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MULTINUCLEAR NMR STUDY OF COMPLEXATION OF SOME
UNIVALENT CATIONS BY CROWN ETHERS IN

NONAQUEOUS SOLVENTS

By

Mojtaba Shamsipur

A DISSERTATION

Submitted to
Michigan State University
in partial fulfillment of the requirements

for the degree of
DOCTOR OF PHILOSOPHY

Department of Chemistry

1979

 

 

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ABSTRACT

MULTINUCLEAR NMR STUDY OF COMPLEXATION OF SOME
UNIVALENT CATIONS BY CROWN ETHERS IN
NONAQUEOUS SOLVENTS

By

Mojtaba Shamsipur

l3C 23Na, 13305, and

Nuclear magnetic resonance of ,

205Tl nuclei were used to study the sodium, potassium,
cesium, and thallium(I) ion complexes with dibenzo—BO-
crown-10 (DBBOClO), dibenzo-ZM—crown-8 (DBZNC8), and dibenzo-
21-crown-7 (DB2lC7) in nitromethane, acetonitrile, acetone,
methanol, dimethylformamide, dimethylsulfoxide, and pyri-
dine solutions.

With the exception of the DB30010-Na+ system, all of
the complexes were formed in 1:1 mole ratios. The presence
of three sodium DB3OClO complexes, Na2(DB3OClO), Na3(DB3OClO)2,
and Na(DB3OClO) in nitromethane and acetonitrile solutions
was deduced from the behavior of the 23Na chemical shift
as a function of DB3OClO/Na+ mole ratio. The NMR data
support the existence of a "wrap around" structure for

the DBBOClO complexes with cesium and potassium ions in

solution.

n3

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3»
a:

a w

. 1 u‘l‘
Q» \~q Shir.
h. n . .f

Mojtaba Shamsipur

The stabilities of the sodium, cesium, and thallium(I)
ion complexes with both DB2HC8 and DB21C7 decrease in the
order Tl+-Crown > Cs+-Crown > Na+-Crown in solvents of low
and medium donicities such as nitromethane, acetonitrile,
acetone, and methanol. The ability of the ligands for
the formation of a three-dimensional "wrap around" complex
with the same cation decreases with decreasing the size of
the ligand, liéi: DB3OC10 > DBZUCB > DB2lC7. In all cases
studied, with the exception of pyridine, there is an ex-
pected inverse relationship between the donor strength of
the solvents and the stability of the complexes. The
extent of the solvent effect on the complex formation

+, and T1+ ion complexes de-

+

constants for the Na+, Cs

+ > Cs

creases in the order Tl+ > Na
The chemical shift of the 133Cs resonance was studied
as a function of the ligand/Cs+ ion mole ratio at various
temperatures in five solvents, iiii: nitromethane, aceto-
nitrile, acetone, methanol, and pyridine. From the result-
ing data AG°, AH°, and AS° values for the complexation
reactions between the cesium ion and DB3OClO, DB2NCB, and
DB21C7 were calculated. It was found that in all cases,
while the stabilities (or the AG° values) of the complexes
are not very sensitive to the solvent, the enthalpy and
the entropy values vary very significantly with the solvent.

In all cases the complexes are enthalpy stabilized but

entropy destabilized. From the results, it seems

Mojtaba Shamsipur

reasonable to assume that the main reason for the negative
entropy of complexation is the decrease in the conforma-
tional entrOpy of the ligands upon the formation of a metal

complex.

+, Cs+, and Tl+ ions by 1,10-

The complexation of Li+, Na
diaza-l8-crown-6 (DA1806) in several nonaqueous solvents
was studied by multinuclear NMR technique. The formation
constants of the resulting 1:1 complexes were calculated
by computer fitting of the mole ratio data. The stabilities
of the complexes decrease in the order DAl8C6-T1+ >
DAl8C6-Li+ > DA1806-Na+ > DAl8C6°Cs+. In order to study
the effect of the substitution of two nitrogen atoms for
the two oxygen atoms in lB—crown-6, the formation constants
were compared with those reported for the l8—crown-6 com-
plexes with the same cations. As expected, the sodium
and the cesium ion complexes are weakened appreciably by
the nitrogen substitution. The sodium and cesium ions
as "hard acids", cannot interact as strongly with the
substituted nitrogen atoms of the ligand as they can with
the oxygen atoms. The effect of the nitrogen substitution
on the lithium and thallium(I) ion complexes is exactly
the opposite, the stabilities of these complexes are
greatly increased. These cations can form partially co-
valent bonds which cause an increase in the strength of

the interaction between the ligand and the cations upon

the nitrogen substitution.

Mojtaba Shamsipur

The kinetics of the complexation reactions of the cesium
ion with DB2lC7, DB2MC8, and DBBOClO in acetone and
methanol were investigated by the temperature dependent
133Cs NMR. The energies of activation for the release
of Cs+ from the CESiUfll complexes decrease with decreasing
donicity of the solvent as expressed by the Gutmann donor
number. They also decrease with the increase in the size
of the ligand. The data show that, first, the transition
state must involve a substantial ionic solvation and,
second, the transition state must be more ordered than
the initial and the final states, iii; the solvated complex

and the solvated cesium ion and the free ligand.

To Nahid

ii

I .

 

ACKNOWLEDGMENTS

The author wishes to express his sincere gratitude to
Professor Alexander I. Popov for his guidance, encourage-
ment, and friendship throughout this study.

Professor Stanley R. Crouch is acknowledged for his
many helpful suggestions as second reader.

The author acknowledges the financial assistance of
the people of Iran, as administered by the Isfahan Uni-
versity of Technology, during the course of this study.
The financial aids of the Department of Chemistry, Michigan
State University, and the National Science Foundation are
also acknowledged.

The help of Mr. Frank Bennis, Mr. Wayne Burkhardt, and
Mr. Tom Clarke in keeping the NMR spectrometers in operat—
ing condition is acknowledged.

Deep appreciation to my wife, Nahid, for her love,
understanding, patience, and encouragement throughout

this study.
To her and to our son, Ali, I dedicate this thesis.

iii

 

 

‘9 ‘
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TABLE OF CONTENTS

Chapter Page

LIST OF TABLES. . . . . . . . . . . . . . . . . . .viii
LIST OF FIGURES . . . . . . . . . . . . . . . . . . xv
LIST OF ABBREVIATIONS . . . . . . . . . . . . . . .xxii

CHAPTER 1. HISTORICAL REVIEW . . . . . . . . . . . 1
1.1. Macrocyclic Crown Ethers. . . . . . . . . 2
1.1.1. Introduction. . . . . . . . . . . 2

1.1.2. Metal Ion Complexes with
Large Crowns. . . . . . . . . . . 3

1.1.3. Thermodynamics of Metal—
Complex Formation in Solution . . 1“

1.1.3.1. Open Chain Ligands. . . 1h
1.1.3.2. Cyclic Ligands. . . . . 21

1.2. Nuclear Magnetic Resonance. . . . . . . . 27
1.2.1. Introduction. . . . . . . . . . . 27
1.2.2. Chemical Shift Measurements . . . 28
1.2.3. Multinuclear NMR Studies of
Complexation of T1+ and Alkali
Ions in Solution. . . . . . . . 33
1.3. Conclusions . . . . . . . . . . . . . . . MO

CHAPTER 2. EXPERIMENTAL PART . . . . . . . . . . . ill

2.1“. Synthesis and Purification of
Ligands . . . . . . . . . . . . . . . . . 42

2.1.1. Synthesis of Dibenzo-30-
crown-100 o o o o o o o o o o o o “2

2.1.2. Purification of Ligands . . . . . H3
2.2. Solvents and Salts. . . . . . . . . . . . All
2.2.1. Solvents. . . . . . . . . . . . . A”

iv

 

 

Chapter

2.3.

2.14.

2.5.
CHAPTER 3.

3.3.

3.“.

3.5.

Page

2.2.2. Salts . . . . . . . . . . . . . . 45
Sample Preparation. . . . . . . . . . . . M6
Instrumental Measurements . . . . . . . . 46
Data Handling . . . . . . . . . . . . . . A8

MULTINUCLEAR NMR STUDY
OF DIBENZO-30-CROWN-10,

DIBENzo-2u-CROWN-8, AND

DIBENZO—21-CROWN—7 COMPLEXES

WITH Na+, K+, Cs+, and Tl+

IONS IN NONAQUEOUS SOLVENTS . . . . . . 50

Introduction. . . . . . . . . . . . . . . 51
Complexation of Na+, K+, and
Cs+ Ions with Dibenzo-30-Crown-10 . . . . 52

3.2.1. DB3OC10 Complexes with
Na+ and K+. . . . . . . . . . . . 52

3.2.2. DBBOCIO Complexes with Cs+. . . . 59

Complexation of Na+, Cs+, and

T1+ Ions with Dibenzo-24-Crown—8. . . . . 67
3.3.1. DB2uc8 Complexes with Na+ . . . . 67
3.3.2. DB2uc8 Complexes with Cs+ . . . . 72

3.3.3. DB2uc8 Complexes With T1+ . . . . 77

3.3.u. Conclusions . . . . . . . . . . . 83
+ + +

Complexation of Na , Cs , and T1

Ions with Dibenzo-2l-Crown-7. . . . . . . 86

3.u.l. DB2107 Complexes with Na+ . . . . 86

3.u.2. DB21C7 Complexes with Cs+ . . . . 91

3.u.3. DB2107 Complexes with T1+ . . . . 97

3.u.u. Conclusions . . . . . . . . . . . 102

Discussion. . . . . . . . . . . . . . . . 105

Chapter Page

CHAPTER A. CESIUM-133 NMR STUDY OF THE
THERMODYNAMICS OF THE COMPLEXA-
TION OF DIBENZO-30-CROWN—10,
DIBENZO-2H—CROWN-8, AND DIBENZO—
21-CROWN—7 WITH CESIUM ION IN

NONAQUEOUS SOLVENTS . . . . . . . . . . 109
4.1. Introduction. . . . . . . . . . . . . . . 110
u.2. DB3OC10 Complexes with 03+. . . . . . . . 111
u.3 DE2uc8 Complexes with Cs+ . . . . . . . . 126
u.u. DB2IC7 Complexes with Cs+ . . . . . . . . 139
u.5. Discussion. . . . . . . . . . . . . . . . 1&9

CHAPTER 5. LITHIUM-7, SODIUM-23, CESIUM-133,
AND THALLIUM—205 NMR STUDY OF Li+,
Na+, Cs+, and Tl+ ION COMPLEXES
WITH 1,10-DIAZA-l8-CROWN-6 IN

VARIOUS NONAQUEOUS SOLVENTS . . . . . . 156
5.1. Introduction. . . . . . . . . . . . . . . 157
5.2. Results . . . . . . . . . . . . . . . . . 158

5.2.1. 1,10—Diaza-18-Crown-6
Complexes with Li+. . . . . . . . 158

5.2.2. 1,10-Diaza—18-Crown-6
Complexes with Na+. . . . . . . . 170

5.2.3. 1,10-Diaza-18-Crown-6
Complexes with 05+. . . . . . . . 173

5.2.U. 1,10-Diaza—l8—Crown-6
Complexes with T1+. . . . . . . . 176
5.3. Discussion. . . . . . . . . . . . . . . . 177

CHAPTER 6. A.STUDY OF DYNAMICS OF CESIUM
ION COMPLEXES WITH DIBENZO-30-
CROWN-10, DIBENZO-2A-CROWN-8,
AND DIBENZO-Zl-CROWN-7 IN
ACETONE AND METHANOL. . . . . . . . . . 183

6.1 Introduction . . . . . . . . . . . . . . . 18A

6 (5.2. Determination and Interpretation
of the Lineshapes . . . . . . . . . . . . 186

vi

Chapter Page

6.3. Results and Discussion. . . . . . . . . . 195
APPENDICES

APPENDIX I - DETERMINATION OF COMPLEX

FORMATION CONSTANTS BY THE NMR TECHNIQUE,
DESCRIPTION OF COMPUTER PROGRAM KINFIT AND
SUBROUTINE EQUATION . . . . . . . . . . . . . . 211

APPENDIX II - DETERMINATION OF COMPLEX
FORMATION CONSTANT WITH ION PAIR FORMA—
TION BY THE NMR METHOD. . . . . . . . . . . . . 216

REFERENCES. . . . . . g . . . . . . . . . . . . . . 223

vii

 

Table

LIST OF TABLES

Experimental Conditions for Metal
Ion NMR . . . . . . .

Thermodynamic Parameters for the
Complexation of Large Crown Ethers
with Cations in Solution. . . .
Thermodynamic Data for Diethylene-
triamine Complexes with Some Metal
(II) Ions in 0.1 M KCl at 25°C.
Thermodynamic Data for Metal (II)-
Polyamine Complexes (Chelate
Effect) . . . . . . .

Thermodynamic Data for Reaction

(2) for Different Metal (II) Ions
Nuclear Properties of Alkali
Elements and Thallium . . . . . . .
Key Solvent Properties and Cor-
rection for Magnetic Susceptibility
on DA-60. . . . . . . . . . . . .
Mole Ratio Study of Dibenzo-30-
Crown—10 Complexes with 0.05 M,
Sodium Tetraphenylborate in Various

Solvents at 30°C. . .

viii

Page

“7

11

13

l7

19

35

“9

53

Table

10

ll

12

13

Page

Mole Ratio Study of Dibenzo-30—

Crown-10 Complexes with 0.005 M

Cs+ Ion in Various Solvents at

30°C. . . . . . . . . . . . . . . . . . . . 60
Formation Constants and the

Limiting Chemical Shifts of

DB3OC10-Cs+ Complexes in Various

Solvents. . . . . . . . . . . . . . . . . . 66
Mole Ratio Study of Dibenzo-2h-

Crown-8 Complexes with 0.025111

Sodium Tetraphenylborate in

Various Solvents at 30°C. . . . . . . . . . 68
Formation Constants and the

Limiting Chemical Shifts of

DBZACB-Na+ Complexes in

Various Solvents. . . . . . . . . . . . . . 71
Mole Ratio Study of Dibenzo-ZM-

Crown-8 Complexes with 0.005 M

Cs+ Ion in Various Solvents

at 30°C . . . . . . . . . . . . . . . . . . 7“
Formation Constants and the

Limiting Chemical Shifts of

+

DB2AC8-Cs Complexes in

‘Various Solvents. . . . . . . . . . . . . . 76

 

 

Table Page

1” Mole Ratio Study of Dibenzo-2H-

Crown-8 Complexes with 0.005 M

T1C10u in Various Solvents at

30°C. . . . . . . . . . . . . . . . . . . . 79
15 Formation Constants and the

Limiting Chemical Shifts of

DBZMCB-Tl+ Complexes in Various

Solvents. . . . . . . . . . . . . . . . . . 82
16 Formation Constants of 1:1 Com-

plexes of Na+, Cs+, and T1+

Ions with Dibenzo-Zu-Crown-8

in Various Solvents . . . . . . . . . . . . 8H
17 Mole Ratio Study of Dibenzo-21—

Crown-7 Complexes with 0.025 M

NaBPhu in Various Solvents at

30°C. . . . . . . . . . . . . . . . . . . . 87
18 Formation Constants and the

Limiting Chemical Shifts of DB21C7

-Na+ Complexes in Various Solvents. . . . . 92
19 Mole Ratio Study of DibenzO-21-

Crown-7 Complexes with 0.005 M

CsSCN in Various Solvents at

30°C. . . . . . . . . . . . . . . . . . . . 9h

20 .Formation Constants and the

Limiting Chemical Shifts of

 

Table Page

DB21C7-Cs+ Complexes in Various

Solvents. . . . . . . . . . . . . . . . . . 96
21 Mole Ratio Study of Dibenzo-21-

Crown-7 Complexes with 0.005 M

T1C10u in Various Solvents at

30°C. . . . . . . . . . . . . . . . . . . . 98
22 Formation Constants and the

Limiting Chemical Shifts of

DB21C7-T1+ Complexes in

Various Solvents. . . . . . . . . . . . . . 101
23 Formation Constants of 1:1 Com-

plexes of Na+, 08+, and T1+ Ions

with DB21C7 in Various Solvents . . . . . . 103
24 Formation Constants of 1:1 Complexes

of Na+, 05+, and T1+ Ions with

DB21C7, DBZ4C8, and DB30010

in Various Solvents at 30°C . . . . . . . . 106
25 Cesiump133 Chemical Shifts

of 0.005 M Cs+ Ion in the

Presence of DB30010 at Various

Temperatures. . . . . . . . . . . . . . . . 113
25 Formation Constants of DBBOCIO

'Cs+ Complex in Nonaqueous Sol—

vents at Different Temperatures . . . . . . 122

xi

Table

27

28

29

30

31

32

33

Thermodynamic Parameters for the

Complexation Of Cs+ Ion by Dibenzo-

30-Crown-10 in Various Solvents
Cesium-133 Chemical Shifts of
0,005 M.CS+ Ion in the Presence
of DB24C8 at Various Temperatures
Formation Constants of DB24C8-Cs+
Complex in Nonaqueous Solvents at
Different Temperatures. . . . . .
Thermodynamic Parameters for the
Complexation of Cs+ Ion by
Dibenzo-24-Crown-8 in Various
Solvents. . . . . . . . . . . . .
Cesium-133 Chemical Shifts of
0.005 Mos+ Ion in the Presence
of DB21C7 at Various Temperatures
Formation Constants of DBZlC7-Cs+
Complex in Nonaqueous Solvents at
Various Temperatures. . . . . .
Thermodynamic Parameters for

the Complexation Of Cs+ Ion

by DBZlC? in Nonaqueous Solvents

at Different Temperatures . . . .

xii

Page

124

127

136

138

140

148

151

Table

34

35

36

37

38

39

40

Entropies of the Complexation of

Cesium Ion by DB3OC10, DB24C8, and
DB21C7 in Various Solvents. . . . . . . .
Mole Ratio Study of 1,10-Diaza-18-
Crown-6 Complex With 0.02 M LiClOu

in Various Solvents at 30°C . . . . . . .
Mole Ratio Study of 1,10-Diaza—
18-Crown-6 Complex with 0.05 M

NaBPhu in Various Solvents at

30°C. . . . . . . . . . . . . . . . . . .
Mole Ratio Study of 1,10-Diaza-
18-Crown-6 Complex with Cs+ Ion

in Various Solvents at 30°C . . . . . . .
Mole Ratio Study of 1,10-Diaza-18-
Crown-6 Complexes with T1+ Ion

in Various Solvents at 30°C . . . .
Formation Constants and the

Limiting Chemical Shifts of
1,10-Diaza-18-Crown-6-Li+

Complexes in Various Solvents

Formation Constants and the

Limiting Chemical Shifts Of
1,10-Diaza-18-Crown—6-Na+

Complexes in Various Solvents . .

xiii

Page

153

159

160

161

163

169

172

Table

41

42

43

44

45

Page

Formation Constants and the

Limiting Chemical Shifts of
1,10-Diaza-18-Crown-6-Cs+

Complexes in Various Solvents . . . . . . . 175
Formation Constants and the

Limiting Chemical Shifts of
1,10-Diaza-18-Crown-6'T1+

Complexes in Various Solvents . . . . . . . .178
Formation Constants of Li+, Cs+,
and T1+ Ion Complexes of 1,10-
Diaza—lB—Crown-6 and 18-Crown-6

in Various Solvents . . . . . . . . . . . . 179
Temperature Dependence of

the Rate Constants for the

Release of Cs+ Ion from

D821C7-Cs+, DB24C8-Cs+, and

DB3OC10-Cs+ Complexes in

Acetone and Methanol. . . . . . . . . . . . 202
Exchange Rates and Thermodynamic

Parameters for Release of Cs+

Ion from Some Large Crown

Complexes in Acetone and

Methanol. . . . . . . . . . . . . . . . . . 206

xiv

Figure

LIST OF FIGURES

Page

Structure of Some Large Crown

Ethers. . . . . . . . . . . . . . . . . . 4
Crystalline Structure of Some Metal

Ion-Crown Complexes. A—(KSCN)2-DB2408,
B—KI-DB30C10, C—(NaSCN)2DB30C10 . . . . . 7
Sodium-23 Chemical Shifts YE: [DB30C10]/

[Na+1 Mole Ratio in Different Solvents.
A-Nitromethane, B-Acetonitrile, C-

Pyridine. . . . . . . . . . . . . . . . . 54
Carbon-13 Chemical Shifts at Various

[Metal IonJ/[DB30C10] Mole Ratios

(MR) in Nitromethane. A-Sodium

Ion at MR=0.0, 1.0, and 2.0; B-

Potassium Ion at MR=0.0, 0.5, and

1.0 . . . . . . . . . . . . . . . . . . . 56
Carbon-13 Chemical Shifts for

the Four Polyether Chain Carbon

Atoms at Various [Na+1/[DB3OC10]

Mole Ratios in Nitromethane . . . . . . . 57
Cesium-133 Chemical Shifts 1s

[DB3OC10]/[Cs+] Mole Ratio in

XV

Figure

10

ll

12

13

Page

Different Solvents. A—Nitromethane,
B-Methanol, C-Acetone, D-Pyridine,
E—Acetonitrile. . . . . . . . . . . . . . 61
Computer Fit of the Cesium-133 Mole

Ratio Data for DB3OC10-Cs+ in

Methanol at 30°C. . . . . . . . . . . . . 65
Sodium—23 Chemical Shifts Mg.

[DB24C8]/[Na+] Mole Ratio in

Different Solvents. . . . . . . . . . . . 70
Cesium—133 Chemical Shifts Kg,

[DBZ4C8j/[Cs+] Mole Ratio in

Different Solvents. . . . . . . . . . . . 75
Thallium-205 Chemical Shifts Kg.
[DB21C81/[T1+] Mole Ratio in Dif-

ferent Solvents . . . . . . . . . . . . . 80
Sodium-23 Chemical Shifts gs.

[DB21C7]/[Na+] Mole Ratio in

Different Solvents. . . . . . . . . . . . 89
Cesium-133 Chemical Shifts Kg.

[DB21C7]/[Cs+] Mole Ratio in

Different Solvents. . . . . . . . . . . . 95
Thallium-205 Chemical Shifts gs.
[DBZlC7]/[T1+] Mole Ratio

in Different Solvents . . . . . . . . . . 99

xvi

Figure

14

15

16

17

18

19

2O

21

Page

Cesium-133 Chemical Shifts 1g.

[DB3OC10]/[Cs+] Mole Ratio in

Nitromethane at Different Tem-

peratures . . . . . . . . . . . . . . . . 116
Cesium-133 Chemical Shifts Kg.

[DB3OC10]/[Cs+] Mole Ratio in

Acetonitrile at Different Tem-

peratures . . . . . . . . . . . . . . . . 117
Cesium-133 Chemical Shifts YE-

[DB3OClO]/[Cs+] Mole Ratio in

Acetone at Different Temperatures . . . . 118
Cesium-133 Chemical Shifts gs.

[DB3OClO]/[Cs+] Mole Ratio in

Methanolat Different Temperatures . . . . 119
Cesium-133 Chemical Shifts gs.

[DB3OClO]/[Cs+] Mole Ratio in

Pyridine at Different Temperatures. . . . 120
Van't Hoff Plots for Complexation of

Cs+

Ion by Dibenzo-30-Crown-10 in

Various Solvents. . . . . . . . . . . . . 123
Cesium-133 Chemical Shifts lg.

[DB2408]/[Cs+] Mole Ratio in Nitro-

methane at Different Temperatures . . . . 130

Cesium-133 Chemical Shifts 1g.
[DB24C8]/[Cs+] Mole Ratio in

xvii

Figure Page

Acetonitrile at Different Tem-

peratures . . . . . . . . . . . . . . . . 131
22 Cesium-133 Chemical Shifts gs.

[DB24C8]/[Cs+] Mole Ratio in

Acetone at Different Temperatures . . . . 132
23 Cesium—133 Chemical Shifts lg.

[0324081/[Csfj Mole Ratio in

Methanol at Different Temperatures. . . . 133
24 Cesium-133 Chemical Shifts Kg.

[DBZ4C8]/[Cs+] Mole Ratio in

Pyridine at Different Temperatures. . . . 134
25 Van't Hoff Plots for Complexation

of Cs+ Ion by Dibenzo-24—Crown-8

in Various Solvents . . . . . . . . . . . 137
26 Cesium-133 Chemical Shifts vs.

[DB21C7]/[Cs+] Mole Ratio in

Nitromethane at Different Tem-

peratures . . . . . . . . . . . . . . . . 143
27 Cesium-133 Chemical Shifts YE-

[DB2107]/[Cs+] Mole Ratio in

Acetonitrile at Different Tem-

peratures . . . . . . . . . . . . . . . . 144
28 Cesium-133 Chemical Shifts YE-

[DBZlC71/[Cs+] Mole Ratio in

Acetone at Different Temperatures . . . . 145

xviii

 

Figure Page

29 Cesium-133 Chemical Shifts 1s.

[DB21C7]/[Cs+] Mole Ratio in

Methanol at Different Temperatures. . . . 146
30 Cesium-133 Chemical Shifts 1g.

[DBZlC7]/[Cs+] Mole Ratio in

Pyridine at Different Temperatures. . . . 147
31 Van't Hoff Plots for Complexation

of Cs+ Ion by Dibenzo-21-Crown-7

in Various Solvents . . . . . . . . . . . 150
32 Lithium-7 Chemical Shifts vs.

[DA18C6J/[Li+] Mole Ratio in

Different Solvents. . . . . . . . . . . . 16“
33 Sodium-23 Chemical Shifts gs.

[DA18C6J/[Na+] Mole Ratio in

Different Solvents. . . . . . . . . . . . 165
34 Cesium-133 Chemical Shifts YE-

[DA18C6J/[Cs+] Mole Ratio in

Different Solvents. . . . ._. . . . . . . 166
35 Thallium-205 Chemical Shifts gs.

[DA18C61/[T1+] Mole Ratio in

Different Solvents. . . . . . . . . . . . 167
36 Cesium-133 NMR Spectra of 0.02 M

CsSCN, 0.01 M_DBZlC7 Solution in

Acetone at Various Temperatures . . . . . 196

xix

Figure Page

37 Cesium-133 NMR Spectra of 0.04 M

CsSCN, 0.02 M DB2107 Solution in

Methanol at Various Temperatures. . . . . 197
38 Cesium-133 NMR Spectra of 0.02 M

CsSCN, 0.01 M DB24C8 Solution in

Acetone at Various Temperatures . . . . . 198
39 Cesium-133 NMR Spectra of 0.02 M_

CsSCN, 0.01 M DB2408 Solution in

Methanol at Various Temperatures. . . . . 199
40 Cesium-133 NMR Spectra of 0.01 M

CsSCN, 0.005 M_DB30C10 Solution

in Acetone at Various Temperatures. . . . 200
41 Cesium-133 NMR Spectra of 0.01 M

CsSCN, 0.005 M DB3OC10 Solution

in Methanol at Various Temperatures . . . 201
42 Arrhenius plots of log kb 172. l/T for

the Release of Cs+ Ion in Acetone

and Methanol with Large Crown

Ethers. A-DB30C10-Cs+ in Methanol,

B-DB2408-Cs+ in Methanol, C-DB3OClo-Cs+

in Acetone, D-DB21C7'Cs+ in Meth-

anol, E-DBZ4C8'Cs+ in Acetone,

F-DB21C7°CS+ in Acetone . . . . . . . . . 206

XX

Figure Page

43 Entropy Profile for the Release

+

'Of Cs Ion from DB3OC10-Cs+ Complex

in Acetone. . . . . . . . . . . . . . . . 209

xxi

Ac

AN

DMF
DMSO
MeOH
NM

PC

Py
TMG
TMO
DB3OC10
DB24C8
DB2107
DA18C6

LIST OF ABBREVIATIONS

Acetone
Acetonitrile
Dimethylformamide
Dimethylsulfoxide
Methanol
Nitromethane
Propylene carbonate
Pyridine
Tetramethylguanidine
Trimethylene oxide
Dibenzo-30-crown-10
Dibenzo-24-crown-8
Dibenzo-21-crown-7

1,10-Diaza-l8-crown-6

xxii

CHAPTER 1

HISTORICAL REVIEW

1.1. MACROCYCLIC CROWN ETHERS

1.1.1. Introduction

Since Pedersen's discovery of macrocyclic polyether
(crown) compounds capable of forming stable complexes with
the alkali ions (1) the studies of these ligands and their
complexes have become a very popular field of research.

A large number of such complexes have been isolated in
crystalline form and many solution studies have been car-
ried out (2,3). Most of the investigations in solution
have been done in water, methanol and/or their mixtures;
such studies in other nonaqueous solutions are quite
sparse.

A variety of physicochemical techniques have been
used for such investigations (3); the choice of favorite
technique being dictated by the systems studied as well
as by the particular expertise of the investigators.

One of the most interesting characteristics of the
macrocyclic compounds is their ability to selectively bind
certain cations in solution in the presence of others.

The selectivity and stability of crown ether complexes
has been reported to be dependent on several important
parameters characteristic of the ligand, the cation and
the reaction medium. These parameters are: the relative
sizes Of cation and ligand cavity (4,5), the type and

the number of donor atoms in the ring (4,6,7,8), substi-

tution on the macrocyclic ring (9-12), type and charge

2

 

 

of cation (13) and the solvent effect (14-16).

Several useful review articles are available on the
study of macrocyclic polyethers and their complexes (2,3,
17-19). In this thesis only the studies on the complexes
of large crowns and the thermodynamics Of the metal-

complex formation in solution will be reviewed.

