Ligand effects on the metal ion catalyzeddecarboxylation of dimethyloxaloacetic acid

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Authors Claus, Kenneth Granger, 1941-

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LIGAND EFFECTS ON THE METAL ION CATALYZED

DECARBOXYLATION OF DIMETHYLOXALOACETIC ACID

by/vO-'Kenneth G„ Claus

A Dissertation Submitted to the Faculty of the

DEPARTMENT OF CHEMISTRY

In Partial Fulfillment of the Requirements For the Degree of .

DOCTOR OF PHILOSOPHY

In the Graduate College

THE UNIVERSITY OF ARIZONA

1 9 6 9

THE UNIVERSITY OF ARIZONA

GRADUATE COLLEGE

I hereby recommend that this dissertation prepared under my

direction by ______________Kenneth G. Claus_______________________

entitled LIGAND EFFECTS ON THE METAL ION CATALYZED__________

DECARBOXYLATION OF DIMETHYLOXALOACETIC ACID_________

be accepted as fulfilling the dissertation requirement of the

degree of ______________DOCTOR OF PHILOSOPHY_____________________

(q S '

Dissertation Director Date

After inspection of the final copy of the dissertation, the

following members of the Final Examination Committee concur in

its approval and recommend its acceptance:*

~ y t 9 - 3 J - 6 ?

This approval and acceptance is contingent on the candidate's adequate performance and defense of this dissertation at the final oral examination. The inclusion of this sheet bound into the library copy of the dissertation is evidence of satisfactory performance at the final examination.

STATEMENT BY AUTHOR

This dissertation has been submitted in partial fulfillment of requirements for an advanced degree at The University, of Arizona and' is deposited in the University Library to be made available to borrowers under rules of the Library0

Brief quotations from this dissertation are allowable without special permission, provided that accurate acknowledgment of source is made. Requests for permission for extended quotation from or reproduction of this manuscript in whole or in part may be granted by the head of the major department or the Dean of the Graduate College when in his judgment the proposed use of the material is in the interests of scholarshipo In all other instances, however, permission must be obtained from the author0

ACKNOWLEDGMENTS

The author is greatly, indebted to Dr6 John V* Rund for his

guidance in this research, the synthesis of many of the substituted

phenanthrolines, and the writing of the computer program for calcu

lating the kinetic results0 Special thanks also go to Dr, M 0 Barfield

for his assistance with the proton magnetic resonance work.

The author also is grateful to The University of Arizona for

a teaching assistantship in 1965, the Office of Education of the De

partment of Health, Education, and Welfare for a National Defense

Education Act Title IV Predoetoral Fellowship during 1966-1968, and

the National Institutes of Health for research funds for the summers

of 1966, 1967, and 1968.

iii

TABLE OF CONTENTS

' Page

LIST OF ILLUSTRATIONS . . . ... . . . . . . . . ........ . . vi

LIST OF TABLES 0 © © © © © © © © © © © © © © » © © © © © © « © « v u i

ABSTRACT © © © © © © © © © © © © © © © © © © © © © © © ©.© © © © 12c

INTRODUCTION. © © © © © © © © © © © © © © © © © © © e © © © © © © I

The Autodecarboxylation of Oxaloacetic Acid © © © © © © © © © 1The Influence of Metal Ions on the Decarboxylation of

Oxaloacetic Acid © © © © © © © © © © © © © © © © © © © © 3Amine Catalysis of the Decarboxylation of Oxaloacetic

Ac id © © © © © © © © © © © © ©. © © © © © © © © © © © o ©. b. The Enzyme-Catalyzed Decarboxylation of Oxaloacetic Acid © © 7The Effect of Added Ligands on the Rate of the Metal-

Ion-Catalyzed Decarboxylation © © © © © © © © © © © © © ©, 8Purpose of This Research © © © © © © © © © © © © © © © © © © 9

EXPERIMENTAL © © © © © © © © © © © © © © © © © © © © © © © © ©© 11

Chemicals © © © © © © © © © © © © © © © © © © © © © © © © © © 11Dimethyloxaloacetic Acid © © © © © © © © © © .© © © © © © 111,10-Phenanthrolines © © © © © © © © © © © © © © © © © © 11

4 9 7-Dimethoxy-19 10-phenanthroline © © © © © © © © © © 13• ■* 2-Phenoxy-l910-phenanthr oline ©. © © © © © © © © © © © . 13

2-Amino-!

V

TABLE OF CONTENTS--Continued

Page

RESULTS............................................................ 31

Aqueous Amine Catalysts .................................... 31Aliphatic Amine-Metal Ion Catalysts ......................... 48Phenanthroline-Metal Ion Catalysts ......................... 48

1,10-Phenanthroline-Metal Ion Catalysts . 48Catalysts Involving Substituted Phenanthrolines ........ 51The Rate-Determining Step . . . . . 54

Rate-Determining Decarboxylation ................... 54Rate-Determining Association ....................... 55Rate-Determining Dissociation ....................... 59

Lowest Lying Phenanthroline TT PTT* Transitions........... 59Oxidation Potentials and Polarography ................ 61Proton Magnetic Resonance ......................... . . . . . 61

DISCUSSION....................................................... 65

Aqueous Metal Ion Catalysts............................... . . 65Phenanthroline--Metal Ion Catalysts . . . . . . . ........... 68

Catalysts Involving Remotely Substituted Phenanthrolines . 68Correlations With Hammett Type Q" Constants andpKa * s ............................. 68A Change in Mechanism............ 76Di s cussion........................ • • • ........... 78Overlap Integrals . . . .............................. 80TF— *7T* Transitions................................ 82

2-Substituted-1,10-phenanthroline-Metal Catalysts . . . . 84Zinc Catalysts . .................................. 86Manganese Catalysts............ 87Catalysts With Amine Groups . . . . . . 88

Conclusion .......................................... 90

REFERENCES 93

LIST OF ILLUSTRATIONS

Figure Page

1 o Reaction Vessel e « o'o o « o o o o o o o o o o » o o 18

20 Observed Kinetic Plot for the Autodecarboxylation ofDimethyloxaloacetic Acid „ * „ * * * ,

vii

LIST OF ILLUSTRATIONS--Continued

Figure Page

13. Logarithms of the Decarboxylation Rate Constants forthe Reaction Catalyzed by 1:1 Zinc-Phenanthroline Complexes as a Function of the Hammett Substituent Constants for the Phenanthrolines ........... . . . . . 74

14. Logarithms of the Decarboxylation Rate Constants forthe Reaction Catalyzed by 1:1 Iron-Phenanthroline Complexes as a Function of the Hammett Substituent Constants for the Phenanthrolines..................... 75

15. A Kinetic Run for Mn(II)--4,7-Dimethoxy-1,10- phenanthroline Catalyst at a Concentration of6.57 x 10~3 M ............................................ 77

16. Values of the Overlap Integrals as a Function ofthe Metal-Nitrogen Distance ............................ 81

17. Frequencies of the Lowest Lying zfT— > TT* Transition of the Phenanthrolines as a Function of the Decarboxylation Rate Constants for the Reaction Catalyzedby 1:1 Zinc-Phenanthroline Complexes ................... 83

18. A Possible Mechanism for the Amine CatalyzedDecarboxylation of Dimethyloxaloacetic Acid ........... 89

LIST OF TABLES

Table

I.

II.

III.

IV.

V.

VI.

VII.

VIII.

Page

Decarboxylation Rates and Rate Constants for Various Catalysts . . . . . 32

Catalysis by Aliphatic Amine Complexes ................. 49

Decarboxylation Rates of 1,lO-Phenanthroline-MetalCatalysts.................................... 50

Wave Lengths of the Lowest Lying 7T— > TT* Transitionof the Phenan thro l i n e s .................................. 60

N.M.R. Chemical Shifts of 1,10-Phenanthrolines . . . . . 62

Relative Carbon Electron Densities From N.M.R.Chemical Shifts ............... . . . . . . 64

Rate Constants for Catalysis by 1:1 Meta1-Phenanthro-line Complexes. Remotely Substituted Phenanthrolines . . 69

Rate Constants for Catalysis by 1:1 Meta1-Phenanthro-line Complexes. Two Substituted Phenanthrolines . . . . 85

viii

ABSTRACT

The decarboxylation of acids is catalyzed by certain

metal ions. Divalent transition metals catalyze the reaction in a

manner related to the degree of interaction of the metal ion with the

substrate; the greater, the interaction, the more active the catalyst.

Several tervalent metal ions are also very good catalysts. When nega

tively charged ligands are added to the catalyst, the activity of the

catalyst is decreased, whereas-neutral ligands can either enhance or

diminish the activity of a metal ion. One ligand which enhances the

activity of metal-ion catalysts is 1,10-phenanthroline.

The effect of various substituted 1,10-phenanthroline--metal

complexes on the rate of decarboxylation of dimethyloxaloacetic acid

was studied by following the evolution of carbon dioxide during the

reaction. Metals studied were Mn(II), Zn(II), and Fe(II). The effect

of several aliphatic amine--metal catalysts was also studied. Several

new derivatives of 1,10-phenanthroline were prepared and their effect

on the metal-ion catalysts studied. Nuclear magnetic resonance spectra

of several phenanthrolines were measured and the relative carbon elec

tron densities calculated. Overlap integrals for phenanthroline with

Mn and Zn were calculated, and several properties of the phenanthrolines

and their complexes were measured.

