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This is to certify that the

thesis entitled

Spectroscopic and Electrochemical Studies of
Ion Solvation and Complexation in Various Solvents

presented by

Richard M. Farmer

has been accepted towards fulfillment
of the requirements for

Ph.D. degree in Chemistry

 

Major professor

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OVERDUE FINES:

25¢ per day per item

RETURNING LIBRARY MATERIALS:

Place in book return to remove
charge from circulation records

 

 

 

 

SPECTROSCOPIC AND ELECTROCHEMICAL STUDIES OF
I ION SOLVATION AND COMPLEXATION

IN VARIOUS SOLVENTS
By

Richard M. Farmer

A DISSERTATION

Submitted to
Michigan State University
in partial fulfillment of the requirements

for the degree of
DOCTOR OF PHILOSOPHY
Department of Chemistry

1981

 

Jr/‘/</

C—;}//

 

ABSTRACT

SPECTROSCOPIC AND ELECTROCHEMICAL STUDIES OF
ION SOLVATION AND COMPLEXATION

IN VARIOUS SOLVENTS

By

Richard M. Farmer

Calcium—A3 nuclear magnetic resonance (NMR) behavior
of armmber of calcium salts was studied in acetone, di-
nwthylsulfoxide, dimethylformamide (DMF), ethylene glycol,
formamide, methanol, propylene carbonate, tetramethylguani—
dine, and water as a function of salt concentration. There
mm evidence for ion—pairing for the chlorides, bromides,
perdflcrates, and nitrates in all solvents except DMF.
Thecomplexation of Ca2+ ion with ethylenediaminetetra—
acetic acid (EDTA), and crown ethers l8-crown—6 (18C6)
andJB—crown—S (lSCS) was also studied. The complexes
wifiiEDTA and l8C6 exhibited a slow exchange of the u3Ca
imibetween the free and complexed state, resulting in two
resonances. There was no evidence for Ca2+ ion complexe-
tion with 1505.

The tetraphenylarsonium (and tetraphenylphosphonium)

 

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..-. s

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l

Tri

Richard M. Farmer

tetraphenylborate single ion assumption was studied using
75As, 31?, 11B, 130, and proton NMR. Due to broad reson—

75As NMR was not sufficiently precise

ance lines (1000 Hz),
to test the validity of the assumption. Phosphorous—31
amiboron—ll NMR data show that the central ion is not
amirely shielded from the solvent, although ion—ion
hmeractions are negligible. Carbon-l3 data indicate that
Hm charge is not localized on the central atom. Both
carbon—l3 and proton data indicate that tetraphenylarsonium,
tetraphenylphosphonium, and tetraphenylborate ion are not
solvated equivalently. Ion-ion interactions, as seen by
13C and proton NMR, are negligible. Care is advised when
ung this assumption to calculate free energies of trans—
ier for single ions.

Complexation of phosphonoacetic acid (PAA) with calcium
ion was studied potentiometrically in aqueous solution.

Th

(D

thermodynamic formation constants for the deprotonated
2+

Q.

m1 monoprotonated PAA complexes with Ca were found to be:

log Kf = 4.68 i 0.03 and log Kf = 2.68 i 0.08, respectively.
Th

('D

thermodynamic quantities

AG° = -6.38 0.07 kcal/mole
AHO = 0.6 0.2 kcal/mole

AS° = 21.4 0.7 cal/OK-mole

were calculated from the temperature dependence of the

formation constant for the deprotonated PAA complex with

 

 

Richard M. Farmer

Ca2+ ion. The formation constant of manganese(II) ion with
PAA was determined using electron spin resonance, while

the complexation constants of copper(II), lead(II), and
thallium(I) with PAA were determined using cyclic voltam—
metry. The silver(l) ion complexation constant with PAA
was measured potentiometrically. The complexation of
calcium(II), manganese(ll), and thallium(I) ion was studied
with the ligands 3—phosphonopropionic acid and phosphono—
formic acid, resulting in the determination of formation

constants for these ligands.

 

 

To Jane

"Love is patient, love is kind, and is not jealous;
love does not brag and is not arrogant,

does not act unbecomingly; it does not seek its own, is
rmt provoked, does not take into account a wrong suffered,

does not rejoice in unrighteousness, but rejoices with
the truth;

bears all things, believes all things, hopes all things,

endures all things.

1 Cor. 13:4—7

 

 

 

 

....

.l..

 

ACKNOWLEDGMENTS

The author wishes to thank Professor Alexander I.

Popov for his guidance, counsel, and patience in the field
of chemical research and scientific writing.

The contributions of Professor Stanley R. Crouch as
second reader are gratefully appreciated.

The author wishes to acknowledge the advice and
assistance of the following individuals: Dr. Thomas V.
Atkinson, in the area of computer programming and usage,
Professor Michael J. Weaver in electrochemistry, and
Professor Gerald T. Babcock in the field of electron spin
resonance.

The friendship and help of Dr. A. I. Popov's group is
greatly appreciated, as it made life at Michigan State
mflyersity infinitely easier and more pleasant.

The help of Frank Bennis, Wayne Burkhardt, and Tom
Clark in the maintenance of the NMR instruments is acknow—
ledged.

Financial aid from the Department of Chemistry, Michigan
State University, the National Science Foundation, and
lhuon Carbide was appreciated.

Finally, to my wife, Jane, whose love, encouragement,

- . . - L ' ompletion
and patience were mayor lactors in the successful 0 _

Of this research project, thank you.

TABLE OF CONTENTS

Chapter

LIST OF TABLES.
LIST OF FIGURES
PART I — Calcium—43 NMR Studies
CHAPTER I - HISTORICAL REVIEW
1.1. Introduction.
1.2. Calcium—43 NMR.
1.3. Ligands
1.4. Conclusions
CHAPTER II — MATERIALS AND METHODS.
2.1. Reagents.
2.1.1. Salts

.1.2. Purification Procedures
.1.3. Drying Procedures .

l\)l\)

2.2. Ligands
2-3- Solvents.
2.4. Molecular Sieves.

2.5- Nuclear Magnetic Resonance.

2.5.1. Varian CF T- 20 NMR
Spectrometer. . .
2.5.2. Bruker WH— 180 Super—

conducting NMR
Spectrometer. - -
2.5.3. Bruker WM— 250 Super—
conducting NMR
Spectrometer. . .
2.5.4. Operating Procedures.

iv

 

Page

ix

xvii

10
12
13
13
14
15
16
17
17

17

18

18
18

 

Chapter

2.

2.

6.
7.

Miscellaneous Instrumentation .

Data Handling

CHAPTER III — Results and Discussion.

3.
3.

PART II —

7.
8.

WU)

Introduction.

General Information on Calcium—43
NMR . . . . . . . . . . .

Aqueous Studies — Chemical
Shifts. . . . . . .

Linewidths.

Studies in Nonaqueous Solutions

.5.1. Chemical Shifts
.5.2. Linewidths.

Gutmann Donor Number Cor-
relation. . . . . . .

Complexation.

Conclusions

AN NMR INVESTIGATION OF THE
REFERENCE ELECTROLYTE ASSUMPTION
FOR THE DETERMINATION OF SINGL
ION MEDIUM EFFECTS.

CHAPTER I - INTRODUCTION AND HISTORICAL

Introduction.
Historical.

Reference Electrolyte
Assumption. .

A Nuclear Magnetic Resonance
Study . . . .

1.4.1. Arsenic—75 NMR.
1.4.2. Phosphorous-31 NMR.

Page

20
2O
21

22

23

26
33
33

33
43

J: 42‘
KO '\1

\J1
U1

56
57
58
61

62

69

69
70

 

Chapter

1.5.

.4
.4

Conclusions

3. Boron-11 NMR.
4 Carbon—13 and Proton.

CHAPTER II — RESULTS AND DISCUSSION

2.

2.

2.

1.

2.

LA)

7.

Introduction.

Phosphorous-31 NMR Results.

Arsenic—75 NMR Results.

Boron—11 NMR Results.

Carbon—13 NMR Results

Proton NMR Results.

Summary and Conclusions

PART III - PHOSPHONOACETIC

CHAPTER I — HISTORICAL

1.

1.

ACID STUDIES

REVIEW

Introduction.

1.1.1. Viral Replication

’_l

.1.2. Inhibition of the

tion Path by PAA.

Toxicity of PAA

Analogs of Phosphono-

acetic Acid

Viral Replica—

Chemical Properties of PAA

and its Analogs

Experimental Techniques

Conclusions

.5.2. Manganese
Resonance

Electron

-5.1. Calcium Ion Selective
Electrodes .

Spin

Page

71
72

73
75
76
77
81
87
88
112
140
142
143
144
146

 

Chapter

CHAPTER
2.

2.

2.

1.

2.

4.

CHAPTER

3.

1.

.5.

II — MATERIALS AND METHODS.
Reagents.
Experimental Methods.

2.2.1. Electron Spin Resonance

Spectroscopy.
2.2.2 Cyclic Voltammetry.
2.2.3 Potentiometry
2.2.4 Coulometer.

Sample Preparation.

2.3.1. Manganese Electron
Spin Resonance.

l\)

.3.2. Cyclic Voltammetry.

.3.3. Potentiometry

m

Data Handling
III — RESULTS AND DISCUSSION.
Complexation Studies of PAA

3.1.1. Manganese Electron Spin
Resonance . . . .

Cyclic Voltammetry.

CA) E\)

1.
3.1. Potentiometric Studies.

Thermodynamics of Calcium
Complexation with PAA

Summary and Discussion.
Phosphono Acetic Acid Analogs

3.4.1. Manganese Electron
Spin Resonance.
3.4.2. Cyclic Voltammetry.

4.3. Potentiometry

Thermodynamic Studies of 3-PAA

Complexes

vii

210
210
213

 

.r

,-
r
,..a'
.-
a

 

Chapter

3.6.
3.7.
APPENDICES.

APPENDIX A —

APPENDIX B —

APPENDIX C —

APPENDIX D —

BIBLIOGRAPHY

Summary and Discussion.

Conclusion.

THE GENERAL USE OF THE NON—
LINEAR CURVE FITTING ROUTINE
KINFIT4. . . . . . . . . .

DATA REDUCTION OF A POTEN-
TIOMETRIC TITRATION FOR THE
DETERMINATION OF COMPLEX FORMA—
TION CONSTANTS USING THE COMPUTER
PROGRAM MINIQUAD

THE USE OF THE PROGRAM FARM2.TSK
FOR THE DETERMINATION OF MONO—
PROTONATED PAA-CALCIUM FORMATION
CONSTANTS FROM POTENTIOMETRIC DATA

SOME CROWN AND CRYPTAND COMPLEX
FORMATION CONSTANTS WITH MANGANESE
(II) ION DETERMINED BY ELECTRON
SPIN RESONANCE . . . .

viii

219
219

222

222

225

238
241

 

Table

II

III

IV

VI

LIST OF TABLES

Properties of Selected Nuclei in

NMR Spectroscopy.

Some Solvent Properties and Magnetic
Susceptibility Corrections.
Linewidths of the Calcium-43
Resonance Signal in Some Solvents

A Comparison of the Transfer Ac—
tivity Coefficients from Water to
Methanol of Silver and Sodium

Based on Different Single Ion
Assumptions

The 31F Chemical Shifts in ppm for
Tetraphenylphosphonium Salts as a
Function of Counterion, Solvent, and
Concentration

The 75As Chemical Shift in ppm

(i7 ppm) for Tetraphenylarsonium

Ion as a Function of Concentration,
Counterion, and Solvent

Some Representative 75As Linewidths

for Tetraphenylarsonium Chloride

Page

24

35

44

63

78

82

 

 

 

Table

VIII

IX

XI

XII

XIII

in Various Solvents (:80 Hz).
ThelJTSChemical Shifts of Sodium
Tetraphenylborate as a Function

of Concentration and Solvent in
Various Solvents.

The Infinite DilutionllB Chemical
Shifts, Corrected for Bulk Mag-

netic Susceptibilities, as a

Function of Solvent

Carbon—13 Chemical Shift Changes with
Concentration for Tetraphenylarsonium
(As(Ph)u+) and Tetraphenylphos-

phonium (P(Ph)u+) Salts in H O.

2
Carbon-13 Chemical Shift Changes
with Concentration for Tetraphenyl-
arsonium (As(Ph)u+) and Tetra-
phenylphosphonium (P(Ph)u+) Salts
in MeOH

Carbon—13 Chemical Shift Changes
with Concentration for Tetraphenyl-
arsonium and Tetraphenylphosphonium
Salts in DMSO

Carbon—13 Chemical Shift Changes
with Concentration for Tetraphenyl—

arsonium and Tetraphenylphosphonium

 

Page

83

84

86

89

90

 

 

V‘.-

 

 

 

 

Table

XIV

XV

XVII

XVIII

XIX

Page

Salts in DMF. . . . . . . . . . . . . . . . 93
Carbon—l3 Chemical Shift Changes

with Concentration for Tetraphenyl—

arsonium and Tetraphenylphosphon—

ium Salts in Nitromethane ... . . . . . . . 95
Carbon-13 Chemical Shift Changes

with Concentration for Tetraphenyl—

arsonium and Tetraphenylphosphonium

Salts in Propylene Carbonate. . . . . . . . 96
Carbon—13 Chemical Shift Changes

with Concentration for Tetraphenyl—

arsonium and Tetraphenylphosphonium

Salts in Acetonitrile . . . . . . . . . . . O7
Carbon—13 Chemical Shift Changes

for Sodium Tetraphenylborate in

Various Solvents as a Function

of Concentration. . . . . . . . . . . . . . 98
Bulk Magnetic Susceptibility Cor-

rections as a Function of Tetra-
phenylphosphonium Chloride Concen—

tration in DMSO . . . . . . . . . . . . . . 103

13

The Infinite Di1ution C Chemical
Shifts for Tetraphenylborate,
Tetraphenylphosphonium, and

Tetraphenylarsonium Ions as a

xi

 

 

 

a
.v

.,--

 

 

 

 

 

 

Table

XX

XXI

XXII

XXIII

XXIV

XXV

Page

Function of Solvent . . . . . . . . . . . 105
The Difference in 13C Infinite

Dilution Chemical of Tetraphenyl—

arsonium (and Tetraphenylphos-

phonium) vs Tetraphenylborate

Ion as a Function of Solvent. . . . . . . 107

The 13

C Chemical Shifts of Tetra—
phenylgermanium and Tetraphenyl—

arsonium Chloride in Dichloro-

methane and Deuterated Chloro-

form. . . . . . . . . . . . . . . . . . . 113
Proton Chemical Shift Data for
Tetraphenylarsonium Chloride,
Tetraphenylphosphonium Chloride

and Tetraphenylborate in Water. . . . . . 114
Proton Chemical Shifts of Tetra—
phenylarsonium Chloride as a

Function of the Total Chloride

Ion Concentration . . . . . . . . . . . . 115
Proton Chemical Shift Data for
Tetraphenylarsonium Chloride and

Some Tetraphenylphosphonium Salts

in Methanol . . . . . . . . . . . . . . , 116
Proton Chemical Shift Data for

Tetraphenylarsonium Chloride,

 

 

 

Table

XXVI

XXVII

XXVIII

XXIX

XXX

Some Tetraphenylphosphonium Salts
and Sodium Tetraphenylborate in
DMSO.

Proton Chemical Shift Data for
Tetraphenylarsonium Chloride,

Some Tetraphenylphosphonium

Salts, and Sodium Tetraphenyl-
borate in DMF

Proton Chemical Shift Data for
Tetraphenylarsonium Chloride,
Some Tetraphenylphosphonium Salts,
and Sodium Tetraphenylborate in
Acetonitrile.

Proton Chemical Shift Data for
Tetraphenylarsonium Chloride,
Some Tetraphenylphosphonium Salts,
and Sodium Tetraphenylborate in
Nitromethane.

Proton Chemical Shift Data for
Tetraphenylarsonium Chloride,

Some Tetraphenylphosphonium Salts,
and Sodium Tetraphenylborate in
Propylene Carbonate

The Infinite Dilution Proton

Chemical Shifts for Tetraphenyl—

xiii

 

Page

117

119

121

122

123

 

 

 

 

 

Table

XXXI

XXXII

XXXIII

XXXIV

XXXV

XXXVI

arsonium, Tetraphenylphosphonium,

and Tetraphenylborate Ion as a

Function of the Solvent . . . . . . . . 134
The Difference in Proton Infinite

Dilution Chemical Shifts.of
Tetraphenylarsonium XE Tetra-

phenylborate Ion as a Function

of Solvent. . . . . . . . . . . . . . . 137
The Proton Chemical Shifts of
Tetraphenylgermanium and Tetra-
phenylarsonium Chloride in Di-

chloromethane and Deuterated

Chloroform. . . . . . . . . . . . . . . 139
Formation Constants of a De—

protonated Manganese-Citrate

Complex . . . . . . . . . . . . . . . . 177
Formation Constants of Manganese—

(II) Ion—PAA Complexes at Dif-

ferent Ionic Strengths and 25°C . . . . 185
The Change in the Emf of a Pb—

PAA Complex as a Function of the

Negative log [PAA] at 25°C,

I = 0.05. . . . . . . . . . . . . . . . 192
The Change in the Emf of a Cu—PAA

Complex as a Function Of the

 

 

 

 

Table

XXXVII

XXXVIII

XXXIX

XL

XLII

 

 

Negative log [PAA] at 25°C,

I = 0.05.

The Change in the Emf of a T1—
PAA Complex as a Function of the
Negative log [PAA] at 2590,

I = 0.05.

Formation Constants of Thallium(I)—
PAA Complexes at Different Ionic
Strengths and 25°C.

Formation Constants for Cal—
cium(II) Ion—PAA Complexes as

a Function of the Ionic Strength
at 25°C

Thermodynamic Quantities for
Metal Ion Complexation with

PAA.

The Solvation Entropy, Radius,
and Formation Constants of Some
Metal Ions with PAA at 25°C and
an Ionic Strength Of 0.05
Formation Constants of Thallium-
(I)—PFA Complexes at Different

Ionic Strengths and 25°C.

XV

 

Page

192

194

201

205

 

 

 

Table

XLIII

XLIV

XLV

Formation Constants of Calcium—
(II)-PFA Complexes as a Function
of the Ionic Strength at 25°C
Formation Constants of
(II)—3-PPA Complexes as a Function
of the Ionic Strength at 25°C
Formation Constants for Some

Metal Ions with PAA,

3—PPA at 25°C and I

Page

217

220

 

LIST OF FIGURES

Figure Page
1 The structure of some crown ii
ethers and a oryptand. . . . . . . . . . . 5
2 Concentration and anion de—

43

pendence of the Ca chemical

shift in aqueous solution measured

with a 60 MHZ NMR spectrometer

(triangles - CaC12, squares —

CaBr2, circles — Ca(010u)2, and

crosses - Ca(N03),). . . . . . . . . . . . 27
3 Concentration and anion de—

A3

pendence of the Ca chemical

shift in aqueous solution measured

with a 180 MHz NMR spectrometer. . . . . . 28
4 Concentration and anion de—
43

pendence of the Ca chemical

shift in TMG, methanol, and

ethylene glycol. 36
5 Concentration and anion de—
pendence of the ”30a chemical
shift in DMSO, formamide, DMF,
37

acetone, and propylene carbonate

xvii

 

Figure

10

11

12

A plot of the Gutmann donor
43

number Kg the

tion chemical 5

chemical shift IE the mole ratio

Ca infinite dilu—

hift.

of EDTA to calcium

Calcium-43 resonance of a solu—

tion 0.4M in CaCl2 and 0.2M in

EDTA at pH mlO.

+
complexed Ca2 ,

2+
Ca

The structure of tetraphenyl—

arsonium, tetraphenylphosphonium,

Peak A —

peak B — free

A plot of the observed ”3

Ca

and tetraphenylborate ion.

The 13

tetraphenylarsonium chloride,
middle — tetraphenylphosphonium

chloride, and lower — sodium

0 NMR Spectra ofzupper —

tetraphenylborate in DMSO.

The 13

phenylarsonium tetraphenylborate

in DMSO and DMF

Some resonance structures of

tetraphenylarsonium ion.

xviii

C NMR spectra Of tetra—

Page

48

5O

64

99

101

 

 

 

Figure

13

14

15

16

17

l8

19

Page

A plot of the difference in chemical shift
between the cation and anion Kg 7

the dielectric constant. . . . . . . . . . 111
A plot of the proton chemical

shift lg the total salt concentra-

tion in water (squares — para protons,
triangles — meta protons, circles —

ortho protons, open symbols — tetra—
phenylarsonium chloride, closed

symbols — sodium tetraphenylborate). . . . 124
A plot of the proton chemical shift

Ki total salt concentration in

DMSO . . . . . . . . . . . . . . . . . . . 125
A plot of the proton chemical Shift

_§ total salt concentration in DMF . . . . 126
A plot of the proton chemical

shift Ki total salt concentration

in acetonitrile. . . . . . . . . . . . . . 127
A plot of the proton chemical shift

Kg total salt concentration in

nitromethane . . . . . . . . . . . . . . . 128
A plot of the proton chemical shift

Kg total salt concentration in

propylene carbonate. . . . . . . . . . . . 129

 

 

 

Figure

20

21

22

23

24

25

26

Page

A plot of the proton chemical

shift XE total chloride ion con—

centration for tetraphenylarsonium

chloride in water. . . . . . . . . . . . . 131
The proton NMR spectra of.sodium
tetraphenylborate, tetraphenyl-

arsonium chloride, and tetra—
phenylphosphonium chloride in

DMSO . . . . . . . . . . . . . . . . . . . 132
The proton NMR spectra of tetra-
phenylarsonium chloride in dif—

ferent solvents. . . . . . . . . . . . . . 135
The structures of phosphonoacetic

acid (PAA), phosphonoformic acid

(PFA), 2-phosphonopropionic acid,

and 3-phosphonopropionic acid. . . . . . . 1A5
DNA synthesis and the proposed PAA

inhibition mechanism . . . . . . . . . . . 1A8
A schematic representation of the

Orion 90-20 calcium ion selective

electrode. . . . . . . . . . . . . . . . . 154
A schematic representation of the

pyrex glass cells used for A -

cyclic voltametric studies, and

B — potentiometric studies . . . . . . . . 164

XX

 

 

 

Figure

27

28

29

3O

31

A calibration plot of the ob-
served ESR intensity as a function
of MnCl2 concentration

A plot of the log formation
constant of Mn—PAA as a function
of the ionic strength (upper curve
— deprotonated PAA, lower curve —
monoprotonated PAA).

A plot of the change in Emf XE

the negative log [PAA] to determine
the formation constant of a Pb-
PAA complex (upper curve — de—
protonated PAA, lower curve —mono—
protonated PAA).

A plot of the change in Emf XE

the negative log [PAA] to determine
the formation constant Of a Cu-
PAA complex (upper curve — de—
protonated PAA, lower curve -
monoprotonated PAA).

A plot of the change in Emf XE

the negative log [PAA] to determine
the formation constant of a mono—

protonated PAA—T1 complex.

Page

184

189

190

191

 

 

 

33

34

35

36

37

A plot of log formation constant
Of deprotonated PAA—T1 complex as
a function of the ionic strength

A sample calibration plot and
titration curve of silver ion
with deprotonated PAA using a
silver ion Specific electrode.

A sample calibration plot and titra-
tion curve of calcium ion with de—
protonated PAA using a calcium

ion selective electrode.

A plot of log formation constant
of Ca—PAA as a function of the
ionic strength (upper curve —
deprotonated PAA, lower curve —
monoprotonated PAA).

A plot of log Kt is 1000 x Tem-
perature-l to determine the entropy
and enthalpy of a deprotonated
Ca+2 complex with deprotonated

PAA (upper curve) and 3-PPA (lower
curve)

A plot of the entropy of solvation
(e.u.) Ki the log formation

constant of some metal ion complexes

xxii

195

200

204

 

Figure

38

39

40

with deprotonated PAA.

A plot of the log formation constant.
of a deprotonated PFA—Tl complex

as a function Of the ionic strength.
A plot of the log formation

constant of Ca-PFA as a function

of the ionic strength (upper curve —
deprotonated PAA, lower curve —
monoprotonated PAA).

A plot of the log formation

constant of Ca-3-PPA as a function
of the ionic strength (upper curve

- deprotonated PAA, lower curve

- monoprotonated PAA).

xxiii

Page

208

211

 

PART I

CALCIUM—43 NMR STUDIES

 

CHAPTER I

HISTORICAL REVIEW

 

 

._u

 

 

 

 

 

 

 

 

 

1.1. Introduction

 

Multinuclear NMR has been shown to be a sensitive
technique for probing the immediate chemical environment
Of metal ions in solution. This sensitivity, coupled
with the non-destructive nature of the method, renders
multinuclear NMR a desirable technique for the study of
a wide variety of chemical and biological systems. AS
the SCOpe of this investigation encompasses a range Of
metal—ligand interactions, e;g;, calcium complexes with
phosphonocarboxylic acids, it was Of interest to determine
the suitability of natural abundance calcium-43 NMR as a
probe of the complexation phenomena. Prior to such a study,
however, fundamental work on the NMR of calcium salt solu-
tions was necessary to characterize such solution inter—
actions as ion-pairing and solvation. These interactions
are important to the extent that they influence complexa-
tion (ngL, ion—pairing will compete for the free calcium
ions in solution, which will have a definite effect on
complexation.)

It was also Of interest to us to investigate a pos-
sible use of M3Ca NMR for studies of calcium complexes in
+

- 2
aqueous and nonaqueous solutions. We chose to study Ca

complexes with the well known EDTA ligand as well as with

 

 

 

 

 

 

 

 

 

 

some crown ethers. The latter ligands have been a
subject of intense interest since their introduction
by Pederson (12,13). The ligands 1505 and 1806 (Figure
l) were chosen because the calcium diameter most closely
matches the ring cavity size (13) for these ligands; 18C6
has a slightly larger cavity size than would be indicated
for a "best fit” with the calcium ion, but studies (16-18)
have shown that ions will also complex with the next larger
ring size.

These ligands should provide a range of environments
which would effect the observed calcium-43 resonance, there—
by indicating the suitability of calcium NMR for the in-

vestigation of complexation reactions.

1.2. Calcium-43 NMR

 

The calcium-43 resonance signal was first observed by
Jeffries (1) in 1953 at a frequency of 2.85 Mc/sec with a
10 RG magnet Operating in a continuous wave mode. A 0.7 M

43

CaBr2 aqueous solution (68% enriched in Ca) was used to
obtain the value of 7/2 for the Spin and a magnetic moment
Of —l.3l52 nm for this nuclei. Further studies were not
conducted until 1969 when Bryant (2) studied calcium—ATP
complexes in water. A continuous wave Varian V-4210A
spectrometer was used to study 0.9 M aqueous calcium solu-

43

tions (31.68% enriched in Ca) with variable ATP concen—

trations. The exchange rate was observed to be fast on the

 

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mum”

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a.
r.

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NMR time scale, resulting in a single resonance with a

typical spin—lattice relaxation time T = 0.85 s. The

2
author concluded that calcium—43 NMR would be.a powerful
tool for investigating calcium interactions with large
biological molecules.

The calcium-43 NMR of solid compounds has been in—
vestigated by Bleich and Redfield (3). It was found that
a double resonance technique with solid state decoupling

43

Ca resonance for a CaF2
4
sample with natural abundance of 3Ca. The full width at

allowed the measurement of a

half intensity was 130 Hz at 770K. A similar study on
CaF2 was conducted by Jacquinot, 3: £1. (4).

The most fundamental study of calcium ion interactions,
though, was published by Lutz, at £1. (5) in 1973 and con-
tinued in 1975 (6). Using a 18.07 kG Bruker BKR32 spec-
trometer at a resonance frequency of 5.178 MHz, Lutz in—
vestigated the aqueous behavior of some calcium chloride,
bromide, nitrate, and perchlorate solutions. The effect
of concentration and counter—ion on the observed calcium
resonance was studied for solutions ranging from 0.5 to
7.2 M. Also of fundamental importance was the determina-

43

tion of the quadrapole moment for Ca by Grundeuik, t al.
(7). An atomic beam magnetic resonance technique with
laser detection was used to determine a quadrapole moment

Of -0.065 b. The bare calcium-43 nucleus was generated

by means of a plasma discharge source.

