69-22,134 GOSSER, Leo Anthony, 1942- an EQUILIBRIUM AND

This dissertation has been microfilmed exactly as received 69-22,134

GOSSER, Leo Anthony, 1942- AN EQUILIBRIUM AND REACTION STUDY OF THE GLYCYLGLYCINATE - GLYOXALATE NICKEL(II) AND ZINC(II) SYSTEMS.

The Ohio State University, Ph.D., 1969 Chemistry, analytical

University Microfilms, Inc., Ann Arbor, Michigan AM EQUILIBRIUM AND REACTION STUDY

OF THE GLYCYLGLYCINATE - GLYOXALATE -

NICKEL(II) AND ZINC(II) SYSTEMS

DISSERTATION

Presented In Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University

Leo Anthony Gossor, 3«S.

The Ohio State University 3.969

Approved by

Adviser epartment of Che ACKNOWLEDGMENTS

The author wishes to express his appreciation to Dr. Daniel

L.Leusslng for his encouragement and advice during the course of this work. Financial aid from the Ohio State University in the form of both research and teaching assistantshlps, from the

National Science Foundation for a Summer Fellowship, and from

National Defense loans Is'gratefully acknowledged. I am especial ly grateful to my wife, Mary Louise, for her continuous patience and assistance in the preparation of this dissertation. This work is dedicated to my parents in appreciation for their gifts of faith and love.

ii VITA

May 30, 1942. . . Born - Shelby, Ohio

1964...... B.S., St. Vincent College, Latrobe, Pennsylvania

1964-1966 .... Teaching Assistant, Department of Chemistry, The Ohio State University, Columbus, Ohio

Summer of 1966. . National Science Foundation Summer Fellow, Department of Chemistry, The Ohio State University, Columbus, Ohio

1966-1969 .... Research AssistantDepartment of Chemistry, ThtkOhio .State.‘Uniy.prtsity, Columbus, Ohio '■ ':■■■■■ *..1 ’■ . ; t i , 1969...... Scifcnt.ist-,,.Scifcnt.is ty. War^'ejr/iliainhjert h\ir,rier>hlambjfir 'Research c ■'K.es eari institute, •• Morria \?icihs/':NewvrjV:rstty • ' • *. •. 'I . > * I

PUBLICATIONS .

None

FIELDS OF STUDY

Major Field: Chemistry

ill TABLE OF CONTENTS

ACKNOWLEDGMENTS

VITA......

TABLE OF CONTENTS

LIST OF TABLES

LIST OF FIGURES . . .

Chapter I. INTRODUCTION

A. Schiff Base Systems...... B. Determination of Equilibrium Constants . . . . C. Glycylglycine-Glyoxalate Mixed Syste\n , . , . D. Statement of the Problem ......

II. EXPERIMENTAL

A. Reagents B. Procedure

III. NUMERICAL TREATMENT ...... ,

IV. RESULTS AND DISCUSSION

A. Glycylglycine and Its Metal Complexes B. Schiff Base and Mixed Systems. . . . C. Schiff Base Reactions...... D. Summary of Results ......

APPENDIX A. DATA

APPENDIX B. FRACTION PLOTS

LIST OF REFERENCES LIST OF TABLES.

Table 0 Page

.1. pK^ Values Reported for Glycylglycine; ...... kO

2 * PK0cl Values of Glycylglycine and Related Compounds. . • UO 3. Formation Constants for Glycylglycine...... 1»8 * It. Near-IR and Visible Spectra...... 30

5 1 Amide Dissociation Constants...... 32

6. Schiff Base Formation Constants...... 33

7* Formation Constants for Glycylglycine-Glyoxalate . . . 62

6. LEF V a l u e s ...... 67

9# Carbon Dioxide Evolution...... 86

10. Equilibrium Constants for Glycine...... ' ...... 112

v LIST OF FIGURES

Figure 0 Page

1, Reactions of the Pyridoxal Schiff Base Metal Complex...... 2

2, Mechanism of Pyridoxal Schiff Base Reactions... It

3, Structures of Schiff Base Complexes ...... 7

1*. Electron Mobility in Aldehydes, ...... 8

5. Titration Curve of Glycylglycine with H C 1 ...... 35

6. Titration Curve of Glycylglycine with NaOH...... 37

7# Formation Curve of Glycylglycine...... 39

8. Titration of Hi(ll)-Glycylglycine ...... U3

9. Titration of Zn{ll)-Glycylglycine ...... U5

10. Formation Curves of Glycylglycine-Metal Complexes . . 1*7

11, Titration of Glycylglycine-Glyoxalate ...... 55

12, Change in pH with Time for Addition of Glyoxalate to Glycylglycine-Metal Solutions...... 58

13. Change in pH for Addition of Glyoxalate to Hi(ll)-Glycylglycinate Solutions...... 60 lU. Stepwise Formation Constants in Ni(II)-Mixed System...... 6h

15.' Stepwise Formation Constants in Zn(II)-Mixed System...... 66

16 . HMR Spectrum of Glycylglycine-Glyoxalate at 25°C. . • 71

17. NMR Spectrum of Glycylglycine-Glyoxalate at 0°C .. . 73

vi ■

LIST OF FIGURES (cont.)

Figure if Page

L8. UV Spectra of Glycinate-GIyoxalate ...... 7T

19. Trans aminated Schiff Base...... • T • . 78

20. UV Spectra of Glycylglycinate-Glyoxalate...... 80

21. UV Spectra of Glycylglycinate-Glyoxalate...... 8 3

22. Effect of Zn(ll) on the UV Spectra of Glycinate-GIyoxalate...... • • • 88

23. Effect of Zn(ll) on the UV Spectra of Glycylglycinate-Glyoxalate ...••• ...... 90

2l», Effect of Hi(ll) on the UV Spectra of Glycylglycinate-Glyoxalate ...... 93

25, NMR Chemical Shift as a Function of n for Glycine...... 96

26, HMR Chemical Shift as a Function of By for Glycylglycine...... 98

27, HMR Spectra of Glycylglycinate-Zn(ll)...... 100

28, HMR Spectra of Glycylglycinate-Glyoxalate...... 103

29* HI© Spectra of Glycylglycinate-Glyoxalate...... 106

30. HMR Spectra of Glycylglyclnate-Glyoxalate-Zn(ll) . • 109

31. HMR Spectrum of Glycinate-GIyoxalate ...... Ill

vii I. INTRODUCTION

Much attention has been given in recent years to the investiga tion of model enzymatic systems. It ‘is hoped that a better under standing of biological processes will develop from this work on model systems. Already significant advances have been made in the elucidation of the mechanisms by which some of the more common enzymes function. An example of the progress toward an understand ing of enzymatic reactions is the recently completed work of Dunathan 1 2 et al. * on the stereochemistry of enzymatic transaminations.

Progress along these lines is dependent upon the knowledge of funda mental chemical reactions. A basic quantitative understanding of model systems is, therefore, of paramount importance in the applica tion of these systems to their more complex enzymatic counterparts.

A. Schiff Base Systems

As a result of their role in transamination reactions, Schiff bases and Schiff base-metal complexes have been the object of con siderable interest in the model enzymatic systems. In addition,to their function as intermediates in transamination reactions, Schiff bases have also been investigated as activated intermediates for decar boxylations, trans-iminations, the reversible cleavage of /3 -hydroxy-

3 amino acids, deaminations and racemizations. Metzler and Snell have contributed a wealth of information concerning the pyridoxal systems.

1 Pyridoxal is of great interest as the vitamin cofactor of enzymatic transamination reactions. From a series of work, on pyridoxal, Metzler

Ikawa and Snell^ have proposed a general mechanism in which the pyri doxal forms a Schiff base Intermediate. In non-enzymatic systems the

Schiff base is bound to a metal ion, whereas, in enzymatic systems there is probably not a metal complex. The following mechanism ex plains the role of pyridoxal in a variety of reactions. Figure 1 shows the structure of the proposed intermediate. It can be seen that this form of intermediate may lead to different types of reactions, depending upon which bond to the amino acid carbon atom is most easily fractured.

H FRACTURE OF PRODUCES I* .0 R — C -Z-C? Bond 1 Cleavage reactions 0 Bond 2 Carbanion formation H-C

Bond 3 Decarboxylation h o h £c

Pyridoxal Intermediate

Figure 1. Reactions of the Pyridoxal Schiff Base Metal Complex Figure 2 Illustrates the mechanism for transamination and racem-

Ization. The components of the pyridoxal system which are essential 4 for Its catalytic properties are the formyl and phenolic groups in the proper orientation and the pyridinium ring.

Work with compounds structurally similar to pyridoxal has in creased the scope of systems which may be applicable as model enzymat ic systems. Ikawa and Snell** found that 4-nitrosalicylaldehyde in the presence of metals possesses some of the catalytic activity of Its pyridoxal analog. The dehydration of serine, the splitting of threo nine to glycine, and the desulfhydration of cysteine are promoted by

4-nitrosalicylaldehyde through the same type of Schiff base intermed iate as that of pyridoxal. This work along with investigations on other salicylaldehyde compounds has shown that nearly all pyridoxal- catalyzed reactions can be catalyzed by other compounds albeit perhaps not as effectively as pyridoxal.

Harada and Oh-hashi^ have studied the condensation reactions of carbonyl compounds with the N-salicylideneglycinatocopper(II) complex.

Under moderate conditions (pH = 7-8, 25 - 1°C, 24 hours in an aqueous solution), they have reported over 80^.' yield of threonine by the con densation of acetaldehyde with glycine. Likewise lesser yields of phenylserine, /3 -hydroxyaspartlc acid, and serine were formed by the condensation of benzaldehyde, glyoxyllc acid and formaldehyde respect ively with the glycine Schiff base complex. The formation of

/3 -hydroxyamino acids as promoted by salicylaldehyde and copper(II) is shown in the following scheme. u

HOHjC

N v i ^ C H , H* Pyridoxal Intermediate I-H+

R - ^ - c f

0 H-C HOH?C

R I R-C-COO' H*- ac-C O O II

HOHjC HOH,C

N CH

hydrolysis hydrolysis

Racemization Transamination Figure 2. Mechanism of Pyridoxal Schiff Base Reactions 5

? " f - . 0

pH T-8 +RCHO > H20

25°C

The similarities between this reactant and the intermediate presented previously for pyridoxal-promoted reactions are most apparent in the chelate rings formed; these are identical for both pyridoxal and sal- 7 icylaldehyde compounds. This also corroborates the work of Cennamo who found that the carboxylate group of the amino acid and the phen olic group of pyridoxal are essential for the chelate formation.

Although the model systems that have been investigated most thor oughly are of the pyridoxal-related type, it is interesting to remem ber that the first non-enzymatic transamination as reported by Herbst 8 9 and Engel ’ was between pyruvic acid and various amino acids. Decar boxylation frequently accompanied the transamination reaction which was carried out in boiling aqueous solution. The earlier works on

a . -keto acids and amino acids were carried out under the conditions of high temperature and high pH, More recently, however, Nakada and

Weinhouse^ have reported the rapid non-enzymatic transamination of glyoxalate with various amino acids at a lower temperature and in the physiological pH range.

A comparison (Figure 3) of the structures of the Schiff base com plexes which have been discussed thus far illustrates several factors.

A notable difference is the change from one five- and one six- membered chelate ring structure for the pyridoxal and salicylaldehyde

Schiff base complexes to two five-membered chelate rings for the glyoxalate 'ind pyruvate Schiff base complexes. Also, substitution of

the carboxylate group of glyoxalate and pyruvate for the resonance stabilized aromatic rings of pyridoxal and salicylaldehyde results In a decrease of the electron mobility and a decrease in the amount of

-bonding within the chelate ring. Thus while glyox^J^te’and pyru vate possess some similarities to pyridoxal in structure and reactiv- 'u v ' ity, the comparison is limited.

The successful promotion of reactions via a Schiff base inter mediate depends on several factors. Among these are the stability of

the Schiff base, the stereochemical restrictions and a position of equilibrium, from a thermodynamic standpoint, between the starting materials and the products.

Perhaps most important, however, is the extent to which the

OC-carbon of an amino acid can be promoted as an electrophilic site by the electron-withdrawing constituents of the Schiff base inter mediate of its metal complex. D o s e ^ has shown that an electron- donating substituent in a position ortho or para to the aldehyde group of benzaldehyde serves to reduce the tendency to transamlnate.

In his discussion of the pyridoxal systems, Melster 12 has empha sized the importance of the delocalization of electrons from the

a -carbon of the amino acid. This is facilitated by a structure which possesses a planar system of conjugated double bonds from the

OC -carbon of the amino acid to the heterocyclic nitrogen atom as shown

In Figure 1. Figure 4 illustrates the similar electronic displacements Pyridoxal Salicylaldehyde

■? F 0 I / H— C C H— C cf'

h - c ^ m ^ 0 h3c - J ^ m / 0 \J W \ V 0

Glyoxylate Pyruvate

Figure 3. Structures of Schiff Base Complexes that are possible for aldehydes which are known to form Schiff bases.

Chelation of the Schiff base formed by the aldehydes of Figure 4 to a

divalent metal ion should increase the withdrawal of electrons from

the (X -carbon and thus increase the reactivity of the Schiff base.

" - v H - C ^ 0 OH HOH.C

c h i\n h 3c — c

Figure U, Electron Mobility in Aldehydes

Metal ions are known to promote, in this manner, a large number

of organic reactions including transaminations and aldol condensations. 13 Williams and Busch demonstrated the labilization of the OC -hydrogen

atoms of glydine upon chelation. They prepared the [Co***(en) 2 (glycine}]

(N0^)2 complex in acidic D2 O. . Neutralization of the solution to a

blue litmus color caused rapid exchange of the (X -hydrogens of glycine

as observed by nuclear magnetic resonance techniques.

Several syntheses have been developed which depend upon this che

lation for their success. A single step synthesis^ of DL-serine from

glycine and formaldehyde in an alkaline aqueous solution containing

catalytic amounts of CuSO^ has been reported. Threonine^-* has been

produced by the reaction of acetaldehyde with the copper(II) complex of

glycine in weakly alkaline solutions. Likewise, pyruvic acid reacts with Cu(Gly) 2 to give / 3 -hydroxy-/# -methyl aspartic acid.*^ The yield

for these aldol condensations has been increased^ by substituting the

[co***(en)2 gly]+2 complex for the Cu(Gly)2 complex. This improved

effectiveness of the metal ion in activating the methylene group of

glycine is due to the two formal positive charges of the

[CoIII(en)2 gly]+^ complex,

Harada and Oh-hashi^ have shown that carbonyl compounds react

rather easily with the N-salicylideneglycinatoaquocopper(II)- complex

to form/^-hydroxy- (X -amino acids in neutral to weakly alkaline aque

ous solutions at room temperature. The facility with which this

reaction proceeds is partly attributable to the increased withdrawal

of electrons from the Of-carbon of glycine by the imine bond of the

Schiff base. Metzler*: Longnecker and Snell^ investigated the rever

sible catalytic cleavage of hydroxyamino acids by pyridoxal and metal

salts; thrie conclusion was that the amino acid probably participates

in these reactions only via the Schiff base. They also observed the

formation-of /3 -hydroxyaspartate from glyoxalate and glycine in the

presence of Al(III) at 30°C. In this case, glyoxalate catalyzes the

reaction by formation of a Schiff base and also serves as the condensing

aldehyde for the aldol condensation reaction.

B. Determination of Equilibrium Constants

Previous investigations of Schiff base complexes have been mostly

of a preparative nature, or for the purpose of qualitatively comparing

similar systems. There is no longer any doubt that'Schiff base com

plexes play an intermediate role in the reactions of non-enzymatic 10 systems. The existence of Schiff base complexes has been verified by such techniques as NMR, absorption spectroscopy, chromatography and elec trophoresls. 19 Nakahara and co-workers have isolated and characterized the

» Cu(II) salicylideneglycinate complex and also the glycylglycinate 20 analog. The Pd(XI) Schiff base complexes of glycine with pyruvate and of glyoxalate with alanine have been prepared; the NMR spectra of both complexes are Identical and rapid exchange of the methylene pro tons occurs upon dissolution in D2 O. An attempt to prepare the Cu(II) 21 complex of the above Schiff bases also resulted in identical com plexes for both sets of ligands. Identification of the amino acid as glycine in both cases gave positive proof of the transamination reaction which occurred between the complexed alanine and glyoxalate.

Equilibrium constants have been determined for the formation of relatively stable Schiff bases and of their metal complexes. For exam- 22 pie, Lane and Kandathil have worked on Schiff bases derived from

O -aminophenol and benzaldehydes; under the conditions employed, none of the Schiff bases underwent hydrolysis. Hovey et al.^ attributed the stability of Schiff bases derived from aromatic amines and alde hydes to an Increase of the resonance ene?'gy in the system brought about by the presence of the aromatic groups. The stability of the

Schiff bases in these cases permits the equilibrium constants to be determined readily by spectrophotometric techniques.

In comparison, only a few investigations have dealt with Schiff 22 bases that dissociate extensively In solution. Lane and Kandathil stated that they could not study Schiff bases derived from oliphatic 11 aldehydes and O -aminophenol due Co Che hydrolysis of Che ligands. A problem arises In Che spectrophoComeCrlc approach due Co Che existence of a number of specCrally similar species In soluClon; Co overcome these difficulties, solution conditions are arranged so Chat only one 24 or two species predominate. Elchhorn and Dowes have studied the partially dissociated Schiff bases derived from pyridoxal and various amino acids by this method. In the subsequent calculation of the form ation constants, the mathematical equations are simplified by the assumption that certain species are negligible. This method is inher ently accurate only to the degree that the assumptions are correct. 27 Watters*0 has shown that the pH method developed by Bjerrum can be applied successfully to mixed ligand systems provided that the acid-base properties are sufficiently different. Several works by 28-35 Leussing et al. have employed this potentiometric method and have coupled It with refinements in the mathematical analysis.

The use of computer techniques enables complex systems to be analyzed rapidly and thoroughly by taking into account all probable species in solution. This approach lends itself particularly suitable 37 to Silien's pitmapping procedure which Leussing has applied to

Bjerrum's pH method with notable success. An added advantage of this pH .method is that the mathematical treatment requires no assumption as to the nature of the ligands in the coordinated state; whether they are condensed as the Schiff base or are coordinated to the metal inde pendently has no effect upon the mathematical equations. Furthermore, from the relative magnitudes of the constants, one can gain signifi cant insight into the nature of the mixed complexes. 12

C. Glycylglycine - Glyoxalate Mixed System

Glycylglycine Is the simplest peptide. In this respect, it has been found useful as a convenient stepping stone from simple acids to more complex biologically significant peptides and proteins.

It is not surprising that a considerable number of works have dealt with the metal stability constants of the glycyl peptide series.