Substituents on 1,10-phenanthroline in metal--phenanthroline

complexes were found to have a significant effect on the catalytic

activity of the complex toward the decarboxylation of dimethyloxalo-

acetic acid when compared to the metal~"l§10-phenanthroline catalyst0

In general9 the phenanthroline--metal catalysts were more active than

the corresponding aqueous metal-ion catalysts9 except when steric ef

fects made the approach of the substrate to the catalyst difficult0

On the other hand9 aliphatic amines were found to decrease the activ

ity of the metal-ion catalysts0 These results suggest that the aro

matic system of 1,10-phenanthroline interacts in some manner with the

metal ion to make the metal ion a better catalyst. When the substitu

ents were on the carbon adjacent to the donor nitrogens' of 1910-

phenanthroline 9 the inductive effect of the substituents appeared to

control the activity of the catalyst, with electron withdrawing groups

activating the catalyst. Substituents on the remote positions caused

increased catalyst activity when they could donate electrons to the

phenanthroline system by means of resonance. For the latter class,

the inductive effect is of somewhat smaller importance than the reso

nance effect, whereas the inductive effect greatly predominates for

the former class. The fact that groups which donate electron density

by resonance activate the catalysts is in the opposite direction from

what would be expected from the function of the metal ion. An expla

nation is suggested for this, based on a mixing of the phenanthroline

ground and excited states as the transition state of the rate-

determining step is approached.

The rate-determining step for the decarboxylation of dimethyl-

oxaloacetic acid is the decarboxylation of the metal-substrate

xi

complex* Evidence is presented for a new mechanism occurring for some

very active catalysts, which may or may not have a rate-determining

decarboxylation step*

INTRODUCTION

Metal ions are known to catalyze many organic reactions. Al

though they do not catalyze the decarboxylation of monocarboxylic,acids

in solution, a large number of metal ions do catalyze the decarboxyla

tion of phenylmaIonic acid (21) as well as several yQ-keto acids (57,

31). However, and yQ-keto monocarboxyl ic acids and Q£-keto dicar-

boxy lie acid decarboxylations are not catalyzed by metal ions (35).

One y^-keto acid, oxaloacetic acid (A), has been the subject of a great

deal of study because it is active in many metabolic systems. In the

^C-C-CH-cf ^ ^ 0;C-t=CH-C(0 (A)HO OH HO OHketo enol

Krebs cycle, oxaloacetic acid is decomposed to carbon dioxide and

pyruvic acid, the reaction being catalyzed by an enzyme.

The Autodecarboxylation of Oxaloacetic Acid

The yQ-keto acids slowly decarboxylate without a catalyst. The

keto form of the acid was shown to be the active species by methylating

the carbon of acetoacetic acid to prevent enolization (44). The

product of the decarboxylation was, however, in the enol form (3).

Oxaloacetic acid and the mono- and dianions undergo first-order

decomposition (17). The monoanion was the most active species in

2

autodecarboxylation (17, 33). The fastest reaction occurred at a pH of

3.6, where the first dissociation was virtually complete (45), and the

keto form of the acid was predominant (32).

When the ̂ -carboxyl group of oxaloacetic acid was labeled with 13 14C of C, the rate of decarboxylation decreased, showing that the

breaking of the carbon-carbon bond was the rate-determining step (64,

65). The mechanisms for the decarboxylation of the undissociated acid

and its mono- and dianion are as follows: for the unionized acid (61);

o 0 o o 0 s0 n OHc ° 2H H *-

for the monoanion (33);

°t-C-tcf %-C-cf1 > Vc-CHj6 &A * 6h H 0 &Vfor the dianion (63);

■» Q;c-ctCH2 + co2 - 0 0 “ d

The Influence of Metal Ions on the Decarboxylation of Oxaloacetic Acid

The addition of a metal ion to solutions of oxaloacetic acid

generally caused an increase in the rate of decarboxylation (35). An

electron pair was transferred from the carboxyl group to the remainder

of the molecule, this transfer being facilitated by the positive charge

on a metal ion coordinated with the molecule. However, if the carboxyl

group was coordinated to the metal, the decarboxylation was prevented.-f- -J- -j-Monovalent cations such as Na , K ’ and Ag did not catalyze

the reaction, but di- and trivalent cations showed a wide range of ac

tivity. Rate constants of the bivalent metal ions of the first transi

tion series followed the Irving-Williams (28) order of complex stability

the more stable the complex, the faster the observed rate (56). The

rate also increased with increasing metal ion concentration, as long as

the concentration was low. The highly active Cu(II) catalyst began to

show a decrease in activity as the concentration became greater than

10 ** M. This type of behavior was also noted for Pb(II) and La(III).

Plots of the observed rate of decarboxylation versus metal ion concen- .

tration for the metals which had become the most activating at high

concentrations |zn(II), Ni(II), Co(II), Fe(II)j showed a plateau begin

ning to form at the highest molar concentration studied, 10 ̂M.

These plateaus indicated that these metals also may have reached an op

timal activity at a higher concentration. The plots of the weakly

activating metals were just beginning to rise at the highest concentra

tion studied.

The decarboxylation of oxaloacetic acid was also catalyzed by

rare earth ions (19). Three diamagnetic rare earth ions showed a

linear relationship between the logarithm of the rate constants and the

corresponding thermodynamic association constants. The paramagnetic

rare earth ions were more active than predicted from this relationship.

Three different chelate complexes with oxaloacetic acid are

shown below, each one being important for the decarobxylation of the

acid.

0. 0+ 0 0 0=C CM M b. (f(a) (b) (c)

Two arguments were presented to show which form of the chelate

was predominant. The probable association constants for the chelates

of oxaloacetic acid were estimated by comparing known association con

stants of similar compounds. Complexes a and b would have constants

similar to that of the metal-pyruvate complex, while complex c would

have a constant similar to the succinate-metal complex. The associa

tion constants for zinc with these ligands are: pyruvate (63),2 2 31.2 x 10 ; succinate (43), 5.0 x 10 ; oxaloacetate (46), 1.6 x 10 .

It was obvious that the metal-succinate complex was stronger than the

metal-pyruvate complex, and that the metal-oxaloacetate complex was

even stronger. This comparison tended to show that chelate c was the

favored form. An alternate argument stated that chelates a and b were

the favored form, with the high association constants attributed to the

stabilization of these complexes by the second carboxyl group (18).

The decarobxylation of another ^Q-keto acid, dimethyloxalo-

acetic acid, was also shown to be catalyzed by various metal ions, and

its kinetics were similar to those of oxaloacetic acid (58, 48). Plac

ing methyl groups on the methylene group of oxaloacetic acid removed

the possibility of enolization and simplified the study of the mechan

ism of the decarboxylation.

The mechanism for the decarboxylation of dimethy1oxa1oacetic

acid and oxaloacetic acid is (58)

°t-l .io-+ m1* If.L.c5'"o / / ) c ' ' (f X xXo

CH3 CH3 CH3pH3

M

, c - c < CH3 -hC02 - i X V ? - c$CH3 + m -oz nCH3 ̂ 0 x c h 3

Evidence presented for the mechanism was: a) the product of the reac

tion, Q(-ketovaleric acid, was isolated as a derivative, b) the enolic

intermediate was identified spectrophotometrically, c) the pH depend

ence of the reaction showed that the complex of the dianion was the

most active species, d) metals which form square planar complexes were

good catalysts, indicating that the metal ion need not be coordinated

with both of the carboxyl groups and the carbonyl oxygen, e) no elec

tron transfer was involved because AI(III) was a good catalyst, and

this is the only stable ion for aluminum.

with primary amines being the most active catalysts, secondary amines

less active, and tertiary amines inactive (25). For the decarboxyla

tion of oxaloacetic acid, aniline was a very good catalyst while

ethylenediamine and ammonia were not as active but still catalyzed the

reaction to a small extent (47). The effects of other amines and amino

acids have also been studied (1).

Amine Catalysis of the Decarboxylation of Oxaloacetic Acid

The decarboxylation of -keto acids is catalyzed by amines,

The mechanism for the amine-catalyzed decarboxylation is (25)

Secondary amines form only the carbinolamine, whereas primary amines

could dehydrate to form the Schiff base. The mechanism accounts for

the product being in the enol form.

The Enzyme-Catalyzed Decarboxylation of Oxaloacetic Acid

An enzyme that catalyzed the decarboxylation of oxaloacetic

acid was isolated from the bacterium Micrococcus lysodeikticus (34) 0

This enzyme required a metal ion, preferably Mn(II) or Mg(II), as a

cofactoro The crude enzyme protein prepared from parsley roots was

also a good catalyst for the decarboxylation of oxaloacetic acid (61)0

The activation of the decarboxylase by several metal ions was also

studied (56), and each metal ion had a concentration at which it acti

vated the enzyme to the greatest extent0 The greatest amount of acti

vation was found with Mn(II) as the cofactor, the rate being 18 times

as fast as the aqueous metal ion rate0 On the other hand, Ni(II) was

only 10 2 times as good a catalyst when complexed by the enzyme0‘ The

rates of the enzyme-complexed metal ion catalysts were generally 10 to

100 times as fast as the rate of the catalysts involving the enzyme

without any added metal ion„. Acetone-dried pigeon liver extract,

another catalyst for the decarboxylation, was activated by Mn(II) but

not by Mg(XI) (13)o This extract was more active than the parsley root

extract, and it was proposed that the pigeon liver extract interacted

with the substrate to a much greater extent (61)0 The action of en

zymes on oxalosuccinic acid was quite similar to that on oxaloacetic

acid (31)o . -

The formation of a Schiff base between the substrate and an

amine group on the enzyme has been suggested as being involved in the

mechanism of the enzyme catalyzed decarboxylation of oxaloacetic

acid (62, 15), No mechanism for the meta11oenzyme-catalyzed

decarboxylation of yjf-keto acids has been suggested, but several data

tell something of this mechanism. When the metal ion was complexed

with an enzyme, the decarboxylation step was probably increased by ag

factor of 10 and was no longer rate determining. The following step,

regeneration of the catalyst, became rate determining in this case (52)

Dime thy1oxa1oace t ic acid which had the ^-carboxyl group labeled with 13C showed no difference in rate compared to the unlabeled acid. The

function of the enzyme was said to be twofold (58): first, it imparts

specificity with respect to the substrate to the enzyme system;

secondly, it complexes the metal ion in such a way as to enhance its

activity. These observations indicated that the enzyme-catalyzed de

carboxylation followed a different mechanistic path than the metal-ion-

catalyzed reaction, or at least had a different rate-determining step.