 

The bulk of calcium—43 NMR investigations since 1970,
however, followed Bryant's suggestion that biological
systems were particularly amenable to this technique.
Robertson, Hiskey, and Koehler (8) studied the binding
of y-carboxylglutamic acid containing peptides to calcium.
The calcium salt used in this study was 79.98% enriched in

43

Ca. The dissociation constant of calcium ion with the

.ligand was found to be 0.6 mM with a total chemical Shift

range upon complexation Of 2 ppm. Calcium binding to par-
valbumins was investigated by Parello, gt £1. (9). The
effects of pH, temperature, and calcium/protein ratio on
the full width at half-height of the resonance band were
studied with a 100 MHz spectrometer using a 61.63% en—
riched calcium salt. The exchange rate between free and
bound calcium was found in all cases to be fast on the NMR
time scale. The interaction of the calcium ion with calf
DNA was studied by Reimarsson (10) using enriched samples.
The average lifetime for calcium ion bound to DNA was
estimated as 9ms at 25°C and 3 ms at 50°C. Marsh, _§ _1-
(11) studied calcium ion—bovine prothrombin fragment 1
interactions with enriched calciume43 NMR, in conjunction
with circular dichroism and fluorescence techniques and
found that calcium ion induces a helical structure in the
prothrombin as a function of pH. The linewidth of the

calcium resonance varied between 2 and 10 Hz.

 

 

 

 

1.3. Ligands

The complexation properties of 1505 in aqueous solu-
tions have been studied calorimetrically by Izatt, _t al.
(19) with the alkali, alkaline earth, and some post transi-
tion elements. It was found that, within experimental
error, there was no evidence for complexation with the
calcium ion. This does not preclude completely the pos-
sibility of such a complex, as calorimetric techniques
require a measurable AH of complexation in order to deter-
mine a formation constant. A low AH would indicate no
complex according to calorimetric studies, although the
AS term and subsequent AG might be favorable to complexa-
tion. A study by Lamb (20) in methanol yielded a Ca2+ —
1505 formation constant Of log Kf = 2,18. Crystallographic
studies of Ca2+ — 1505 complexes have not been done due to
the difficulty Of obtaining crystals. However, structures
of Ca2+ - Benzo 1505 (21) and Ca2+ - Dibenzo 1505 (22)
complexes indicate the calcium ion is not positioned in the
plane Of the ring. This suggests that the complexation
with the benzo derivatives of 1505 is weak. These results,
together with Izatt's work, indicate that in aqueous solu—
tions, at best only a weak complex can exist.

The toxic effects of 1505 have been studied by Hen—
drixson 33 g1. (23). The acute and chronic effects Of oral
administration of 1505 was studied in mice. The LD (the

50

lethal dose required to kill 50% of the mice) was found

 

 

 

 

to be 1.02 g/kg body weight of the mouse. There was no
cumulative effect from daily ingestion of small amounts

of 1505 over a period of two weeks. The ligand is a strong
topical irritant which is readily absorbed through the
skin.)

Complexation studies of 1806 with calcium ion have been
conducted by Izatt in water (24), 70% aqueous methanol (25),
and pure methanol (20). There is no evidence for complexa—
tion in water, while log Kf (formation constant) values
are 2.51 in 70% methanol, and 3.86 in methanol. These
data, along with those for calcium - 1505 complexes, indi-
cate that water solvates Ca2+ ion much stronger than
methanol. In water, therefore, the solvent effectively
competes with the ligand for the calcium ion, leading to
little or no complex formation. In methanol, the stability
of the complex is greater than the solvation enthalpy (26). A
crystallographic study of the Ca2+-18C6-SCN complex (27)
revealed a disordered structure in which calcium lies in
the mean plane of the ring. The thiocyanate anions are
associated with the calcium above and below the plane of
the ring. The inclusion of calcium within the ring is
indicative Of a stronger interaction than that of Ca2+-
(di)benzo 1505 where calcium is outside the cavity. These
data indicate that Ca2+ ion forms a stronger complex with
the next larger crown than would be predicted from the

best fit Of an ion into the crown cavity.

 

 

10

Toxicological studies of 1806 on mice (23) have led to
the determination of an LD50 of 0.71 g/kg body weight.
Again, no cumulative effects were Observed with daily
ingestion over a period of two weeks. This study concluded
that the toxic effects of crown ethers increased with in—
creasing ring size.

The above brief survey is by no means an attempt at a
complete compilation of crown ether chemistry. The reader
is referred to several books (28—30) and review articles
(14,31-33) for a more complete study of macrocyclic chem—
istry.

Complex compounds with EDTA are too well known to
warrant an historical survey. It seems sufficient to
note that the formation constant of the calcium ion complex
with EDTA was found to be 5.0 x 108 at 25°C and an ionic

strength of 0.1 (15).

1.4. Conclusions

Thus far, the use of calcium—43 NMR for the investi—
gation of complexation and ion—ion interactions has been
restricted to aqueous solutions, usually using an iso—
topically enriched calcium salt. The behavior of calcium ‘
salts in nonaqueous solvents has not been investigated
by calcium—43 NMR. The purpose of this portion of
the thesis was to study the technique of natural abundance

calcium-43 NMR, especially as it applies to the solution

 

A,"

 

Tie——

11

chemistry of calcium salts. This work should serve as a
basis for future studies involving complexation reactions

in aqueous and nonaqueous solvents.

 

 

 

CHAPTER 2

MATERIALS AND METHODS

 

12

 

 

 

 

 

2.1. Reagents

2.1.1. Salts

The chemical manufacturer of each compound has been
coded as follows: Aldrich (A), Alfa (A1), J. T. Baker (B),
Eastman Organics (E), Fischer Scientific 00. (F), K & K
(K), Mallinckrodt (M), Matheson, Colleman, and Bell (MCB),
Merck & CO. (Mc), Ozark Mahoning (O), Pfaltz & Baver (PB),
Parish (P), Richmond Organics (R), and G. F. Smith CO. (S).

The following chemicals were of reagent grade and were
used without further purification (except for drying)
calcium bromide dihydrate (MCB), calcium chloride (MCB),
calcium nitrate tetrahydrate (F), calcium perchlorate tetra—
hydrate (S), cupric chloride dihydrate (MCB), ferrous
chloride hydrate (MCB), lead nitrate (F), lithium per—
chlorate (K), manganous chloride hexahydrate (B), nickel
perchlorate hexahydrate (B), potassium bromide (M), potas-
sium hexafluorophosphate (PB), potassium hydroxide (MCB),
potassium nitrate (MCB), potassium phosphate dibasic tri—
hydrate (M), silver nitrate (MCB), sodium chloride (M),
sodium hexafluoroarsenate (0), sodium para-toluene sul—
fonate (E), sodium perchlorate (F), thallium nitrate (A1),
3-trimethylsi1yl propionate sodium salt (known as TSP,

the water soluble analog of TMS for use as an external

13

 

 

 

14

standard in D20), and zinc chloride (M). Tetramethyl-
ammonium hydroxide (E) was obtained as a 10% solution in
water. A tris buffer was made with gold label tris-(hydroxy—

methyl)amino methane (Al) and perchloric acid (B).

2.1.2. Purification Procedures

Tetraethylammonium perchlorate (TEA?) was recrystallized
twice by dissolving in a minimum amount of boiling water.
The solution was filtered hot, then recrystallized upon
gradual cooling to 0°C. The crystals were filtered and
washed with cold water. Tetraphenylarsonium (S) and tetra-
phenylphosphonium chloride (Al) were precipitated from an
iSOpropyl alcohol-acetone solution upon the addition of
benzene (92). Tetraphenyl phosphonium bromide (A1) and
iodide (A1) were recrystallized from water following the
procedure used for tetraethylammonium perchlorate. Tetra—
phenylarsonium tetraphenylborate was synthesized by mixing
aqueous solutions of tetraphenylarsonium chloride and gold
label sodium tetraphenylborate (A) in equal proportions.
Tetraphenylarsonium iodide was synthesized by mixing
aqueous solutions of tetraphenylarsonium chloride and re—

agent grade sodium iodide (M) in equal proportions.

2.1.3. Drying Procedures

 

The following chemicals were dried in a vacuum ove
o n

(except where noted), with the drying time and temperature

 

 

 

 

 

 

15

in parenthesis: calcium chloride (*), calcium nitrate
(96 h, 140°C) (32), calcium perchlorate (96 h, 140°C)
(32), lithium perchlorate (*), manganese chloride tetra-
hydrate (48 h, 210°C, regular oven), tetraphenylarsonium
salts (24 h, 100°C), tetraphenylphosphonium salts (24 h,

100°C), and tetraethylammonium perchlorate (24 h, 50°C).

2.2. Ligands

Ethylenediamine tetraacetate disodium salt (B) and
citric acid (M) were of reagent grade, requiring only dry-
ing (*) prior to usage. Crown ether 1806 (A or P) was
recrystallized from acetonitrile by forming the 1806
acetonitrile complex (33) at an acetone—ice bath tempera—
ture. The solution was filtered and dried (48 h, ambient).
The melting point of the purified product was approxi-
mately one degree below that of the literature value (33).
The 1505 (A) was fractionally distilled under reduced pres—
sure and vacuum dried. The 1204 (A) was used without
further purification.

Phosphonoacetic acid (PAA, R) was dissolved in a
minimal amount of glacial acetic acid near its boiling
point. The solution was gradually cooled to room tem-

perature. A minimal amount Of ethyl ether was then added

*(24 h, 110°C, regular oven)

 

 

 

 

16

to aid in recrystallization. The crystals were filtered
and rinsed with ethyl ether, followed by vacuum drying
(48 h, 60°C). The melting point of 141-142 agrees well
with the literature value Of 142—143°C (34). Phosphono—
propionic acid (3—PPA, R) was recrystallized similarly.
The melting point was l76—l78°0 compared to a literature
value of l78-l80°0 (35). The phosphonoformic acid (PFA)
was donated by Dr. Boezi's group and was used without

further purification.

2.3. Solvents

The following solvents were used during the course of
experimentation: acetone (F), acetonitrile (M), dimethyl—
formamide (DMF, M), dimethylsulfoxide (DMSO, F), ethylene
glycol (MCB), formamide (F), methanol (M), nitromethane
(A), propylene carbonate (PC, A), pyridine (M), tetra—
methylguanidine (TMG, E), and water.

The following purification and drying procedures were
Common to all solvents: the solvent was refluxed and
fractionally distilled for 48 hours over a chemical drying
agent, with only the middle fraction retained. The solvent
was stored over molecular sieves (Davison Chemical 00.,

38 pore Size) under a nitrogen atmosphere. Unless stated,
the chemical drying agent used was calcium hydride.
Acetone, ethylene glycol, propylene carbonate, and

. ' ent.
TMG were distilled using molecular Sieves as a drying ag

 

17

Methanol was distilled over Mg turnings, while pyridine
was distilled using KOH and molecular sieves as drying
agents. Conductance water was obtained through the
courtesy Of Dr. Weaver's laboratory and was used without

further purification.

2.4. Molecular Sieves

 

Molecular sieves of pore size 3A were dried at 110°C
for 24 h prior to activation. Activation was accomplished
by drying at 500°C for 24 h under nitrogen (dried by passage

through sulfuric acid).

2.5. Nuclear Magnetic Resonance

The nuclear magnetic resonance spectra were collected
on three instruments, all operating in the Fourier trans—

form mode.

2.5.1. Varian CFT—20 NMR Spectrometer

13
The OFT-20 spectrometer was used to measure some 0

l .
Spectra. The resonance frequency of 3C is 20 MHz at a
field strength of 18.7 kG. Spinning tubes of 10 mm OD
With a 3 mm OD capillary containing 0.3 M TSP in D20 as

a reference held in the center with teflon spacers.

 

 

s

 

18

2.5.2. Bruker WH—18O Superconducting NMR Spectrometer

Phosphorous-31, arsenic-75, and calcium—43 spectra
were obtained using this instrument. Their resonance
frequencies are 72.88 MHz, 30.82 MHz, and 12.12 MHz,
respectively at a field of 42.28 kG. Calcium—43 and ar—
senic—75 were run in 20 mm OD non-spinning tubes with a
5 mm 0D reference tube held in the center. The loss of
sample homogeneity was negligible when comparing spectra
obtained using spinning Kg non—spinning tubes. The phos—
phorous-31 studies were conducted using a 10 mm OD spinning
tube with a 5 mm reference held in the center with teflon

spacers .

2.5.3. Bruker WM—250 Superconducting NMR Spectrometer

Proton and carbon-13 spectra were obtained primarily
with the WM-25O spectrometer at resonance frequencies of
250 MHz and 62.894 MHz in a 58.72 kG field. The sample

preparation is the same as that for the CFT-20 Spectrometer.

2-5.4. Operating Procedures

All chemical shifts were measured with respect to an

external standard and are as follows: Proton and carbon—

' A.
13 - a 0.3 A solution of TSP in D20. arseNlC-75 - a 0-1 A

KASF6 SOlution in D20, phosphorous—31 — A 85% phosphoric

 

 

19

acid in D20 solution, boron-11 — A 0.1 M (CH30)3B solution
in D20, and calcium—43 — an infinitely dilute aqueous solu-
tion (a 5.071 M 0a012 was used as a standard)- The solutes
were weighed under normal atmospheric conditions, then
transferred to a dry box for dilution with a given non—
aqueous solvent.

Chemical shifts downfield from the reference (paramag—
netic are designated as positive. All chemical shifts have
been corrected for the bulk magnetic susceptibility dif—
ferences between pure water and the given non-aqueous sol-
vent. The correction due to the solute is assumed to be
negligible except in the case of proton NMR (36). The
corrections have been calculated based on the equations

of Live and Chan (37) for normal magnets (Equation (1))

and superconducting magnets (Equation (2)),

N 4N 1

COP? : 6obs _ I? (Xref _ Xsample) ( )

— 3 2i _ <2)
6corr — sobs + 3 (Kref Xsample)

where scorr is the actual chemical shift (independent of

N - S -
the t.Vpe of magnet used), Sobs and dobs are the observed

chemical shifts for normal and superconducting magnets,

and Xrea and X are the magnetic susceptibility of

I sample

the reference and sample-

 

20

2.6. Miscellaneous Instrumentation

Melting points were determined with a Fischer—John
melting point apparatus. The water content of the non—
aqueous solvents was measured with an Aqua Test 11 Karl-
Fischer apparatus. The purity of selected solvents was

determined using a Varian Aerograph Model 920 gas chromato—

graph.

2.7. Data Handling

Extensive use was made of the CDC—7501 computer for
equilibrium calculations. The programs KINFIT4 (Appendix
A, 38) and MINIQUAD (Appendix B, 39) were used for
equilibrium and thermodynamic parameter calculations.
Some complex equilibrium data required further treatment

using a PDP—ll minicomputer and the program in Appendix

C.

 

CHAPTER 3

RESULTS AND DISCUSSION

21

 

;

 

 

3.1. Introduction

The purpose of this investigation was to determine the
suitability of calcium—43 NMR Spectroscopy to the study of
Ca2+ ion solution chemistry in aqueous and nonaqueous sol-
vents. It has been shown that alkali metal NMR is a very
sensitive probe of the immediate environment around an
alkali ion in solution (40,41), yielding both structural
and thermodynamic information (17,18). Calcium is an
ion of both biological and chemical interest. Its chemical
behavior is more representative of the chemistry of alkaline
earth metals than magnesium, while its role in living organ—
isms is well known. It was desirable, therefore, to use the
sensitivity of multinuclear NMR in the study of calcium
interactions in solution. Due to the high cost Of enrich—
ment, natural abundance studies seemed to be of more
practical value as a preliminary effort.

The system which generated interest in this technique
was the aqueous solution complexation of calcium with phos—
phonoacetic acid (PAA). However, basic solution inter—
actions, such as ion—pairing, were studied first, as they
influence the complexation behavior. The applicability of
calcium-43 NMR to the complexation phenomena itself was

then tested with three known systems. Finally, a study

22

 

23
Of the calcium—PAA system was then attempted.

3.2. General Information on Calcium—43 NMR

The properties of the calcium—43 nuclei (the only
stable Odd spin calcium nuclei) are not as favorable as
some of the alkali metal nuclei for the use as a probe
nucleus in NMR. The natural abundance Of the calcium—43
isotope is 0.13%. The sensitivity of NMR towards calcium—
43 is 0.064 (at a constant field), relative to a proton NMR
sensitivity of one. A more accurate reflection of the dif-
ficulties associated with calcium NMR is a comparison of the
product of sensitivity and natural abundance with other
nuclei, as shown in Table I. Calcium-43 has a lower overall
NMR sensitivity than any of the alkali metals. In fact, it
has a lower sensitivity than most nuclei studied today.

The nucleus has a spin of 7/2 and a quadrapole moment of
-0.065b; these factors lead to a fairly narrow natural line
width (N10 Hz). The resonance frequency is somewhat low

at 12.12 MHz (at a field of 42.88 kG), which is experi—
mentally exascerbated by the fact that it is on the border—
line between the frequency ranges of two probes we used to
Observe this nucleus.

Calcium salts do have an advantage (as Opposed to Li
or Ms) in that they do not have a high affinity for water.
The attainment of water—free salts for nonaqueous studies

is then rendered a difficult, but not an impossible task.

 

24

Table I. Properties of Selected Nuclei in NMR Spectroscopy.

 

 

 

Quadrupole Natural Overall
Nuclei Spin Moment Abundance Sensitivity Sensitivity
7Li 3/2 —0.040 93 0.29 27
23Na 3/2 0.11 100 _ 0.093 9.3
39x 3/2 0.09 93 0.0005 0.047
87Rb 3/2 0.14 27 0.18 4.9
l330s 7/2 -0.003 100 0.047 4.7
25Mg 5/2 0.22 10 0.027 0.27
”30a 7/2 V -0.003 0.13 0.064 0.0083
67Zn 5/2 0.16 41 0.0029 0.12
75As 3/2 0.3 100 0.025 2.5

31E 1/2 --— 100 0.066 6.6

 

 

 

 

25

Once dried, the salts are sufficiently non—hygroscopic to
assure that water is not producing a measurable contamina—
tion effect when solutions of these salts are studied in
nonaqueous solvents.

This study has elucidated general trends in the solu—
bility of calcium salts in nonaqueous solvents. The trends
tend to follow the ”hard, soft” concept of solvent donicity
(42), rather than a Gutmann donor number (43) or dielectric
constant relationship. That is, the "hard” calcium (rela—
tively high charge to mass ratio) will prefer to bind to the
”hard” oxygen donating solvents rather than the "soft"
nitrogen donating solvents. This trend can be seen when
considering the solubility of calcium nitrate in acetone,
pyridine, and acetonitrile. Calcium nitrate is more sol—
uble in acetone (donor number = 17.0, dielectric constant
3 = 20.7) than in pyridine DN = 33.1, e = 12.3) or aceto—
nitrile (DN = 14.1, e = 38.0), even though pyridine has a
hiSher donor number and acetonitrile a higher dielectric
constant than acetone. Within a given solvent the nitrates
and perchlorates are more soluble than the halides, due to
their bulky size and lower charge—tO—mass ratio. With respeCt
to calcium salts, it may be generalized that the lower
surface charge density counterions (0104' and N03-) tend

to be more soluble due to less ion—ion interaCPlon-

 

26

3.3. Aqueous Studies — Chemical Shifts

 

The study by Lutz (5,6) represents the only previous
natural abundance calcium—43 NMR investigation of the be—
havior of some calcium salts in aqueous solutions. The
results are shown in Figure 2, where the units of concen—
tration have been changed from the-author's original units
Of moles solute/mole water to the more familiar molarity
scale. The results of aqueous NMR studies of calcium
salts in this work are shown in Figure 3. It is important
to note that in Lutz' study the lowest limit of detection
is 0-56 M calcium salt, while the lower limit of detection
for this study is 0.15 M Ca<ClOy)2. This was obtained after
approximately 15,000 scans (N4 h) at 8K memory and 5000
Hz sweepwidth. The minimum signal—to—noise ratio neces—
sary before a chemical shift was recorded Was 3. The
SiSnal-to—noise ratio was calculated by measuring the av-
erage noise around the baseline and comparing it to the
height of the M3Ca resonance peak from that baseline.

The difference in detection limits is due to the use
Of a different magnetic field strength in the two studies.
Lutz used a 1.807 T field whereas this study was done on a
4.288 T instrument. The increase in sensitivity of ap-
proximately 3.5, as measured by comparing the lower limits
Of detection, was due to the higher field, as well as the
increased field stability of a superconducting magnet 12

an iron core magnet. The greater stability OI the

 

N'TI

27

 

 

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28

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29

magnetic field in this study also resulted in a lower error
associated with a given measurement. The estimated error
for any given chemical shift by repetitive measurement,
in this study was i0.2 ppm, while Lutz found errors of ap—
proximately i2 ppm.

Aside from these differences, the experimental data
of the two studies agree very well; For example, the con—
centration dependence of the chemical shift was identical
within the experimental error. The conclusion drawn from
this comparison is that the behavior of the resonance sig—
nal was reproducible, even if a different instrument was
used. The demonstration of such reproducibility was neces—
sary before the technique could be accepted as scientif—
ically and analytically valid.

The relationship between the observed resonance signal
and concentration in a system characterized by fast ex-
Change of the cation between the two sites (the free ion and

the ion-pair) is given by the expression
=x<S+x.cS. (3)

where 6 5

. a]? h S 1. d 1? , and. l .

p:
paired chemical shifts, respectively, while Xf and Kip
are the mole fraction of calcium ions in the free and ion—

paired state. In the case of the chloride and bromide

1 t n water.
salts, the halogens are better eleCtron donors cha

 

4 .

.1

 

30

An ion—pair with a halogen will have a 610 further down—

I.

field than that of the water—solvated, free calcium ion,
due to the increase in electron density around the nucleus
when the halogen replaces water. At infinite dilution the
Xip is essentially zero (no ion—pairing), increasing with
increasing salt concentration. The result is a gradual
shift in the resonance signal downfield with increasing
concentration. The nitrate and perchlorate curves can be
similarly predicted with the exception that the nitrate and
perchlorate ions are worse electron donors than water,
leading to a gradual upfield shift with increasing concen-
tration.

Aqueous conductance studies of 0aCl2 (44,45) and Ca(NO3)2
(45), as well as a Raman study of Ca(N03)2 (46) tend to con—
firm the calcium—43 NMR evidence for ion—pairing. The
conductance of aqueous CaCl2 and Ca(NO3)2 solutions de-
creases with increasing concentration of salt. This be—
havior has been attributed to the formation of a neutral
ion‘pair as the salt concentration increases. The Raman
evidence for ion—pairing is more direct, as the intensity
of a band assigned to the ion-pair was observed to increase
in intensity (as a free NO3_ signal decreased) With in-
creasing concentration.

While the NMR, conductance and Raman studies prediCt
the presence of ion—pairing, the Ca012 conductance studies

- . 1 , w .t t tbsp
indicate that ion-pairing occurs to a greatei ex en 'a

. _ . a - Le ion-pairing
seen oy NMR studies. Conductance data -ndlcat o

 

I

u

 

31

of Ca012 in solutions with a salt concentration of 0.01
M, whereas the NMR results indicate ion—pairing becomes
significant only at concentrations of approximately 2.5 M
and larger. This apparent contradiction can be rationalized
through the realization that NMR and conductance techniques
measure different phenomena. Conductance data measure the
total amount of ion-pairing, while NMR data will only
measure ”contact" ion—pairing (47,48). Nuclear magnetic
resonance spectroscopy, observing predominantly nearest
neighbor effects, would be insensitive to solvent-separated
ion—pairs. Hence the extent of ion-pairing in Ca012 as
Observed by calcium-43 NMR is less than the extent of total
ion-pairing observed by conductance experiments. A similar
argument applies to Ca(N03)2, although the agreement
between conductance, Raman, and NMR data is better.
Equation (3) may be rewritten through the use of

equilibrium and mass balance equations, to yield

M 1/2
_ u
6 = [_II:I:1EEEIZ.__] (5 _a_ ) + 5_ (4)
obs 2k CM f 1p lp
“a T

1 M
where Ka is the ion—pair formation constant and CT the

total salt concentration. As the total salt concentration

is increased, the term in brackets approaches zero, or

= (5)
Sobs 6ip

 

32

This is somewhat at Odds with the aqueous data, as there
is no sign Of a limiting behavior in the Observed chemical
shift. However, since the Ka for these ion—pairs is
assumed to be small (46), the attainment of limiting be—
havior would not occur until a concentration exceeding the
solubility of the salt.

Another consequence Of Equation (4) is that the con—
ductance data are not directly comparable to the NMR data.
In Equation (4) the Ka is the contact ion-pair formation

constant; hence there are two unknowns, K and 6i . At—

p
tempts to fit Equation (4) to the data using a non—linear

8.

curve fitting computer program (38) with two unknowns have

 

failed due to a lack of points in the limiting region.
This is why only qualitative comparisons can be made be-
tween the methods.

The behavior of Ca(010u)2 in aqueous sOlutions is
anomolous, with two processes apparently occurring in solu-
tion. From infinite dilution to ml.5 M, ion—pairing occurs,
with the chemical shift remaining constant from about 1—2 M.
The resonance signal then shifts upfield, apparently not
reaching a constant value. A possible explanation is the
formation of a Ca(ClO,_,)2 ion triplet at N2 M, This would
shift the resonance further upfield with increasing con~
centration, just as an ion—pair did initially. Aqueous
conductance data are lacking to support such a hypothesis.

The conclusion reached from this study is that CaCl2

 

 

 

 

33

is the salt of choice in aqueous solutions when it is
desired to minimize ion—ion interactions. The influence
Of the chloride ion on the calcium resonance signal be—

havior is less than that Of the other salts studied.

3.4. Linewidths

 

The linewidths (the full width at half-height of a

43A . . .

ca resonance Signal) do not vary appreCIably Ior the
aqueous calcium salts studied. The bromide and chloride
salts exhibited a linewidth of 5:2 Hz over the entire
concentration range. This agrees, within the experimental
error, with the values Of Lutz (6) and Parello (9). The
linewidth for the nitrate and perchlorate salts increases
with increasing concentration from 5 Hz to approximately
15 Hz. This increase is probably due to viscosity broaden—

ing, as the solutions above 3 M tended to be viscous. The

same behavior was observed by Lutz (6).

3-5. Studies in Nonaqueous Solutions
3-5.1. Chemical Shifts

The behavior of calcium salts in nonaqueous solvents
was studied by “30a NMR in the following media: acetone,
dimethylformamide (DMF), dimethylsulfoxide (DMSO), ethylene
glycol, formamide, methanol (MeOH), propylene carbonate

(PC), and tetramethylguanidine (TMG). Some pertinent

 

 

34

properties of the above solvents are shown in Table II.
The solvents chosen represent a wide range of solution
properties such as, dielectric constant (ETMG ; 11.0 to
aformamide = 109.5), viscosity ( l5 cp for E0 to 0.32 Cp
for methanol at 25°C), and the Gutmann donor number (15.1
for P0 to 29.8 for DMSO). The Gutmann donor number (43)
is an empirical scale of the donor ability of solvents.
It is based on the enthalpy of the 1:1 complex formation of
the solvent molecule, S, with antimony pentachloride, both
dissolved in 1,2—dichloroethane as the solvent. The degree
of contact ion—pairing, as measured by multinuclear NMR
Spectroscopy, has been found to be dependent on bulk solu-
tion properties such as the donor number and dielectric
constant (49—51). Ions in solvents with low donicity and/
or low dielectric constants tend to show an increasing
tendency towards ion-pairing. ‘

The chemical shift results in the calcium—43 NMR
studies of nonaqueous systems are presented in Figures
4 and 5. The studies were limited in the number of salts
investigated due to solubility restrictions. The study
Of some calcium salts in methanol is shown in Figure 4.
The behavior of u30a chemical shifts for the chloride,
nitrate, and perchlorate is qualitatively similar to that
Observed for these salts in water. An increase in concen—
tration produces a downfield shift for chloride salt, but

an upfield shift for nitrate and perchlorate salts. As

 

 

35

 

 

 

 

 

 

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38

the concentration increases, perchlorate shifts further up—
field than the nitrate which is similar to its aqueous
behavior. Methanol is generally a weaker donor solvent
than water, although there is some disagreement as to the
donor number of water (43,52). The extent of ion-pairing
Should be greater in methanol, and this can be seen by the
behavior of the chloride salt. The chemical shift of cal—
cium chloride moves downfield even at low concentrations,
indicating a greater tendency towards ion—pairing than in
water.

In aqueous solutions, the chemical shifts for all salts
upon infinite dilution tended to the same value, which is
indicative of the free hydrated species (54). Similarly,
the infinite dilution chemical shifts in methanol should
converge to one value. This can be qualitatively seen in
Figure 4. It seems likely that, near infinite dilution,
the calcium nitrate and perchlorate shifts will vere
sharply downfield, much the same as in water. In methanol,
this shift would occur below the limit of detection. The
linear extrapolation of the calcium chloride line to
infinite dilution will probably more nearly represent the
infinite dilution chemical shift in methanol (estimated
at 1 ppm).