Li and Chen^® have reported the stability constants of glyclnamide, glycylglycine, triglycine and tetraglyclne with Cd(II), Zn(II) and

Ni(II). They have observed that the four ligands are about equally stable; this implies that the metal ion is bound only to the terminal

-NH2 and the adjacent amide group for all four ligands. Murphy and 39 Martell have investigated the Cu(II), Mn(II) and Mg(II) stability constants for the glycyl series and reported that the stabilities decrease for the metals in that order. It has been shown that a lin- 40 ear relationship exists between the pKa of the amino group in the series of glycyl peptides and the log of the metal complexes. This

Indicates that the initial step' in complex formation is probably inter action between the terminal NH2 group and the metal ion, followed by interaction at the peptide bond. Additional support for this mechan ism of complexation is the effect of acetylatlon of the NH2 group of glycylglycine upon its affinity for Cu(II) ions; Koltun, Fried and

G u r d ^ report that the acetylated glycylglycine shows the same affinity for Cu(II) as does acetate, so the amide groups do not appear to con

tribute to the stability in-this case. 13

The manner In which the amide group participates In complexatlon has been a point for speculation. LI, Doody and Vhlte^2 postulate that Zn(II) and Co(II) probably coordinate with the amide oxygen, whereas Cu(ll) and Ni(II) may complex with the amide nitrogen. Some metal Ions have been reported to Induce Ionization of the amide hydro-

/ gen by their electron-withdrawing ability. Dobbie and Kermack ob served the titration of the amide hydrogen In the presence of Cu(II).

At this time, it was thought that only Cu(II) ions possessed suffi cient electron-attracting ability to induce the reaction. Subsequent

ly, an investigation of the Ni(II)-glycyl peptide complexes has shown 44 that the amide ionization occurs for Ni(II) but at a higher pH, and accompanied In some cases by a change from octahedral to square planar coordination of the Ni(II). Gillard, Harrison and McKenzie^"* have prepared and characterized the optically-active Co(III) complex of doubly ionized glycylglycine; a determination of the crystal structure*^

TTT of the NH^(Co (GG)2)'2H20 complex by x-ray diffraction has revealed

the amide nitrogen is bound to the Co(III) ion. An x-ray diffraction

study ^ of the K2 Cu**(GG= ) 2 complex likewise has shown the amide nitrogen bound to the metal ion.

Glyoxalate, as has been mentioned previously, has been the object -

of .considerable Interest In the study of non-enzymatic systems. Dixon

in particular has done some highly significant work on the transamina

tion of peptides and proteins. In a series of works, Dixon and co- workers^ have developed a mild method of converting the N-terminal

residue of a protein into an oxo-acyl group. This oxo-acyl group may

then be easily removed by treatment with an aromatic diammlne. It is 14 hoped that this method may be applicable to the selective identifica tion of N-terminal residues of peptides and proteins.

D. Statement of the Problem

This investigation is a continuation of a series of work done in these laboratories which deal with the determination of the equilibrium constants of Schiff base complexes.

The Schiff base system consisting of glycylglycine, glyoxalate and metal ion should be a significant step in the application of simple amino acid Schiff base systems to the more complex biological reactions which involve peptides and proteins. Glyoxalate is active in transam ination reactions and in aldol condensation reactions; in some respects,

it is comparable to pyridoxal. Glycylglycine possesses the amide func

tion which is the primary distinction between the individual amino acids and amino acids that evist in the peptide configuration. In this sense,

the reactivity of the glycylglycine Schiff base system should more closely parallel that of complex biological reactions than should the reactivity of amino acid Schiff base systems.

The objectives of this research were threefold: to investigate

the Zn(II)- and Ni(II)-glycylglycinate complexes in terms of their sta bility constants and the ionization of the amide hydrogen; to charact

erize the glycylglycine-glyoxalate Schiff base system in terras of its

formation constants and its stability constants with Zn(II) and

Nl(II), and to compare this system with similar systems; and-to in

vestigate the possible reactions of the Schiff base formed by glycyl

glycine and glyoxalate. II. EXPERIMENTAL

A. Reagents

All solutions unless otherwise specified were prepared with dou bly-distilled demineralized water. The ionic impurity level of this water is specified at less than 0.1 ppm.(expressed as NaCl).

Glycylglycine

Glycylglycine obtained from Sigma Chemical Company was vacuum- dried for 24 hours. The dried glycylglycine without further purifica tion was 100.0^0.2% pure as shown by titration in the presence of 52 formaldehyde ("formal1' titration). Solutions were prepared as needed.

Sodium Glyoxalate Monohydrate

Anhydrous sodium glyoxalate as obtained from Sigma Chemical Com pany was converted to the monohydrate form. This was accomplished by dissolving the anhydrous material in a minimum amount of water

(^0.4 g/ml) and recrystallizing the monohydrated sodium glyoxalate 53 from a 50% acetone solution. After the recrystallized monohydrate had been vacuum-dried for 24 hours, the purity was checked by passing a weighed amount of the monohydrate through a strong cation exchange column (amberlite IR - 120). The purity was found to be 100.0— 0,17.,

Solutions were prepared as needed.

15 16

Glycine

Glycine obtained from Sigma Chemical Company vas recrystallised from an ethanol-vater solution and vacuum-dried for Zk hours. The purity as determined by "formal” titration**2 was 100.0 ± 0.2JS. Sol utions were prepared as needed.

Zinc Chloride

Reagent grade sine oxide powder from Baker and Adamson was dis solved in an amount of hydrochloric acid calculated to produce a min imal excess of hydrochloric acid. The dissolved Zn(Xl) ion was then diluted to a final concentration of~0.2M. The sincchloride concen tration was determined accurately by titration with a standard solution of the disodium salt of ethylene-diammine tetraacetic acid (EDTA) ac cording to the procedure described by Flashka.^ A known volume of the Zn C ^ solution was passed through a strong cation exchange column

(Amberlite XR-120) and the total acid content of the eluent solution was determined by titration with base to a phenolphthlein end point.

The two methods agreed to within 0.3/5. From the difference between the total acid and the acid released by zinc exchange, the excess acid vas found to be less than O.OOIM} this agreed with a measured pH of

3.0 for the ZnCl2 solution.

Hickel Chloride

A calculated amount of "Baker Analyzed" reagent grade nickel chloride hexahydrate was dissolved in sufficient water to prepare a stock solution ~ 0.050M in NiClg. The solution vas standardized by titration with a standard EDTA solution according to the method de scribed by Vogel*55

Potassium Chloride

Reagent grade potassium chloride from the J« T. Baker Chemical

Company vas recrystallized from vater prior to use* A stock solution of 2.50M KC1 vas prepared* This vas used to maintain the ionic strength of solutions at a constant level* *

Sodium Hydroxide and Hydrochloric Acid

Standard solutions of hydrochloric acid and carbonate-free sodium hydroxide vere obtained from the Ohio State University Reagent

Laboratory.

Buffers

Three standard buffer solutions vere prepared from National Bur

eau of Standard reference materials. These solutions vere used to

standardize the pH meter* A solution of 0.01M potassium acid tartrate

is stated to have a pH of 3.639 at 25°C. A solution of 0.025M pot

assium dihydrogen phosphate is stated to have a pH of 6*862 at 25°C.

A 0.01M solution of borax (Nag 0^.10 H^O) is stated to have a pH

of 9.130 at 25°C. For the intermediate pH range, the tartrate and

phosphate buffers vere used Jointly. For the higher pH range (>pH=9)*

the borax and phosphate buffers vere used Jointly*

Deuterium Oxide

Deuterium oxide (DgO) 99.9# pure from Isotopes Specialties Com pany vas used for the preparation of some of the solutions required in 16

the nuclear magnetic resonance spectra portion of this investigation* «

Deuterated Glycine

Previously recrystallized glycine from Sigma Chemical Company vaa dissolved in D^O 13.5g/50ml), evaporated, and vacuum-dried for

Zk hours* To determine the extent of deuteration, a weighed portion of the deuterated glycine was titrated in the presence of formalde hyde* Based upon theoretical molecular weights of 75*07 and 7 8 . 0 9 respectively for non-deuterated and deuterated glycine, B0% of the

exchangeable hydrogens were replaced by deuterium*

Deuterated Sodium Olyoxalate

Anhydrous sodium glyoxalate from Sigma Chemical Company was

dissolved in D^O (~ 5g/25*l)» The concentration of sodium glyoxa-

late was determined by passing 1 ml of the solution through a strong

cation exchange column (Amberlite IR-120) and titrating the eluent with 0«1N NaOH to a phenolphtholein end point.

Sodium Deuteroxide

A sodium deuteroxide solution was prepared by dissolving 8g of

sodium hydroxide pellets from the J* T. Baker Chemical Company in 100ml

of DgO* This solution was standardized by titration with 1.0N HC1 to

a phenolphtholein end point*

Formaldehyde

Formaldehyde solution from Fisher Scientific Company containing

36.9% formaldehyde was used in the "formal" titrations. Blank titra

tions vere run to determine the amount of formic acid present in the sol

ution. 19

B. Procedure

Solutions were prepared and used at 25°C unless otherwise stated.

The ionic strength of solutions when so stated was maintained at 0.50 by the addition of sufficient quantities of 2.5M stock KC1 solution.

All titrations used for the determination of equilibrium con stants were carried out in a water-jacketed titration vessel with water maintained at 25.0°C circulating through the jacket. The water temperature was regulated at 25.0 i 0.1°C by a "Heto" circulating water bath. The titration vessel was covered with a layer of "Parafilm" to exclude air and contaminants. Nitrogen gas was continuously passed through the cell during the course of a titration; to insure uniform humidity and temperature of this nitrogen gas which passed through the cell, the nitrogen was initially passed through a water trap at 25.0°C,

Tltrants were added by means of either a Manostat 3-ml piston- type buret or a Metrolim 10-ml piston-type buret. The pH measurements were made using either a Corning "Model 12" pH meter or a Radiometer

"PHM 25a" pH meter with a "PHA-925" scale expander. To monitor the pH measurements as a function of time, a Heathkit Servo recorder or a Houston T-Y recorder was used. Unless otherwise indicated, the pH was recorded when the pH had reached a steady value; this criteria was then used as an indication of equilibrium in the solution.

For the determination of the acid dissociation constants of gjycylglycine, 25ml aliquots of 0.003M and O.OOhM glycylglycine were titrated, with 0.1N HC1; alqo, 50ml aliquots of 0.01M glycylglycine and 20ml aliquots of 0.02M glycylglycine were titrated with 0.1N

NaOH. Graphical methods were used to evaluate K^a and Kga from the 20 straight line plots of 5^ functions as explained in Chapter 3*

The formation constants for the simple complexes of glycyl glycine vitb Ni(XX) and vith Zn(ll) were determined from titrations of the glycylglycine-metal solutions vith NaOH, For the Ni(XX) system, 50ml aliquots of solution containing 0.002M NiClg and either

0,008M or 0.012M glycylglycine vere titrated vith 0.1N NaOH. Xt vas found that if the ratio of glycylglycine to Ni(ll) vas 2:1, a pre cipitation of Ni(OH) occurred tetveen a pH of 8-9; therefore, a m m higher ratio of ligand to metal vas used in this case to avoid precip itation and to Investigate more thoroughly the amide dissociation constants of the metal complexes. For the Zn(ll) system, 20-H0ml aliquots of solution containing 0 .0 6 m glycylglycine and from 0 .0 1 M to 0,02M ZnClg vere titrated vith 1,0N NaOH, Both the Ni(ll) and

Zn(Il)-glycylglycine solutions vere stable over the entire range of pH investigated and equilibrium vas attained vithin a matter of seconds.

The stability constants and Bqq o f the glycylglycine-metal complexes vere determined graphically from straight line plots of

functions. The value of Bq^ for Zn(XX) vas evaluated by means of a one-dimensional pitmap procedure. This is explained in more detail in Chapter 3*

In the glycylglycine-Ni(ll) system, Ni(ll) has an electron vith- draving effect sufficient to induce ionisation of the amide hydrogen of glycylglycine. This ionization vas evaluated from the same titra tion as above through the use of a tvo-dimensional pitmap procedure.

The Schiff base condensation reaction between glycylglycine and glyoxalate vas investigated by titrations ©f 1 0 ml aliquots of solution 21 containing 0.01M glycylglycine and from 0.005M to 0.10M sodium glyoxalate vith 0.1N NaOH and by titrations of 20ml aliquots of solution containing 0.1M glycylglycine and from 0.05M to 0.30M sodium glyoxalate vith 1.0N NaOH, The Schiff base association constant (Bq A) vas evaluated by means of a one-dimensional pitmap procedure. Equilibrium vas established vithln a matter of seconds in the glycylglycine-glyoxalate system and the pH of the solutions vas stable for at least a period of several hours at pH values of less than 9* At pH values above 9 t the pH gradually increased; this Indicated that additional reactions vere occurring. To investigate these additional reactions further vork vith this system was carried out. This included an ex periment with thin-layer chromatography to determine the possibility of glycine formation, experiments vith nuclear magnetic resonance to detect nev species, and experiments vith ultraviolet spectroscopy to follow the course of these additional reactions.

For the determination of the "mixed” constants of the glycylglycine-glyoxalate-metal systems, it vas necessary to prepare a separate solution for each point in a titration. This vas as a result of the non-equilibrium conditions encountered. Both the glycylglycine- glyoxalate and the glycylglycine-metal systems vere stable as men tioned previously; however, when all three components vere present in solution, there vas a gradual but continuous decrease in pH observed.

At the intermediate pH values of 5-7» this decrease continued for

^-12 hours (depending upon the pH) and vas followed by the evolution of carbon dioxide and a more gradual increase in pH. 22

The amount of COg evolved vas estimated by means of a NaOH trap.

Nitrogen gas vas bubbled through the enclosed titration vessel and subsequently bubbled through a trap containing & known amount of standard

NaOH solution* The COg from the titration vessel reacted vith an equivalent amount of NaOH* Titration of the NaOH trap vith standard

HC1 to a phenolphthalein end point yielded the amount of unreacted

NaOH* This value, vhen subtracted from the titration of an equal amount of unused standard NaOH solution, yielded a measure of the COg.

Previous work by Leus 6 ing and Hanna^ on the glycine-glyoxalate- metal system indicated that the complex formation equilibrium vas rapid and stable. On this basis, it vas assumed that the equilibrium for the formation of the mixed complexes in the glycylglycine-glyoxalate-metal system vould likewise be rapidly attained. The addition of a volume of sodium glyoxalate solution to an aliquot of glycylgly cine-metal solution produced a rapid decrease in pH followed by a more gradual decrease in pH as mentioned above. This initial decrease in pH vas assumed to be caused by the rapid formation of the mixed com plexes and vas taken as the point of equilibrium for this formation*

This initial decrease in pH vas, therefore, used to determine the mixed formation constants*

For the Zn(ll) mixed system, a series of 100ml of solution con taining 0*02M glycylglycine and O.OO^M ZnClg, and also a series of

100ml of solution containing 0.0l*M glycylglycine and 0.008m ZnClg vere prepared vith various amounts of sodium hydroxide added to give

25-T5JS neutralization of the glycylglycine amino group. From each

1 0 0 ml of solution in the series, 2 0 ml aliquots vere withdrawn for each separate point in a titration. This aliquot vas placed in the titration vessel and the pH of the solution recorded. A volume of

0.2M or O.HM sodium glyoxalate vas added to the titration vessel and the change in pH recorded on a Houston T-Y recorder. The pH of the solution upon completion of complex formation vas determined hy ex trapolation of the recorded pH changes to the point at vhich the initial rapid pH change vas complete. This procedure vas repeated for four 2 0 ml aliquots from each of the 1 0 0 ml of glycylglycine- metal solution prepared. Volumes of 1, 2, 5 and 10ml of the gly oxalate solution vere added respectively to the four aliquots and the change in pH recorded for each addition. The same procedure vas followed in order to obtain the changes in pH for the addition of glyoxalate to the Ni(ll)-glycylglycine system.

The mixed formation constants for these systems are defined in the Numerical Treatment chapter in terms of the metal ion and the two ligands. They are represented as B^, B^* ^21* ®22 resPec'*'ively vith the number of ligands per metal ion indicated by the subscripts.

The first subscript is for the number of glyoxalate ligands and the second subscript for the number of glycylglycinate ligands. A pre liminary evaluation of the mixed systems vas made by the systematic variation of the mixed constants in a four-dimensional pitmap procedure.

Further refinement of the constants vas accomplished by means of a self-converging modification of the general pitmap procedure. This refinement program involves the calculation of a minimum for a multi dimensional paraboloid sum-squares surface and is explained in Chapter 3* 2k

For all of the above types of titrations, the ionic strength (^) vas maintained at 0 . 5 0 by the addition of appropriate amounts of

stock 2.5M KC1, The KC1 vas added to the solutions to be titrated as they vere prepared* No KC1 vas added to the titrant solution except

for the sodium glyoxalate titrant solutions. This vas considered necessary since there vas a 50% change in the titration volume for the final point of the titration.

In mass balance equations, the hydrogen ion or hydroxyl ion con

centration may contribute a significant amount to the total hydrogen

In a Bystera; for this reason it vas necessary to determine the free

hydrogen ion and hydroxyl ion concentration as a function of the

hydrogen ion activity. This vas accomplished by titrating 0.50M KC1

blank solutions vith standard solutions of HC1 and NaOH.

Several experiments vere carried out to elucidate the secondary

reactions found in the glycylglycine-glyoxalate and glycylglycine-

glyoxalate-metal systems.

Thin layer chromatography (TLC) vas used to determine the

presence of glycine vhieh could arise in these systems by hydrolysis

or by transamination. Glass plates (8 "x8 ") vere covered vith a

slurry of Avarin (35g/150ml) 1 mm thick and air-dried overnight prior

to use. Samples vere applied vith a lambda pipet and the plates devel

oped in a 702 ethanol solvent. To detect the spots, a 0,32 ninhydrin ‘

solution in absolute alcohol vas sprayed on the plates and the plates

vere set in a dark place overnight. values of 3 2 . 5 and 2 k ,5 vere

obtained for glycine and glycylglycine respectively. The samples used

for TLC experiments vere taken directly from titration solutions; to 25 reduce the amount of excess salts (KCl) present, the sample vas first passed through an ion-retardation column ( 5 0 cm long, 1 cm 2 cross area) containing Ion Retardation Resin AG-11A8 (50-100 mesh) from

Bio-Rad Laboratories.