The Effect of Added Ligands on the Rate of the Metal-Ion-Catalyzed Decarboxylation

The addition of negatively charged ligands such as citrate and

acetate diminished the catalytic activity of the metal ions (48, 58).

This was the expected result, because the negative charge of the ligand

reduced the net positive charge on the metal, thereby reducing its

ability to attract electrons. At the same pH, metal ion concentration

and ionic strength, the rate of decarboxylation in the presence of

chloroacetate buffer, was nearly twice as large as the rate in acetate

buffer. This is reasonable, because acetate complexes are much stronge

than chloroacetate complexes. The negatively charged 8-hydroxy-

quinoline-5-sulfonic acid decreased the activity of Ni(II) and Mn(ll)

catalysts (48)0 Ethy1enediaminetetraacetic acid decreased the decar

boxylation rates, presumably by occupying all of the coordination sites

on the metal and preventing the metal from complexing with the sub

strate (48) o

Coordinating agents which did not destroy the charge on the

metal ion did not significantly decrease the activity of the metal ion6

Added pyridine caused a rate enhancement, which was attributed to the

fact that the uncharged pyridine molecule complexed the metal but did.

not reduce the effective charge, thereby making the observed rate

faster than it was in the acetate buffer (58)e Added terpyridyl en

hanced the activity of Mri(II) but had little effect on the activity of

Ni(II) (48)o Added 1,10-phenanthroline enhanced the activity of both

Ni(II) and Mn(II) (48)0 The activity increased with added 1,10-

phenanthroline until a maximum was reached at a two-to-one phenanthro-

1ine to metal ratiOo At higher phenanthroline concentrations, coordina

tion of the substrate was prevented by the formation of tris-1,10-

phenanthroline complexes \ "

Purpose of This Research

The decarboxylation of dimethyloxaloacetic acid is catalyzed by

various metal ions, primary and secondary amines, arid enzymes„ The pres

ent work is concerned with finding ligand-metal complexes which cata

lyze this reaction, measuring the catalytic rate constants and observing

how these catalysts compare to the enzyme catalysts.

] Because several derivatives were available, 1,10-phenanthroline

was chosen for the ligand, and the effect of substituents on the

10

overall rate could be studied. The substituted phenanthrolines can be

placed in two classes: those substituted in the remote positions

(positions 3 through 8), and those substituted in the 2 or 2 and 9

positions, where the substituents could affect not only the electronic

properties of the phenanthroline but also could react with the sub

strate.

5 6

3

1,10-Phenanthroline

It was also of interest to look at some aliphatic amines where

the possibility of 'ff bonding between the ligand and metal does not

exist.

The metal ions to be studied were chosen for several reasons.

Mn(II) is an activating cofactor in many enzyme reactions and should

behave in a similar fashion with an enzyme model. Zn(ll) is a good

catalyst, and because it has a full 3d subshell, it should behave in a

manner different from Mn(II). Fe(II) was chosen because its oxidation

to Fe(III) can be followed, allowing the possibility of electron trans

fer to be studied.

When all of the reactions are run in the same solvent and buf

fer at the same temperature and pH, substituent effects should be the

only influence on changes in the rate for a given metal.

EXPERIMENTAL

Chemicals

Dimethyloxaloacetic Acid

Dime thyloxaloacetic acid was prepared by the method of Stein-

berger and Westheimer (58)0 The product melted at 106*, lit, 105,5°-

106,5°, It was stored under refrigeration and two gram portions were

removed and recrystallized as needed. Analytical grade benzene (100

ml)9 which had been stored over sodium, was used to recrystallize the

two gram portions,

1*10-Phenanthrolines

The following phenanthrolines were purchased from, the G, Fred

erick Smith Chemical Company and used as received: l/lO-phenanthroline

and its 4,7-dimethyl, 2,9-dimethyl, 3,5,6,8-tetramethyl, 3,4,7,8-

tetramethyl, 4 ,7-diphenyl, and 4 ,7-diphenyl sodium sulfonate deriva

tives, The 5,6-dimethyl derivative was also purchased from the G,

Frederick Smith Chemical Company and recrystallized from aqueous ethanol

prior to use,

Phenanthroline derivatives prepared according to methods in the

literature were: 5-nitro-l,10-phenanthroline (54), mp 200°, lit* 199°;

4 ,7-dichloro-l,10-phenanthroline (55), mp 248°, lit, 249-250°; 4,7-

dibromo-1,10-phenanthroline (6), mp 233°, lit, 236°; 2-chloro-l,10-

phenanthroline (24), mp 130-130,5°, lit, 129-130°; 2-cyano-l,10-phenan-

throline (9), mp 233-235°, lit, 233-234°; 2-carbonamide-

■ ■ „

; 12

1910-phenanthroline (4), rrip 302-303*, lit0 304*5-305.5°; 2-carboxy-l, 10-

phenanthroline (9), mp 210.5-211.5°, lit. 209-210°. The 3,8-dicarboxy-

4,7-dihydroxy and 4,7-dihydroxy derivatives (55). were not obtained

. pure. .

The method of Halcrow and Kermack (24) was used to prepare

2-piperidino-l, 10-phenan thro line and 2 - d i e.thy 1 aminoe thy 1 amino -1,10-

phenanthroline. For 2-diethylaminoethylamino-1,10-phenanthroline, the

final ether solution was saturated with hydrogen bromide, causing a

yellow oil to form* This mixture was placed under refrigeration over

night, and the solid was filtered off and dried under vacuum. ‘ The

crude product was sensitive to both heat and moisture. Recrystalliza

tion was done first from methanol and then from ethanol. The last re

crystallization was very slow. The compound was isolated as the yellow

dihydrobromide, while the literature method gave the 3,5-dinitrobenzo-

ate derivative. Anal. Galcd for 2H^N2NHC2H^N(^2^5^2” ^9 47*3;

H, 5.27; N, 12.29. Found: C, 47.16; H, 5.32; N, 12.21.

For 2-piperidino-l,10-phenanthroline the final ether solution

was put under refrigeration until approximately one-fourth of the ether

had evaporated, and then was saturated with hydrogen bromide. The re

sulting yellow solid was filtered off, recrystallized from ethanol, and

ether was used to flush out the product. The product was isolated as

the yellow hydrobromide, whereas the literature preparation gave the

picrate derivative. Anal. Calcd for ^^2^7^2^

13

4g7-Dimethoxy-l,10-phenanthroline. The preparation for 4,7-

dimethoxy-T,10-phenanthroline reported by Zacharias and Case (66) pro

duced poor results, so the following procedure was used0 To A* R 0

grade methanol (65 ml) were added 4,7-dichloro-1,10-phenanthroline

(2*0 grams, 0*008 mole) and a 10 percent sodium methoxide solution in

methanol (35 ml, 0*065 mole), and the mixture was refluxed for 20

hours* The precipitated sodium chloride was removed by filtration and

the solvent evaporated * The residue was recrystallized first from

aqueous ethanol and then from benzene* Final recrystallization was

done by dissolving the crude product in methanol and adding water until

the solution became cloudy* The final product was white and melted at

205** Anal* Calcd for C, 70*00; H, 5*00; N, 11*66*\ •Founds C, 70*23; H, 4*99; N, 11*67*

2-Phenoxy-l,10-phenanthroline* Finely powdered potassium hy

droxide (4*4 grams, 0*079 mole), phenol (12*2 grams, 0*079 mole), and

2-chloro-l,10-phenanthroline (8*35 grams, 0*026 mole) were well mixed*

After the mixture had been heated at 115° for 40 hours, 100 ml of 20

percent potassium hydroxide were added and the mixture was heated for

one additional hour* After cooling to room temperature, the mixture

was filtered and the isolated product was washed several times with

water by decantation* The crude product was recrystallized from a

small amount of aqueous ethanol and was allowed to dry in air* Final

drying was done at 110° for one hour, giving a product with a melting

point of 160-161° * Anal* Calcd for C^^H^N^OC^H^s. C, 79 *4; H, 4*44;

N, 10*29* Founds . C, 79*48; H, 4*38; N, 10*07*

" ' . 14

2-Amino-lgIQ-phenanthroline« A mixture of 2-phenoxy-l510-

phenanthroline (206 grams, 0.0096 moles) and ammonium chloride (15

grams, 0.28 moles) was heated at 320-340° for 40 minutes in a Wood1s

Metal bath. The resulting solid was treated with 50 ml of water and

filtered. Upon addition of 50 ml of aqueous ammonia, a resinous

brown mass precipitated. After the mixture was cooled in ice for one

hour and the liquid decanted, the remaining solid was dissolved in 50

ml of ethanol with decolorizing charcoal and was boiled for 10 minutes.

The charcoal was filtered off and the solution was dried.over molecular

sieve. The dry solution was saturated with HBr and the yellow dihydro

bromide which precipitated was dried under vacuum, giving a product

which melted at 281°. . Anal. Calcd for 2HBrs C, 40.4; H,

3.08; N, 11.76. Found % C, 40.52; H, 3.40; N, 11.46. When the di

hydrobromide was placed under vacuum for two days at 110°, one mole of

HBr was lost, giving a yellow monohydrobromide which melted at 302-304°.

Anal. Calcd for HBrs C, 52.3; H, 3.62; N, 15.2. Founds

C, 52.17; H, 3.42; N, 15.14. ,

2-Br omo -1,10- phenan thr oline. A mixture of 1-me thyl-2~o - phenan-

throlone (9.0 grams, 0.043 mole), PBr^ (220 grams, 0.81 mole), and

POBr^ (13.5 grams, 0.047 mole) was heated with stirring for 16 hours at

120°. The PBrQ was removed under vacuum, ice was added to the residue,3 • •• • • • ■ ■ -and the solution was filtered. Upon addition of aqueous ammonia, the

crude product precipitated from the filtrate. The product was recrys

tallized from water and then from methanol. The product was yellow and

melted at 161-162°. Anal. Calcd for C H N Br: C, 55.6; H, 2.70;

N, 10.80. Found: C, 55.59; H, 2.72; N, 10.85.