The variation in the chemical shift with concentra—

tion for calcium nitrate and perchlorate in ethylene gly—

1- ' —Oairs.
001 solution indicates the formation Of stable ion 1

 

39

There is an upfield chemical shift with increasing con—

centration, with an apparent limiting value of the chemical

shift reached at approximately 1 M. While variation in the

chemical shift with concentration is too slight to computer

fit to Equation (4) (the computed errors in the calculated

Ka and dip are much larger than the computed values), it

is apparent that most of the calcium ions are ion—paired

by 1.0 M. Similar behavior has also been observed in

calcium nitrate solutions in PC. Again, the trend is

not large enough to analyze the data quantitatively, but

it appears that the calcium nitrate is completely ion-

paired at concentrations 3 0.25 M. From a consideration

of the low donor number Of P0 and low dielectric constant

Of ethylene glycol (donor number of which has not been

measured), the extent of ion—pairing is not unexpected.
Significant ion—pair formation is also indicated by

studies in DMSO solution. A limiting chemical shift was

obtained at approximately 1 M salt concentration. The

chemical shift data have been more amenable to computer

fitting to derive Ka and 51 The x§03 was found to be

p.
ClOu

0.8i0.2 and 6ip is -9i1 ppm, while the Re = 4:2 and

Sip = —2i1 ppm. The errors are large, but it can be seen
qualitatively in Figure 5 that the perchlorate reaches a
limiting behavior before the nitrate. A secondary source

of error, not considered in the computer analysis, is in

the estimation of of, with the total error rendering the

 

 

 

40

values trustworthy only to an order of magnitude. It
must be emphasized that in these studies the Ka determined
is a concentration ion-pair formation constant.

The study in dimethylformamide (DMF) and formamide Solu—
tion are interesting in that interactions with the calcium
ion should be present in each solvent, but in DMF these are
modified by the presence of the two methyl groups. In
the case of formamide, contact ion-pair formation is indi-
cated, although the extent Of formation is small. The
overall chemical shift range for the concentration study
was 1 ppm for the perchlorate and 2 ppm for the nitrate.
The counter-ion dependency is also indicated by the pres-
ence of two distinct curves representing the concentration
dependence of the chemical Shift, separated by 2 ppm.

This behavior is due to the strong solvating ability Of
the formamide. It has a high dielectric cOnstant, which
would tend to minimize ion—ion interactions. The study

in DMF produced somewhat unexpected results. There was
found to be no concentration or counter-ion dependence of
the chemical shift in the 0.2 - 1.5 M concentration range.
Therefore, either there is a quantitative formation of
contact ion-pairs, or complete dissociation throughout the
concentration range.

The co-linearity of the lines representing the con—
centration dependence of the chemical shift would tend to

eliminate the first possibility, as the electron density

 

 

 

 

41

around the nitrate and perchlorate ions are different;
this difference would lead to a different environment
experienced by the calcium ion for each ion—pair. The
second possibility is somewhat at odds with the current
theories of solvation. Bulk solvent properties, such as
the dielectric constant and Gutmann donor number, would
tend to indicate that DMF will not solvate the calcium ion
as we 1 as water or DMSO. Hence, one would expect more
ion—ion interactions as the solvating ability decreases.
However, neither the Gutmann donor number nor the di—
electric constant is, of itself, the best indication of
the solvating ability of a given solvent. The anomalous
behavior in DMF could indicate that the solvent forms a
strong complex with the calcium ion. The keto group of DMF
would have a higher electron density than formamide, due to
the inductive effects of the methyl groups. This could
lead to a strong complex which would be resistant to re-
placement of a DMF molecule by the nitrate or perchlorate.
Calcium nitrate in tetramethylguanidine solution also
shows somewhat anomalous behavior. The linear behavior of
the calcium—43 chemical shift with concentration can be
explained by ion—pairing, and it has been seen previously
in formamide and methanol. However, the direction Of the
Chemical shift change with increasing concentration is
anomalous. Replacement of the solvent by the nitrate

ion is expected to produce an upfield shift, yet in TAG

 

 

 

 

42

the shift is downfield with increasing concentration.

This may be explained by the structure of TMG Kg that of
the nitrate ion. TMG is a pure nitrogen donor (as Opposed
to a nitrogen and oxygen donor such as formamide), which
would have a lower electron density around the coordinating
group than an oxygen donor. Hence, although the negative
charge is Spread over three oxygen atoms, the donating
oxygens of the nitrate group have a higher electron den—
sity than the solvent. The result is a downfield shift
with increasing concentration.

Calcium ion solutions in acetone were also studied.
There is evidence for ion-pairing, as seen by the concen—
tration dependence of tne chemical shift, but the overall
change is slight. Counterion dependency is seen in the dif-
ferent lines obtained for the perchlorate and the nitrate.

Conductance studies on calcium perchlorate solutions in

IJ.

acetone (55) indicate the format on

U
i—IJ
}_J
O
:3

I
C
so
P
hS
m
p)
|._._J
Cr
D"
O
C

09
D‘

such studies predict a greater degree of ion-pairing than
. 4 3
is seen by Ca NMR.

In conclusion, calcium-43 NMR studies Show the tech—

nique to be very sensitive to the immediate environment of

 

ffects are minimal, which

(D

the calcium ion. Near-neighbor

renders the technique sensitive to contact ion-pairin

()9

r

only. This lack of sensitivity to long rang

‘4

(D

interactions

(.

is confirmed by comparing the extent of ion-pair formation

as seen by conductance (studies are sadly lacking in most

 

 

43

nonaqueous solvents) and NMR. Conductance studies fre—
quently indicate a greater extent of ion—pairing than

NMR studies. Raman spectroscopy studies of Ca.(NO3)2 solu-
tions (46) indicate that this technique is sensitive only

to contact ion—pairing. The degree of ion—pairing predicted
from Raman spectroscopy studies is closer to that predicted

Ll. .
from '3Ca NMR than conductance ion—pairing data.

3.5.2. Linewidths

Representative linewidths for all solvents studied are
presented in Table III. The linewidth is related to the

Spin—spin relaxation time through
: i6
and
= '26 " '—
l/T2 l/T2 + yAHO/Z . (7)

where T3 is the apparent spin—spin relaxation obtained

from the linewidth, T2 is the actual spin—Spin relaxation
time, Avl/Z is the full width at half height of the reson-
ance signal, y is the gyromagnetic ratio for a given nuclei,
and ARC is the contribution to the linewidth due to field
inhomogeniety. The contribution due to the field inhomo-

geneity is usually negligible for super conducting magnets.

 

44

Table III. Linewidths of the Calcium—43 Resonance Signal
in Some Solvents.

 

 

Full Width at l/2 Height (Hz)

 

Conc.
Solvent (M) Ca(NO3)2 Ca(010u)2
Acetone 0.25 22 14
1.25 ' 49 53
2.7 205 ——
DMF 0.3 7 7
1 1 10 15
DMSO 0.4 14 12
1 3 28 51
Ethylene glycol 0.4 40 13
l 2 53 26
Formamide 0.2 17 8
2 0 34 25
Methanol 0.4 6 6
2 2 15 10
PC 0.4 30 ‘—
0 7 36 “
TMG 0.4 37 --
Water 0.3 5 5

 

 

 

 

 

 

 

 

 

45

Since extreme narrowing conditions exist for this nuclei
under the conditions of this study (which implies T1 =
T2), the observed linewidth is then related toTl (the spin—

lattice relaxation time) by

1/2 = l (8)

Of the five principle relaxation processes: dipole-
dipole relaxation, chemical shift anisotropy, scalar coupl-
ing, spin rotation, and quadrapole relaxation, the latter

has been found to be dominant for nuclei with a spin

greater than l/2 (56). The rate of relaxation is then
given by
3 21+ n2
.L
: —— __._._.___.___.__ + _...._ f
31 (,0) < 12(21fl fl ><13><e2Q9/h>2C (9)
where R1 is the relaxation rate, I the spin of the nucleus,

v the assymetry parameter, Q the quadrupole coupling con—
stant, and Tc the translational correlation time. A change
in linewidth due to ion—pairing is usually caused by Chang-
ing the molecular symmetry (n) and quadrupole coupling
constant. The correlation time is related more to a dif-
ference in solvation of the ion from one solvent to another.
It can be thought of as the time required for a transla—
tional movement through the distance of a molecular diameter

and is hence related to tne solvent structure as well as ion

 

 

 

46

solvation. Solvents which are strongly bound to an ion
tend to have low correlation times with a subsequent narrow
linewidth (57).

The lack of ion—pairing information from chemical
shift data in acetone is compensated by a wealth of line-
width information. Acetone will weakly solvate the calcium
ion leading to fairly large linewidths at low concentrations,
where little ion-pairing is expected. A contact ion-pair,
however, drastically alters the symmetry and quadrupole
coupling constant, leading to a large increase in line—
width with increasing concentration. Hence, while chemi-
cal shift data do not seem to indicate significant ion—
paiiing, changes in the linewidth with concentration appear
to indicate that it is substantial. Such studies under-
line the importance of obtaining both linewidth and chemi-
cal shift information in an NMR study.

The narrow linewidth of the calcium resonance signal
in DMSO, methanol, DMF, and water solutions all indicate
a symmetrical, as well as relatively tight, solvation.
The linewidth data are further evidence for a strong
solvation/complexation of DMF with calcium. The study in
tetramethylguanidine indicate weak solvation (resulting in
a broad resonance signal) with an influence of ion—pairing
as the concentration increases. Propylene carbonate,
formamide, and ethylene glycol are also weakly solvated,

although how much this effect contributes to the linewidth

 

 

(I!

l

 

 

 

 

is unknown. The other contribution in the case of these
three solvents is line broadening due to viscosity. The
viscosity of ethylene glycol is approximately 10 cp, that

of PC 2.5 cp, and that of formamide 3.3 cp at 25°C. When
salt solutions are made on the order of 0.5 M in these
solvents, the viscosity is further increased. In these
three cases, viscosity is certainly a dominant factor in the

resultant broad linewidth.

3.6. Gutmann Donor Number Correlation

 

Previous studies in this laboratory (51,58,59) have

shown a correlation between the solvent donicity scale of

 

Gutmann (43) and the infinite dilution chemical shift of
given nuclei. Since the infinite dilution chemical shift
represents the solvation of the bare calcium ion, its
value would be a measure of the donating ability of a sol-
vent. The infinite dilution chemical shifts for a series
of solvents should then reflect the relative donating
ability of a given solvent with respect to others. The
correlation between the relative donicities expressed by
NMR and those formulated by Gutmann are shown in Figure
6.

The relatively large errors associated with the estima—
tion of infinite dilution chemical shifts in acetone, DMSO,
ethylene glycol, formamide, PC, and TMG are due to an

uncertainty in the behavior below the limits of

flUallifih

 

 

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49
sensitivity. The concentration dependence of the 43Ca
chemical shift for CaCNO3)2 and CatCIO4)2 solutions in
water shows that the chemical shift may change by 2 0r 3
ppm between the infinite dilution and the lower detection
limits. In fact, the infinite dilution 43Ca chemical shift
could only be obtained by studying CaCl2 solutions in water.
The downfield (higher electron density) trend in the
chemical shift at infinite dilution with increasing donor
number is logical, since a higher donating ability implies
a greater electron donation to the calcium ion. However,
the scatter of the points (a correlation coefficient for
the two parameters of 0.83), would indicate that the Gutmann
donor number is perhaps not the best indication of solvent
donicity. The donor numbers of TMG and ethylene glycol
are predicted graphically to be 44 and 23, respectively.

These numbers are qualitative, but useful for the predic—

tion of chemical shift behavior for other nuclei.

3.7. ComplexatiOn

 

Complexation studies were exploratory, being limited
in both SCOpe and depth- The calcium—ethylenediamine
tetraacetate system was studied using 43Ca NMR in aqueous
solution at a pH of 10. A plot of ligand—to—metal mole

ratio XE the chemical shift of the calciume43 resonance

is shown in Figure 7. There is a slight downfield shift

 

 

 

 

50

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51

of the calcium resonance between a mole ratio of zero to
0.4. In the 0.4 to 0.6 range a new resonance signal ap-
pears approximately 18 ppm downfield. This signal increases
in intensity, while that of the free calcium signal de—
creases, but the resonance frequency remains constant.
Beyond a mole ratio of 0.6, the free calcium signal is no
longer detectable. The linewidth of the free (7 Hz) and
complexed (8 Hz) calcium resonance remained fairly constant
over the entire mole ratio range. The presence of two
distinct signals indicates that the exchange process between
solvated and complexed calcium is slow on the NMR time

4.

scale. Similar results were obtained by Heubel (51) in the
study of the magnesium—EDTA system.

In principle the formation constant of the complex
could be calculated from the areas under the curves, which
are proportional to the relative concentrations of the
free and complexed calcium ion. In this case, however,
the signal-to-noise ratio was not high enough to allow the
accurate determination of the area under the peak (see
Figure 8). The areas can be approximated by a function of
the height of the peak, with the ratio of areas equal to a
ratio of peak heights (for approximately equal linewidths).
From the ratio of peak heights, the formation constant must

10
be large (it is approximately 10 under these conditions

(15)). Within the experimental error, all of the EDTA was

complexed by calcium below a mole ratio of one.

 

52

.+
me he <enm ca 2m.o age

so mote I m xmcm n+mmo comeQEoo I < xmcm .oaé
mfiomo :H S:.o soapsfiom m to cocmcomct szESHono

 

 

.m osswflh

 

 

53

The complexation of lSCS with calcium ion was investi—
gated in aqueous solution. A mole ratio study showed no
evidence for the formation of a lSCS-calcium complex. There
was no change in the linewidth or the chemical shift over
the mole ratio range O—2. This result confirms the results
of a calorimetric study of the calcium—1505 system by Izatt

(l9), and the emf measurements by Hoviland (60).
. 2+ 9
In the case OI the 18C6-Ca complex, the exchange Ol

2+ 0 y 1- O 0

Ca ion between the two Sites - free solvated ion and
. 43
complex — is slow, and two Ca resonances are observed
. 2+ .
in solutions containing an excess of Ca ion. The reson—
l 2+ . _ . . _ 4 .

ance of the Ca ion was apprOXimately 0 ppm upfield from
the free calcium resonance. This signal increased in
intensity with increasing ligand—to-metal mole ratio, while

showing no linewidth or chemical shift changes over the

mole ratio range 0 to 1. Slow exchange between the complex

and solvated calcium was also indicated by Dale and
Kristiansen (61), who conducted a proton NMR study of the

same system. No attempt was made to calculate the forma—
tion constant since the signal-to-noise ratio precluded
either an accurate peak area or height measurement.

The chemical shifts of the complexed resonances for
calcium—EDTA and calcium 1806 can be qualitatively ex—
plained by the electron density of the ligand. The EDTA
has negatively charged oxygen atoms chelating the calcium,

which have a higher electron density than the water. The

1806, on the other hand, has only ether atom bonding,

 

 

#—

54

with a lower electron density than the solvent methanol.
EDTA would then have a complex resonance downfield from
the free aqueous calcium, whereas the 18C6 calcium complex
has a resonance far upfield from the solvent methanol.

Calcium ion forms a sparingly soluble precipitate upon
the addition of the deprotonated form of PAA. The solu-
bility of this precipitate is less than 0.1 M, which means
that the sensitivity of calcium—43 NMR precludes its use
for this system.

Routine determination of complexation formation con-
stants using calcium—43 NMR is limited by the low sensi—
tivity of the method. Since the lower limit of detection

is 0.2 M, any complexation study with calcium would have

—

to be carried out at a very high ionic strength. Any con-
centration formation constants determined would have little
physical significance, as the ion—ion interactions

would abound, and these will complicate the system

beyond a simple complex formation model. Although the
chemical shift is a very sensitive probe of the immediate

environment around the calcium, limits of detection render

 

other techniques, e.g., potentiometry with an ion selective

electrode, more applicable to the study of calcium ion

complex.

 

 

 

55

3.8. Conclusions

4 . . .
The range of 3Ca chemical shifts due to complexation,
ion—pairing,or nonaqueous solvent interactions is approxi-
mately 60 ppm. This indicates the sensitivity of the tech—

V6

(1)
‘ _|

nique to the environment of the calcium ion. Exten
use of this method for the measurement of thermodynamic
prOperties is inhibited chiefly by the low natural abun-
dance of calcium—43. Enrichment is presently very expen—
sive, while natural abundance studies do not have reSpect-
able detection limits. However, the trend towards higher
magnetic fields will result in a dramatic improvement of
detection limits for this technique. A comparison between
the data of Lutz (5) and the present work reveal an increase
in precision and sensitivity attained by higher fields.

It is hoped that, as larger and more homogeneous fields
evolve, the technique of calcium-43 NMR will prove to be

useful in dilute solutions.

 

PART II

AN NMR INVESTIGATION OF THE REFERENCE ELECTROLYTE
ASSUMPTION FOR THE DETERMINATION OF SINGLE

ION MEDIUM EFFECTS

 

56

CHAPTER I

INTRODUCTION AND HISTORICAL

 

57

 

 

1.1. Introduction

 

Historically, most solution chemistry studies have
been conducted in aqueous media. Convenient standard
states evolved, e.g., an ideal one molal (or one molar)
solution of an electrolyte for activity, the standard
hydrogen electrode for electrochemistry, etc., which were
used as references for solution properties when comparing
them with another solution similarly referenced. There has
been increasing interest, however, in the solution chem-
iSth that occurs in nonaqueous solvents. Nonaqueous sol-
vents are advantageous when compared to water in that the
physical properties of the media (the acidity or basicity,
the dielectric constant, the electrochemical potential
window, temperature range, etc.) can be varied by changing
solvents to facilitate the observation of an interested
chemical phenomenon. The use of nonaqueous solvents
necessitated the adoption of standard states in that sol-
vent. In most cases for a given solvent, the standard
state was adopted analogous to that in water, e.g., the
hypothetical one molal solution, the standard hydrogen
electrode in that solvent, etc. Such a system leads to
as many standard states as there are solvents, with no

‘ ' . ,. 1 . 4 - ._ of
E 83123; TGlationsnip between them. Since it is

58

,--v

 

.4-

..'-

4. ,-
..—'

..,.

...‘

 

 

 

59

interest to compare the behavior of a system in different
solvents, a means of relating the standard states of various
solvents is necessary. If this could be done, fundamental
relationships such as a solvent-independent pH scale or
electrochemical potential scale, would be obtainable.
Water is usually chosen as the standard solvent, all
other solvents being referred to the analogous aqueous
solution. The problem then reduces to relating the stan—
dard states in water to those in the nonaqueous solvent.
The chemical potential of an electrolyte in a nonaqueous
solvent at a given concentration is given by
51 = 36; + aTthSai (l)
where 51 is the chemical potential, SE; is the standard
chemical potential, and sai is the activity using, the sol-
vent standard state (17). Similarly, the same equation

may be written using an aqueous standard state

- = To
Gi Wci + RTQnWai (2)

w ‘o l
here wGi and wai refer to an aqueous standard state and

aetiVitty‘. While Si is clearly the same in the two cases,
this does not ' ‘0 _ ‘0 = F‘ -
imply that wGi — 5G1, or Wa.l Sa1. squat
mg (1) and (2) yields
< co 6°) — ‘0 — "’ai <3)
S i - W i — AGEI‘_ RTQU T .Jl

s51

 

 

 

The activities can be expressed in terms of activity
coefficients and molalities. Since we are dealing with

one solution, Equation (3) becomes

 

(4)

But sYi approaches unity at infinite dilution, yielding

AG%P= Rith (in) (5)

This factor in is the activity coefficient of the electro—
lyte dissolved in a nonaqueous solvent, but based on an
aqueous standard state. Its value depends on the dif-
ference in the molal standard chemical potentials between
water and the nonaqueous Solvent. This term is given the
symbol in and called the medium effect, or transfer
activity coefficient. The quantity A§%I,iscalled the stan—
dard free energy of transfer. Conceptually, the transfer
activity coefficient is a conversion factor which relates
a solute‘s properties in water to those in some nonaqueous
solvent.

Equation (5) is valid for neutral molecules as well
as neutral combinations of electrolytes. This is as far
as a rigorous thermodynamical treatment will allow the
theory to proceed. However, Equation (5) is not in a

form which is too terribly useful, as many experiments

 

 

 

(e.g., the measurement of pH, the determination of the SHE
emf, etc.) require the knowledge of transfer activity co—
efficients of single ions. To obtain values for the trans—
fer activity coefficients of a single ion requires the use

of an extra—thermodynamic assumption.

1.2. Historical

The estimation of single ion activity coefficients has
been attempted using a number of different methods. Bjerrum
and Larsson (62) and later Iowa (63), and Parker et g1.

(64) estimated single ion activity coefficients electro—
chemically assuming a negligible liquid junction potential
at an aqueous and nonaqueous interface. Extrapolation tech-
niques, based on ion size parameters, were used by Izmaylou
(65), Feakins and Watson (66), and Deligny, gt g1 (67) to
calculate transfer free energies from aqueous to nonaqueous
solvents for some alkali metal ions. Similar extrapola—
tion studies were conducted, based on Buckingham‘s theory
of hydration (68), to determine transfer free energies from
water to N-methylformamide (69) and to some mixed aqueous—
nonaqueous solvents (70—72)-

A major area of interest has been in the use of a
standard reference, whose properties are solvent inde-
pendent. Pleskov (73) was the first author to propose
the use of this assumption when he proposed that the emf

Of a Rb+/Rb half—cell was solvent independent. Other

 

 

62

proposals for reference standards have been: iron(III)—
phenanthroline/iron(II)-phenanthroline (74,75), tetraphenyl-
arsonium—tetraphenylgermanium (78), tetraphenylborate—
tetraphenylmethane (79), Hammett indicator acids (80), tri—
iodide—iodine (81). The most popular of these assumptions
is that proposed by Koepp, g3 g; (76), that the emf of a
ferricene—ferrocinium (or cobaltocene—cobalticinium) half—
cell is solvent independent.

A more complete review and critique of these assumptions
may be found in the reviews by Popovych (77,82), Parker (83)
and Kolthoff (84). The most disturbing aspect of these
assumptions is their lack of agreement in calculated medium

effects, as seen in Table IV.

1.3. Reference Electrolyte Assumption

The most popular of the standard reference assumptions
is the reference electrolyte assumption (82). It postu—
lates that the transfer free energy for a given 1:1 elec— ,
trolyte can be equally apportioned between the cation and
anion. Obviously, the choice of the appropriate salt is of
critical importance. To date, three such reference elec-
trolytes have been introduced: triisoamyl-n—butyl ammonium
tetraphenylborate (85), tetraphenylphosphonium tetraphenyl—
borate (86), and tetraphenylarsonium tetraphenylborate (87)
(the latter two being the most popular, see Figure 9).

For a solid reference electrolyte in equilibrium with

 

 

 

 

63

Table IV. A Comparison of the Transfer Activity Coefficients
from Water to Methanol of Silver and Sodium Based
on Different Single Ion Assumptions.

 

 

 

Log

Assumption Ion in Ref.
MYAs(Ph)u+ = MyB(Ph)u— Ag+ 1.2 A
MYAs(Ph)4+ = Mvc(Ph)u 2.1 ,
MvB(Ph)u = MyC(Ph)u 0.4 B,C
MY = MYARX 2.6 B
My = MyCH3l 2.0 B
Ferrocene—Ferricinium -2.2 B
Ejunction = O 1.5 B
Extrapolation N_2+O 3.5 D
Extrapolation of R‘1+0 N Na+ 2.2 E
Extrapolation N_2+O 1.8 D
Extrapolation R_l+0 —2.23 F
Extrapolation f(R)+O —l.34 G
Ion—Quadrapole Extrapolation 0.4 H
Ion—Quadrapole Extrapolation —2.0 I

 

 

aPopovych, 0., A. Gibofsky, and D. H. Berne, Anal. Chem.,
33, 811 (1972).

bAlexander, R., A. J. Parker, J. H. Sharp, and W. E. Waghorne,
J. Amer. Chem. Soc., 95, 1148 (1972).

cCox, B. G., and A. J. Parker, J. Amer. Chem. Soc., 93,

3674 (1972).

dizmaylov, N. A., Dokl. Akad. Nauk. sssa, 149, 884, 1103,
1364 (1963).

elhid, 126, 1033 (1959).
Andrews, A. L., H. P. Bennetoo, D. Feakins, K. G. Lawrence,
R.P.T. Tomkins, J. Chem. Soc., A, 1968, 1486.

gAlfenaar, M., and C. L. Deligny, Rec. Trav. Chim., 8g
929 (1967).

hBax, D., C. L. Deligny, M. Alfenaar, Rec. Trav. Chim., 91,
_452 (1972). 7‘
lSalomon, M., J. Electrochem. Soc. lgg, 1609 (1971).

,

 

 

64

 

.COH ohmsocfizcoSQ

nonpop ccm WESHQOSQmOSQahgoLQoLpop .ESflcomLmflhcoceoLpop to cannostpm o£B .m onsmflm

=o_ owacoa_~=o=a~..oh

o©7©

=o_ =e_
_==_=e;amo==_»=ogaacfloh ==:=em.~_.=o;n~.~oh
Q. O Q 10
a
°=~ue
@ Q 32:

acea

 

 

65

a saturated solution of reference electrolyte in solvent

5, the chemical potential is

—
_ ~—
—

G . — G = O T
s solid 5 saturated 5G saturated + R1 insa+sa

(6)

and in water

. = G l = —0
8 solid w saturated wG saturated + RT inwa+wa~

(7)

Eouating the two yields

“'0 _.
AGt — RT 2n SKSD/ K (8)

wher T I re . - l .i “7 i ‘ ' -
e sésp and wKsp a the thermodynam_c SOIublllty pro

ducts in solvent 5 and water, respectively. If the sol-

ubility is low, activity corrections are negligible and
4G8? is obtained from a ratio of the concentration sol~

ubility products. From Equation (5), since the A53 is f

L)

or
the complete electrolyte, the medium effect is then

0 .... m w
AG‘CI’ — RI Q11 Y+Y_

"\
\O
\v’

where 7+ and v_ are the medium effects for the cation and
anion respectively. According to the reference ele~
assumption, the medium effects are equal, such that the

medium effect for a single refere ce ion is the geometric

mean of the cation and anion. By then using a different

 

 

 

 

 

 

 

 

 

66

 

anion (OT cation), a Y_ (or 7+) may then be calculated for

the new anion. This leads to an entire series of transfer

free energies for single ions.

The AG of a single ion may be broken down into

tr

D

Q
do
"5

= AGel + AG + AGSI (10)

where AGel is the contribution due to electrostatic inter—

actions, AGv is a neutral term, commonly associated with
.L

the energy required to produce a solvent cage the si

[\1
P1)

60

the ion, and AGSI is a specific term which accounts for
all interactions of a non-electrostatic nature (77). Since

the assumption involves the equivalent solvation of a cation

and anion (implied by the assumption of equivalent aGf

cl”

for each ion) the AGe should be negligible or minimal. Un—

1

less the specific interaction terms are equivalent for

cation and anion, these too should be minimal, leaving the

 

AGN term dominant.

The validity of the reference electrolyte assumption
has been tested both theoretically and experimentally.
The first experimental test of this hypothesis was by
Coetzee and Sharpe (88) in 1971. NMR investigations of

primarily the solvent chemical shift in the presence of the

 

reference electrolyte led Coetzee and Sharpe to conclude

that caution is advised when using this assumption. Ex—

(‘1‘

V
V
I4

perimental data from this study cast doubt on the validi

 

 

 

 

 

 

67

of assuming equivalent solvation of the cation and anion.
This method of investigation was continued by Caruso gg_gl.
(89-91).

The possibility of ion-pairingijisolutions of tetra—
phenylarsonium chloride in water, acetonitrile, ethylene
dichloride, 70% dioxane, and 95% dioxane was investigated
spectroscopically (UV) and conductometrically by Popov and
Humphrey (92). There was no evidence for ion—pairing in
water or acetonitrile, while the other three solvents ex-
hibited slight ion—pairing. This behavior can probably be
extended to the case of tetraphenylborate salt in water
and acetonitrile, with no ion—pairing expected.

The assumption was tested both theoretically and ex—
perimentally by Kim gt a}: (93-95). The electrostatic con-
tribution was calculated for interactions between ions and
multipoles of solvent molecules. The neutral term is de~
rived from experimental AGfiPfor the neutral analogs of the
ions: tetraphenylgermanium for tetraphenylarsonium and
tetraphenylmethane for tetraphenylborate. The sum of these
contributions was then compared to experimentally determined
Agn,for complete electrolytes, with substantial agreement
between the values within experimental error. An impor-
tant conclusion reached was that the AatnShOUld not be ap-

portioned equally between cation and anion, rather that the

cation should have a slightly greater portion.

