
in ethanol and in DMSO are shown in Figure (65). The
absorption band at N350 nm (28,500 cm-l) with molar ab-
sorptivity (c = 11,000) corresponds to a transition to
a charge-transfer state within the coordinated ligand.
Similar absorption band at 350 nm was observed in the
spectra of the ligand in acidic medium which is again a
good clue for the involvement of the azomethine nitrogen
in chelation. In metal chelates of these model Schiff

(118)

bases,it is well known , that coordination through the
azomethine nitrogen produces an absorption band around

365 nm. No absorption bands due to f-f transitions were
observed because of their very low molar absorptivity

(a Z 3)(119).

Figure (66) shows the emission spectra of the Eu
complex in ethanol and DMSO at room temperature and in
ethanol at 77°K. The emission at H50 nm (22,200 cm'l)
is assigned to a fluorescence from the coordinated ligand
at room temperature in both solvents. The 2610 nm (l6,HOO

cm-l) narrow fluorescence is assigned to an f-f transition

within the metal ion. Usually the transition 5D - 7F

O 2
at 216,3HO cm"1 is very intense in europium chelates and
(120,

its fine structure has received considerable attention

121). Salicylaldehyde Schiff bases chelate with Zn(II)

156

 

Absorbance

 

 

l I I T
\U/\
. \
0.6 - “ .1
\
\
l" l ‘
l
l
,_ l ' \ ..
0.4 ‘ ,r \
l / \
/ \
h- / \ d
, \
/ . \
/ \
0.2 - I \ ..
/ \
/
/ \\
I' \
\
\ \
oo- . 7‘

 

L l
300 350 400
Wavelength (nm)

Figure 65. Absorption spectra of dilute solution (3 x 10’5 M)
of K[Eu(C12H13NO3)2] in ethanol (----) and DMSO
) .

(———.

157

 

 

 

 

 

 

 

I I l I | l

3>~
4...
'fi
c 0.
~9- \.
5 ll
a) ,‘ l.
o l
C: ' .
CD \
Q) .
m l.
g l
2 \‘\
L’— .\.J

I’ ----l-

350 450 550 650
Wavelength (nm)
Figure 66. Emission spectra of dilute solutions (3 x 10'5

M) of KEEu(C 2H13NO3) ] in ethanol at room
temperature 1----), e hanol at 77°C ( ----- ),
and DMSO at room temperature ( ).

 

 

 

158

and phosphoresces around (20,500 cm-l) which corresponds

(122)

to the energy of the lowest triplet state in these

chelates. In Figure (67) we constructed an energy level
diagram which shows the important radiative and radiation-
lesswtransitionswithin the complex. Intramolecular energy
transfer occurs from the ligand lowest triplet (as an

5

l or D0 of the metal (as an energy

acceptor) followed by radiative transition (fluorescence)

energy donor) to the 5D

to the 7F. Two necessary conditions are required for
sharp fluorescence from the rare-earth chelates through
indirect excitation of the rare-earth ion via energy trans-
fer mechanism: (1) at least one acceptor level should lie
below the lowest triplet (donor) level of the isolated
ligand, (2) the lifetime of triplet state should not be
too short to achieve a good inversion and not too long to
retain the inversion while waiting for the action to com-
mence.

In water, the metal complex is completely hydrolyzed
to give both absorption and emission spectra characteris-
tic of the free ligand. In DMSO, DMF,and acetonitrite
red emission from the metal ion was observed at room tem-
perature; whereas,in ethanol this emission was only ob-
served at 77°K. This can be due to the difference in
the stability of the metal chelate in these solvents. In
highly coordinating solvents DMSO, DMF, and acetonitrile,
the complex is more stable, maylxsthrough solvent coordina-

tion which increases the coordination number to 8 or 9.

 

159

 

__ D.
_L 50

 

5.7

 
  

3:32:09:— co. 335+ _ ___lI =

653|
7:: 420

.0:
OquLonuO££ 7
an.“
2

 

20-

Arc. 11..

, 02.9.3933.» 1:00:

l0“-
0

EB 335

Energy level diagram for KEEu(C12Hl3NO3)2].

Figure 67.

CHAPTER VI

CONCLUSION AND FUTURE WORK

In the carbonyl compounds, the emission spectra showed
an excited state proton transfer. No ground state proton
transfer occurs in these compounds as evidenced by their
absorption spectra. Both ground-state and excited-state
proton transfer were observed in salicylaldehyde Schiff base
derivatives with n-butylamine and valine. Depending on
the solvent, intramolecular or intermolecular proton trans-
fer was observed. In nonpolar solvents, e.g., 3MP, intra-
molecular proton transfer occurred whereas,in polar sol-
vents,intermolecular proton transfer to or from solvent
molecules was observed. Such observation demonstrates the
higher basicity of the azomethine nitrogen compared with
the carbonyl oxygen in the ground state, and the enhanced
basicity of both functional groups in the excited state.

Salicylaldehyde anion interacts strongly with some
monovalent cations resulting in the formation of chelated-
type complexes. Salts of para-substituted phenols form
contact ion-pairs and solvent-separated pairs depending
on the solvent and on the presence of crown ether molecules.

The spectral blue shift observed by decreasing the cationic

160

161

radius indicate stronger interactions between the ions form-
ing the ion pair and may be a factor in determining which
cation is a better enzyme activator. Evidence for excited-
state dissociation of ion pairs was given. The formation
of these ion pairs may play a key role in the function of
monovalent cations as activators of some PLP-enzymes.

The use of lanthanide ions as probes in biological
systems is becoming interestingly important. Lanthanide
ions, especially Eu(III) and Tb(III),are used extensively
to probe the microenvironment at the active site and to
measure the distances between the different binding sites
in enzymes such as thermolysin. We have prepared a
Eu(III) chelate with a Schiff base model compound. We
have studied~its magnetic and optical properties. The
complex is more likely to have an octahedral configuration.
Our demonstration of energy transfer resulting in a red
emission at 610 nm offers new opportunities in probing the

active site of PLP-enzymes.

Future work in this area is very important. Thus, time-
resolved picosecond studies will undoubtedly help in iden-
tifying excited-state species resulting from proton trans-
fer. The study of sterically hindered model compounds,
which are expected to have slower rates of proton transfer
(higher energy barrier), is particularly interesting. More
attention should be given to the study of possible con-
formers that may exist in equilibrium in the ground state of

a particular species.

 

162

The use of luminescence techniques offer a unique
opportunity to study ion-pair association in the excited
state. Our study opens the door for an extensive study of
ion-association where one of the ions has convenient emission
properties, e.g. fluorenyl anion. The study of time-
resolved spectra of ion pairs is very interesting and im-
portant in order to obtain rate constants of excited-
state sodium ion or potassium ion transfer.

Moreover, the study of the interaction of tryptophanase

 

holoenzyme (tryptophanase-PLP) with Eu(III) is an obviously
important experiment to measure distances between trypto-
phenyl residues and the active site. An experiment where

a tunable dye laser is used as an exciting source to moni-
tor the excitation spectrum of Eu(III) emission and life-
time measurements in the microsecond range of the holo-
enzyme in aqueous media and in the presence of D20 provide
unique information regarding the microscopic environment
(water presence) at the active site. The use of other

lanthanide ions such as Tb(III) should also be explored.

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