• : , ' ■ ■ ■ . ■ v ■ 15Z-Methoxy-l,10-phenanthroline, A 10 percent solution of sodium

methoxide in methanol (20 ml, 0,037 mole), 2-chloro-l,10-phenanthroline

(2,00 grams, 0,011 moles), and 65 ml of methanol were refluxed for 20

hours, Approximately one-half of the solvent was pumped off under

vacuum and the precipitated sodium chloride was filtered off. The ad

dition of water caused a white product .to precipitate out. This crude

product was recrystallized from 75 ml of water. The crystallization

had to be carefully induced or an oil formed. The white product had a

melting point of 74-76°, Anal, Calcd for C^.^H^N^OGH^6H^O t C, 68,5;

H, 5,37; N, 12,28, Founds C, 68,48; H, 5,26; N, 12,41, * .

Non-phenanthroline Ligands

N,N,N?,N?-Tetramethylethylenediamine, N,Nfdimethylethylenedi-

amine, N-methyl picolylamine, picolylamine, biuret, and tris(diethyl-

aminomethyl)phosphine were purchased from the Aldrich Chemical Company

and used as received. Weighed amounts of the liquid ligands were di-

- luted with water to form standard solutions for kinetic runs, Solid

ligands were weighed out as needed, "

Buffer

The buffer for all kinetic runs was made by diluting 0,05 moles

of glutaric acid (6,6 grams), and enough of a standard sodium hydroxide

solution to give 0,075 moles per liter. The glutaric acid was purchased

from Matheson, Coleman, and Bell Chemical Company and recrystallized

from chloroform before use. Sodium hydroxide solutions were standard

ized against;reagent grade potassium acid pthalate which had been

.■ v - \ -'/: v . i6

dried at 100° for two hours prior to use. The pH of the buffer was

5.3.

Standardization of Metal Ion Solutions

Metal salts were purchased from the Mallinkrodt Chemical Works

and were reagent grade.

MnCl^o solution was standardized by treating the sample

with an excess of ferrous sulfate and back titrating with standard po

tassium permanganate. The chloride ion is known to often cause prob

lems in permanganate titrations, so samples were tested containing

various amounts of added chloride ion. No significant error was found

with the excess chloride. The end point was indicated by the appear

ance of the purple color of the permanganate ion.

NiCl0° 6H 0 solution was standardized by titration with ethy1 ̂

enediaminetetraacetic acid using murexide as an indicator and aqueous

ammonia to adjust the pH to 7. The color change was blue-black to

blue-violet.

ZnCl^ solution was standardized by titration with ethylene-

diaminetetraacetic acid with xylenol orange as the indicator in a

sodium acetate-acetic acid buffer, at pH 5. At this pH the color

change was very distinct, whereas it was not at a higher or lower.pH.

CoCl0° 6H 0 solution was treated with an excess of ethylene-’I ■; I

diaminetetraacetic acid and the excess acid back titrated with standard

zinc(II) solution at pH 10 using erichrome black T as the indicator.

CuCl °xHo0 solution was standardized with sodium thiosulfate ; : 2 2 - \ . / ■ ■■ - ■ \

using iodine as a rough indicator. Starch solution was added as the

end point approacheds and the color change was blue to colorlesse

FeSO^ solution was standardized against potassium dichrornate

using an acid medium and sodium diphenyl amine as the indicator0 Care

was taken to exclude oxygen and the color change was clear to green to

purple, the purple color being the end point. When reagent grade fer

rous ammonium sulfate was used in place of the standard Fe(II) solution

in the kinetic runs, the kinetic results were the same, and the need to

prepare the kinetic runs under nitrogen was eliminated. When phenan-

throline derivatives were added, the possibility of oxidation was re

moved because phenanthroline stabilizes the.ferrous ion.

Kinetic Apparatus

Kinetic runs were carried out in the reaction vessel shown in

Figure 1. The decarboxylation of dimethyloxaloacetic acid produces

carbon dioxide which is followed with the mercury-filled manometer

tube. The entire assembly was placed in a constant temperature bath

which could hold temperature to zt0 .01D, and was large enough to accom

modate six different sets of reaction vessels at one time.. Submersible

magnetic stirrers were placed in the bath so that they lay directly

under the reaction flask. Timers were placed so that they could be

easily read while watching the mercury columns.

Kinetic Procedure

All solutions to be used for a given kinetic run were placed in

the constant temperature bath several hours before use to allow them, to

TO VACUUM

STOPCOCK

STOPCOCK 2

MANOMETER

REACTION FLASK

Figure 1. Reaction Vessel

" - /- .." ' -- i9

attain proper temperature. Before use, each reaction flask was. treated

with Siliclad? purchased from Clay, Adams, Inc0, to place a smooth,

inert surface on the glass walls and prevent catalysis by metal ion/

adsorbed on the glass, The catalyst was placed in the top compartment

and the buffer and substrate placed in the reaction flask, A small,

teflon-coated magnetic stirring bar was placed in each flask and the

flask attached to the reaction vessel0

When the catalyst was a metal complex, small changes were made

in the procedure. If the complexing agent was a solid, it was placed

in the reaction flask before attaching it to the complete vessel. If

the complexing agent was a liquid, a standard solution was made and

added to the catalyst solution. When the complexing agent reacted with

the dimethy1oxaloacetic acid, it was placed in the top compartment with

the metal ion solution. In this case the reaction had to be started

immediately, as metal complex aging has a large effect on the reaction

rate (48), Several checks were made to determine the difference be

tween placing the complexing agent in the top compartment or in the re

action flask. As long as the run was started immediately after being

set up, there was no significant difference between the two.

From four to six runs were.made at one time. The mercury level

was read at appropriate intervals to give between 15 and 25 readings

over a period of time covering one to two half lives of the reaction.

The reaction vessel was allowed to stand overnight to allow the reaction

to go to completion, and an infinite time reading taken. The observed

20

rates were averaged and the standard deviation found. The kinetic ap

paratus was carefully cleaned and dried before reuse.

Treatment of Data

For a single run of the decarboxylation of dimethyloxaloacetic

acid, the observed rate is given by the expression

Rate = kobs B

where is the total amount of substrate present, either complexed or

uncomplexed. The carbon dioxide produced during the reaction was fol

lowed as an indication of the amount of substrate reacting, giving

dt ^obs[^J dt

Integrating,

o

When the reaction is complete, the total amount of carbon dioxide

evolved is equal to the initial amount of substrate present, S^, and

the rate expression can be written

-in [(C02)00- ( C O p J = kobst

The amount of carbon dioxide is directly proportional to the pressure

change as measured in the manometer, giving

21

-ln(P - P ) (const) = k , t oo t obs

For convenience, this was changed to

-2.303 login(P - P ) (const) = k t 10 oo t obs

A plot of logj^P^- P̂ ,) gave a straight line of slope - k^^/2.303.

A sample run of an aqueous metal ion is shown in Figure 2.

In the case of autodecarboxylation, no corrections were nec

essary for the observed rate. Figure 3 shows a sample run of this

type.

For a metal ion-substrate system the autodecarboxylation rate

must be subtracted from the observed rate in order to obtain the cata

lytic rate constant, k^. The formation constants for the various metal

ions with dimethyloxaloacetic acid were estimated from similar com

pounds, the amount of uncomplexed substrate was calculated, and its

contribution was subtracted out, giving the rate of the metal ion cata

lyzed decarboxylation, k^.

When the catalyst was a metal-phenanthroline complex, two cor

rections were made. There are three decarboxylating species; the sub

strate, the metal-substrate complex, and the phenanthroline-metal-

substrate complex. The amounts of the phenanthroline-metal complex

and aqueous metal ion were calculated from known or estimated associa

tion constants, and the contribution of the latter to the rate was

eliminated. A sample run of this type is shown in Figure 4.

Figure 2. Observed Kinetic Plot for the Autodecarboxylation of Dimethyloxaloacetic Acid

Total substrate concentration = 7.81 x 10 ̂M

TIME

(Sec

. x

10"3

)22

2 -

4 -

8 "

9 -

0.5 0.4 0.3

Log (Pco-P)Figure 2. Observed Kinetic Plot for the Autodecarboxylation

of Dime thyloxaloace tic Acid

Figure 3. Observed Kinetic Plot for the Aqueous Zinc Catalyzed Decarboxylation of Dimethyloxaloacetic Acid

Total metal concentration = 7.00 x 10 ̂M

Total substrate concentration = 7.81 x 10 ^ M

23

-10

-J5

-20

-25

-30

350.6 0.5 0.4 0.3 0.2

Figure 3. Observed Kinetic Plot for the Aqueous Zinc Catalyzed Decarboxylation of Dime thy1oxaloacetic Acid

TIME

( S

econ

ds

x I0

~

Figure 4. Observed Kinetic Plot for the Z n ( 1 0 - P h e n a n t h r o l i n e Catalyzed Decarboxylation of Dimethyloxaloacetic Acid

Total phenan thro line concentration = 7.00 x 10*"'* M

Total metal concentration = 7.00 x 10” ̂M

Total substrate concentration = 7.81 x 10”^ M

TIME

(Sec. x

10-2

)24

15 -

20 -

0.8 0.7 0.6 0.5 0.4 0 .3Log (P^-P)

Figure 4. Observed Kinetic Plot for the Zn(II)-1,10-Phenanthroline Catalyzed Decarboxylation of Dimethyloxaloacetic Acid

When the phenanthroline itself caused the substrate to de-

carboxylate, the rate of this reaction was used in place of the auto

decarboxylation rate when the correction for the uncomplexed substrate

was made. Calculations were then done as in the phenanthroline-metal-

substrate case.

Calculation of the Intrinsic Rate Constant

The intrinsic rate constant, k, for the aqueous metal ion cata

lyzed decarboxylation of dime thy1oxaloacetic acid was calculated from

equation (1), where is the observed rate

constant which has been corrected for autodecarboxylation, Ka is the

concentration of metal ion not complexed by the substrate, K is the

equilibrium constant between the metal ion and substrate where the sub

strate is in a position favorable for decarboxylation, and K 1 is the

equilibrium constant between metal ion and substrate for complexation

in a manner unfavorable for decarboxylation (48). Because the present

work was concerned only with metal-substrate complexes which were

favorable for decarboxylation, K 1 was left out of equation (1) giving

equation (2).