The assumption was theoretically tested by Treiner (96)

 

 

 

 

 

68

using the scaled—particle theory (97). The scaled-particle
theory estimates the free energy required to form a cavity
the size of the reference ion, assuming equivalent solvent—
ion interactions in all solvents. He concluded the assump-
tion was not valid for the transfer of tetraphenylarsonium
tetraphenylborate from water to dimethylformamide (DMF),
dimethylsulfoxide (DMSO), sulpholane, propylene carbonate
(PC), and N-methyl-pyrrolid ne (NMP). These conclusions
were later refuted by Abraham and Nasehzadeh (98).

The reference electrolyte assumption was critically
compared to the ferrocene-ferricinium assumption by Kolthoff
and Chantooni (99,100). The conclusion, based upon experi—
mental data, was that the reference electrolyte is
preferred, since the values derived are chemically signif—
cant. Some data based upon the ferrocene-ferricinium as-
sumption leads to conclusions which are contrary to generally
accepted theories of solvation (i.e., the medium effect for
a proton transferred from water to methanol is approximately
zero, while all other evidence indicates that water is a
stronger base).

The general consensus of these studies was that the
standard reference electrolyte assumption (particularly
the tetraphenylarsonium tetraphenylborate assumption) is
a valid hypothesis for the determination of single ion
activity coefficients. Many studies then ensued (131—106,

tr

e ion activity coe--icients using

[*1

etc.) which measured sing

this assumption.

 

 

 

 

 

69

1.“. A Nuclear Magnetic Resonance Study

It was of interest to study the tetraphenylarsonium
tetraphenylborate single ion assumption using nuclear mag-
netic resonance. The NMR technique is very sensitive to
the immediate chemical environment around the probe nucleus,
which should yield valuable information about the solvation
and any interactions of these ions with the solvent or one
another. The only previous NMR study (88) investigated
primarily the solvent, whereas this study focused directly
on the ion. The nuclei 75As, 31?, 118, 130 and 1H of the
tetraphenylphosphonium, tetraphenylarsonium and tetraphenyl-
borate ions were studied as a function of salt concentra—

tion, counter—ion, and solvent.

1.“.1. Arsenic-75 NMR

 

 

Despite the fact that 75As has a 100% natural abundance
and a relative sensitivity of 0.025 compared to the proton
(overall sensitivity of 2.5, see Table I), this nucleus has
been little studied by NMR. This lack of interest is prob—
ably due to the high quadrupole moment (0.3 dB), which
produces a broad signal. Nonetheless, chemical shift and
linewidth data should provide some information on tne elec—
tronic environment around the nucleus.

The 75As resonance in liquid ASP; was observed b

V

Jones and Vehling (107) using a Knight high—level

 

7O

spectrometer. The linewidth was found to be mlS gauss at

a frequency of 8.8 MHZ, with no evidence of structure. The
magnetic moment and linewidth in aqueous solution were found
to be +1.4347 nuclear magnetons and W8 gauss by Dharmatti and
Weaver (108) in l95l. However, the first and only aqueous
study of various alkyl—arsonium compounds using Fourier
transform NMR was done in 1977 by Haliman and Pregosin
(lO9,llO). These authors obtained linewidth and chemical
shift data for a variety of compounds, including several
tetraphenylarsonium salts. The linewidth of the tetra-
phenylarsonium salts was found to be on the order of 1300
Hz, which renders the chemical shift data accurate to no
better than :7 ppm. It is no wonder that interest has been

lacking when such large errors (due to the high quadrupole

moment) are evident.

l.M.2. Phosphorous—31 NMR

Phosphorous-31 NMR has been studied to a much larger

extent than arsenic-75 NMR, due to its l/2 spin (no quadru—

pole moment) and overall sensitivity of 6.6 (relative

sensitivity of 0.06% and 100% natural abundance), as well

as its large biological significance. Indeed, several

volumes have been written on this subject (lll—ll3) as W611

. J
as extensive compilation of coupling constant data (ll4).

. . a 4. en
The tetraphenylphosphonium ion, however, has not be

 

71

well studied. The first study of this ion was in 1966 by
Grim, et al. (115), with the observed resonance appearing
23-2 ppm downfield from the 85% H3POM reference in deuterated
chloroform. The behavior of phenyl-phosphorous compounds

of penta—coordinated phosphorous were then studied by
Latscha (116). An attempt was made to derive an empirical
relation between the substituents and the observed chemical
shift. This line of inquiry was also pursued by Schmid—
peter and Brecht (117). A relationship between the chemical
shifts of substituted trioxyphosphorous and substituted tri—
phenylphosphorous compounds based on empirical data was
developed. The spin—lattice relaxation time T1 of organo-
phosphorous compounds was studied by Koole et al. (118)

and Wille (119). In the former study it was concluded that
dipole-dipole contributions to the relaxation rate were the
dominant factor at normal temperatures (25-6000). The

. - . l 31
latter study reached the same concluSions uSing DObh P

l

and 3C data.

1.4.3. Boron-11 NMR

There are two isotopes of boron, boron-ll and boron-10,
Which have an odd spin nucleus rendering them suitable ior

NMR study. Boron—10 has a relative NMR sensitivity of 1.99
X 10-2 with respect to the proton and a natural abundance

” * ' ' 't'i'tv of 0.165
01 19.58%. Boron-ll has a relative NMR senSitiJi y 1

 

72

and a natural abundance of 80.42%. The overall sensitivity

of 11B is 34 times that of 10

11

B, which led to the choice of
B as the nucleus to study.

Boron-ll studies of sodium tetraphenylborate were first
conducted by Onak, gt gt (130) in 1959. The authors studied
a number of alkyl and aryll boron compounds and published a
large compilation of chemical shift values. Later studies
by Noeth, gt gt (131), Thompson and Davies (132), and
Gragg, gt gt. (126) were concerned with a correlation
between the observed llB chemical shifts and the substi—

tuent electronegativity for a series of ions with the struc-

ture B<X)u—.

1.M.h. Carbon-13 and Proton

Carbon—13 NMR studies of tetraphenylarsonium and tetra—

phenylphosphonium salts have been few. The aforementioned

study by Wilke (119) was used for the assignment of the
13 ’ L w ‘ honium salts while
0 resonance peaks in tetlaphenylpnosp - ,

the work of Baliman, gt gt. (109,110) was used to assign

'1 . h
the ~L3C resonance peaks of tetraphenylarsonium Salts.

Contrary to an apparent lack of interest in the cations,

sodium tetraphenylborate has been studied rather exten-

SiVely using 130 NMR. All of the authors, however, were

interested primarily in 118—130 coupling constants, although

' ' ' lvents.
they did study the ion in a few nonaqueous so

73

Weigert and Roberts (122) were the first to study the llB—
130 coupling constants for this anion in D20 solution. Their
work was later followed by Axelson, gt_gt. (123), Hall, gt gt
(124), and Odom, gt gt. (125), who obtained essentially

the same results as did the original investigators. The
chemical shifts of the boron and Cl carbon as a function

of the substituent group (including tetraphenylborate)

were studied by Gragg, gt gt. (126). Finally, the effect

1 l

1
of simultaneous decoupling of 1B and H from the ‘30

spectra of tetraphenylborate was studied by Mazurek, gt gt.
(127).

The proton NMR behavior of tetraphenylarsonium, tetra-
phenylphosphonium, and tetraphenylborate ion has not been
studied in any detail. The proton NMR spectra of tetra—
phenylarsonium chloride, tetraphenylphosphonium chloride,
and sodium tetraphenylborate were observed by Coetzee and
Sharpe (88), but the authors were unable to distinguish
separate peaks for the different protons.
Spectra (121) of tetraphenylphosphonium iodide and sodium

tetraphenylborate are also lacking in resolution, with the

Signals appearing only as a complex multiplet.

1-5. Conclusions

_ x. (5 31p .L 72'
Very little work has been done on the As, -, t,
130: and proton NMR behavior of tetraphenylarsonium,
Such

tetraphenylphosphonium, and tetraphenylborate salts.

74

studies should yield useful information about the solva-
tion and solution interactions of these ions. This in-
formation, in turn, will help assay the validity of the

reference electrolyte assumption for the calculation of

medium effects.

 

 

 

 

 

RESULTS AND DISCUSSION

75

 

2.1. Introduction

 

The ions chosen as reference electrolytes (tetraphenyl-
phosphonium and tetraphenylarsonium as the cations, and
tetraphenylborate as the anion) have nearly equivalent
molecular structures (Figure 9). All have a central
atom, where the preponderance of charge is presumed to
reside, surrounded by bulky phenyl rings. These phenyl

rings serve two purposes: to ”shield” the central charge
0 L b’ o

 

from the solvent environment and to constitute the bulk of

the volume associated with the ion. Because of the ex-

 

tensive shielding, a minimal surface charge can be assumed.
This would tend to imply that the ion is solvated equi~
valently regardless of the charge on the central atom.

The large volume of the rings assures that both the cation
and anion have approximately the same size. Together,
these assumptions imply that, in the absence of ion—ion
interactions, the cation and anion should be equivalently
solvated.

Multinuclear NMR is a suitable technique for testing
the validity of the above assumptions. The sensitivit
of this method to the immediate environment of a nucleus
(both electrons and nearest neighboring atoms or mole—

cules) renders it a powerful tool for probing ion-ion

f ‘9

and ion~solvent interactions. The assumption 0 a central

76

 

_—

 

atom shielded from the solvent and other ions was tested
by observing the 75As, 31F, and 11B resonance signal as

a function of the solvent, concentration, and counter—ion.
The amount of surface charge was qualitatively estimated
by observing the 130 and proton chemical shift differences
between the cation and anion. The relative position of the
13C and proton chemical shifts of the cation with respect
to the anion can be related to the relative electron den—
sity at that carbon or proton. The extent and nature of
the ion—ion and ion-solvent interactions were investigated
by the observation of the relative positions of the 130

and proton resonances, as well as their behavior as a func—

tion of solvent, concentration, and counter-ion.

2.2. Phosphorous—31 NMR Results

The phosphorous-31 NMR results of the study of tetra—
phenylphosphonium salts in water, methanol, DMSO, DMF,
nitromethane, acetonitrile, propylene carbonate, and
pyridine solutions are presented in Table V. The values
are close to the chemical shift value obtained by Grim,
et al. (115) in deuterated chloroform. A comparison of

the results for different halide counter-ions (at the same

|_..

salt concentration), using a student t test (‘28), led
to the conclusion that there was no dependence of the ob-
served chemical shift on the counter—ion. Due to the low

solubility of tetraphenylarsonium (or tetraphenylphosphonium)

 

 

 

 

78

 

 

 

 

 

 

 

 

 

 

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80

tetraphenylborate, which fall below the NMR limits of
detection in most cases, halide salts of the cations were
studied. As such, any ion-ion interactions detected would
be of a ”worst case" sort, since the bulky, shielded tetra-
phenylborate anion would be expected to interact to a lesser
extent than a halide ion with the tetraphenylarsonium (or
tetraphenylphosphonium) cation.

The fact that there is no apparent counter-ion de—
pendence for the series chloride, bromide, and iodide in-
dicates that the ion-ion interactions have not penetrated
to the central atom. If such interactions were significant,
the effect of the higher surface charge density of the
chloride ion on the observed 31? chemical shift would be
the most pronounced, while that of the iodide ion would be
the least. This reasoning can be extended to the tetra-
phenylborate salt, with the lack of ion-ion interactions

between the cation and halide ions leading to the concluSion

. - i. - '0' J— '14.
that ion—ion interactions are not Slgnlllcanb for most salts

Of tetraphenylphosphonium ion.

There is a small concentration dependence to the ob—

served chemical shift which is probably due to changes in

, ‘ . o o -J_- +1 | «L—i n 86"
the bulk magnetic susceptibility of the so_ut o ( e

' ‘ ' '4‘ 'ts ‘ lent
Section 2-5-)Wlth concentration, with luo subsequ

- '-.'- k ' o~ '—
effect on the observed chemical shift. The average 00

. . l1 e e m. . ria‘ion
served chemical shift varies With the solvent. the va b

81

is 0.8 ppm (pyridine to acetonitrile) compared to a typical
standard deviation within a given solvent of 0.1 ppm.

The solvent dependence is significant, but not large enough
to attempt to correlate the average chemical shift for a
solvent with such solvent properties as the Gutmann donor
number or the solvent dielectric constant. It is con-
cluded that the phenyl rings do shield the cation from the
ion-ion interactions, but do not shield the central atom
(and hence the charge) entirely from the solvent, although

the effects of the solvent on the central atom are minor.

2.3. Arsenic—75 NMR Results

The results of an 75As NMR study of tetraphenylarson—
ium salts in water, methanol, DMSO, DMF, acetonitrile,
nitromethane, propylene carbonate, and pyridine solution
are given in Tables VI and VII. Due to a very broad line—
width, the accuracy of the chemical shifts are estimated
at i7 ppm. Within this error range, our values are
eQuivalent to those obtained by Baliman and Pregcsin
(109,110). There is no apparent dependence of the ob-
counter-ion,

served chemical shift on the concentration,

or solvent. The average chemical shift is 225il ppm
downfield from the reference with all values occurring
Within i2 ppm of this average.

If the tetraphenylarsonium salts exhibit an NMR be—

. ii +
havior Similar to the tetraphenylphosphonium salts, the

 

82

Table VI. The 75As Chemical Shift in ppm (:7 ppm) for Tetra-

phenylarsonium Ion as a Function of Concentra—
tion, Counterion, and Solvent.

 

 

 

 

 

 

 

 

 

Gone. in Gone. in

Molarity Chloride Iodide Molarity Chloride Iodide

Water Acetonitrile

0.305 227

0.100 225

0.050 225 0.010 224 225

Methanol Nitromethane

0.276 224

0.100 224

0.050 225 224 0.010 226 226

0.010 225 224

DMSO Prqpylene Carbonate

0.000 224 0.101 226

0.050 226 0.050 ' 224

0.010 224 224 0.010 225 224
i DMF Pyridine
i 0.102 224 0.100 225

0.050 224 0.050 225

0.010 224 223 0.010 224

 

 

 

 

 

 

 

 

83

Table VII. Some Representative 75As Linewidths for Tetra-

phenylarsonium Chloride in Various Solvents
(:80 HZ).

 

 

Conc. in Av1/2 in
Solvent molarity HZ

 

2 0.3 3830
0.1 1860

0.05 990

Methanol 0.3 740

0.01 450

DMSO 0.1 1670

0.01 1300

DMF 0.01 800

Pyridine 0.1 2100

. 0.01 1360
i Propylene Carbonate 0.1 1850
l 0.01 1150

 

 

 

 

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» 4

~-..-_

h I
- .33.)“
“4\. I|

I\)

 

84

total extent of change in the chemical shift would be 0.8
ppm. These changes would not be apparent when considering
the large error associated with any given chemical shift
measurement. The conclusion is that 75As NMR is not an
appropriate technique for such precise measurements of
chemical shifts. No conclusions on the validity of the
tetraphenylarsonium tetraphenylborate single ion assumption

75As

can be drawn from NMR measurements.

The trend in linewidths (full width at half peak
height), shown in Table VII, generally follows solution
viscosity trends. As per the discussion in Part I, Sec-
tion 3.5.2., the increase in viscosity, whether due to a
change in solvent or an increase in the solute concentra-
tion, will lead to an increase in linewidth. The trend
towards increasing linewidth with increasing solute con—
centration is evident in all solvents. A comparison of the
linewidths between solvents, at the same solute concentra-

tion, shows that the more viscous solvents will tend to

have broader linewidths.

2.4. Boron—11 NMR Results

 

The results of the 11B NMR study of sodium tetraphenyl~
borate in water, DMSO, DMF, acetonitrile, nitromethane,
propylene carbonate, and pyridine, conducted by Dr. Yukuo
Sasaki, are presented in Tables VIII and IX. There is a

slight concentration dependence of the observed chemical

 

 

 

 

 

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85

Table VIII. TheJJTSChemical Shifts of Sodium Tetraphenyl—
borate as a Function of Concentration and
Solvent in Various Solvents.

 

 

Gone. in 6 Gone. in 0 Cone. in 6
Molarity in ppm Molarity in ppm Molarity in ppm

 

H29
0.7 -26.26. 0.05 —26;18 0 0025 —26.29
0.5 -26.16 0 025 —26.21 O'OOlu —26.29
0.25 —26.19 0.01 —26.23 5X10" -26.3u
0.1 —26.20 0.005 -26.2u

2M§9
0.7 —25.72 0.05 —25.70 0.0025 —25.70
0.5 —25.70 0.025 —25.70 0.001 ~25.76
0.25 —25.68 0.01 —25.68 5x10“ -25.65
0.1 —25.65 0.005 —25.70

Nitromethane
0.4 —25.00 0.025 —2u.9u 0.001 —2u.89
0.25 —24.98 0.01 -2u.9u 5x10' —2u.85
0.1 —2u.96 0.005 —2u.98
0.05 —2u.9u 0.0025 —24.91

Acetonitrile
0.7 -25.u0 0.05 -25.3u 0.0025 —25 36
0.5 ~25.u0 0.025 ‘25.33 0.001 -25.37
0.25 -25.38 0.01 -25 35
0.1 —25.38 0.005 -25.36

Pyridine
0.u —25.38 0.025 -25.29 0.001] —25.26
0.25 —25.36 0.01 —25.27 5x10" ~25.27
0.1 -25.29 0.005 ~25.27
0.05 -25.31 0.0025 —25.21

 

 

 

 

fl

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."v'..v . -
r s. - .I_
I“ 9“.
“no-an .. J

 

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9 I -

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p‘

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a

1 ~
,A

- I _

N -

x. g _

 

 

86

Table VIII. Continued.

 

 

 

 

CODC- in 5 Gone. in 5 Gone. in 5
Molarity in ppm Molarity in ppm Molarity in ppm
Propylene Carbonate
0.7 -25.87 0.05 -25.89 0.0025 -25.86
0.5 —25.89 0.025 —25;84 0.001 -25.89
0.25 —25.8L1 0.01 -25.88 5><10‘LL —25.76
QME
0.7 —25.64 0.05 —25.57 0.0025 -25.57
0.5 -25.63 0.025 —25.55 0.001 -25.55
0.25 -25.59 0.01 -2555 5x10" -2551
0.1 —25.59 0.005 —25.56

 

 

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87

Table IX. The Infinite DilutionlJISChemical Shifts, Cor—
rected for Bulk Magnetic Susceptibilities, as

a Function of Solvent.

 

 

 

.Solvent 5m in ppm Solvent 5m in ppm
H20 -26.23:0.06 Pyridine -25.72i0.06
DMSO —26.l8:0.05 Propylene —26.23:0.09

Carbonate
Nitromethane —26.3li0.04 DMF —26.17i0.04

Acetonitrile -26.l3:0.07

 

 

 

 

 

ur‘l '_
C;i_.- J
r:.”‘ ’5

( 1'
(I)

(I)

(h

A.
1
R 0
~-_ 1f‘_
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-AW“- ,
.. - ,
J»...'~d L
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Nf‘D— (
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88

shift which may be attributed to changes in the bulk mag—
netic susceptibilities of the solution with changes in
salt concentration (see Section 2.5.). As can be seen in
Table IX, there is a slight solvent dependency of the in-
finite dilution chemical shift. As in the case of the

31P chemical shift data, the trend is too slight to at-
tempt a correlation between the infinite dilution chemical
shift and such solvent properties as the dielectric constant
or Gutmann donor number. The conclusion reached is that
the phenyl rings do not shield the central atom (and hence
the charge on the central atom) entirely from solvent—

solute interaCtions.

 

2.5. Carbon—l3 NMR Results

 

The results of the 13C NMR studies in water, methanol,

DMSO, DMF, acetonitrile, nitromethane, propylene carbonate,
. . . . . -1 1
and pyridine are presented in Tables X — XVli. The 3C

NMR data on sodium tetraphenylborate were provided by Dr.

 

Yukuo Sasaki. A sample spectra of each ion is shown in
Figure 10. Studies of tetraphenylphosphonium chloride,
bromide, iodide, and thiocyanate salts show that, within
the experimental error of i0.03 ppm, there is no effect
Of the counter-ion on the observed chemical shift. This
absence of counter—ion dependency indicates that in the

solvents studied, the anion-tetraphenylphosphonium inter—

actions tend to be minimal.

 

 

89

 

 

 

 

 

Table X. Carbon—13 Chemical Shift Changes with Concentra—
tion for Tetraphenylarsonium (As(Ph) +) and
Tetraphenylphosphonium (P(Ph)u+) Salts in H20.

Cone.

in

M01. Cl meta ortho para

As(Ph)th

0.750 122.81 133.33 135.19 137.04

0.498 122.94 133.36 135.36 137.04

0.252 123.13 133.33 135.52 137.04

0.098 123.32 133.26 135.69 137.01

0.047 123.45 133.27 135.79 136.92

P<Ph>u£l

0.400 120.05i0.04 132.73i0.03 136.91i0.03 137.84i0.05

0.301 120.1 io.l 132.66i0.03 136.94i0.03 l37.88i0.03

0.207 120.19i0.02 132.63i0.04 136.97i0.03 137.85i0.07

0.100 120.30i0.03 132.62i0.04 137.1 i0.7 137.76:0.05

0.050 120.37i0.02 132.60i0.04 137.21i0.04 137.8 i0.1

0°1M A5(Ph)u01 with Variable NaCl Concentration

0.101 l23.27:0.02 133.26:0.01 l35.65:0.02 136.97i0.02

0.196 123.25i0,03 i33.27:0.04 135.62:0.02 136.98i0.03

0.299 123.22i0.02 133.26i0.04 135.59i0.02 136.96i0.01

0.405 l23.20i0.01 133.29i0.02 135.60i0.02 136.98i0.03

0.489 123.20:0.03 133.29i0.04 135.60:0.02 136.94i0.03

 

 

 

 

 

90

 

 

 

 

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mo. H . .o
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mo.oeom.w H mo.onmm.smfl mo.os::.mmfl mo.oema.amfi mme.c

me mo.oe:o.smfl mo.oem:.mmfl c.0hmH.HmH a
Hzflnmvm
mm.WMWJ.QMH MO.OHmF.NMH :O.OHN:.MMH NO.OH©N.HNH mwm.m
Ho.ohme.wMfi so.ohm©.smfi mo.osom.mme HO.OH:N.HNH mom.o
HO.O| :.® H NOoOHm®.FMH MOoOHMJ.MMH HO.OHJH.HNH ©M.O
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mm.mmme.wmfi mo.osos.sma Ho.ohm:.mmfl No.0hmm.ema mom.m
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mmesmmem
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we smfi :H.mmfl Ho.:mH sH.:mH mom 0
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[somemHSCmndmepoe Lop goeumppcoocoo Que: mowcmso pmegm HMOH50QO mfileonnwo .Hx ofihme

w|JIII|L

 

 

91

 

 

 

Table XII. Carbon—l3 Chemical Shift Changes with Concentra—
tion for Tetraphenylarsonium and Tetraphenyl-
phosphonium Salts in DMSO.

Gone.

in

M01. C meta ortho para

ASfithug

0.401 124.42 134.45 136.66 137.80

0.299 124.51 134.46 136.70 137.86

0.176 124.53 134.41 136.64 137.73

0.100 124.60 134.54 136.74 137.92

0.050 124.54 134.45 136.89 137.77

0.040 124.66 134.57 136.84 138.04

0.030 124.55 134.60 136.66 137.90

0.026 124.54 134.42 136.80 137.77

130301)th

0.400 121.07i0.02 133.97:0.04 138.00:0.05 138.78:0.02

0.299 121.13i0.03 133-97i0.04 138.00i0.02 138.76:0.O4

0.201 121.15:0.02 134.00i0.05 138 l :0-1 138-8 :0-1

0.102 121.22:0.01 134.08:0.03 138.15:0.06 138.80i0.04

0.054 121.26:0.02 134.04:0.06 138.2 10.1 138-87i0.03

Eigfllugi

0.400 121.12i0.03 134.1 i0.1 138.05i0.06 138.72i0-Ol

0 302 121.13i0.01 134.07+0 02 138.1530.01 138.77:0.02

0.200 121.17:0.02 133.98i0 03 138.1130.05 138.83i0.04

0.101 l2l.25i0.02 134.0 :0.1 138.08:0.02 138.7410-01

0 050 121.25i0.03 134.04i0.05 138.1430.03 138.8 20.1

 

 

 

 

 

 

 

 

92
Table XII. Continued.
Gone.
in
M01. 01 meta ortho para
P<Ph>ul
sat. 121.1 i0.1 133.98i0.06 137.96i0.06 138.75i0.04
0.200 121.14i0.02 134.04i0.0 l38.04i0.02 138.73i0.01
0.102 121.2 i0.l 134.1 i0._ l38.11i0.02 138.8 i0.l
0.051 121.2 i0.l 134.0 +0.1 138.20i0.05 138.86i0.06
P(Ph)uSCN
0.106 121.22:0.01 134.0030.02 138.1630.03 138.8830 01
0.052 121.24i0.02 l34.01:0.01 138.14i0.02 138.9 i0.1
0.026 12l.24:0.01 134.00i0.03 138.15i0.04 l38.88i0.01
+
As(Ph)uB(Ph)u As(Ph)u 0 in ppm
sat
(W0.02g)124.61 134.58 136.74 137.86
B(Ph)u-
ortho meta para
139.08 128.85 125.07

 

 

 

93

 

 

 

 

Table XIII. Carbon-l3 Chemical Shift Changes with Concen-
tration for Tetraphenylarsonium and Tetra-
phenylphosphonium Salts in DMF.

Cone. 5 in ppm

111

Mol. 0 Meta Ortho Para

As(Ph)th

0.402 124.51 134.29 136.66 137.64

0.301 124.60 134.30 136.69 137.66

0.249 124.42 134.13 136.70 137.61

0.201 124.5 134.20 136 64 137.55

0.150 124.67 134.27 136.73 137.63

0 100 124 54 134.16 136.66 137.51

0.052 124.54 134.13 136.66 137.57

0.040 124.57 134.10 136.60 137.63.

0.024 124.54 134.13 136.72 137.54

P<Ph>49l

0.402 l21.07i0.02 133.60i0.01 137.79i0.06 138.40:0.03

0.298 l21.15:0.02 133.63:0.03 137.87i0.02 138.36i0.05

0.204 121.2l:0.02 133.62i0.02 137.97i0.06 138.57:0.03

0.097 121.26+0.01 133.6110.01 l38.05:0.05 138.65:0.02

0.052 121.30i0.02 133.6330.04 138.02i0.04 138.50:0.04

BLEElMBZ

0.398 l2l.07i0.02 133.56:0.01 137.81:0.02 138.48i0.03

0.199 121 20:0 O2 133.59i0.02 137.91iO-02 138-50i0°02

0.051 121.28:0.03 133.59:0.03 137.97io 03 138.48i0~08

21210.11

0.201 121.20:0.01 133.5930.O2 137.94:0.01 138.47i0.01

0.098 121.26i0.01 133.61:0.01 138.10i0.06 138.51:0.04

0.047 121.29:0.01 133.58:0.04 137.99:0.03 138 O i9~1

 

 

 

 

 

 

 

 

 

 

 

 

 

 

9 4 '
Table XIII. Continued.
Conc.
in
M01. Cl meta ortho para
+

As(Ph),lB(Ph),l As(Ph)u . in ppm

B(Ph)u— 134.11 136.64 137.52

Saturated 128.32 139.11 124.58

 

 

 

 

 

 

 

 

 

 

 

r_—_f

 

 

 

 

 

95

Table XIV. Carbon-13 Chemical Shift Changes with Concen—

tration for Tetraphenylarsonium and Tetraphenyl—

phosphonium Salts in Nitromethane.
Gone. 6 ,

in in ppm

M01. 01 meta ortho para
591321491
0.399 124.72 134.67 136.89 138.24
0.299 124.65 134.65 136.83 138.18
0.200 124.66 134.57 136.98 138.24
0.151 124.76 134.70 137.00 138.29
0.026 124.70 134.60 136.92 138.24
P<Ph)u_1
0-399 l2l.53:0.03 133.91i0.02 l38.19:0.01 138.93i0.03
O 300 121.59i0.01 133.94i0.04 138.2810.01 138.96:0.01
0.200 121.65i0.01 l33.92:0.04 138.29:0.02 l39.00:0.05
0.099 121.73:0.01 134.01i0.03 -38.39:0.05 139.10:0.04
0.052 121.76:0.01 l34.05:0.05 l38.41:0.01 139.1 :0.1
P(Ph)ugr;
O 299 121.58:0.01 133.96:0 01 138.2830.01 139.00:0.02
0.201 121.64i0.01 133.97i0 02 l38.31:0.01 139.0330.01
0.149 121.66i0.01 133.98i0 04 138.32i0.02 139.0330.03
0.099 121.66i0.01 133.96i0 05 138.33i0.03 139.05:0.03
0.051 121.75i0.01 133.99i0 01 l38.41i0.02 l39.12i0.02
P<Ph>ul
0.124 121.70i0.01 133.99i0.03 138.32:0.01 139.0330 05
0.094 121.69i0.02 134.00i0.02 l38.36i0.02 139 06:0.02
0.050 121.71:0.01 l33.98i0.05 138.33i0.03 139.13:0.04

 

 

 

 

 

 

 

 

 

 

 

 

 

96

 

 

 

 

Table XV. Carbon-13 Chemical Shift Changes with Concen—
tration for Tetraphenylarsonium and Tetraphenyl-
phosphonium Salts in Propylene Carbonate.