A strong anion exchange column in the acetate form vas used to separate glyoxalate from glyoxalylglycine , the product obtained by the transamination of glycylglycine vith glyoxalate. A 100ml column

(Item cross area) vas packed vith Bio-Rad Laboratories AG2-X8 (20-50 mesh) anion exchange resin vhich had been converted to the acetate form. The column vas equilibrated vith 1.0M sodium acetate -1.0M acetic acid buffer containing O.ljS toluene. The column vas Jacketed and maintained at 25 -.0,5°C by means of vater circulated from a

"Heto".circulating vater bath. A Buchler fraction collector vas used to collect fractions of eluent. The fractions vere then tested for the presence of aldehydes vith 2 , l»-dinitrophenylhydrazine according ho to the method of Dixon. ?

Ultraviolet, visible and near-IR spectra vere obtained using a

Cory Model 1^ Recording Spectrophotometer vith matched quarts cells

(0 .1 - 1 0 cm) and neutral density filters.

Nuclear magnetic resonance (NMR) data vas obtained by mean 3 of a

Varian Model A-60 NMR Spectrometer and a Varian Model HA-100 NMR

Spectrometer equipped vith a variable-teraperature probe. For spectra obtained on the A- 6 o, tetramethylammonium chloride (TMAC) vas used as an internal reference. For spectra obtained on the HA-100, an external neat tetramethylsilane (TMS) reference vas used in a precision coaxial insert from Wilmad Glass Company. To correct for the differences in the bulk magnetic susceptibilities! vith this insert the method of

"static" peak splitting^ vas used. A plot of the peak splitting for non-spinning reference samples versus their known magnetic suscepti bilities vas employed as a calibration chart to determine the unknown magnetic susceptibilities of the samples in which the external TMS reference vas used. NMR spectra vere obtained at 0°C for the glycine-glyoxalate system on the Model HA-100 NMR Spectrometer. A meth anol sample vas employed to determine the temperature of the probe before and after the spectra vere obtained.

Potentiometric data vas obtained at 0°C to aid in the interpre tation of NMR data at 0°C. The temperature of the solutions vas maintained at 0°C by means of a "Heto" circulating vater bath con taining a mixture of ethylene glycol and vater. The titrations vere carried out in a "coldroom" maintained at U°C in order to decrease the thermal differences between the solutions and the surroundings. At room temperature (~ 25°C) the large thermal difference (0 to 25°C), due to the difference in the temperature of that part of the electrodes in solution from that part which vas out of solution, caused a notice able effect on the stability and reproducibility of the potentiometric data. By carrying out the titrations in a coldroom at U°C, this effect vas minimised. Known volumes of sample vere withdrawn during the course of a titration and used for obtaining NMR spectra. In the anal ysis of the potentiometric data the additional change in volume vas taken into account. III. NUMERICAL TREATMENT

The equilibrium constants used in this work ore defined as follows!

(HA)a (A-)a (Gx-)a 1C. * H «K_ * H *K « H — ■ * 2a — TuTT «- — (HgA+) * * * ' (HGx)

B * (MGx+ ) « B „ ■ (MG3C2 * .B a (AGx"2) W W W > » o a

B - (MA+) ,B » (MAg) *B * (MA^") 01 (""HA") 02 („«-}(A-)2 03 (M++)(A-)3

B • (MAGx) ,B ■= tKApGx~) -B _ (MAGxg- )

H (M++)U-HGx-) “ (M++)(A-)2(Gx-) 21 (M++)(A")(Gx -)2

(MAGX-? (MDjajj (HDA-Jaj, 22 1------* all ■ -■ ■ ■ % * a21 ■ (K^+)(A-)2 (Gx -)2 (MA+) {MAa)

K » (MDA -2)^ (MD -2)a^ a31 ------* o22 — ;------;-- * a32 ------(MA3") (MDA“) (MDA2-2) where A" represents glycylglycinate (end in some cases glycinate),

-2 - D represents doubly-ionised glycylglycinate and Gx represents glyox- _2 alate* For the constant Bq ^, the species represented as AGx is the

Schlff base as follows: "00C - CH = N - CH 2 - C - X 9 where X “ 0“ for the glycinate Schiff base and X - N-CHgCOO" for the glycylglycinate Schlff base* The mixed constants (B^, B12* B21 B22^ as they are defined above contain the metal complex species MAGx,

2 7 28 — — —2 — - MA^Gx , MAGXg and MAgGxg respectively. The ligands A and Gx for

these complexes may exist combined as the Schiff base or they may be

bound independently; this is not stipulated by the definition of the

constants. The constants K fl1 ^ *^ 3 1 * ^a22 81114 ^a32* as ^ e y 8 1 , 6

defined above, contain metal complex species of the ligands A“ and •2 - -2 D , vhere A and D represent glycylglycinate in tvo different

stages of protonaiion. It is stipulated that A and D are bound

independently to the metal ion in these complexes. The quantity afl

as used here represents the hydrogen ion activity, Quantities in par

entheses represent concentrations.

The simple protonation constants and the simple metal formation

constants are evaluated from straight line plots of n functions as 27 derived by Bjerrum and used previously in this laboratory by 28-31* , Leussing et al. (See Appendix A for the derivation of pertinent

equations)•

For the evaluation of the amide ionization constants of the Ni(ll)-

glycylglycinate complexes and also for the Schiff base formation con

stant, a pit-mapping procedure is used. This method involves the cal

culation of a theoretical titration (pH versus volume of titrant) curve

of a specified solution for a given set of constants. In order to solve

the three multiple-order simultaneous equations involved for each point

of the titration, a computer program utilizing matrix inversion and a

Hewton-Raphson approximation subroutine is employed.

f 2 9

For example, in the Ni(ll)-glycylglycinate amide dissociation

system, the three unknown quantities are a^, (M++) and (A“ ) and the

appropriate mass balance equations are:

1. ■ (H+ )-(0H**)+2(H2A+ )+(HA)-(MD)-(!ffiA“ )-2(MD2“2 )-(MDA2‘ 2 )

-2(MD2A” 3 )

2. A a (H A*)+(HA)+(A“)+(MA+ )+(MD)+2(MDA")+2(MDo"2 )+2(MA0 )+3(MA “ ) T 2 2 <; 3 +3(MDAg*2)+3(MDgA"3)

.3. M « (M+ 2 )+(MA+ )+(MD)+(MAo )+(M0A")+(MD "2 )+(MA “)+(MDA "2)+(MD A“3) T 2 2 3 2 2

These mass balance equations may be expressed in terms of the three

unknowns and the appropriate equilibrium constants and rearranged as

functions of the three unknowns as follows:

Hj H *h (»h . a “ , m + 2 )

5. A^ a A", M+ 2 )

6 . Kp = fA (aH , A”, M+2)

Upon rearrangement these give:

7. Fh « A", M+2 )-H^ * 0

8. Fa = fA(an, A“, M + 2 ) ^ = 0

9 * rM " fM(V A"« “0

Now if one assumes initial values (0 ^ , A“^, and M*2^) for the

unknowns, more accurate values of the unknowns will be equal to the 30 initial guesses plus the appropriate correction terms indicated aB deltas:

10* a ■* a + A a H2 HI HI

H* A g " A i A ^

12. » M+2x + A M+2x

These corrected values (a^2» A~2' and M+2g) may also he improved by a second set of correction terms (Aa , A A" , a n d A M +2 ) to obtain U2 2 2 still more accurate values of the unknowns A~3» and M+2j).

This process may be repeated until the desired accuracy is obtained.

In order to evaluate the above correction terms, equations 10-12 are substituted into equations 7-9* Upon expansion into a Taylor series (neglecting terms which contain cross products or squares of the correction terms) one obtains the following equations:

c /f H1 c/ FH1 13. ^ * H 1 + A A”i U A“ - ^ ® H 1 A", M + 2

+A m +2i *= -F. c/M+2j HI aH,

'c/ FA1 c/ FA1 i*t. A *^1 + A A“i +2 P^H l . A” , M + 2 VM

r^FAi + A m +2i * -F A1 aH ’ A ’ 3 1

a -F Ml

The partial derivatives and the original functions may be evaluated

in a straightforward manner by using the initial values for the unknowns. The above equations (13-15) are therefore reduced to three

simultaneous linear equations with the three unknowns A a , A A*V , * HI* 1 and A m +^ , These are easily solved by means of a preprogrammed matrix inversion subroutine.

The constants to be evaluated are then varied in a systematic vay and the difference between the calculated pH and the experimental pH for each point in a titration curve is evaluated in a least sum-

squares manner. The sum-squares associated vith a particular set

of constants is taken as an indication of the correctness of these

constants. That set of constants which has the smallest sum-squares, '

therefore, is considered to be the "best fit" or the correct set of

constants which most accurately describes the equilibrium.

For the Ni(ll) mixed system, the pH of the solutions used for the

determination of the mixed constants never exceeded a pH of 8,1.

Under these conditions, no ionization of the glycylglycinate amide

proton vas expected; therefore, the masB balance equations for the

Hi(ll) mixed system did not include doubly-ionized glycylglycinate

species. 32

The computer program used for the above pit-mapping procedure is f modified slightly to evaluate the "mixed" constants. For the mixed systems, the experimental quantity which describe the equilibria is the change in pH (ApH) observed upon addition of glyoxalate to a metal-glycylglycinate solution. The program was modified, therefore, to calculate a theoretical A. pH and the sum-squares of the difference between this theoretical and the experimental A pH.

In the case of the mixed systems, the program is designed to vary either one, two, three or all four of the mixed constants

Bjp, B21* 611(1 B22^ upon coauiland» After the initial guesses for the constants have been refined in this manner, the program is adapted to a method by which the constants undergo further refinement by con vergence to a minimum sum-squares. For this final refinement, it is assumed that the sum-squares pit-map surface in the region of the min imum becomes analogous in shape to that of a paraboloid surface.

Accordingly, if the sum-squares are known for sets of constants near the minimum, it is possible to predict the constants which will give the minimum value of sum-squares by calculating the minimum of the paraboloid surface. Also from the shape of the paraboloid surface, the standard deviation of the experimental points and of the formation constants can be calculated. IV. RESULTS AND DISCUSSION

A. Glycylglycine and its Metal Complexes

Figures 5-6 show several titration curves of glycylglycine and

Figure 7 shows the change in the as a function of pH. The undergoes two rapid changes corresponding to the two pH ranges in which the titrations occur. Between the pH values of 5-7* the value of remains almost constant at 1.0. A good approximation to the pK 's may he obtained from Figure 7 by reading the pH at n_ = 0.5 and H By * 1.5. These points of course correspond to the midpoints of the titration.

The values of the acid dissociation constants of glycylglycine vere determined from linear plots of n„ functions as described in the ii Numerical Treatment chapter. These values compare favorably to liter ature values as shown in Table 1.

It is well known that glycylglycine exists in aqueous solutions as the zwitterion, as do all of the amino acids. Upon titration with acidt therefore, the carboxylate group functions as the proton acceptor.

Likewise, upon titration vith base, the terminal amino group functions as the proton donor. It is worth comparing the pKa*s of glycylglycine relative to those of similar compounds (see Table 2).

33 Figure 5

Titration of Glycylglycine with Hydrochloric Acid

Curve 1 — Solution: 3.00x10"3m glycylglycine 0.50M KC1 25ml initial volume

Titrant: 0.1001M HC1

Curve 2 -- Solution: H,00x 1O~^M glycylglycine 0.50M KC1 25ml initial volume

Titrant: 0.099H5N HC1 5.0

4.4

3.8

3.2

0 Q 2 0.4 0.6 0.8 1.0 TITRANT ADDED M) u> V/l Figure 6

Titration of Glycylglycine vith Sodium Hydroxide

Curve 1 Solution: 0.0200M glycylglycine 0.U8M KC1 20ml Initial volume

Titrant: O.O9 8 8 5 N NaOH

Curve 2 — Solution; 0,01001-1 glycylglycine 0,1*9M KC1 90ml initial volume

Titrant: 0,0991*2N NaOH

36 t @ to- j — A

3.0 4.0 5 0 ADDED (ml) U> Figure 7

Formation Curve for Glycylglycine

38 39

O CD

X P-

O LO

O rO

Q IT) o in O OJ C>

IP 1(0

TABLE 1

pK Values Reported for Glycylglycine CL

This work a b c d pK 3.215 2.91 3.12 3.1(9 3.33 la pK£ft 8.232 8.23 8.12 8.27 8.27

This work, 25°C, ionic strength **0.50

ftEvans and Monk^, 25°C, ionic strength » 0

^Gilbert, Otey and Hear on**®, 26°C, ionic strength « 0.03

cSmith59, 25°C, ionic strength *= 0.058M KC1

^Datta and Rabin^, 25°C, ionic strength = 0*02

TABLE 2

pK Values of Glycylglycine and Related Compounds CL

Compound Structure pK^ * PK2a glycylglycine8, ^NCHgCONHCHgCOO" 3 .2 1 5 8.232

glycine* . +H3NCH2C00" 2,3h 9 . 6 0

acetic acid0 CH COOH U.53 3 methylamine^ ^H^NCH^ 10.72

®This work, 25°C, ionic strength =0,50

bKroll61, 25°C, ionic strength * 0,1M KC1

e 6 2 Yasuda, Ysasaki and Ohtaki , 25°C, ionic strength = 0.1

^BJerrum^( 25°C, ionic strength »0.5 hi

The presence of an -NH^* group in glycine exerts an electron- withdrawing effect vhich facilitates the removal of a proton from the carhoxylate group; this is demonstrated by the low pK of glycine as a compared to that of acetic acid. For glycylglycine, this effect is reduced by the substitution of -NHCOCHgNH^* for the -NH^* group; thus, the pKlft of glycylglycine is correspondingly greater than that of glycine, but lower than that of acetic acid. Also, the positive charge (which would tend to lover the pKft) is closer to the carboxylate group in glycine than it is in glycylglycine.

The same type of argument may be used to account for the relative pK *s of the amino groups. The electron-withdrawing effect of the CL carhoxylate group in glycine is sufficient to lower the pK_ in comparison to that of methylamine. The -CONHCHgCOO"’ group of glycylglycine has a greater electron-withdrawing effect, thus lowering the pK0 still more. In this case, the negative charge (which would tend to increase the pK&) is once again closer to the amino group in glycine than it is in glycylglycine.

Figures 8 and 9 show several titration curves of the metal-glycylglycine solutions and Figure 10 shows the change in n as a function L of -log(glycylglycinate). The formation constants of the Ni(ll) and

Zn(ll) complexes were determined from titrations of the metal-glycylglycinate solutions with base. Values for and BQ2 for both Ni(ll) and Zn(ll) complexes are consistent vith literature values as shown in

Table 3. These constants were evaluated from linear plots of n^ functions as described in the Numerical Treatment chapter. The values of B q ^ for both Zn(ll) and Ni(ll) were obtained directly from the Figure 8

Titration of Ni(ll) - Glycylglycine vith NaOH

Curve 1 ~ Solution: 0,800x10"^M glycylglycine OMli KC1 p 0.2105x10 M NiCl2 5 0 ml initial volume

Titrant: 0.09939N NaOH

Curve 2 — - Solution: 1.200xl0'“2M glycylglycine 0,i»9M KC1 0.2105x10 M NiClg 50ml initial volume

Titrant: 0.09939N NaOH

kz 10.0 pH 8.0 2

6.0 J______L 0 2.0 4.0 6.0 8.0 TITRANT ADDED M) Figure 9

Titration of Zn(ll) - Glycylglycine with HaOH

Curve 1 — Solution: 0.0600M glycylglycine 0 A 5 M KC1 0.02108M ZnCl .25ml initial Volume

Titrant: 1.00OH NaOH

Curve 2 — Solution: 0.0600M glycylglycine 0,1*5M KC1 0.012108M ZnCln 2 5 ml initial volume

Titrant: 1.000N NaOH 7.0

5.0

0 0.4 0.8 1.2

TITRANT ADDED (ml) -*r- vn Figure 10

Formation Curves for Metal(ll) - Glycylglycinate Complexes *»T

O LO

o (- 9 9 )

o 1 0 9 rO

O c\3

O c\* kb y-lntercept of the linear plot used to determine BQ2.

TABLE 3

Log Formation Constants Reported for Glycylglycine with Ni(ll) and Zn(ll)

Ni(ll) Zn(ll) B01 B02 B01 B02

This work 3.961 7.12° 3.312 5.91*2

a k.^9 7.91 3 . 8 0 6.59

b 3.93 7.18 - -

c - 7.22 -

d - - 3.6 -

This work, 25°C> ionic strength « 0,50

“Honk61*, 25°Ct ionic strength = 0

^Martin, Chamberlain and Edsall^*1, 25°C, ionic strength - 0.016

°Lyons^5# 25°C, ionic strength = 0.058M KC1

S ’erkins^, 2 5 °Ct ionic strength « 0.01

It can be seen from Figure 10 that both Ni(ll) and Zn(ll) form

complexes with a 3:1 ligand to metal ratio in addition to the 2:1 and

111 complexes, Glycylglycine is potentially a tridentate ligand as shown below: CHi~"r=»0 O ^HrCOO

/ I A / MH* A A / ° H *C \ N - C H j - C O O I 0=C\ /M+ < \ CH^— NHj. HJM— *M+

I I M For a 3:1 complex to form vlthln the octahedral coordination sphere

of Hi(ll) and Zn(ll)t each glycylglycinate molecule could coordinate at only two of its possible sites. As mentioned in the Introduction, the question of how glycylglycine coordinates with metals in solution

Is still unresolved. The configuration which is considered to be the most probable for the complexes in solution is II above. For the

3:1 complex, the metal ion would be coordinated with six nitrogen

atoms (two from each glycylglycinate molecule) and the carboxylate

groups would remain uncoordinated.

Another point of interest regarding the glycylglycine-metal 4 complexes is the acidity of the amide hydrogen. Several authors have

previously investigated the dissociation of the amide proton of complexed glycyl peptides. It has been concluded that the Cu(ll) com

plexes do undergo amide dissociation relatively easy at a pH range li3 lib of 5-6 • Martin, Chamberlain and Edsall have postulated that a sim

ilar dissociation occurs with the Ni(ll) complexes in the pH range of

10-11, Other authors have suggested the possibility of Ni(ll)-

hydroxy species.