_1kc

K! + K kK (1)

second ionization constant of dimethyloxaloacetic acid, M J is the

26

From equation (2) it is obvious that at very high metal ion

concentrations the observed catalytic rate will approach the intrinsic

rate, as all of the substrate will be in the form of the metal ion-

substrate complex. A plot of 1/k^ versus l/^nj will be a straight

line, which, when extrapolated to 1/ M = 0, will have an intercept

equal to the intrinsic rate constant, k. Figure 7 shows plots of this

type (see page 56).

The calculation becomes more complicated as M cannot be cal

culated until K is known. The concentration of metal ion is found by

considering two points, (1/k^,1/ [m ]^) and (l/k2,1/ [m ]2), which lie on

the straight line plot of 1/ M versus 1/k. Equation (2) indicates

that, since the extrapolation of 1/k^ to 1/ [m ] = 0 gives 1/k,

1 1 1 1 1̂ kl " 1/,k2(Ml I/Mi - i/[m]2 (3)

The equilibrium metal concentration can be expressed (48) as

[Mj = - k^S^/k. Substituting this into equation (3) gives

1/k, - 1/k,k,. / kn k _ -

st j " i - V St, - 1/ M2 ■ r st

(4)

Rearranging gives a quadratic in k

K2 f(M1k2/k1) - m J + k [k2(M2 - M p ] + [(k2k2 - k22)St] = 0 (5)

This equation is solved for k, which is then substituted into equation

(2) to calculate K.

This treatment is extended to the ligand-metal catalyst by mak

ing additional corrections for the incomplete association between the

ligand and metal0 Few association constants are known for metal-

substituted phenanthroline complexes* Those which are known are simi

lar to phenanthroline itself, and since corrections involving this

constant are minor, the literature values for 1,10-phenanthroline it

self (26) are used in most cases * When the phenanthroline was substi- .

tuted in the 2 position, the values for 2-methyl-1,10™phenanthrolihe

(27) are used* The association constants for 2,9-dimethy1-1^10-

phenanthroline1 are known (65) and are used only for that derivative*

Using the association constant, the correct amount of ligand-metal

complex and uncomplexed metal is calculated, and the contribution to

the observed rate of the aqueous metal-ion catalyst as well as that of

autodecarboxylation can be subtracted out*.

Observed rate constants were calculated by hand, while all

other calculations were done on either a C0D 0C* 6400 or I0B 0M* 7072

digital computer using a program written by Dr* J* V* Rund * Each time

a new supply of.. dimethyloxaloacetic acid was synthesized, a new auto-

decarboxylation rate was determined and that value used for all runs

made with: the particular batch of acid*.

Proton Magnetic Resonance

The proton magnetic resonance spectra were recorded on a Varian

A60 spectrometer, using tetramethylsilane as an external standard*

Samples were prepared:as five percent solutions in deutrochloroform

and run at room temperature * Peak position could be read to it 0*5 cps *

. • . ■ : 28

Final assignment of the chemical shifts were done on an I0B»M0 7072

digital computer using the program LA0C00N II, part 1 (7)0

Ultraviolet Spectra

Ultraviolet spectra were recorded on a Cary Model 14 recording

spectrometer between 200 and 400 microns at room temperature using one-

centimeter cellSo Cyclohexane, methanol, and water were used as sol

vents o Concentrations were varied over a hundredfold range whenever

permitted by solubility* .

Acid Dissociation Constants

. The pKa’s of the substituted phenanthrolines which have not

been reported in the literature were measured by the method of Schi It

and Smith (51) * A known amount of the phenan thro line was. mixed with a

measured amount of strong acid (HCl) and diluted to a given volume*

The pH was measured at 25° with a Beckman Model G pH meter, care being

taken to.keep the solution free of carbon dioxide contamination* Meas

urements were made in a water-dioxane mixture, as solubility prevented

the measurements from being made in pure water* The water-dioxane

ratio was varied over a wide range, and the resulting pKa values

plotted as a function.of percent dioxane and extrapolated to pure

water* A sample plot is shown in Figure 6 (see page 53), Measurements

were reproducible to i 0*02 pH units* The dioxane was purified by re-

fluxing over molten sodium and distilling as needed* With the concen

tration of added phenanthroline known, as well as the total amount of

acid and the resulting pH, the pKa of the monoprotonated phenan throline

could easily be calculated*

29

Polarography

Polarograms of the phenanthrolines and their 1:1 complexes were

run on a Seargent Model XV recording polarograph, using 0.1 M KC1 as a

supporting electrolyte and Triton X as a maximum suppressor. A drop

ping mercury electrode with a drop time of 2.6 seconds was used, and

all runs were made in a jacketed beaker at 25°. Runs were also made in

the absence of KC1, using instead the gluterate buffer solution used in

the kinetic runs. The half-wave potentials were determined as being at

the equivalence point on the resulting polarogram.

Oxidation Potentials

Formal oxidation potentials were determined for 1:1 complexes

of various phenanthrolines and iron(II). Potentiometric titrations

were done on a Leeds and Northrup student potentiometer equipped with a

calomel reference electrode and a glass indicator electrode. Known

amounts of ferrous ammonium sulfate were weighed out, dissolved in

water, and titrated with standard eerie ion solution in 1 F sulfuric

acid. Oxidation potentials were determined from the point halfway to

the equivalence point, where the last term in the equation

o o 0.0591 [le+3] [ce+3]cell Ce+4-Ce+3 Fe(phen)+2 -Fe(phen) 1 [Fe+ ]̂ [ce+^

becomes 0, and the E° of the iron-phenanthro1ine complex is found from

the difference between the known eerie ion potential and the measured

cell potential. Occasional checks were made by running the aqueous

ferrous ion, whose potential is known.

30

Overlap Integrals

Slater overlap integrals for the manganese and zinc complexes

of 1,10-phenanthroline were calculated using the tables in the litera

ture (29, 30, 37, 40). This is done by solving the equations

p = *Kpa + Jib) R/aH

(pa ~ V

for p and t, using various values of R, the distance between the metal

(a) and the nitrogen on the phenanthroline (b), and where a^ is the

Bohr radius. The values and were calculated by the method of

Slater (53). Values for p and t were calculated for all orbitals and

the total amount of overlap obtained by adding the individual orbital

contributions for the Q~ bond and for the zn~ bond. The reported value

for the Zn-N distance in crystals is 2.05 A (43). The N atom of phe-2nanthroline is considered to be sp hybridized and treated by the

method of Mu11iken (40).

RESULTS

The results of the kinetic measurements for the catalyst sys

tems studied are listed in Table I. is the total metal ion concen

tration in the reaction, and [cat] is the concentration of the catalyst

whether it is a metal ion, metal-phenanthroline or metal-amine complex,

or an amine. The observed rate constant is kQ, O" is the standard de

viation. K is the corrected rate constant, and K and k are the intrin- 7 c 1sic reaction constants which were discussed in the Experimental section.

When the rates at one or two concentrations were all that were measured

for a catalyst, K and k could not be calculated. The few cases where

two rates are given for the same catalyst at the same concentration

will be discussed in a later section.

Aqueous Amine Catalysts

Amines have been shown to catalyze the decarboxylation of oxalo

acetic acid without any metal ion being present (25). For this reason,

when the ligand used in the metal-ligand complex contained an amine

group, the ligand itself was run as the catalyst. Ethylenediamine was

was the only ligand which showed a large rate enhancement. The remain

ing ligands that contain either primary or secondary amine groups

showed a slight rate enhancement, with primary amines being somewhat

more active than secondary amines. Aniline, which is a good catalyst

for the decarboxylation of oxaloacetic acid (15), also was a good

catalyst with dime thy1oxaloacetic acid.

31

32

Table I. Decarboxylation Rates and Rate Constants for VariousCatal*ys tsa

k k kM. feat]

°cx I0b x 105 c5 x 10^ x 105 K„ .3 3 -1 -1 -1 — 1x 10 x 10 sec sec sec sec

1. Aqueous Manganese(II)

7.25 7.25 8.0 0.39 6.5 199.48 8.48 8.6 0.38 7.5 ±114.5 14.5 10.3 0.97 9.621.7 21.7 12.1 1.00 11.743.4 43.4 14.2 1.12 14.072.5 72.5 16.7 0.23 16.6108.6 108.6 18.8 0.76 18.720.8 20.8 11.9 0.51 11.5

2. Manganese(II)- -1,10-Phenanthroline

11.3 10.1 14.0 0.62 11.5 8415.0 13.7 16.9 0.84 14.6 ±822.5 20.9 22.3 1.19 19.945.0 42.7 34.9 1.44 32.175.1 72.0 45.9 2.78 42.5

112.6 108.8 62.8 2.30 58.820.8 19.2 21.6 1.06 19.3

3. Manganese(II)- -2(1,10-Phenanthroline)

7.90 5.88 13.5 0.85 9.3 1650°11.8 9.31 20.6 0.32 16.417.8 14.6 26.9 0.89 22.425.7 21.8 40.9 1.74 35.839.5 34.6 62.5 2.07 56.4

4. Manganese(II)- -4,7-Diphenyl-1,10-phenanthroline

7.5 6.57 11.8 1.00 8.8 3011.3 10.1 14.6 0.69 12.2 ±415.0 13.7 17.7 1.41 15.422.5 20.8 20.5 3.03 18.145.0 42.7 26.5 3.42 23.775.1 72.0 27.6 3.60 24.2

116±13

18±3

107±40

33

Table I. Decarboxylation Rates and Rate Constants Catalysts3-(Continued)

for Various

k k k

Mt°5[cat] x 10 x 105 x 105 x 105 K

1\3 1 r\ 3 - 1 — 1 — 1 -1x 10 x 10 sec sec sec sec

5. Manganese(II)--3,8-Dicarboxy-4,7-dihydroxy-1,10-phenanthroline

7.50 6.57 14.36 0.09 1.45 23 1211.3 10.1 14.94 0.16 2.49 ±22 ±715.0 13.7 5.22 0.37 2.93 -822.5 20.8 6.18 0.14 4.16