Gone. 0 in ppm

in

M01. 01 meta ortho para

82132139;

0.3625 123.67i0.01 l33.75i0.02 135.9210.02 137.23i0.01

0.1812 123.74:0.02 133.74:0.02 135.97i0.02 137.25i0.03

0.0906 123.77i0.02 133.76i0.01 135.99i0.01 l37.28i0.02

0.0493 123.76:0.01 133.76i0.01 136.00t0.02 l37.26:0.03

0.0227 123.81i0.02 133.78:0.02 136.02i0.01 137.2710.02

P(Ph>u_g__

0-3879 l20.61i0.02 133.17i0.02 137.34i0.02 138.16i0.02

0.3020 l20.65:0.02 133.17i0.02 l37.37:0.02 138.17i0.01

0.2019 120.71i0.02 133.15i0.03 137.38i0.03 138.18i0.02

0.0998 120.73:0.02 133.16:0.02 137.43:0.01 138.18i0.02

0.0528 l20.74i0.02 133.16i0.02 137.43i0.03 138.19i0.03

l37.36:0.02 138.15i0.02
137.42i0.02 138.18i0.02
137.43i0.01 138.18i0.02
137.43i0.01 l38.l8i0.02
137.40i0.01 l38.15i0.02
137.42i0.02 l38.l8:0.02

 

 

 

 

 

 

 

 

 

97

 

 

 

Table XVI. Carbon—l3 Chemical Shift Changes with Concen—
tration for Tetraphenylarsonium and Tetra-
phenylphosphonium Salts in Acetonitrile.

Gone. 8 in ppm

in

M01. Cl meta ortho para

148(1)?”qu

0.412 124.16 134.25 136.54 137.77

0.299 124.14 134.20 136.55 137.74

0.198 124.23 134.29 136.67 137.91

0.149 124.36 134.30 136.69 137.86

0.100 124.34 134.25 136.60 137.68

0.048 124.28 134.25 136.60 137.83

0.024 —————— 134.22 136.57 137.80

P(Ph)u01

0.300 121.13i0.07 133.66i0.03 137.9510.04 138.68i0.01

0.201 121.18: 01 133.66i0.03 l37.95i0.01 l38.68:0.01

0.100 121.22i0.01 133.65i0.03 137.97i0.02 138.61i0.08

0.051 l21.28i0.02 133.72:0.04 138.10:0.01 138.70:0.01

P(Ph)u§£

0-303 121.1 i0.1 133.6630.05 137.92:0.05 l38.63¢0.01

0.250 l21.19:0.03 l33.72:0.05 l38.00¢0.02 138.7 i0.1

0.200 121.18i0.02 133.64i0.03 l38.01:0.05 138.73i0.01

0.100 121.16i0.0l 133.68i0.05 138.03i0 01 l38.75i0.03

0.051 121.23i0.01 133.71i0.02 l38.10:0.02 138.73i0.03

0.099 121.25i0.0 133.73i0.03 138.04i0.02 138.68i0.01

0.051 121.27i0.0 133.65i0.06 138.10i0.03 138.67i0.01

 

 

 

 

 

98

 

 

 

 

 

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mm.mmH mm.HmH mm.mmH mo.o mm :NH HH.mmH sm.mmH mm.o
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99

OTthOfl meta

 

 

 

 

 

 

 

 

 

 

 

 

 

 

C.
para 1
L1 . a. . .
ortho‘
, m ta
1 Ci
para
14 L L 1 i 1 a L a 1 a L; 1 a 1
ortho "1713
para
C1

 

 

 

 

A 1 AL I ‘J+14;1-¥A4#_L_1 L__A L+J

: - l '
lgure 10. The 30 NMR spectra of: Upper - tetraphenylar-

sonium chloride, Middle - tetraphenylphos-
phonium chloride, and Lower — sodium tetra—

phenylborate in DMSO.

 

 

 

 

 

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100

As a test of this hypothesis, the concentration of
chloride ion was increased in the aqueous solution (while

the tetraphenylarsonium chloride concentration was kept

constant) by the addition of potassium chloride. The

 

resulting changes in chemical shift (see Table X) are
smaller than when the total concentration of tetraphenyl—
arsonium chloride is increased. The changes in chemical
shifts are within the experimental error of a given measure-
ment, with minor corrections for bulk magnetic suscept-
ibility changes with concentration (see next paragraph).

A further test was to measure the chemical shifts of tetra- ‘

phenylarsonium tetraphenylborate in the DMSO and DMF

 

(Figure 11). The low solubility of tetraphenylarsonium

 

tetraphenylborate in most solvents precludes further NMR
studies of this salt. The chemical shifts of the ions
together are the same, within the experimental error, as
those for the sodium salt of the anion and the chloride of
the cation. These data lead to the conclusion that, in

our systems, the ion—ion interactions are either absent,

 

1
or have no effect on the ‘30 chemical shift.

There is a slight concentration dependence (a maximum

 

of 0.3 ppm) of the observed tetraphenylarsonium, tetra—
phenylphosphonium, and tetraphenylborate chemical shift.
A portion of this change in chemical shift is due to the
change in bulk magnetic susceptibility of the solutions as

the solute concentration is increased. This change in

 

   

 

orthc

orthc

 

101

meta ' DMSO

 

 

 

 

 

  
   
 

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l3
tetraphenylborate in DMSO and DMF.

Figure 11. The C NMR spectra of tetraphenylarsonium

 

 

 

 

 

 

 

 

 

 

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102

the bulk magnetic susceptibility manifests itself as a
change in chemical shift through the Live and Chan (37)
relationship for a superconducting magnet:

= 68 + 2t/3 (Xref _ x ) (11)

corr obs sample

where the symbols have been defined in Part I, Section
2.5.4. The chemical shift correction due to the bulk
magnetic susceptibility for an iron-core magnet is

I

6corr = CSobs

4w/3 (Xref — ) (12)

(sample
The same sample was studied using both an iron-core and a
super—conducting magnet. The resulting difference in

chemical shifts is

0 — 08 = 2v(x

obs obs ref — Xsample) (13)

from which the quantity (X ) was determined.

ref _ Xsample
Using Equation (2), the correction to the observed chemi—
cal shift was then calculated. The results of such a study
are presented in Table XVIII. Bulk magnetic susceptibility
corrections account for an approximately 0.06 ppm change

in the observed chemical shift with concentration in

range from 0.4 to 0.05 M. The remaining change in chemi—

cal shift (0.09 ppm in DMSO) is probably due to solute—

 

 

 

 

 

 

 

 

 

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104

solvent interactions.

A change in the chemical shift upon dilution has been
observed previously for some aromatic systems (129),with
a changing solvation shell upon dilution used as an ex-
planation. However, the chemical shift change is slight,
indicating that the forces responsible for the change in
chemical shift are weak. Consistent with this explana-
tion is the shift of the para carbon resonance upon dilu-
tion. This position is the most sensitive to the solvent,
while being the least sensitive to electron density shifts
in the phenyl rings, since it is furthest from the charge
center. There is a small shift, indicating some sclute-
solvent interaction, although it is slight. As has been
shown previously (l29),the meta position is relatively
invariant with concentration, while the ortho position
will shift similar to the para.

The infinite dilution chemical shifts of tetraphenyl-
arsonium, tetraphenylphosphonium, and tetraphenylborate
ions in different solvents are shown in Table XIX. There
is a definite solvent dependency of the infinite dilution
chemical shifts, which implies that a given ion is sol-
vated differently in each solvent. These data are con—
sistent with the non—zero 05;? determinations of several
authors (33,40,42-46) for these ions, since an equivalent
solvation in two different solvents implies a transfer free

energy of zero. However, the infinite dilution chemical

 

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105

The Infinite Dilution 130 Chemical Shifts*for

 

 

 

 

 

Table XIX.

Tetraphenylborate, Tetraphenylphosphonium, and

Tetraphenylarsonium Ions as a Function of Sol-

vent.

Infinite Dilution 0 in ppm

Solvent C7 meta ortho para
P(Ph)u:
H20 120.40 132.58 137.20 137.75
Methanol 120.41 132.61 136.89 137.62
DMSO 120.78 133.54 137.69 138.37
DMF 120.71 132.99 137.45 137.94
Nitromethane 120.38 132.64 137.02 137.74
Acetonitrile 120.49 132.92 137.30 137.92
Propylene
Carbonate 120.41 132.80 137.09 137.83
As(Ph)u:
H20 ' 123.42 133.27 135.78 136.98
Methanol 123.36 133.16 135.32 136.70
DMSO 124.12 134.04 136.32 137.40
DMF 123.95 133.49 136.05 136.93
Nitromethane 123.37 133.27 135.64 136.86
Acetonitrile 123.68 133.54 135.91 137.12
PrOpylene
Carbonate 123.44 133.41 135.66 136.92
B<Ph)4: ortho meta para
H20 137.95 129.22 125.63
DMSO 138.54 128.38 124.54
DMF 138.39 127.65 123.86
Nitromethane 138.16 128.00 124.12
Acetonitrile 138.19 127.98 124.21
Propylene Carbonate 138.40 128.04 124.38
Pyridine 138.98 127.59 124.08

 

 

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106

shifts of these ions do not correlate with such solvent
property scales as the dielectric constant and the Gut-
mann donor number.

The crux of the tetraphenylarsonium (or —phosphonium)
tetraphenylborate assumption is the equivalency of solva-
tion of the cation and anion which implies that the charge
density at any given carbon on the ring should be the same
for the anion and cation, since a positive charge on the
cation surface would be solvated differently than a negative
charge on the anion surface. In terms of 13C NMR measure—
ments, equivalent solvation would result in the cation and
anion resonance peaks occurring at the same frequency for
a given carbon. The difference between the infinite dilu—
tion 130 chemical shifts as a function of the solvent are
shown in Table XX. This difference may also be seen in the
13C 1 g , . fi .

spectrum 01 tetraphenylarsonium tetraphenylborate in
DMF and DMSO, as shown in Figure 11. There is a significant
difference in the chemical shifts of correSponding ortho,
meta, and para carbons. The difference between the ortho
positions between the cation and anion are attributable to
the different charges on the central atom in each case.

The negative charge on the boron imparts a higher electron
density around the ortho carbon of the anion than that
around the cation. This model is incomplete, since it
does not consider electron density shifts within the ring

(86).

 

 

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107

Table XX. The Difference in 130 Infinite Dilution Chemical
of Tetraphenylarsonium (and Tetraphenylphos-
phonium) vs Tetraphenylborate Ion as a Function
of Solvent.

 

 

45 (5As(Ph)A+ - dB(Ph)u- in ppm

 

 

Solvent ortho meta para

H20 -2.2 ' 4.1 11.4
DMSO -2.2 5.7 12.9
DMF —2.3 5.8 13.1
Nitromethane -2.5 5.3 12.7
Acetonitrile -2.3 5.6 12.9
Propylene Carbonate ~2.7 5-4 12.5

A6 (6P(Ph)a+ — 68(Ph)u_) in ppm
H2O —O.8 3.4 12.1
DMSO -O.9 5.2 13.8
DMF —O.9 5.3 14.1
Nitromethane -1.1 4.6 13.6
Acetonitrile —O.9 4.9 13.7
Propylene Carbonate -l.3 4.8 13.5

 

 

 

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108

The tetraphenylarsonium (and tetraphenylphosphonium)
ion may be represented by the resonance structures in
Figure 12. The electrons on the phenyl rings of tetra-
phenylborate would not be expected to resonate with the
electrons on the boron atom, as opposed to the arsenic
case, since all the electrons on the boron are bonding
electrons. Hence, the resonance frequency of the ortho
carbon on the anion is further downfield (higher electron
density) than that of the ortho carbon on the cation.

The para carbons are the least sensitive to such electron
shifts, due to their distance from the charge center,

while being the most sensitive to the solvent environment.
This carbon also exhibits the largest chemical shift dif—
ference between the cation and anion ( 13 ppm). The con—
clusion is that solvation is a dominant factor in the
resonance frequency of the para carbon, and that the solva—

- n
I

tion of the cation and anion are somewhat d: ferent. Since

the cation para position has a partial positive character,
it is reasonable to assume that it would attract the nega-
tive dipole of the solvent which would lead to a higher
electron density at the para carbon of the cation (with
respect to the para position on the anion), resulting in
the resonance signal occurring further downfield than the
anion signal. The meta position would be subject to both
electron shift and solvation forces, with the interpreta-

tion of the relative chemical shift rendered less definite.

 

 

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The C1 carbon chemical shifts (not shown) are dominated

by the effects of the charge on the central atom; the anion
chemical shift occurred the furthest downfield of any of
the other positions, while the cation occurred furthest

upfield.

The difference between the meta carbon chemical

shifts of the cation and anion were solvent dependent to
approximately the same extent as the para carbon, while the
ortho carbon exhibited a much smaller solvent dependence.

As mentioned previously, the interpretation of the para car-
bon chemical shifts are more straight forward than for

the metal carbon, therefore, although similar arguments

hold for the meta carbon, only the para position will be
discussed. The range of chemical shift differences between
the cation and anion was 1.72 ppm for tetraphenylarsonium

"\

phenylborate and 1.96 ppm ior tetraphenylphosphonium

‘1.

(1‘

1"

cr
m

e

(‘l‘

e h orate. The difference between the two salts

(D

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CT
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is t mall to make a statement on the preference of one

0

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cation over the other. There is a relationship between

the dielectric constant of the solvent and the difference

"l’\
in the *3C chemical shifts for the para carbon (Figure 13),
although the correlation is slight.
-+ , a . W 11 13p ~11
10 was also 01 interest to compaie tne c.chemicai

shifts of tetraphenylgermanium with those of the tetra—

l
-

>heny1arsonium ion since tetraphenylgermanium 1’ often

5;, v

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112

Table XXI shows the results of such a comparison in di-
chloromethane and deuterated chloroform solutions. The
two molecules are clearly not solvated equivalently, as
the peaks for the ortho, meta, and para carbon do not
coincide.

The conclusions of the carbon-l3 study of the tetra-
phenylarsonium (or tetraphenylphosphonium) tetraphenyl—
borate single ion assumption are as follows: in the solu-
tions studied the ion-ion interactions are negligible, and
that tetraphenylarsonium (or tetraphenylphosphonium), tetra—
phenylgermanium, and tetraphenylborate are not solvated

equivalently.

2.6. Proton NMR Results

 

The proton NMR results from a study of-tetraphenyl-
arsonium, tetraphenylphosphonium, and tetraphenylborate
ion in water, methanol, DMSO, DMF, acetonitrile, nitro-
methane, propylene carbonate, and pyridine solution are
presented in Tables XXII—XXIX. The proton NMR data on
sodium tetraphenylborate was furnished by Dr. Yukuo Sasaki.
The studies of sodium tetraphenylborate and tetraphenyl-
arsonium chloride in a variety of solvents are presented
graphically in Figures 14—19. The concentration dependence
3f the observed chemical shift of the tetraphenylphosphon—

o .H I mfi
ium ion as a function of the counter—ion is shown in iables

{XII-XXIX. No apparent dependence o: the chemical shlit

 

 

 

 

 

 

 

 

 

 

 

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130

(within the experimental error) was observed for the
chloride, bromide, iodide, and thiocyanate counterions.
Figure 20 shows a plot of the observed chemical shift
of tetraphenylarsonium chloride in water as a function of
the total chloride ion concentration (adjusted with KCl).
The concentration dependence is much less than when the
total salt concentration was varied. The variation in
chemical shift is due predominantly to magnetic suscept—
ibility changes (see Section 2.5.)with concentration.
Tetraphenylarsonium tetraphenylborate was studied in
DMSO and DMF, with no significant difference seen between
the observed chemical shift of the chloride and tetraphenyl-
borate salt of tetraphenylarsonium, or between the ob-
served chemical shift of the sodium and tetraphenylarsonium
salt of tetraphenylborate (see Figure 2l). These data
indicate that cation—anion interactions are negligible with—
in the experimental error of the proton NMR measurements.
Similarly to the 13C data, the observed proton chemical
shifts exhibit a concentration dependence. In DMSO solu—
tion a large portion of this chemical shift change (050%
of the chemical shift from 0.05 to 0.5 M) for tetraphenyl—
arsonium chloride was due to bulk magnetic susceptibility
corrections (see Section 2.6). The rest of the chemical
shift change may be attributed to a change in solvation

with the concentration of the solute.

The behavior in water (Figure 1H) was anomolous in

 

 

 

 

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DMSO

Sodium Tetraphenylborate

Tetraphenylarsonium

Tetraphenylborate M U b

 

 

 

 

 

 

 

 

 

 

 

 

 

Tetraphenylarsonium chloride

 

 

 

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09

ure 21. The proton NMR spectra of sodium tetraphenylbor-
ate, tetraphenylarsonium chloride, and tetra-
phenylphosphonium chloride in DMSO.

 

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133

that there was a pronounced curvature in the chemical shift
below a concentration of 0.5 M (the behavior of tetraphenyl-
phosphonium and tetraphenylborate ion were similar, although
the curvature was less pronounced). The extrapolation of
an infinite dilution chemical shift for tetraphenylarsonium
chloride was subject to a large error, but the shifts of
the ortho, meta, and para proton were generally among the
furthest downfield of all the solvents. The downfield
chemical shifts tend to indicate a strong interaction be—
tween the negative dipole of the water and the ion. Such
an interaction might cause a rotation about the As - C
bond, leading to steric hinderance effects for the meta
and ortho position (greater for the ortho proton). The
chemical shift of the ortho proton then shifts the greatest,
followed by the meta, with the para proton the least sus-
ceptible to such effects.

The infinite dilution chemical shifts of the ortho,
meta, and para proton of tetraphenylarsonium, tetraphenyl-
phosphonium, and tetraphenylborate ions are shown in Table

xxx. As with the 13c chemical shift data, there is a sol—

vent dependence of the observed chemical shift. However,

as shown by a comparison of the tetraphenylarsonium chloride

spectra in a number of solvents (Figure 22), the actual

shapes of the peaks do not vary much from solvent to solvent,

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134

Table xxx. The Infinite Dilution Proton Chemical Shifts*
for Tetraphenylarsonium, Tetraphenylphosphonium,
and Tetraphenylborate Ion as a Function of the

 

 

 

Solvent.
5 in ppm

Solvent Ortho Meta Para
man)“:
H2O 7.78 7.76 7.96
Methanol 7.635 7.679 7.898
DMSO 7.822 7.90“ 8.061
DMF 7.779 7.825 -----
Acetonitrile 7.747 7.800 7.991
Nitromethane 7.689 7.699 7.836
Propylene Carbonate 7.623 7.669 7.8H3

+

WM
H20 7.79 7.75 7.88
Methanol 7.663 7.684 7.770
DMSO 7.857 7.883 7.970
DMF 7.858 ——————————
Acetonitrile 7.787 7.816 7-939
Nitromethane 7.66“ 7.682 7.780
Propylene Carbonate 7.653 7.679 7.76
BgPh)u:
H20 7.45 7.20 7.07
DMSO 7.251 7 036 6.862
DMF 7.206 6.844 6.699
Acetonitrile 7.32Ll 7 050 6.897
Nitromethane 7-195 6 857 6.694
Propylene Carbonate 7.194 6 852 6.712
Pyridine 7.63M __________

 

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Figure 22. The proton NMR spectra of tetraphenylarsonium
chloride in different solvents.

 

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solvent dependency, while the metal and para position
exhibit approximately the same shift with solvent. There
is no apparent correlation of the infinite dilution chemi—

cal shifts with the dielectric constant or the Gutmann

donor number of the solvent.

The difference between the infinite dilution chemical
shifts of the cations and anion are shown in Table XXXI.
The difference may also be seen in Figure 21, where a
spectrum of tetraphenylarsonium tetraphenylborate in DMSO
is shown. The relative position of the ortho, meta, and
para proton shifts of the cations and anion are dominated
by solvation forces, as opposed to the carbons where ring
current effects were also significant. The proximity of
the negative dipole of the solvent causes the chemical shift
of the cation protons to occur further downfield (higher
electron density) than those of the anion. The para proton
has the highest electron density, as it is the proton
nearest the solvent, followed by the meta and ortho protons
respectively. A similar rational was used to explain the
position of the anion protons, except that the positive
dipole of the solvent is oriented towards the protons.

The relative chemical shift of the anion with respect to
the cation has been observed by Coetzee and Sharp (28),
although their resolution did not allow them to distinguish
between the ortho, meta, and para protons.

The difference between the infinite dilution proton

 

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137

Table XXXI. The Difference in Proton Infinite Dilution Chemi—
cal Shifts of Tetraphenylarsonium vs Tetraphenyl-
borate Ion as a Function of Solvent.

 

 

A5 (0As(Ph)u+ - 0B<Ph)u-) in ppm

 

 

Solvent A ortho A meta A para
H20 0.33 0.56 0.9
DMSO 0.57 0.87 1.20
DMF 0.57 0.98 —————
Acetonitrile 0.92 0.75 1.09
Nitromethane 0.M9 0.89 1.1M
Propylene Carbonate 0.U8 0.82 1.13
Ad (6P(Ph)u+ — 08(Ph)u_) in ppm
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DMSO 0.61 0.85 1.11
DMF 0.65 __________
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Nitromethane 0.M7 .83 09
Propylene Carbonate 0.51 0.83 06

 

 

 

 

 

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chemical shift of the cation and anion exhibits a definite

’2
solvent dependence. However, contrary to the 1JC case, the

ortho proton exhibited the largest solvent dependency,
meta less, and para the least. This relative order is
probably due to rotational changes about the central atom

- C bond. Solvent-solute interactions are probable causes

1
for such rotational changes; the extent of change is then
solvent dependent. This rotation was manifested in a change
in the solute chemical shift, with the ortho proton ex-
hibiting the greatest change. This effect was less pro—
nounced for the meta proton and probably negligible for

the para proton. The difference in the chemical shifts of
the protons with solvent is then a composite of two effects;
a direct solvent dipole — ion interaction and a rotational
phenomena caused by solvation. The former is dominant at
the para position, while the latter is dominant at the

ortho position. The major conclusion derived from this
proton NMR study is that there is a difference in solva—
tion of the cation and anion which is solvent dependent.

As in the 130 study, it was of interest to determine

the proton NMR chemical shifts of tetraphenylgermanium as
compared to tetraphenylarsonium ion. The results are shown
in Table XXXII. Clearly, the two molecules are not sol-

vated equivalently, since the corresponding peaks diifer

by an average of 0.29 ppm.

7n

The conclusion from the proton NMR study of the

 

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Table XXXII. The Proton Chemical Shifts of Tetraphenyl-
germanium and Tetraphenylarsonium Chloride in
Dichloromethane and Deuterated Chloroform.

 

 

 

6 in ppm
Solvent Molecule ortho meta para
Dichloromethane Ge(Ph)u 7.387 7.508 7.533
As(Ph)uCl 7 639 7.770 7.868
Deuterated Ge(Ph)u M.681 H.82M H.8M8
Chloroform As(Ph)uCl H.856 5.033 5.082

 

 

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140

tetraphenylarsonium (or tetraphenylphosphonium) tetra-
phenylborate system are the same as that from the 130

study: ion—ion interactions are negligible and that tetra—
phenylarsonium, tetraphenylphosphonium, tetraphenylgermanium,
and tetraphenylborate molecules are not solvated equiva—

lently.

2.7. Summary and Conclusions .

The ij, 11B, 13C and proton NMR data all lead to the

conclusion that, in the solvents studied, ion-ion inter-
actions between tetraphenylarsonium or tetraphenylphos-
phonium cation and tetraphenylborate anion are negligible.
Likewise, all halide or alkali salts of these ions will

.. o . o . o o I“(_ 0 ('1'! 31p
also exhibit negligible ion—ion interactions. -he -

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shielded completely irom the solvent, while ate in it te
that the phenyl groups undergo electron density shifts
within the rings so that the charge is distributed over

the entire molecule. Contrary to the original assumption

stated in Section 2.1, the phenyl rings do not effectively

insulate the central atom from the solvent. The most

1 L 3
important consideration is that the 3C and proton data

indicate that the cation and anion are not solvated
l 130 d t indicate that tetra—
valently. Also, the H and 0 a a l- l tn

1 - .'_ wan, ' 1m . Q 6. ‘1" t Sfll—
pflenylgermanium and tetraphenylalsonium ion ar 0 0

vated equivalently.

 

141

This study indicates that caution is advised when using
the tetraphenylarsonium tetraphenylborate reference
electrolyte single ion assumption. Several key assumptions
in this hypothesis have been shown to be questionable.

It should be emphasized, though, that NMR data provides
only qualitative information on the assumption, i.e., the
NMR data do not indicate to what extent (in terms of kilo—
calories) it is in error. The difference in solvation
evidenced by NMR measurements may be negligible in the
measurements of A0§r within the estimated experimental

error, i0.3 kcal (93),of such determinations.

PART

H
H
H

PHOSPHONOACETIC ACID STUDIES

142

 

CHAPTER I

HISTORICAL REVIEW

143

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1.1. Introduction

During a random testing of compounds with a cell cul~
ture screen, it was found that phosphonoacetic acid (PAA)
(Figure 23) was an effective inhibitor of some forms of
Herpes simplex viruses (1339139). This discovery led to
an interest in this compound as a possible therapeutic
agent for the treatment of other forms of Herpes virus.
Herpes is the genus of a deoxyribonucleic acid core virus
group which replicates itself in the nucleus of the host
cell (135). Some common forms of this virus which cause
disease in man are: herpes simplex (which causes blister—
ing of the lips, external nares, glans, prepuce, and vulva),
herpes labialis (commonly called cold sores), herpes—zoster

(the cause of shingles), varicella—zoster (chicken pox),
Epstein-Barr virus (responsible for infectious mononucleosis),
and cytomegalovirus (causes pneumonitis, cytomegalouvirus
mononucleosis). The list is far from complete when it is
considered that there are many more forms of Herpes virus
specific to animals. It is estimated that 10% of all
Americanssufftm’recurrent herpes infections, usually in

its milder forms of cold sores, while fully 70% have anti—
bodies which would indicate a past exposure to Herpes.

The therapeuti value of PAA in the treatment of

Herpes viruses was of interest because few drugs exist

140

 

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146

which have an antiviral action against Herpes. Most of
these drugs are nucleoside analogs (136,137), which hold
a potential for teratogenic or carcinogenic behavior. As
such, only one is commercially used in the U.S., while two
or three are available elsewhere in the world. Some of
these have been compared to PAA in their inhibitory ef-
fects. The compound 9-B-D—arabinofuranosy1 adenosine (Ara-
A) was found to be as effective as PAA against Herpes
Simplex virus type 1 (138), while showing no therapeutic
effect in clinical trials for Herpes Genitalis (139,140).
The therapeutic effects of Ara—A and PAA against Shope
fibroma in rabbits were the same (141), as were the ef-
fects of idoxuridine and PAA in the treatment of Herpes
Keratitis (142,143). There was a greater therapeutic ef—
fect by PAA than Ara-A against cytomegalovirus. The drug
tilorone hydrochloride had no inhibitory effect against
Herpes Simplex virus type 1, as compared to PAA, in mice
(144). Hence, PAA is as effective as any of the drugs
presently available for the treatment of Herpes virus, and
it does not exhibit a teratogenic or carcinogenic behavior

in biological studies.

1.1.1. Viral Replication

 

The Herpes virus are capsid protein enclosed DNA
strands. When the capsid encounters a cell membrane it

will merge with the cell untillfimanaked DNA is released.

 

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147

The DNA then migrates to the nucleus where it attaches it-
self to the DNA of the cell. The virus DNA will then manu—
facture ribonucleic acid (RHAHV) which leaves the nucleus
and produces a DNA polymerase enzyme in the cytoplasm.

This DNA polymerase then migrates back into the nucleus
where it replicates the viral DNA (DNAHV). The entire cycle
repeats itself, increasing the amount of DNAHV until the
cell begins to lyze. The encapsulated DNAHV then leaves

the cell to propagate the disease further in other cells.