The titration curves of the Ni(ll)-glycylglycine solutions

(Figure 8) show conclusively that, in addition to the base required

for neutralization of the glycylglycine amino group, two equivalents

of base per mole of metal ion are taken up by the complexes. This could

be by dissociation of the amide proton or by the formation of a hydroxy

complex. No additional ionization was detected for the Zn(ll)-gly-

cylglycinate solutions. 50

Near-IR and visible spectra of Ni(ll)-glycylglycinate solutions in the pH range of 10-11 have provided more information. The four absorbance bands of Ni(ll) are denoted as iS j and. An enhancement of the 1 /f absorbance relative to that of3£(an increase in the ratio£ y ./S v ) is characteristic of the Schiff base complexes studied in these laboratories and is taken as evidence for the pres- 29 ence of an imine structure in these systems , Table U shows the spectral changes which occur as the pH of a Ni(ll)-glycylglycinate solution is increased from 9-10. The ratioincreases from 1.6 / 3 to 5.^3. It is concludedf therefore, that dissociation of the amide proton does occur in the Hi(II)-glycylglycinate complexes and further more, that these complexes exist in the following configuration in solution*

H2C N-CHrCOO' \ / H i N — Ni+2 TABLE U

Near-IR and Visible Spectra of Ni(ll)-Glycylglycinate

pH ^ ( £ ) 7 $ (£) - ^ ( e )

9 * 1 0 6 0 0 A (5 .2 ) 6 5 0 0 A (3 .2 ) 3900A (6 .6 ) 1 . 6

1 1 1010QA (37.5) 6 0 5 OA (6.9) 3670A (1 6 .7 ) 5.^3 51

The pKa*B for the amide dissociation of the various species in

solution vere determined by a pit-mapping procedure as discussed in

the Numerical Treatment chapter. In this system there are six species vhich can possibly have an amide dissociation:

NiU)+ Kall Ni(D)°

NitA)® Ka21 Ni(DA)"

|jT Ni(DA)“ a22 Ni(D)g”2

Ni(A)3" *a31 Ni(DA2)-2

Ni(DA2 )~2 Ka32 NifDgA)*3

NitDgA)-3 Ka33 NitD)^1*

_ o vhere A represents glycylglycinate and D represents doubly-lonited

glycylglycinate. As a necessary simplification of the mathematical

treatment, several valid assumptions (or approximations) vere made.

First of all, it was assumed that glycylglycine coordinates at the two

nitrogens and not at the carboxylate group; this has previously been

postulated for these complexes• Then it was assumed that K ,,— all K^2^— Neglecting the statistical factor, this is reasonable

Bince for each case, the amide proton would be dissociating from nearly

identical sites within the coordination sphere. For each of these dis

sociations, the coordination sphere ha3 a net charge of +2, as the car

boxylate group (by the first assumption) Is not within the coordination

sphere. Likewise, the same’argument was used to assume that K^gg^

Ka32* the net charge of the coordination sphere 1b +1 for these two 52 cases, A fourth assumption vas that ^ 3 3 is negligible. This is verified experimentally by the fact that the titrations require only two equivalents of base per metal ion in excess of that required for the amino group. « With these assumptions, it vas possible to determine ^ and

*^a22 ^ means a two-dimensional pit-mapping procedure. The values obtained by this method are compared to those reported by Martin,

Chamberlain and Eds all** ^ in Table 5»

TABLE 5

Amide Dissociation Constants of the Ni(ll)-Glycylglycinate Complexes

This work8, Martin et al.fe

pKall 9.W8 9.35

p K ^ 9.62 9.95

°This work, 25°C, ionic strength “ 0.50

^Martin et al.**\ 25°C, ionic strength *= 0.16

For the glycylglycine and the glycylglycine-metal systems, the experimental pH values in nearly all cases fell within 0.01-0,02 pH units of the theoretical curves. The theoretical curves were gen erated by means of a computer program which used the experimentally- determined constants. This same program vas used to calculate the fractions of the various species in solution for several representative titrations (see Appendix B). A summary of the constants determined in this portion of work appears in Table 7 along with the Schiff base and mixed constants. 53

B. Schiff Base and Mixed Systems

Figure 11 shovs the titration curves of glycylglycine with var ious amounts of glyoxalate present. Equilibrium was attained rapidly and solutions below a pH of 9 were stable over a period of several hours as indicated by a monitoring of the pH and the UV spectra. To assure the stability of the solutions even at low pH values, it was necessary to deaerate the solutions thoroughly prior to the addition of baBe. This was accomplished by bubbling nitrogen gas through the solution for approximately five minutes. The os-carbonyl group of glyoxalate condenses with glycylglycine in solution to form a weakly- associated Schiff base. The magnitude of this Schiff base formation constant (Bq A) agrees well with that expected from a comparison of similar Schiff base systems (see Table 6).

TABLE 6

Log Bq a Constants for Glyoxalate-Amino Acid Schiff Bases

& 1) Glycylglycine Glycine

1.778 1.75^ 1.273 1.00

®Thi3 work, 25pC, ionic strength « 0.50

^Leussing and Hanna^, 25°C, ionic strength - 0.50 i

Figure 11

Titration of Glycylglycine - Glyoxalate with NaOH

Curve 1 — Solution; 0.100M glycylglycine 0.0100M sodium glyoxalate 0 #1*8M KC1 10ml Initial volume

Titrant: 0.09882N NaOH

Curve 2 -- Solution: O+OIQOM glycylglycine 0«0300M sodium glyoxalate O.U6M KC1 10ml initial volume

Titrant: 0.09882N NaOH

Curve 3 — Solution; Q*0100M glycylglycine 0a100M sodium glyoxalate 0.39M KC1 10ml initial volume

Titrant: 0.09882N NaOH

Stability proved to be a greater problem when metal ions were present in the Schiff base solutions. As mentioned above, glycylglycinate-glyoxalate solutions are stable for several hours below a pH of 9; the addition of Zn(ll) or Ni(ll) ions, however, caused an initial rapid (within the time of mixing) decrease in pH (due to formation of metal complexes) followed by a more gradual decrease in pH over an extended period of time (U-12 hours).

It has been reported^ that the mixed complexes of glycinate and glyoxalate form rapidly (within the time of mixing for the Zn(ll) system) and are stable for as long as six hours. It is expected that the mixed complexes of glycylglycinate and glyoxalate would form as rapidly as those of its. glycinate analog. Accordingly, the initial decrease in pH upon addition of the third component of the mixed system was taken as an indication of the rapid and complete formation of the mixed complexes.

Figure 12 shows the change in pH as a function of time for several titration points. The erratic behavior following the rapid decrease in the lower curves is a function of the rate of mixing. By the proper extrapolation of the linear portions, the pH of the solu tions, after complexation is complete and before the secondary reactions become appreciable, may be determined to within + 0.01 pH units or within the accuracy of the pH meter. This point-by-point method was therefore used to obtain the data.

Figure 13 shows the change in pH upon addition of glyoxalate to several glycylglycinate-Hi(II) solutions. The standard deviation of

Individual data points for the Hi(II) mixed system is 0.028 pH units. Figure 12

Change in pH with Time for Glyoxalate Addition

to a Glycylglycine - Metal Solution

Curve 1 -- Solution: 0.0200M glycylglycine 0.00398N NaOH 0.00l*03M ZnCl2 0.U7M KC1 20ml initial volume

Titrant: 0.200M sodium glyoxalate 0.30M KC1

Curve 2 — Solution: 0.0200M glycylglycine 0.00998N NaOH 0.00l»056M NiCl 0.1»7M KC1 2 20ml initial volume

Titrant: 0.200M sodium glyoxalate 0.30M KC1

Titrant added: a - 1.00ml b - 2.00ml c - 5«00ml d -10,00ml

57 56

7.2

6.8 2 c.

6.4 I CL

6.0

5.6

0 4 8 12 Time after Addition, of Glyoxalate (minutes) Figure 13

Change in pH Upon Addition of Glyoxalate

to Ni(ll) - Glycylglycinate Solutions

Curve 1 — Solutions 0.0200M glycylglycine 0.00503N NaOH O.OOU056M NiC1 0.U7M KC1 d 20ml initial volume

Titrants 0.200M sodium glyoxalate 0.30M KC1

Curve 2 ~ Solution: 0.0200M glycylglycine 0.01002N NaOH 0.00U056M NiCl 0.U7M KC1 20ml initial volume

Titrant: 0.200M sodium glyoxalate 0.30M KC1

Curve 3 — Solution: 0.0200M glycylglycine 0.01501N NaOH 0.0014056M NiClo 0.H7M KC1 20ml initial volume

Titrant: 0.200M sodium glyoxalate 0.30M KC1

This Is a good fit of the data and that of the Zn(ll) mixed system is better 0*020 pH units). The standard deviation of both systems

is vithin the expected tolerance for this method. Another indication

of a good fit is the small standard deviations of the constants (see

Table 7)* This reflects on the sensitivity of the sum-squares sur

face to the individual constants, that is, on the steepness of the

pit.

The mixed constants for the Ni(ll) and Zn(ll) systems are tabu

lated in Table 7* The magnitude of the mixed constants provides

valuable insight into the nature of the mixed complexes formed by

glycylglycinate and glyoxalate. Figures lb and 15 show the logs of

the stepwise formation constants for the various Hi(ll) and Zn(ll)

complexes. If the two different types of ligands were complexed

independently by the metal ion, one could calculate the theoretical

log (Bi;l) by adding the logs of B and B1Q. For a more realistic

result, a statistical correction must be applied to the constant as

follows;

log (B^) 85 iog (BQ i ) + log (B1Q) - correction

The correction is based oa the fact that if one ligand is already

present, the second ligand will have a reduced number of sites at

which it can become complexed to the same metal ion. For comparative

purposes between similar systems, this statistical effect may be

ignored.

If there is an interaction between the ligands which are coord

inated to the same metal ion, one would expect log (B^) to increase 62

TABLE 7

Formation Constants for Glycylglycinate-Glyoxalate Simple and Mixed Complexes at 25°C and Ionic Strength =0.50

3 .2 1 5 £ o.Ol 5 "logs of glycylglyclne { dissociation constants 8.232 t 0 .0 0 8 * 2 . 9.

M Zn(ll) Hi(ll)

2.05 - 0.02 X 103 9 .1 k ± 0.03 x 103 B oi 8.73 0.08 X 105 1.32 ± 0.02 x 107 B 02 ±

± 0.3 X 107 2.0 £ 0.1* x 109 B03 3 .k

Bioa U.37 8.70

1.7^ i 0.38 X 105 6 . 0 6 - 0 . 6 2 X 10 5 B 1 1

1.33 £ 0.28 X 108 . k . 9 0 £ 0 . 1 2 x 109 B12

108 6 . 9 9 B22 1.15 £ 0.13 X t o.ii* x 1010

2 . 2 1 £ 0 . 2 9 x 105 B21 -

0.0093^35 0 . 0 1 5 6 6

<7~c 0.020 0 . 0 2 7

§ Points 28 2 k

MM ®Leussing and Hanna t 250C, ionic strength = 0.50

*= Sum-squares of the experimental data

*2 7 -' « Standard deviation of the individual data points Figure lU

Log Formation Constants for Stepwise Ccmplexation

in the Hi(ll) - Mixed System

63 r MA

MA;

MA MAoGx <6> Of /b '8 +2 M MAGx M A 2G ^ D* MGx

STEPWISE FORMATION CONSTANTS FOR NICKEL(IT)

LEF = 0.8S -C"ON Figure 15

Log Formation Constants for Stepvise Complexation

in the Zn(ll) - Mixed System

6 5 MA

'<<9 MAt MA?Gx

+2 -2 M MAGx MAjGx^

<£ & MGx

STEPWISE FORMATION CONSTANTS FOR ZINC(H)

LEF= 1.29 CN C\ by an amount corresponding to the extent cf interaction* This in crease is known as the ligand enhancement factor (LEF) and may he represented by the following expression:

LE*! ° log (Bi;l) - log (BQ1) - log (B10)

As it is defined above, the LEF^ is uncorrected for statistical ef fects but is adequate for the purpose of comparison with other sys tems. A statistical correction would increase the kEF^ by a constant amount for similar systems. A second LEF (kEFg) Bay be defined for the addition of a glyoxalate ligand to the 2:1 glycylglycinate-metal complex: LEFg = lOg (Bgi) - log (BQ2) - lOg (B10)

TABLE 8

j LEF^ Values for Glyoxalate Addition to MA Complexes

A" Zn(ll) Ni(ll)

Olycylglycinate 1.29 0.86

Glycinate0, 1*23 l.k9

Of - Alan at ea 1,12 1,31

0 - Aminoisobutyratea 1.79 1.30

0.0 0.0 H2°a

30 aLeussing and Hanna

The LEF^ values for the Ni(ll) and Zn(ll) glycylglycinate-gly oxalate mixed systems are compared with similar systems in Table 8.

The LEF^ for Zn(ll) is of the expected magnitude. The value for Ni(ll) 66

Is slightly less than expected; this is of no serious concern, how ever , as the LEFg value of 1,63 for Ni(ll) is comparable to that expected. The uncertainty in LEF values may be appreciable since they are obtained by the difference of two experimentally-determined constants•

The positive LEF values for Ni(ll) and Zn(ll) are indicative of an Interaction between the complexed ligands. This interaction is assumed to be the formation of a Schiff-base as illustrated below:

? H 1 H N-CHrCOCr H2C N-CHrCOO- N— h-c^ y Ho-ct^y % f Imlne Carbinolamine

The complexed Schiff base may exist as the imine or the carbinolamine.

The extent of carbinolamine is determined primarily by the amount of strain produced in the chelate ring by the > C=JI- group. Formation of the carbinolamine serves to reduce this ring strain.

The extent of carbinolamine formation has been estimated for the

Ni(ll)-glycinate-glyoxalate complexes from the ratio of G - y - / £ y .

If one assumes values of 1.1 and t.7 as normal values of for saturated and imine nitrogen respectively, the Ni(ll)-glycylglycinate-glyoxalate complexes exist as the carbinolamine to the extent of

75#* This is comparable to 7255 reported for the analogous glycinate

complex. Another approach to the investigation of carbinolamine 69 6k formation has employed NMR; Leussing and Stanfield have reported peaks in the NMR spectra of the Zn(ll)~ and Ca(ll)-pyruvate-glycinate complexes which ere due to a small amount of the carbinolamine.

An attempt was made to determine the extent of carbinolamine, if any, for the uncomplexed Schiff base by the use of NMR techniques.

It was found, however, that the NMR spectra of glycylglycinate-glyoxalate solutions were too complicated in the methylene region to be of any practical value for this type of investigation. The glycinate- glyoxalate Schiff base system was chosen as on appropriate substitute; this choice was based on the strong similarities in the structural region of interest. In general, NMR investigations of methylene pro tons in aqueous solutions are difficult because of the proximity of the water peak and the annoyance of spin coupled sidebands associated vith this peak. The use of deuterium oxide (DgO) solutions reduces this difficulty considerably but does not eliminate it completely.

At room temperature (25°C), the NMR spectra of glycinate-glyoxalate solutions have five peaks as shown in Figure 16, The sharp peak at -558cps from TMS corresponds to the glyoxalate methine proton and the peak at -386cps is due to the methylene protons of glycinate. The 68 -olOcps peak is in the frequency range of a proton attached to the carbon of an imine double bond; therefore, the methylene protons of the

Schiff base are assigned to the -li67cps peak. The remaining large peak is due to HDO. Figure 16

NMR Spectrum of Glycylglycine - Glyoxalate

Solution at 25°C

Solvent: DgO

Solution: 1.00M glycylglycine 1.00M sodium glyoxalate 0.80N sodium deuteroxide

Instrument: Vaxian HA-100 NMR Spectrometer

Instrument Conditions:

Freq. Response 5

Output 0.5

Sweep Freq* 0.02

Man. Osc. Freq, 0.08

Spect. Amp. 2000

Sweep Time 500

Sweep Width 500

Sweep Offset -360cps from TMS

External Lock TMS

70 X X M -800 -550 - 5 0 0

CHEMICAL SHIFT vs TMS (c p s ) Figure IT

NMR Spectrum of Glycylglycine - Glyoxalate

Solution at 0°C

Solvent: D^O

Solution: 1.00M glycylglycine 1.00II sodium glyoxalate 0.80N sodium deuteroxide

Instrument: Varian HA-100 NMR Spectrometer

Instrument Conditions:

Freq, Response 5

Output 0.5

Sweep Freq. 0.02

Spect, Amp. 2000

Man. Osc, Freq, 0,08

Sweep Time 500

Sweep Width 500

Sweep Offset -B^Qeps from TMS

External Lock TMS

T2 -800 -550 -500 -450 -400 CHEMICAL SHIFT vs TMS (c?s) -55ficps -306cps

Hydrated Glyoxalate Glycinate

HO H I I "QOC - HC = N - CH„ - COO" "00C - HC - N - CH - COO" t t [ t - 8 1 0 cps -U6Tcps -502cps -U 2 6 cpa

Schiff base (imine) Schiff base (carbinolamine)

Upon cooling the samples to 0°C (Figure 17), the -H67cps and

-8lOcps peaks became leas broad and were more easily observable. In

addition, two new peaks became visible at -li26cps and -502cps. Due

to the spin-coupled sidebands of the HDO peak and to the general

quality of the spectra, the ratio of these two peaks could only be

estimated as 2:1 for the -l42(Jcps and -502cps peaks respectively.

These peaks axe probably due to the carbinolamine form of the Schiff base. A rough estimate of the relative areas of the -^26cps and the

-l*67cps peaks Indicates that about 10-20# of the Schiff base exists

In the hydrated form at 0°C.

The possibility of a diamine species was also considered. A

diamine^2 could be formed at low temperatures by the condensation

of two glycinate molecules.with one glyoxalate. To investigate this ■

possibility, the mathematical equations describing the glycinate-

glyoxalate system were expanded to include a diamine species. The 75

H COO* I I "OOC - CH - N - CH - NH - CH„ -COO" 2 2

Diamine subsequent evaluation of the formation constant for this species,

along vith the Schiff base formation constant (B ), vas handled by VA means of a two-dimensional pit-mapping procedure. A minimum sura-

squares was obtained only by setting the diamine formation constant \ equal to zero; this was taken as evidence tyat no diamine species was

formed.

C. Schiff Base Reactions

Solutions containing glycylglycinate and glyoxalate, and solu

tions containing glycinate and glyoxalate both exhibit an absorbance

shoulder in the UV spectra which begins at ^ 320rau and merges with

the carboxylate end absorbance. Reactions of the Schiff base systems may be studied by observing the changes which occur in their spectra.

For the glycinate-glyoxalate system at a pH of > 10, this

shoulder absorbance decreases to less than 20% of the initial absorb

ance in two days (see Figure 18). The decrease is probably due to an

aldol condensation of glyoxalate at the methylene carbon of glycinate

to form/? -hydroxyaspartate. This type of condensation may be cat

alyzed by glyoxalate through the intermediate formation of a Schiff

base followed by the aldol condensation reaction with a second mole- 6 19 cule of glyoxalate. Several people • have reported this type of

reaction in glycinate-glyoxalate solutions under various conditions Figure 18

UV Spectra of Glycinate - Glyoxalate Solution

Solution: 0.10M glycine 0.020M sodium glyoxalate 0.12N NaOH 1.0cm cell

Curve 1 — 5 min. after mixing

Curve 2 — 20 hrs. after mixing

Curve 3 — 2 days after mixing

Curve U — 3 days after mixing

Curve 5 — U days after mixing

Curve 6 — Theoretical spectra assuming no Schiff-base formation

T6 2S0 300 340 WAVELENGTH (rR/U.) and have Isolated end identified the resulting product as the postu lated $ -hydroxyaspartate• Transamination probably also occurs in this system, but would result in no net change in the chemical com position of the solution. The change in UV spectra is, therefore, due to the aldol condensation.