1.1.2. Metal Ion Complexation with Large Crowns

Despite the very interesting properties of large
crowns (143;) larger than 18-crown-6), not much attention
has been focused on the study of the complexation of metal
ions with these ligands. In comparison with approximately
three hundred scientific papers on small size crowns, and
in particular on different 18—crown-6 derivatives, such
investigations on large crown complexes are sparse. The
structure of some of these ligands are shown in Figure 1.

Pedersen (1) was the first to report the isolation of
cesium ion complexes with dibenzo-24-crown-8 and dibenzo-
30-crown-10 in methanol solution and to study the effect
of complexation on the ultraviolet spectra of large
crowns. He likewise (20) studied complexation of alkali
and alkali earth cations with 21-crown—7 and 24-crown-8
derivatives using ultraviolet spectroscopy and solvent

extraction methods.

The effect of dicyclohexy1-24-crown—8 on the ionic

permeability of natural membranes such as HK and LK sheep

 

DIBENZO—Ql-CROWN—7

Figure 1. Structure of Some Large Crown Ethers.

red cells has been studied by Tosteson (21). He showed
that this ligand notably increases potassium permeability
more than sodium permeability, while the behavior is
opposite in the case of cyclohexyl-15-crown-5. It has
been reported (22) that cyclic polyethers exhibit specific
influences on cation transport in rat liver mitochondria.
Dicyclohexy1—2l-crown-7 was most effective in the presence
of potassium or rubidium ions. Dicycloxyl-BO-crown-lo

was displayed striking specifity in that it was very
active with rubidium present and much less active with
potassium.

The stoichiometry of crystalline complexes of alkali
cations with large crowns has been investigated (23-25),
and the crystalline structure of some of these complexes
has been reported (26—29). The stoichiometry of such
compounds depends on the relative sizes of the cation and
the cavity size of the macrocycle, the flexibility of the
crown molecule, and the nature of the anion and of the
solvent (17). For dibenzo-24-crown-8—K+ and dibenzo-30-
crown-lo—Na+ systems the ratio of two metal ions to one
crown ether has been found so far (24).

Mercer and Truter (27) have reported the structure
of the 2:1 (metal to ligand) complex between KSCN salt
and dibenzo-24-crown-8, isolated from methanol solution.
According to them, each one of the K+ ions, located in—

side the ring, are bound to the five oxygen atoms of the

ligand, two nitrogen atoms of bridging thiocyanate groups,
and two carbon atoms of a benzene ring (Figure 2a). In
another publication (25). Truter and co-workers have
shown that when potassium tetraphenylborate was used
instead Of potassium thiocyanate, only the 1:1 complex
was formed since the tetraphenylborate anion cannot be
coordinated with the metal ion.

The crystalline structure of the 1:1 complex between
KI and dibenzo-30-crown—10 has been studied by Bush and
Truter (26). They showed that the potassium ion is com-
pletely located inside the cavity which is created by
the twisting of the ligand around the cation, all the
oxygen atoms are coordinated with the cation to form a
so called "wrap around" complex (Figure 2b).

Owen and Truter (29) recently have reported the crystal—
line structure of the 2:1 (metal to ligand) complex of
sodium isothiocyanate with dibenzo-30-crown-1 (29),
determined from x-ray diffraction measurements. Each
ligand complex with two Na+ ions, and each cation is co-
ordinated to six oxygen atoms in the ring and also to
one isothiocyanate anion, through the nitrogen atom.

The ligand is in an extended conformation, twisted so
that the molecule assumes a figure eight configuration,
with the Na+ ions at the center of two loops (Figure 2c).

The stability constants for the 1:1 complexes of large

crown ethers (i.e., 21, 24, 30, and 60-membered rings)

 

‘73"
’fih 1W? x

‘ a 1 '
@53c 75);} U“... ‘

C

Figure 2. Crystalline Structure of Some Metal Ion-Crown
Complexes. A-(KSCN)2-DB24C8, B-KI-DBBOClO,
C-(NaSCN)2DB30010.

with alkali ions has been determined by Frensdorff (4)

in water and methanol solutions using potentiometric
measurements with cation-selective electrodes. He point-
ed Out that the selectivity order is governed by the rela-
tive sizes of the cation and the cavity of the ligand.
From the increase of K+ stability constants between 24-
crown-8 and 30-crown-10 it was suggested that a "wrap
around" complexing might occur with polyethers, such as
that observed for K+-valinomycin (30).

Rechnitz and Eyal (31) constructed liquid membrane
electrodes with dibenzo—30-crown-10 in the membrane phase
(nitrobenzene) to measure the crown's electrochemical
activity, and to determine the crown-metal ion complex
formation constant. The potentiometric selectivity ratio
for the crown ether in nitrobenzene for the Rb+, 03+,

Na+, and NH”+ ions with respect to K+ ion has been de-
termined, and formation constants of rubidium and potassium
ion complexes with dibenzo-30-crown-1O have been reported.
The AG° and AH° parameters for complex formation of di-

benzo-30-crown-10 and Li+, Na+, K+, Rb+

, Cs+, NHu+ and
Tl+ in methanol has been determined by Chock (32), who
has studied the ultraviolet spectra of the ligand at
different temperatures. With the exception of the am-
monium ion, a good correlation between the stability

constant of the complex and the size of metal ion was

observed.

In an attempt to elucidate the solution structure Of

a number of crown ethers, Live and Chan (33) carried out
a careful measurement on the 1H and 13C NMR spectra of
the ligands and their metal complexes in water, water-
acetone, acetone, and chloroform solutions. From the
data, the authors have concluded that the structure of

K+ dibenzo-30-crown-10 in solution is similar to that in
crystalline state (26) where K+ ion is completely surrounded
by the ligand to form a "wrap around" complex. 0n the
basis of the proton NMR spectrum of Na+ dibenzo-30-crown-
10 complex in acetone, which did not resemble the one for
+

K , they proposed a twisted configuration for the complex

in solution. Their 13C NMR results support these conclu-

sions. The stability constants for 1:1 complexes of the

+

ligand with Na+, K , and Cs+ in acetone have been deter-

1

mined from the H NMR data.

Izatt and co-workers (34) have reported thermodynamic

+, Rb+, and Cs+ ions

parameters for complexation Of Na+, K
with dibenzo-24-crown-8 and dibenzo-27-crown-9 in methanol-
water mixtures which have determined calorimetrically.

They pointed out that the entropy changes for a 1:1 re-
action of a given cation become more negative with in-
creased size of the polyether ring which seem to indicate
an increased conformational change of the ligand upon

complex formation.

The stability of thallium and alkali metal ion

C)
.9

w-
:i

I‘ir

 

10

complexes with dibenzo-24-crown-8, dibenzo-30—crown-10
and their benzo group derivatives in acetonitrite and
methanol has been investigated by Mittal and co-workers
using polarographic methods (35). The nature of the sol-
vent was found to affect the stability constants of the
complexes.

Recently, Lehn and co-workers (36) reported formation
of stable and selective complexes between guadinium and
imidazolium ions and chiral hexacarboxylate-27-crown-9
in aqueous solution. The stability constants of the
complexes were measured potentiometrically by a competition
method using a NHu+-selective electrode. They have con-
cluded that the presence of anionic carboxylate groups
in the ligand markedly increases the stability of the
crown complexes and that the selectivity of the complexa-
tion arises from central discrimination of the macrocyclic
ring.

The literature reported thermodynamic values for the
complexation reaction of metal ions with large crowns

are listed in Table 1.

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11

 

 

 

 

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l2

 

 

 

 

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mm IIII IIII os.s woooaod z< +om
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n...

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We
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Ir ..
LIV

14

1.1.3. Thermodynamics of Metal-complex Formation in

Solution

1.1.3.1. Open Chain Ligands - The thermodynamics
of metal ion complexes of a number of polyamine systems
have been investigated calorimetrically by Paoletti g£_a1.
(37-40) in aqueous solutions. The ligand studies were
diethylene triamine (dien) (30), n=l; triethylene tetra-
mine (trien) (38), n=2; tetraethylenepentamine (tetren)
(39), n=3; and N,N,N',N'-tetra-(2-aminoethyl)-ethylene-
diamine (Penten) (40). Some of the thermodynamic data
for these complexes are listed in Table 2.

H

2

Polyamines' Structure

It is seen that there is a tendency for successive
enthalpy changes to become more exothermic, while the
entropy changes for 1:1 complexes are positive, those
for the second step reactions are negative. The overall
entropy changes, however, are positive in all cases and
the complexes are both enthalpy and entropy stabilized.

The nature of the chelate effect can be examined in
more detail using the available thermodynamic data for

the polyamine complexes for the following reaction.

15

Table 2. Thermodynamic Data for Diethylenetriamine Com-
plexes with Some Metal (II) Ions in 0.1 M KCl

 

 

 

at 25°C.
2+ AG AH AS
M (Kcal/mole) (Kcal/mole) (cal/mole deg)

Co2++dien -10.90 —8.15 9.0
Codien2++dien -8.00 -10.25 -7.5
Niz++dien -14.45 -11.85 8.5
Nidien2++dien -10.90 -13.45 —8.5
Cu2++dien -21.55 -l8.00 12.0
Cudien2++dien -7.10 -8.15 -3.5
Zn2++dien -12.00 -6.45 18.5

anien2++dien -7.50 -10.15 -9.0

 

 

i.

I;

:x

AV

16

2+

M(NI-IBM

+ L z ML2+ + n NH3 (1)
Some of the results obtained are shown in Table 3. It
can be seen that the complexes are stabilized with
respect to the corresponding ammonia complexes not only
by enthalpy changes but also by the large positive en-
tropy changes for the reactions. From the data, it is
clear that the large negative free energy changes can be
attributed to the positive entropy changes accompanying
the release of increasing numbers of ammonia molecules.
Extensive calorimetric studies have been made Of the
formation of metal complexes with aminopolycarboxalate
ions such as ethylendiamine tetraacetic acid and its
homolagues (42-48). Schwarzenbach (41) has shown that
in ligands with more than two donor atoms the presence
of a nitrogen atom meets the requirements for the forma-
tion of relatively strain-free chelate rings. These
ligands have been found to be stabilized by positive
entropy changes resulting from charge neutralization and
subsequent solvent release from the solvation shells of
interacting ions. Thermodynamics data are obtained from
those available for methyliminodiacetate (mIDA) (42)

and ethylendiamine tetraacetate homologues:

17

 

 

 

Table 3. Thermodynamic Data for Metal (II)-Polyamine
Complexes (Chelate Effect).
AS

2+ AG AH (cal/mole)
M L (Kcal/mole) (Kcal/mole) deg) Ref.

2+
Ni dien -5.33 -1.35 13.4 37
Cu2+ dien -7.24 -3.00 14.2 37
Ni2+ trien -8.09 0 27.1 38
Cu2+ trien -lO.20 -l.55 29.0 38
N12+ tetren -12.10 -1.40 35.9 39
Cu2+ tetren -14.65 —1.75 43.3 39
Ni2+ penten -15.44 +1.35 56.0 40

 

 

 

 

o. .. .
in «a I? v . w r ..
I... o .I.
I I C
to F. In.

 

 

l8

OOCCH2 l/CH2COO
_ N<CH2)nN\\ -
OOCCH2 CHZCOO

in which n=2 for EDTA (42); n=3 for TMTA (44);
n=4 for TETA (44); n=5 for PETA (45); n=6 for HDTA
(45); and n=8 for ODTA (45).

The existence of thermodynamic data on mIDA and EDTA
metal complexes (mIDA constitutes one half of EDTA) makes
it possible to examine the chelate effect in more detail

considering the following reaction.

4- 2-

M(mIDA)§’ + EDTA : MEDTA + 2mIDA2‘ (2)
The data for the above reaction is given in Table 4. The
positive values for AS clearly show the chelate effect
which reflects the increase in the number of solute
particles in the reaction (2). The endothermic enthalpy

2+ and Mn2+

changes for some of the metal ions such as Ca
are probably due to the strain involved in fitting another
chelate ring around the metal ions. The chelate effect
is smallest with these metal ions.

Another interesting point to consider is to investi-
gate the variation of the chelate effect with n, the

number of methylene groups in the central chain of EDTA

homologues. The following reaction should be considered:

19

Table 4. Thermodynamic Data for Reaction (2) for Dif-
ferent Metal (II) Ions.

 

 

 

M2+ (Kcalgmole) (Kcalgmole) (cal/m8Ie deg)
Mg2+ -3.65 1.31 17.1
Ca2+ -5.27 -4.62 2.7
Mn2+ -5.68 -u.33 3.1
002+ -3.21 1.28 15.3
Ni2+ -3.57 0.10 12.5
Cu2+ -1.17 3.96 17.4
Zn2+ -3.23 0.98 14.6
Cd2+ -5.27 -1.78 11.9

 

 

2O

m(mIDA)§’ + L“' : ML2' + 2mIDA2'

(3)
Anderegg in a series of publications (41-45) has reported
the thermodynamic data for several metal (II) ion com-
plexes with EDTA and its homologues. He has pointed out
that an increase in the number of methylene groups causes
the stability of ML to decrease rapidly. The entropy
changes were not appreciably different, indicating that
the coordination in the complexes is approximately the
same.

The factors involved in determining enthalpies and
entropies of metal chelate formation in aqueous solution

(100) are given below:

Enthalpy Effects

Variation of bond strength with electronegatives of
metal ions and ligand donor atoms.

Ligand field effects.

Steric and electrostatic repulsions between ligand
donor groups in the complex.

Enthalpy effects related to the conformation of

the uncoordinated ligand.

Other coulombic forces involved in chelate ring

formation.

21

Entropy Effects

Number of chelate rings.

Size of the chelate ring.

Changes of solvation on complex formation.
Arrangement Of chelate rings.

Entropy variations in uncoordinated ligands.

Effects resulting from differences in configurational

entropies of the ligand in complex compounds.

1.1.3.2. Cyclic Ligands - It has been shown that
cyclic ligands form much more stable complexes than their
corresponding open chain analogs with metal ions. This
enhancement in stability was called the "macrocyclic
effect" first by Cabbiness and Margerum (46) who Observed
an appreciable enhancement in the stability of copper
complex with a tetramine macrocyclic ligand compared to
similar non-cyclic tetramine ligands. According to them
the important factors responsible for this effect are the
solvation and the configuration of the ligand.

On the basis of the results Obtained from a study Of
copper (II)-tetramine complexes, Paoletti e£_a1. (47)
proposed that both entropy and enthalpy terms contribute
to the macrocyclic effect, in contrast to the chelate ef-

fect which results only from an entropy factor. In a

later publication on the study of Ni(II)-tetramine complexes

22

in water, Hinz and Margerum (48) concluded that the en-
hanced stability with the cyclic vs. the open-chain ligand
is almost entirely due to a more favorable enthalpy change.
They have pointed out that the release of water from the
metal ion and the ligand results in a positive entropy
contribution because the number of independent particles
would increase, but this positive contribution is offset

by a negative contribution due to the loss of configurational
entropy of the ligand upon complexation. The noncyclic
ligand would be expected to undergo a much larger loss

of configurational entropy than the cyclic ligand. Dei
and Gori (49) also emphasized the enthalpy stabilization
to explain the macrocyclic effect in a study of AH of Cu
(II)-tetramine complexes in water.

Kodoma and Kimura (50,51) studied polarographically
the complexation of copper (II) with 1,4,7,10-tetraza-
cyclododecane in water. They found that the stability
constant for the 1:1 complex is more than lou-fold greater
than that for corresponding open-chain tetraamine. Despite
the unfavorable enthalpy term, the large increase in the
stability of the copper complex with the cyclic ligand was
entirely attributed to the favorable entropy term. The
same authors (52) determined thermodynamic values for the
complexation of 18-crown-6 and tetraglyme with Pb2+
and T1+ ions using polarographic measurements. Again,

the stabilities of macrocyclic complexes were much greater

23

than those of tetraglyme complexes, and this macrocyclic
effect was explained in terms of entropy. Frensdorff

(4) has also compared the stability Of Na+ and K+ complexes
with 18-crown-6 and pentaglyme in methanol. He has noted

a 103 to 10” enhancement of the stability constant of
cyclic ligand complexes.

Recently Arnud-neu g£_§1. (52) have reported results
of calorimetric and radiocrystallographic studies of
Cu (II) and Pb (II) complexes in aqueous solution with
1-Oxa-4,13,dithia-7,lO-diaza-cyclopentadecane and its
corresponding open-chain homolog. According to the above
authors macrocyclic effect is both enthalpy and entropy
dependent.

One Of the first thermodynamic studies on the metal
ion-crown complexes by calorimetric titration has been
done by Izatt and co-workers (53,54) who determined the
stability, the enthalpy, and the entropy values for the
complexation reaction between different metal ions and
dicyclohexyl-lB-crown-6 in aqueous solution. The alkali
metal ion stability order was found to be identical with
the permeability order for these metal ions with the anti-
biotics valinomycin and monactin. All complexes were
enthalpy stabilized but entropy destabilized. Thermo-
dynamic properties Of alkali complexes of various carrier
antibiotics in methanol and ethanol by Simon g£,a1. (55-

57). They used a computerized microcalorimeter for the

24

thermodynamic measurements.

Schori and Jagur-Grodzinski (58) have studied the
thermodynamics of the complexation of dibenzo-l8-crown-6
and its derivatives with sodium ion in dimethylformamide
and dimethoxyethane solutions at different temperatures
by electrical conductance measurements. In dimethoxy—
ethane as solvent, the complexation reaction was found
to be both enthalpy and entropy driven. Hogen-Esch and
Smid (59) also used conductometric measurements to study
the thermodynamics of the dissociation of fluorenyl salts
and their complexes with dimethyldibenzo-18-crown-6 in
tetrahydrofuran and tetrahydropyran. The complexes were
found to be entrOpy destabilized in both solvents.

Izatt and co-workers (60) have reported Log K, AH,

and AS values for the 1:1 complexes of Na+, K+, Rb+, Cs+,

+ 2+ 2+ 2+

Ag+, T1 , NHK, CH NH+, Ba , Ca2+, Sr2+, Pb

3
ions with 15-crown-5, 18-crown-5, 18—crown-6 and dif-

, and Hg

ferent isomers of dicyclohexo-lB-crown-6 in water which
were determined by calorimetric titration. For most of
the complexation reactions they studied both the entropy
and the enthalpy were negative. The authors could not
find any reproducible trends in AH or AS among the reac-
tions studied to be able to explain the macrocyclic
effect.

Izatt t 1. (61) synthesized new crown ethers by

introducing a pyridine ring in l8-crown-6 capable of

25

forming stable complexes with many metal ions. They
studied the thermodynamics of complexation of these

2+ ions in methanol using

ligands with Na+, K+. As+ and Ba
calorimetric measurements. The ligands were found to
form unusually stable complexes with the above ions in
which the entropy term favors complexation of the ligands
over l8-crown-6 (more negative AS) but the enthalpy term
does not (more positive AH). The same authors (62)
synthesized two ligands analogous to 18-crown-6 having
carbonyl groups available for cation complexation as does
valinomycin. They determined log K, AH, and TAS values

+, and Ba2+

for the reaction of these ligands with Na+, K
ions in methanol. The decreased stabilities of the cation
complexes with these ligands, compared with those with
l8-Crown—6, were found to be due to less exothermic AH
values, but more favorable TAS values.

Mei SE El- (63,64) studied the thermodynamics of

+ ion with 18-crown-6 in various non-

complexation of Cs
aqueous solvents by cesium—133 NMR. The behavior of the
133Cs chemical shift as a function of l8-crown-6/Cs+

mole ratio indicates a two step reaction, first formation
of a stable 1:1 complex followed by the addition of a
second molecule of the ligand to give a 2:1 sandwich com-
pound. It was found that the stability of the 2:1 complex

in pyridine increases with decreasing temperature and

that the complex is enthalpy stabilized but entropy

26

destabilized.

Izatt gt,§;. (65) determined calorimetrically the free
energy, enthalpy, and entropy of complexation of ammonium
and substituted ammonium cations with l8-crown-6 again in
methanol solution. The complexes with secondary and ter-
tiary ammonium ions were found to be much weaker than those
with the primary amines. The binding of all these ions
involves hydrogen bonding with ether oxygens in addition
to the ion-dipole forces. The thermodynamics of 18-crown-
6 complexes with arenediazonium and anilinium salts in
methanol have also been studied by the same authors (66).

Izatt 32,21. (67) have compared the stabilities
of several sulfur derivatives of 9-crown-3, 12-crown-4,
15-crown-5, 18—Crown-6, and 24-crown—8 complexes with Ag+,
Tl+, Pb2+, and Hg2+ ions with their polyether analogs.
Values of log K, AH, and AS have been reported for the
different complexes in water or water-methanol sol. Many
of cyclic thio ethers were found to form 1:2 (cation to

2+ ions but only 1:1

b2+

ligand) complexes with Ag+ and Hg
complexes were found for T1+ and P ions.

Despite existence of much thermodynamic data on the
various complexes with cyclic polyethers and their non-
cyclic analogs, the source of the macrocycle effect has
not been completely defined, and further investigation

seems to be needed to have a general understanding of

this effect.

27

1.2. NUCLEAR MAGNETIC RESONANCE
1.2.1. Introduction

Since the discovery of nuclear magnetic resonance
(NMR) spectroscopy in the 1940's (68,69) this technique
has had a spectacular growth and a wide application to
many chemical problems. Because of its steady develop-
ment, both theoretical and instrumental, nuclear magnetic
resonance spectroscopy has now reached such an advanced
stage that it is a nearly indispensable tool with chemists.

In recent years the use of nuclear magnetic resonance
has become quite popular in studies of electrolyte solu-
tions.- This is because Of the presence of extremely rapid
and generally random molecular motions in electrolyte
solutions which can result in narrow resonance lines even
for nuclei with quadropole moment. Proton NMR is useful
for the studies of ionic interactions as well as for the
determination of solvation numbers. Resonance frequencies
Of metal ions are very sensitive probes of the immediate
chemical environment of metal cation and therefore can
detect very weak ion-ion, ion-ligand, and ion solvent
interactions. Despite this tremendous sensitivity to en-
vironment, metal ions give a generally weak resonance
signal which necessitates special instrumentation for

measurement Of the interactions.

28

1.2.2. Chemical Shift Measurements

Only nuclei possessing angular momentum P (nuclear spin
I) are suitable probes for nuclear magnetic resonance.
The angular momentum imparts a magnetic moment u = 7P
to these nuclei. The ratio Y between the magnetic moment
and the angular momentum is designated as the magneto-
gyric ratio.

When the spinning nuclei are placed in a magnetic field,
Ho’ their spin axes precess about the field direction.
By increasing the strength of the magnetic field, the nuclei
cannot be forced to become aligned with the field direc-
tion, but they only precess faster. They can be made to
flip over, however, by applying a second, much weaker
magnetic field, H1, at right angles to H0, and causing this
second field to rotate at exactly the precession frequency
V0' The magnetic vector of a radio-frequency field supplied
by a transmitter coil with its axis perpendicular to H0,
is used for this purpose conveniently. When the frequency
of this field equals v0, the nuclei flip over and induce
a voltage in a receiver coil placed at right angles to both
H0 and the transmitter coil. This voltage can be ampli—
field and recorded conveniently. Alternatively, both
coils can be combined in one and the flipping of the nuclei
can be observed as an absorption Of energy from the radio—

frequency field.

29

According to quantum mechanics there are 2I + 1 pos-
sible orientations for a magnetic nucleus in an external

field of strength H The energy level separations are

0.
given by:

where h is Plank's constant divided by 2N and u is the
magnetic moment. The transitions occur between the levels
by irradiation of the Larner frequency, equal to the energy
separation of the nuclear spin states. Actually, resonance
occurs at the field experienced by the nucleus, H, which

is different from HO, given by

H = H°(l-O)

where a is known as screening (or shielding) constant.
According to Ramsey equation (70-72) the screening
constant, C, is the sum of various diamagnetic (shielding)

and paramagnetic (deshielding) contributions:

U=Ud+0p

where Up and 0d are the paramagnetic and diamagnetic com-

ponents, respectively. He expressed these terms theo-

retically from perturbation theory as follows:

 

 

r
e K K K
o= {<12 I}? |w>+
2mc2 O K r3 O
K
2(E -E )'1[< lzi X |z 2K] >1
0 m. ll’0 me wm ‘3 YO
m K K rK
where
e : electronic charge
n1: mass of the electron
c : velocity of light
EK : angular momentum of the Kth electron
th electron from the Kth

rK : radial distance Of the K

electron from the origin at the nucleus.

The diamagnetic component, Gd, depends on ground state
electronic wave functions and is a function of the sym-
metry of the electronic distribution and the density of
circulating electrons. The magnitude of the paramagnetic
contribution, up, is zero for ions with spherically sym—
metrical S states but it is substantial for atoms involved
in chemical bonding. It is determined by several factors.
(1) The inverse of the energy separations 1E between ground
and excited electronic states of the molecule. (ii) The
relative electron densities in the various p, d, and
higher states involved in bonding, i;g;, upon the degree

of asymmetry in electron distribution near the nucleus.

31

(iii) The value of <1/r3>, the average inverse cube dis-
tance from the nucleus to the orbitals concerned. Usually
the contributions to ad and CD for a nucleus are considered
only for the electrons immediately neighboring, or local
to, that nucleus. More distant electrons give rise to

long range effects on both op and 9d which are large but
cancel to make only a small net contribution to 0. Gen-
erally, downfield shifts are referred to as paramagnetic
and upfield shifts as diamagnetic.

Kondo and Yamashita (73) have proposed the theory of
paramagnetic interaction. They suggested that the para-
magnetic shift of cations and anions in alkali halide
crystals is due to the short range repulsive forces between
the closed shell of the ions. These forces can excite p
orbital electrons of the alkali nuclei to the higher states,
so that the net result would be a decrease in the shield-
ing of the nucleus.

The success of the Kondo-Yamashita theory in inter-
preting chemical shifts in solids suggested that it may
also provide some way for interpretation of the chemical
shifts in solution. In this case, however, the problem
is more complex. In solids the relative positions and
distances of separation of the ions are known, but in the
solution the environment Of the nucleus will vary randomly
with time because of the diffusion of the ions and solvent

molecules through the solution and the observed chemical

32

shift will result from an average of many instantaneous
values.

Deverell and Richards (74) applied Kondo-Yamashita
theory to provide a qualitative interpretation of the
cation chemical shifts in aqueous solutions. They sug-
gested that at infinite dilution, where the only inter-
actions present are between the ion and water molecules,

the contribution to the paramagnetic chemical shift is

°=— 2_1_ 0l-0
an 16“ (r3>np A A ion-water

where a is the fine-structure constant, <l is the

-—o
r3 np
average over the outer p orbitals Of the central ion,
O ..

A is the mean excitation energy, and Alon-water is an ap
proximate sum Of the overlap integrals of the orbitals of
the central ion and surrounding water molecules.

By increasing the concentration of the solution the
interactions between the ions during the collisions will

also contribute to the chemical shift. The chemical shift

at concentration of c can be expressed as:

=_ 21 .1. (c) (‘3)
0 16a <;3°np A I:[tion-water'mion-ion:I

33

(c)

A(c)
ion-water

ion-ion represent the ion-water and

where A and
ion-ion, respectively. Ikenberry and Das introduced a
more exact equation by including the effects of overlap
and charge transfer covalency. The magnitude of para-

magnetic screening for an alkali nucleus is proportional

to <3L> ' 1. Since the <3;> and l-both increase with
1.3 np K r3 A
increasing atomic number (75), the magnitude and the range

'I-

of 0 increases from Li+ to Cs ions. Therefore, the range

P
of chemical shift varies from about 10 ppm for Li+ to several

hundred ppm for Cs+ ion.

1.2.3. Multinuclear NMR Studies of the Complexation of

T1+ and Alkali Ions in Solution

During the past decade the use of multinuclear NMR
for the studies of the thermodynamics and kinetics of
reactions in solutions has been expanded very rapidly.

In particular, multinuclear NMR has been widely used to
study the behavior of alkali ions in solutions.

Since the advent Of new macrocyclic ligands, such as
crowns, discovered by Pedersen (l), and cryptands, dis-
covered by Lehn (76), capable Of forming strong complexes
with alkali ions, the studies of the alkali ion complexes
with these ligands have become a very popular field of
research. Among a variety of physicochemical techniques
used for such investigations, alkali metal NMR has been

shown to be very popular.

34

The nuclear properties of alkali elements as well as
of thallium are shown in Table 5. All alkali nuclei have
at least one isotope with non-zero spin. In all cases I
is equal or greater than 1/2 and, therefore, the nuclei
have a quadropole moment. Thus it should be expected that
the alkali resonances would have broad lines. In prac-
tice, however, the natural line width of 23Na and specially
of 7Li, 133Cs and 205T1 nuclei are quite narrow so that in
most cases chemical shifts can be measured precisely.
Relative sensitivities are generally adequate for all of
them except 39K. Thallium-205 is shown to be an ideally
suited NMR probe for potassium in biological systems.

The chemical properties and ionic radii (1.44 A and 1.54 A)
Of K+ and Tl+ are similar so that T1+ can replace K+ in
several enzymes without loss Of activity.

Lithium-7 NMR has been used for determining formation
constants of lithium complexes with pentamethylentetrazole
in nitromethane (77). It was found that lithium ion forms
a fairly strong complex with a convulsant tetrazole in
nitromethane.

Lithium ion complexes with cryptands C222, C221 and C211
in water and in several non-aqueous solvents have been
studied by Cahen 92,3l. using 7Li NMR technique (78).