Inhibition of the Viral Replication Path by

l—’
|_J
l\)

PAA

The basic pathway involved in the -replication of Herpes
virus specific DNA is shown in Figure 24. The enzyme DNA
polymerase reversibly binds to the DNAHV, which is
separated, one nucleic acid at a time, into two strands.
Deoxynucleoside triphosphate (dNTP) then binds to another
site on the enzyme. An internal reorientation occurs such
:hat the dNTP binds to the complementary nucleic acid on

:he template DNA forming an inorganic pyrophosphate in

HV’
:he process (145). The inorganic pyrophosphate is then
eversibly released, followed by the movement of the poly—
erase to the next nucleic acid link in the DNA chain.

he whole process is repeated until the entire strand of

JA has been replicated into two identical daughter strands.

There have been two mechanisms postulated (see Figure

 

 

 

 

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149

24) for the inhibition of Herpes virus of turkey induced

DNA polymerase by PAA (146). The phosphonoacetate binds

to the polymerase at the inorganic perphOSphate site (postu-

lated via a complexation with magnesium ion) and acts as a

competitive inhibitor of inorganic pyrophosphate in the

exchange reaction. The PAA can then act either to dis—
DNA(N+1)

sociate from the enzyme complex EPA , or it may under-

go a reorientation with the nucleotide at the 3' position
of the DNA template to form the nucleotide deoxynucleo-
side monophOSphate-phosphono acetate (dNMP—PA) and EDNA.

A kinetic study of the system by Leinbach, et al.
(145), yielded results which were not consistent with the
mechanism in which the PAA only binds and dissociates from
the EDNA<N+1> complex (a deadend inhibitor). The results
are consistent with the second mechanism, although the iso-
lation of dNMP-PA, which would confirm this hypothesis, has
so far eluded investigators.

From the proposed mechanism it is evident that pyro-
phosphate and PAA act similarly to inhibit DNAHV replica—
tion. That is, an increase in the PAA or pyrophosphate
concentration should inhibit viral replication. Studies
(145) have shown this to be the case, although the PAA
Was two or three orders of magnitude more effective than
the pyrophosphate. This is seen from the dosage require—

ments necessary to produce similar inhibitory effects,

with 100 to 1000 times more pyrophosphate required than

PAA.

 

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.2. Toxicity of PAA

'_1

While it is evident that PAA shows great promise in
the treatment of Herpes viruses, it is not free from side
effects. Studies by Meyer et al. (143) indicate that
intravenous injections of PAA, at the level of 300 mg/kg,
in rabbits often produced fatal tetanic muscular spasms,
although the same dose given orally or intraperitonally to
mice was tolerated (133,138).

It was found by Roboz 33 al. (146) that subcutaneously
administered doses of 500 mg/kg PAA were lethal to mon—
keys within two days. The resultant concentration of PAA
in the blood was 0.5 - 3 mg/ml. If the total disage is
reduced to 100 mg/kg (subcutaneously injected in 25 mg/kg
doses every two hours), the blood level of PAA is reduced
to the 1 mg/ml range, which is easily tolerated. For thera-
peutic trials the recommended blood level is a continuous
50 mg/ml, which results from a dosage of approximately

230 mg/kg.

Carbon-14 studies of labeled PAA by BOpp EE._1- (147)
have shown significant bone deposition of this drug in
rats, rabbits, monkeys, and dogs. The concentrations
found in dry femur after seven days were: 55 mg/g in
rats, 62 mg/g in rabbits, 24 mg/g in monkeys, 4 and 57
mg/g in adult dogs and puppies, respectively. Studies
on rabbits indicated that the drug, or a metabolite of the

drug, was retained in the bone for over two hundred days.

 

 

 

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151

These studies tend to implicate a secondary complexation
with calcium (aside from the primary inhibitory effect

postulated to occur via magnesium complexation), which

causes the observed side effects.

1.3. Analogs of Phosphonoacetic Acid

The antiviral properties of structural analogs of PAA
were investigated to understand the drug mode of action
and to attempt to minimize the toxic side effects. Studies

by Herrin at _l. (148) and Lee at al. (149), indicate that
substitution of functional groups for the carboxylic or
phosphono—moiety drastically reduces or eliminates the
inhibitory effect. The only two analogs among the many
tested which exhibit an inhibitory effect of the magnitude
shown by PAA are 2—phosphonopropionic acid and phosphono—
formic acid (150). The phosphonoformic acid (PFA) had
the same inhibitory effect as PAA, while 2-phosphonopropionic
acid was 50 times less effective than PAA. The analog 3-_
phosphonopropionic acid (3—PAA) has no therapeutic effect
against Herpes virus (150) and hence it is used as a con-
trol for comparison with the physicochemical properties of
PAA and PFA.

The PFA analog decomposes at low pH values (151,152),
which would tend to recommend it over PAA. It is as ef—

fective as PAA against Herpes virus of turkey and Herpes

simplex virus, but not effective against mutant strains

 

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152

resistant to PAA (150). This implies that it has a mode
of action similar to PAA as well as similar biochemical
prOperties; its physiological decomposition properties tend

to favor its use.

1.4. Chemical Properties of PAA and its Analogs

The compounds PAA and PFA, as well as various salts of
these ligands, were first synthesized by Nylen in 1924
(34); the 3-PPA molecule was synthesized in 1926 (35).
Originally, PAA was used primarily as an extracting agent
for rare earth ions (153), with particular interest shown
by Elesin (154,155) in its complexing ability with ameri-
cium, curium, and promethium ions.

The pKa values for this triprotic acid were determined
by Elesin (155), Mao gt gt. (156), Heubel and Popov (51),
and Stunzi and Perrin (157). The values of Elesin and
Mao have been criticized by Heubel (158), while the values
obtained by Stunzi gt gt., were reported simultaneously
with the article by Heubel (51). The study by Stunzi
was conducted at 37°C and I = 0.15 (in order to approximate
physiological conditions) and was therefore not applicable
to this work (25°C and varying ionic strengths less than
0.15). Heubel and Popov determined the pKés as a function
of both temperature and ionic strength, which led to the
use of those values in the present investigation. Stunzi

, , + +2 +2
and Perron studied PAA complexation With Mg 2, Ca 3 Cu

 

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153

+2 .
and Zn , although, again, at 37°C and I = 0.15. Heubel
conducted investigations of the complexation of PAA pri—
marily with Mg2+, while complexations with Ca2+, Sr2+, Ba2+,

and Cd2+ were briefly studied.

1.5. Exgerimental Techniques
1.5.1. Calcium Ion Selective Electrodes

The prototype of the commercially available calcium
ISE was deveIOped by Ross (159) in 1967. The composition of
such electrodes (Figure 25) are similar to glass electrodes
with the exception that the potential is developed across
a liquid ion exchanger rather than a glass membrane. The
behavior of the electrode, 312;: the sensitivity, selec—
tivity, stability, lifetime, etc., is highly dependent upon
the specific liquid ion exchanger used (160—163).

The calcium ion selective electrode has been used to
determine the dissociation constant of calcium sulfate di-

hydrate (164) as well as the formation constants for Ca2+

complexes with the following ligands: thylenediamine
tetracetate (EDTA) and nitrilotriacetate (NTA) (165),
malate, citrate, and trans—l,2-d,aminocyclohexane-N,N,N',N'—
tetracetate (166), tri and tetrametaphosphate (167), adeno-
sine triphOSphate (ATP) (168), ATP, EDTA, and ethylene
glycol-bis(2-aminoethylether)—N,N,N',N'-tetraacetate (EGTA)

(169), and citrate, malate, malonate, oxalate, EDTA, NTA,

 

 

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igure 25. A schematic representation of the Orion 90-20
calcium ion selective electrode.

 

 

 

 

 

 

 

 

 

sulfate, orthophosphate, tripolyphosphate, and pyrophos-
phate (170). In all cases the authors compared the values
obtained using a calcium selective electrode to literature
values obtained using other techniques (primarily poten-
tiometry using a pH electrode). The general consensus was
that the calcium ion—selective electrode was a valid analyti-
cal tool for the determination of calcium(II) ion aqueous

complexation constants.

1.5.2. Manganese Electron Spin Resonance

 

Although manganese(II) ion is of biological and chemi—
cal interest, its complexation solution chemistry has not
been investigated too thoroughly due to technical dif—
ficulties of studying this ion in solution. Direct electro-
chemical observation of this ion is difficult due to the
very negative electrochemical potential of the Mn2+/O
couple and its generally irreversible behavior in aqueous
solutions. The difficulty in obtaining an organic ion
exchanger that is highly selective for manganese, yet
suitable for a membrane, has hindered the development of an
ion selective electrode for potentiometric studies.
SpectroscOpic studies in the UV and visible region are
only plausible in cases where highly colored Mn2+ complexes
are formed, and hence, are not universally applicable.

Most other common techniques are either difficult destruc-
- 9

tive, not applicable, and/or time-consuming.

 

 

 

 

 

 

 

 

 

 

156

One technique which shows great promise in the study
of manganese(II) ion solution chemistry is electron spin
resonance (ESR). Similar to NMR spectroscopy, ESR responds
to species with an odd spin, except in this case it is the
odd spin of an unpaired electron. With five unpaired

electrons in a Mn(H20)g+ complex, the manganese ion is

 

ideal for observation with ESR. The similarity to NMR
spectroscopy also extends to the wealth of information
available from linewidth, intensity, and chemical shift
data on an observed manganese resonance.

The application of manganese ESR to the study of
aqueous solution chemistry was pioneered by Townsend (171)
in 1954. The complexation formation constants of manganese
with malonic acid, glycylglycine, glucose, and histidine
were determined. The behavior of manganese salts in
aqueous solutions was then investigated by Hayes gt gt.
(172) and Flato (173), while the mixed solvent behavior of
salt solutions was studied by Bard gt gt. (174), Vishnevskay
_e_t ~i. (175), and Burlamacchi (176). The linewidth and '
chemical shift of the observed manganese resonance were
found to be independent of the chloride and perchlorate
ion concentration below 0.05M, while the sulfate ion led to
a decrease in intensity and an increase in the linewidth
even at low concentrations. The conclusion is that, in

order to minimize effects due to the counterion, a chloride

or, preferably, a perchlorate salt of the manganese(II)

 

 

 

 

 

ion should be used.

Complexation studies since Townsend's have dealt
primarily with biomacromolecules. Blankenship and Saver
(177) studied the environment of chloroplasts washed in a
tris buffer and estimated the dissociation constants for
the binding sites to be 1.2 x lO-Ll M. Complexes of manganese
with nucleobases, nucleotides, nucleosides, and DNA were
studied by Basosi gt _t. (178,179), to obtain structural
information on these systems. The effects of substrates,
cofactors and substrate analogs on the complexation of man—
ganese(II) isocitrate dehydrogenase was qualitatively in—
vestigated by Levy and Villafranca (180). Correlations
between the appearance of fine structure and the binding
of individual substrates were found. Nicotinamide adenine
dinucleotide complexes with manganese were studied by Green
and Kotowycz (181), who found two metal binding sites with
formation constants of 640:90 and 88:13. Armstrong gt gt.
(182) investigated the kinetics and formation constants
of manganese binding to adenosine-3',5'—monophosphate de-7
pendent protein kinase from bovine heart. A rapid equi—
librium was found with the manganese(II) ion binding to a
nucleotide site and a protein site.

While most studies have been restricted to biomacro-
molecules, there is no g priori reason why this method

should not be applicable to the study of smaller ligand

complexation reactions with manganese ion.

 

 

 

 

 

 

 

 

1.6. Conclusions

The biological side effects attributed to the use of

PAA as a therapeutic agent would tend to implicate a com—

plexation reaction with calcium. It is then of interest

to investigate the solution chemistry of the calcium—PAA
system to determine the validity of this hypothesis, as

well as to elucidate thermodynamic information about the

complexation process. Since this drug is of therapeutic

value, it is also of interest to study its complexation

properties with other biologically important ions, espec—

ially manganese(II) ion. Together with the information

derived from a study of analogs of PAA it is hoped that

the tg vivo behavior of this drug may be better understood

from the derived tg vitro models.

 

 

 

 

 

CHAPTER II

MATERIALS AND METHODS

159

 

 

 

2.1. Reagents

All reagents and purification schemes are as given in

Section 1.2.1.

2.2. Experimental Methods
2.2.1. Electron Spin Resonance Spectroscopy

The aqueous solution chemistry of manganese(II) ion
was studied using a Varian E—4 Electron Paramagnetic Reson—
ance Spectrometer at a resonance frequency of 9.407 Giga—
hertz. The modulation frequency was 100 kHz with a peak—
to—peak amplitude of 12.5 gauss. A sweep width of 1000
gauss was used with the midpoint of the range set at 3550
gauss. Ten decibels of power were output to the sample
at a detector current of 300 mA, while the first derivative
time constant was set at one second. The total sweep time
of this continuous wave instrument was 8 minutes. These
parameters were the optimized conditions for the maximum
sensitivity with a minimal loss due to saturation of the
resonance signal.

A Wilmad aqueous EPR cell (catalog No. WG-812) was

used for all measurements. The cell is made of quartz,

with a small path length to minimize dipole absorption

160

from the solvent. Ten solutions were studied to yield one

formation constant.

2.2.2. Cyclic Voltammetry

Cyclic voltammetric studies were conducted with a PAR
Model 174—A polarographic analyzer.and recorded on a
Hewlett—Packard x—y recorder. The cell used is shown in
Figure 26. The main compartment of the water-jacketed cell
was thermostated at 25.0i0.1°C, while the reference side-
arm was allowed to equilibrate with the atmosphere. All
solutions were degassed for 20 minutes with deoxygenated
nitrogen. The nitrogen was deoxygenated by passage through
two solutions of ammonium vanadate in H01 (183), which is
oxidized in the presence of oxygen. The vanadate was re—
generated by reaction with a Zn—Hg amalgam in the solution.
The nitrogen was then washed by passage through water.

The working electrode was a hanging mercury drop elec-
trode (HMDE), the counter electrode, a platinum wire, and
a standard calomel electrode (SCE) was used as a reference.
The sweep rate was 100 mV/s. The half-wave potential of a
metal ion for a given solution was obtained by averaging
the maximum current for the anodic and cathodic waves.
Five such determinations were conducted on each solution
and the results averaged to yield the reported half—wave

potentials. Five solutions were required to obtain a

formation constant.

162

2.2.3. Potentiometry

Potentiometric measurements were conducted with a line;
powered Analogic (#AN2546) voltmeter with a range of i2
volts and a stability of iO.l mV. The voltmeter was con—
nected to the sensing electrode via a high impedance buffer
built locally in the Chemistry Department electronics
shop. The sensing electrode was an Orion (93-20) calcium
ion-selective electrode with an Orion (90-01) single junc-
tion reference electrode. The normal lifetime of the cal-
cium ion—selective electrode (ISE) is approximately 6
months due to the deterioration of themembrane ion—exchanger.
Normal symptoms of a deteriorating electrode were unstable
readings and a decreasing calibration slope. It was also
necessary to monitor the pH of all solutions, which led to
the use of a Beckman Expanded Scale pH meter and an Orion
(91—05) combination electrode.

Silver complexation studies were conducted using
a silver wire electrode. This electrode was pretreated in
nitric acid to roughen the surface. A calibration curve
yielded a slope of 63.3:O.6 mV per decade change in the
concentration of AgNO3, which is close to the theoretical
slope of 59.2 mV expected for a reversible system. It
was then concluded that this silver wire electrode ex—
hibited Nernstian behavior and was suitable as a specific

' t ' Ole
ion electrode. The counter—eleCtrode was the Slno

 

163

junction reference mentioned above.

Titrations were carried out using the cell shown in
Figure 26. The temperature was maintained by circulating
water through the outer jacket, with a degassing of the
solution also possible through the inlet arm. A Teflon
cap was used to enclose the cell, cutting down the solu—
tion evaporation. Stirring was accomplished with an air-
driven magnetic stirrer. The entire apparatus was en—
closed in a Faraday cage to reduce the "noise” detected
by the voltmeter. The "noise" is random fluctuations in
the observed potential due to static charges built up on
the equipment (and operator) and 60 cycle line noise from
house voltage sources. The fluctuations were on the order
of 80 mV outside the cage and i0.1 mV in the Faraday cage.
The use of the coulometric equipment (described in the
next section) in a pH stat mode caused major fluctuations
in the voltage readings from the pH and Ca ISE when a cur—
rent was passed through solution. The coulometric cables
themselves (due to leakage from the current source) were
also a source of noise, even though the current was not
switched on. Hence, whenever voltage readings were made
(pH and calcium ISE) the coulometric cables were isolated
by grounding them.

Thermal equilibrium was established initially by wait—
ing 20 minutes prior to the first measurement. Upon

. . - 1
additions of titrant, the solution pH was coulometrlcal—y

164

Reference
Sidearm

1

Water Inlet

 

 

 

 

 

 

 

 

 

 

l
Gas -
Water Scintered-glass
Jacket Inlet Frit

 

 

 

 

 

 

t ,
l

0-

Water‘ Outlet

 

Figure 26. A schematic representation of the pyrex glass
cells used for A — cyclic voltammetric studies,
and B — potentiometric studies.

165

adjusted/maintained (approximately two minutes), the solu—
tion allowed to mix and equilibrate (two minutes), and

the emf was observed and averaged over 30 seconds. The
total time associated with one titration point was then
approximately five minutes. At least 30 points were taken
to form a calibration curve, and 30 points for the titra—

tion curve.

2.2.4. Coulometer

The pH of the solution was recorded at the same time
the calcium ISE emf was recorded. The pH was maintained
at a constant value by the coulometric generation of protons
or hydroxide ions. The coulometric electrodes consisted of
a platinum rectangle (ml7 x 10 mm) anode and a platinum
grid (N30 mm in length, 32 mm in diameter) as the cathode.
A 12 mm OD glass tube, 15 cm in length, acted as the anodic
compartment. The end of the tube in contact with the solu—
tion was closed off by an anion transfer membrane which was
held in place with an ”O" ring and Teflon cap (158). The
anion transfer membrane was impermeable to the solvent, but
allowed the passage of the supporting electrolyte. The
electrodes were cleaned prior to usage by immersion in 6 M
HNO3 with a current passed through the electrodes for 60
then the other electrode, as the

seconds using first one,

cathode.

 

 

 

 

166

The electronics of the coulometer have been mentioned
previously (158) and will not be dealt with here. The use
of this equipment by Heubel (158) was for the generation
of hydroxide ions as a titrant for determinations of the
pKa's of PAA. The use of the equipment to adjust and main—
tain the pH at a constant value required less stringent
demands in performance, with the conclusion that this

equipment is more than adequate for this study.

2.3. Sample Preparation
2.3.1. Manganese Electron Spin Resonance

A calibration curve of resonance line intensity Kg

— —6
concentration in the range of 10 3 - 10 M was prepared by

dilution of stock solutions of Mn012 and measuring the
intensity of the resonance lines. The effect of a tris
buffer at a pH of 8 and the supporting electrolyte tetra—
ethylammonium perchlorate on the manganese resonance were

then tested, with no change in linewidth or decrease in

the intensity of the manganese signal resulting. Complexa—

tion studies were then performed at a pH of 8 and 4.5 (ad—

justed with tetramethylammonium hydroxide) at various

ionic strengths. The manganese chloride concentration

was approximately 10"3 M; the ligand concentration was

. 7 l
varied. In the case of strong complexation (log Hf > 4)

- 4 " . l.
the mole ratio of ligand—to-manganese varied irom 0 to

 

 

167

With weak complexes (log Kf < 4), the mole ratio was in—
creased to a maximum of 5:1 and when no complexation
was evident, the mole ratio covered the range from 0 to
10.

Since the manganese ion can be oxidized in basic solu—
tions by dissolved oxygen (184) (although the rate is slow),
the solutions were degassed prior to the measurements of the
ESR intensity. A comparison of these results to those ob—
tained when solutions were not degassed showed no sig—
nificant difference in the data sets within the experimental
error. The conclusion is that over the course of a single
experiment, approximately 3 hours, the oxidation of man—
ganese proceeds at a negligible rate.

The intensity of an ESR resonance due to the aquo
manganese complex was obtained by measuring the peak—to—
peak (minimum to maximum) height of the fourth peak from
the low field for a first derivative plot of magnetic field
Kg absorption and dividing by the receiver gain of the
instrument. The error associated with one measurement,
which is composed of solution preparation errors, sample
placement errors, sample fill factor errors, the change
in the cavity resonance with a change in sample, instru-
mental errors, etc., is estimated by repetitive measure-

ment to be 4% of the intensity.

 

 

 

 

 

 

 

 

 

 

168

2.3.2. Cyclic Voltammetry

 

Stock solutions of the metal ion and the ligand were
used to prepare solutions of varying concentration. The
metal ion concentration was on the order of 0.1 — 0.05 mM,
and the ligand concentrations was held in excess by a
factor of 2 to 50 times the metal ion concentration to
satisfy the restrictions of the Lingane and Buck equations.
The ionic strength was maintained with TEAP, while the pH
was monitored and adjusted with either tetramethylammonium
hydroxide or perchloric acid. The alkylammonium cation
tends to adsorb on the mercury surface of the HMDE. This
creates a positive sheath in the double layer of the elec-
trode which tends to repell the positive metal ions. The
net effect of this adsorption is to slow electrode kin—
etics such that some reactions appear irreversible. For
those reactions where this is not a problem, TEAP was used
because of its lack of complexation properties. When the
TEAP was found to cause irreversible behavior, the sup-
porting electrolyte chosen was then LiClOu. It is doubt—
ful that the lithium ion would form a complex with the
ligand, as it is very small and highly solvated.

In cases where metal ions exhibited irreversible be~
havior even in LiClOu solutions, the double layer was
altered using para—toluene sulfonate (185). The para-
toluene sulfonate (PTS) adsorbs on the mercury surface

much as the alkyl ammonium ion does, except that in this

 

 

 

 

 

 

 

 

169

case the ion is negatively charged. The double layer is
then negatively charged, with the positive metal ions
attracted to it. This increases the rate of charge trans—
fer to produce reversible behavior. In all cases, with
TEAP, PTS, LiClOu, tetramethyl ammonium hydroxide, and
perchloric acid, there was no shift in the half—wave po—
tential of the metal ion upon addition of these species.
This implies that these species are inert with respect to
complexation with the metal ion, as well as implying no
alteration in the charge transfer kinetics of the elec—

trode reaction.

2.3.3. Potentiometry

The potentiometric cell initially contained 25 ml of
supporting electrolyte solution, to which was added 0.05
ml of 0.01 M CaCl2 from a precision buret. Thermal
equilibrium was established, the emf recorded, and the
next increment of the calcium solution was added as des-
cribed in Section 2.2.3. The procedure was continued a
total of 4 ml of calcium solution was added. The cali-
bration range of emf IE calcium ion concentration was ap—
proximately 10"5 to 10_3 M. The metal ion was then titrated
with the ligand (approximately 0.005 M) following the pro-
cedure of Section 2.2.3. The titration was stopped when

the ligand to metal ion ratio was approximately two. In

 

 

 

 

 

 

 

 

170

this manner, changes in the liquid junction potential
were minimized.

The volume correction for solution expansion or con-
traction (186) when the system was studied at tempera-
tures other than 25°C was found to have a negligible ef—
fect on the calculated free calcium ion concentration.

The supporting electrolyte chosen was K01. Tetra-

 

alkyl ammonium ions tended to interfere with the measured
emf due to interactions with the organic membrane of the
calcium ISE, leading to somewhat large errors in the
calculated formation constant. Perchlorate ions also inter—
fere due to their penetration into the organic membrane.
Potassium chloride is the recommended supporting electro—

lyte for use with a calcium ISE (187), even though it may

 

potentially compete with the calcium ion for the ligand.
A comparison of the results using tetramethylammonium

chloride and K01, however, showed that these values agreed

 

within the experimental error of the K01 data. These
results suggest that a competition between calcium and

potassium ions for the ligand is small or negligible.

 

.0
.L

Conductance water obtained rom the laboratory of Dr.

Weaver was used in most cases. Some experiments were con—

ducted using water distilled from 2 x 1072 M KMnOu and

2

2 x 10' M KOH. When calibration curve results using

 

conductance water and permanganate distilled water were

1

compared to house distilled water, the nou e distilled

U)

 

 

 

 

 

 

 

 

 

171

water was found to contain residual calcium on the order

of 1076 _ 1077 M. Since the typical free calcium concentra—
tion in complexation studies ranges from 10-3_to 1077 M,
this represents at worst an error of 1%. The use of con-
ductance water is recommended to minimize this source of
error.

The contribution to the ionic strength from ions other
than the supporting electrolyte will cause an error of 5%
(worst case with IKCl = 0.02) in the total ionic strength.
During the coursecfi‘a single experiment, however, the
ionic strength will vary as the concentration of ions in
solution changes by less than 1%. The total ionic strength
may then be corrected for the complexation reaction by
summing the ionic strength due to the supporting electro—
1yte and the average ionic strength due to all other
species in solution (as measured over the course of the
experiment).

The use of the term "constant ionic strength", then,

while not strictly valid, is a good approximation.

2.4. Data Handling

The computer program KINFIT4 (38) was used to fit the
calcium ISE calibration curve data. This is a general
purpose non—linear curve fitting routine. The calibration

data were then fed to the program MINIQUAD (188) which

Cl‘

calculated the complexation formation constan s for the

 

 

 

 

 

 

deprotonated ligands with the calcium ion. When mono—

protonated ligands were studied, the calibration data were
given to the computer program shown in Appendix C, which
then calculated the formation constants for these complexes
with the calcium ion. Cyclic voltammetry, AG25°C’ AH°,

and A80 data were processed using KINFIT4.

 

 

 

 

 

 

CHAPTER III

 

RESULTS AND DISCUSSION

i_.l
~l
UK)

 

3.1. Complexation Studies of PAA
3.1.1. Manganese Electron Spin Resonance

A calibration curve of the ESR resonance intensity IE

 

Mn2+ concentration, resulting from 67 measurements, is

shown in Figure 27. The data were computer fitted to ob—
tain a slope of 3.04:0.02 x 106 and an intercept of 0.53:
0.03. These values were used to determine the concentra-
tion of free manganese(II) ions upon the addition of PAA.

he supporting e-ectrolyte or the tris buffer

Cl‘

No effect of

on the intensity of the observed manganese ion resonance

was found.

Prior to studies of the manganese—PAA system, the

 

general applicability of this technique to the investiga-

 

tion of aqueous solution complexation was tested. _t was
desired to determine the formation constant of a manganese
complex with a ligand that was both similar to PAA and
well studied. The manganese citrate complex fulfilled
these objectives very well. Citric acid is a triprotic
acid (as is PAA) with successive pH's (negative lOg of

the protonation constant) of pEi3 = 5.78, pK2 = 4.32, and
9K1 = 2.89 (63). The low value of pK3 allows the study of

the deprotonated citrate at a pH of % However, the

 

 

 

 

 

m

Hos: mo compocsu m

00.N~
I _

00.09

175

no zpflmcopCH mmm po>gomoo och 0o

+9 x 028 S 08265:
00.0 00.0 004

—
—
h

e-Ib

—.

S

.coapmspcoocoo

poad coaumgoflflmo < .sm ogswam

1... 9o...
Ll 98.
Ill. Odor

0.00N

AUSNEUM 883

1.... 0.00m

ll. 0.00m

 

n: cos

 

 

 

 

 

similarity in pK3 and pK2 render the determination of a
monOprotonated complex with manganese subject to a large
degree of error. Therefore, only the deprotonated complex
with manganese was studied.

The results, alon with values obtained lrom the litera—
ture, are shown in Table XXXIII. The manganese—citrate

formation constant was determined using both TEAP and

tetramethylammonium cLlo-ide (TMAC). The ion-pair forma—

tion constant for manganese chloride was found to be 3.7
from pH studies by Grzybowski gt gt (189). With the excep-
tion of reference d, all previous citric acid—manganese

(11) ion studies were conducted using chlorides as the

upporting electrolyte. A formation constant was determined

U)

with TJAC as the "inert” electrolyte so as to compare our
results with the above values, while a formation constant
using TEAP was compared to the data of Grzybowski gt gt.
(189). The values agree well when the difference in ionic
strength of the two solutions are taken into account. The

data of Grzybowski et 1 (189) are considered to be the

_——__-——-

most accurate, since ion-pairing and manganese hydroxide

formation were considered. Our work agrees with that of

l. (189) to 0.09 log K units when corrected

Grzybowski gt

to an ionic strength of 0.1 using the Debye—Huckel equa-

tion.