Ultraviolet spectra of glycylglycinate-glyoxalate solutions at a pH > 9 likewise undergo changes. As with glycinate, the shoulder due to the absorbance of the Schiff base decreases; however, the develop ment of several new absorbance bands between 2h0mu and 360mu accompanies

the secondary reactions. From the results of the glycinate-glyoxalate system, one might expect the same types of reactions, that is, aldol condensation and transamination.

Figure 20 shows the spectral changes observed in a glycylglycinate- glyoxalate solution containing an excess of glycylglycinate. The initial shoulder absorbance decreased with time but was accompanied by the appearance of two new bands, one at 265mu and one at 330mu.

These are thought to arise from the transominated Schiff base (see

Figure 19)• The chromophoricity of these species results from the

0 I! "00C - CHg - N « CH - C - NH - CHg - COO" Keto form

”00 C - CH2 - N*=CH-C = N- CHg - COO" Enol form

Figure 19* Transominated Schiff Base Figure 20

UV Spectra of Glyeylglycinate - Glyoxalate Solution

Solution: 0.050M glycylglycine 0.0050M sodium glyoxalate 0.060N NaOH 1.0cm cell

Curve 1 — upon mixing

Curve 2 — 2 hrs. after mixing

Curve 3 — 5.5 hrs. after mixing

Curve h — 8 hrs, after mixing

Curve 5 — 2U hrs. after mixing

Curve 6 -- lij.hrs. after mixing

79 Bo.

1.4

0.6

0.2

0.0 260 3 0 0 3 4 0 WAVELENGTH [lU/*} 81 conjugated double bonds. Both of these bands are probably due to

*jf—\*ffm transitions.

It vas noted previously that glyoxalate condenses with the methylene carbon of glycinate (via the Schiff base); likewise, an aldol condensation reaction Is expected to occur with glycylglycinate.

In solutions where the free glyoxalate concentration is relatively high, the UV spectral behavior is different from that shown in

Figure 20 for the transaminated Schiff base. A high concentration of glyoxalate would favor the aldol condensation; the spectra in Figure 21 should, therefore, correspond to the product of the aldol condensation, or 3 -hydroxyaspartylglycinate:

HO H N 0 i 2i ii ”00C -CH-C-C-N- CHg - COO"

From the structure of this species, it becomes apparent that it cannot have an absorbance bandat ^ 285mu, A logical alternative is that this species has undergone transamination with glyoxalate to give the corresponding Schiff base:

"00C I HOCH 0 H I H I "ooc - ciig - n = c - c - n - ch2 - coo”

This species possesses the same type of chromophoric grouping a3 the ' transaminated Schiff base of Figure 19 and could possibly account for the ~285rau band. Figure 21

UV Spectra of Glycylglycinate - Glyoxalate Solution

Solution: O.I4 OM glycylglycine 0,1*01-1 sodium glyoxalate 0.1»ON NaOH 0,1cm cell

Curve 1 — 1 hr, after mixing

Curve 2 2 hrs, after mixing

Curve 3 — 5 hrs, after mixing

Curve U — 7 hrs, after mixing

81*

The spectral behavior of glycylglycinate-glyoxalate solutions at ■ alkaline pH values can be explained satisfactorily by the two com peting reactions described above. When free glyoxalate is present in solution to an appreciable extent, the aldol condensation is the pre ferred reaction followed by transamination. If there is not a suffi cient amount of free glyoxalate in solution, the transamination reaction predominates.

In the Zn(ll)-giycylglycinate-glyoxalate solutions used for the determination of the mixed formation constants, the pH decreased over a period of several hours after the initial rapid change in pH. It has been assumed that the initial change in pH is due to a rapid formation of the Schiff base complexes as vas found to be the case vlth glycinate-glyoxalate-metal solutions. The mixed formation con stants calculated on this basis are of the magnitude expected and therefore the assumption i3 apparently valid, A reaction (or reactions) which occurs subsequently to Schiff base formation must then cause the gradual changes in pH which are observed.

Several possible reactions, including the previously mentioned aldol condensation and transamination, could cause a decrease in pH.

A look at the UV spectra of these solutions (see Figure 23) reveals that here also the absorption barids at 2 6 0 mu and 330mu develope along with the decrease in the Schiff base shoulder absorbance. After l*-5 days, however,- the 260mu and 330mu bands also decrease. This is in accord with the decrease in pH followed by on increase in pH,

For the transamination reaction, glycine and glyoxalylglycine would be the products expected. A check for glycine by thin-layer chromatography did verify the presence of glycine; however, glycine could also he produced by hydrolysis of glycylglycinate or by the transamination of the aldol condensation product of glyoxalate and llQ glycylglycinate. The procedure used by Dixon to detect the alde hyde produced in a series of transamination reactions was used to determine whether glyoxalylglycine was produced in these glycylglycinate-glyoxalate-Zn(ll) solutions. Although a good separation was not achieved on the column used, a second aldehyde was detected in the eame relative position of eluent as that reported by Dixon^^ for glyoxalylglycine. It is probable, therefore, that the gradual de crease in pH is due to this transamination reaction.

The increase in pH which follows the transamination reaction is accompanied by the evolution of carbon dioxide (COg). A semi-quanti tative (see the Procedure chapter) determination of the COg produced

as a function of time is shown in Table 9. This is compared with the pH as a function of time. It is apparent that COg evolution accompanies the reaction which causes the increase in pH; calculations verify this

correlation.

Glyoxamide is a relatively unstable species; the analogous

• HC - C - NH0 + H O 2 2 ■> HgCO + COg + NH Glyoxamide

glyoxalylglycine would likewise be expected to decompose into CO2 , formaldehyde and glycinate. • This provides additional evidence for the proposed transamination of the Zn(II)complex, of the Schiff base to 8 6

TABLE 9

Correlation of COg Evolution with Changes in pH for a Glycylglycinate-Glyoxalate-Zn( II) Solution3,

time after mmole of C0g £H preparation (hr.) evolved

0 0.000 5.19

12 0.002 li.78

2 k 0 . 0 0 9 U.78

36 0 . 0 1 7 U.82

U8 0.025 1*.91

S o l u t i o n contains 0.025M glycylglycine, 0.05014 ZnClg, 0.050M sodium glyoxalate and 0.010N NaOH. Total volume » 20.0ml.

0 0 H II li I HC - C - N - CHg - COO" + HgO ------* HgCO + C0g + HgN - CHgCOO"

Glyoxalylglycinate ■ produce glycinate and glyoxalylglycinate.

Figures 22 and 23 show the effect that Zn(ll) ions have upon the tvo reactions vhich occur in the glycinate-and glycylglycinate-glyoxalate systems. In Figure 22, the reaction vhich causes the decrease in the' Schiff base shoulder absorbance is the formation of /3 -hydroxyaspartate by the aldol condensation. Transamination in this case vould cause no change in the spectra or in the chemical composition. The aldol condensation reaction is 75# completed vithin one hour as compared to one day (Figure 20) for the same solution in the absence of Zn(ll) ions. Figure 22

Effect of Zn(ll) ions upon the UV Spectra

of a Glycinate - Glyoxalate Solution

Solution: 0.10M Glycine 0.010M sodium glyoxalate 0.0010M ZnCl 0.120N NaOK 1.0cm cell

Curve 1 — Upon mixing

Curve 2 — 1 hr, after mixing

87 8 8

1.4

0 A

0.6

0.2

0 260 300 340

WAVELENGTH {t^M) Figure 23

Effect of Zn(ll) ions upon the UV Spectra

of a Glycylglycinate •* Glyoxalate Solution

Solution: Q.050M glycylglycine 0.0050M sodium glyoxalate 0.00050M ZnCl 0,060N HaOU d 1,0cm cell

Curve 1 — Upon mixing

Curve 2 — 20 min, after mixing

Curve 3 — 2 hrs. after mixing

Curve H — .20 hrs, after mixing

Curve 5 -- hrs, after mixing

89 90

1.4

i »

i/o A

o .s

0.2

0.0 240 260300 340

WAVELENGTH (wjtt) 9 1

Figure 23 shows the effect of Zn(ll) ions upon the glycylglycinate-gly oxalate system. Although the initial decrease of the shoulder absorbance occurs at a faster rate in the presence of Zn(ll) ions than in the absence of Zn(ll) ions (Figure 20), the appearance of the

260mu and 330mu peaks of the transaminated Schiff base is actually slower in the presence of Zn(ll) ions. This is due to the fact that the Zn(ll) complex of the transaminated Schiff base would exist, to a great extent, as the carbinolamine form.

It 1b interesting to note that Ni(ll) ions appear to have a negative effect upon the transamination reaction of glycylglycine, but at the same time, a positive effect upon the aldol condensation reaction. Figure 2k shows the UV spectra of a Ni(ll)-glycylglycinate- glyoxalate solution three days after the solution had been prepared.

The only absorbance peak vhich has developed is one at 270mu. In

addition, the color of the Ni(ll) solution changed from deep-blue to

a pink in 2-3 days upon the appearance of this 270mu peak, and finally

to a deep-yellow color in 7-8 days. It would appear that the string

ent octahedral configuration of the Ni(ll) ion causes a significant

emount of strain in the Schiff base complexes. It was noted previ

ously that the Ni(ll)-Schiff base complexes exist predominantly as

the carbinolamine, which is another indication of steric constraints

imposed by the octahedral configuration. To reduce this strain, the

Hi(II) changes from an octahedral to a square-planar configuration.

To further reduce the strain, the transaminated Schiff base of the

aldol condensation reaction may dissociate to glycine and / 3 -hydroxy-

aspartylglyeinate in its complexed state. This would also account Figure 2U

Effect of Ni(ll) ions Upon the UV Spectra of a Glycylglycinate - Glyoxalate Solution

Three Days After Mixing

Solution: O.OiiOM glycylglycine O.OllOM sodium glyoxalate 0.010M NiCl2 0.030N NaOH 0.1cm cell

92 93

1.4

0.6

0.2

0.0 260300 340 WAVELENGTH (771^] 9h for the position of the peak at 270mu as compared to the previously reported peak at 285mu.

As discussed previously, the portion of the NMR spectra of glycylglycine in the region of the vater peak vas of little use; however» the NMR spectra of glycylglycine do provide some useful

Information in the frequency range of its methylene protons as shown toy the literature. Mathur and Martin^0 have investigated the effects of charge on the chemical shift of a series of glycyl peptides. 71 Takeda and Jardetzky have also reported on the NMR of simple amino acids and dipeptides in aqueous solution. The nonequivalence of the methylene protons of some glycyl dipeptides has toeen reported toy 72 73-76 Morlino and Martin' • Sheintolatt , in a series of works, has investigated the protolysis kinetics of amino and amide protons for glycine and glycyl peptides. The chemical shift of the Of~CH protons of amino acids with pH has toeen shown to toe related linearly to the degree of dissociation of the group attached to it*^.

The linear relationship between chemical shift and the degree of dissociation of the amino group has toeen verified in this work for glycine and glycylglycine. Figure 25 shows the chemical shift of the methylene protons of glycine as a function of Hy. The chemical shift of the methylene protons adjacent to the amino group of glycylglycine

(the & -methylene protons) as a function of n is shown in Figure 26. H Figure 27 shows the effect of Zn(ll) upon the NMR spectra of glycylglycine. Glycylglycine has two peaks (Figure 27, spectrum l), one at -36cps from TMAC (tetramethylammonium chloride) for the

Cf -methylene protons and one at -llcps for the /3 -methylene protons. Figure 25

Linear Relationship Between the NMR Chemical Shift

of the Methylene Protons of Glycine

and the n^ Function

Instruments Varian HA-100 NMR Spectrometer

95 Chemical Shift vg TMS 390 430 410 0 l O 0.3 H n 0.5 0.7 0.9 VO a\ Figure 26

Linear Relationship Between the NMR Chemical Shift

of the /£? -Methylene Protons of Glycylglycine

and the n Function H

Instrument: Varian HA-100 NMR Spectrometer

97 Chemical Shift vs TMS

o CJ

e 6 Figure 2T

NMR Spectra of Glycylglycinate - Zn(ll)

Solutions at 25°C

Solvent* Hg0

Solution: 1.0M glycylglycine 1.2N NaOH

Spectrum 1 - O.OOM ZnClg (2 hours after prep.)

Spectrum 2 - 0.05M ZnCl2 (15-30 minutes after prep.)

Spectrum 3 - 0.10M ZnClg (15-30 minutes after prep.)

Spectrum ^ - 0.20M ZnCl2 (15-30 minutes after prep.)

Spectrum 5 - 0.30M ZnCl2 (15-30 minutes after prep.)

Spectrum 6 - O.^OM ZnCl2 (15-30 minutes after prep.)

Instrument: Vorian A-60 RMR Spectrometer 1 Instrument Conditions:

Filter bandwidth 2cps R. F. Field 0.l6mG Spectrum Amp. 8 Sweep time 250 nee. Sweep width 250eps

S9 0 52- 05- o S 2 - OS1* 1 i __ L i _L __ Li. _L1

9*ou

os- OS'-

gx?u

$2- 0S~ 1______Li

OOT 101

The third peak at Ocps from TMAC is due to a small amount of glyclnate produced hy the hydrolysis of glycylglycine« Addition of Zn(ll) ions

0 II H N - CH - C - HH - CH - COO H N - CH - COO" T /* / -Heps -36cps Ocps

(Figure 27, spectra 2-6).causes both glycylglyeinate peaks to broaden as a result of ligand exchange.

In spectra H-5 of Figure 27, the rate of ligand exchange is sufficiently slow so that the $ -methylene protons adjacent to the amino group appear as two peaks, one for the complexed and one for the uncomplexed ligand.. In spectrum 6 of Figure 27, essentially all of the glycylglyeinate exists as the Zn(lX) complex; in addition, some Zn(0H)g precipitate had formed. t It Ib interesting to note that the -36cps peak (due to the

Heps upon complexation (the -llcps peak shifts downfield by 13cps).

This would indicate that the electron environment of the

In these complexes.

The addition of glyoxalate to glycylglyeinate is shown in •

Figure 28, From these Bpectra it can be seen that a new peak appears

■^lcps downfield (at -37cps) from the -36cps peak of glycylglyeinate, Figure 28 NMR Spectra of Glycylglyeinate - Glyoxalate Solutions at 25°C

Solvent* HgO

Solutions 1.00M glycylglycine 1.20N NaOH

Spectrum 1 - 0.0M sodium glyoxalate (7 days after prep.)

Spectrum 2 - 0.25M sodium glyoxalate (1-2 hrs. after prep.)

Spectrum 3 - 0.50M sodium glyoxalate (1-2 hrs. after prep.)

Spectrum b - 1.00M sodium glyoxalate (l-2 hrs. after prep.)

Instrument: Varian A-60 NMR Spectrometer

Instrument Conditions:

Filter Bandwidth 2cps R.F. Field 0.2mG Spectxmm Amp. 6.3 Sweep time 250 sec. Sweep width 250cps

102 103

710.1 no.3

T -so o

o 10l* and that a series of small peaks appear in spectra 3-^4 of Figure 28.

The new peak at -37cps is probably due to the (X -methylene protons of a Schiff base species; however, the absence of any additional peaks • precludes a more definite explanation. The series of small peaks which appear in spectra 3-H of Figure 28 are probably due to the product of the aldol condensation reaction. A similar series of peaks vas found in the NMR spectra of glycinate-glyoxalate solutions

(see Figure 31) for which the aldol condensation product is/3-hydroxyaspartate.

HO NH0 0 HO NH„ I I 2 ii I I 2 ”00C - HC - CH - C - NH - CHg - COO” “00C - HC - CH - C00“

$ -hydroxyaspartylglycinate / 3 -hydroxyaspartate

Figure 29 shows the spectra of glycylglycine-glyoxalate solutions at various pH values shortly (l-2 hours) after preparation and one week after preparation. Spectrum la of Figure 29 is essentially unchanged after one week (spectrum lb) and there is no indication of

Schiff base formation at this pH (pH » 6.0). The / 3 -methylene peak of glycylglycine overlaps the downfield half of th^cX -methylene

7 I1 doublet. Sbeinblatt has shown by spin-decoupling experiments that the (X -methylene protons are spin-coupled to the amide proton. At higher pH values, the amide proton*s rate of exchange increases and this doublet collapses into a singlet.

At a pH of 8.0 (Figure 29, spectrum 2a), there are still only two peaks present; a comparison of the relative areas of the peaks, however, indicates that the C( -methylene peak of glycylglycine is Figure 29

HMR Spectra of Glycylglycine - Glyoxalate

Solutions at 25°C as a Function

of pH and Tine

Solvent: Hg0

Solution: 1.0M glycylglycine 1.0M sodium glyoxalate

Spectra 1 -pH » 6.0 a. 1-2 hours after prep, h, 1 week after prep.

Spectra 2 -pH = 8.0 a. 1-2 hours after prep., h. 1 veek after prep,

Spectra 3 -pH = 10,0 a, 1-2 hours after prep. b. 1 veek after prep.

Instrument: Vorian A-60 NMR Spectrometer

Instrument Conditions:

Filter Bandwidth lcps R. F, Field 0.16-0.20mG Spectrum Amp. Varied Sweep time 250 sec. Sweep vidth 250cps

105 106

•4 '0

^ 1 i'^1 V 1

1 — ~r ~T~ ■so - z s o -so

no. 2 b.

T -&S ■3c - z s O

no.3b.

■ -5 0 o s o ■ zs 107 overlapping another peak (or peaks). After a veek (Figure 29, spectrum 2b) it can be seen that tvo other peaks have now become resolved from the C( -methylene peak of glycylglycine. Also, glycine is present as shovn by the peak at -l8cps.

At a pH of 10,0 (Figure 29, spectrum 3a), the peak at -37cps which has previously been assigned to the < X -methylene protons of a

Schiff base species has appeared as a shoulder on the OC-methylene peak of glycylglycine. In addition, the same series of peaks which have been attributed to an aldol condensation product are easily observable. After a veek (Figure 29, spectrum 3b),'the cluster of three peaks a t ^ -36eps have become resolved, the glycine peak has developed, and the peaks due to the aldol condensation reaction have grown. Once again, positive assignment of the peaks a t ^ -36cps to specific species is difficult because of the numerous possible species in solution* It is probable that all three of these peaks are due to

OJ-methylene protons of N-glycinate derivatives (X - NH - CH^ - COO").