They showed that the first two ligands form weak 1:1 com-
plexes with Li+ ion in solvents of low donicity such as

nitromethane. On the other hand, cryptand 211 was found

 

 

35

mmH.o N\H w=.os mam.:m Hamom

 

NIOH x 35.: m\~ ooa 3mm.» .a..omm._H
uwa.o m\m m.m~ omm.mH nmwm
:IoH x mo.m _ m\m mo.mm oom.m Mam
NIOH x pm.m m\m ooa www.ma ome
:mm.o m\m nm.m: mam.mm HAN
oaofim unopmcoo on ma :HQm ANV mmsomx msoaosz
on o>HumHom mpa>aufimcom oocoocsn< mo.:H no Aumzv
Annapmz mocoooonm mzz

 

 

.ESHHHmnB oco mucoEon HmeH< mo mofiunoqonm nmoaosz .m manoe

36

to form much more stable complexes and two 7Li resonances
(corresponding to the free and the complexed Li+) were
observed for solutions containing excess of the Li+ ion.
The resonance of the Li+ ion inside the cryptand cavity
was found to be completely independent of the solvent in-
dicating that the ligand completely insulates the cation
from the solvent.

The kinetics of the complexation reaction of the Li+
ion with cryptand C211 in water and several non-aqueous
solvents have been investigated by temperature—dependent
7Li NMR (79). The activation energy for the release of
Li+ from the complex was found to be larger in solvents
with higher Gutmann donor number. The exchange rates and
thermodynamic parameters of lithium cryptate exchange in
various solvents were determined from the 7Li NMR tempera-
ture-dependent data.

Hourdakis and Popov (80) have used 7Li, 23Na, and 133Cs
NMR to study alkali complexes with cryptand C222—dilactam
in various solvents. They found out that the complexing
ability of the dilactam is similar to, but weaker than,
that of the cryptand C222. Recently, Smetana and Popov (81)
have studied complexes of Li+ ion with several crowns in
various solvents using 7Li NMR.

Sodium 23 NMR measurements have been obtained on many
antibiotic ionophores in chloroform and methanol solutions

(82). In all cases, addition of ionophores to the sodium

37

ion solution broadens the 23Na resonance lines. Despite
the similar nature of the complexes, the 23Na chemical
shifts were found to be very different for different anti-
biotics. The complexation of Na+ ion with pentamethylene
tetrazole in nitromethane has also been studied by 23Na
NMR (83).

Addition of the crowns, such as 18-crown—6 derivatives,
to the sodium salt solutions of various solvents (84,85)
has been shown to result in a very appreciable broadening
Of the sodium-23 resonance so that the resonance line
could not be detected. This is because most crown ethers
tend to form two-dimensional complexes with the alkali
ions which could distort the spherically symmetrical
electric field around the solvated sodium ion and, there-
fore, broaden the 23Na resonance line.

GrandJean 22.21: (86) have determined the thermo-
dynamic parameters for complexation of a heptadentate non-
cyclic crown ether with Na+ ion in pyridine by 23Na-NMR
spectroscopy. Using 23Na resonance line broadening measure-
ments, they have determined the relaxation rate, and by
measurement of relaxation rate as a function of composition
have calculated complex formation constants at different
temperatures. From the resulted thermodynamic parameters
they have concluded that the complex has a "wrap around"
structure.

Sodium-23 NMR has been extensively used to study the

38

exchange kinetics of Na+ ion with crowns (87,88) and cryp-
tand (89,91) in different solvents. Shchori g£_al,
(87,88) have investigated the kinetics of Na+ complexes

of dicyclohexyl-lB-crown-6 and dibenzo-l8-crown-6 and its
derivatives in various solvents. The life times of free
and complexed sodium ion and the pseudo first-order rate
constant for the decomplexation reaction have been found
from line shape analysis as a function of temperature.
Different substituent groups on the ligands had a sig-
nificant effect on the decomplexation reaction.

Dye and co-workers (89,91) Obtained two 23Na resonance
signals for Na+-C222 cryptate solutions with the excess
of the sodium salt in various solvents. One resonance
corresponds to the sodium ion inside the cryptand cavity,
and the other corresponds to the uncomplexed solvated
sodium ion. The rate constants, activation energies and
thermodynamic parameters for the decomplexation reaction
were obtained from line shape analysis Of the 23Na NMR
temperature-dependent data.

Shih and Popov (92) studied complexation reaction
between K+ ion and several crowns and cryptands in various
non-aqueous solvents by 39K NMR spectroscopy. They found
evidence for formation of an inclusive complex between K+
ion cryptand C222 but an exclusive one for K+-C221 cryptate
in solution. The K+-18-crown-6 complexes were found to

be quite stable in non-aqueous solvents. The lS-crown-S's

39

were found to form both 1:1 and 2:1 sandwich complexes
with K+ in all non-aqueous solvents used.

Shporer and Luz (93) studied the longitudinal relaxa-
tion time T1 of the potassium-39 nucleus as a function of
temperature in methanol solutions in the presence of di-
benzo-18-crown-6 using 39K-NMR technique. The rate of the
decomplexation reaction and the activation energy for the
reaction were then calculated from the resulting data.

Popov and co-workers (63,96,94-96) have obtained
interesting results from a cesium-133 NMR study of both
the kinetics and thermodynamics of crowns and cryptands
complexes with Cs+ ion in non-aqueous solvents. From
13303 chemical shift measurement as a function of ligand
to metal ion mole ratio, they have obtained evidence of
a two-step complexation reaction between Cs+ ion and 18—
crown-6's ligands (63,64,94). The formation of a 1:1
complex is followed by the addition of a second molecule
of crown to give a 2:1 sandwich complex. The thermodynamic
and kinetic parameters for the two step reaction of Cs+
ion with 18-crown—6 in pyridine have been obtained from
the mole ratio data at different temperatures (64).

The chemical shift of the 1330s resonance was studied
as a function of cryptand C222 to Cs+ at various tempera-
tures in different solvents (95). The formation constants
for Cs+-C222 cryptate were then determined from the

resulting data. The limiting chemical shifts for Cs+-0222

40

complex in dimethylformamide, propylenecarbonate and acetone

were found to approach to the same value at about -100°C.

From these results the authors suggested that the complex

is exclusive at higher temperatures but inclusive at lower

temperatures. The thermodynamic parameters were calculated

by analysis of the data Obtained at different temperatures.
Thallium-205 NMR has been used to study the metal ion

binding to biological macromolecules such as pyruvate

kinase (97). It also has been used to investigate com-

plexation of T1+ ion with crowns (98) and antibiotics (99).

1.3. CONCLUSIONS

From the above discussion, it is evident that multi-
nuclear NMR provides a very powerful tool for studies of
complexation reactions in solution. The full potential
of this method has not been realized; however, with con-
tinuous improvement in NMR instrumentation its increasing
usefulness for different kinds of studies can be predicted.
Information on the thermodynamics and kinetics of complexa-
tion reaction can also be obtained by this method. The
subject of this thesis is a multinuclear NMR study of thermo-
dynamics and kinetics of large crowns in various non-

aqueous solutions.

CHAPTER 2

EXPERIMENTAL PART

41

2.1. SYNTHESIS AND PURIFICATIONS OF LIGANDS
2.1.1. Synthesis of Dibenzo-30-Crown-10

The method for synthesis of dibenzo-30-crown-10 was
essentially based on Pedersen's (1) published method for
the synthesis of benzo-lS-crown-S. By reacting catechol
with l,ll—dichloro-3,6,9-trioxaundecane both benzo-15-
crown—5 and dibenzo-30-crown-10 are formed. The main
product of the reaction is benzo-15-crown-5. The procedure
is slightly modified for separation of dibenzo-30-crown-10
from benzo-15—crown-5.

Catechol (88 g, 0.8 mole) was dissolved in 800 ml Of
dried n-butanol containing 0.8 mole of KOH (53 g, 85%
pure) in a 2.5 liter three necked flask, equipped with a
thermometer, a dropping funnel, and a condenser. The
equipment was under a nitrogen atmosphere. The mixture
was heated to reflux, when solution became muddy-white-
greenish in color.

l,1l-Dich1oro-3,6,9-trioxaundecane (93 g, 0.4 mole),
prepared according to Pedersen's method (1), was then added
dropwise over four hours with continued heating, and the
mixture was further heated for an additional three hours.
The mixture was then cooled to room temperature, and 0.8
mole of KOH (53 g, 85% pure) was added. The mixture was
again heated to reflux. Another 0.4 mole (93 g) of 1,11-

dichloro-3,6,9-trioxaundecane was then added dropwise

42

43

over four hours. The mixture was refluxed for about 24
hours, while the color of the solution changes to dark
brown, and then was cooled and evaporated under vacuum in
a rotO-vac apparatus.

After cooling the mixture, a solution of 20 ml of
concentrated HCl in 150 ml water and 500 ml of chloroform
was added. The chloroform layer (lower layer) was then
separated and was washed three times with 400 ml of each
Of saturated NaCl-50% NaOH solutions, and once with 400 m1
of saturated NaCl solution (separation of phases was slow).
The chloroform solution was then dried Over Nazsou.

The chloroform was then evaporated using a rote-vac.

A soft solid was then obtained which was suspended in a
warm (80—90°C) n-heptane solution in a continuous liquid-
liquid extractor for several days. After cooling the
yellow n-heptane extract to about 40-50°C, two layers

were obtained. The upper layer containing benzo-lS-crown-
5 was separated. The lower layer was recrystalized from
pure acetone to give white plate crystals of dibenzo-30-
crown-10 with a melting point of 106.5°C. The yield of
reaction for benzo-lS-crown-S and dibenzo-30-crown-10

were found to be 37% and 3%, respectively.

2.1.2. Purifications of Ligands

Some dibenzo-BO-crown-lo was Obtained through the

courtesy of DuPont company. The ligand was recrystallized

44

from acetone and vacuum dried. Dibenzo-2l—crown-7 and
dibenzo-24-crown-8 (Parish Chemical Company) were re-
crystallized from n-heptane and dried under vacuum for
three days. The melting points of well defined crystals
were found to be 107 and 103°C, respectively, which are
the same as the reported values (17). 1,10-Diaza-18-
crown-6 (Merck Company) was recrystallized from reagent
grade n-heptane and dried under vacuum for several days.
Benzo-lS-crown-S was recrystallized from pure n-heptane

and vacuum dried for 72 hours.

2.2. SOLVENTS AND SALTS
2.2.1. Solvents

Reagent grade acetone (Mallinkrodt) was refluxed over
calcium sulfate, fractionally distilled, and dried over
4A Linde molecular sieves. Spectrophotometric grade nitro-
methane (Aldrich) was refluxed over night over phosphorus
pentoxide, fractionally distilled, and dried over activated
molecular sieves for 24 hours. Dimethylsulfoxide (Fisher)
was dried over Linde 4A molecular sieves and vacuum dis-
tilled. Dimethylformamide (Fisher) was vacuum distilled
over phosphorus pentoxide. Propylene carbonate (Aldrich)
was dried for 48 hours over Linde 4A molecular sieves
followed by vacuum distillation. Acetonitrile (Mallin-

krodt) was refluxed over calcium hydride, fractionally

45

distilled and dried over molecular sieves. Trimethylene
oxide (Aldrich) was dried over activated molecular sieves
for 48 hours. Tetramethylguanidine (Eastman) was refluxed
over granulated barium oxide and fractionally distilled.

The molecular sieves used were activated by heating them
at 600°C for 12 hours under a nitrogen atmosphere. Analyses
for water in solvents were carried out with an automatic
Karl Fischer Aquatest (Photovolt Corp.) titrator. The
water content Of the sOlvents after drying was found to

be less than 100 ppm.

2.2.2. Salts

Lithium perchlorate (Fisher) was dried at 190°C for
several days. Sodium tetraphenylborate (J. T. Baker)
was dried under vacuum at 60°C for 72 hours. Potassium
hexafluorophosphate (Pfaltz & Bauer) was purified by
recrystallization from water and dried under the vacuum
at 110°C for 72 hours. Cesium thiocyanate (Pfaltz & Bauer)
was recrystallized from reagent grade methanol and dried
under vacuum for several days. Cesium tetraphenylborate
was prepared by mixing tetrahydrofuran solution of sodium
tetraphenylborate with equimolar amount Of concentrated
aqueous solution of cesium chloride (Ventron Alfa Product).
The resulting fine white precipitate was collected, washed
several times to remove any adhering sodium salt and

dried under vacuum at 70°C for 72 hours. Thallium (I)

46

perchlorate (K & K) was recrystallized from water and

dried at 110°C for 24 hours.

2.3. SAMPLE PREPARATION

All solutions were prepared in a dry-box under a nitro-
gen atmosphere to maintain their water content at the
lowest possible level. Dilute solutions of the salts were
prepared by appropriate dilution of a stock solution. The
ligand solutions were prepared by proper dilution of a stock
solution (if the ligand was soluble enough) or by weighing
out in the desired amount into a 2 ml volumetric flask

followed by dilution with the solvent or the solution.

2.4. INSTRUMENTAL MEASUREMENTS

Lithium-7, sodium-23, cesium-133, and thallium-205
NMR measurements were carried out on a modified Varian
Associate DA-60 spectrometer equipped with a wide-band
probe capable of multinuclear operation (101), Operating at
a field of 14.09 kgauss in a pulsed Fourier transform
mode. The spectrometer is equipped with an external proton
lock to maintain the field stability. The NMR spectrometer
is interfaced to a Nicolet 1083 computer for pulse genera-
tion, data collection, and data treatment. A previously
described program (102) was used for pulse generation and

collection of the resultant free induction decay (FID)

47

signal. Data treatment was-performed using the Nicolet
FT-NMRD program (NIC-80/S-7202-D) (103). The experimental

conditions are given in Table I.

Table I. Experimental Conditions for Metal Ion NMR.

 

 

 

Resonance External Reference
Nucleus Frequency (MHz) Solution
7L1 23.32 4.0 M LiClOLl in H20
23Na 15.87 3.0 M NaCl in H20
133Cs 7.87 0.5 M CsBr in H20
205Tl 34.61 0.3 M T1NO3 in 320

 

 

All the chemical shifts for sodium-23, cesium-133,
and thallium-205 reported in this thesis are referred to
infinitely dilute aqueous Na+, Cs+, and Tl+ solutions,
and the chemical shifts for lithium-7 are referred to a
4.0 M LiClOu aqueous solution. Downfield (paramagnetic)
chemical shifts from the reference are indicated as ne a-
tlzg, In order to keep the chemical shift of external
reference constant, an insulated reference tube (64) was
used in the measurements of cesium-133 Chemical shifts as
a function of temperature.

The reported chemical shifts are also corrected for

the differences in the bulk diamagnetic susceptibility

48

between the sample (nonaqueous) and the reference (aqueous)
solutions according to the Live and Chan (104) equation for

non-superconducting spectrometers

sample)

V

_ 2n ref
6corr " 6obs + I? (XV - X

where xgef and Xsample are the volume susceptibility of
the reference and the sample solutions, respectively.
The magnitude of correCtions, calculated on the basis of
the published susceptibilities (105), and the physical
properties (106,107) for the solvents used in this study
are shown in Table 6.

Carbon-13 NMR spectra were obtained on a Varian CRT-20
spectrometer operating at a field of 18.68 kgauss in a
pulsed Fourier transform mode. Acetone was used as an
external reference and D20 was used to lock the system.
All carbon-l3 chemical shifts are reported with respect

to TMS.

2.5. DATA HANDLING

The complex formation constants were obtained by
fitting the chemical shift-mole ratio data to appropriate
equations (which will be discussed in detail later) using
the least squares program KINFIT (108) on a CDC-6500
computer. A linear least squares program was used to

obtain enthalpies and entropies.

49

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mom.oI oom.o o.om s.om ooaeoewooaacooeHo
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mam.oI ooa.o o.sH s.om osoooo<
“EaavomI<o AmoaxxIv mnonasz pcmpmcoo pcm>Hom
so coapooppoo n:flflwmmhwmmuwsm hocoa ofinuooamfim

 

 

.omI<Q so :pHHHpHuaoomsm capocwoz pom coauoonnoo ocm mofiunoaonm

uco>aom mom .m manna

CHAPTER 3

MULTINUCLEAR NMR STUDY OF DIBENZO—30-CROWN-10,
DIBENZO-24-CROWN-8, AND DIBENZO-21-CROWN—7

+

COMPLEXES WITH Na+, K , 08+, and T1+ IONS IN

NONAQUEOUS SOLVENTS

50

3.1. INTRODUCTION

Previous studies in our laboratories (109—112) and
elsewhere (113—117) have shown that the nuclear magnetic
resonance Of thallium and alkali nuclei Offers a very
sensitive technique for the studies of changes in the im-
mediate chemical environment of the thallium and alkali
ions in solution. The chemical shifts and line width of
the resonances can given information about ion-ion, ion—
solvent, and ion-ligand interactions. During the past
decade alkali metal NMR has been used extensively to study
the thermodynamics and kinetics of the complexation reac-
tion between alkali metal ions and crowns and cryptands
(109-118).

Among the crown ethers, large molecules such as dibenzo-
30-crown-10 and dibenzo-24-crown—8 have some interesting
properties. These are very flexible molecules with enough
oxygen atoms in the ring so that they can twist around a
metal ion with a suitable size to envelope it completely
and form a three-dimensional "wrap around" complex (26).

Alkali complexes of dibenzo-30-crown-10, dibenzo-24-
crown-8, and dibenzo-21-crown-7 have been studied by
potentiometry (4,31), polarography (35), calorimetry (34),
spectrometry (32), and proton and carbon-13 NMR (33) in
different solvents. The purpose of the study described

+, 03+, and Tl+

in this chapter was to investigate Na+, K
ion complexes of the above mentioned ligands in a number

of nonaqueous solutions by the multinuclear NMR technique.
51

52

3.2. COMPLEXATION OF Na+, K+, AND CS+ IONS WITH DIBENZO-

30-CROWN-10
3.2.1. 0330010 Complexes with Na+ and K+

Sodium-23 chemical shifts were determined as a function
of dibenzo-30-crown-lO/sodium ion mole ratios in nitro-
methane, acetonitrile, and pyridine solutions. The result-
ing data are given in Table 7 and the mole ratio plots are
shown in Figure 3. In all cases only one population-
average resonance was observed indicating that the ex—
change of the metal ion between the bulk solution and the
complex is faster than the NMR time scales. In the case

of pyridine solutions, the addition of the ligand to the

Na+ solution produces a gradual diamagnetic shift Of sodium-
23 resonance which begins to level off at a mole ratio of
about 1, which indicates the formation Of a 1:1 complex of
the sodium ion with the ligand.

On the other hand, in the case of nitromethane and
acetonitrile solutions, the chemical shift 1g. mole ratio
plots show three distinct inflection points at the ligand/
metal mole ratios of about 0.5, 0.7, and 1 indicating the
formation of three complexes with the respective stoi-
chiometries DB30010-2Na+, 2DB30C10-3Na+, and DB3OClO-Na+.
The synthesis and isolation of crystalline 1:2 and 1:1
(ligand to metal ion) complexes of sodium tetraphenyl—

borate and dibenzo—30-crown-10 have been previously

53

 

 

 

 

 

 

msa 0m.0H m=.m

00H 0H.0H m0.m

mm: no.0: mm.H

mm: mm.0a 0a.:

m0 m0.0 mm.H 00H m0.0H 00.0

mm No.0 0m.a mm: 00.0: 00.0

0mm ma.» 0H.m mm 00.0 0a.: 00m mo.HH 0s.0

mom mm.s e0.H 00 se.0 00.0 H00 HA.HH 0s.0

0mm 00.0 00.: mo oa.0 00.0 mam H0.HH 00.0

m0m ms.0 0H.H 0s 00.A H0.0 we: 00.HH 00.0

00H m0.0 00.0 mm ss.s as.0 00m 0m.0 sm.0

Hod mH.m mw.o a» mo.m Hm.o 23H H~.m mm.o

00H Hm.m 00.0 as 0H.0 00.0 mm: ss.0H mm.0

AHA ss.a 0m.0 m0 00.s 0m.0 0H: m0.HH mm.0

mm ms.0 0H.0 mm sm.s 0H.0 00 Hm.NH 0H.0

a: ss.0I 00.0 HH m0.e 00.0 0: 00.0: 00.0

AN:VN\H>0 Asaave +:\o AN:VN\H>< Asaave +o2\o Asmvmxaao Asaave +o2\o
. as 2: :z
I .ooom no muco>aom macano> CH ouononamcoca

Iosooa season :00.0 can: aoonano 0HIcsos0I0mIoscooao 0o soaom capo: oHo: .s oases

54

 

 

 

 

 

IS:
A
Io - . ‘r i
8 8
(ppm) M
I
C
5 ..
O _
1
1 l I I l
(15 L0 L5 2£> 25

[DB SOCIO]/[No+]

Figure 3. Sodium-23 Chemical Shifts Kg. [DB3OClO]/[Na+]
Mole Ratio in Different Solvents. A-Nitromethane,

B-Acetonitrile, C-Pyridine.

55

reported by Truter and coworkers (25). These three species
were isolated from acetonitrile solutions in the form of
well defined crystals. The melting points of 1:1 and 1:2
(ligand to metal ion) complexes were found to be very close
to the reported values. The crystalline structure of the
third species which we supposed to be 2:3 complex were
determined. The results showed that it is not the 2:3
species, but instead it is the 1:2 complex associated with
four molecules of acetbnitrile of crystallization. It
seems that 2:3 complex is quite unstable and exists only

in solution.

It has been shown previously that carbon—l3 chemical
shifts Of the carbons in the ether region of cyclic poly-
ethers are sensitive to the conformational change of the
ligand upon complexation (119). In order to get further
information about dibenzo-30-crown-10 interactions with the
sodium ion, we studied the chemical shift of the polyether
chain carbons as a function of sodium and potassium ion
concentrations relative to the concentration of the ligand.
The results, obtained in nitromethane solutions, are shown
in Figures 4 and 5.

In the case of the DB3OCIO-K+ ion system. the addition
of potassium hexafluorophosphate to a DB3OC10 solution
results in a gradual coalescence of the four resonances
and only one signal is obtained at equimolar concentrations

of the potassium ion and the ligand. This behavior seems

56

 

R: 25,,

All.
11

, l
21) -quJ LO

1 [J I
73 7| 69 67 73716967

8

 

 

 

 

 

 

Figure 4. Carbon-13 Chemical Shifts at Various [Metal IonJ/
[DB3OC10] Mole Ratios (MR) in Nitromethane.
A-Sodium Ion at MR=0.0, 1.0, and 2.0; B—Potassium
Ion at MR-0.0, 0.5, and 1.0.

56

 

5:5

 

, 1
2;) L] Aaequ I!)

J l l l l l
73 7| 69 67 73 7l6967
A B

 

 

 

 

 

 

Figure u. Carbon-13 Chemical Shifts at Various [Metal IonJ/
[DB30C10] Mole Ratios (MB) in Nitromethane.
A-Sodium Ion at MR=0.0, 1.0, and 2.0; B—Potassium
Ion at MR-0.0, 0.5, and 1.0.

57

72-

 

 

 

’OC3

C4

< ~—IC2
69-
68'-

CI

 

l l l l I l
(15 L0 L5 2£> 25 30

[No+]/ [083000]

Figure 5. Carbon-13 Chemical Shifts for the Four Poly-
ether Chain Carbon Atoms at Various [Na+]/
[DB3OClO] Mole Ratios in Nitromethane.

58

to indicate an essentially equal interaction between the

ten oxygens of the polyether ring and the potassium ion.

Naturally, such equal interactions is only possible if

in solutions the ligand is "wrapped around" the cation as
postulated by Live and Chan (33) in solution and shown by
Truter and coworkers (26) for the solid complex.

Considerably different behavior is observed in the
case of the sodium ion. The details are given in Figure
5. Between mole ratios of 0.0 and 2.0 all four carbons
behave quite differently. While the initial addition of
the sodium ion results in a chemical shift of carbons 2, 3,
and H, the chemical shift of carbon 1 remains unaffected.
0n the other hand, between mole ratio of 1.0 and 2.0 the
chemical shifts of carbons 2 and 3 remains constant while
carbon N and especially carbon 1, show a significant down-
field shift. Beyond mole ratio of 2.0, the resonance fre-
quencies are constant.

The results show that the addition of the sodium ion
to DB3OC10 results in at least two conformational changes
of the ligand molecule following the formation of DB3OC10°Na+
and DB3OC10'2Na+ complexes. No evidence for the forma-
tion of the 2DB30010-3Na+ complex was observed; however,
it is to be expected that the 23Na chemical shift is a
much more sensitive probe of the sodium ion interaction
than the 13C chemical shift. Once again, these data

support the conclusions of Live and Chan (33) that the

59

DB30ClO'Na+ complex has a different configuration from the
DB3OC10 complex with potassium ion (and presumably with

2+ ions).

the Cs+ and Ba
An attempt was made to calculate the formation con-
stants of DB3OC10 complexes with the potassium ion from
the variation of the carbon-13 chemical shift as a func-
tion of K+/ligand mole ratio. It was found that in nitro-

methane and acetonitrile solutions log Kf was greater than

5, while in acetone solutions log Kf = “.3 i 0.1.

3.2.2. 0830010 Complexes with Cs+

The variation of cesium-133 chemical shift as a func-
tion of the ligand/Cs+ mole ratio in nitromethane, aceto-
nitrile, acetone, methanol, and pyridine solutions at 30°C
are shown in Figure 6 (the data are given in Table 8).

It is seen that the shift is diamagnetic in acetonitrile
solutions and paramagnetic in all others. The shifts begin
to level off at a mole ratio of about 1, indicating the
formation of a 1:1 complex. It is interesting to note

that the limiting chemical shifts for the complexed cesium
tend to approach each other indicating that in the complex
the cation is largely insulated from the solution and,

once more, confirming the "wrap around" structure.

The formation constants for the DB3OClO°Cs+ complex
in different solvents were determined from the variations

of the 133Cs chemical shifts with the ligand/Cs+ mole

60

 

 

 

 

 

 

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me.HH mH.m mm.m oo.m mm.mH mo.m
m:.HH es.fi 3H.ea ee.a mm.mm mm.H :m.w om.H oo.mH ow.H
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mm.HH NN.H Hm.ea em.H mm.mm mm.H NH.: mH.H oe.mH m=.H
NH.NH 0H.H mm.ma mm.H Hm.mm mm.H me.m NH.H o=.mH om.H
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+

61

 

50

3O

8
(ppm)

 

 

 

 

I

0% To _ IJS— 2.0 2.5
[DBBOCIO]/[Cs*]

Figure 6. Cesium-133 Chemical Shifts E. [DBBOClO]/[Cs+]
Mole Ratio in Different Solvents. A-Nitromethane,
B—Methanol, C-Acetone, D-Pyridine, E-Acetonitrile.

62

ratio according to the following method. Assuming that
only cation-ligand interactions are important, for the two
chemical environments of the metal ion, freely solvated
and complexed, in which the exchange rate is fast as com-
pared to the NMR time scale, a population average shift

is observed

6obs = XMGM + XMLGML (1)

where Gob is the observed chemical shift in ppm, XM

S

and XML are the respective chemical shifts for the two

sites. Equation 2 is easily derived from Equation 1 where

O

6 =fl

obs t (5M ‘ 6ML) + 6ML (2)

O
3

CM is the concentration of free metal ion and 0% is the

total concentration of metal ion. Assuming formation of

a 1:1 complex, we have the equilibrium

where L is the ligand. The concentration equilibrium

constant can be written as

63

where CML is the concentration of the complex, CM is the
concentration of free metal ion, and CL is the concentra-
tion of free ligand. By a simple algebraic substitution,

Equation 3 can be written as

 

t
CM ‘ CM
Kf g t t (u)

where CE is the total concentration of the ligand.

From Equations 2, 3, and u we can obtain Equation 5
which relates the observed chemical shift to the formation
constant, the total concentrations of the metal ion and the
ligand (0; and 0E respectively), the limiting chemical
shift of the complexes metal ion (GML), and to the chemical

shift of the free metal ion (6M).

2 2
_ t t 2 t 2 t 2 t t
Gob, - [(KfcM - Kch - 1),: (KfCL + KfCM - 2KfC CL
+ 2K 0t + 2K 0t + 1)1/2].(Efl:ifl£. + 6 ( )
r L r M t ) ML 5
2KfCM

In Equation 5, 0; and CE are known, 6M can be easily

determined from the measurement on solutions of free metal

ion. The equation then contains two unknowns Kr and 5ML'

t t
The procedure then is to input the Sobs’ CM’ CL’ and 5M

parameters and vary Kr and 6ML using a computer least

6“

squares program KINFIT (108). The procedure will continue
until the calculated chemical shifts correspond with the
experimental chemical shifts within the error limits. A
typical fit of the data for DB30C10-Cs+ in methanol at
30°C is shown in Figure 7.

The results for the DB3OC10°Cs+ complexes in various
solvents are shown in Table 9. It is seen that in acetone
solutions the stability of the complex is not affected by
a change in the concentration of salt or of the counter
ion. It is evident, therefore, that at low concentrations
of the cesium salts, which we used, the formation of the
complex is unaffected by ion pairing. It is reasonable to
assume that the same situation will exist in solvents with
higher donicities and/or higher dielectric constants such
as nitromethane, acetonitrile, and methanol. Comparison
of our values with those reported in the literature (and
obtained by different techniques) show a satisfactory
agreement (Table 9).

Since the cesium ion is rather weakly solvated be-
cause of low charge density of the cation, it is not
surprising that the stability of the complex is only

marginally dependent on the nature of the solvents.

65

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66

Table 9. Formation Constants and the Limiting Chemical
Shifts of DB3OClO-Cs+ Complexes in Various

 

 

 

Solvents.