6. Manganese(II)--5-Nitro-l,10-phenanthroline

7.50 6.57 10.8 0.75 7.5 27 8211.3 10.1 12.3 0.35 9.5 ±2 ±1915.0 13.7 14.2 0.57 11.722.5 20.9 16.8 0.91 14.445.0 42.7 23.1 0.97 20.275.1 72.0 27.0 1.60 23.6

7. Hanganese(II)--4,7-Dihydroxy-1,10-phenanthroline

7.50 6.57 6.6 0.19 3.7 16 6711.3 10.1 7.8 0.22 5.4 ±1 ±1015.0 13.7 8.5 0.16 6.222.5 20.9 10.7 0.45 8.445.0 42.7 14.0 0.51 11.275.1 72.0 16.0 0.92 12.6112.6 108.8 18.3 0.65 14.3

8. Manganese(II)--4,7-Dimethyl- 1,10-phenanthroline

1.88 1.44 7.9 0.03 3.62.81 2.26 9.6 0.71 5.63.75 3.11 10.3 2.21 6.5 d5.63 4.83 22.5 3.15 19.27.50 6.57 34.8 5.07 31.8

9. Manganese(II)--2,9-Dimethyl- 1,10-phenanthroline

7.50 3.89 6.0 0.51 0.3 280°10.5 6.28 7.3 0.22 1.615.0 9.40 8.9 0.31 2.722.5 15.3 11.1 0.15 4.4

34

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts5--(Continued)

Mt 3X 10[cat]

3x 10

kx 105

-1secx 105

-1sec

kx 105

-1sec

kx 105

-1secK

10. Manganese(II)--4, 7-Dimethoxy -1,10- phenanthroline

4.50 3.79 21.0 1.00 16.56.00 5.17 35.5 3.17 31.69.01 7.98 40.7 2.10 37.6 d12.0 10.8 62.9 7.84 60.215.0 13.7 83.0 8.45 80.5

11. Manganese(II)--4, 7-Dichloro- 1,10-phenanthroline

9.01 7.98 11.1 0.44 8.0 21 15312.0 10.8 12.7 0.65 9.9 ±1 ±2515.0 13.7 14.8 0.21 12.221.0 19.4 16.5 0.08 14.130.0 28.1 17.6 0.37 15.145.0 42.7 18.9 0.51 16.0

12. Manganese(II)--4. 7-Dibromo-l ,10-phenanthroline

9.01 7.98 12.3 0.56 9.21 44 4515.0 13.7 16.8 0.73 14.3 ±13 +2722.5 20.8 20.5 2.52 18.1 -1745.0 42.7 29.2 6.27 26.3

13. Aqueous Zinc(II)

0.431 6.3 0.14 111 1030.718 8.0 0.19 ±10 ±151.44 1.44 12.6 0.20 9.22.87 2.87 20.3 0.98 17.64.31 4.31 26.7 0.90 24.67.18 7.18 35.8 1.84 34.714.4 14.4 .55.3 2.95 55.028.7 28.7 74.7 4.52 74.6

35

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts3--(Continued)

% x 10[cat]

3x 10

k°5 x 10-1sec

x 105 -1sec

kc5 x 10-1sec

kx 105

-1secK

14. Zinc(II) — 1,10i -Phenan thr o1ine

0.718 9.7 0.33 119 1121.01 0.94 11.6 0.20 7.2 ±30 ±401.44 1.36 14.0 0.18 9.72.01 1.92 17.6 0.29 13.52.87 2.76 22.8 0.46 19.04.31 4.18 29.9 1.20 26.6

15. Zinc(II)--2(l, 10-Phenanthroline)

0.718 0.601 8.8 0.20 3.8 350 221.44 ■ 1.27 12.8 1.34 7.72.87 2.63 22.1 1.40 17.24.31 4.00 27.9 0.91 23.1

16. Zinc(II)— 4,7-Diphenyl-1,10-phenan thr o1ine

1.44 1.36 13.4 0.31 9.1 164 662.87 2.76 22.9 0.93 19.1 ±18 ±104.31 4.18 35.3 4.25 32.07.18 7.01 50.5 5.10 48.014.4 14.1 63.0 11.67 61.0

17. Zinc(II)— 4,7-Diphenyl-1,10-phenanthroline, Sulfonated,Sodium Salt

1.44 1.36 12.61 1.33 8.32.01 1.92 15.31 1.69 11.22.87 2.76 14.21 0.96 10.43.44 3.33 15.18 0.17 11.64.31 4.18 16.13 1.68 12.85.74 5.59 18.81 2.68 16.0

18. Zinc(II)— 4,7- Dihydroxy-1,10-phenanthroline1.44 1.36 12.2 0.74 7.9 250 312.87 2.76 19.1 0.41 15.3 +325 +304.31 4.17 26.8 0.64 23.5 -100 -207.18 7.01 38.3 2.09 35.914.35 14.1 60.8 10.25 58.8

36

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts5--(Continned)

M t 3x 10feat]

3x 10

kx 105

“ 1secx 105

-1sec

kC5 x 10-1sec

kx 105

-1secK

19. Zinc(II)— 4,7-■Dimethyl-1, 10-phenanthroline

0.718 0.665 5.9 0.31 0.11.44 1.36 6.1 0.23 0.7 k2.87 2.76 6.8 1.18 2.14.31 4.18 6.7 1.28 2.7

20. Zinc(II)— 4,7- Dimethoxy-1 ,10-phenanthroline

0.574 0.528 10.0 0.78 4.30.861 0.804 16.3 1.56 10.71.15 1.08 22.8 0.28 17.3 d1.44 1.36 30.8 4.92 25.41.72 1.64 39.0 5.36 33.8

21. Zinc(Il)— 4,7-Dichloro-1, 10-phenanthroline

1.44 1.36 . 12.0 0.73 7.7 112 712.87 2.76 14.7 0.79 10.9 ±21 ±224.31 4.18 21.0 0.73 17.75.74 5.59 25.6 1.18 22.77.18 7.01 31.8 1.47 29.314.4 14.1 51.9 1.46 49.8

22. Zinc(II)— 4,7-Dibromo-l,0-phenanthroline

1.44 1.36 14.2 0.42 9.9 80 1962.87 2.76 20.6 0.71 16.8 ±7 ±354.31 4.18 25.8 1.08 22.57.18 7.01 37.7 0.71 35.3

23. Aqueous Magnesiura(II)

10.0 10.0 5.1 0.12 3.1 8.5 9315.0 15.0 5.7 0.27 4.2 ±0.3 ±1022.5 22.5 6.3 0.24 5.245.0 45.0 6.8 0.20 6.275.0 75.0 7.6 0.28 7.3150.0 150.0 8.6 0.25 8.4

37

Table I, Decarboxylation Rates and Rate Constants for VariousCatalysts5--(Continued)

% x 10

k°5[cat] x 10

3 -1 x 10 secx 105

-1sec

kx 105

- 1sec

kx 105

-1secK

24. Magne s ium(II)--1,10-Phenanthroline

45.0 6.7 0.1675.0 7.8 0.29

25. Manganese(II)--3,4,7,8-Tetramethyl-1, 10-phenanthroline

2.20 2.20 23.8 0.78 20.8 b3.84 3.19 37.2 4.3 34.45.37 4.59 57.3 3.5 54.77.69 6.75 82.7 4.8 80.4

22.5 20.8 125.0 6.0 123.0

26. Manganese(II)--5,6-Dimethyl- 1,10-phenanthroline

3.84 3.19 10.8 2.2 7.9 45 1897.69 6.75 25.4 2.0 23.1 ±8 ±70

11.53 10.36 28.3 2.5 26.115.37 14.0 27.0 0.8 24.922.5 20.8 33.5 3.6 31.3

27. Zinc(II)--5-Nitro-l,10-phenanthroline

1.44 1.36 15.1 1.0 12.0 162 962.87 2.76 28.5 2.0 25.7 ±80 ±704.31 4.18 33.5 3.6 30.95.74 5.59 41.8 2.9 40.07.18 7.01 53.0 9.9 50.9

28. Aqueous Iron(II)

2.55 2.55 18.2 0.9 15.6 60 2275.11 5.11 20.9 0.4 19.27.65 7.65 26.2 2.1 25.210.2 10.2 34.0 2.5 33.412.8 12.8 39.6 0.5 39.2

38

Table I. Decarboxylation Rates and Rate Constants for VariousCatalys tsa--(Continued)

Mt [cat] * *■ o

o Ul

x 105k

x 105k

x 105 K3x 10 3x 10 -1sec -1sec -1sec

29. Aqueous Nickel(II)

2.38 2.38 23.7 1.2 21.1 138 1434.76 4.76 39.0 2.0 37.47.14 7.14 59.2 1.7 58.49.52 9.52 65.1 1.0 64.714.3 14.3 73.1 0.6 72.9

JO. Nickel(II)— N,N,N' ,N' -Tetramethyle thylenediamine

2.38 2.37 23.6 1.0 23.4 115 1203.33 3.32 29.9 0.9 29.74.76 4.75 39.8 1.5 39.67.14 7.12 48.5 0.9 48.3

14.3 14.2 72.0 2.3 71.8

Jl. Nickel(II)— N,N ’-Dimethylethylenediamine

2.43 2.41 2.22 1.3 21.9 112 1134.87 4.84 37.0 1.2 36.67.30 7.26 47.9 2.5 47.59.74 9.70 61.2 2.6 60.712.0 12.0 56.5 2.1 56.314.6 14.6 68.7 2.2 68.1