The NMR spectra of glycylglycinate-glyoxalate-Zn(ll) solutions are shown In Figure 30, The solutions were prepared such that only

80Jt of the glycylglycine was neutralized to glycylglyeinate. All of the spectra in Figure 30 show the presence of glycine (-itaps) which is formed by the transamination reaction; hydrlysis of the glycylglycinate should not occur in these solutions. The glycylglyeinate peaks undergo broadening as before; also, the new peak at -37cps which appeared in Figure 28 broadens, indicating that this species may undergo ligand, exchange. Figure 30

NMR Spectra of Glycylglyeinate - Glyoxalate - Zn(ll)

Solutions at 25°C

Solvents HgO

Solutions 1.0M glycylglycine 0.5M sodium glyoxalate 0.8N NaOH

Spectrum 1 -0.25M ZnCl2 (l day after prep.)

Spectrum 2 -0.50M ZnClg (l day after prep.)

Spectrum 3 -0.75M ZnCl2 (1 day after prep.)

Spectrum U -1.00M ZnCl2 (l day after prep.)

Instruments Varian A-60 NMR Spectrometer

Instrument Conditions:

Filter Bandwidth 2cps R. F. Field 0.l6mG Specturm Amp. 8 Sweep time 250 sec. Sweep width lOOcps 109

T\0.( m 3

i I “I i - 5 b -zs o --SO -as'- o

T\a2 tlo4

-*i— __J— ~T~ ~I— n r -S'O -as* O -•SO -as* O Figure 31

NMR Spectrum of Glycinate - Glyoxalate

Solution at 25°C

Solvent: HgO

Solution: 1.0M glycine 1.0M sodium glyoxalate 1.2N NaOH

Instrument: Varian A-60 NMR Spectrometer

Instrument Conditions:

Filter Bandwidth 2cps R. F. Field 0.2mG Spectrum.Amp. 6.3 Sweep time 250 sec. Sweep width lOOcps

110 "1

- 2,0 TMA H P These NMR spectra of glycylglyeinate systems (Figures 27-30) are insufficient for the purpose of positive identification of specific

Schiff "base species; howeverf the methylene region of the NMR spectra has provided some information on the Schiff hase reactions* The presence of glycine in Figure 30 confirms the previous thin-layer chromatography results and verifies the Schiff base transamination.

The series of peaks in Figure 28 (spectra 3-U) and Figure 29 (spectra

3a-b) and the similar series of peaks in Figure 31 for the glycinate- glyoxalate solution verify that a second reaction occurs in the presence of sufficient glyoxalate and that this reaction is probably an aldol condensation. Table 10 contains a summary of the equilibrium constants determined in conjunction with the NMR experiments.

TABLE 10

Equilibrium Constants0, for Glycine Determined in Conjunction with NMR Experiments

pK2a (at 25°C) 9.7V2 .

pK2a (at 0°C) 10.1*93

pK (at 25°C in DO) 10.36° £■& b 0 ro —j

(at 0°C) + 1 OA

a Ionic strength = 1,0 1 1 3 D. Summary of Results

The simple constdnts of the glycylglycinate-Zn(ll) and -Ni(ll) systems vere evaluated from potentiometric titrations "by linear plots of 5^ and ^ functions. It was verified that the Hi(II)-gly cylglycinate complexes undergo dissociation of the amide proton; an imine structure vas proposed for the doubly-ionized glycylglyeinate based upon the near-IR and visible spectra of the Ni(ll) complexes.

The mixed constants of the Ni(ll)- and Zn(ll)-glycylglycinate- glyoxalate Schiff base systems vere evaluated from potentiometric data by means of a least-sum-squares pit-mapping procedure. The complexed Schiff base aas found to exist predominantly es the carbinolamine; whereas, the uncomplexed Schiff base of glycinate and glyoxalate va3 found by NMR techniques to exist as the imine at room temperatures. Some carbinolamine vas detected for the glycinate- glyoxalate Schiff base at 0°C.

Both the glycylglycine- and glycinate-glyoxalate Schiff base systems vere found to undergo an aldol condensation with glyoxalate in alkaline solutions. The presence of Zn(ll) ions accelerated this reaction. In addition, glycylglycine transaminates with glyoxalate in alkaline solution and at lower pH values in the presence of Zn(ll) ions. Nickel(ll) ions favor the aldol condensation reaction. NMR experiments have verified the transamination reaction and the aldol condensation reaction of the Zn(ll)-glycylglycinate-glyoxalate system. APPENDIX A

POTENTIOMETRIC TITRATION DATA

AND CALCULATIONS ■ 115

Volumes as written have been corrected for buret calibrations.

The quantities used for n^ and n^ linear plots are defined as follow:

(2 - H )(a )2 (n - l)(a ) Y1 » ----- 5-- 2— XI = — S— ---- S- nH 5H

(l — IL.) Hri Y2 = — X2 ------

Sl (l - n^) Y3 * ------— X3 = (2 - u )(A") (2 - n )(A” ) L ' b

Qy + (nL - 1)B01(A“ ) {2 - n ) yl» a Xl» = — (3 - nL )(A")2 (3 - nL )(A~) 116

K and of Glycylglycine at 25°C

Solution: 0.00300 M Glycylglycine 0-.50 M KCl 25 ml initial volume ml titrant pH Y1 XI SH added

Run A (Titrant: 0.10015 N HCl) k 0.120 4.074 1.128- 5.500 x 10*9 9.557 X 10'6 0.200 3.810* 1.207- 1.575 x 10-8 2.660 X 10"5 0.280 3.628- 1.283 3.101 • 5.191

0.360 3.484* 1.353- 5.144 8.566 / 0.440 3.362 1.418 7.743 1.282 X 10_A 0.520 3.255 1.477 1.094 x 10"7 1.796 0.600 3.160 1.530 1.470- 2.397

Run B (Titrant: 0.10015 N HCl)

0.160 3.932 1.169 9.730 x 10"9 1.687 X io“5 0.240 3.722 1.247 2.171 3.759

0.320 3.562 1.321 3.863 6.663 § 0.400 3.430 1.390 6.063 1.042 X 10" 0.480 3.319 1.454 8.646 1.498 0.560 3.219 1.511 1.179 x 10"7 2.044 0.640 3.130 1.564 1.534 2.672

Run C (Titrant: 0.10015 N HCl)

0.120 4.093 1.129 5.025 x 10"9 9.236 X 10" 6 0.200 3.819 1.208 1.507 x 10"8 2.618 X 10"5 0.280 3.634 1.284 3.008 5.138

0.360 3.487 1.354 5.063 8.523 i 0.440 3.366 1.420 7.573 1.273 X 10"4 0.520 3.258 1.479 1.074 x 10’7 1.787 0.600 3.163 1.532 1.442 2.386 117

Kia and K£a of Glycylglycine at 25°C (cont.)

Solution: 0.00400 M Glycylglycine 0.50 M KC1 25 ml initial volume ml titrant pH ?1 XI “H added

Run A (Titrant: 0.09940 N HCl)

0.200 3.927 1.165 1.004 x 10‘8 1.673 X 10'5 0.280 3.754 1.227 1.S54 x 10" 8 3.263 X 10"5 0.360 3.616 1.287 3.243 5.406 0.440 3.498 1.345 4.917 8.145 0.523 3.391 1.401 7.063 1.163 X 10“4 0.600 3.302 ] .450 9.439 1.548 0.680 3.218 1.498 1.229 x 10“ 7 2.012 0.760 3.140 1.542 1,561 2.545 0.840 3.069 1.582 1.921 3.139

Run B (Titrant: 0.09940 N HCl)

0.240 3.826 1.195 1.500 x 10"8 2.440 X 10“5 0.321 3.671 1.257 2.688 4.364 0.401 3.546 1.316 4.207 6.823 0.480 3.438 1.371 6.109 9.864 0.560 3.341 1.423 8.431 ■j 1.356 X 10 4 0.640 3.253 1.472 1.113 x 10 7 1.791 0.720 3.172 1.517 1.441 2.294 0.800 3.100 1.560 1.778 2.852 0.881 3.028 1.598 • 2.214 3.507

Run C (Titrant: 0.09940 N HCl)

0.160 4.048 1.133 6.131 x 1°'* 1.053 X 10“5 0.240 3.837 1.197 1.423 x 10“8 2.390 0.321 3.682 1.259 2.547 4.275 Q.400 3.556 1.317 4.009 6.686 0.480 3.447 1.373 5.831 9.703 t. 0.560 3.350 1.426 8.035 1.334 X 10"4 0.640 3.261 1.475 1.070 x 10"7 1.766 0.720 3.181 1.521 1.367 2.259 0.800 3.106 1.563 1.713 2.824 0.880 3.037 1.602 2.093 3.452 118

K^a and K^a of Glycylglyclrie at 25°C (cont.)

Solution: 0.01 M Glycylglycine 0.50 M KCl 50 ml initial volume ml titrant pH Du Y2 X2 added

Run A (Titrant: 0.09942 N NaOH)

0.847 7.523 0.832 7.913 X 1014 4.805 X 106 1.047 7.640 0.792 1.249 X 1015 7.521 1.446 7.829 0.712 2.518 1.506 X 10 7 1.846 7.987 0.633 4.362 2.605 2.244 8.140 0.554 7.299 4.258 2.642 8.282 0.475 1.141 X 1016 6.592 3.040 8.427 0.396 1.763 1.007 X 10 8 3.540 8.604 0.296 2.809 1.659 4.040 8.867 0.197 5.932 3.278

Run B (Titrant: 0.09942 N NaOH)

0.409 7.157 0.919 1.751 X 1014 1.080 X 106 0.648 7.381 0.871 4.461 2.744 1.197 7.709 0.762 1.612 X 1015 9.836 1.796 7.958 0.643 3.905 2.388 X 10 7 2.393 8.180 0.524 8.139 1 A 4.879 2.990 8.392 0.406 1.547 X 1016 9.193 3.590 8.626 0.287 2.987 1.560 X 108 4.240 8.956 0.158 6.988 4.132

Run C (Titrant: 0.09936 N NaOH)

0.847 7.512 0.832 7.523 X 10141 C 4.683 X 106 1.247 7.726 0.752 1.707 X 1015 1.057 X 10 7 1.646 7.894 0.673 3.112 1.931 2.045 8.049 0.594 5.290 3.234 2.. 39 8 8.173 0.524 7.866 4.806 2.741 8.295 0.455 1.147 X 1016 6.954 2.940 8.367 0.416 1.423 8.584 8.441 0.376 1.766 ---- 1.060 X 1083.140 3.340 8.518 0.336 2.198 1.315 119

K^fl and K2a of Glycylglycine at 25°C (cone,)

Solution: 0,02 M Glycylglycine 0.50 M KCl 20 ml initial volume ll titrant PH Y2 X2 h added

Run A (Titrant: 0.09885 N NaOH)

0.598 7.452 0.852 5.952 x 1014 3.645 x 10! 0.997 7.727 0.754 1.720 x 1015 1.054 x 10- 1.396 7.936 0.655 3.627 2.213 1.796 8.123 0.556 6.788 4.080 2.194 8.299 0.458 1.177 x 1016 6.998 2.592 8.481 0.360 2.008 1.182 x 10 8 2.990 8.688 0.261 3.572 2.071 3.390 8.957 0.163 7.262 4.128

Run B (Titrant: 0.09885 N NaOH)

0.399 7.253 0.901 2.631 x 1014 1.607 x 0.797 7.605 0.803 1.088 x 1015 6.626 1.197 7.840 0.704 2.601 1.579 x 10 7 1.596 8.035 0.606 5.103 3.066 1.995 8.213 0.507 9.058 5.392 2.394 8.391 0.408 1.554 x 1016 9.144 2.790 8.581 0.311 2.671 1.555 x 108 3.190 8.813 0.212 5.010 2.865 3.590 9.152 0.113 1.211 x 1017 6.669 120

Derivation of XI, Yl, X2 and Y2

The constants K, and K_ are defined as: la 2a

(aUHA) - (aH)(A-) K B — i±------K b 11 I . lR (HgA+ ) 2 a (HA) vhere (A") represents glycylglycinate. The mass balance equations for this system may be vritten as follows vhere the subscript t denotes the total concentration of a substance in solution:

H. ** (H+ ) - (0H“ ) + 2 ( H A + ) + (HA) « 2

At «* (HgA+ ) + (HA) + (A")

The value of n„ is evaluated from the known quantities in these n mass balance equations.

2(HCA+) + (HA) H. - (H+) + (OH") n. b ■ B ■ 1 ■ 11 11 ■■■ — — ^ H2A+) + (HA) + (A") At

The quantities (H+ ) and (OH“ ) represent the hydrogen ion and hydroxyl ion concentrations respectively. Substitution for (HgA+ ) and (HA) in terms of (A“ ) in the above definition of n^ yields:

2(aH)2(A“) + ( a ^ K ^ A ”)

“ < ^ ) a(A-) + (aH)Kla(A-) + KlaKga(A-). 121

Derivation of XI, Yl, X2 and Y2 (cont.)

Elimination of (A~) from this equation and rearrangement gives:

(a«) iL » (1 - rL) ---- + (2-n )------*• . “ V U V#> 2a la 2a

For values of ii = 1.1-1,9, this can he arranged into the form of a H linear equation (Yl = n^Xl + t^), where

(2 - HuJfa^)2 ‘ (rL - X){au) n ---- 2 _ 2 _ xi =--2---- 2 _ nH “h

“i ‘ Kla bl " KlaK2a

For values of rLj » 0.1-0,9, this can he arranged into the form

Y2 « ®gX2 + hg, where

(i - 5jj) nH Y2 « ------2— X2 <2 - y < V (2 - nH )(aH)2

»2 - 1/K2a b2 = l/(KlaK2a) 122

Formation Constants for Zn(II) - Glycylglycinate Complexes at 25°C

Solution: 0.0600 M Glycylglycine 0.01009 M ZnCl2 0.41 M KCl 40.0 ml initial volume

Titrant: 1.0002 N NaOH ml titrant pH Y3 X3 5l added

0.000 5.047 0.079 2.862 x 10* 1.263 x 10* 0.100 5.643 0.254 6.805 x 106 2.923 x 103 0.150 5.837 0.362 4.417 1.740 0.200 5.989 0.474 3.217 1.110 0.250 6.134 0.584 2.295 6.939 x 10z 0.299 6.257 0.691 1.749 4.301 0.349 6.368 0.799 1.391 2.417 0.399 6.473 ■ 0.906 1.122 1.005

—

ml titrant PH Y4 X4 nL added

0.548 6.760 1.205 2.856 x 108 2.887 x 102 0.598 6.849 1.299 2.322 2.252 0.648 6.934 1.389 • 1.914 1.755 0.698 7.024 1.470 1.506 1.342 0.748 7.107 1.548 1.229 1.027 0.797 7.185 1.620 1.019 ? 7.833 x 101 0.847 7.262 1.689 8.465 x 10 5.851 0.897 7.335 1.753 7.155 4.269 0.947 7.406 1.812 6.068 3.003 0.997 7.473 1.868 5.245 1.967 1.047 7.539 1.918 4.522 1.142 123

Formation Constants for Zn(II) - Glycylglycinate Complexes at 25°C (cont.) Solution: 0.0600 M Glycylglycine 0.02018 M ZnCl, 0.38 M KCl 25 ml initial volume

Titrant: 1.0002 N NaOH ml titrant PH Y3 X3 added “L

0.050 5.268 0.120 1.689 x 107 7.597 x 103 • 0.100 5,540 0.205 9.137 x 106 3.956 0.150 5.735 0.297 6.079 2.439 0.200 5.900 0.390 4.261 1.590 0.250 6.040 0.483 3.191 1.078 0.299 6.167 0.575 2.444 7.340 x K T 0.349 6.281 0.668 1.964 4.925 0.399 6.390 0.761 1.596 3.106 0.449 6.493 0.853 1.324 1.705 0.499 6.595 0.945 1.101 5.829 x 101

ml titrant pH Y4 SLJu X4 added

0.598 6.790 1.122 3.755 x 108 3.688 x 102 0.648 6.884 1.210 3.155 2.974 0.698 6.980 1.295 2.618 2.376 0.748 7.083 1.378 2.087 1.858 0.797 7.181 1.457 1.714 1.458 0.847 7.281 1.535 1.413 1.127 0.897 7.380 1.611 1.183 8.588 x 10l 0.947 7.479 1.685 1.003 6.394 0.997 7.578 1.756 8.634 x 10 ' 4.583 1.097 7.780 1.894 6.686 1.766 1 2 h Formation Constants for Zn(II) - Glycylglycinate Complexes at 25°C (cont.)

Solution: 0.0600 M Glycylglycine 0.012108 M ZnCl, 0.40 M KCl • 25 ml initial volume

Titrant; 1.0002 N NaOH ml titrant PH Y3 X3 added L

0.050 5.470 0.183 1.032 x 10? 4.556 x 103 0.100 5.780 0.326 5.117 x 106 2.061 0.150 5.997 0.477 3.253 1.107 0.200 6.187 0.627 2.135 5.877 x K T 0.250 6.348 0.775 1.530 2.851 0.299 6.495 0.918 1.131 8.788 x 10l

ml titrant pH UL Y4 X4 added

0.399 6.769 1.195 3.016 x 108 3.001 x 102 0.449 6.896 1.325 2.262 2.124 0.499 7.024. 1.445 1.656 1.473 0.548 7.145 1.555 1.236 1.014 0.598 7.264 1.657 9.307 x 10/ 6.764 x 10 0.648 7.374 1.753 7.328 4.323 0.698 7.480 1.839 5.847 2.515 0.748 7.580 1.920 4.810 1.140 125

Formation Constants for Nl(II) - Glycylglycinate Complexes at 25°C

Solution: 0.00800 M Glycylglycine 0.002105 M NICI 2 0.50 M KCl 50 ml Initial volume

Titrant: 0.09942 N NaOH ml titrant PH n Y3 X3 added L

0.100 5.467 0.110 3.345 x 108 3.574 x 10A 0.200 5.731 0.189 1.867 1.893 0.299 5.923 0.273 1.227 1.174 0.399 6.085 0.358 8.488 x 107 7.724 x 103 0.499 6.218 0.443 6.372 5.356 0.598 6.337 0.526 4.914 3.769 0.698 6.447 0.610 3.872 2.635 0.797 6.547 0.692 3.137 1.815 0.897 6.644 0.773 2.553 1.180 1.047 6.781 0.891 1.923 4.822 x 102

ml titrant pH SL Y4 X4 added

1.346 7.031 1.113 2.237 x 1010 1.446 x 103 1.546 7.186 1.251 1.626 1 . 0 0 0 7 1.746 7.331 1.380 1.234 6.989 x 10 1.946 7.471 1.499 9.524 x 10y 4.858 2.145 7.610 1.604 7.299 3.329 2.343 7.741 1.702 5.901 2.234 2.542 7.867 1.799 5.039 1.387 2.741 7.990 1.897 4.556 6.748 x 101 126

Formation Constants for Ni(IX) - Glycylglycinate Complexes at 25°C (cont.)