Solvent log Kf 611m(ppm)
Nitromethane H.30i0.05 18.5810.0h
Acetonitrile 3.39:0.09 13.5630.95
Acetonitrile 3.50a ----
Acetone - 3.99:0.08b lu.69:o.01
Acetone u.ouio.05° 15.05i0.02
Acetone 3.96:0.07d 15.89ao.o7
Acetone 4.05:0.06e 15.28i0.05
Acetone 14.23f ----
Methanol “.18i0.07 l6.SOiO.ll
Methanol 14.23g ----
Pyridine “.4120.10 11.3510.02

 

 

aReference 35.
bo.oosl~_a_0ssCN.
°o.oosg Cs Picrate.
d0.005fl’CsBPhu.
e0.0025M|CsBPhu.
fReference 33.

gReference A.

67

3.3. COMPLEXATION 0F Na+, Cs+, and Tl+ IONS WITH DIBENZO-

2U-CROWN-8

3.3.1. 032u08 Complexes with Na+

The complexation between sodium ion and dibenzo-2u-
crown-8 was studied in nitromethane, acetonitrile, dimethyl—
formamide, dimethylsulfoxide, and pyridine. The measured
sodium-23 chemical shifts at various ligand to sodium
ion mole ratios and the line widths of the resonance lines
at the half height at 30°C are listed in Table 10. The
mole ratio plots are shown in Figure 8. It is seen im-
mediately that the solvent plays an important role in
the complexation reaction. In solvents of high donicity,
such as DMF and DMSO, the sodium-23 resonance is almost
independent of the ligand/sodium ion mole ratio. This
behavior indicates that in these solvents the immediate
environment of the sodium ion is not changed upon the
addition of the ligand indicating formation of a very weak

complex at the best.

On the other hand, in nitromethane, acetonitrile, and
pyridine solutions the sodium-23 resonances shift upfield
or downfield with some indication of a break at a mole
ratio of about 1 implying the formation of a 1:1 DB21408-Cs+
complex. The formation constants and the limiting chemical
shifts of the complex in different solvents are given in

Table 11. It is obvious that in solvents with low solvating

\

68

 

 

 

 

 

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m: om.: mm.a m: mm.m mm.H mm mm.~ :N.H
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m: m~.: :o.a mm Ha.m mm.o mm mm.m mm.o
o: m>.: No.0 mm wo.m mw.o mm mm.m mm.o
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mm on.: 00.0 Ha om.» oo.o NH no.=a oo.o
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69

 

 

 

 

 

 

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mos oe.e OH.H om am.o oa.a
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em mm.m mm.o Hm mm.o mm.o
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7O

 

 

 

 

 

 

 

 

 

 

14
12 -
10 -
PY
AN
8" ‘nur sot i;._ a_4n
5 NM
(ppm
6-
DMF
_.__
4..
2..
DMSO
__.___.__._'_______.__ ‘fit— 0“
°7
J l l I 11
0.0 0.5 1.0 1.5 2.0 2.5

[DB24C8]/[No+]

Figure 8. Sodium—23 Chemical Shifts y_s__. [DB24C8j/[Na+]
Mole Ratio in Different Solvents.

71

Table 11. Formation Constants and the Limiting Chemical
Shifts of DB2hCB-Na+ Complexes in Various

 

 

 

Solvents.

Solvent log Kf 611m(ppm)
Nitromethane 3.74:0.12 7.82:0.01
Acetonitrile 2.95:0.07 8.33:0.01
DMF m0 -_--
DMSO - m0 ----

Pyridine 2.89:0.10 9.20:0.09

 

 

72

abilities the complex is more stable. Once again pyridine
seems to be an exceptional case.

The limiting chemical shifts of the complexed sodium
ion in nitromethane, acetonitrile, and pyridine solutions
are close to each other indicating that in these solvents
the complexed sodium ion is no longer exposed to the sol-
vent molecules. This is only possible if the sodium ion
is located inside the cavity created by the twisting of
the ligand around the Cation. Another evidence which sup-
ports the formation of such a "wrap around" complex between
DB2HC8 and the sodium ion is the existence of the rela-
tively narrow signals for sodium-23 resonance in all solu-
tions. The width of the 23Na resonance lines in this case
are less than one half of those observed for Na+-DB30C10
system in the same solvents (Table 7). For nuclei such
as 23Na with appreciable quadrupole moment, we expect to
have a broader resonance line for a more unsymmetrical
environment around the nucleus. Thus existence of rela-
tively narrow 23Na resonance lines in this case indicates
a symmetrical environment around the sodium ion which can
only be obtained by the formation of a three-dimensional

complex.

3.3.2. DB2HC8 Complexes with Cs+

Cesium-133 chemical shifts were determined as a func-

tion of the ligand to the cesium ion mole ratio. The data

73

are given in Table 12 and the mole ratio plots are shown in
Figure 9. In all cases, the cesium-133 resonance has a
gradual shift, paramagnetic or diamagnetic depending on

the nature of the solvent, until a mole ratio of about 1

is reached and then begins to level off. This behavior
indicates formation of a 1:1 DB2u08-Cs+ complex in all
solvents.

The formation constants and the limiting chemical
shifts of the complex in various solvents are shown in
Table 13. From the results it is obvious that solvent
plays an important role in the complexation. In strong
solvating solvents such as DMF and DMSO the complex is
much weaker than in solvents of low and medium donor strength
such as nitromethane, acetonitrile and acetone. This behavior
shows the existence of a competition between the ligand and
the solvent molecules for the cesium ion. The only ex-
ception is the case of pyridine where despite the high
Gutmann donor number of the solvent the complex is un-
expectedly stable. This exception probably results from
the relatively weak interaction between the "soft base"
nitrogen atom of pyridine and a "hard acid" (cesium ion).

Three different cesium salts at different concentra-
tions were used to investigate the effects of the anion
and of the concentration on the complexation of cesium
ion with dibenzo-24-crown-8 in acetonitrile solution.

The results are also given in Table 13. As seen in all

7h

Table 12. Mole Ratio Study of Dibenzo-2u—Crown-8 Complexes
with 0.005flCs+ Ion in Various Solvents at 30°C.

 

 

 

 

 

 

 

 

 

 

NMa ANa ACa MeCHa
+ + + +
L/Cs 6(ppm) L/Cs 6(ppm) L/Cs 6(ppm) L/Cs 6(ppm)
0.00 57.57 0.00 -33.55 0.00 19.27 0.00 u6.15
0.21 53.31 0.25 -23.2u 0.25 21.19 0.20 nu.76
0.60 u5.62 0.62 -7.50 0.60 23.82 0.5a A1.50
0.70 uu.38 0.72 -3.70 0.77 25.15 0.78 u0.03
0.9a 91.37 0.8u -0.20 0.98 26.69 0.9a 39.02
1.05 no.59 0.98 5.85 1.12 27.70 1.02 38.63
1.10 u0.u2 1.09 8.32 1.25 28.16 1.18 37.87
1.18 39.86 1.13 8.55 1.81 28.u8 1.3u 37.39
l.u0 39.0u 1.38 12.12 2.19 28.79 1.60 37.09
1.87 38.56 1.78 13.68 2.66 29.02 1.76 36.92
2.39 38.u2 2.32 lu.1u 29.0u 2.07 36.79
2.88 38.u2 2.78 1M 22 2.78 36.70
DMFa DMSOa Pyb
+ + +
L/Cs 6(ppm) L/Cs 6(ppm) L/Cs 6(ppm)
0.00 0.26 0.00 -67.26 0.00 23.99
0.3a 3.12 0.u2 —63.27 0.38 22.82
0.60 5.23 0.71 -60.63 0.56 22.uu
0.80 6.68 0.86 -59.02 0.83 21.75
0.8a 6.86 1.00 -58.53 0.92 21.36
0.96 7.78 1.1u -57.15 1.05 21.19
1.12 8.78 1.20 -57.29 1.26 20.98
1.30 9.71 1.38 -55.90 1.52 20.88
1.10 10.11 1.85 -52.79 2.03 20.81
1.85 12.19 2.30 -u9.92 2.39 20.79
2.60 19.20 2.80 -u7.52 2.83 20.79
aCsSCN.
b

CSBPhu .

75

 

50b

 

 

 

 

 

 

I l I l J J

0.0 0.5 3.0 2.5 3.0

1.0 1.5
[DBZAC8]/[C 5*

+
Figure 9. Cesium—133 Chemical Shifts 1g. [DB214C8J/[Cs J
Mole Ratio in Different Solvents.

 

76

 

 

 

Table 13. Formation Constants and the Limiting Chemical
Shifts of DB214CBoCs+ Complexes in Various
Solvents.

Solvent Log Kf 611m(ppm)
Nitromethane M.lli0.08 38.2610.03
Acetonitrile 3.9uao.07a lu.86:0.02
Acetonitrile 3.89:0.07b lu.5010.07
Acetonitrile 3.90:0.011c lh.8210.06
Acetonitrile 3.97:0.011d 19.2310.06
Acetone 3.71:0.09 29.2“10.06
Methanol 3.65:0.05 36.h710.03
Methanol 3.78:0.08e ----
Dimethylformamide 2.10:0.0“ 2U.9330.96
Dimethylsulfoxide 1.61:0.0M -8.32i3.h6
Pyridine 9.00:0.03 20.7510.07

 

 

a0.00SM_CsSCN.
b0.0101310sSCN.
c0.005M CsBPhu.
d0.005110s1.

eReference A.

77

cases the limiting chemical shifts and the stability of
the complex are not affected either by a change in the salt
concentration or by a change in the anion. It is evident,
therefore, that at low concentrations used in this study,
the formation of the complex is unaffected by ion pairing.
We can assume that in solvents of higher donicities and/or
higher dielectric constants the same situation will exist.
The large divergence in DB2HCB'Cs+ complex limiting
chemical shifts (a range of about N6 ppm) probably indi-
cates that the cesium ion is not insulated from the solvent
by the ligand. The complex formation constant in methanol
solution is in satisfactory agreement with the previously

reported value (A).

3.3.3. DB24C8 Complexes with T1+

The thallium nucleus, 205Tl, has very favorable prop-
erties for NMR studies. It has a spin of I = 1/2 and its
chemical shift is very sensitive to small changes in the
chemical environment. The solvent dependence of the
thallium-205 chemical shift is over 2600 ppm (116,120)
in comparison to a chemical shift range of about 8 ppm for

7L1 (123,12u), 30 ppm for 23Na (123,121), and 130 ppm for
133Cs (125,126). The thallous ion is a useful probe
because its chemistry is very similar to that of the alkali
ions (127). In particular, the chemical properties and

ionic radii of Tl+ and K+ ions are very close (1.5“ X and

78

1.11 1, respectively) (128) so that thallium(I) ion can be
used as a NMR probe for potassium. Because of these
interesting properties, we were interested in studying

the complexation of thallous ion with dibenzo-2h-crown-8
in various nonaqueous solvents.

The thallium-205 chemical shift was measured as a
function of DB2HC8/Tl+ mole ratio in nitromethane, aceto-
nitrile, acetone, methanol, DMF, DMSO, and pyridine solu-
tions at 30°C. The thallous ion concentration was main-
tained constant at 0.005 M_in all cases. The mole ratio
data are given in Table 1A and the variations of the 20511
chemical shift as a function of the ligand/Tl+ mole ratio
are shown in Figure 10. In poor solvating solvents such
as nitromethane, acetonitrile, and acetone, the addition of
the ligand to the thallous ion solution causes a quite
linear change in the 20511 chemical shift (upfield or
downfield) until a mole ratio of l is reached. After the
mole ratio of 1, further addition of the ligand does not
affect the thallium-205 resonance. This behavior indicates
the formation of a very stable l:l complex between the
thallium(I) ion and the ligand in the above solvents. 'On
the other hand, in methanol solution the mole ratio plot
shows some curvature around the mole ratio of l and a
limiting value is obtained at mole ratios greater than
2.5, indicating formation of a weaker 1:1 complex than

that in the previous solvents. In solvents of high

79

Table 11. MOle Ratio Study of Dibenzo-21-Crown-8 Complexes
with O-OO5E.T10101 in Various Solvents at 30°C.

 

 

 

 

 

 

 

 

NM AN Ac MeOH
L/Tl+ 6 + 6 + 6 +
(ppm) L/Tl (ppm) L/Tl (ppm) L/Tl (ppm)
0.00 366.6 0.00 209.8 0 00 206.1 0.00 -lo7.5'
0.73 281.9 0.31 222.1 0 36 226.1 0.16 -7.7
0.97 250.1 0.70 237.8 0 72 212.8 0.75 60.7
1.11 216.9 0.85 213.1 0 82 219.9 0.86 78.6
1.22 216.9 0.96 215.6 0 97 253.1 1.06 118.2
1.65 216.9 1.09 219.1 1 10 257.0 1.15 131.8
1.88 216.7 1.21 219.8 1.31 258.3 1.28 113.1
2.15 216.9 1.52 219.8 1.56 258.9 1.59 166.3
2.68 216.9 1.88 219.9 1.87 259.7 1.79 176.7
2.19 219.9 2.23 259.6 2.06 187.6
2.72 250.0 2.80 260.1 2.70 192.3
DMF DMSO Py
+ + +
L/Tl 6(ppm) L/Tl 5(ppm) L/Tl 6(ppm)
0.00 -l90.2 0.00 -319.3 0.00 651.1
0.35 -l78.3 0.15 -320.7 0.38 603.1
0.78 -163.5 0.80 -321.7 0.67 569.1
0.91 -156.3 0.87 -322.0 0.87 511.9
1.12 -151.5 0.98 -322.3 0.99 529.3
1.22 -118.0 1.10 -322.7 1.08 516.7
1.16 -139.3 1.29 -323.1 1.22 198.8
1.80 -126.2 1.61 -321.2 1.59 161.1
2.08 -115.2 1.83 -321.8 1.81 111.5
2.19 -113.6 2.37 -325.7 2.35 398.5
2.77 -102.1 2.81 -326.6 2.80 360.2

 

 

80

 

600

400

 

 

 

 

 

 

 

 

NM
200 MeOH
0.—
DMF
-20 l I I I I
((1.0 0.5 1.0 1.5 2.0 is 3.0
[DB24C81/[TF]

Figure 10. Thallium-205 Chemical Shifts XE: [DB21081/[T1+]
Mole Ratio in Different Solvents.

81

donicities, such as DMF, DMSO, and pyridine, upon addition
of the ligand there is a gradual shift of the 205T1 reson-
ance which does not reach a limiting value even at a mole
ratio of 3 (Figure 10), indicating formation of a weak
complex in these solvents.

The formation constants and the limiting chemical shifts
of the l:l complex of the ligand with the thallous ion in
various solvents are given in Table 15. It is immediately
obvious that the stability of the complex is very much
dependent on the nature of the solvent. The stability of
the complex increases with decreasing Gutmann donor number
of the solvents. It is interesting to note that in con-
trast to the sodium and cesium ion complexes with DB21C8,
in pyridine solution the DB21C80T1+ complex is expectedly
weak. The thallous ion as a "soft acid" can strongly
interact with the nitrogen atom of pyridine which is a
"soft base" so that the resulting DB2HC8oTl+ in this
solvent is weak. As shown in Table 15, the formation
constants of the complex in acetonitrile and methanol
solutions are in a good agreement with the values reported
by Hofmanova gt él' (35) who used polarographic technique
for the measurements.

The limiting chemical shifts of the complexed thallous
ion in solvents of low and medium donicities such as
nitromethane, acetonitrile, acetone, and methanol are

close to each other (Figure 10). This behavior indicates

82

Table 15. Formation Constants and the Limiting Chemical
Shifts of DB21C8-Tl+ Complexes in Various

 

 

 

Solvents.

Solvent Log Kf 611m(ppm)
Nitromethane >5 215.00
Acetonitrile 1.81:0.05 250.110.l
Acetonitrile 1.80a ----
Acetone - 1.15:0.05 260.510.l
Methanol 3.19:0.07 215.2:3.8
Methanol 3.10a ----
Dimethylformamide 1.1610.2l 336 :91

Dimethylsulfoxide <l.0 ——~—
Pyridine 1.61 0.01 —150 116

 

 

aReference 35.

83

that the thallous ion is insulated from the solvent by the
ligand so that there is only a weak interaction between

the solvent molecules and the cation. In solvents of
strong solvating ability such as DMF, DMSO, and pyridine,
however, the limiting chemical shifts are solvent dependent
indicating a strong interaction between the solvent mole-

cules and the cation.

3.3.1. Conclusions

The formation constants of the 1:1 complexes of Na+,
Cs+, and T1+ ions with dibenzo-21—crown—8 in various
solvents are compared in Table 16. From the results it
is immediately obvious that solvent plays an important
role in the complexation reactions. With the exception of
the sodium and cesium ion complexes in pyridine solution,
in all cases the stabilities of the complexes are de-
creased with increasing solvating abilities of the sol-
vents. In poor solvating solvents such as nitromethane,
acetonitrile, and acetone the stabilities of the Na+,

Cs+, and Tl+ complexes decreases in the order DB2108-Tl+
> 082108-Cs+ > DB2108-Na+.

For large crown ethers such as dibenzo-30-crown-10
and dibenzo—21-crown-8 which are capable of formation of
three-dimensional "wrap around" complexes with cations (33),
we expect the size of the cation to play an important role

in the complexation reaction. The complexation of a cation

81

Table 16. Formation Constants of 1:1 Complexes of Na+,
03+, and Tl+ Ions with Dibenzo-21-Crown-8 in

Various Solvents.

 

 

 

 

Log Kf

Solvent Cs+ Tl+ Na+
Nitromethane 1.11:0.08 >5.0 3.7110.12
Acetonitrile 3.91:0.07 U.8110.05 2.95:0.07
Acetone 3.71:0.09 U.lSi0.0S ----
Methanol 3.60:0.05 3.1910.05 ----
Dimethylformamide 2.10:0.01 1.1610.2l 10.0
Dimethylsulfoxide 1.61:0.01 <l.0 ~0.0
Pyridine 1.00:0.03 1.61:0.01 2.89i0.10

 

 

85

and the ligand with appropriate relative sizes results in
the formation of a very strong complex, for example
complexation of potassium ion by dibenzo-30-crown-10
(26,33). In the case of the complexation of the ligand
with larger cations, the three-dimensional structure cannot
be complete and, therefore, a weaker complex will be formed.
0n the other hand, in the case of the complexation of the
ligand with smaller cations, the ligand can still twist
around the cation to form the three—dimensional structure,
but in this case the oxygen atoms of the ligand will have
repulsive forces on each other which, consequently, result
in the formation of a weak complex between the cation and
theligand.

The formation constants and the limiting chemical shifts
obtained in this study support the above conclusions for
the complexation of Na+, Cs+, and Tl+ ions by dibenzo—21—
crown-8. Cesium ion is too large to form a complete
"wrap around" complex with the ligand. Thus the cation
remains exposed to the solvent molecules which causes
the existence of quite scattered limiting chemical shifts
of the complex in different solvents (Figure 9). In the
case of the sodium ion, despite the smaller size of the
cation, the "wrap around" structure can be formed, but
because of the repulsive forces of the oxygen atoms of the
ring on each other the resulting complex is weak. In

this case, the cation is insulated from the solvent by

86

the ligand and, therefore, the complex limiting chemical
shifts in different solvents are close (Figure 8). Thallous
ion seems to have the most suitable size for the cavity
since the DB2108-Tl+ complex is the strongest among the

series and its limiting chemical shifts are very close.

3.1. COMPLEXATION OF Na+, Cs+, and 11+ IONS WITH DIBENZO-
2l-CROWN-7

3.1.1. DB2lC7 Complexes with Na+

The complexation of sodium ion with dibenzo-2l-crown-7
was investigated in nitromethane, acetonitrile, acetone,
dimethylformamide, dimethylsulfoxide, and pyridine solu—
tions by sodium-23 NMR. The concentration of sodium tetra-
phenylborate was maintained at 0.025 M while the concentra-
tion of the ligand was varied from zero to about 0.075 M.
The sodium-23 chemical shifts were determined as a function
of DB2lC7/Na+ mole ratios. The data are given in Table 17
and the mole ratio plots are shown in Figure 11. In DMF
and DMSO solutions the 23Na chemical shift is very little
affected by the addition of the ligand, although the
resonance line shows some broadening (Table 17). The
results seem to indicate the formation of very weak com-
plex at the best in these solvents. This is not surpris-

ing since DMF and DMSO are solvents with strong solvating

87

 

 

 

 

 

mm om.m mo.m mm om.0H mo.m omH o:.m mm.m
on mm.m mm.m mm NH.OH mm.m mma m:.m om.m
so mm.m sm.H om mo.oa Hm.H smH ma.m ms.fi
mm om.m mm.H as em.m Hm.H Hma mm.m mm.H
mm me.m HH.H m: mm.m MH.H mma ms.m NH.H
mm =m.m mm.o ms so.m mm.o smfi om.m mo.H
mm me.m em.o m: me.m mm.o Hmfl mm.oa om.o
we ma.m 05.0 mm Hm.m ms.o NHH ma.oa =s.o
Hm m>.m m=.o em :m.m m=.o we mm.HH H=.o
ow om.m oo.o ma sa.s oo.o ma mm.mH oo.o
Anmvmxaoo Agoove +mzxd Anmvmxaoo Asodve +mzxd Anmvmxaoo “shove +m2\q
ca za :2
.000m um mpcm>aom mSOfipm>
ca sammmz_ammo.o spas moxoaosoo euszoaonamnonsooao co sosom oaomm ofloz .sH oHome

88

 

 

 

 

 

 

mmm 00.0 00.m
00m 00.0 0m.m 00H 0H.0 00.0 00 00.0 mH.m
mam mm.0 m0.H 03H 0H.0 00.H mm 00.0 00.H
000 00.5 mm.H 00H 01.0 mm.H 00 H0.: 0m.H
0:0 0:.» 0H.H NHH 0H.0 mm.fi as m0.0 0m.H
000 00.0 00.0 00 0H.0 ~0.H 00 00.0 m0.H
0H0 as.m 00.0 mm mH.0 00.0 H0 00.: 00.0
00H 00.: 00.0 mm 00.0 00.0 mm 00.0 H0.0
maa m=.m mm.0 m0 m0.0 0m.0 0m 00.: 00.0
0H 0H.Hu 00.0 m: H0.0- 00.0 00 00.: 00.0
Asavm\flea Aecaos +12\a AN00N\H90 Aeaavc +12\a A600~\H>< Aecave +QZ\4

00 0020 0:0

.0cssfieoo0 .NH magma

89

 

 

 

 

 

 

 

 

 

6-
o
DMF
4..
2..
0' DMSO
-2 1 1 J, I I I
00 0-5 1.0 1.5 2.0 2.5 3-0

[0821C 7]/[No+]

Figure 11. Sodium-23 Chemical Shifts E- [DB2lC7]/[Na+]
Mole Ratio in Different Solvents.

9O

abilities which can strongly compete with the ligand for
the cation.

0n the other hand, a considerable change in 23Na
chemical shifts was observed in nitromethane, acetonitrile,
acetone, and pyridine solutions upon the addition of the
ligand indicating a change in the chemical environment of
the nucleus. The diamagnetic shift of the sodium-23
resonance in acetonitrile, acetone, and pyridine solution
(paramagnetic in nitromethane) begins to level off at a
mole ratio of about 1, indicating formation of a 1:1 complex
between the ligand and the sodium ion. In all cases, the
line width of the 23Na resonance at the half height
(Avl/Z) increases almost linearly with the ligand/Na+
mole ratio until a mole ratio of about 1 is reached, and
then it remains almost constant upon further addition of
the ligand (Table 17). This is not surprising because for
a nucleus such as sodium with an appreciable quadrupole
moment we expect the line width of the resonance signal to
be sensitive to the electrical field gradient around the
nucleus. The solvated sodium ion has a symmetrical environ-
ment and, therefore, the line width of the resonance signal
is small (about 10 to 20 Hz depending on the solvent used).
Complexation of the sodium ion with the ligand creates an
unsymmetrical environment around the cation which results
in the broadening of the resonance signal. As the DB2IC7/Na+

mole ratio increases, the amount of the complexed sodium

91

23Na resonance signal

ion increases and, therefore, the
becomes broader. After a mole ratio of l is reached,
essentially all of the cation is complexed and further
addition of ligand has no effect on the sodium-23 resonance.
This behavior, once more, confirms the formation of a stable
1:1 complex between DB2lC7 and sodium ion.

The formation constant and the limiting chemical
shifts of DB21C7-Na+ complex in various solvents at 30°C
are shown in Table l8.’ With the exception of pyridine
there is an inverse relationship between the donicity of
the solvent and the stability of the complex. It has been
pointed out previously, however, that being a soft nitrogen
donor, pyridine does not strongly solvate a hard ion such
as sodium ion (80) and, therefore, a stable complex can
be formed between the ligand and the sodium ion in this
solvent. The limiting chemical shifts of the complexed
sodium ion in nitromethane, acetonitrile, acetone, and
pyridine solutions are very close, indicating that the
sodium ion is enclosed inside the ligand's cavity so that

the solvent molecules can barely interact with the cation.

3.1.2. DB21C7 Complexes with Cs+

Cesium-133 NMR was used to study the complexation of

the cesium ion with dibenzo-2l-crown-7 in various non-

aqueous solvents. The 133Cs chemical shift was determined

:15 a function of the ligand to cesium ion mole ratio.

92

Table 18. Formation Constants and the Limiting Chemical
Shifts of DB2lC7-Na+ Complexes in Various

 

 

 

Solvents
Solvent Log Kf 611m(ppm)
Nitromethane 3.11:0.05 9.35:0.01
Acetonitrile 2.7810.08 10.2910.02
Acetone 2.2810.08 9.88:0.06
Dimethylformamide- ~0.0 ----
Dimethylsulfoxide ~0.0 ----

Pyridine 2.56:0.05 9.1710.06

 

 

93

The results are shown in Table 19 and in Figure 12. The
cesium ion concentration was maintained at 0.005 M, Upon
addition of the ligand, the line width observed shows a
slight increase from about 1 Hz to 11 Hz. The complexation
was studied up to a mole ratio of about 3, where a limiting
value for the chemical shift of the complexed cesium ion

is reached in nitromethane, acetonitrile, acetone, methanol,
and pyridine solutions. The variations of the cesium-133
chemical shift as a function of DB2lC7/Cs+ mole ratio in
the above mentioned solvents show a single inflection point
at a mole ratio of about 1, indicating formation of a 1:1
complex between the ligand and the cesium ion.

No evidence for formation of a 2:1 (ligand to metal
ion) complex was observed in any of the solvents used.
While the existence of a 2:1 complex of the ligand with
cesium ion in the solid crystalline state has been demon-
strated (23), it does not follow necessarily that it also
exists in solution. Frensdorff (1), however, has reported
formation of a very weak 2:1 complex between DB2lC7 and
cesium ion in methanol solution.

The formation constant and the limiting chemical shifts
obtained from the computer analysis of the mole ratio
data in various solvents are shown in Table 20. The large
difference in the limiting chemical shifts of the complex
in different solvents is a good indication that the metal

ion remains exposed to the solvent. The cesium ion is

91

Table 19. Mole Ratio Study of Dibenzo-21-Crown-7 Complexes
with 0.005! CsSCN in Various Solvents at 30°C.

 

 

 

 

 

 

 

 

NM AN AC MeOH
+ + + +
L/Cs 6(ppm) L/Cs 6(ppm) L/Cs 6(ppm) L/Cs 6(ppm)
0.00 57.89 0.00 -3l.93 0.00 19.18 0.00 16.97
0.26 50.13 0.35 -21.79 0.30 16.16 0.30 39.53
0.50 13.37 0.61 -l9.06 0.56 13.82 0.51 32.86
0.77 35.22 0.82 -15.19 0.79 11.65 0.68 29.92
0.91 30.75 0.92 —l3.17 0.89 10.72 0.90 21.61
1.11 27.37 1.01 -ll.76 1.03 9.67 1.03 22.10
1.28 25.11 1.22 -9.U7 1.12 8.9“ 1.1“ 21.55
1.18 21.86 1.33 -9.36 1.20 8.78 1.28 20.62
1.80 21.21 1.50 -9.05 1.13 8.12 1.16 19.90
2.25 21.15 1.83 -8.7N 1.90 8.2“ 1.85 19.66
2.67 21.07 2.52 -8.52 2.37 88.20 2.12 19.53
3.21 -8.10 3.16 8.10 3.01 19.31
DMF DMSO Pya
+ + +
L/CS 6(ppm) L/CS 6(ppm) L/CS 6(ppm)
0.00 2.51 0.00 -66.59 0.00 21.53
0.37 2.20 0.30 ~63.18 0.17 19.80
0.69 1.88 0.72 -58.19 0.50 11.12
0.86 1.61 0.93 -55.89 0.77 3.91
0.99 1.50 1.06 -51.27 0.91 -0.67
1.19 1.91 1.20 -53.25 0.99 -2.30
1.31 1.36 1.90 -50.92 1.1“ -5.U7
1.95 1.27 1.82 -17.36 1.29 -5.63
2.35 1.19 2.37 -13.72 1.51 -5.78
2.9“ 1.11 3.06 -39.73 1.88 —5.82
2.37 -5.91
2.97 -6.05

 

 

60

95

 

 

 

 

 

 

 

 

 

 

 

NM
MeOHI
f 4
. a
—.> AIPY
AN’“
J J L
0.0 0.5 1.0 1.5 2:6 285 380
[DB21C7]/[Cs"]
Figure 12. Cesium-133 Chemical Shifts vs. [DB2107]/[Cs+]

Mole Ratio in Different SolVEhts.

96

Table 20. Formation Constants and the Limiting Chemical
Shifts of DB2107°Cs+ Complexes in Various

 

 

 

Solvents.