(a

The value obtained using TMAC is within the range

the other literature values. As

1‘.)

rather broad range) 0

 

 

 

 

177

Table XXXIII. Formation Constants of a Deprotonated Man-
ganese—Citrate Complex.

 

 

 

 

 

 

Supporting
log K Electrolyte Method Ref.
3.6 K01 Potentiometry (pH) a
and visible spectros-
copy'(competition)
3.69 NaCl Radio isotope ion- b
exchange
3.83 NaCl Radio isotope c
ion-exchange
3.74:0.03 (Me)uNCl Mn ESR this
' work
4.28:0.04 (Me)uN010u Mn ESR this
' work
4.15i0.02 (Me)uNCl Potentiometry (pH) d
a
I. E. Kalinichenko, Ukrain. khim. Zhur., gt, 92 (1970).
b

N. C. Li, A. Lindenbaum, J. M. White, J. Inorg. Nucl. Chem.,
12, 122 (1959).

8J. S. Wiberg, Arch. Biochem. Biophys., 73, 337 (1958).

dA. K. Grzybowski, S. S. Tate, S. P. Data, J. Chem. Soc.
(A), 241 (1970).

 

 

 

 

 

 

a test of self—consistency, the formation constant of the
manganese citrate complex using TMAC as a supporting elec—
trolyte was corrected for the formation of theMnCl+ ion-
pair using the ion—pair formation constant of Grzybowski
gt gt. (189). This calculated formation constant agrees
with the value obtained using TEAP as a supporting elec-
trolyte to within 0.04 log K units. Therefore, the tech-
nique is at least self—consistent under different eXperi-
mental conditions. This self consistency, together with the
agreement between literature values and those obtained in
this study, was considered sufficient evidence to indicate
the validity of this technique in determining manganese
formation constants.

In all of the complexation studies it was important to
establish that the signal intensity is due only to the
free manganese<II) ion. That is, there is no contribution
to the observed intensity from the complex. When
an excess of ligand was added so that manganese is almost
entirely complexed, the resulting intensity was due only
to the residual free manganese(II) ion. There was no
evidence for a contribution from the complexed manganese
ion resonance to the observed free manganese ion reson—
ance.

The manganese(II) ion can complex with both the de-
protonated and monoprotonated forms of PAA (as shown by

Heubel (29)) according to the equilibria

 

 

 

 

 

 

 

 

 

 

 

 

179
Mn2+ + L3‘ : ML (1)
Mn2+ + HL2_ : MHL . (2)

— 2‘- 4. i i—
where L3 and HL represent tne deprOtonated and mono-

protonated forms of PAA. The ligand is a triprotic acid

 

and under oes the following equilibria
g c -

 

-+ -- -—
H3L Z A + H2L (3)
_ + —
H2L r H + HL2 (4)
_ _ 3-
HL2 2 H+ + L“ (5)

 

The thermodynamic values for the negative log of the acidity
constants are: pKB = 8.69iO-O5, ng = 5.11i0.04, 9K1 =
2.0 (158). The values at a given ionic strength were cal—
culated using the Guntelburg (128) equation. Clearly, the

a . n + I o o 4. e
complexation o1 Mng ion with PAA is pH dependent, with

equilibrium (1) dominating at high pH and equilibrium (2)

Hi

dominating in a pH range of 3.5 — 4.5. The contribution o
equilibrium (2) to the overall complexation is small (<1t)
at the experimental pH of m8. From the pH and manganese
ESR intensity data (which was used to calculate the free

t

2+ . L. v ''r d p 1... ,
Mn concentration) at pH 0: NB, the lormation conStant

for equilibrium (1) was calculated. Equilibrium (2, was

 

 

 

 

180

then studied at a pH of 4.0, where equilibrium (1) had a
small but still significant contribution. The formation
constant was calculated for equilibrium (2) from pH and
ESR data, correcting for equilibrium (1) using the pre—
viously determined formation constant.

The constant for the deprotonated complex can then be
refined by incorporating the value for the monoprotonated

l gand. The cycle is iterative, although convergence is

H-

attained in one or two iterations, since the monoprotonated
equilibria is almost negligible at higher pH's. There is
no evidence for other complexation equilibria occurring in

solution.

The concentration formation constant for equilibrium

x = [MnL’l/[Mng+lIL3‘l (6)

cd
while the relationship between the concentration formation
constant, Kc, and the thermodynamic formation constant,

Kt’ is
th = chYML‘/YMn2+YL3‘ (7)

b

the Debye-Huckel equation

C
U)
1"
:3

09

7.1
O
09
.<
II
I
p
N
l\
in
\
13
+
111
I {D
l—l
I}

 

 

 

 

 

181

where

6

A = 1.823 x 10 /(eT)3/2 . (9)

and
—l/2

B = 50.3(8T)

Equation (7) may be rewritten as

_ 2 2 2 —
1°g ch ‘ l°g th + (ZMnL‘ ' ZMn2+ ' ZPAA3')(I/(l+B§/T)
(10)
or
log ch = log th _ 12A/I/(1+Bg/I) (11)

where at 25° in aqueous solution, B = 0.328 x 107 M71,

A = 0-510, and g‘is an adjustable parameter which is con—
sidered to be the solvated radius of the metal ion (190,
191). Similarly, the thermodynamic formation constant
for equilibrium (2) is,

log Kcm = log Ktm = 8A/I/(l+Bg/I) (12)
The Debye—Huckel equation is generally considered to be

valid to approximately 0.1 M ionic strength. which led to

 

 

 

 

182

the use of ionic strengths equal to or lower than this value.

The presence of other ions in solution, e.g., Mn2+, MnL-,

L3_, contributed to the ionic strength to the extent of

about a 7% error at I = 0.03. During the course of a

single experiment, the change in ionic strength with a change
in ligand concentration was approximately 3%. The former
may be corrected for, while the latter is inherent to the
method. These considerations set a lower limit of 0.03

for the ionic strength.

The formation constants for the manganese—PAA complex
were determined at seven different ionic strengths. The
data are listed in Table XXXIV while the plot of log K as
a function of the ionic strength is shown in Figure 28.

The theoretical slopes for the deprotonated and monopro—
tonated complexes at 25°C are —6.11 and —4.07, respectively.
The data were fitted to Equations (11) and (12) using these
slopes, with Kt and g as the unknowns. This resulted in
the average _ of 6i2 A for the complexes. While the error
is large, this value for g is comparable to the values of

6 ii estimated by Kielland (191) and 4.7 A determined by
Stokes (190). The log Kt values for the mono— and de—
protonated complexes were found to be 3.5i0.2 and 6.3:0.l,
reSpectively.

The errors of 0.1 and 0.2 pK units are larger than
those normally attained from potentiometric measurements,

although they are on the order of the experimental error

 

 

 

183

Formation Constants of Manganese(II) Ion-PAA

Table XXXIV.
Complexes at Different Ionic Strengths and

 

 

 

25°C.

Ionic Strength log KgL log KMHL
0.03 5.5i0.1 3.11i0.08
0.04 5.39i0.07 3.04:0.06
0.05 5.25:0.06 3.0iO.1
0.06 5.0i0.1 2.79:0.08
0.07 5.00i0.07 29:0.1
0.08 4.99:0.08 3.07i0.06
0.09 --------- 2.80:0.08

0.10 4.96:0.05 __________

 

 

 

 

 

 

 

 

poomCOpogdocoe i o>szo Lozoa a

<<a pomeOpopoop
oHCOH one no coflpocsw o no <<aicz

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. <
i o>a50 Lodasv com A <a

. :cspm
Log aoH one no poHQ < .mm oedema

A kdm + F o \ t.
3&6 oouo ow to
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ONTO
_ _ . _ _ ”

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184

Ll.an

 

9» 901

ll.OQv

ll.OQn

 

L100.0

 

 

 

 

 

 

 

 

 

185

associated with cyclic voltammetric determinations. The
technique of manganese electron spin resonance for the
determination of formation constants has been shown to

be accurate, relatively precise, and suitable for these

studies when compared to other techniques. This method
should prove to be a powerful tool in the study of both
aqueous and nonaqueous solution chemistry, especially as

newer instrumentation with greater computer capabilities

further enhances the precision of measurements.

3.1.2. Cyclic Voltammetry

Cyclic voltammetry was used to study the complexation

of PAA with Zn2+, Cu2+, N12+, Co2+, Fe2+, Mn2+, Pb2+, and

Tl+ ions. The above transition metal ions have similar
sizes, of t0.8 A radius (192) while the Pb2+ ion and Tl+
ion have radii of 1.32 A and 1.49 A, respectively. The
sizes of these ions are intermediate between that of mag—
nesium(II) (0.78 A radius) and calcium(II) (1.06 A radius).
with the exception of lead and thallium, which are larger.
If size alone were the criterium for predicting the extent
of complexation with PAA, the transition metals would have
a formation constant value between those obtained for Mg2+
and Ca2+, whereas lead and thallium would have a much lower
Ks. It was also possible that, if the formation constant of

a Tl-PAA complex was large enough, competition studies would

lead to the determination of a formation constant with other

 

 

 

 

 

 

 

 

 

 

186

metals (as has been done with the thallium-cryptate com-

plexes by Weaver gt gt. (193)). A value obtained for the

Mn—PAA system would also provide a cross—check with the
electron spin resonance determinations in Section 3.1.1.

+ . .
Unfortunately, N12 , Co2+, and Mn2+ exhibit irrever-

sible behavior at the hanging mercury drop electrode (HMDE).
That is, no anodic wave was observed after cathodic reduc-

tion. In the case of iron(II), quasi—reversible behavior

D

was observed, which became irreversible upon addition 01

the ligand Quasi—reversible in this sense means that an

anodic peak was observed, but the separation between the

cathodic and anodic peaks was much greaterlflnnlthe theo-
retical 29.5 mV. Zinc was found to be reversible, but the
electron transfer kinetics became irreversible upon the ad-
dition of PAA. Hence, formation constants were only de-

1 . + 2+
termined for Cu2 , Pb , and

rd

1*.
The well-known Lingane (194) equation was used to fit

the data for strong complexes;

C «’7 — - e-

E _ at = — RT/nf (8n xi + p 2nLrAAl) (13)

1/2 1 = :

where EC and Ef are the half-wave potentials of the
1/2 1/2 ‘

free and complexed metal ion, Kf is the formation constant

of the complex, p is the stoichiometric coefficient of PAA
and [PAA] is the total concentration of PAA. The other

symbols have their usual meaning. Where weak complexation

 

 

 

 

 

187

was found, the modification of this equation formulated by

Buck (195) was used;

Ei/2 = Ei/g = —RT/nF(£an+p£n[PAA] + tn(l+l/xf[PAAl) (14)
with the symbolism the same as in Equation (13). The Buck
equation differs from the Lingane equation in the last
term on the right, which is negligible when Kf is large.
The use of the Lingane and Buck equations requires
that several experimental conditions be met. The elec-
trode kinetics must be reversible, which is experimentally

difference in the anodic and cathodic peaks of

‘

seen as a
29.5 mV for a two electron transfer and 59 mV for a one
electron transfer reaction. The thallium-PAA system ex—
hibited a peak separation of 59:2 mV, while the lead and
copper systems had a separation of 30i4 mV, indicating

L

reversible behavior. The ligand—to—metal concentration

W

ratio must be such that the concentration of ree metal
ion is small in comparison to total metal ion concentra—
tion. The diffusion coefficients for the free and com—

plexed ions should be approximately equal and the transi-

tion state for the electron transfer should be ap-roximately

half—way between the two oxidation states. The former

0

y be checked ior

{1)

condition is assumed, while the latter m
gross errors by observing the symmetry of the anodic and

cathodic peaks (symmetry indicates that the condition is

 

 

 

 

 

188

fulfilled).

2+

2+
, , and a

The data for PAA complexes with Pb Cu
sample graph for Tl+ are plotted in Figures 29-31 and tabu—
lated in Tables XXXV - XXXVII. The slopes of the lines are
li0.3, indicative of a stoichiometric ratio of one ligand

to one metal ion for each of the complexes. The formation

constants derived from such a treatment are: Log KgUI =

.L P .
8.0:0.l, log xguq = 4.2:01, log KEbL = 7.0:1, log xfde
= 4.1iO.l, and log KglL = 2.3:0.l at 25°C and an ionic

strength of 0.05. The monoprotonated thallium—PAA complex
formation constant was too weak to measure using cyclic

voltametry.
The thermodynamic formation constants was determined
for the deprotonated thallium-PAA complex by using the

Debye—Hfickel equation

log KC = log Kt — 6A/I/(l + Bg/I) (15)

Where the symbols have been defined in Section 3.1.1.
The results are given in Table XXXVIII and plotted in
Figure 32. The theoretical slope of —3.05 was attained

when the ion size parameter g was 9.3, much larger than

the value of 2.5 calculated by Kielland (l9l)- This is

- 1 i‘n ' . . ovalert
probably due to a large contribution from the c .

character of thallium to the Complexation With PAA, Since

 

 

189

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. 02s: 00.. ..
03. 8... one 86 03 88
willlfielllfllillnllllnllI03
1.. 00a
.
1.. ll 00.».
3
J- 35:
15 (
1.. /
ll 00* .0
.i 0
7w
1.. 6
1+ l x
l. 80
Li

 

006

 

 

 

190

A

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ii UmpfiCOpopgmw I 95:0 .8915 xwaquo 53:59 m go pcwpmcoo COprELA¥

93 «525.966 on 3.3: moa 9306mm: map a (HEM CH mwcwCQ 93 mo uOHQ <

 

 

.om mhswflm

3&8..-
00* 0Q* 006 00% CON CON
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{8a
II 356“
_
11033
i 2...:
1| (
.3 /
llooé.0
1: 0
Z
-- m
+ 1186
I“ II
a! 090

191

 

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map @CHEhmme ow h<<mg moa m>flpwwmc map m> mEm :fl wmcmgo map mo poam < .Hm wgzmfih

 

7)}: 004 I
on». 8.» 93 8a on;
—|L _ — _ — _ . _ _ — — _ . _ _ — . _ _ OP‘O
— _ _ _ — — . q _ _ — . _ _ H _ a _ _ —
In one
: any
.
Ln -1 3
II and 3.3
-1 .0
11 08 m
w
9
II 8.0

 

 

192

Table XXXV. The Change in the Emf of a Pb—PAA Complex as
a Function of the Negative log [PAA] at 25°C,
I = 0.0

 

 

 

 

 

 

 

 

ML C F MHL

(Bi/TEi/E) <E1/2‘Ei/2
0.02958 -1og [PAA] 0.02958 —1og [PAA]
u.63 2.91 1.83 3.09
u.u0 2.59 1.81 3.19
3.79 3.11 1.76 3.31
3.58 3.21 1.67 3.54
3.52 3-3“ L”? 3.95
3.31 3.52
2.99 3.93

39.2223: :22 :2: 12121.: ig:1§.:32§3.22,

1 = 0.05.
‘ MHL

(EC -31" > ML (Bi/2132.)

“‘%¥%§9%8g— —1og [PAA] "0702958_‘ —log [PAA]
5.73 2.97 1.91 3.17
5-53 2.57 1.89 3.23
5.91 2.69 1.79 3.36
5.38 2.73 1.69 3.28
5.27 2.71 1.30 3...
5.14 2.86
9.911 2.91
4.83 3.10

 

 

 

 

 

 

 

 

 

 

 

Table XXXVII.

The Change in
a Function of

the Emf of a
the Negative log [PAA] at 25°C,

Tl—PAA Complex as

 

 

 

1 = 0.05.

(E1/2‘31/2) ML (fin—41m) ML

"“UTU§§TE_‘ ~log [PAA] __6765916—_ —log [PAA]
0.20 3.00 0.55 2.u0
0.23 2.92 0.66 2.30
0.25 2.85 0.70 2.26
0.27 2.80 0.79 2.22
0.29 2.74 0.78 2.19
0.31 2.70 0.82 2.15
0.35 2.62 0.86 2.12
0.39 2.55 0.90 2.10
0.97 2.99

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Table XXXVIII.

Formation
Complexes

25°C.

199

Constants of Thallium(1)

- PAA

at Different Ionic Strengths and

 

 

Ionic Strength

0.

O.

030

040

.050
.060
.070
.090

.110

(\J
l\)
\51

l\)
I\)

ML

l+

H-

H-

|+

H

|+

H-

0.

.07
.07
.09

}._l

.08

 

 

 

 

 

195

 

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m we onQEoo HBI<<m oomeOpopooo @o pcwpmcoo QOflmeLom moa @o pOHQ <

Abdm + F o \ I;
ONTO 8—30
F . — _ —

_ . . . _ n l n u l l 1

 

 

oo fio owod

.mm mgswfle
8.0
oo. _
1
8a w
8...”

00+

 

l

 

 

196

the Debye—Huckel equation considers only electrostatic

interactions. The thermodynamic formation constant for the

deprotonated complex was 2.63:0.09. The low stability of

the deprotonated PAA-thalliumCI) complex does not lend it—

which were subsequently aban—

F")

sel to competition studies,

doned.

 

The value obtained for the formation constant of the

deprotonated and monoprotonated Cu-PAA complexes agree
fairly well with the values obtained by Stunzi and Perrin

(log KEUL = 7.19 and log KEUHD = 3.85) when considering
that they used an ionic strength of 0.15 and made the meas—

urements at 37°C (157)-

3.l.3. Potentiometric Studies

 

 

A silver specific ion electrode was used to determine

+ .
the formation constants of Ag ion complex With the mono—

he calibration curve o

|—]

and deprotonated forms of PAA.

ration curve for the de-

(‘1‘

the electrode, as well as the ti

protonated complex, is shown in Figure 33. As can be seen,

#4)

the calibration curve is relatively linear with a slope o
63.310.6 mV, close to the ideal Nernstian lepe of 5

mV. The formation constants determined were: Log 1

= 3.11i0.05 and log KfigHL = 2.7:O.l at 25°C and an ionic

strength of 0.05.

The formation constants of the calcium ion with FAA

were measured using a calcium ion selective electrode.

 

400-0

3750

3500

3250

EMF IN MILUVOLTS

2750

2500

3800

3750

3700

3850

3600

EMF IN MILUVOLTS

e 33.

197

A CALIBRATION CURVE OF A
SILVER ION SPECIFIC ELECTRODE

 

 

 

 

“I
—I
._
::
“ I . I 1

fi— I T j 1" fifi
~50 ~40 -3.0 -2-0

LOGIAg‘J
A TITRATION CURVE 0? Ag’ wrm
PAA AT 25°C AND I - 0.05

«L.

.11.1.1.111.I.L.I.111.Ivlhlnflllj
‘T'I {TIT—FTITW‘ITI‘I'I'l—‘I I’Tti
0.00 4.00 8.00 12.00 woo
umuunnscr’wu

A sample calibration plot and titration curve

of silver ion with deprotonated PAA using a
silver ion specific electrode.

 

 

198

A typical calibration curve (25°C, I = 0.05) is shown in
Figure 39 along with the titration curve. The data were
processed assuming the formation of a mono- and depro-
tonated complex. There is no evidence for the formation
of other complexes in solution. Formation constants were
determined at different ionic strengths and extrapolated
to zero ionic strength to determine the thermodynamic
formation constants using Equations (11) and (12). The
formation constant and ion size parameter a were used as

the fitted parameters. The resulting plots for the mono—

and deprotonated complexes of PAA with calcium are shown
in Figure 35 using the values given in Table XLIX. The

logarithm of the thermodynamic formation constant of de-
protonated PAA—Ca is 9.68:0.03, with a value of 2.68:0.08
for the monoprotonated formation constant. The ion size
parameter was found to be 5.2 E, which compares to the

Stokes determination of approximately 5 3 (190) and that

of Kielland 6 fi (191).
The ionic strength was extended to 0.14 to more pre—
cisely compare the values obtained by this technique to

those of Stunzi (157). The graph shows that linear be—

havior is exhibited even to this ionic strength when the

Debye—Huckel equation is used. Thus, when enough data

points are available to calculate an ion-size parameter,

this equation is the preferred description of solution

complexation behavior.

 

I,
I

 

 

 

 

EMF IN MILLNOLTS

EMF IN MILUVOLTS

gure 3A.

199

A CALIBRATION CURVE FOR A
CALCIUM ION SELECTIVE ELECTRODE

100
O0
-100

-200

 

LOG tec*21

A THRAHON CURVE OF CALOUM WWH
PAA AT 25°C AND 1 -001

00

1 I . 1 l 1 l . J
| ‘ I ' I ,. I I I ‘ I
80 12.0
MILULITERS OF PAA ADDED

-, 1.1.1
5°"I'I

titration curve of

A sample calibration plot and
PAA using a cal—

calcium ion with deprotonated
cium ion selective electrode.

 

 

 

 

200

.A<<m
oonCOQOoncoE 1 0>L50 pesoa .<<g oomeOpOLQmo 1 o>pzo pogozv Camcoppm

OHCOH esp mo coepocsm m mm <<m1mo mo pcmpmcoo coemeaom woa mo poaa < .mm opsmflh

Csdm+:\.§

83 83 83 83 . 3 _.o 8 to 08.0
TITiIlITTITIITIiITITITITITITIIITII 8.—

II.Oo+

gl 090

 

501901

 

 

 

 

 

 

 

 

 

 

 

 

 

 

201

Table XXXIX. Formation Constants for Calcium(II) Ion—PAA
Complexes as a Function of the Ionic Strength

 

 

 

at 25°C.
Ionic Strength log KgL log KIf/IHL
0-02 ——————————— 2.21 1 0.07
0.03 3.83 1 0.02 2.16 1 0.03
0.04 3.73 1 0.02 2.08 1 0.09
0.05 3.67 1 0.02 2.10 1 0.06
0.06 3.57 1 0.01 2.07 1 0.06
0.08 3.u6 1 0.01 1.96 1 0.05
0 10 3 33 1 0.01 1 89 1 0 05
0 12 3.28 1 0 01 -----------

 

 

 

 

 

 

 

202

3.2. Thermodynamics of Calcium Complexation with PAA
____________l_________________l___1_______1______

The thermodynamic parameters AH0 and A80 were determined

from the relationship
—RTtnKt = AGO = AHO - TASO (16)
which can be rearranged to yield the van't Hoff relation
znxt = —AH°/RT + ASO/R (17)

Hence, a plot of AnKt XE 1/T will have a slope of —AH°/R
and an intercept of ASO/R.

The formation constants Kc were determined at an ionic
strength of 0.05, with the Kt then calculated using the
Debye—Hdckel equation with an ion size parameter of 5.2 A.
As can be seen in Equations (8) and (9), the Debye—Hfickel
equation itself has the temperature dependent quantities
A and B. Both are a function of the dielectric constant
(a temperature dependent quantity) and the temperature.
Values for these quantities were obtained at a given tem-
perature by interpolating values given by Earned and Owen
(196). The volume of the solution will also change with
temperature (186), although the effect of such changes on
the concentrations involved was negligible.

The AH0 and 03° of complexation were not obtained for

t ‘ I experimental
the monoprotonated complex, due to a large .

203

error. Generally, the error increases as weaker complexes
are studied, because the relative error increases as the
change in emf between two titration points decreases. The
vs l/T for the calcium-PAA complex and the
calcium—3-PPA complex are shown in Figure 36. The com—
puted thermodynamic quantities are given in Table XL along
with values for the complexation of Mg2+ ion with PAA.
The formation constants for Ca2+ and Mg2+ complexes with
pyrophosphate are also shown for comparison in Table XL.

Complexes of calcium and magnesium ion with PAA and
pyrophosphate anion are seen to be entropy stabilized and
enthalpy destabilized. This behavior is attributable to the
”chelate" effect, which will partially result from the
release of solvent molecules upon complexation leading to
an increase in the entropy of the system. The ligand PAA
is more selective with respect to magnesium XE calcium
complexation as the magnesium formation constant is ap—
proximately an order of magnitude greater than that of the
calcium ion. This is different from the pyrophosphate case
in which calcium and magnesium show approximately
equivalent formation constants.

The effect of pyrophosphate on enzyme kinetics may
be two—fold (200). The primary effect is an enhancement
of activity via magnesium complexation (107): while a
secondary effect is an inhibition due to calcium complexa—

tion (198). Hence, FAA would effect enzyme kinetics

 

 

 

 

 

 

204

.Ao>Lso gmsofiv <mm1m ocm Ao>Lzo

Leggsv <<m pwpmcouomgmp nus: onQEoo +mmo ommeOposdoo m mo zoflwcucm

pcm moospco one ocflELopop on HImLSmemoEop x coca MM ex moa ho pOHQ < .wm mpswfle
P1 h x OOOF
on.» on“; on“ or.»
TIIIITIIIITIIII+IIII+IIII+IIII+IIII+IIII+IIII+IIII1IIII1IIIIII an
# 11 86
1;! I
-- m
11 oos 9
-- .wfl

 

gl 00.:

 

I

 

 

 

205

 

 

.Awsmfiv mmopm ESCmHm .xsow 302

.mecmaaq oscmmsocH .1 .Ho> .wpcmpmcoo spHHAQMpm Hmoapaso .Hamptmz .m .< .cpasm .2 .mo

A

 

geomav ..epq smaooo a museum: .coecoq .mpcmpmcoo psaapmpm

astute: .m .< .cmaaaw .e .nn

.Amsmav .spamsm>flc: mpmpm :swazoaz .mammss .Q.:m .prsmm .o .m1.m

 

 

 

 

 

 

 

 

 

 

 

 

o m: m m:.m1

s 11 0.1 m.s-

n ma mm s.o m s.onm.s1

xsoz mass ma 2H m.owH.H m.oH ©.:1

xmos mesa s.oflz.fim m.onw.o so.onwm.w1
om< 03< 00<

 

 

 

 

 

 

w
. s m . m
m: m I: o a +N 2
1.m scam. so
1: +m
- I o 0 w
mo o+mm m <1; +m z
mo.oamm.m <mm1m. mo
+m
mo.owwe.1 <<1. so
+m
ex mom

uHHwnHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHHH

.<<m Sufi: :oepmmeQEoo COH Hmpoz pom mmflpflpsmsa owEdChooELmse

.qx mflnt

 

206

more dramatically than pyrophosphate, for, although the
magnesium complexation is approximately the same, the cal—
cium complexation is weaker. This leads to an overall en—
hancement of enzyme kinetics of PAA vs pyrophosphate. The
analogy is extended to the in yiyg PAA case of replication
inhibition, with PAA competing as effectively, and perhaps
better, than pyrophosphate for the binding site in the
scheme proposed in Section 1.2.2.

The extent of complexation of PAA with calcium would
suggest that an in give calcium complex with this drug is

a possible explanation for the observed side effects.

3.3. Summary and Discussion

The complexation formation constants for some metal
ions with PAA at 25°C and an ionic strength of 0.05 are
given in Table XLI, along with the A80 of solvation for
each ion (196). The plot of the entropy of solvation XE
the log K¥L is shown in Figure 37. There appears to be
two linear relationships between the solvation entropy and
log K¥L. For an entropy of solvation greater than ten
e.u., a linear relationship is seen with the formation
constant, although there are not enough points to reach a
conclusion. Another linear relationship is seen for the
remaining metal ions (correlation coefficient of 0.83)
between the entropy of solvation and the formation constant.

If it is considered that the calcium and magne31um

 

 

 

 

 

207

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Table XLI. The Solvation Entropy, Radius, and Formation

Constants of Some Metal Ions With PAA at 25°C

and an Ionic Strength of 0.05.

R°(A) ASgolv. ML MHL

(66) in e.u. log Kf log Kf Source
002+ 0.82 —14.5 8.0:0.1 4 2 10.1 This Work
Zn2+ 0.83 —11 9 5 3510.03 3.4410 03 (31)
Mg2+ 0.78 -11.7 4 5010.02 2 5610 05 (30)
Mn2+ 0.80 — 9.5 5.2510.06 3 0 10.1 This work
Cd2+ 1.03 - 2.6 3.9 +0.1 (30)
Ca2+ 1.06 1.8 3.6710.02 2.1010 06 This work
Sr2+ 1.27 3.4 3 6710.02 2.5610 06 (30)
Ag+ 1.13 3.7 3.1110.05 2 7 10 1 This work
Na.+0 0.98 5.1 1.4310.02 0 7910.05 (30)
Ba2+ 1.43 14.1 3.6710.02 2.5010 06 (30)
T1+ 1.49 14.6 2.5110 03 ————————— This work
9132+ 1.32 15.2 7.0 10.1 4 1 10.1 This work
a1 = 0.15, 37°C.
bl: = 0.4.
CI = 0.078.

 

208

.5; oopmcouosoop Spas moonQEoo com 139: mEom cHo . .
pcmpmcono conumciom woa one. a 75.3 soapy/How Lo aaoppco one no moan < .sm oezmfla

.D.u Z. ZO_._.<>1.Om no >Qomhzw NIH
oodN oo.N F 004 0061

 

8&7 8.81

IIIIlIITlIlIlII- s...
11 8.~

-1 .l

O

N+om 9

II OO..v N

O

-- w.