Solution: 0.01200 M Glycylglycine 0.002105 M N1C12 0.50 M KCl • 50 tnl initial volume

Titrant: 0.09939 N NaOH ml titrant PH Y3 X3 "L added

0.299 5.757 0.280 1.151 x 10® 1.112 x 104 0.349 5.832 0.322 9.754 x 107 9.119 x 103 0.449 5.967 0.405 7.204 6.282 0.548 6.092 0.487 5.336 4.363 0.650 6.201 0.572 4.184 3.062 0.748 6.299 0.653 3.351 2.145 0.847 6.393 0.733 2.702 1.443 0.947 6.480 0.812 2.235 9.035 x 10* 1.047 6.565 0.890 1.849 4.764

ml titrant PH Y4 X4 added t

• 1.346 6.799 1.109 2.055 x 1010 1.412 x 103 1.546 6.938 1.246 • 1.511 9.808 x 102 1.746 7.074 1.369 1.086 6.782 • 1.946 7.199 1.481 8.126 x 109 4.726 2.145 7.316 1.582 6.202 3.290 2.343 7.425 1.671 4.850 2.273 2.542 7.526 1.753 3.931 1.526 2.741 7.621 1.829 3.262 9.639 x 101 027

Derivation of Y3, X3, YU and XU

The constants B^g and BQ3 ore defined as:

(MA+) (MAg) (MA3“)

01 (m + 2 )(a ~) 02 (m + 2 )(a -)2 03 (M+ 2 )(A“ )3

The mass balance equations for this system may be written as follows vhere the subscript t denotes the total concentration of a substance in solution:

Ht * (H+ ) - (OH") + 2(HgA+) + (HA)

At » (HgA+ ) + (HA) + (A") + (MA+) + 2 (MAg) + 3(MA3~)

Mt « (M+2) + (MA+ ) + (MAg) + (MA3")

For the Ni(ll)-glycylglycinate system, species containing doubly- ionized glycylglycinate (D”2) are not significant in the pH range used to evaluate and B.., 01 02 The values of nT are evaluated from the mass balance equations ii as follows:

(MA+ ) + 2 (MAg) ♦ 3(MA3“ ) At - (HgA+ ) - (ILA) - (A“) ^ (M+ 2 ) + (MA+ ) + (MAg) + ( M A ^ ) Mt Derivation of Y3, X3» YU and XU (cont.)

Substitution for (H_A+ ), (HA) and (A*") in terms of K_ and K0 from 2 la 2 a the equation gives an equation containing known quantities:

.!slKls * V ! S ,a. - (H+ ) + (OH-)) _ ** * V Kla

* * ~ •

By definition,

(MA+ ) + 2(MAg) + 3(MA3")

^ (M+2) + (MA+) + (MAj + (MA “)

Substitution of B^, BQ2 and BQ3 gives,

3Bq 3(A")3 + 2B (A“ )2 + B01(A") 5 » ■■■' ■ — ...... -■ ■ - * b 03(a -)3 + b 02(a ")2 + b 0 1 (a -) + 1

This equation may be simplified to

=L “ (3 " SL)B03U_>3 + 12 ~ V B02(A_)2 + (1 "

For 5^ values = (0«l-0.9), the BQ3 term of the above equation may be assumed negligible and the equation can be rearranged into the form Y3 ■ m„X3 + b„, where 3 3 129

Derivation of Y3,X3V YU and XU (cont.)

For values ** (l,l-1 .9 )t the equation can he rearranged into the form YU * m^xU + h^t vhere

EL + (ST - l)B (A") (2 - n _ ) xi, . 3 --- £---- 21-- xU - ^ (3 - 5JU-)3 (3 - nJtA")

m, n B . h, *= B k 02 uk 03 130

Amide Dissociation Constants of Ni(II) - Glycylglycinate Complexes at 25°C

Solution: 0.00800 M Glycylglycine 0.002105 M NiCl, 0.50 M KCl 50 ml Initial volume

Titrant: 0.09940 N NaOH ml titrant pH ml titrant pH ml titrant PH added added added

4.240 9.100 4.840 9.538 5.439 9.942 4.440 9.254 5.039 9.676 5.538 10.015 4.640 9.400 5.239 9.806

Solution: 0.01200 M Glycylglycine 0.002105 M NiCl2 0.50 M KCl 50 ml Initial volume

Titrant: 0.09940 N NaOH ml titrant pH ml titrant pH ml titrant PH added added added

6.437 9.380 7.037 9.731 7.636 10.100 6.637 9.499 7.237 9.847 7.736 10.166 6.837 9.615 7.436 9.967 Evaluation of K and K all a22

The mass balance equations for the Ni(ll)-glycylglycinate amide dissociation system sre:

H. - (H+ ) - (OH” ) + 2(H A+ ) + (HA) - (MD) - (MDA~) - 2(MD ”2 ) t 2 2

-(MDA2”2) -2(MD2A"3)

A ® (H_A+ ) + (HA) +(A” ) +(MA+ ) + (MD) + 2(MA ) + 2(MDA") t 2 2

+ 2(MD”2) + 3(MA “ ) + 3(MDA ”2) + 3(MD A*3) c j c 2

M. » (M+2) + (MA+ ) + (MD) + (MA ) + (MDA” ) + (MD ”2) (MA ”) t 2 2 3

+ (MDAg"2) + (MDgA*3) vhere (A*“) represents glycylglyclnate and D"2 represents doubly- lonized glycylglyclnate. The amide dissociation constants which describe the various species in terms of known quantities are defined

_ (MD)(aH )

K&11 " (MA+)

(MDA”)(a ) (MD *2)(a^) £ b K k ^ ■ n _ a21 (MA ) 822 (MDA“) 2

(MDAg^Jtay) (MD2A”3)(aH)

831 (MA3”) a32 (MDAg"2)

It has been assumed that K » K « K and that K __ ** K all a21 &31 a22 a32 With these assumptions, the mass balance equations can be written in 132

Evaluation of K „ and K 00 (cont«) all ac<; toms of and the quantities (a^)v (A“) and (M+2)t and the knortm constants Kla# Bq i » Bq 2 and B q ^» These mass Balance equations are then used in a pit-mapping routine which varies Kwl1 and K no to obtain the best fit of the data, a22 133

Schlff Base Formation Constant for Glyoxalate with Glycylglycinate at 25°C

Solution: 0.100 M Glycylglycine O.OSO M Sodium Glyoxalate 0.35 M KCl 20 ml initial volume

Titrant: 1.000 N NaOH ml titrant pH ml titrant pH ml titrant pH added added added

0.299 6.952 0.897 7.754 1.496 8.424 0.399 7.131 0.997 7.860 1.596 8.563 0.499 7.280 1.097 7.964 1.696 8.723 0.598 7.410 1.197 8.076 1.796 8.934 0.698 7.530 1.297 8.186 0.797 7.644 1.396 8.301

Solution: 0.100 M Glycylglycine 0.100 M Sodium Glyoxalate 0.30 M KCl 20 ml initial volume

Titrant: 1.000 N NaOH ml titrant pH ml titrant pH ml titrant pH added added added

0.200 6.491 0.797 7.369 1.396 8.032 0.299 6.704 0.897 7.476 1.496 8.161 0.399 6.871 0.997 7.582 1.596 8.305 0.499 7.016 1.097 7.690 1.696 8.477 0.598 7.143 1.197 7.799 1.796 8.700 0.698 7.259 1.297 7.909 1.896 9.060 13fc

Schlff Base Formation Constant for Glyoxalate with Glycylglycinate at 25°C (cont.)

Solution: 0.100 M Glycylglycine 0.200 M Sodium Glyoxalate 0.20 M KCl 20 ml initial volume

Titrant: 1.000 N NaOH ml titrant pH ml titrant pH ml titrant pH added added added

0.200 6.213 0.797 7.046 1.396 7.665 0.299 6.415 0.897 7.146 1.496 7.785 0.399 6.576 0.997 7.245 .1.596 7.922 0.499 6.710 1.097 7.346 1.696 8.091 0.598 6.828 1.197 7.447 1.796 8.303 0.698 6.937 1.297 7.552 1.896 8.635

Solution: 0.100 M Glycylglycine 0.300 M Sodium Glyoxalate 0.10 M KCl 20 ml initial volume

Titrant: 1.000 N NaOH ml titrant PH ml titrant pH ml titrant PH added* added 4 dded

0.200 6.072 0.797 6.879 1.396 7.512 0.299 6.267 0.897 6.978 1.496 7.641 0.399 6.997 0.997 7.083 1.596 7.788 0.499 6.555 1.097 7.184 1.696 7.975 0.598 6.672 1.197 7.288 1.796 8.249 0.698 6.779 1.297 7.396 1.896 8.815 135

Schiff Base Formation Constants for Glyoxalate with Glycylglycinate at 25°C (cont.)

Solution: 0.0100 M Glycylglycine 0.0100 M Sodium Glyoxalate 0.48 M KCl 10 ml initial volume

Titrant: 0.09882 N NaOH ml titrant PH ml titrant pH ml titrant PH added added added

0.000 5.572 0.349 7.748 0.648 8.297 0.150 7.260 0.449 7.928 0.748 8.498 0.250 7.534 0.548 8.114 0.847 8.750

Solution: 0.0100 M Glycylglycine 0.0200 M Sodium Glyoxalate 0.47 M KCl 10 ml initial volume

Titrant: 0.09882 N NaOH ml titrant pH ml titrant pH ml titrant pH added added added

0.000 5.700 0.349 7.731 0.648 8.189 0.150 7.126 0.449 7.822 0.748 8.392 0.250 7.411 0.548 8.002 0.847 8.655 136

Schiff Base Formation Constants for Glyoxalate with Glycylglyclna'te at 25°C (cont.)

Solution: 0.0100 M Glycylglycine 0.0300 M Sodium Glyoxalate 0.46 M KCl 10 .ml Initial volume

Titrant: 0.09882 N NaOH ml titrant PH ml titrant pH ml titrant PH added added added

0.000 5.736 0.349 7.532 0.648 8.076 0.150 7.035 0.449 7.725 0.748 8.295 0.250 7.315 0.548 7.900 0.847 8.560

i Solution: 0.0100 M Glycylglycine 0.100 M Sodium Glyoxalate 0.39 M KCl 10 ml initial volume

Titrant: 0.09882 N NaOH

ml titrant PH ml titrant pH ml titrant PH added added added

0.150 6.637 0.449 7.317 0.748 7.877 0.250 6.912 0.548 7.498 0.847 8.140 0.349 7.131 0.648 7.679 0.947 8.550 137

Evaluation of B_. QA

The mass balance equations for the glycylglycine-glyoxalate system are:

H - (H+ ) - (OH") + 2(H A+ ) + (HA) + (HGx) » z

At « (HgA+ ) + (HA) +(A") + (AGx**2 )

Gx^ * (HGx) + (Gx“) + (AGx”2) vhere the species (AGx"2) represents the Schiff base formed by glycylglycinate and glyoxalate* These equations can be written in terms of B ^ , the quantities (a^)* (A” ) and (Gx“ ), and the known constants and K&. The equations ore then used in a pit- mapping routine which varies B . to obtain the best fit of the data* OA 138

Zinc(II) Mixed Formation Constants at 25°C

Solution: 0.0200 M Glycylglycine 0.004030 M ZnCl2 Concentration of base as indicated for each run 0.42 M KCl

Titrant: 0.200 M Sodium Glyoxalate 0.30 M KCl ml titrant pH ml titrant pH added added

Run A (0.001996 N NaOH)

1.00 0.196 5.00 0.422 2.00 0.292 10.00 0.460

Run B (0.003994 N NaOH)

1.00 0.146 5.00 0.406 2.00 0.248 10.00 0.502

Run C (0.009982 N NaOH)

1.00 0.106 5.00 0.406 2.00 0.208 10.00 0.556

Run D (0.14972 N NaOH)

1.00 0.112 5.00 • 0.396 2.00 0.200 10.00 0.550 139

Zinc(II) Mixed Formation Constants at 25°C (cont.) » Solution: 0.0400 M Glycylglycine 0.008060 M ZnCl2 Concentration of base as indicated for each run 0.39 M KCl

Titrant: 0.400 M Sodium Glyoxalate 0.10 M KCl ml titrant pH ml titrant pH added added

Run A (0.02994 N NaOH)

1.00 0.138 5.00 0.522 2.00 0.264 10.00 0.712

Run 8 (0.01996 N NaOH)

1.00 0.156 5.00 0.516 2.00 0.284 10.00 0.672

Run C (0.00998 N NaOH)

1.00 0.208 5.00 0.480 2.00 0.324 10.00 0.554 \

Nickel(II) Mixed Formation Constants at 2S°C • , Solution: 0.0200 M Glycylglycine 0.004056 M NiClo Concentration-of base as indicated for each run 0.42 M KCl

Titrant: 0.200 M Sodium Glyoxalate 0.30 M KCl ml titrant pH ml titrant pH added added

Run A (0.01497 N NaOH)

1.00 0.084 5.00 0.358 2.00 0.162 10.00 0.514

Run B (0.00998 N NaOH)

1.00 0.102 5.00 0.382 2.00 0.198 10.00 0.516

Run C (0.00499 N NaOH)

1.00 0.236 5.00 0.534 2.00 0.368 10.00 0.582 ll»l

Nickel(II) Mixed Formation Constants at 2S°C (cont.)

Solution: 0.0400 M Glycylglycine 0.008112 M NiCU Concentration ot base as indicated for each run 0.39 M KCl

Titrant: 0.400 M Sodium Glyoxalate 0.10 M KCl ml titrant pH ml titrant pH added added

Run A (0.02994 N NaOH)

1.00 0.116 5.00 0.510 2.00 0.248 10.00 0.714

Run B (0.01996 N NaOH)

1.00 0.126 5.00 0.510 2.00 0.274 10.00 0.662

Run C (0.00998 N NaOH)

1.00 0.346 5.00 0.676 2.00 0.526 10.00 0.680 Evaluation of Mixed Constants

The mass balance equations for the Ni(ll)- and Zn(ll)-glycylglycinate-glyoxalate systems are:

« (H+ ) - (OH*) + 2(HgA+ ) + (HA) + (HGx)

At * ^H2A+* + +^A”* + + ^MA+^ +2(MAg) + 3(MA3")

+ (MAGx) + (MAGXg") + 2(MAgGx“ ) + 2(MAgGXg*2)

Gxt » (HGx) + (Gx") + (MGx+ ) + (MAGx) + (MAgGx”) + 2(MAGXg")

♦2(MA2Gx2“2)

Mt * + *MGx+> + ^MA+^ + ^ 2 * + MA3") + (MAGx) + (MAgGx")

+ (MAGx 2") + (MAgGXg“2 ).

These equations can he written in terms of the previously-discussed constants plus the mixed constants B , B , Bp_ and B , and the iJL 12 22 four quantities (&jj)t (A- ), (Gx“ ) and (M+2). The equations were used in a multidimensional pit-mapping procedure to obtain good estimates of the mixed constants* Refinement of the mixed constants is accomplished by meens of a self-converging pit-map procedure using these same mass balance equations. 1U3

K of Glycine at 0°C cQl Solution: 0.1012 M Glycine 0.90 M KC1 20.0 ml initial volume ml titrant pH Y2 X2 added H

Run A (Titrant: 0 .2 * 1 8 2 N NaOH)

1 .1 * 9 6 9.826 0.817 3.096 X i o 1 9 1 . 0 3 8 X 109 1.995 10.000 0.755 6 . 0 7 0 1 . 9 6 5 2.1*92 10.133 O.6 9 U 9.815 3 . 1 7 9 2.990 1 0 .2 U 5 0.633 1.1*32 x 1020 1*.715 3.1*90 10.358 0.572 2.08U 6.833 3.990 10.1*66 0.511 2.933 9.605 10 1*.1*90 10.571* 0.1* 1*9 U . 0 7 6 1.331 X 10 1*.990 10.679 0.388 5.1*93 1 . 8 1 2 5.1*89 10.799 0.327 7.7U7 2.532 5.988 10.928 0.2 66 1.100 x 1 0 2 1 ■3.587 6.1*88 11.071 0.201* 1.580 5.217 6.987 11.21*7 0.11*3 2.1*09 8.11*8

Run B (Titrant : 0.21*96 N NaOH) ,8 0.997 9.628 O . 8 7 7 1.1*08 X 1 0 1 9 1*.61*7 X 10 1.1*96 9.833 0 . 8 1 6 3.191 1 . 0 6 0 X 101 1.995 9.990 0.751* 5.780 1.929 2.1*92 10.125 0.693 9.1*21* 3.131* 2.990. 10.251* 0.631 1.1*86 x 1 0 s 0 l*.83l* 3.1*90 10.365 0.570 2.139 6.971 0 . 5 0 8 3.021 9.821 3.990 10.1*71* 10 1*.1*90 10.585 0.1* 1*6 1*.250 1.370 X 10 U.990 10.670 0.385 5.212 1.781 5.1*89 10.800 0.323 7.676 ol 2.5 !*6 5.988 10.916 0 . 2 6 2 1.023 x 1 0 2 1 3.500 6.1*88 11.070 0.200 . 1.535 5.221 6.987 11.250 0.139 2.351* 8.229 ii*i*

Kga of Glycine at 25°C

Solution: 0.1000 M Glycine 0.90 M KC1 20.0 ml initial volume

. titrant pH 72 X2 added SH

Run A (Titrant: 0.21*85 N NaOH)

1.21*7 8.991* 0.81*5 7.120 x 10lJ 1.322 x 108 1 . 7 W 9.171* 0.783 1.1*35 x 10iO 2.659 2.2l*l* 9.318 0.721 2.1*1*1 U.531 2.71H 9.1*1*1* 0 . 6 6 0 3.801* 7.056 3.2l*0 9.558 0.598 5.570 1.036 X 109 3.7^0 9.668 0.536 7.935 1.1*76 l*.2l*0 9.776 0.1*71* 1.107 x 10iy 2.058 l*.7l*0 9.881* 0.1*12 1.521 2.835 5*239 9.996 0.350 2.081* 3.902 5.738 10.120 O . 2 8 9 2.931 - 5.1*80 6.238 10.256 ' 0.227 l*.l6 2 7 . 8 6 1 10 6.738 1 0 .1 * 1 6 0 . 1 6 6 6.137 1.185 x 10