Solvent Log Kf 611m(ppm)
Nitromethane 1.11i0.07 23.78i0.05
Acetonitrile 3.95:0.01 -8.1720.02
Acetone 3.93:0.06 7.98:0.02
Methanol 3.96:0.06 19.07:0.01
Methanol 1.20a ----
Dimethylformamide 2.81:0.10 0.92:0.01
Dimethylsulfoxide 1.72:0.01 -0.96:3.18
Pyridine 1.27:0.07 -6.2110.03

 

 

aReference 1.

97

known as a large cation with a low charge density which
cannot be strongly solvated by the solvent molecules. Thus
it is not surprising if the stability of the complex in
nitromethane, acetonitrile, acetone, methanol, and pyridine
is only marginally dependent on the nature of the solvent.
In DMF and DMSO, however, the DB2lC-Cs+ complex is much
weaker than that in the above solvents. Dimethylformamide
and dimethylsulfoxide are solvents of high donicities, and
their interactions even with a large cation such as cesium
ion is still strong enough to compete with the ligand for
this cation and to prevent formation of a strong complex.
The value obtained for the stability constant of the
DB2lc7-Cs+ complex in methanol solution at 30°C is in a
satisfactory agreement with the literature reported value

in the same solution at 25°C (1) (Table 20).

3.1.3. DB2107 Complexes with T1+

In order to determine the stoichiometry and the
stability of the thallium(I) complex with dibenzo-2l-
crown-7 in various solvents, the thallium-205 NMR chemical
shift was measured as a function of DB21C7/Tl+ mole ratio.
The concentration of thallium(I) perchlorate was maintained
at 0.005 M, while the concentration of the ligand was
varied from zero to about 0.015 M. The data are given in
Table 21 and the mole ratio plots are shown in Figure 13.

In nitromethane, acetonitrile, and acetone solutions, the

98

Table 21. Mole Ratio Study of Dibenzo-21-Crown—7 Complexes
with 0.005M’T1C10u in Various Solvents at 30°C.

 

 

 

 

 

 

 

 

NM AN Ac

+ + +
L/Tl 6(ppm) L/Tl 6(ppm) L/Tl 6(ppm)
0.00 319.6 0.00 213.9 0.00 150.3
0.71 319.2 0.37 231.1 0.12 202.8
0.86 315.1 0.85 260.5 0.51 216.8
1.01 310.1 1.05 272.2 0.77 216.5
1.11 309.0 1.15 271.1 0.91 262.1
1.28 308.8 lL32 271.3 1.06 271.6
1.13 308.5 1.50 271.6 1.21 275.0
1.88 308.1 1.98 271.7 1.51 277.3
2.37 308.3 2.55 271.6 1.63 277.9
3.11 308.1 3.21 271.5 2.03 278.1

2.57 278.2
3.09 278.3

MeOH DMF DMSO

+ + +
L/Tl 6(ppm) L/Tl 6(ppm) L/Tl 6(ppm)
0.00 -112,o 0,00 -l18.5 0.00 -212.9
0.28 -39.u 0.10 -9u.3 0.27 -218.3
0.52 32.0 0.70 -55.1 0.51 -223.1
0.76 103.7 0.86 -36.0 0.82 —230.0
0.87 130.0 1.03 -15.2 0.92 -231.9
0.96 151.1 1.13 -6.7 1.01 -233.1
1.10 182.9 1.33 15.3 1.16 -236.2
1.23 188.0 1.85 11.1 1.38 -210.3
1.18 200.2 2.12 75.5 1.60 -216.0
1.75 201.1 3.11 101.1 1.80 -218.5
2.20 207.5 2.17 -261.0
2.90 209.1 3.09 -273.3

 

 

99

 

 

 

 

 

 

 

 

 

 

N1
3mJ- .Ac
1 - 7:.-
aMeOH
zoo *
DMF
mo -
8
(PM)!
0 .-
-]00 .—
-2oo
DMSO
1 l L l 1 L
00 05 10 LS 10 25 10
[Dismal/[n+1

Figure 13. Thallium—205 Chemical Shifts gs. [DBZlC7]/[T1+]
Mole Ratio in Different Solvents.

100

thallous ion resonance shifts either upfield or downfield
as the concentration of the ligand increased. Once again
mole ratio plots show a break at a mole ratio of 1, indi-
cating formation of a 1:1 complex. The break is more pro-
nounced in the nitromethane and acetonitrile curves compare
to that in acetone and methanol plots which is indicative
of formation of a stronger complex in the former solvents.
No other break is found, indicating that if any 2:1 (ligand
to metal ion) complex formation occurs it is negligible.

In solvents of high donicities such as DMF and DMSO
only a gradual shift of thallium-205 resonance is observed
upon addition of the ligand which does not reach a limit-
ing value even at a mole ratio of about 3. This behavior
indicates formation of a weak complex between the ligand and
the thallous ion in these solvents. The formation constant
and the limiting chemical shift of the 032107-11+ complex
in various solvents are given in Table 22. It is obvious
that the nature of the solvent has a large effect on the
stability of the complex. There is an inverse relationship
between the Gutmann donor number of the solvent and the
complex formation constant. Except for the acetonitrile
and acetone cases the complex limiting chemical shifts
are different in different solvents, indicating that the
complexed thallous ion remains exposed to the solvent.

For stable complexes with Kf > 105, the 205T1 Chemical

shift-mole ratio plots consist of two straight lines

101

Table 22. Formation Constants and the Limiting Chemical
Shifts of DB2lC7-T1+ Complexes in Various

 

 

 

Solvents.

Solvent Log Kf 611m(ppm)
Nitromethane >5 308.1
Acetonitrile >5 271.5
Acetone 1.71:0.08 278.5i0.1
Methanol - 3.97:0.03 212.710.3
Dimethylformamide 2.18:0.02 233.8:8.1

Dimethylsulfoxide 0.63:0.15 ----

 

 

102

intersecting at 1:1 mole ratio. These kinds of plots cannot
be analyzed by our technique, and in such cases we can only
conclude that log Kf > 5. This behavior is observed for

the DB2107-Tl+ complex in nitromethane and acetonitrile
solutions which is not surprising since both nitromethane
and acetonitrile are poor solvating solvents with respective

Gutmann donor number of 2.7 and 11.1.

3.1.1. Conclusions

The formation constants of the 1:1 complexes of Na+,
03+, and Tl+ ions in various solvents are compared in
Table 23. In solvents of low and medium solvating abilities
such as nitromethane, acetonitrile, acetone and methanol,
the stabilities of DB21C7 complexes with sodium, cesium,
and thallium(I) ions decrease in the order DB21C7-T1+ >
DB21C7-05+ > DB21C7-Na+. The ionic radii of the above
cations vary in the order Na+ < T1+ < Cs+. Cesium ion
has a diameter of 3.68 1 (128) which is Just the right
size to fit conveniently inside the cavity of dibenzo-2l-
crown-7 with the size of 3.1-1.3 K (1), while sodium ion
(diameter 2.21 3) is too small for the ligand's "hole".
Thus the DB2lC7-Cs+ complex is more stable than the
DB21C7-Na+ complex. In the case of DB2lC7-Cs+ complex the
crown ether only occupies the equatorial coordination sites
of the cation. Thus the complexed cesium ion remains

exposed to the solvent molecules from the axial positions

103

 

 

 

 

Table 23. Formation Constants of 1:1 Complexes of Na+,
03+, and T1+ Ions with 032107 in Various
Solvents.
Log Kf
Solvent Na+ Cs+ 11*
Nitromethane 3.1110.05 1.11:0.07 >5
Acetonitrile 2.78:0.08 3.95:0.01 >5
Acetone 2.28:0.08 3.93:0.06 1.71:0.08
Methanol ---- 3.96:0.06 3.97:0.03
Dimethylformamide ~0.0 2.81iO.10 2.18:0.02
Dimethylsulfoxide ~0.0 1.72:0.01 0.63:0.15
Pyridine 2.5610.05 1.27:0.07 ---—

 

 

101

and, therefore, the complex limiting chemical shift is
strongly solvent dependent (Figure 12).

In the case of DBZlC7-Na+ complex, however, because of
the small size of the cation, the ligand is able to twist
around the sodium ion to form a three—dimensional array
and thus insulate it from the solvent molecules. Conse-
quently, the limiting chemical shift of the complexed
cation is essentially solvent independent (Table 18).

In the case of the thallium(I) complex, despite the more
inconvenient relative sizes of the cation and the ligand
than the cesium ion case, the DBZlC7-Tl+ complex is more
stable than the DB210-Cs+. This is not surprising since
the thallous ion is known to bond to the oxygen atoms of
the macrocyclic ligands by an ion-dipole interaction with
a covalent contribution (129) which results in the forma-
tion of a very strong complex between the cation and the
ligand.

In solvents of strong solvating abilities such as DMF
and DMSO no evidence for the formation of a complex between
the ligand and the sodium ion was observed. The cesium
and thallous ions, however, form weak complexes with the
ligand in these solvents, but in this case the DB21C7-CS+
complex is more stable than the DB21C7-Tl+ complex. The
smaller stability constant of the thallous ion complex
inay be related to the preference of thallous ion as a

soft acid for the nitrogen and sulfur atoms as soft bases

105

in the DMF and DMSO structures respectively. Pyridine is
an exceptional case where despite the high donicity (Gut-
mann donor number of 33.1) of the solvent, the resulting
complexes of the ligand with sodium and cesium ions are
unexpectedly stable. The reason for this behavior has

been discussed previously.

3.5. DISCUSSION

The stability constants of Na+, 08+, and Tl+ ion com-
plexes with dibenzo-2l-crown-7, dibenzo-21-crown-8, and
dibenzo-30—crown-10 in various solvents at 30°C are shown
in Table 21. It is immediately obvious that in all cases
the nature of the solvent has a great effect on the
stabilities of the complexes. The magnitude of the solvent
effect on the complex formation constants decreases in the

4.

order of Tl+ > Na > Cs+. Cesium ion is a large cation with

a low charge density which cannot strongly interact with
the solvent and, therefore, the stabilities of the cesium
ion complexes are only marginally dependent on the nature
of the solvent. In the case of the sodium ion complexes
the solvent effect is more pronounced, which can be related
to the smaller Size and consequently to the higher charge
density of the sodium ion than that of the cesium ion.
Since the thallium(I) ion is bonded to the solvent

Inolecules by an ion-dipole interaction with a covalent

106

 

 

 

 

 

 

 

 

:o.ou:m.a IIII oa.ouaa.: mo.onoo.= No.0wpm.: oa.onmm.m mo.o«mm.m maficfipmm
oo.Hv mfi.ommm.o IIII :o.o«am.a zo.ommw.a 0.08 0.08 omza
Hm.o«mH.H mo.o«ma.m IIII :o.owoa.m oa.ow=m.m 0.08 0.08 min
mo.owma.m mo.ou~m.m wo.owma.z mo.owom.m mo.owmm.m IIII IIII Hocmnpmz
mo.onma.: mo.oaan.= no.0Hmm.m mo.ona>.m mo.oumm.m IIII mo.osmm.m ocopoo«
mo.onam.: mA mo.owmm.m no.0«zm.m no.owmm.m uo.o«mm.m mo.o«mu.m mafippacoumo<
omA mA mo.ouom.n wo.owaa.= No.ouza.: ma.ou:w.m mo.ouza.m ocmnpothpdz

mozmmo scamma oaoommo mozmmn noammo mozmma woamma pco>aom

+HB +mo m2
ox moo
.ooom pm mucm>aom msofihm> ca oHoome cam .mozmma

.woamma spa: mCOH +HB new .+mo .+mz mo moonQEoo Hua no mpcmumcoo coapmapom .zm magma

107

contribution, it is not surprising that the stabilities
of the complexes depend strongly on the nature of the sol-
vent.

With the exception of the sodium and cesium ion com-
plexes in pyridine solution, there is an inverse relation-
ship between the donicity of the solvent and the stability
of the complex. In the case of pyridine, however, we have
a solvent with the highest donor strength (142;, Gutmann
donor number of 33.1) and yet the sodium and cesium ion
complexes are strong in this medium. It has been pointed
out previously that pyridine as a "soft" nitrogen donor
does not solvate strongly "hard" cations such as an alkali
ion (80,130).

As can be seen, sodium ion forms a more stable complex
with DBZ1C8 than with DB21C7. While the cavity sizes of
both ligands are too large for the small sodium ion, the
increased number of the oxygen atoms, as binding sites,
in DBZ1C8 seems to play an important role in the complexa-
tion. As discussed in previous sections (3.3.1 and 3.1.1)
the limiting chemical shifts of the complexed cations are
relatively solvent independent in both cases indicating
that the sodium ion is insulated from the solvent by the
jligands. According to this picture it is evident that
iDBZ1C8 with more binding sites can form a more stable
Icomplex with the sodium ion than DB21C7. In the case of

the thallous ion complexes the opposite behavior is

108

seen: 032107-11+ is more stable than DB21C8°T1+. In this
case the relative sizes of the ligand to the cation seem
to be the key factor in the complexation reactions. The
thallous ion with the size of 3.08 K has a more convenient
fit inside the cavity of 002107 (with the size of 3.1—1.3 1)
than with DB21C8 which has a cavity size greater than 1 X.
With the exception of the DB3OClO-Cs+ complex in aceto-
nitrile solution the stabilities of the cesium ion complexes

with the ligands decreases in the order DB3OC10-CS+ >

+ in all solvents used. Dibenzo-30-

002107-Cs+ > DB21C8-Cs
crown-10 is a large molecule with enough oxygen atoms in
the ring which can form a stable "wrap around" complex with
cesium ion (Sec. 3.2.2). The Sizes of the dibenzo-21-
crown-7 cavity and the cesium ion are very close so that
the cation can be held by the ligand to form a stable
complex, which is expectedly weaker than DB3OC10-Cs+
complex because of its two-dimensional structure. Dibenzo-
21—crown-8 has neither a long enough chain to form a three-
dimensional complex with the cesium ion nor a convenient
cavity size to hold the cation as tight as dibenzo-21-

crown-7 can. Thus the resulting complex is the weakest

in the series.

CHAPTER 1

CESIUM-133 NMR STUDY OF THE THERMODYNAMICS OF THE
COMPLEXATION 0F DIBENZO-30-CROWN-10, DIBENZO-21-CROWN-8,
AND DIBENZO-21-CROWN-7 WITH CESIUM ION
IN NONAQUEOUS SOLVENTS

109

1.1. INTRODUCTION

In order to have a deeper understanding of the thermo-
dynamics of the complexation reactions, it is useful to
consider separately the enthalpic and the entropic contribu-
tions to the reaction. The change in free energy of complexa-
tion, AG°, can be divided into two parts: the change in
enthalpy, AH°, and the change in entropy, A80, of the

reaction:
AG° = AH° - TAS° (l)

The complex can be stabilized by meeting one of the follow-
ing requirements: (a) AH° < 0 and dominant, TAS° < 0
(enthalpy stabilized); (b) AH° > 0, TAS° > 0 and dominant
(entropy stabilized); and (0) either AH° < 0 and dominant,
TAS° > 0 or AH° < 0, TAS° > 0 and dominant (both enthalpy
and entropy stabilized).

Presently available thermodynamic data for the com-
plexation of alkali metal ions with uncharged ligands such
as crowns and cryptands are not detailed enough to give a
clear picture of the thermodynamic behavior of the com-
plexes. The results of the thermodynamic studies, reported
so far, Show that most of the alkali ion complexes with
crowns and cryptands (18) are enthalpy stabilized but

entropy destabilized. In particular, the reported

110

111

thermodynamic data for the complexation of the cations
with large crown ethers (112;: 2l-crown-7 and larger) are
quite sparse. To the best of our knowledge, there are
Just four papers available (96,33,31,l3l), reporting the
thermodynamic parameters fortflmecomplexation of large
crowns.

In previous work, Lehn and coworkers (132) and Cahen
_§|al. (110) both noted that the observed chemical shift
of the complexed metal ion varied with temperature. Cesium-
133 and lithium-7 NMR were used by E. Mei (112) and by A.
Hourdakis (133) to determine the thermodynamic parameters
for the complexation of the cesium and lithium ions with

some crowns and cryptands in nonaqueous solutions.

The results of the study of the stoichiometries and

+, Cs+, and T1+ ion complexes with

stabilities of Na+, K
dibenzo-30-crown-10, dibenzo-21-crown-8, and dibenzo-2l-
crown-7 in various solvents were discussed in chapter 3.
In this chapter we will attempt to evaluate the thermo-

dynamic parameters of the complexation reaction between

cesium ion and the above mentioned crown ethers.

1.2. DB3OC10 COMPLEXES WITH Cs+

The method for determining of the thermodynamic values
‘was based on the temperature dependence of the complex

‘formation constant. The complex formation constant is

 

112

related to the relevant thermodynamic parameters by the

following relationships:

 

AGO = AHO - TASO (l)
o o
in Kf = - %gr-+ A: (3)

Thus a plot of in Kf vs. l/T gives a straight line with a

AH° AS° o
slopt of - R and an intercept of —R—, provided that AH

 

is independent of temperature over the temperature range
considered.

The variation of the cesium-133 chemical shift was
measured as a function of DB3OClO/Cs+ mole ratio in nitro-
methane, acetonitrile, acetone, methanol, and pyridine
at different temperatures. In all cases studied, only
one population average signal was observed indicating that
the exchange of the metal ion between the two sites (112;:
free ion in the bulk solution and the complexed ion) is
faster than the NMR time scale. The data are given in
Table 25 and the mole ratio plots at different tempera-
tures are shown in Figures 11-18.

In all five solvents used, i;g;, nitromethane, aceto-
nitrile, acetone, methanol, and pyridine, the curvature
in the mole ratio plots increases with increasing the

temperature. This behavior indicates the existence of

113

 

 

 

 

 

 

 

Table 25. Cesium-133 Chemical Shifts of 0.005M_Cs Ion

in the Presence of DB3OC10 at Various Tempera-

tures.
Solvent: Nitromethane
DB3QClO/Cs+

Temperature, °C
10 30 15 60 70
0.00 58.21 59.71 61.71 63.01 61.81 66.51
0.23 18.61 50.61 . 52.11 53.81 55.91 57.21
0.50 37.71 38.21 39.72 11.30 13.21 15.12
0.71 27.10 28.51 30.61 32.81 31.80 36.11
0.81 23.31 21.22 25.70 27.91 30.72 32.71
0.96 19.82 21.11 22.00 21.22 27.00 29.62
1.15 17.70 19.12 20.11 21.80 21.81 26.10
1 30 17.11 18.82 19.10 21.22 23.71 21.71
1.19 17.11 18.30 19.10 20.21 22.60 23.52
1.68 17.11 18.22 19.11 20.21 22.00 22.11
1.80 17.11 18.30 19.00 19.82 21.81 22.20
2.08 17.12 18.22 18.92 19.61 21.30 22.11
Solvent: Acetonitrile
0030010/Cs+
Temperature, °C

18 30 15 60 77
0.00 -35.10 -33.65 -32.11 -30.58 -26.00
0.25 -26.75 -25.21 -23.10 -22.33 —l9.00
0.50 -17.13 -16.20 -l1.73 -l1.01 ~12.33
0.59 -l1.81 -13.57 -12.71 -11.39 -9.85
0.70 -10.62 -9.77 -9.07 -8.36 -7.91
0.81 -6097 -6028 -60u3 -5082 -507“
0.98 -1.08 -0.85 -1.70 -1.62 -2.25
1.12 3.11 2.79 2.02 1.98 -0.10
1.19 1.12 1.12 3.26 2.18 1.10
1.38 7.37 6.98 6.13 5.27 1.19
1.60 8.61 8.51 8.00 7.29 6.05
2.00 9.55 9.55 9.71 9.28 8.02

 

111

 

 

 

Temperature, °C

20 30 50

 

29.13 30.87 33.16
25.78 26.87 29.73
21.52 22.38 21.77
20.20 20.97 23.38
18.19 19.66 22.13
17.31 18.26 21.06
17.03 17.87 20.58
16.10 16.87 19.53
15.70 16.18 18.80
15.65 16.11 18.26
15.63 16.25 18.03

 

 

Temperature, °C

30 15 60

 

Table 25. Continued.
Solvent: Acetone
0030010/Cs+

-10 10
0.00 25.13 28.11
0.30 21.91 21.16
0.61 17.87 20.20_
0.75 16.55 18.57
0.88 15.31 17.26
0.99 13.72 16.17
1.07 13.26 15.86
1.23 13.01 15.21
1.36 12.91 11.93
1.53 12.91 11.93
1.86 12.91 11.93
Solvent: Methanol
0030010/Cs+

0 15
0.00 12.10 11.37
0.22 35.05 37.69
0.56 25.05 27.38
0.80 17.69 20.71
0.86 17.00 19.97
0.98 13.97 17.30
1.07 13.19 16.25
1.26 13.15 15.11
1.37 13.11 15.28
1.60 13.11 15.15
1.71 13.11 15.08

16.00 17.23 50.25
39.75 11.18 15.38
30.02 32.26 37.16
23.71 27.30 32.11
23.31 26.25 31.25
20.17 21.71 28.95
19.15 23.85 27.99
18.26 21.50 26.15
17.61 20.57 25.67
17.11 19.93 21.13
17.11 19.75 23.97

 

115

Table 25. Continued.

 

 

Solvent: Pyridine

 

 

0030010/Cs+
Temperature, °C
0 10 3O 50 65 85

0.00 27.17 30.05 31.39 38.12 11.68 11.21
0.22 23.69 25.50 28.57 31.91 31.70 37.99
0.17 18.65 20.28 _ 23.77 26.09 27.95 30.82
0.66 11.61 16.56 18.57 20.82 22.52 25.17
0.80 11.99 12.60 11.77 16.18 17.71 21.59
0.86 11.81 11.82 13.69 15.50 16.86 20.13
1.03 9.02 9.80 11.11 13.06 11.16 17.25
1.10 9.25 9.73 11.20 12.76 11.32 17.22
1.27 9.19 9.72 10.90 11.83 13.22 16.01
1.15 9.17 9.62 10.66 11.75 12.81 15.38
1.77 9.19 1.61 10.51 11.11 12.28 11.78
2.18 9.18 9.59 10.51 11.10 12.15 11.37

 

 

116

 

 

 

 

 

NITROMETHANE
60k
50”-
E‘
o.
9:
co
40-
30-
700
so.
20 I- 450
it"
L l 1 L 1 l 1 L 1— 1 1
013(12 CMQ O£5<18 I!) L2. h4 Iii L8 2£>12
[ DB3OC10]/[Cs*]

Figure 11. Cesium-133 Chemical Shifts vs. [DB3OClO]/[Cs+]

Mole Ratio in Nitromethane at Different Tempera-
ures.

117

 

ACETONITRILE u;
m» f, £0
502
77
i,

7.2..) /

 

 

l l

0.0 0.5 Ito Ls 2.0
[0830001161

 

Figure 15. Cesium-133 Chemical Shifts vs. [DB3OC10]/[Cs+]

Mole Ratio in Acetonitrile EE'Different Tem-
peratures.

 

 

 

 

ACETONE
1
30
l
1
1
6
(.0me
20 '-
50°
30°
20°
- 4 : —: 10:
i t -10
10 '-
l l l i
as 1.0 1.5 20
[0830001165

 

 

0.0

Chemical Shifts v_s_. [DB3OClO]/[Cs+]
Mole Ratio in Acetone at Different Temperatures.

Figure 16. Cesium-133

119

 

 

 

 

5O
METHANOL
40 I
E‘
a.
2:
6°
30*-
60'
20" 45‘
30‘
ISP
0.
IO-
L l L J l I l l I
CLO 022(L4I(16 CHE Ix) L2 64 M5 L8
[08300011116]

Figure 17. Cesium-133 Chemical Shifts 1g. [DBBOClO]/[Cs+]
Ratio in Methanol at Different Temperatures.

120

 

 

 

 

PYRIDINE
4o1~
30
O
E:
a.
m
20-—
as-
659
so-
30-
IC-- 197
l l l l 1 l L l L l J
0.0 0.2 0.4 0.6 0.8 to L2 L4 l.6 L8 2.0 2.2
[moon/[cg]

Figure 18. Cesium-133 Chemical Shifts vs. [DB3OClO]/[Cs+]
Mole Ratio in Pyridine at DIITerent Temperatures.

121

an exothermic reaction between the cesium ion and the ligand.
The formation constants were computed by the KINFIT program
as described previously and the results are listed in Table
26. It is readily seen that lowering the temperature in-
creases the stability of the complex. In most cases at

lower temperatures the formation constants were greater

than 105 and their precise values could not be determined

by our technique.

Plots of in Kf vs. l/T for the five systems are shown
in Figure 19. The slopes and intercepts of the straight
lines are calculated by linear least squares fitting of
the data, and the corresponding thermodynamic parameters
are listed in Table 27. It is seen that the results ob-
tained in methanol solutions agree reasonably well with
the results of Chock (32). It is interesting to note that
while the stability (or AG° values) of the complex is not
very sensitive to the solvent (at least in the case of the
five solvents studied here), the enthalpy and the entropy
values vary Significantly with the solvent. Since the
cesium ion is rather weakly solvated because of the low
charge density of the cation, it is not surprising that
the free energy, AG°, of the complexation is only marginally
dependent on the nature of the solvent. However, the
solvation of the ligand would be different in different
solvents and because of that the enthalpy of the complexa-

tion is solvent dependent. The entropy of the complexation

122

Table 26. Formation Constants of DB30010-Cs+ Complex in

Nonaqueous Solvents at Different Temperatures.

 

 

 

SolvenSolvent Temp. (°C) Log Kf
Nitromethane 70 3.65:0.01
60 3.70:0.02
15 3.99:0.10
3O 1.3020.05
10 1.67:0.11
0 >
Acetonitrile 77 2.85:0.07
60 3.01:0.05
15 3.20:0.08
30 3.39i0.09
18 3.19:0.10
Acetone 60 3.1010.08
15 3.96:0.07
30 1.31:0.11
15 1.92:0.20
Methanol 60 3.36:0.08
15 3.70:0.05
30 1.18:0.07
15 1.65:0.02
0 >5
Pyridine 85 3.52:0.02
65 3.81:0.05
50 1.1320.01
30 1.1110.10
10 1.81:0.07

 

 

 

123

 

 

 

 

u—
IO _
Ian v
9 .—
P)’ NM
8" MeOH .
Ac
7 #-
Ahk
l l
2.5 3.0 3.5
IO'3/ T

Figure 19. Van't Hoff Plots for Complexation of Cs+ Ion
by Dibenzo-30-Crown—10 in Various Solvents.

121

Table 27. Thermodynamic Parameters for the Complexation
of Cs+ Ion by Dibenzo-30-Crown-10 in Various

 

 

 

Solvents.
AG°(30°C) AH° AS°

Solvent (kcal/mole) (kcal/mole) (cal/mole °K)
Nitromethane -5.97i0.07 -7.95:0.39 —6.66il.25
Acetonitrile -1.71i0.13 -5.1310.28 -1.5310.89
Acetone -5.5020.10 -l3.1810.51 r26.19il.69
Methanol -5.8liO.10 -l2.72iO.31 —22.82¢1.11
Methanola -5.77 -11.2 —18.2
Pyridine -6.1320.11 -7.9110.36 -5.93tl.13

 

 

aReference 32.

125

in various solvents increases in the order of acetone <
methanol < nitromethane < pyridine < acetonitrile. In
all cases the complexes are enthalpy stabilized but
entropy destabilized.

It should be noted that Similar behavior was previously
observed in nonaqueous solutions. For example, in the
case of the cryptate C222 exclusive complex with cesium

1

ion, the AH° and AS° values are -l2.9 kcal mole' and

1 and -13.7 e.u.

1

-26.8 e.u. in acetone, -8.6 kcal mole-
in propylene carbonate and —5.7 kcal mole' and -11.2 e.u.
in N,N-dimethylformamide solutions (96). Entropy destabiliza-
tion was also observed by Izatt 33 31. for the complexa-

tion of sodium and potassium ions by benzo-lS-crown—S

and 18—crown-6 in water—methanol mixtures (31) and by
Kauffmann gt_a1. for the complexation of potassium and
rubidium ions by cryptand C221 and Na+, I<+, Rb+, and

Cs+ ions by cryptand C222 in aqueous solutions (131).

In the last case, it was assumed that the decrease in entropy
was largely due to the rearrangement of the water structure
upon the metastasis of a small inorganic cation into a

large hydrophobic organic cation. While this explanation

is quite feasible for aqueous solutions, it cannot be

carried over to much less structured organic solvents used

in this study.

It seems reasonable to assume that the main reason for

the negative entropy of the complexation is the decrease

126

in the conformational entropy of the ligand upon the
formation of a metal complex. Large macrocycle ligand

such as DB3OC10 should be rather flexible in the free
state. The degree of flexibility would vary with the
solvent, i;g;, with the extent of ligand-solvent inter-
action. The formation of a rigid three-dimensional complex
should decrease the conformational entrOpy of the ligand
and thus, perhaps, give rise to negative entropy of the
complexation. At the present time, however, thermodynamic
data on the formation of macrocyclic complexes in non-
aqueous solvents are quite sparse. Additional work is

very necessary before the entropy destabilization of macro-
cyclic ligands in nonaqueous solvents can be explained

satisfactorily.

1.3. 002108 COMPLEXES WITH Cs+

In order to study the thermodynamic behavior of the
complexation reaction between cesium ion and dibenzo-21—
crown-8, extensive chemical shift-mole ratio studies were
made of the Cs+-DB21C8 system in nitromethane, acetonitrile,
acetone, methanol, and pyridine solutions over a wide
temperature range. The measured chemical Shifts for dif-
ferent ligand to metal ion mole ratios at various tempera-
tures are given in Table 28 and the corresponding mole

ratio plots are shown in Figures 20-21. In all systems

Table 28. Cesium-133 Chemical Shifts of 0.005M 0s+ Ion

127

in the Presence of DB21C8 at VariouS'Tempera-

tures.