$2. Zinc(II)— N,N, N 1,N!-Tetramethylethylenediamine

2.40 2.21 11.5 1.5 7.5 134 553.06 2.84 18.6 1.7 14.7 ±100 ±413.82 3.58 22.0 1.4 18.27.64 33.7 1.6- 30.38.61 8.24 41.2 1.5 37.911.5 11.0 40.3 3.0 37.0

13. Zinc(II)--N,N' -Dimethyle thylenediamine

2.87 2.77 • 17.8 1.3 14.7 103 954.31 4.19 23.4 0.9 20.7 ±50 ±605.74 5.60 27.4 1.8 25.07.18 7.03 33.6 1.9 31.5

39

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts3--(Continned)

k k kM [cat]t 3 3 x 10 x 10

0 5 x 10-1sec

5 5 5 x 10 x 10 x 10-1 -1sec sec

K

34. Manganese(II) --N,N,N‘,N'-T etramethylethylenediamine

7.69 2.59 7.2 0.2 0.3 b11.5 4.69 7.1 0.6 0.015.4 7.00 8.0 0.4 0.119.2 9.48 7.9 0.4 0.023.1 12.1 10.0 1.0 1.3

35. Manganese(II) --N,N'-Dimethylethylenediamine

7.69 3.50 6.3 0.2 0.2 b11.5 6.04 7.5 0.6 1.215.4 8.75 9.4 0.3 2.719.2 11.6 7.6 0.7 0.5

36. Zinc(lI)--2,9-Dimethyl-1, 10-phenanthroline

7.00 6.73 4.96 0.50 0.0

37. Zinc(II)--5,6-Dimethyl-1, 10-phenanthroline

4.31 4.18 9.67 0.18 7.144.31 4.18 31.4 2.0 28.9

38. Zinc(XI)--3,4 ,7,8-Tetramethyl-l,10-phenanthroline

4.31 4.18 17.7 2.0 14.9

39. Zinc(II)--3,5,6,8-Tetramethyl-l,10-phenanthroline

4.31 4.18 12.5 1.4 9.7

40. Mangane se(II)--3,5,6,8-Tetramethyl-l,10-phenanthroline

20.8 19.2 202 5.6 20022.5 20.8 309 49.8 307

41. Manganese(II)--5,6-Dimethyl-1,10-phenanthroline

22.5 20.8 10.6 1.2 9,5

40

Table I. Decarboxylation Rates and Rate Constants for VariousCatalystsa--(Continued)

k° 5[cat] x 10

x 10"* x 10"* sec *"x 105

-1sec

kc 5 x 10-1sec

42. Zinc(II)--2-Ani1ino-1,10-phenanthroline

7.00 6.73 33.8 2.8 30.8

43. Zinc(II)--2-Amino~l,10-phenanthroline

7.00 6.73 79.3 9.8 76.0

44. Zinc(II)--2-Amido-l,10-phenanthroline

7.00 6.73 69.3 4.1 65.9

45. Zinc(II)--2-Carboxy-1,10-phenanthroline

7.00 6.73 32.4 1.2 29.4

46. Zinc(II)--2-Bromo-l,10-phenanthroline

7.00 6.73 36.9 3.5 33.9

47. Zinc(II)--2-Carboe thoxy-1,10-phenanthroline

7.00 6.73 40.1 2.7 37.1

48. Zinc(II)--2-Cyano-l,10-phenanthroline

7.00 6.73 47.4 4.1 44.4

49. Zinc(II)--2-Chloro-1,10-phenanthroline

7.00 6.73 59.4 5.6 58.2

50. Zinc(II)--2-Phenoxy-l,10-phenanthroline

7.00 6.73 6.14 0.67 3.2

51. Zinc(II)--2-Methoxy-1,10-phenanthroline

7.00 6.73 47.2 2.9 44.2

41

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts5--(Continued)

k° 5M [cat] x 10 x 105

kc 5 x 10

3 3 -1 x 10 x 10 sec -1sec -1sec

52. Zinc(II)--2-Piperidino-l,10-phenanthroline

7.00 6.73 63.8 4.1 60.8 .

53. Zinc(II)--2-Diethylaminoethylamino-1,10-phenanthroline

7.00 6.73 297 8.7 294

54. Manganese(II)--2-Anilino-l,10-phenanthroline

2.08 1.67 11.6 0.4 7.1

55. Manganese(II)--2-Amino-1,10-phenanthroline

7.00 4.81 14.5 1.3 10.320.8 16.7 35.9 1.3 31.3

56. Manganese(II)--2-Amido-l,10-phenanthroline

7.00 4.81 9.0 0.4 4.220.8 16.7 13.4 0.5 8.7

57. Manganese(II)--2-Phenoxy-l,10-phenanthroline

20.8 16.7 10.2 0.6 5.8

58. Manganese(II)--2-Carboxy-l,10-phenanthroline

7.00 4.81 7.2 0.2 3.020.8 16.7 8.9 0.3 4.4

59. Manganese(II)--2-Carboethoxy-l,10-phenanthroline

7.00 4.81 9.6 0.4 5.420.8 16.7 14.0 0.2 9.5

60. Manganese(II)--2-Bromo-l,10-phenanthroline

20.8 16.7 9.3 0.6 4.8

42

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts5--(Continued)

M [cat]3 3 x 10 x 10

k° 5x 10 -1sec

x 105 -1sec

kx 105

-1sec

61. Manganese(II)--2-Methoxy- 1,10-phenanthroline

20.8 16.7 13.3 1.3 8.8

62. Manganese(II)--2-Chloro-1,10-phenanthroline

20.8 16.7 14.1 0.6 9.6

63. Manganese(II)--2-Cyano-l, 10-phenanthroline

20.8 16.7 10.2 0.2 9.3

64. Manganese(II)--2-Piperidino-l,10-phenanthroline

20.8 16.7 13.4 0.8 . 8.9

6 5. Manganese(II)--2-Diethy1aminoe thy1amino-1,10-•phenanthroline

20.8 16.7 27.6 1.2 22.8

66. Zinc(II)--Biuret

4.35 4.31 23.7 2.6 21.17.00 6.95 31.0 1.8 29.3

67. Zinc(II)--Ethylenediamine

5.24 5.14 35.2 1.5 27.6

68. Zinc(II)--Picolylamine

4.45 4.31 31.6 0.8 27.8

69. Zinc(II)--N-Methylpicolylamine

4.54 4.31 27.2 1.8 23.7

70. Zinc(II)--tris(Dime thy1aminome thy1)Ph o s ph ine

4.31 22.8 2.6 *■ —

43

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts3--(Continued)

k° 5M [cat] x 10

3 3 -1 x 10 x 10 secx 105 k= 5 x 10

-1sec -1sec

71. Manganese(II)--Biuret

9.43 7.25 11.4 0.7 7.5

7 2. Manganese(II)--Ethylenediamine

11.1 7.25 31.8 3.2 23.2

73. Manganese(II)--Picolylamine

11.1 7.25 15.0 0.8 9.5

74. Manganese(II)--N-MethyIpicolylamine

12.6 7.25 11.6 0.1 5.5

75. Manganese(II)--tris(Dime thylaminomethyl)Phosphine

24.7 -- 0 - - 0

76. Iron(II)--l,10-Phenanthroline

7.00 6.90 22.3 0.8 20.5

77• Iron(II)--4,7-Dimethy1-1,10-phenanthroline

7.00 6.90 26.3 0.8 24.5

78. Iron(II)--4,7-Diphenyl-l,lOphenanthroline

7.00 6.90 30.4 4.1 28.6

79. Iron(II)--4,7-Dichloro-l,10-phenanthroline

7.00 6.90 36.8 4.6 35.0

80. Iron(II)--4,7-Dibromo-l,10-phenanthroline

7.00 6.90 39.7 4.1 37.9

44

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts5--(Continued)

k° 5M [cat] x 10

1A3 ln3 -1 x 10 x 10 secx 105

-1sec

kx l o 5

- 1sec

81. Iron(II)--4,7-Dihydroxy-1,10-phenanthroline

7.00 6.90 52.4 1.5 50.6

82. Iron(II)--4,7-Dimethoxy-1,10-phenanthroline

7.00 6.90 77.4 4.3 75.6

83. Iron(II)--5-Nitro-l,10-phenanthroline

7.00 6.90 48.7 5.2 46.9

84. Iron(II)--5,6-Dimethy1-1,10-phenanthroline

7.00 6.90 33.0 0.5 31.2

85. Iron(II)--3,4,7,8-Tetramethyl-l,10-phenanthroline

7.00 6.90 42.2 8.2 40.4

86. Iron(II)--2-Chloro-l,10-phenanthroline

7.00 6.90 38.8 5.8 33.6

87. Iron(II)--3,5,6,8-Tetramethy1-1,10-phenanthroline

7.00 6.90 28.8 1.6 27.0

88. Iron(II)--Ethylenediamine

7.00 6.43 48.9 9.2 40.0

89. Iron(II)--Biuret

7.00 6.74 32.2 5.5 29.3

90. Iron(ll)--N,N,N',N'-Tetramethylethylenediamine

7.00 6.92 35.9 4.8 34.2

45

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts^--(Continued)