11 88 .

.0

0

N+QQ 11 C.
m+so H 11 88
I¥I 00.5:

 

 

209

complexes are entropy stabilized, it may be generalized
that these other cations might also exhibit an entropy
stabilized complexation. (In the case of calcium and mag—
nesium complexation, the enthalpy of complexation detracts
less than 10% from the overall free energy of complexation.
It is expected that the enthalpy will also play a minor
role in all complexes with PAA.) This leads to the
conclusion that stabilization of the complex is derived
from the release of solvent molecules (thereby increasing
the entropy) upon complexation. This phenomena is normally

associated with the "chelate effect” (201).

3.4. Phosphonoacetic Acid Analogs

Microbiological studies (150,199) of phosphonoformic
acid (PFA) and 3—phosphonopropionic acid (3—PPA) indicate
that the former inhibits the replication of some Herpes
viruses while the latter does not. Complexation studies
of PAA, PFA, and 3-PPA have shown that the biological in-
hibition trends of these ligands are paralleled by the
extent of complexation with the magnesium(II) ion (158).
To draw further analogies between the biological effects
and complexation, the complexing ability of these drugs
were tested with manganese(II), thallium(I), and calcium-

(II) ions.

3.4.1. Manganese Electron Spin Resonance

Manganese electron spin resonance was used to study the
complexation of manganese with PFA and 3-PPA at an ionic
strength of 0.05. Both PFA and 3—FPA are triprotic acids
which will protonate according to equilibria (3), (A) and
(5). The protonation constants were determined at I =
0.05 by Heubel (158) and are given in Table XLV, along with
those of PAA at I = 0.05. The formation constants for
complexes of manganese with these ligands were determined

according to the procedures outlined in Sections 2.3.1 and

1:1
3.1 1. The resultant values were: log KgnP‘A = 5.3410.05,
log KEHHPFA 2.5710.05, log xgn‘3‘PPA = 3.1510.04, and
J.
log KMHH‘3‘PAA = 1.610.2.

f

3.4.2. Cyclic Voltammetry

The pKa’s of PFA and 3—PPA were calculated at a given
ionic strength from the appropriate Guntelburg equation
using the data of Heubel (158). Formation constants were

then determined for a deprotonated thallium—PFA complex

at different ionic strengths. There was no evidence of

thallium complexation with 3—PPA or the protonated form

then used to calculate the thermo-

of PFA. Equation (14) was

dynamic formation constant from a linear extrapolation to

- - ' ' n' . T * lues of
zero ionic strength, shown in rigure 38. -he va

each point are given in Table XLII. The ion size parameter

211

 

.meCohpm QflCOfl one mo COApocsg m we

onQEoo fi31<mm oomeOQOLQoU n no pcwumgoo coepwesom woa oso mo pOHQ < .wm osswfim

A.§.mm + Fv \ ms

ONNd ow v.0 9v _..0 co to

I
I
O
O
N
O>1 001

II 09M.

 

 

 

 

212

Table XLII. Formation Constants of Thallium (I)—PFA Com-
plexes at Different Ionic Strengths and 25°C.

 

 

 

Ionic Strength log xii”?
0.030 2 3 1 0 1
0.040 2 27 1 0 09
O. 50 2 2 1 O 1
0.062 2 1 1 0 1
0.070 2 1 1 0 1
0.092 2 1 1 0 1
0.131 2 0 1 0 2

 

 

 

213

was found to be 9.0, again in poor agreement with other
calculated values (190,191). The thermodynamic formation

constant is 2.65i0.07.

3.4.3. Potentiometry

Thermodynamic formation constants for Ca-PFA and Ca—3—
PPA were calculated from potentiometric data using a
calcium ion selective electrode. The formation constants
were determined at different ionic strengths as in Section
3.1.3, then extra_olated to zero ionic strength using Equa-
tions (11) and (12). The data are plotted in Figures
39 and 40, with the values shown in Tables XLIII and XLIV.
The resulting ion-size parameter of 5il A compares well
with the values of Stokes (190) and Kielland (191), as well
as with the value obtained from studies of the calcium—PAA
system. The significance of the a parameter is seen in
the self-consistency of its determination for a series of
analogous ligands (PAA, PFA, and 3—PPA) and the agreement
in values obtained by very different methods (190,191).
While the interpretation of this parameter as the solvated
radius may be debatable, these and other studies indicate

that it does have a physical significance. The thermo—

log KiiA = 4.4010 03, log

dynamic formation constants are:
K3—PPA _

PFA ,3—PPA = 3910.09, and log q
KMHL — 2.610.1, log hML 3 MhL

2.3:0.1.

 

 

 

214

.A<<m oopGCOposo
1ocoE 1 o>s50 LoSOH A<hm popmcoQopdoo 1 o>LSo Loomsv newcogpm OflCOH

oz» mo cowpocze m we <mm1mo mo peemeOO coemesoe EOH one mo poad <

“L5.mm + F v \ Ls

OOTO ONTO omOO

O+WO OOWO

  

 

.mm oezmflm

OQF

OQN

OQn

00*

OQn

°>1 001

 

 

 

 

215

.A<<m poemsouosg
locoE 1 o>L:o Losofi A<<m popmcouopdop 1 o>s30 Locasv Qumconpm oHCOH

one go coepocsm o no <mm1mlmo no pampmcoo :oemeLom moa one go pOHQ < .o: oszmfle

Atwm + : \1;
3N6 8N6

00 «.0 ON —.O 0906
TITITIITIIITIITIITITITIII1 8. .

3x001

 

I‘I 001v

 

 

 

 

 

216

Table XLIII. Formation Constants of Calcium(II)-PFA Com-
plexes as a Function of the Ionic Strength

 

 

 

 

 

 

at 25°C.

Ionic Strength log KfL log KIf/IHL
0-02 ----------- 2.09 1 0.02
0.03 3.69 1 0.06 1.97 1 0.03
0.04 ——————————— 1.90 1 0.02
0.05 3.55 1 0.05 1.84 1 0.04
0.06 3.47 1 0.04 ___________
0.07 3 42 1 0.05 1.72 1 0.06
0.08 3.37 i 0.06 1.78 1 0.02
0.10 3.28 1 0.03 1.54 1 0.06
0.12 3.24 1 0.03 -----------
0.14 3.18 1 0.03 -----------

 

 

 

 

 

 

 

 

 

 

 

217
Table XLIV. Formation Constants of Calcium(II)-3-PPA
Complexes as a Function of the Ionic Strength
at 25°C.
s§§:§§1h log K¥L log KMHL
0.02 ——————————— 1.81 i 0.09
0.03 2.60 1 0.02 ' 1.79 1 0.02
0.035 2.46 1 0.02 1.74 1 0.04
0.04 2.38 1 0.02 1.75 1 0.07
0.05 2.38 1 0.02 1.64 1 0.07
0.06 2.28 1 0.02 1.57 1 0.02
0.07 2.19 1 0.02 1.61 1 0.08
0 08 2.16 1 0 07 -----------
0.10 2.06 1 0.02 -----------

 

 

 

 

 

 

218

3.5. Thermodynamic Studies of 3—PPA Complexes

 

 

 

 

Since 3—PPA plays no role in inhibition of Herpes
virus and exhibits a weak complex with calcium (as com—
pared to FAA and PFA), it was of interest to contrast the
enthalpy and entropy of complexation with that of PAA.

The formation constant of a deprotonated calcium-3—PPA
complex was determined for different temperatures at an
ionic strength of 0.05. Again, due to the larger error,
the mono-protonated complex was not studied. Using the
Debye—Hdckel equation with an g of 5 A, the thermodynamic
formation constant was then calculated at a given tempera-
ture. Equation (17) was then used to calculate AH0 and
08° of complexation, shown in Table XLV.

Similarly to the PAA case, the calcium—3—PPA complex
is entropy stabilized and enthalpy destabilized. While
the AH0 values for FAA and 3—PPA calcium complexes are
approximately equal, the 08° of complexation is much lower
in the case of 3-PPA. The stability of these complexes
with calcium is then a function of the entropy of complexa—

tion, with the enthalpy contributing only a small de—

stabilizing effect. Hence, the chelate effect is likely

the dominant mechanism in most complexes with PAA, PFA,

and 3—PFA.

 

 

 

 

 

219

3.6. Summary and Discussion

The formation constants of some metal ions with PAA,
PFA and 3-PPA are shown in Table XLV at 25°C and 1 = 0 05.
The complexing ability of FAA and PFA are approximately
equivalent for the deprotonated complexes, which will pre-
dominate at a biological pH of 7. -The formation constants
of FAA and PFA complexes with calcium and manganese ions
agree well with each other, while that for the magnesium—
PFA is closer to PAA than 3—PPA. Since the biological ef—
fects of FAA and PFA are also approximately equivalent, it
seems that there may be a correlation between the biological
effects of these drugs and their complexing ability. This
conclusion is supported by the data from complexation studies
of 3-PPA. This ligand has no biological inhibitory activity
against Herpes Virus and shows a complexing ability that is
at least an order of magnitude less than that for PAA or

PFA when considering a deprotonated complex.

3.7. Conclusion

. + + 2+
Complexation studies of PAA w1th Mn , T1 , Cu ,

Pb2+ Ag+, and Ca

2+, together with data from other sources
(51,157) seem to indicate that there is a general rela—
tionship between the entropy of solvation and the strength

t ' f t.
of complexation. Temperature dependence Studies 0. «he

Calcium—PAA and 3-FPA systems have_shown that complexation

 

 

 

 

 

 

220

Table XLV. Formation Constants for Some Metal Ions with PAA,
PFA, and 3-PPA at 25°C and I = 0.05.

 

 

PFA PAA 3-PPA

 

 

 

 

log K¥gL <30) 3 5910.05 4.50:0.05 2.28:0.05
log K¥gHL (30) 1.7 10.3 2.6 10.1 1.7 10.1

log KgaL 3 5510 05 3 6710 02 2 3810 02
log KgaHL 1.8410.04 2.1010.06 1.6410.07
log KgnL 5.3410.05 5.2510.06 3.1510.04
log KMnHL 2.5710.05 2.9710.06 1.6 10.2

log KElL 2.6310 07 2.6210.04 ________

0K1 1.7 :0.1 2.18:0.04 2.2610.04
pK2 3.5910.02 4.9410.02 4.6310.02
pK3 7.5610.02 8.14:0.02 7.75:0.02

 

 

 

 

 

 

221

is entropy stabilized and enthalpy destabilized, with the
entropy term dominating. These studies, together with
similar studies using magnesium, have shown that entropy
stabilization, or the chelate effect, is the most likely
mechanism which explains complex stability. Studies on PAA,
PFA, and 3—PPA led to the conclusion that there may be a
correlation between the complexing ability of these ligands

and their biological activity.

 

 

 

 

 

 

 

 

 

APPENDICES

 

 

 

 

 

APPENDIX A

THE GENERAL USE OF THE NON-LINEAR CURVE

FITTING ROUTINE KINFIT4

.. .{N-J—

 

 

 

 

 

 

 

 

THE GENERAL USE OF THE NON-LINEAR CURVE
FITTING ROUTINE KINFIT4 ‘

A.l. Scope of KINFIT4 Usage

The program KINFIT4 was used to curve fit the data for
all linear relationships (Lingane equation, calibration
data, van't Hoff plots, thermodynamic Kf plots, etc.) as
well as some non—linear relationships (some calibration
curves, the Buck equation, etc.). The experimentally
determined quantities were expressed in such a manner
that one quantity was always expressed as a function of

the other, e.g.,
y = f(x) (1)

with the function f(x) containing the unknown parameters
to be determined by a computer fit. The resultant cal—
culation of the unknowns was accompanied by such statis—
tical data as the standard deviation of a given calculated
value, the coupling (correlation) between the calculated

values, and a plot of the actual data XE that calculated

from the computed unknowns.

A.2. Data Input

For a more complete summary of the data input instruc-

tions, see Reference (204)-

 

 

 

 

A.2.1. Equation Card / 6X,74A

This card describes the relationship between experi—

mentally measured variables and is of the form
CALC = f(XX(1)) (2)

A.2.2. Residual Calculation Card / 6X,74A

This card is of the form
RESID = CALC = xx<2> (3)
where XX(2) is the dependent variable from Equation (2).

A.2.3. Control Card / 1015 NOPT, IMETH, ITMAX, IWT,
IRX, ISMIN, IPLT, NCST,
TEST, KVAR

Unless specified, all values are set at the default

zero. Only those quantities which were not set at the
default will be mentioned.
NOPT is the number of data points.

ITMAX is the maximum number of iterations permitted if
convergence is not attained.

TEST is the convergence tolerance.

A.2.4. Descriptive Title / 704

A.2.5. Constant Card / 8E10.5 CONST(i)

\O

224

Initial Estimate Card / 8F10.5 U(i)

Data / 8E10.5 XX(1),
Blank Card

6
7
89

2(1) ,xx<2),

2<2)

 

 

 

 

 

 

 

APPENDIX B

DATA REDUCTION OF A POTENTIOMETRIC TITRATION FOR
THE DETERMINATION OF COMPLEX FORMATION
CONSTANTS USING THE COMPUTER

PROGRAM MINIQUAD

 

 

 

 

 

 

DATA REDUCTION OF A POTENTIOMETRIC TITRATION FOR THE
DETERMINATION OF COMPLEX FORMATION CONSTANTS USING
THE COMPUTER PROGRAM MINIQUAD

8.1. The Scope of MINIQUAD

The computer program MINIQUAD is a versatile program
which enables the user to determine simultaneously up to

20 unknown formation constants with five reactants from

 

potentiometric data. The program can accommodate up to
three known concentrations (the pH, free metal ion con-
centration, free ligand concentration, etc.) determined

from potentiometric data.

The equations describing the equilibria involved are

 

 

of the form

B
aA + bB + cC Aa bcc (l)

with a formation constant of

K, = [AasbcCl/[AlaEBlbIClC <21

In the specific case of complexation with PAA, all equilib—

ria can be described in terms of the free metal ion concen—
tration, [M], the concentration of the free deprotonated
ligand, [L], and the hydrogen ion concentration, [H], with

the complex concentration calculated by the program. Hence,the
formation constants for equilibria (l) and (2) from Part III,

Section 311.1.

225

 

 

 

 

 

 

226

Kf(ML) = [MLl/IMIIL] (3)
and
Kf(MHL) = [MHLl/[M][H][L] = Kf(MHL)/K3 (4)

where K3 is the third acidity constant for the ligand.

The acidity constants are written as cumulative B's, i.e.,

 

K = [HJLl/[H]2[Ll (5)

2

The computer program will accept the stoichiometric
coefficients of the reactants for all equilibria involved.
Along with initial estimates (or values obtained from other
sources) for the respective formation constants. The user
then has the option of refining the formation constants or
holding them constant. In the PAA study, the acidity
constants were known.(l58% leaving the complex formation

constants to be determined by Equations (3) and (4). Both

 

the pH and the free calcium ion concentration were monitored
potentiometrically, while the total ligand concentration I
was known. Each data point input to the program consisted
of a pH and volume of ligand added.

The program was modifiad(l58) to include the KINFIT4
program described in Appendix A. This program is a general
purpose non-linear curve fitting routine that was used to
calculate the slope, intercept, and residual calcium from

calibration data for the calcium ion selective electrode.

 

III

 

 

 

227

The data were fit to the equation

+2

E = b + M log [(Ca )f + R] _ (6)

obs

where EObs is the observed emf, b is the intercept, m the

slope, (Ca+2)f the free calcium concentration, and R a

 

residual to account for non—linearity in the calibration

curve at low concentrations.

8.2. Data Input Format for MINIQUAD
B.2.l. KINFIT Calibration Input

Control card/815, F10.0, IS/see Appendix

03
I\.)
[—1
H

A, Section
8.2.1.2. Descriptive title/20A4

B.2.l.3. Initial Estimates/8F 10.5

 

This card will contain the initial estimates of the

slope, intercept, and residual.

2

0.2.1.4. Data/8F 10.5/log (Ca+ )f, OCa’ Emf, o 1

emi

The dataare:1nput as the log of the calcium ion con-
centration, its estimated variance, the emf, and its esti—
mated variance. The program internally converts the log

of the concentration to a concentration, which is then used

 

 

 

228

in Equation (6). There are two data points (X, OX, Y, by)

per card.

B.2.1.5. Blank card

B.2.l.6. 789 card

 

3.2.2. MINIQUAD Titration Data input
8.2.2.1. Descriptive title/20A4

B.2.2.2. Control Card/8I5/LARS, NK, N, MAXIT,

IPRIN, NUMBEO, NCO, I COM

LARS indicates the data points to be used in the cal-
culation of an unknown formation constant: LARS = 1 means
all points are used, LARS = 2 every other point, LARS = 3
every third point, etc.

NK is the total number of formation constants input,

 

both refined and constant.

 

N is the number of unknown formation constants to be
calculated. I
MAXIT is the maximum number of iteration cycles to be i
performed.
IPRIN will allow the monitoring of the program at each
iteration. A value of 1 will print the results at each
iteration, 2 will also print each data point and its resi-
duals, while 0 is the default with no monitoring done.

NMBEO is the total number of reactants for a system.

 

 

 

 

229

NCO is the number of free concentrations unknown.

ICOM will allow the elimination of spurious results.
A value of 0 will include all results and a value of 1 will
eliminate a point if the normal equation matrix is not

positive definite at this point.

8.2.2.3. Temperature compensation/3F10.6, 8X,

I2/TEMP, ADDTEMP, ALPHA, NOTAPE

TEMP is the temperature of the bulk solution in °C.
ADDTEMP is the titrant temperature in °C
ALPHA is the coefficient of cubical expansion for the

solvent in 0C71.

NOTAPE interfaces KINFIT4 to MINIQUAD. NOTAPE = 1
reads the m and b values from the KINFIT4 program, while

NOTAPE = 0 will use values given by EZERO and SLOPE.

8.2.2.4. Formation constant cards/F10.6, 7I5/

BETA(I), JPOT(I), JQRO(J,I), KEY(I)

The formation constants are input in the form 8i =
BETA(I) . lOJPOT(I)

The JQRO(J,I) values are the stoichiometric coefficients

A.

of the ibh species with the formation constant Bi' While

.5.

the order is arbitrary, the reactants which are measured

potentiometrically must come last. In the case of a mono-

protonated Ca-PAA complex, the input would be 1 l 1

 

 

 

 

 

230

where the first number is the stoichiometric coefficient

of the ligand L, the second of the measured calcium ion,

and the third of the measured hydrogen ion concentration.
KEY(I) indicates if the value is to be refined (=1)

or if it is a constant (=0).

3.2.2.5. Control Card/l2I5/NMBE, JNMB(I), NC,

JP(I)

NMBE is the number of reactants in the equilibria.

The JNMB(I) values are the integers assigned to those
reactants. The equilibria in Section 0.2.2.4 would have
the reactants labeled

1 2 3

NC is the number of unknown free concentrations at any
given point in the titration. This can be calculated by
NMBE—NCO = NC where NCO is the number of electrodes.

JP contains the reactants (as numbered in JNMB) which
will be processed further by the subroutine STATS. This
subprogram will calculate the relative percentages of those
reactants specified by JP and print them out at each point

in the titration curve.

3.2.2.6. Reactant titles/5A10/REACT(I)

REACT contains the names of the reactants in the order

specified by JQRO(I).

 

 

231

8.2.2.7. Electrode parameters/4I5/JEL(I), JCOUL

JEL is the number of electrons transferred at each
electrode. When pH values are to be read, JEL = 0. The
values are given in the same order (lowest number first)
as in JQRO(I).

JCOUL specifies a coulometric titration if JCOUL = 1
(no volume change). Dilutions will be calculated if JCOUL

= 0.

8.2.2.8. Concentrations/8F10.6/TOTC(I), EZERO(I),

ADDC(I), VINIT

TOTC(I) is the initial millimoles of species I in the
order of JQRO(I).

EZERO(I) is the intercept of a calibration plot for the
1th electrode (ignored if JEL(I) = 0).

ADDC(I) is the concentration of the titrant solution.
In all cases, a number must be entered for each species,
even if it is zero.

VINIT is the initial volume of the solution in milli-

liters.

B.2.2.9. Slope/8F10.6/SLOPE(I)

Slope (I) contains the slope from a calibration curve

for the ith electrode (ignored if JEL(I) = 0).

232

8.2.2.10. Data input/I5, 8F8.3/LUIGI, TITRE(I),

EMF(I).

LUIGI is a data marker for the end of a data set.

If LUIGI = 0, the next card is read. Another data set
to follow is indicated by LUIGI = l. LUIGI = 2 that the data
are from coulometric measurements, with current and frac—
tional efficiency read instead of an experimental point.
LUIGI < O signals the end of all data sets.

TITRE is the volume titrant added (or the time of cur-
rent passage in sec).

EMF is the measured potential (or pH) for each electrode.

8.2.2.11. Statistics/I5/JPRIN

JPRIN controls the statistical output as follows:

JPRIN Statistical Analysis Tables Graphs
0 no no no
1 yes no no
2 yes yes no
3 yes no yes
4 yes yes yes

8.2.2.12. Termination/I5/NSET

NSET = l for another set of formation constants, NSET

= O for another complete set of data, and NSET = —l for the

 

233

termination of a run.

3.2.2.13. 78

B.2.2.1U. 6
78
9

 

 

 

 

APPENDIX C

THE USE OF THE PROGRAM FARM2.TSK FOR THE
DETERMINATION OF MONOPROTONATED PAA-CALCIUM

FORMATION CONSTANTS FROM POTENTIOMETRIC DATA

. _.. _. ___—V;

 

 

 

 

THE USE OF THE PROGRAM FARM2.TSK FOR THE
DETERMINATION OF MONOPROTONATED PAA—CALCIUM
FORMATION CONSTANTS FROM POTENTIOMETRIC DATA

C.l. Scope of FARM2.TSK

The program FARM2.TSK is a limited-utility computer
program designed to solve repetitive formation constant
calculations based on potentiometric data. The program
evolved to fill the need for the calculation of mono—
protonated PAA—calcium formation constants from potentiom—
etric data. The program MINIQUAD was found to be unsuitable
for these calculations. The program will solve five simul—
taneous equilibrium equations to yield an exact solution for
the mono—protonated formation constant (with the deproton—
ated complex input as a known value) and the resulting

standard deviation of that value.

0.2. Program Derivation

PAA is a triprotic acid symbolized by H3L in its com—

pletely protonated form. The three acid equilibria are
+
' "J" ‘* = ' H H (1)
H2]; + . H3L Kl [H3L1/[ 2th l
— -2 +
-— + _ _ 2
HI, 2 + H : H2L K2 _ [H2L J/EHL ][H ] ( )
-2 + _
L‘3 + H+ 2 HL‘ K3 = [HL l/[H ML 3] (3)

' ' f ll arise
When PAA complexes With a metal ion, two -orms usua y

234

 

 

M+2 + L'3 : ML" K = [ML‘][/[M+2][L'3l (A)

M+2 + HL‘2 : MHL K = [MHL]/[M+2]£HL'2] <5)

The K1 is such that this equilibria may be ignored above

a pH of 4, resulting in the mass balance equations

[thotal = [H2L-]+[HI-2]+[L_3]+[MHL]+[ML_] (6)

[Mltotal = EM+21 + [MHL] + [ML—1 (7)
Subtracting (7) from (6) gives
[LlT — [MJT + IM+21 = [H2L_l + [HL‘2] + [L‘31 <8)

The right hand side may be restated in terms of only [HL]

from Equations (2) and (3) to yield

[LIT - [MlT + £M+21 = K2EH+JEHL'J + [HL‘] + [HL-J/

K3EH+1 (9)
OZ“

EHLl = ([LlT - [MJT + [M+2])/

(K2[H+l + 1/K3[H+1 + 1) (10)

From Equation (A) the concentration [ML] is

 

 

 

 

 

[ML] = KdEM+21[L‘3J (11)
and from (3)

[ML] = KdtM+2JEHL‘2J/K3[H+J (12)
Then from Equation (7)

[M]T = [M+21 + Kdtm+2JEHL']/K3EH+] + [MHL] (13)
or

[MHL] = [MJT _ [W2] - KdEM+2][HL'J/K3[H+] (1A)

The concentration EM+21 was known at any given point in the
titration from emf data, with the pH (and hence the [H+])
also measured at each point. Since K3, K2, and Kd were
known, as well as the total amount of ligand and metal ion,
the equations were exactly solvable at each point. The

average Km was then calculated over all the data points,

with the standard deviation calculated for this average.

Data Input File - the data file was created, named,

The

C.3.
and edited using the programs RXl (202) and TECO (203).

reader is referred to these for further information.

 

 

 

237

C.3.l. Descriptive title

 

0.3.2. Control Data/FlS.7/SLOPE, INTERCEPT, RESID,

TOTM, CONCL

SLOPE, INTERCEPT, and RESID were the Values derived from

a KINFITA plot of the calibration data (see B.l.).

TOTM was the total mmoles of metal ion, while CONCL was the

concentration of L in molarity.

0.3.3. Data/FlS.5/MLL, EMF, PH

MLL was the milliliters ligand, EMF the measured emf,

and PH the measured pH.

0.4. FARM2.TSK

The program FARM2.TSK was then run. Upon the inquiry
"input file?", the file name created in C.3.2. was given.

When "output file?" appeared, the desired name of the out—-

put file was given.

The program may be modified through FARM2.FTN to suit a

particular purpose. The file is then compiled and task

built to give a new FARM2.TSK.

 

 

APPENDIX D

SOME CROWN AND CRYPTAND COMPLEX FORMATION CONSTANTS
WITH MANGANESE (II) ION DETERMINED BY

ELECTRON SPIN RESONANCE

 

 

 

 

 

 

 

 

SOME CROWN AND CRYPTAND COMPLEX FORMATION CONSTA
_ NTS WITH
MANGANESE(II) 10N DETERMINED BY ELECTRON SPIN RESONANCE

D.l. Crown Complexes with Manganese (II) Ion.

The complexation properties of the crown ethers 1204,
1505, and 1806 (Figure l) with manganese (II) ion were
studied using electron spin resonance. The data were col—
lected following the procedures outlined in Sections 2.2.1.
and 2.3.1.

There was found to be no significant formation of a
complex within the experimental error. This was evidenced
by no change in the intensity, linewidth, or saturation
properties of the free manganese (II) ion resonance upon
the addition of the ligand. This behavior was exhibited
even when the mole ratio of ligand to metal was as high
as 10:1.

Since the error in the measurement of the free manganese
ion concentration was estimated at iA%, approximately 4%
of manganese could complex with no effect on the signal
intensity (it was probable that there would be no effect
an the linewidth or saturation properties either). With
1 total concentration of manganese 10—3 M and a ligand
zoncentration of 10‘2 M, this puts an upper limit of ap-
>roximately A for the formation constants of l2CA, lSCS,

1nd 1806 with manganese. That is, if a complex exists,

rhich is doubtful, it would be very weak.

238

 

 

 

 

239

D.2. Cryptand Complexation with Manganese (II) Ion

The complexation Of C211 (Figure l) with manganese (II)
ion was also studied. The experimental procedures were
as outlined in 2.2.1 and 2.3.1 with one exception. It
was found that the manganese (II) ion oxidation is greatly
enhanced by the addition of the cryptand (the rate is
dependent on the ligand concentration). However, this prob-
lem was alleviated by degassing the solution with deoxy—
genated nitrogen (183)for 15 minutes prior to the addition
of the ligand. In this manner, the oxidation of manganese
(II) ion was negligible over the course of the experi-
ment. Such behavior (oxidation enhancement) has been ob-
served previously (28), although not specifically for manganese.

The cryptand C2ll is a very strong base, with a pKl(BH+)
of 11.3 and a pK2(BH2+2) of 8.1 (14). The pH of the solu—
tion was adjusted to 8.5 to partially deprotonate the ligand.
The formation constant was then found to be log Kf = 1.6
i 0.1. This value is comparable to formation constants ob-
tained for other divalent cations (28). The presence of
complexation would tend to indicate that the oxidation of
manganese (II) ion is occurring through an intermediate
:omplex. The cryptand would tend to stabilize a higher
>xidation state for manganese, with subsequent oxide forma-
;ion upon decomplexation.

Hence, it has been shown that electron spin resonance

'or manganese (II) ion is a versatile technique. BObfl

 

 

____-_.._ —_. - m-nd— v—z— ~..

2ND

protic and aprotic ligand complexation systems may be
studied using this technique, which should lead to increased

interest in the complexation properties of this biologically

important ion.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

1
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10.

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16.

 

 

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