 

 

Solvent: Nitromethane
DB21C8/05+

 

Temperature, °C

 

 

 

 

20 30 15 60 75 90

0.00 56.87 57.57 58.17 61.15 62.16 61.55
0.21 52.16 53.31 51.32 56.18 57.11 59.12
0.60 15.09 15.62' 16.61 17.61 18.88 51.06
0.70 13.51 11.38 11.91 16.17 17.59 18.82
0.91 10.13 11.37 12.29 13.16 15.09 16.87
1.05 39.60 10.59 11.55 12.70 11.21 11.55
1.10 39.31 10.12 10.90 12.38 11.00 15.00
1.18 39.11 39.86 10.59 11.83 13.19 11.13
1.10 38.12 39.01 39.66 10.66 11.79 13.38
1.87 38.03 38.58 39.29 39.97 10.78 11.98
2.39 37.96 38.12 39.07 39.71 10.28 11.60
2.88 37.96 38.12 39.00 39.53 10.11 11.15
Solvent: Acetonitrile

002108/Cs+

Temperature, °C
5 30 10 50 60 75

0.00 -37.35 -33.55 -32.08 -30.15 -29.29 -27.27
0.25 -26.65 23.21 -22.51 -21.11 -20.15 -19.52
0.62 -8.97 -7.50 -7.26 -7.12 -7.03 -6.57
0.72 -1.17 -3.70 -3.17 -3.85 -3.69 -3.62
0.81 0.72 -0.20 0.56 0.10 0.09 —0.21
0.98 6.16 5.85 5.00 1.36 3.71 2.73
1.09 9.25 8.32 7.51 6.77 6.16 5.69
1.13 9.88 8.55 7.93 7.08 6.96 5.90
1.38 12.51 12.12 11.73 10.61 10.80 9.87
1.78 13.30 13.68 13.05 12.93 12.81 11.57
2.32 13.11 11.11 13.90 13.85 13.75 13.67
2.78 13.11 11.22 13.95 11.10 11.21 11.35

 

128

 

 

 

 

 

 

 

Table 28. Continued.
Solvent: Acetone
DB21C8/03+
Temperature, °C

5 15 30 10 55
0.00 15.61 17.80 19.27 20.30 22.12
0.25 18.78 20.25 21.11 21.81 21.06
0.60 22.13 23.12 23.82 21.56 26.32
0.77 23.98 21.51 . 25.15 25.13 26.92
0.98 26.16 26.53 26.69 27.22 27.91
1.12 27.39 27.60 27.70 28.21 28.18
1.25 28.01 28.21 28.16 28.50 29.03
1.13 28.16 28.15 28.18 29.06 29.10
1.81 28.17 28.68 28.79 29.60 29.72
2.19 28.20 28.72 29.02 29.83 30.10
2.66 28.21 28.81 29.01 29.98 30.29
Solvent: Methanol
DB21C8/03+

Temperature, °C

10 20 30 10 50 65
0.00 11.11 15.31 16.15 17.00 17.83 18.90
0.20 12.10 13.56 11.76 15.12 16.13 17.10
0.51 39.81 10.70 11.50 12.63 13.57 11.72
0.78 38.21 39.11 10.03 10.77 11.70 12.79
0.91 37.28 38.13 39.02 39.77 10.77 11.93
1.02 37.12 37.80 38.63 39.15 39.87 11.62
1.18 36.56 37.20 37.87 38.15 39.21 10.17
1.31 36.35 36.65 37.39 38.16 38.67 10.06
1.60 36.00 36.15 37.09 37.90 38.26 39.16
1.76 35.96 36.35 36.92 37.60 37.83 38.81
2.07 35.85 36.25 36.79 37.30 37.13 38.20
2.78 35.85 36.20 36.70 37.02 37.10 37.80

 

129

 

 

 

 

Table 28. Continued.
Solvent: Pyridine
DB21C8/08+
Temperature, °C

30 15 50 75 90
0.00 23.99 26.93 30.50 31.22 37.71
0.38 22.82 25.21 27.63 30.66 33.85
0.56 22.11 21.56 26.63 29.26 31.80
0.83 21.75 23.22 ‘ 21.69 27.17 29.81
0.92 21.36 22.83 21.25 26.70 28.72
1.05 21.19 22.13 23.68 25.61 27.87
1.26 20.98 21.59 22.77 21.07 26.21
1.52 20.88 21.28 22.60 23.16 21.92
2.03 20.81 21.20 22.36 23.01 21.25
2.39 20.79 21.18 22.28 22.88 21.13
2.83 20.79 21.13 22.20 22.80 23.96

 

 

 

130

 

 

 

 

 

as
NITROMETHANE

45 -
OC F
:1

40" so 1
4s
20

35 - .

, .1 I 14 11 L 1

0.0 0.5 In 1.5 2.0 2.5 3.1
[0824C81/[CJ]

Figure 20. Cesium-133 Chemical Shifts vs. [DB21C8]/[Cs+]
Mole Ratio in Nitromethane 56 Different Tem—
peratures.

20

Figure 21.

131

 

 

 

 

- ACETONITRILE

J J

 

 

L
0.0 0.5 l 0 1.5

. zfo 2.5
[0824C8]/[C§ ]

Cesium-133 Chemical Shifts E. [DB21C8]/[CS+]
Mole Ratio in Acetonitrile at Different Tem-
peratures.

10

132

 

30 - 33
/ an
a“"’

15

 

 

 

 

ACETONE
25
8

PPM
20

Is

‘ I I I I I
0.0 0.5 III 1.5 2.0 2.5
[D824C81/[Cffl

Figure 22. Cesium-133 Chemical Shifts vs. [DB21C8]/[Cs+]
Mole Ratio in Acetone at Different Tempera-

tures.

10

 

133

 

 

 

 

 

 

 

0
5 METHANOL
45 ‘
5
ppm
40 -
°c
55
A 30
:.2
A 13
35 -
I L J . 1
0-0 0.5 1-0 1.5 2.0 2.5 3.0
[0824C81/[CJ]

Figure 23. Cesium-133 Chemical Shifts vs. [Dszucej/[Cs+]

Mole Ratio in Methanol at Di?ferent Tempera-
tures.

134

 

35

30

Ppm

25

 

PYRIDINE

O)
K I
3 no

 

 

 

75
50
45
se— {o—————-030
20 I I J I J
an &5 L0 L5 LB 15 10
[0324C81/ [CQ‘ ]

Figure 2D. Cesium-133

tures.

Chemical Shifts 15. [D32u08J/ECS+]
Mole Ratio in Pyridine at Different Tempera-

 

135

studied, the curvature in the mole ratio plots is more
pronounced at lower temperatures. This behavior implies
that the complex formed is more stable at lower tempera-
tures and that, of course, the complexation reaction is
exothermic.

The calculated formation constants of DBZDCB-Cs+
complex in different solvents at various temperatures are
listed in Table 29. It is seen that such as with DB3OC10
the stability of the complex drastically increases as the
temperature is lowered. Plots of fin Kf XE' the reciprocal
of absolute temperature gave straight lines in all solvents
used (Figure 25) and the AH° and AS° values were determined
in the usual manner for the slope and the intercept of the
plots. The calculated thermodynamic parameters for the
complexation in various solvents are listed in Table 30.
The data clearly show that in all cases studied the complexa-
tion reaction is enthalpy stabilized but entropy destabi-
lized. While the stability of DB2llD8-Cs+ complex (or AG°)
is only marginally dependent on the nature of the solvent,
the changes in enthalpy and particularly in the entropy of
the complexation are very solvent dependent. The entropy
change for the complexation in various solvents decreases in
the order of pyridine > nitromethane > acetonitrile >
methanol > acetone.

Izatt and coworkers have studied the thermodynamics of

the complexation of the cesium ion with DB2HC8 in 70%

136

Table 29. Formation Constants of DB2H08-CS+ Complex in
Nonaqueous Solvents at Different Temperatures.

 

 

 

Solvent Temp. (°C) Log Kf
Nitromethane 90 3.37:0.05
75 3.52:0.03

6O 3.68:0.0h

MS 3.91:0.03

30 h.11i0.08

20 U.26i0.06

Acetonitrile 75 3.19:0-03
6O 3.u5eo.ou

50 3.5710.05

no 3.77:0.07

30 3.9hio.07

5 M.50i0.05

Acetone 55 3.07:0.03
no 3.37:0.07

30 3.71:0.09

15 H.15i0.10

5 u.37:o.o7

Methanol 65 2.86:0.06
50 3.11:0.0u

HO 3.36:0.10

3O 3.65:0.05

20 3.85:0.06

10 h.0hi0.10

Pyridine 9O 3.27:0.09
75 3.uuio.o7

60 3.60:0.0u

us 3.76:0.08

30 u.OOi0.03

 

 

 

137

 

 

 

 

 

IIb
10‘
9 .—
{2‘
.E
8 .-
NM
PY
AN
7—
MeOH
J l l l L
as 23 33 12 3A 16

103/T

Figure 25. Van't Hoff Plots for Complexation of Cs+ Ion
by Dibenzo-2u-Crown-8 in Various Solvents.

138

Table 30. Thermodynamic Parameters for the Complexation
of Cs+ Ion by Dibenzo-2u-Crown-8 in Various

 

 

 

Solvents.
AG°(30°C) AH° AS°
Solvent (kcal/mole) (kcal/mole) (cal/mole °K)
Nitromethane -S.7liO.ll -6.25i0.10 -l.7910.3u
Acetonitrile -5.h710.10 -8.12i0.16 -8.6610.U8
Acetone -5.15i0.13 -ll.20t0.52 -20.08il.7l
Methanol -5.06:o.07 -9.87io.46 —l6.lo:l.u9

Pyridine -5.55i0.05 -5.97i0.16 -l.39i0.u6

 

 

139

methanol solution by calorimetric titration. They also
found negative values for the changes in the enthalpy
and the entrOpy of the reaction (i.e., AH° = -8.09 kcal

l and AS° = -lu.l e.u.). Since dibenzo-Zh-crown—lo

mole-
is a rather flexible ligand in free state, it is reason-
able to expect that a change in conformation of the ligand,
from a "loose" structure in the free state to a "rigid"

structure in the complex, contributes to the large negative

entropy changes of the reaction as a dominant factor.

u.u. DB2lC7 COMPLEXES WITH Cs+

The thermodynamics of the complexation of the cesium
ion with dibenzo-21-crown—7 was investigated in nitro-
methane, acetonitrile, acetone, methanol, and pyridine by
cesium-133 NMR. The variation of the chemical shift with
the changing temperature was studied for the complexation
reaction. At each temperature the cesium-133 chemical shift
was monitored as a function of DBZlC7/Cs+ mole ratios. The
data are given in Table 31 and the mole ratio plots are
shown in Figures 26-30. As the temperature increases, the
mole ratio plots show less curvature indicating the for-
mation of a weaker complex at higher temperatures. This
trend is evidence of the existence of an exothermic reac-
tion between the cesium ion and the ligand. The calculated

formation constants show the same trend (Table 32).

1“0

 

 

 

 

 

 

 

Table 31. Cesium-133 Chemical Shifts of 0.005M’Cs+ Ion
in the Presence of DB21C7 at Various Tempera-
tures.

Solvent: Nitromethane
DB2lC7/Cs+ Temperature, °c

15 30 “5 60 75 90
0.00 56.95 57.89 60.86 63.38 6“.39 65.81
0.26 “9.65 50.“3 53.01 5“.75 56.08 57.57
0.50 “2.0“ “3.37 ““.79 “6.12 “8.16 “9.33
0.77 33.57 35.22 36.86 38.51 “0.00 “2.59
0.9“ 29.33 30.75' 32.5“ 3“.59 36.23 38.58
1.11 25.80 27.37 29.“9 31.13 33.72 35.68
1.28 23.68 25.“l 27.05 28.78 30.82 33.02
1.“8 23.“5 2“.86 26.27 28.07 30.0“ 31.92
1.80 23.1“ 2“.2“ 25.72 27.06 28.“7 30.27
2.25 23.06 2“.15 25.02 26.67 27.1“ 28.55
2.67 23.05 2“.07 2“.95 26.27 26.90 28.00
Solvent: Acetonitrile

+
DB21C7/Cs Temperature, °C

20 30 “0 50 60 75
0.00 -33.2“ -31.93 -30.05 -28.““ -27.“3 -2“.9l
0.35 -25.96 -2“.79 -23.02 -21.86 -20.91 -19.13
0.6“ -l9.99 -l9.06 -l7.7“ -l6.72 -l6.02 -l“.79
0.82 -l6.33 -15.“9 -l“.63 -l3.70 -l3.3“ ~12.“7
0.92 -l“.32 -l3.“7 -l2.6“ -l2.23 -ll.53 -ll.l5
1.0“ -l2.30 -11.76 -ll.00 -10.76 -10.1“ —9.98
1.22 -9.98 -9.“7 -8.73 -8.35 —7.83 -7.57
1.33 -9.7“ -9.36 -8.27 -7.80 -7.“2 -7.19
1.50 -9.67 -9.05 -8.05 -7.“9 -6.96 -6.“9
1.83 -9.“5 -8.7“ -7.73 -7.11 —6.25 -5.71
2.52 -9.20 -8.52 -7.“2 -6.“1 -5.6“ -5.00
3.21 -9.1“ -8.“0 -7.18 -6.15 -5.27 -“.30

 

1“l

 

 

 

 

 

 

 

Table 31. Continued.
Solvent: Acetone
DB21C7/Cs+
Temperature, °C

10 20 30 “0 55
0.00 16.67 17.86 19.“8 21.26 22.7“
0.30 13.83 1“.98 16.“6 18.16 19.56
0.56 11.10 l2.“3 13.82 15.60 17.16
0.79 9.02 10.26 11.65 13.35 15.22
0.89 8.01 9.“0' 10.72 12.58 l“.06
1.03 7.00 8.32 9.67 11.3“ 13.36
1.12 6.37 7.5“ 8.9“ 10.80 12.58
1.20 6.15 7.39 8.78 10.65 12.35
1.“3 6.06 7.23 8.“2 10.13 11.73
1.90 6.00 7.07 8.2“ 9.67 ll.“2
2.37 5.98 7.03 8.20 9.“5 11.27
3.16 5.96 6.99 8.10 9.35 10.90
Solvent: Methanol
DB2lc7/Cs+

Temperature, °C

20 30 “0 50 60
0.00 “5.“3 “6.97 “8.76 “9.92 50.50
0.30 38.13 39.53 “1.01 “1.85 “2.55
0.5“ 31.85 32.86 3“.3“ 35.65 36.82
0.68 28.60 29.92 31.07 32.55 33.86
0.90 23.17 2“.6“ 26.66 27.75 29.60
1.03 21.00 22.“0 2“.65 25.81 27.82
1.1“ 20.30 21.55 23.“0 2“.88 26.82
1.28 l9.“5 20.62 22.16 23.32 25.26
1.“6 19.00 19.90 21.39 22.“7 2“.33
1.85 18.83 19.66 20.82 21.70 23.10
2.“2 18.60 19.53 20.“5 21.23 22.80
3.0“ 18.53 19.3“ 20.30 21.00 22.“0

 

1“2

 

 

 

 

Table 31. Continued.
Solvent: Pyridine
DB2lCZ/Cs+
Temperature, °C

30 “0 50 60
0.00 2“.53 27.08 29.18 31.90 3“.92 38.50
0.17 19.80 22.““ 2“.77 27.18 30.20 31.97
0.50 11.12 12.83 15.07 17.“0 20.19 23.06
0.77 3.91 5.61, 7.2“ 9.10 12.20 l“.00
0.91 -0.67 1.03 2.“3 “.1“ 7.39 10.93
0.99 -2.30 -0.60 0.“9 2.89 5.5“ 8.“9
1.1“ -5.“7 -3.77 -2.68 -0.75 1.55 6.00
1.29 -5.63 -3.93 -3.39 -l.60 0.95 3.““
1.51 -5.78 -“.63 -3.93 -2.30 -0.36 1.66
1.88 -5.82 -5.06 —“.50 -2.69 -1.07 1.0“
2.37 -5.9“ -5.16 -“.72 -3.07 —1.36 0.15
2.97 -6.05 -5.32 -“.96 -3.“0 -2.03 -0.61

 

 

l“3

 

 

 

 

NITROMETHANE
so
so -
6
Ppm
4o -
an - .C
so
75
so
45
so
15
0.0 0.5 {a is in 275 3.0
[banal/[Cf]

Figure 26. Cesium-133 Chemical Shifts 13. [DB2107]/[Cs+]

Mole Ratio in Nitromethane at Different Tem-
peratures.

1““

 

 

 

 

351
-‘-‘75
4360
~—£50

 

 

1_._,l *rOZO

 

 

ACETONITRILE
-35 h-
I l i 1 L l
80 05 L0 L5 to 25 30
[0321014qu
Figure 27. Cesium-133 Chemical Shifts vs. [DB21C7]/[Cs+]

Mole Ratio in Acetonitrile at Different Tem-
peratures.

1“5

 

 

 

 

 

 

 

 

 

25‘
ACETONE
20
I5 -
5
ppm
°c
~—055
H)-
——o40
——030
420
as: at e—clo
5..
l l J I l l
00 05 l 0 l 5 20 25 3 0
[banal/[CU

Figure 28. Cesium-133 Chemical Shifts vs. [DB2107]/[Cs+]
Mole Ratio in Acetone at Different Tempera-
tures.

1“6

 

 

 

 

 

 

 

 

 

 

 

 

501
‘ METHANOL
40-
5
ppm
so -
°c
r-OBO
A 50
20" ' *r'40
+ —o 30
‘—'—F* ——020
l J l 1 J I
50 05 10 L5 21| Z5 10
[0521C71/[C9‘]

Figure 29. Cesium-133 Chemical Shifts vs. [DB21C7]/[Cs+]
Mole Ratio in Methanol at DIfferent Tempera-
tures.

1“7

 

 

 

 

 

 

 

 

PYRIDINE
30
20 '-
8
PP"‘
10 '-
_ °C
0 so
r-o75
50
50
1 1 I 4 3L
0.0 0.5 1.0 1.5 2.0 2.51 3.0
[ 0521C 7] /[Cs*]

Figure 30. Cesium-133 Chemical Shifts vs. [DB2107]/[Cs+]

Mole Ratio in Pyridine at Different Tempera—
tures.

l“8

Table 32. Formation Constants of DBZlC7-Cs+ Complex in
Nonaqueous Solvents at Various Temperatures.

 

 

 

Solvent Temp. (°C) Log Kf
Nitromethane 90 3.21:0.03
75 3.39:0.05
60 3.66:0.0“
“5 3.81:0.05
30 “.1“io.07
15 “.“0i0.10
Acetonitrile 75 3.15:0.05
60 3.“1i0.03
“0 3.73:0.05
30 3.95:0.0“
20 “.11i0.05
Acetone 55 3.36:0.06
“0 3.6“:0.03
30 3.93:0.06
20 “.l9i0.0“
10 “.52:0.0“
Methanol ' 60 3.5“:0.0“
50 3.68:0.02
“0 3.83:0.02
30 3.96:0.06
20 “.1“i0.02
Pyridine 90 3.39:0.05
75 3.5710.07
60 3.78:0.05
50 3.89:0.0“
“0 “.07:0.05
30 “.27i0.07

 

 

l“9

Plots of 1n Kf vs, 1/T for the five systems are shown

in Figure 31. The enthalpies and the entropies of the
complexation were obtained in the usual manner from the
slopes and the intercepts of the plots and the results

are listed in Table 33. As is seen, in all cases the com-
plexes are enthalpy stabilized but entropy destabilized.
The enthalpy and entropy values vary very significantly
with the solvent, but in all cases they compensate each
other resulting in a nearly identical free energy for the
complexation. The sequence of the increase in the entropy
change in different solvents is acetone < acetonitrile <

nitromethane < pyridine < methanol.

“.5. DISCUSSION

A deeper understanding of the thermodynamics of the
complexation of the metal ions with macrocyclic crown
ethers can be provided by the evaluation of the enthalpy
and the entropy changes of the reaction. The magnitudes
of enthalpy values are indicative of the metal ion-ligand
interaction providing information about the type and the
'number of binding sites. The magnitude of the entropy
values are indicative of the solvent-solute interactions.
It can supply information about the relative degrees of
solvation of the particles involved (lLSLfi the metal ion,
the ligand, and the complex), the less of degrees of

freedom of the ligand upon complexation, and the Charge

150

 

 

 

 

IO—
9.—
O
¥Hu
15
0
8- MeO
P)’
Ac
NM
AN
7 I l I l
2.8 3.0 3 3.2 3.4 3.6
IO/T

Figure 31. Van't Hoff Plots for Complexation of Cs+
Ion by Dibenzo—Zl—Crown-T in Various Solvents.

151

Table 33. Thermodynamic Parameters for the Complexation
of Cs+ Ion by DBZlC7 in Nonaqueous Solvents
at Different Temperatures.

 

 

 

Solvent (figzlggggg) (kc533;51e> (cal/figie °K)
Nitromethane -5.75i0.10 -7.61i0.25 -6.25i0.78
Acetonitrile -5.“9i0.05 -8.2“i0.20 -9.1910.66
Acetone -5.“6i0.08 -11.13i0.“1 —18.6“il.35
Methanol -5.50i0.08 -6.6liO.1“ -3.5“iO.““

Pyridine -5.93i0.10 -7.22i0.18 -“.38i0.56

 

 

152

types involved in the reaction.

In all three complexation reactions studied in this
thesis, 1,24, complexation of the cesium ion with DB30010,
DB2“C8, and DB2107, the contributions of the metal ion
solvation and its charge type to the entropy changes in
the same solvants are the same. Whereas, because of the
difference in the size and in the number of the binding
sites of the ligands, the contributions of the conformational
change of the ligand and of the solvation of the ligand and
the complex to the entropy changes are different. Thus in
order to interpret the thermodynamic data, these factors
must be considered.

In all cases studied, the complexes are enthalpy
stabilized but entropy destabilized. The entropies of the
complexation of the cesium ion with DB30010, DB2“08, and
DB21C7 in various solvents are compared in Table 3“.

The solvation of the free and of the complexed cesium ion
could be very different so that the complexation can in-
fluence the structure of the solvent which could contribute
to the entropy of the complexation. Since the solvents

we worked with are less structured than water, the contribu-
tion of this factor to the entropy would not be very im-
portant. However, at the present time we have no knowledge
about the extent of solvation of either the ligands or

the complexes.

A more important contribution to the negative entropy

 

 

153

 

 

 

 

Table 3“. Entropies of the Complexation of Cesium Ion by
DB3OC10, DB2“C8, and DB21C7 in Various Solvents.
AS°(ca1/mole °K)

Solvent DB3OC10-Cs+ DB2“C8-Cs+ DB21C7-Cs+
Nitromethane -6.66il.25 -1.79i0.3“ -6.25i0.78
Acetonitrile -1.53i0.89 -8.66i0.“8 -9.19i0.66
Acetone -26.19i1.69 -20.08il.71 -l8.6“i1.35
Methanol -22.82¢1.ll -l6.lOil.“9 -3.5“i0.““
Pyridine -7.9“10.36 -5.93il.13 -“.38:0.56

 

 

 

15“

values would be the change in conformation of the ligands,
from a flexible structure in the free state to a more rigid
form, upon complexation. It is interesting to note that,
with the exception of acetonitrile, in all solvents used,
first, the entropy values of the DB2“C8°Cs+ complex are
about “ to 6 e.u. less negative than those for DB3OC10-Cs+
complex, and second, the sequence of the solvent effect

on the entropy values is the same for both complexes.
According to these results, it seems reasonable to assume
that among the various factors contributing to the negative
entropies of the complexation of the cesium ion with DB30010
and DB2“C8, the decrease in the conformational entropy of
the ligand upon complexation is the dominant one. Large
macrocyclic ligands such as DB3OC10 and DB2“C8 are rather
flexible in free state. The degree of flexibility would
vary with the size of the ligands and with the solvent
(Elia: with the extent of ligand-solvent interaction).
DB30010 is more flexible than DB2“C8 in free state because
of its larger size, and also DB3OClO-Cs+ complex is probably
more rigid than DB2“C8-Cs+ complex because of its complete
"wrap around" structure (Chapter 3). Therefore, the more
negative entropy values for DB3OC10-Cs+ complex than those
for DB2“C8-Cs+ are not unexpected. Acetonitrile as solvent,
however, is an exceptional case probably because it can
form a complex with the ligands, such as that reported

with 18-crown-6 (135), and, therefore, can deviate the

.-___.-___—--— "

 

 

155

results from the expected trend.

0n the other hand, in the case of the complexation of
the cesium ion with DB2lC7, neither the sequence of the
solvent effect on the entropy values nor the trend of the
entropy values, i;gL, decreasing the entropy with increasing
the size of the ligand, agree with the DB3OC10 and DB2“C8
cases. Since the cesium ion, with diameter of 3.68 A (128),
has a very close size to that of DB2lC7 cavity, with the
size of 3.“-“.3 K (23), it seems reasonable to assume that
the DB21C7-Cs+ complex has a two-dimensional structure.
Therefore the large conformation changes involving the
ligand "wrapping around" the cation such as one observed
of DB30010 are not expected in this ligand. Thus, unlike
the cases of DB30010 and DB2“08, the conformational change
of the ligand upon complexation is not necessarily the
dominant factor in this case. Such an example can be found
in the literature. Dibenzo—30-crown-10 (32) and 18-crown-6
(62) form complexes with the potassium ion in methanol
solution with about the same entropy values of -l7.7 e.u.
and -17.3 e.u. respectively. If the decreased conforma-
tional change of the ligand upon complexation was the
dominant contribution to the entropy values in both cases,
we would expect the dibenzo-30-crown-10°Cs+ complex to
have a much more negative entropy value than 18-crown-6-Cs+.
It has been shown that potassium ion forms a three-dimen-
sional "wrap around" complex with dibenzo-30-crown—10 (26,33)

but a two-dimensional one with l8-crown-6 (136).

 

CHAPTER 5

LITHIUM-7, SODIUM-23, CESIUM-133, AND THALLIUM-205
NMR STUDY OF Li+, Na+, 03+, and Tl+ ION COMPLEXES
WITH 1,10-DIAZA-18-CROWN-6 IN VARIOUS NONAQUEOUS

SOLVENTS

156

5.1. INTRODUCTION

The macrocycle 1,10-diaza-18-crown-6 was first

synthesized by Dietrich gt_a1. (137).

1.10-DIAZA-Is-CROWN-s

Frensdorff (“) has shown that the substitution of two
nitrogens for two oxygens in 18-crown-6 reduces the affin—
ity of the ligand for alkali metal ions while enhancing the
stability of the complexes of transition metal ions of the
same size. Thermodynamic parameters for the complexation
of some alkali earth and transition metal ions with this
ligand in aqueous solution have been reported by Andereg
(138) who used pH-metric titration for the measurements.
Structural properties of some metal ion complexes with
1,10-diaza-l8-crown-6 in solution have been studied by
proton NMR (l39,1“0) and the crystalline structure of the
isolated complexes of the ligand with copper (II) and
potassium ions have been determined (l“l,l“2). Copper (II)
was shown to be located inside the cavity of the macro-
cycle and it is bonded to the two nitrogen atoms and to
the two oxygen atoms. The potassium is bonded to the four

oxygen atoms of the ring in the same plane and to the two

157

158

nitrogen atoms from the top and the bottom of the macro-
cycle plane.

The present work was undertaken to determine the
stoichiometry and the formation constants of 1,10-diaza-
18-crown-6 complexes of Li+, Na+, 05+, and Tl+ ions in

various nonaqueous solvents.

5.2. RESULTS

Lithium-7, sodium-23, cesium-133, and thallium-205
chemical shifts were measured as a function of the
ligand/metal ion mole ratio in various solvents and the
results are given in Tables 35-38. In all cases only one
population average resonance of the metal ion was observed.
Generally, this is only possible if a fast exchange exists
between the two sites ($LEL, free and complexed ion) whose
rate is larger than J27wAv (Av is the difference between
the resonance frequency of each site). The chemical shift-
mole ratio plots for different metal ions are shown in

Figures 32-35-

5.2.1. 1,10-Diaza-18—Crown-6 Complexes with Li+

The frequency of the lithium-7 resonance in DMF, DMSO,
and TMG was found to be independent of the ligand/lithium
ion mole ratio (Figure 32). This behavior shows that the

immediate environment of Lithium ion is not changed upon

159

Table 35. Mole Ratio Study of 1,10-Diaza-18-Crown-6 Com-
plex with O-O2M.L1010“ in Various Solvents at

 

 

 

 

 

 

 

 

30°C.
NM AN PC Ac

+ + + +
L/Li 5(ppm) L/Li 5(ppm) L/Li 5(ppm) L/Li 5(ppm)
0.00 0.u3 0.00 2.56 0.00 0.63 0.00 —1.00
0.31 0.28 0.31 1.8“ 0.35 0.u3 0.39 -0.87
0.69 0.06 0.76 0.81 0.69 0.26 0.77 -0.73
0.88 -o.07 0.88 - 0.56 0.90 0.12 0.90 -0.69
1.00 -0.10 1.02 0.27 1.0“ 0.09 0.99 -0.67
1.10 -0.12 1.11 0.20 1.23 0.08 1.1“ -0.6&
1.33 -0.10 1.33 0.16 1.38 0.05 1.23 -0.62
1.60 -0.11 1.77 0.18 1.66 0.05 1.73 -0.52
2.1“ -0.11 2.u1 0.16 1.88 0.0“ 2.18 -0.u8
3.36 -0.11 3.2“ 0.16 2.31 0.06 3.“3 =0.u5

2.85 0.05
3.““ 0.05

DMF DMSO TMG Py

+ + + +
L/Li 5(ppm) L/Li 5(ppm) L/Li 5(ppm) L/Li 5(ppm)
0.00 -0.63 0.00 1.03 0.00 -0.u5 0.00 -2.“1
0.38 -0.62 0.38 1.0“ 0.38 -o.uu e.uu -2.2u
0.75 -0.6u 0.75 1.0“ 0.75 —0.u6 0.67 -2.16
0.86 -0.62 0.87 1.06 0.88 -0.“7 0.91 -2.06
1.02 —0.63 1.08 1.0“ 1.00 -0.u5 1.02 -2.01
1.11 -0.62 1.16 1.06 1.11 -o.uu 1.10 -1.98
1.22 -0.63 1.u0 1.03 1.23 -0.“6 1.32 -l.93
1.57 -0.6“ 1.66 1.05 l.“3 -0.“6 l.“3 -l.88
2.07 -0.63 1.99 1.05 1.90 -0.uu 2.05 -1.66
3.13 -0.6“ 3.35 1.0“ 2.90 -0.u5 3.10 -1.u3

 

 

Id

 

160

Table 36. Mole Ratio Study of 1,10-Diaza—18-Crown—6 Complex
with 0.05M.NaBPhu in Various Solvents at 30°C.