Run B (Titrant: O.2 U8 5 N NaOH)

0.997 6.888 0,876 U.655 x 101T 8.510 X 1 0 l 1.1*96 9.010 0.8ll* 7 .1 9 1 ,8 1.603 X 108 1.995 .9.258 0.752 1.978 x 10A 3.595 2.1*92 9.391* 0.691 3.238 5.852 2.990 9.511* 0.629 U.693 8.839 3.1*90 9 . 6 2 6 0.567 7.067 1.277 x 109 3.990 9.731* 0.505 9.921 1Q 1.795 1*.1*90 9.81*2 0.1* 1*3 1.371* x 1 0 ^ 2.1*86 1*.990 9.952 0.381 1.887 3.1*23 5.1*89 1 0 . 0 6 8 0.319 2.599 1*.736 5.988 1 0 . 1 9 6 0.258 3.61*9 6.690 6 .5 1 * 8 10.360 0.189 5.1*77 1 . 0 2 6 x 1010 lJ*5 K2a of Glycine at 25°C in D^O

Solution: 1.0M Glycine 5.0ml Initial volume

Titrant: 2.56H NaOH ml titrant ml sample pD *2a added withdrawn

0.0 0.50 6.768 1.00

0.20 1.00 9.538 0 . 8 7 1 14.28 x 10-11 o.Uo 1.50 9.938 0 . 7 2 6 k .3 6 0 . 6 0 2.00 10.250 0 . 5 6 2 U.38 0 . 8 0 2.50 10.576 0 . 3 7 6 U.Ui 1.00 3.00 11.06U 0 . 1 6 2 U.U7 1,20 3.50 13.150 0.0 “ 11*6

Bq A of Glycine with Glyoxalate at 0°C

Solution: 0*1000 M Glycine 0.90 M KC1 20*0 ml initial volume

Titrant: 0*290 M Sodium Glyoxalate 0.75 M KC1

ml titrant pH ml titrant pH ml titrant pH added added added

Run A (0.08796 N NaOH)

0.1*99 1 1 .3 1 * 6 2 . 9 9 0 11.200 5.1*89 11.022 0.997 11.310 3 .1 * 9 0 1 1 . 1 6 2 5.988 1 0 . 9 8 9 1.1*96 1 1 . 2 9 2 3 . 9 9 0 11.127 6.1*88 10.958 1.995 11.263 1*.1*90 11.092 6 . 9 8 7 10.931 2.1*92 11.230 1*.990 1 1 . 0 6 0 7.1*86 10.903

- Run B (0. 05871 N NaOH)

0.1*99 10.61*0 2.990 1 0 .1*1 * 0 5.1*89 10.223 0.997 10.580 3.U90 10.368 5.988 10.187 1.1*96 10.530 3.990 10.331* 6.1*88 10.156 1.995 10.1*85 1* .1*90 10.300 6.987 1 0 . 1 2 6 2.1*92 10.1*1*1* 1**990 1 0 . 2 6 2 7.1*86 10.096

Run C (0.0389UN NaOH)

0.1*99 1 0 . 2 6 2 2.990 9.996 5.1*89 9.782 0.997 10.213 3.1*90 9.91*6 5.988 9.71*8 I.U 9 6 1 0 . 1 6 0 . 3.990 9.901 . 6.1*88 9.717 1.995 1 0 . 1 0 6 1».1*90 9.861* 6.987 9.697 2.1*92 10.01*2 U.990 9.822 7.1*86 9.669 lUT

Calculation of the Theoretical pH Change Caused by COg Evolution as Reported in Table 9

The pH of a glycylglycinate-glyoxalate-Zn(ll) solution 1b a function

of the ratio of basic fora to a acidic fora as follows:

pH ■ pKga + log [(A” )/(HA)]

and the change in pH depends on the change in the logarithm of this

ratio: A pH - A log (jA-)/(HA)]

The solution from Table 9 contains:

0.025 M Glycylglycine (HA) 0.050 M ZnClg 0.050 M Glyoxalate (Gx“ ) 0.010 N NaOH

Zt 1b valid to assume that essentially all of the HA which has been

neutralized by base (0.010 M) will exist as the Zn(ll) mixed complex.

The approximate concentration of the various species is therefore:

0.015 M HA 0.03U3 M Zn+2 0.03U3 M Gx* 0.010 M complexed Zn(ll) - as the MAGx species 0.005T M complexed Zn(ll) - as the MGx+ species

The concentration of A" can now be evaluated from the relationship:

BX1 * (MAGx) « 1.7U x 105 (m +s H a -Jg x -

or (MAGx) 3.59 x 10-5 (M+2)(Gx -) x 1.7U x 105 lUQ

Calculation of the Theoretical pU Change Caused by COg Evolution as Reported In Table 9 (cont,)

The decrease in pH in Table 9 (from 0 hours to 12 hours)is dueto the transamination reaction* Transamination has the effect of reducing

(HAGx) and correspondingly (A“ ). The decrease of 0,1+1 pH units amounts to a decrease of approximately 60% in (A") and correspondingly in (MAGx). It may be assumed for these calculations that this is the position ofequilibrium for the transamination reaction.The value of

(A") after this equilibrium has been reached can be calculated from the previous equation by including a factor of 0.1+ as follows:

(A") « 0.1* x (Zn(II)complexed) = x 10-5 (M+2)(Gx -) x 1.71* x 105

The reaction in which COg evolution occurs is (neglecting inter mediate steps) as follows:

A" + Gx“ ► 2 Gly“ + C02 + HCHO where Gly“ represents glycinate. The glycinate reacts with excess HA to produce A" and glycine. For each mmole of COg evolved, two addition al mmole of HA are converted to A*.

After 1+8 hours (in Table 10), 0,025 mmole of COg were evolved and

0.0'50 mmole (0.0025 M) of HA neutralized to A". At this point the solution contains: 0.0125 M HA 0.03875 M Zn+2 0.03750 M Gx 0.01125 M complexed Zn(ll) lU9

Calculation of the Theoretical pH Change Caused by COg Evolution as Reported in Table 9 (cont.)

Therefore, A-« O A x 0.01125______- 1 . 7 8 x10-5' 0.03875 x 0.0375 x 1.71* x 10? and the change in pH from 12 hours to H8 hours is

A pH - log [(A")/(HA)J

a pH * log I"-1.'.78 * 1Q"5] - log [... = 0,17 L 0 , 0 1 2 5 J L 0 . 0 1 5 0 J

This is the maximum increase in pH which could be expected from the amount of C02 evolved. - APPENDIX B

DISTRIBUTION CURVES

.150 For the following distribution curves, the various species are represented in this manner:

HGG ** glycylglycine

GG" « glycylglycinate

Gx” « glyoxalate

0G“^ » doubly-ionized glycylglycinate

GGGx“2 = Schiff base of glycylglycinate and glyoxalate

Zn+^ = zinc(IX) ion

Ni+2 « nickel(XX) ion

151 Fractions of Glycylglycine Species as a Function of pH

Solution 1: 0,0100 M Glycylglycine 0,0100 M Sodium Glyoxalate 10 ml Initial volume

Solution 2: 0,0100 M Glycylglycine 0,0300 M Sodium Glyoxalate 10 ml initial volume

Solution 3: 0.0100 M Glycylglycine 0,100 M Sodium Glyoxalate 10 ml Initial volume

Titrant: 0,09882 N NaOH 153

O C>

O CT>

O 00 X CL-

O s

o CD

CM O

Solution 1 Ll \ j t 10.0 9.0 8.0 7.0 i 6.0 HGG 1.0 0.6 0.0 0.2 Solution 2 O oo

Solution 3 Li- Fractions of Zinc(II) Species as a Function of pH

Solution U s 0.0600 M Glycylglycine 0.01009 M ZnCl 1*0 ml initial volume

Solution 5 : 0,0600 M Glycylglycine 0.012108 M ZnCl2 25 ml initial volume—

Solution 6 : 0,0600 M Glycylglycine 0,02018 M ZnClg 2 5 ml initial volume

Solution 7 : 0.0500 M Glycylglycine 0.02018 M ZnCl£ 20 ml initial volume

Titrantt 1.000 N NaOH

156 Q oo O

Solution ^ +99UZ (99) “Z 51 65 9.0 Zn(GG) 8.0 H P

Solution 6 8.0 pH 7.0 9.0 ZnGG 6.0 ( 0.2 Q 0 5.0 Solution 7 Fractions of Nickel(ll) Species as a Function of pH

Solution 8: 0.00800 M Glycylglycine 0.002105 M UiCl2 50 ml initial volume

Solution 9t 0.01200 M Glycylglycine 0.002105 M HiClg 50 ml initial volume

Titrant: 0.09939 N NaOH 162

O C5

o CO X P-

o N

O isi

O C\J

Solution 8 I l 163

O CF5

Q 00 IE P-

CD,

o iD

Solution 9 li_ Fractions of Zinc(II) Species as a Function of CGx/CGG

Solution 10: 0.01*00 M Glycylglycine 0 . 0 0 8 0 6 0 M ZnClp 0.00998 N NaOH

Solution 11: O.OUOO M Glycylglycine 0.008060 M ZnClp 0.01996 N NaOH

Solution 12: 0.0U00 M Glycylglycine 0.008060 M ZnCl 2 0.0299^ N NaOH

Titrant: O.UOO M Sodium Glyoxalate

16U 165 o CS/C6G

ra tio

Solution 10 166

CD CO

CD RATIO OF RATIO CGVCGG

Solution 11 167

00 CD

CD. RATIO OF RATIO CGVC66

CNJ

Solution 12 U_ Fractions of Nickel(ll) Species as a Function of CGx/CGG

Solution 13* 0.01*00 M Glyeylglycine 0.000112 M NiClg 0.00998 N NaOH

Solution ll*j 0.01(00 M Glyeylglycine 0.000112 M NiCl 0.01996 N NaOH

Solution 15: O.OHOO M Glyeylglycine 0,008112 M NiClg - 0.0299^ N NaOH

Titrant: 0,1(00 K Sodium Glyoxalate

168 * 00

Solution 13 Ll- O CM

Solution lU /CGG x CG RATIO RATIO OF

O O

Solution 15 LIST OF REFERENCES

. 172 LIST OF REFERENCES

1. Dunathan, H. C., Davis, L., Kury, P. G., and Kaplan, M., Biochem., 7, 4532 (1968).

2. Ayling, J. E., Dunathan, H. C., and Snell, E. E., Biochem, _7» 4537 (1968).

3. Snell, E. E., Braunstein, A. E., Severln, E. S., and Turchlnsky, Y. M., "Pyridoxal Catalysis: Enzymes and Model Systems," Inter science Publishers, New York, 1968.

4. Metzler, E. E., Ikawa, M . , and Snell, E. E., J. Am. Chem. Soc., 7b, 648 (1954).

5. Ikawa, M . , and Snell, E. E., J. Am. Chem. Soc., 76, 653 (1954).

6 . Harada, K., and Oh-hashi, J., J. Org. Chem., 32(4), 1103-1106 (1967). '

7. Cennamo, C., G. Biochim, 16(3), 182-197 (1967).

8 . Herbst, R. M., Adv. Enzymology, 4_, 75 (1946).

9. Herbst, R. M., and Engel, L. L., J. Biol. Chem., 107, 505 (1934).

10. Nakada, H. I., and Weinhouse, S. J., J. Biol. Chem., 204, 831 (1953).

11.. Dose, K., Nature, 179. 734-735 (1957).

12. Meister, A., "Biochemistry of the Amino Acids," Academic Press, New York, London, 1965.

13. Williams, D. H., and Busch, "ST H., J. Am. Chem. Soc., 87, 4644- 4645 (1965).

14. Akabori, J., Otani, T. T., Marshal, R,, Winitz, M . , and Green- stein, J. P., Arch. Biochem. Biophys., 83, 1 (1959).

15. Sato, M., Okawa, K., and Akabori, S., Bull. Chem. Soc. Japan, 30, 937 (1957).

16. Benoiton, L., Winitz, M,, Colman, R. F., Birnbaum, S. M., and Greenstein, J. P., J. Am. Chem. Soc., 81, 1726 (1959).

17. Murakami, L., and Takahashi, K., Bull. Chem. Soc. Japan, 32, 308 (1959). 1 7 3 X7*»

18. Metzler, D. E., Longnecker, J. B.( and Snell, E. E., J. Am. Chem. Soc.. 76, 639-644 (1954).

19. Nakao, Y., Sakurai, K., and Nakahara, A., Bull. Chem. Soc. Japan, 38, 687-688 (1965).

20. Yoneda, H., Morimoto, Y., Nakao, Y., and Nakahara, A., Bull. Chem. Soc. Japan. 41, 255-256 (1968).

21. Nakao, Y., Sakurai, K., and Nakahara, A., Bull. Chem. Soc. Japan. 38, 687-688 (1965).

22. Lane, T. L., and Kandathil, A. J., J. Am. Chem. Soc., 83, 3782 (1961).

23. Hovey, R. J., O'Connell, J. J., and Kartell, A. E., J. Am. Chem. Soc., 81, 3189 (1959).

24. Elchhorn, G. L., and Dawes, J. W., J. Am. Chem. Soc., 76, 5663 (1954).

25. Watters, J. I., J. Am. Chem. Soc., 81, 1560 (1959).

26. Watters, J. I., and Dewitt, R., J. Am'. Chem. Soc., 82, 1333 (1960).

27. Bjerrum, J., "Metal Ammtne Formation in Aqueous Solutions," P. Haase and Son, Copenhagen, 1941.

28. Leussing, D. L., and Schultz, D. C., J. Am. Chem. Soc., 86, 4846 (1964).

29. Leussing, D. L., and Hanna, E. M,, J. Am. Chem. Soc., 88, 693 (1966).

30. Leussing, D. L., and Hanna, E. M,, Ibid., 696.

31. Buzzelli, 0. R., M. S. Thesis, The Ohio State University,Columbus, . 1966.

32. Hilton, A. S., Ph. D. Thesis, The Ohio State University, Columbus, 1968.

.33. Tallman, D. E., Ph. D. Thesis, The Ohio State University, Columbus, 1968.

34. Felty, W. L., M. S. Thesis, The Ohio State University, Columbus, 1968.

35. Leussing, D. L., J. Am. Chem. Soc., 85, 231 (1963).

36. Leussing, D. L . , Talanta, 11, 189 (1964). 175

37. Sillen, L. G.t Acta. Chem. Scand., 16, 159 (1962).

38. LI, N. C,, and Chen, M. C. M., J. Am. Chem. Soc., 80, 5678-5680 (1958).

39. Murphy, C. B., and Martell, A. E., J. Biol. Chem., 226, 37 (1957).

AO. Datta, S. P., Leberman, R . , and Rabin, B. R., Nature, 183, 745- 746 (1959).

41. Koltun, W.L., Fried, M., and Gurd, F., J. Am. Chem. Soc., 82, 238 (1960) .

42. Li, N. C., Doody, E., and White, J. M . , J. Am. Chem. Soc., 79, 5879-5863 (1957).

43. Dobbie, H . , and Kermack, W. 0., Biochem. J ., 59, 246-257 (1955).

44. Martin, R. B., Chamberlin, M., and Edsall, J. T . , J. Am. Chem. Soc., 82, 495-498 (I960).

45. Gilliard, R. D., Harrison, P. M., and McKenzie, E. D., J. Chem. Soc., 1967, 618-620 (1967).

46. Robertson, G. B., Gilliard, R. D . , McKenzie, E. D., and Mason, R., Nature, 209, 1347 (1966).

47. Nakao, Y., Sakurai, K., and Nakahara, A., Bull. Chem. Soc. Japan, 39, 1608-1610 (1966).

48. Dixon, H. B. F., Biochem J., 90, 2c (1964).

49. Dixon, H. B. F., ibid., 9 2 , 661.

50. Dixon, H. B. F., and Moret, V., Biochem J ., 94, 463 (1965).

51. Dixon, H. B. F. , ibid., 103, 38P (1967).

52. Dunn, S., J. Biol. Chem., 113, 359 (1936).

53. Metzler, D. E., Olivard, J., and Snell, E. E., J. Am. Chem. Soc., 76, 644 (1954).

54. Flaishka, H. A., "EDTA Titrations," Pergamon Press, 1959.

55. Vogel, A. I., "A Textbook of Quantitative Inorganic Analysis Includ ing Elementary Instrumental Analysis," John Wiley & Co., New York, 1961.

56. Deutsch, J. L., Lawson, A. C., and Poling, S. M., Anal. Chem., 40, 839 (1968). 176

57. Evans, J. I., and Monk, C. B., Trans. Faraday Soc., 51, 1244 (1955).

58. Gilbert, J. B., Otey, M. C., and Hearon, J. Z., J. Am. Chem. Soc., 27, 2599 (1955).

59. Smith, J. C., Dissertation, Kansas State University, (1961),

60. Datta, S. P., and Rabin, B. R., Trans. Faraday Soc., 52, 1117- 1130 (1956).

61. Kroll, H . , J. Am. Chem. Soc., 74, 2034 (1952).

62. Yasuda, M., Yamsaki, K., and Ohtaki, H., Bull. Chem. Soc. Japan, 23, 1067 (1960).

63. Bjerrum, J., Chem, Rev., 46, 381 (1950).

64. Monk, C. B., Trans. Faraday Soc., 47, 292-297 (1951). —

65. Lyons, T. D., Dissertation, Kansas State University, 1957.

66. Perkins, D. J., Biochem. J ., 57, 702 (1954).

67. Leussing, D. L., and Stanfield, C. K., J. Am. Chem. Soc., 86, 2805 (1964).

68. Bovey, F. A., Chem. & Eng. News,43 , 100 (1965).

69. Kalian, R. G., and Jencks, W. P., J. Biol. Chem., 241, 5864-5878 (1966).

70. Mathur, R., and Martin, R. B., J. Phys. Chem., 69, 668 (1965).

71. Takeda, M., and Jardetzky, 0., J. Chem. Phys., 26, 1346 (1957).

72. Morlino, V. J., and Martin, R. B., J. Am. Chem. Soc., 89, 3107- 3111 (1967).

73. Sheinblatt, M . , and Gutowsky, H. S., J. Am. Chem. Soc., 8 6 , 4814- 4820 (1064).

74. Sheinblatt, M., J. Am. Chem. Soc., 87, 572-574 (1965).

75. Sheinblatt, M., ibid., 88, 2123-2126 (1966).

76. Sheinblatt, M., ibid., 88, 2845-2843 (1966).

77. Taddei, F., and Pratt, L., J. Chem. Soc., 1964, 1553 (1964).