 

 

 

 

 

 

 

 

NM Ac
+ +
L/Na 6(ppm) Avl/2(Hz) L/Na 6(ppm) Avl/2(Hz)
0.00 1“.2l 13 0.00 6.93 18
0.32 11.97 11“ 0.“1 7.52 “1
0.70 9.“6 201 0.80 8.1“ 67
0.91 7.89 260 0.90 8.26 68
1.02 7.31 273 1.00 8.29 73
1.19 6.69 268 1.11 8.32 76
l.“7 6.88 27“ 1.22 8.32 79
1.8“ 6.83 276 1.“6 8.35 83
2.“8 6.79 276 2.11 8.“0 85
2.97 6.75 298 3.10 8.56 93
DMSO Py

L/Na+ 5 Av (Hz) L/Na+ 5 Av (Hz)

(ppm) 1/2 (ppm) 1/2
0.00 0.0“ “6 0.00 -1.15 21
0.“O 0.5“ 55 0.29 1.“3 55
0.75 0.86 68 0.70 “.81 10“
0.88 0.90 67 0.86 6.70 121
0.96 1.02 7“ 0.9“ 7.08 137
1.09 1.11 7“ 1.09 7.““ 135
1.23 1.20 78 1.23 7.62 137
1.“9 1.26 80 1.55 7.70 1“6
2.02 1.50 96 2.09 7.66 1““
2.68 1.81 105 2.57 7.67 155

 

 

 

161

 

 

 

 

 

ma.:mu um.»

mm.mmn mm.s sm.mmu om.s mo.mmn os.s

os.Hm- ew.a am.Hmu om.m aa.mmu mo.m

:m.mmu mm.» Hm.san mm.m os.omu sm.m mm.omu nm.m
mo.mmn mm.: mm.oan mm.m so.man mm.H H~.man m=.m
mm.wan om.m H=.:u mm.H mm.sau ma.H s:.mau mm.m
ms.smu ma.~ mm.m m=.H we.mau om.H mo.eau mm.a
we.mmu =m.a mm.a mm.a as.:=- mo.H NH.HHu mm.a
eo.smu sa.a Hm.m ea.fl ms.men no.0 mH.mu o=.H
mm.mo- mm.H OH.OH mm.o ms.men mm.o mo.m om.a
sm.ms- mo.H Hm.mH mm.o mm.m=u ms.o me.a mo.a
am.msu mm.o =m.mH ms.o mo.mmu mm.o mm.m ma.o
no.2mn No.0 os.ma mm.o ea.mmu m=.o mm.mH mm.o
sa.mm- am.o am.mm mm.o mH.omu om.o sm.oa mm.o
mo.ooau oo.o om.am oo.o He.mmu 00.0 mm.:m oo.o
Asaavm +mo\q Aeddvo +m0\q AEqavm +mo\q Aeaavo +m0\q

ooze mom ozg
.ooom om mucosaon
aaoaso> ca soH no spas onoEoo mucsonoumauousaouOH.H eo aospm ofioom oaoz .sm oases

 

162

.300000.030.0o

.20000_ma0.00

 

 

 

 

 

 

00.03: 00.0 00.00- 00.0 03.00: 03.0
00.03- 00.3 30.00- 00.3 03.33- 00.0 00.00: 00.3
03.03- 03.0 30.00: 00.0 00.00: 03.0 00.0- 03.3 00.30: 00.0
00.03: 00.0 03.00: 00.3 00.00: 00.0 00.0- 03.0 00.03- 30.0
00.03: 00.0 00.00- 30.0 00.00- 03.0 00.0: 00.0 00.03- 33.0
00.00: 00.3 03.00- 30.0 00.00- 00.0 00.3: 03.0 00.30: 00.3
00.00: 30.3 00.00: 00.3 00.00- 00.0 00.3: 00.3 00.03- 00.3
30.00: 00.3 00.00-: 00.3 30.00: 00.3 00.0- 00.3 00.0: 00.3
00.30: 30.3 03.003: 03.3 30.00: 03.3 00.0: 00.3 00.0- 03.3
00.03: 00.0 30.003- 00.0 30.00: 30.3 00.0: 00.0 00.0- 00.0
30.03: 00.0 00.303: 00.0 00.00- 00.0 00.3- 30.0 00.0: 00.0
00.0: 00.0 «0.003-. 30.0 00.00- 00.0 00.3: 00.0 33.3 00.0
03.0 33.0 00.303: 00.0 03.00- 03.0 00.0- 00.0 03.0 03.0
00.00 00.0 03.003- 00.0 00.00: 00.0 03.0: 00.0 00.03 00.0
Asaavo +m0\q Aeaavm +mo\q Aegavm +mo\q “agave +m0\q Aegavo +mo\a
000 0020 00020 0020
.000030000 .00 03000

163

Table 38. Mole Ratio Study of 1,10-Diaza-18-Crown-6 Com-
plexes with T1+ Ion in Various Solvents at 30°C.

 

 

 

 

 

 

 

 

NMa ANb Aca
+ + +

L/Tl 6(ppm) L/Tl 6(ppm) L/Tl 6(ppm)
0.00 36“.9 0.00 213.6 0.00 171.3
0.93 273.6 0.85 “5.1 0.50 112.8
1.12 251.8 1.05 12.0 0.70 83.“
1.36 252.9 1.19 12.0 0.88 66.8
1.52 251.3 1.60 12.2 1.12 32.“
1.86 251.9 1.91 12.8 1.32 31.9
2.67 251.8 2.81 12.“ 1.62 32.0
“.26 251.2 “.03 12.8 1.85 32.1
6.05 251.9 7.17 12.“ 2.82 31.9
“.15 32.“

7.“l 32.6

DMFb DMSOb Pyb
+ + +

L/Tl 5(ppm) L/Tl 5(ppm) L/Tl 5(ppm)
0.00 -129.9 0.00 -32“.1 0.00 -u.9
0.36 -85.1 0.“1 -230.8 0.36 -l3.l
0.8“ -10.9 0.72 -l72.2 0.83 -25.u
1.00 1“.“ 0.88 -l“7.5 0.91 —27.0
1.21 “0.1 1.00 -l30.9 1.01 -29.“
1.“8 60.1 1.09 -11“.7 1.10 -30.6
1.7“ 66.0 1.20 -108.7 1.16 -30.8
2.80 71.3 1.55 -77.0 1.56 -3l.7
5.u8 72.7 1.85 -53.2 1.89 -31.8
2.9“ -l9.0 2.90 -3l.6

5.29 13.1 “.67 -31.3

7.80 21.“ 6.58 -30.9

 

 

a0.01M'T1C10u.
b0.02110ch0“.

l6“

 

 

 

 

 

 

2 -

DMSO

I L *3

8

AN
(PPm) a PC
0 - —0NM

WMGr c

 

 

 

 

 

 

0.0 0:5 7:0 {.5 2:0 25 5.0 3.5
[0A18C51/[LV]

Figure 32. Lithium-7 Chemical Shifts yg. [DAl8C6]/[Li+]
Mole Ratio in Different Solvents.

165

 

l4

5

 

2' DMSO

 

 

l L *1

0.0 6.5 LO 1:5 20 2.5 3.0
[0A15C51/[No*]

Figure 33. Sodium-23 Chemical Shifts vs. [DAl8C6J/[Na+]
Mole Ratio in Different SoIvents.

166

 

 

DMF

 

 

 

fin
TMO

 

 

0

in"

aAN

gyso

109

 

 

0.0 to 2.0 3.0 40 5.0 6.0
[DAIBC61/[Cf]

Figure 3“ .
Mole Ratio in Different Solvents.

7.0 8.0

Cesium-133 Chemical Shifts E. [DAlBC6]/[Cs+]

 

167

 

 

 

 

 

 

 

 

 

 

l J

0.0 LO 2.0 3.0 4.0 ~ 50 ab 7.0 8.0
[DAIBCOJIUI’]

 

Figure 35. Thallium-205 Chemical Shifts 15. [ransom/[n+3
Mole Ratio in Different Solvents.

168

addition of the ligand, evidence of formation of a very
weak complex at the best. These are solvents with high
donicities and strong solvating abilities and, therefore
can compete with the ligand for the lithium ion.

On the other hand, in solvents of weak and medium
donor ability such as nitromethane, acetonitrile, propylene
carbonate, and acetone (with respective Gutmann donor
numbers of 2.7, 13.1, 15.1, and 17.0) the lithium-7 chemi-
cal shift is strongly affected by addition of the ligand,
indicating existence of an interaction between the lithium
ion and the ligand. As is seen in Figure 32, the lithium-7
resonance has a linear paramagnetic shift in nitromethane,
acetonitrile, and propylene carbonate (diamagnetic in
acetone) upon addition of the ligand which begins to level
off at mole ratio of about 1, indicating formation of a
complex with 1:1 stoichiometry between the lithium ion
and the ligand. A gradual diamagnetic shift was observed
in pyridine solutions which could be attributed to the
formation of a weak complex. The limiting chemical shifts
and the formation constants of the complexes were obtained
by computer fitting the chemical shift-mole ratio data to
an equation (discussed in Chapter 3) which relates the
observed chemical shift to the complex formation constant
using KINFIT program. The results are given in Table 39.

While the cavity size of the macrocycle is larger than

the size of the lithium ion, the resulting complexes seem

169

Table 39. Formation Constants and the Limiting Chemical
Shifts of l,lO-Diaza-l8-Crown-6-Li+ Complexes
in Various Solvents.

 

 

 

Solvent Log Kf 611m(ppm)
Nitromethane >5 -O.lli0.02
Acetonitrile 4.39:0.Nl 0.16:0.00
Propylenecarbonate 3.67:0.25 0.05:0.00
Acetone - 2.13:0.08 -0.36t0.02
Dimethylformamide ~0.0 -—--
Dimethylsulfoxide No.0 ----
Tetramethylguanidine ~0.o ---—

Pyridine O.h3:0.08 0.6t0.3

 

 

170

unexpectedly strong, particularly in the solvents with
low donicity. The formation constants are even greater
than those reported for lB-crown-6 - Li+ complex in the
same solvents (81). The stability of the complexes in
different solvents decreases in the order of nitromethane
> acetonitrile > propylene carbonate > acetone > pyridine
which is the order of increase in Gutmann donor number of
the solvents. Despite the close range of the limiting
chemical shifts in nitromethane, acetonitrile, propylene
carbonate, and acetone, chemical shift of the complexed
lithium ion is solvent dependent, indicating incomplete

insulation of the cation from the solvent by complexation.

5.2.2. l,lO-Diaza—18-Crown-6 Complexes with Na+

Sodium-23 NMR study of the sodium ion complexes with
the ligand in a number of solvents such as propylene car-
bonate and DMF was limited by the quadropolar broadening
of the 23Na resonance (Avl/2)500 Hz) which makes the
precise measurements of the chemical shift impossible.

The change in 23Na chemical shift upon addition of the
ligand to the sodium salt solution in acetonitrile was

found to be less than 0.5 Ppm which is in the range of

the error of the chemical shift measurement and, therefore,
cannot be used in such studies. Thus, the complexation of
sodium ion with the ligand was studied only in nitromethane,

acetone, DMSO, and pyridine solutions, where the linewidths

171

are narrow enough to measure the chemical shifts accurately.

The addition of the ligand to the sodium tetraphenyl-
borate solution in the above four solvents results in an
upfield or a downfield chemical shift which begins to level
off at mole ratio of about 1 indicating formation of a 1:1
complex between the cation and the ligand. The formation
constants and the limiting chemical shifts for the complexes
are shown in Table H0. The limiting chemical shifts of the
complex in nitromethane, acetone, and pyridine are close to
each other which possibly indicates that the cation is
mostly covered by the ligand. In DMSO solutions, however,
the complex seems to be more solvent dependent. It should
be noted that DMSO has a strong solvating ability and,
therefore, can compete with the ligand for the cation.

With the exception of pyridine, the complex formation
constant increases with decreasing Gutmann donor number
of the solvents. The results in pyridine are even un-
expected because it should be a good solvating solvent as
indicated by the magnitude of its sodium-23 chemical shift
(1H3) and the high Gutmann donor number of 33.1. This is
possibly because of existence of nitrogen atom as a soft
donor which cannot strongly solvate a hard ion such as
sodium. The sodium ion complexes with l,lO-diaza-l8-
crown-6 are much weaker than those with lB-crown-6 in the
same solvents (1&4) because of the introduction of two

nitrogen atoms into the lB-crown-6 ring. According to the

172

Table “0. Formation Constants and the Limiting Chemical
Shifts of l,lO-Diaza-18-Crown-6-Na Complexes
in Various Solvents.

 

 

 

Solvent Log Kf 611m(ppm)
Nitromethane 3.37:0.13 6.7Hi0.02
Acetone 1.96:0.18 8.83:0.09
Dimethylsulfoxide 1.19:0.08 2.86:0.21

Pyridine ' H.12i0.30 7.69:0.02

 

 

 

173

Pearsons Hard-Soft-Acid—Base (HSAB) theory (1&5), the‘
interaction of the sodium ion as a hard acid with the
nitrogen atom as a soft base should be weaker than that
with the oxygen atom as a hard base.

Another interesting point to note is that the line
width of the sodium-23 resonance at half height (Avl/z)
increases almost linearly with increasing the ligand to
sodium ion mole ratio, indicating creation of a more un—
symmetric environment of 23Na nucleus because of the
complexation, and begins to level off after mole ratio
of 1, showing a 1:1 stoichiometry for the complex (Table
36).

5.2.3. 1,10-Diaza-18-Crown—6 Complexes with Cs+

In all solvents used, with the exception of DMSO,
the addition of the ligand to the cesium ion solution
produces a large but gradual paramagnetic or diamagnetic
shift of the cesium-133 resonance (Figure 34). In none
of the solvents used does the chemical shift of the com-
plexed cesium ion reach a limiting value even at ligand
to metal ion mole ratio of about eight. The results seem
to indicate formation of a weak 1:1 complex between the
cesium ion and the ligand.

Using cesium-133 NMR, Mei 22.2l- (9H) have shown that
cesium ion forms both 1:1 and 2:1 (ligand to metal) com-

plexes with l8-crown-6. In pyridine, acetone, DMF, and

17“

PC solutions they observed a downfield shift of 13303
resonance followed by a relatively sharp break at the
mole ratio of l and an upfield shift which gradually ap-
proaches a limiting value. This behavior was explained
by formation of a 1:1 complex followed by addition of a
second molecule of the ligand to form a 2:1 sandwich com-
plex. It is reasonable to assume that the large cesium
ion will also form 1:1 and 2:1 complexes with l,lO-diaza-
18-crown-6 which has a cavity of about the same size as
l8-crown-6. However, no clear evidence of the formation
of a 2:1 complex between cesium ion and l,lO—diaza-18-
crown-6 was observed in this study.

The formation constant and the limiting chemical
shifts for 1:1 complexes of the cesium ion with the ligand
in different solvents are given in Table “1. A large
difference in the limiting chemical shifts of the complexes
(about 65 ppm) in different solvents clearly indicates
that the cesium ion is too large to fit into the ligand's
cavity, therefore the cation mostly remains under the
influence of the solvent. The complexes are much weaker
than the 1:1 cesium ion complex of 18-crown-6 reported
by Mei gt al, (9“). This is simply because of the exist-
ence of the two nitrogen atoms as a soft base in the ligand
which cannot strongly interact with the cesium ion, known
as a hard acid (1&5). Therefore, the resulting complex
should be weaker than that with 18-crown-6. This instability

175

Table H1. Formation Constants and the Limiting Chemical
Shifts of l,lO-Diaza-lB-Crown-S-Cs+ Complexes
in Various Solvents.

 

 

 

Solvent Log Kf 611m(ppm)
Nitromethane 2.79:0.02 -28.06iO.ll
Acetonitrile 2.26:0.02a -55.93¢o.08
Acetonitrile 2.30:0.01b -56.50so.02
Propylene Carbonate 1.95:0.02 -36.7620.36
Acetone 1.89:0.01a —h3.88i0.l9
Acetone l.92:0.01° -u5.56:0.l7
Dimethylformamide 0.61:0.07 -6l.O 18.6
Dimethylsulfoxide m0.0 _ ----
Trimethyleneoxide l.9fli0.02 1.91:0.72
Tetramethylguanidine 1.55:0.02 -79.8210.38
Pyridine 2.62:0.01 -52.D2:0.06

 

 

amoniesscm.

b0.03I_vI_tsSCN.

CO.OlM_CsBPhu.

176

indicates that if the 2:1 complexes are formed, the second
formation constant would be «K1 and, therefore, probably
would not show up on the mole ratio plots.

It is seen that the complex stability is not affected
either by a change in the concentration of the salt, in
acetonitrile solutions, or by changing the anion, in
acetone solutions. It is evident therefore, that at low
concentrations of the cesium salts, used in this study,
the formation of the complex is unaffected by ion pairing.
It is also interesting to note that, with the exception
of pyridine solution, the stability of l,lO-diaza-lB-crown-
6 ° Cs+ complex increases by decreasing the donicity of

the solvent.

5.2.“. l,lO-Diaza-18-Crown-6 Complexes with T1+

In nitromethane, acetonitrile, acetone, and pyridine
solutions the thallium-205 resonance shifts nearly linearly
with the ligand to metal ion mole ratio until a mole ratio
of 1 is reached. Further addition of the ligand does not
have any further effect on the 205T1 resonance (Figure 35).
This behavior indicates formation of a strong complex
(Kf > 105) between Tl+ ion and the ligand in the above
solvents. On the other hand, in DMSO and DMF, solvents
with high donicity, a gradual diamagnetic shift of the
thallium—205 resonance was observed upon addition of the

ligand which tends to level off after the mole ratio of

177

1, indicating formation of weaker complexes in these Sol-
vents.

The formation constants and the limiting chemical
shifts for the complex in various solvents are shown in
Table 42. The formation constants of the complexes in
nitromethane, acetonitrile, acetone, and pyridine are
greater than 105 and their precise values could not be
determined by our technique. In DMSO and DMF, solvents
with high donor ability, the stability of the complex is
still significant. The scattered values of the complex
limiting chemical shifts shows that the solvent molecules
can still interact with the complexed thallium ion, pos-
sibly from the top and the bottom of the ligand's plane.

5.3. DISCUSSION

The formation constants of Li+, Na+, 03+, and Tl+ ion
complexes of 1,10-diaza-lB-crown-6 and 18—crown-6 in various
solvents are compared in Table H3. The most probable ionic
diameters, calculated by using a combination of deduction
from rO values and experimental electron density maps by
Ladd (128), for the above cations are 1.72 A, 2.2a A,

3.68 A, and 3.08 A, respectively. The cavity size of the
ligand is about 2.6-3.2 A (23). The stabilities of the
l,lO-diaza—lB-crown-6 complexes decreases in the following

4.

order Tl+ > Li+ > Na > Cs+. Among the various factors

contributing to the complex stability, the two following

178

Table U2. Formation Constants and the Limiting Chemical
Shifts of l,lO-Diaza-l8-Crown—6-Tl+ Complexes
in Various Solvents.

 

 

 

Solvent Log Kf 611m(ppm)
Nitromethane >5 251.8
Acetonitrile >5 12.4
Acetone >5 32.4
Dimethylformamide - 3.22:0.08 7U.0i0.3
Dimethylsulfoxide 2.1310.03 “0.6i1.l
Pyridine >5 -31.5

 

 

 

 

179

. QG Ouflmhfluflflv . QQH ”Ur—OH 0H0-U .HQ OUGOHOUUM—D . NQH QUGOHUHOKO
nx ex so.os~o.o ma Ho.oa~o.~ on.oa~H.« mo.cans.o nx H.nn an
nun- nun- o.oa .nu- ~o.oann.H a.-- o.oe u..- u--- use
un.n u--- n-.. gnu: ~o.oa«o.a u-.. u-.. n.-- n--- oza

«o.ow«c.n so.owa«.a o.c;. Ho.ow~a.a 9.08. no.6waH.H c.c€ nc.cHnH.~ «.mm Gaza

 

 

 

ma.cwa.n «o.o«m~.~ lulu wo.owmn.n no.cwao.o null c.ce. we.ow-.n c.o~ man
mA «A No.0won.fl nA Ho.onom.a oa.oucm.u mo.cwna.u nA o.sa . o<
ma.ow«u.« «A Ha.owmo.~ null No.owna.a III: n~.ow~o.n lulu H.ma on
«A o~.owom.m «c.ow«n.~ mA ~o.owc~.u nail H«.owmn.« mA H.«H z<
III! «A «A mA No.c«ms.u mu.cwun.n an mA h.~ :z
v+uo o+nz A+wn o+HH +10 +uz +AA +H9 2: uoo>Hom
oassosouaa osssosoumanosoaonofi.a
«a mo;

 

 

.mu=o>aom maowuo> :« claaououwa

new olozou0ImHIonoualoH.H mo moonano com +HH new +mo .+wz .+Aq mo mucmuoaoo nouuoauom .n« manna

180

factors are especially important: first, the relative

sizes of the cation and the ligand, and second, the strength
of the interaction between the cation and the coordination
sites of the ligand. In order to interprete the data,

these factors should be considered.

In the case of the cesium ion both factors are against
the complexation. Cesium ion is a hard acid (1&5) which
cannot strongly interact with the soft nitrogen atoms of
the ligand. It is also too large to fit inside the ligand's
cavity. Therefore, the cesium ion complex in different
solvents is the weakest one in the series. Despite the
hard character of the sodium ion and consequently its
weak interaction with the soft nitrogen atoms of the macro—
cycle ring, this cation has a suitable size for the ligand.
Thus it is not surprising that l,lO-diaza-18-crown-6
forms stronger complex with sodium ion than that with cesium
ion. Although the size of the lithium ion is a little
smaller than that of the sodium ion for the ligand and it
is harder than the sodium ion, it forms an unexpectedly
more stable complex with the ligand than sodium ion.

Among the alkali metal ions, lithium ion is known to have
a great tendency toward covalent bond formation because

of its great polarization power (1A6). Therefore, the
existence of such covalent bond in the lithium ion complex
with the ligand could possibly result in such a great

stability of the complex. In the case of the thallium (I)

181

ion both factors are in the favor of the complexation and,
therefore, the thallium (I) ion complex is the strongest
among the series.

The substitution of two nitrogen atoms for two oxygen
atoms in 18-crown-6 greatly influences the complexation of
the ligands with the metal ions, as shown by the data in
Table “3. The sodium and the cesium ion complexes are
weakened appreciably by the nitrogen substitution. This
is Just as expected: as the negative charge on the co—
ordination sites drops by the substitution, the electro-
static interaction between the ligand and the cation
would diminish and the resulting complex would be weaker.
The effect is more pronounced in the case of the cesium
complex (iLEL, weakening of the complex by about two
orders of magnitude) than that of the sodium complex.

This is simply because the cesium ion is a harder acid
than the sodium ion which results in a weaker interaction
between the substituted nitrogen atoms, as a soft base,
and the cesium ion than that with the sodium ion.

The effect of the nitrogen substitution on the lithium
ion complex is exactly the opposite: the stability of
the complex is greatly increased. Because of its small
size and large polarizability, lithium is known to have
some chemical behavior that resembles the chemistry of
magnesium (146). Because of this unusual behavior, lithium

ion has a great ability to form covalent bond. It is

182

evident, therefore, that the increased stability of the
lithium ion complex upon the substitution is because of the
existence of both electrostatic and covalent bonding, in
contrast with the cesium and the sodium ion cases where

the bonding is only electrostatic in nature. In the case
of the thallium (I) ion complex, in most of the solvents
used the stability of the complexes are not exactly cal-
culated and, therefore, they cannot be compared. we can
expect, however, the thallium (I) ion complex with l,lO-
diaza-l8-crown-6 to be more stable than that with 18-crown-
6, such as that observed in DMSO solution. This is be—
cause the interaction of the soft thallium (I) ion with the

soft nitrogen atoms of the ring is quite strong.

CHAPTER 6

A STUDY OF DYNAMICS OF CESIUM ION COMPLEXES WITH

DIBENZO-30-CROWN-10, DIBENZO-ZH-CROWN-8, AND

DIBENZO-Zl-CROWN—7 IN ACETONE AND METHANOL

183

6.1. INTRODUCTION

While there are several scientific reports available
in the literature on the study of thermodynamics of the
complexation of alkali ions by crown ethers, the complexa-
tion kinetics of the cyclic polyethers have received less
attention. Kinetic studies of alkali metal complexation
with crown ethers are usually impeded by several factors:
complexes are not usually strong enough and therefore must
be studied at high metal ion concentrations; the reaction
rates are usually high and experimental difficulties are'
encountered especially if work at high concentrations is
required; and finally, the complexes are usually color—
less, so that spectrophotometric measurements of rates are
rarely possible.

In an early sodium-23 NMR study of complexation of
sodium ion by dibenzo-l8-crown-6 in DMF solution, Shchori
_t _1. (87) concluded that the strong ionic strength
dependence of the kinetics indicates that the dynamic
equilibrium involves interaction of solvated sodium ion
with uncoordinated ligand rather than a simple bimolecular
exchange equilibrium.

The same authors (88) studied the effect of solvents
and aromatic ring substituents on the kinetics of complexa-
tion of sodium ion by dibenzo-l8-crown-6. They used the

sodium-23 NMR technique and worked with DMF, methanol,

18“

185

and dimethoxyethane solvents. They found that the activa-
tion energy for decomplexation of sodium ion by dibenzo-
18-crown-6 is the same in all solvents (”12.6 kcal/mole)
but is considerably less (8.3 kcal/mole) for dicyclohexyl-
l8-crown—6 in methanol. Shporer and Luz (93) have used

39K and 87Rb NMR to determine the rate constants for potas-
sium and rubidium ion decomplexation by dibenzo-l8-crown-6
in methanol solution at low temperatures. They found

that the decomplexation of rubidium is much faster than

of potassium ion.

Using spectrophotometric measurements, Chock (32)
determined the rate constants for complexation of several
monovalent cations by dibenzo-30-crown-10 in methanol
solution. From the resulting data he suggested the exist-
ence of a fast crown ether conformational transition

preceding the complexation reaction:

CR
1 + K12 +
fast M + M K: MCR
CR2 21

in which CR1 and CR2 denote different conformations of the
uncomplexed ligand.

Eyring and coworkers (1A8,lu9) have determined the
rate constants for cation decomplexation by l5—crown-5

and 18-crown-6 in aqueous solutions by ultrasonic

186

absorption method. They concluded that the slow decomplexa-
tion of 18-crown—6-Cs+ complex is the reason for the high
selectivity of l8-crown-6 for potassium ion over the other
alkali ions.

The purpose of the work described in this chapter is
to study the kinetics of the complexation of the cesium
ion by DB3OClO, DB2AC8, and DB21C7 in acetone and methanol
solutions by cesium-133 NMR lineshape analysis at various

temperatures.

6.2. DETERMINATION AND INTERPRETATION OF THE LINESHAPES

The modified Block equations proposed by McConnell

(150) which describe the motion of the X and Y component"

of magnetization in the rotating frame are shown as fol—

lows:
dG
A _ -1 —1
dt + GAGA ‘ '1YH1M0A + T13 GB ‘ TA GA (1)
dG
B _ . -1 -1
dt + O‘BGB ‘ ’ lYHlMoB + TA GA " TB GB (2)
where

dA = l/T2A - 1(wA-w) and GB = l/T2B - 1(wB-w)

G = u + iv (3)

187

u and v are the transverse components (i;g;, absorption and
dispersion mode lineshapes) of magnetization along and
perpendicular to the rotating field, H1.

The solutions of Equations (1) and (2) appropriate for

slow passage are obtained by

dG dG

atfl=at§=0 (u)

The equations can be solved for GA and GB. Noting that

MOA = PAMO and MOB = PBM (5)

O

the total complex moment is

T + T + T T (a P + a P )
= -1YH1MO A B A B A A B B_ (6)
(1 + aAtA) (1 + aBTB) - l

 

G=GA+GB

Fast Exchange. In the limit of rapid exchange, TA

 

and TB are small and Equation (6) reduces to

 

G = -in M = - (7)

 

The imaginary part is

2 (8)

v=..'YH 0 2 2
v
1 + T2 (PAwA + P

1M w w)
B B -

188

representing a resonance line centered on a mean frequency
of

wmean = PAwA + Paws (9)

with a linewidth given by

T17=TPL+TEE (10)
2_ 2A 2B
If the exchange is not quite rapid enough to give
complete collapse, the central signal centered on ”mean
will appear to have a larger width than that given by Eq.
(10). A corrected form of Equation (10) can be obtained

by putting w = w in Equation (6) and expanding in

mean
powers of T. This gives an effective transverse relaxa-
tion time

P PB

._ A 2 2 2

_1_
T2 2A 2B

where PA and PB are relative populations at sites A and B,
respectively, the quantities “A and “B are the resonance
frequencies at the two sites at a given temperature in

the absence of exchange, and T2A and T2B are the respective
relaxation times at each site at a given temperature.

The lifetime of interaction is defined as

189

T = --——-- (12)

If at a given temperature the I value is greater than

/2 , where Aw = IwA - mBI, two separate resonance lines

nAw
are observed for the two respective sites. If the T value

 

is less than :Eé’ only one population average resonance
is observed.

1H NMR, where the trans-

In some cases such as 13C and
verse relaxation times are long enough, the following

assumption has been made (151-153):

fi§;.= Tl—.= 0 (13)
2A 2B

Since cesium-133 is a nuclei with narrow natural line-
width of about 1 Hz (quadrupole moment of q= 0.003 barns)
and consequently long enough transverse relaxation time,
it seems reasonable to assume that the above assumption is
also valid for 13303 NMR case.

In all cases studied, the ligand to cesium ion mole
ratio was 0.5 and also the formation constants of the

complexes were greater than 103 (Chapter 3). Therefore

PA = PB = 1/2 (in)

The relationship between the frequency in radians per

second, w, and that in cycles per second, v, is defined

190

 

 

as
w = 2flv (15)
In general the relaxation time is given by (15“):
‘3; = rate of removal of molecule from 1th state by exchange
Ti number of molecules in the ith state
(16)
If we consider the following complexation reaction
13f
M + L + ML (17)
kb
then
d /dt
?1_ = ._U—ML (18)
B ML
0 /dt
Ti -- MC (19)
A M
Since
d
ML _
‘ ‘d‘t" kb CML (20)
dM
" Ht" = kf CMCL (21)

191

Therefore

we have

with Equation (23)

For equal population that we used; Equation (1A)

.=_1_
2kb

(22)

(23)

(2“)

(25)

(26)

By substitution of Equations (12), (1A), (15), (2A), and

(26) in Equation (11) we will come out with the following

equation:

 

It is known that

(27)

192

g; = NW (28)

where W is the linewidth at half height. By substitution

of Equation (28) in Equation (27) we have

2
b - 2W

 

(28)

where VA and VB are the resonance frequencies of the two

sites in the absence of exchange, and W = A033: - Aviig.

The AvEBZ and Avis; are the linewidths at half height in
the presence and absence of exchange, respectively.

Thus above the coalescence temperature, the linewidth
of the sample solution (with the ligand/cesium ion mole
ratio of 0.5) was measured and corrected for natural
broadening by subtracting the linewidth of the resonance
signal of a solution of pure cesium salt in the same solvent
at the same temperature from it. The resonance frequencies
of the solutions containing free and complexed cesium ion,
VA and VB, respectively, in the same solvent at the same

temperature were also measured. The kb value at each

temperature was then calculated from Equation (29).

Slow Exchange. If the lifetimes TA and TB are suf-

 

ficiently large compared with the inverse of the separa-
tion (wA - mB)_l, the spectrum will consist of two distinct
signals in the vicinity of the wA and wB frequencies. For

example, if the frequency of w is close to “A, and thus far

193

away from ”8’ GB is effectively zero and the solutions of

Equations (1) and (2) become

 

P T
A
G s GA z inlMO ———5———— (31)
1 + “ATA
The imaginary part is
P T'
A
v = -in1MO ' 2 2A 2 (32)
A broadened signal centered at “A with width given by
parameter
-1 l -1
Y =
T2A T2A + TA (33)

There will be a corresponding signal centered on “B“

This shows that the exchange leads to an additional broad-
ening of the individual signals. If T2A is known, measure-
ments of the width of these broadened signals provide a
means of estimating TA'

Thus below the coalescence temperature the kb value

can be calculated from

kb = WW (3“)

where W is the corrected linewidth of the resonance signals,

i.e., w = Avggg - Avggg.

 

19“

The Arrhenius activation energy, E is given by

a,
81m k
b 2
( ) = E /RT (35)
BT V or P a
or
. Ea

The activation energy was obtained from the slopes of the
activation plots, (i.e., 2n kb vs. 1/T) by using a linear
least squares program. The thermodynamic parameters of

activation were calculated from the following equations:

 

# _
AHO — Ea — RT (37)
# kT AH:
ASO = R in kb - R in Tr + T (38)
¢_¢ a
AGO — AHO - TASO (39)
where AGé, AH:, and A8: are the standard free energy of

activation, the standard enthalpy of activation, and the
standard entropy of activation, respectively, and k and h
are Boltzmann and Planck constants, respectively. The

rate constants for the forward reactions (i.e., complexation)

195

were calculated from the following equation:

kf = K kb (110)

where K is the complex stability constant at 25°C.

6.3. RESULTS AND DISCUSSION

In order to study the kinetics of the complexation re-
actions between cesium ion and dibenzo-2l-crown-7, dibenzo-
2A-crown-8, and dibenzo-30—crown-10 in acetone and methanol
solutions, the cesium-133 NMR spectra of the cesium thio-
cyanate in the presence of the ligands (with a ligand/Cs+
mole ratio of 0.5) at various temperatures were obtained.
The results are shown in Figures 36-Al. The required

information in the absence of exchange (i.e. ”A: v8, and

ref
1/2

NMR spectra of the solutions of salt (site A) and of the

Au ) was obtained by lineshape analyses of the cesium-133

completely complexed cesium ion (site B), collected at
exactly the same conditions as the exchange case.

The rate constants for the release of the cesium ion
from the complexes, kb, at temperatures above coalescence
were calculated from Equation (29), below the coalescence
temperature from Equation (3A). The data are listed in

Table AA. The Arrhenius plots, log k vs. l/T are shown

b
in Figure A2. Activation energies (Ea), rate constants

(kb), and values of AH:, ASS, and AG: for the release of

 

 

-42
4000 2500 Hz

.Figure 36. Cesium-133 NMR Spectra of 0.02 M CsSCN, 0.01 M
DBZlC? Solution in Acetone at Various Tempera-
tures.

197

 

 

 

Cs+
Cs C+ 0C
-84
J 12

‘_-

-27

 

Figure 37. Cesium-133 NMR Spectra of 0.01: M CsSCN, 0.02 13

DB21C7 Solution in Methanol at Various Tem-
peratures.

Cs+ CsC

-54

MMWM MVW

-40
-20
4000 2500 Hz
Figure 38.

Cesium-133 NMR Spectra of 0.02 M CsSCN, 0.01'M

DB2hC8 Solution in Acetone at Various Tem-
peratures.

199

 

 

w-” o ._-

 

 

-55
, -i_
4000 2500 Hz

Figure 39. Cesium-133 NMR Spectra. of 0.02 M CsSCN, 0.01 M
DB2ACB Solution in Methanol at Various Tem-
peratures.

200

 

 

 

 

 

 

 

 

 

J L -28

'I500 A zsunJHz1

Figure 40. Cesium-133 NMR Spectra of 0. 01 M CsSCN, 0. 005 M
DB3OC10 Solution in Acetone at various Tem-
peratures.

201

WM A‘WA’AA ANNA fiiflMALWMNAAWwWVA/A MA;

.. ANAL, -au
nflflfifimmquA AfihfiflhnMVUAJ%fiflqu~%vaWANNA

LANA/AWAAVAWWA ”WW-MA AAA
WMWMWWAT

 

 

Figure Al. Cesium-133 NMR Spectra of 0. 01 M CsSCN, 0. 005 M
DB30010 Solution in Methanol at Various Tem-
peratures.

 

202

Table nu. Temperature Dependence of the Rate Constants
for the Release of 03+ Ion from DB21C7-Cs+,
DB2HC8 Cs+, and DB3OClO°Cs+ Complexes in Acetone

and Methanol.

 

 

Ligand Solvent Temp. (°C) log k

 

b

DB21C7 Acetone -82* 0.93
'76* 1.2“

-70* a

-6u* a

-57 2.18

-50 2.U8

-U2 2.77

Methanol -8H* 1.66
-72* 2.07

-65* 2.32

—61 2.35

-H8 2.85

-27 3.U3

DB2HCB Acetone -98* 0.75
-88* 1.36

—77* 1.8H

-67* a

-5u 2.84

-MO 3.30

-28 3.73

Methanol -92 1.93
—8U 2.24

-78 2.H7

-73 2.62

-55 3.32

 

203

Table uu. Continued.

 

 

 

Ligand Solvent Temp. (°C) log k
DBBOClO Acetone -80 1.84
-68 2.32
—56 2.66
-Ul 3.12
-28 3.28
Methanol -90* a
—80* a
-6M 2.95
-55 3.1M
-u9 3.34
-h5 3.H2
-H2 3.50
-30 3.77

 

 

*
Below the coalesence temperature.

aThe k

D

values are not calculated.

201A

 

 

 

 

l I l

4.0 4.5 5.0 5.5 6.0
103/T

 

Figure 42. Arrhenius Plots of log kb vs. l/T for the Release
of Cs+ Ion in Acetone and Methanol with Large
Crown Ethers. A-DB3OClO'Cs+ in Methanol,
B-DB2uC8-0s+ in Methanol, c—DB3001o-Cs+ in
Acetone, D-DB2lc7-Cs+ in Methanol, E—DB2uc8-0s+
in Acetone, F-DB2lC7-Cs+ in Acetone.

205

the cesium ion from the complexes, as well as the rate
constants for the complexation reactions (kf) are given in
Table H5.

It is immediately obvious that the nature of the solvent
plays an important role in the kinetics of the complexation
reactions. In all cases, the energy of activation for the
release of the cesium ion from the complex in methanol is
less than that in acetone. It should be noted that methanol
is a solvent with higher donicity than acetone as expressed
by their respective Gutmann donor number of 25.7 and 17.0
(106). The same effect was observed by Cahen gt_al. (78)
for the release of lithium ion from cryptate C2ll.Li+ in
pyridine, water, DMSO, DMF, and formamide solutions. By
contrast Shchori et al, (88) found that the activation
energy for decomplexation of sodium ion by dibenzo-18-
crown-6 is the same in methanol, DMF, and dimethoxyethane.
However, two of these solvents, i;g;, DMF and methanol,
have about the same donicity, while the donicity of di—
methoxyethane is not known.

The activation energy for the release of the cesium
ion from the complex is also dependent on the size of the
ligand. In the same solvent, there is an inverse relation-
ship between the activation energy for decomplexation of
the cesium ion by the ligand and the size of the ligand
(Table #5). These results seem to indicate that the major

barrier to removal of the cesium ion by the complexes is the

206

.COHumH>op cpmocmpmp
.ponss: Locop acmepzcm

 

 

 

 

 

AH.oVN.HH Ao.avm.mmn A:.ovo.: Am.HVm.e Am.ovm.m Az.ovm.m oaoommo
AH.ovs.oH Am.mvm.oan Am.ovs.m Ao.Hv:.z As.avs.m Am.ovm.m wozmmo
AH.ove.HH Ae.Hve.eHu Am.ova.o Ao.0vm.m Am.ovm.m Am.ovs.m eofimmo s.mm Hocooooz
AH.ovm.HH Am.va.omn Am.ovm.m Am.ovo.m A:.ova.m Am.ova.m oaoommm
AH.ovo.OH AN.NVH.OHI Ao.ovo.e Aw.Hve.o As.va.oa Am.ovm.m moemmo
AH.ovm.OH Am.mvm.sn Ae.ovs.w Am.HVH.s AH.HVm.m oAm.ovm.m aofimmo o.eH ooooooa
mHoE\HmOx mHoE\HmOx oHoE\Hmox Hloom HIS Hlomm mHoE\Hmox pcmwfiq mzo uco>aom
O O O .m
Axommmv noa nmo xmo mIOonx agoaxox m
.Hocmnpoz ocm oCOpoo< CH moonQEoo czoho mmpmq
mEom Eopw COH no mo ommmfiom pom mpmuoemhmm OHEmsmUoELmne 6cm mmumm owcmnoxm .m: magma

+

207

energy required to affect a conformational rearrangement.
The lower activation energy for the release of the cesium
ion from the larger crowns would then be a result of the

greater flexibility of the ligand.

In both solvents, the rate constants for decomplexation
of the cesium ion by DB2AC8, kb, is greater than those for
the release of the cation by DB3OClO and DB21C7. The rate
constants for the decomplexation reactions seem to be
determined by the rigidity of the resulted complexes.
Cesium ion with diameter of 3.68 A (128) has a convenient
size for the cavity of DB21C7 with the size of 3.u—u.3 A
(23) to form a firm "cation in the hole" complex. We have
already shown (Chapter 3) that DB3OC10 is large enough to
form a rigid "wrap around" complex with the cesium ion,
while due to differences in the sizes of DB2UC8 cavity and
of the cesium ion either a complete "wrap around" or a
"cation in the hole" complex. The rate constants for both

complexation and decomplexation of the cesium ion by

DB3OC10, i424, kf = 7.5 x lo8 M“Asec'l and kb = 3.5 x 10‘A
sec-l, are in a satisfactory agreement with the values
reported by Chock (32), iLEL, kf = 8 x 108 M.1 sec-1 and
kb = u.7 x 10LA M‘1 sec‘l.

As is seen in Table 45, while the enthalpy and the
entropy of activation for the release of the cesium ion
from the ligands are very sensitive to the solvent and

to the size of the ligand, the free energies of activation

208

in all cases are about the same. Such a compensation of
the enthalpy with the entropy is not an uncommon occurr-
ence (155). It is interesting to note that the entropy
of activation for decomplexation of the cesium ion by
large crowns (Table #5) are in opposite direction with
the corresponding overall entropies (Table 3”, Chapter M).
This is in contrast with the entropy values reported for
the release of the cesium ion from cryptate C222 in DMF
solution by Mei §£.§l-'(95): where the activation and the
overall entropies have the same sign and are close in
value. From the results the authors concluded that the
activated complex should resemble the final state, i424)

the solvated cesium ion and the free cryptand. The op-

g

o and AS° values for decomplexation of

posite sign for AS
the cesium ion by large crowns, however, suggests that

the conformational entrOpy changes may play an important
role (156).

Since the activation energy for decomplexation of the
cesium ion by the ligands increases with increasing the
donicity of the solvent, the transition state must involve
a substantial ionic solvation. A sample entropy profile
for decomplexation of the cesium ion by DB3OClO in acetone
is shown in Figure 43. It clearly shows that the transi-
tion state must be more ordered than both the initial and

the final states, i.e., the solvated complex and the

solvated cesium and the free ligand. Thus the possible

 

209

 

 

 
  
 
 
  

 

 

 

 

 

 

 

08300055“
A i
AS°=26.2e.u.
>~
0. t
g vDB30C10'Cf A5 =-46.5e.u.
.5 A
A$*=—20.3e.u.
V , V
(DBaoCm-Cf)
——>

Reaction Coordinate

Figure “3. Entropy Profile for the Release of Cs+ Ion
From DB3OClO-Cs+ Complex in Acetone.

210

mechanism for the release of the cesium ion from the

complex can be shown by

# #
AS AS
DB3OClO°CS+ <—————- (DB3OC10’CS+)# <r—————- DB3OC10 + CS+
-20.3e.u. -U6.5e.u.

where the overall entropy is AS° = 26.2 e.u.

APPENDICES

APPENDIX I

DETERMINATION OF COMPLEX FORMATION CONSTANTS BY THE
NMR TECHNIQUE, DESCRIPTION OF COMPUTER PROGRAM
KINFIT AND SUBROUTINE EQUATION

Let's consider the following equilibrium for a one

to one complex,

M + L 1 ML (1)
with the concentration formation constant K

K = CML/CM'CL (2)

C, stands for concentration.

The observed chemical shift of M (5 ) is a mass

obs
average of the characteristic chemical shift of M at each
site (M in the bulk solution, and M complexed), assuming
that a fast exchange occurs between these two sites with

respect to the NMR time scale.

6 = X

obs MGM + X

Where: 5M is the characteristic chemical shift for M in
the bulk solution, "ML is the characteristic chemical

shift for M complexed (ML), XM is the fraction of M (CM/
(CM + cML)), XML is the fraction of ML (CML/(CM + CML)),

then

obs XMSM + (l’XM)5ML

0)
ll

obs = XM(5M‘5ML) + 6ML (u)

211

 

212

cg = CM + CML (the analytical concentration of M)
CM
Gobs = EB'(6M'6ML) + GML
M
t -
CL - CML + CL (the analytical concentration of L)

using (5) and (7), CL = CL - (CIVI — CM)

t
CM ‘ CM
t

K = t
(CM)(CL = CM + C

 

M)

C is solved in (8)

M

t t _ t
CMCCL - CM + CM)K - CM - CM

2

 

 

t t t _
KCM + (KCL - KCM + 1)Cm - CM - 0
-(Kct — Kct + 1) + VQKct — Kct + 1)2 + cht
c L M ‘ L M M
M

2K

the positive root is

(5)

(6)

(7)

(8‘

213

 

t_ t \/2t22t2 2tt t t
(KCM KCL - 1) + K CL +K CM -2K CLCM+2KCL+2KCM+1

 

2K
(9)

Substitution of CM from (9) in Equation (6)

_ t t
Gobs -[[(KCM ‘ "CL ’ l) +

 

/ 2 2 ’
2 t 2 t 2 t t t t ]
K CL + K CM — 2K CLCM + 2KCL + 2KCM + 1]

[(5M - SML)/2CtK] + GML (10)

We assume a constant value for 5M and that 5ML and K
are unknown. In order to fit the calculated shift (the
right hand side of Equation (10) to the observed chemical
shift, the program may vary the values of 6ML and K.
Hence, the number of unknowns, NOUNK, equals two as does
the number of variables, NOVAR.

The first card contains the number of experimental
points (columns 1-5 (F.15)), the maximum number of itera-
tions allowed (columns 10-15 (F.15)), the number of
constants (columns 36-AO (F.15)) and the maximum value of
(A parameter/parameter) for convergence to be assumed
(0.0001 works well) in columns Al-SO (F10.6). The second

data card contains any title the user desires. The third

 

214

data card contains the value of CONST(l) (0%) columns 1—10
(F10.6) in M, CONST(2) (6M) columns ll-20 (F10.6) other
constants can be listed on columns 21-30, 31-“0, etc.

The fourth data card contains the initial estimates of the
unknowns U(l) = 6ML and U(2) = K, in columns l—lO and ll-20
(F10.6), respectively. The fifth through N data cards
contain XX(1) = C§ in columns l-lO (F10.6) variances on
XX(l) in columns ll-20, XX(2) = the chemical shift at XX(1)
in columns 3l-A0 (F1036) followed by the same parameters
for the next data point. Each card may contain two data
points. If no further data are to be analyzed the next
card after the last data point(s) should be a blank card
followed by a 6789 card. If more data sets are to be
analyzed, the next card after the last data point(s) is

the first data card of the next set.

215

 

‘66.
NNN
000 0
EE E
CE o. P
X! 0 0
‘09 o 5
VT?! 3 9!.
cl OD ()x
1L0 UOR
OZL Filo!
U o O 9‘ 0
Opt \IVI’
X 09 0 00
9U0 0’5
KPH. 30‘
N O 9 (0|.
”0.! 0.1.!
GPA F‘X
N 90 0T0
'pJ ,5.
no C ox,
Apr-l o 90 ‘l
v 0‘ 3’5 II
010 (0‘ l.
NIN 029 U
.x 9 0", 0
TD?! FLT—.3 I.
P 05 QAXI. C
OPN .IVRI. Pry
NYO “G Ox ‘0
OTC (I‘lx \I,
RX 9 Xonp 2’
CR! x 95 9V! 1|,
V .5 .‘ll"! U 0
Ixc DON“! ‘l
XXN gill. I O
O 9 O 01‘Ml\ o! IE.-
0.9.! 305T, I. '0
AYC 9115?. 1‘ [U
LIE 1* 9 0‘]. .l (O
91v ("0 C S TIC
T 9 o XOOCI N 50
WSY 705 9.! 0 NR
.1an xu31|ol 9 C 00
9E O x 9"...00 \v. CA
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PD 9 OOZC( O 3 I) | OS 2
AAJ 1'2 ONY9 ’0' \r 08
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909 J926 '0 O‘X‘vu) IT
El“: H "l 9"? o ’vll‘IlS‘l PR o
959 T100 909 G ‘ISQCNC ’0 6
‘EL ED. 0205.! (N’XOI’AS
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N1... 0" 1K QTalvls ul VAC. ((N‘”- I.
OOVD //4CQ.D. 9 . ((OODOX. 0‘- .
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"07 DNXVPSLQ E ul)7.l\rl\.7.o\ E
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N0 QTYR0N301 9 o o o U((I).IBX o
IK'SSFPO‘CO?F°I E E EH E. O (UU)?SO. E E EH E E E E E
.l SAN//I.26.5‘U6b?7 U U UT U”. (ItaliN‘g U U U... U U U U U
UHF. OONNS 0’... N: : = : NHNHNHENquZOO 0 (HOA: NHNNNNENNNNNNNNNNH‘
BTLCOUNOQIOIEEKPPIDIFINpol‘l. oOOUCCl‘n—«AC IQIMCl-PIPID 110.19
Du 0A 0M Mr?! PYTPPHAUTUT UT? UTUU-C o 0‘ .(lIUT UT UTIUTHIUTUTUTH.
H.“ IVYM MN‘IAU .NflA UVTN'NTN‘IN“. 22: .. :(Slhrl NIN‘TN‘NTHT N... HT
HOIBYOGIPYHOOTYOOEOEOEOFEC:::=:AHC:EEQEOECFEOEOEOEOEOE
SCWN...” CCD o 0L GCIJM NPCRCFCIC CA PCDFABCSQQCD CQCIRCQCRCD CQCD
II II?
.ul 7 8 2 R: 3 A :1 O Q. C .n! 2
1.. at I. 1. I.

(MU

APPENDIX II

DETERMINATION OF COMPLEX FORMATION CONSTANT WITH

ION PAIR FORMATION BY THE NMR METHOD

 

Definition of symbols,

[M] = cation

[A] = anion

[L] = ligand

[MA] = ion pair

[ML] = metal complex

CM = analytical concentration of metal ion
CL = analytical concentration of ligand
K1p == ion pair equilibrium constant

Kf = formation constant of complex

Ion pair equilibrium,

[M] + [A] 2 [MA]

- [MA]
Kip ‘ mm (11)

Complexation equilibrium,

"1?
[M] + [L] [ML]

CL = [ML] + [L] (13)

216

217
CM = [ML] + [M] + [MA]
CM = [MA] + [A]
From (11),
[MA] = Kip [MJEA]
From (12),
[ML] = KfEMJEL]

Substitute in (13)

 

cL = KEMJEL] + [L] = [L](Kf[M] + 1)
C
° _ L
° '[L] ‘ Kf[M] + 1

Substitute in (1H),
CM = KfEMJEL] + [M] + KiptMJEA]
Substitute in (15),

CM = [A] + KiPEMJEA] = [A](l + KiPEMJ)

.'[A] = 1 + K

(1“)

(15)

(16)

(17)

(18)

(19)

(2o)

218

Substitute (18) and (20) in (19),

Kf[M]CL

= K1p[M]CM
M Kf[M] + l

1 + KipEM] (21)

c + [M] +

 

 

Multiply (21) across by (l + Kf[M])(l + Kip[M]),
CM(1 + KfEM])(l + KiprJ) = KfCLEM](1 + KipEMJ)

+ [M](l + Kf[M])(l + Kip[M]) + KipCM[M](l + Kf[M]) (22)

2
CM + KipCMEM] + KfCM[M] + KfKipCMEM]

2 2
= KchEM] + foichEM] + [M] + KipEM]

2 3 2
+ Kf[M] + KfKipEM] + KipCMEM] + KfKipCMEM] (23)

Collecting terms,

3 2
KfKip[M] + (KfKich + Kip + Kf)[M]

+ (KfC + 1 - KfcM)[M] — CM = 0 (2“)

L

219

(KfK CL + K + Kf)

 

 

 

 

 

 

[M33 + 19 1p [M]2
foip
+ L [M] — M = o (25)
KfKip KfKip
Solution to cubic equation,
y3 + py2 + qy + r = 0
+
p (KfKipCL Kip + Kr)
KfKip
KfKip
-CM
r’ =
KfKip

(3q - p2)/3

93
ll

(2p3 - 9pq + 27r)/27

0‘
ll

220

 

 

A =\\]% +\’ + a
B =\q/% VT 27
2 3
Case I b a
1r-+ 27 > 0 + 1 real root
b2 a3
II 1f-+ 57 — 0 + 3 real roots
b2 a3
III 7T' + 27 < 0 + 3 real roots

Case I, x = A + B

Case II and III, use trigonometric form

2“ — % Cos (% + 120°)
24 — % Cos (% + 2A0°)

221

Now, solve for [M] in (25).

Then substitute in following equations,

 

 

 

K C [M]
[ML] = f L
Kf[M] + 1
K C [M]
[MA] = ip M
Kip[M] + 1
sobs = XMGM + XMLGML + XMA‘SMA
_ LE1. [ML] [ML]
6obs ' CM. 5M + CM 5ML + -Cfi—6MA

Use final form of 6 in EQN subroutine.

obs
Coding symbols in EQN,

a = AA p = PP
b = BB q = QQ

A = AAA r = RR

B = BBB ¢ = FE

y = R Cos ¢ = CFE
CONST(l) = 1p

CONST(2) = CM

CONST(3) = 61p

CONST(A) = 6M

U(l) = SML XX(1) = 0L
U(2) = Kf xx<2) = sobs

(26)

(27)

(28)

(29)

nnnnnnnnn

2222

sUEonuTxNE En-

C'Nucu «aunt.ITADEo.;TAo£.IwToLADoxthp.NopT.N VAR.~ouux.x.n.tTwa. Eons
IBTA. TEST I.Av.9E§lD.IA°~ EPSolTYDoXXoRXTvpcnxl oFODan-Fu.°.ZIoTO E EDNA
210:025321g onTOL0“9IIJOYODY0VEC70NCSToCONSToNDAToJDAToHODToL0° T Eons

VT

corona/FDEDT

f‘zuCN/DOIM

D} "(NSION I

UW'vx
daox
on:

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*40K‘CCWJOD<'M

"-
~Odlxn~3333490

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WM
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q u—«o
'hn UKDU

n
[’2 Cflflx

444430"

T
0
S
:
t F {nu PAID
¢ TA

r
upng UFTAL
T F COMPLEX

IGAND
«IFT
Ttla
) I

-—,-
vad

W
)0

—mv—-—wmmw30\
0<MOOIISDFWZOV4
.3$IDZ—..'-— (V100
FOKOWWWWOD ocx
0c» fl-u-«qolr vom-

oonnnqnmn
'4
NWWFZ

a r0fl71NUE
O'Iiflpc
2 C(I“VH
52%!“F'”'~5°"’ Go TO 35
I ‘
as CDuTlIHE
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5:? 136 THE coerr c
99:13‘ 110(|./U( T
Qo=(I./(U(?TOCO S
50=-CON§T(?TI(U P
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i- o
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9 Z

‘1'"!

A2=¢ °n0-99-
9H8. .O(90001
CHrgg DOINT F0:

12 753(OHH/?fOSOQT(T

78=conH/2.-STGT(TESTTI
adB=(AflS(TRTT°'CHRR

cagcwrwG THE SIGN or 1H: cuaEROOT TERMS

(7"179179 l8

‘1PAMFA“

I“ 7F (IH’190I9925

q and:

5 qfisazlonflq
R:Np-(DD/3. )

jxflngéggnNo“EIplc FORM OF VHE SOLUTION IS USED WERE
QAu§2o°Acns¢-|.b
ANGLE=RAO/1.
grist-3n/2. .l/(SORT(ARS(C-AAOOJDIZT.ITT

xF(¢C7E. LT -|.T.08.(EFE. 6;. If? TI GO TO ‘0

r =pA0-3coscc"
1A §.:;° :g§3§1:Ae§(-AA/3.vnTocos¢rE/3.oANGLE-rLAGI
' 0
CHECK FOR VOLLOITT or THE CALCULATED «(VAL cowc
!F(Q '110130 3
A: unit (JTAOE. 2
2: EOSHAITO COSINE or rc: OUTSIDE or o1 Ano -..,

cosegrx AND rxwo ANOTHER ROOT
[1 E%AG:E>‘GO‘

LAoiis.ts.la
IS weliéchAo 5.2::

2i rggvATC 0 ALL THREE ROOTS FLUVKEDOT

s
SUDST TUTE cau-cn ATED “ETA couc. ~ on n DUAT n
30 0|9§5‘:Egggg?gfgl?:0‘285§;%;::é:;;gtléi%l{febtgb'wiitgou5TTItoo°cnus
O
QESYD’DELTA-XX(2T ' 0 2’
Rf ITJDN
3 cc~rluut
6 ZSerguE
s CONTINUE
éééyflttfloflE.-|! Go To 20
20 com [NU
8E1U PM E
o ONTINUE
Wu
‘3 ggfiflflufi
ii cngINUE
VIP"
12 ro' UxNu:

EN

REFERENCES

10.

ll.

12.

13.

1“.

15.

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