M [cat]3 3 x 10 x 10J

kx 105

-1secx 105

-1sec -1sec

91. Iron(II)--N-Methylpicolylamine

7.00 6.43 74.9 6.6 70.1

92. Aqueous Iron(III)

7 .00 157 16 —

93. Aqueous Cobalt(II)

7.00 7.00 25.7 0.4 24.6

94. Aqueous Copper(II)

7.00 7.00 256 15 2567.00 7.00 502 13 502

95. Aqueous Aluminum(III)

7.00 7.00 483 37 - -

96. Nickel(II)--1,10-Phenanthroline

7.00 6.99 67.6 2.3 66.7

97. Nickel(II)--4,7-Dimethyl- 1,10-phenanthroline

7.00 6.99 75.8 2.9 74.9

98. Cobalt(Il)--l,10-Phenanthroline

7.00 -- 37.1 0.7

99. Cobalt(II)--4,7-Dimethyl-1,10-phenanthroline

7.00 -- 40.2 3.3

46

Table I. Decarboxylation Rates and Rate Constants for VariousCatalystsa--(Continued)

k° 5M [cat] x 10

3 3 -1 x 10 x 10 secx 105

-1sec

kc 5 x 10-1sec

100. Copper(II)--1,10-Phenanthroline

7.00 — 221 9.5 ■ w7.00 — 603 47 --

101. Copper(II)--4,7-Dimethyl-1,10-phenanthroline

7.00 — 336 17 — **7.00 -- 582 28 - -

102, Aluminum(III)--1,10-Phenanthroline

7.00 • -- 495 24 —

103. Aluminum(III)--4,7-Dimethy1-1,10-phenanthroline

7.00 — 488 32 —

104. Magnesium(II)--2-Amino-1,10-phenanthroline

7.00 — 6.8 0.2 - -

105. No Catalyst

0.00 0.00 4.6 0.3 4.6

106. Aniline

0.00 7.00 19.5 0.9 —

107. Ethylened iamine

0.00 7.00 20.2 1.8 —

108. 2-Amino-1,10-phenanthroline

0.00 7.00 . 6.0 0.2 —■ —

47

Table I. Decarboxylation Rates and Rate Constants for VariousCatalysts5--(Continued)

M (cat]3 3 x 10 x 10

k° 5 x 10-1sec

x 105 -1sec

kc 5x 10 -1sec

109. 2-Amido-l,10-phenanthroline

0.00 7.00 6.5 0.4 --

110. Picolylamine

0.00 7.00 6.9 0.3 - -

111. N-Methylpicolylamine

0.00 7.00 4.7 0.2 - -

a. The initial concentration of dimethy1oxaloacetic acid was 7.81 x 10""3 in every case.

b. Results are too erratic to allow a meaningful calculation.

c. Calculated on the assumption of an association-limited mechanism.

d. Possible change in mechanism (see Discussion).

; ' ' ^' v" Y - ' - - " - ... ' Y " " , . -- ' ^ : . ' Y ' / ' 4 8 ''

Aliphatic Anine-Metal Ion Catalysts

The corrected rates of decarboxylation by several aliphatic

amine-metal ion catalysts are listed in Table II0 These catalysts show

an inhibiting effect when compared to the aqueous metal ion catalysts0

This effect is easily rationalized on the basis of the function of the

metal ion as an electron sinko Because amines are more basic than

water9 they should decrease the net positive charge on the metal ion

and correspondingly decrease the reaction rate0 The amine-iron cata

lysts are not included because they gave very erratic results0 This

results from the formation of hydroxo complexes in the strongly basic

amine solutions (11)0 Results for metal complexes with ethylenedia-

mine are also not included because ethy1enediamine by itself is a good

catalyst for the reaction, and the contribution to the overall rate of

complexed and uncomplexed ethylenediamine could not be separated0 In

general, the;ethylenediamine-metal ion rates were slightly faster than

the aqueous ethylenediamine rate0 Picolylamine shows a rate enhancer

ment, but it contains an aromatic nitrogen as well as an aliphatic

nitrogen,

Phenanthroline-Metal Ion Catalysts

1910-Phenanthroline-Metal Ion Catalysts

The effect of added 1,10-phenanthroline on the rate of the

meta1-1on-catalyzed decarboxylation of dimethyloxaloacetic acid is

shorn in Table III, No clear pattern is apparent from these rates ex

cept that the activity of the catalysts again follow the Irving-Williams

49

Table II. Catalysis by Aliphatic Amine Complexes

Catalyst3 k x 10"*,C -lv(sec )

Zn(I )--Aqueous ' 24.6

Zn(I )--N,N'-D ime thy1e thy1ened iamine 21.2

Zn( I )--N,N,Nl,N'-Tetramethy1ethylenediamine 20.0

Zn(I )--Biuret 21.1

Zn(I )--tris(Dime thy1aminomethy1)Pho s ph ine 22.8a

Zn(I )--Picolylamine 27.8

Zn(I )--N-MethyIpicolylamine 23.7

Mn(I )--Aqueous 6.5

Mn( I )--N,N*-Dime thyle thylened iamine 1.9

Mn( I )--N,N,N*,N’-Tetramethylethylenediamine 0.1

Mn(I )--Biuret 7.5

Mn( I )--tris(Dimethy1aminomethyl)phosphine ob

Mn( I )--Picolylamine 9.5

Mn( I )--N-MethyIpicolylamine 5.5

Ni( I )--Aqueous 58.4

Ni(I )--N,N'-Dimethylethylenediamine 48.2

Ni(I )--N,N,N’,N'-Tetramethylethylenediamine 47.0

a.« Catalyst concentrations: Zn(II), 4.31 x 10 Mn(II),7.25 x 10 ; Ni(II), 7.14 x 10 .

b. Observed rate constants because association constants are not known for this ligand.

50

Table III. Decarboxylation Rates of 1,10-Phenanthroline-Metal Catalysts

[Metal] = 7.00 x 10 ^ M [l, 10-phenanthroline] = 7.00 x 10 ̂M

[substrate*] = 7.81 x 10 ^ M k x 105

Metal ° -1 k (phen)/k (aq) assocsec o o n

Mn(II) 8.1 1.6 2.8

Fe(II) 23.5 1.0 - -

Co(II) 25.7 1.4 3.1

Ni(II) 53.3 1.3 3.5

Cu(II) 256 0.9 4.9

Zn(II) 37.2 1.4 3.2

Mg(II) 4.6 1.0 - "

Al(III) 483 1,0 ■ —

ordero One thing which does become apparent is that the metal ions

which have the lowest attraction for the substrate are most highly ac

tivated by l,10-phenanthroline

X 10

80

70

60

MANGANESE

50

in40

U30

20

20 30 40 50 60 70 80 90 100(Numbers refer to Table I)

[CATALYST] x I0 3Figure 5. Corrected Rate Constants as a Function of Catalyst Concentration

for Manganese Catalysts

70

60

Z I N C50

40 —

2030

20

2020 30 40

(Numbers refer to. Table I)

[CATALYST] x i o 3Figure 6. Corrected Rate Constants as a Function of Catalyst Concentration

for Zinc Catalysts

the favored form with iron (26). However, the formation of this com

plex was slow enough to maintain a large amount of mono complex in

solution during the kinetic runs. The observed kinetic plots with iron

do show very erratic results after the reaction has proceeded for a

while. For this reason the observed rate constants were calculated

from the first several points on the plot. The concentrations listed

were calculated from the first association constant only, and are ap

proximate. The concentration of mono complex should be reasonably

consistent for the various phenanthrolines to allow correlation of the

rates with electronic properties.

The Rate-Determining Step

In the mechanism of the metal-ion-catalyzed decarboxylation of

dimethyloxaloacetic acid, any one of three steps may control the rate;

the association of the catalyst and substrate, the decarboxylation of

the catalyst-substrate complex, or the dissociation of the decarboxy-

lated catalyst-substrate complex. The first two of these possibilities

can be distinguished by their kinetic plots.

Rate-Determining Decarboxylation. When the decarboxylation

step is rate determining, the reaction is shown by the following equa

tions:

55

where |LMJ is the catalyst concentration, £sj is the substrate concen

tration, and P is the product of the reaction. A plot of 1/k^ versus

1/[l m J will be a straight line with non-zero intercept for this type of

reaction. This relationship was derived in the experimental section.

Most catalysts showed this type of behavior. Sample reciprocal plots

are shown in Figure 7. Substituents on the phenanthroline influence

the rate of the reaction to a fairly large extent and should give in

sight as to what controls the ability of an added ligand to enhance the

catalytic activity of a metal ion.

The calculated intrinsic equilibrium constants are smaller than

the ones previously reported between the metals and similar yQ-keto

acids (20). This has been pointed out previously as arising from the

difference in what is being measured (22). The present method calcu

lates only the association constant of the acid in a manner where de

carboxylation is favored, while the potentiometrie method measured

coordination through any possible atoms. The difference, which is

about an order of magnitude, suggests that the substrate spends most of

its time coordinated in a manner unfavorable for decarboxylation. If

the added ligand were capable of causing the substrate to approach the

metal only in the manner favorable for decarboxylation, the rate would

be greatly enhanced.

Rate-Determining Association. When association is rate deter

mining, the decarboxylation will be first order in catalyst concentra

tion. A plot of k^ versus catalyst concentration, and also 1/k^

56

lbO

20

15

10

5

- + — +

00 5 10 15

/ [ c atalys t ]X I02 for Zn

-IX 10 for Mn

Figure 7. Some Representative Plots for Catalysts With Rate-Determining Decarboxylation Steps

(Numbers refer to Table I)

■> LMP + CO

57

Rate k kobs

kc k [ m ]

versus 1/j^LMJ will be straight lines passing through the origin.

One catalyst which appears to have its activity controlled by

the rate of its coordination to the substrate is the manganese cata

lyst with 2,9-dimethyl-1,10-phenanthroline. A plot of k^ versus

catalyst concentration for this catalyst is shown in Figure 8. Of all

the catalysts for which concentration dependence was studied, this one

would be the most likely to exhibit rate determining association, be

cause molecular models show that the methyl groups on the phenanthro-

line extend out past the metal ion and must hinder the approach of the

substrate. The intrinsic rate constant for this catalyst is 208 x 10 ̂

sec \ which is much smaller than the product kK for mono-phenanthro-

line catalysts which have a rate-determining decarboxylation step.

Several other 1,10-phenanthrolines that are substituted at position two

will be treated separately.

The catalyst which involves two molecules of 1,10-phenanthro-

line for each manganese ion may also fit into this category because a

plot of k^ versus catalyst concentration for this catalyst also fits

the expected pattern. When a reaction is fast, as this one is, the

intercept on a plot of 1/k^ versus 1/ |l m J will be small and could be

confused with rate-determining decarboxylation.

58

46

4 4

42

40

38

36

34

32

30

O 28x 26 U

24

2220

08

06

04

02

(Numbers refer to Table I)

(CATALYST] xIO3Figure 8. Catalysts Conforming to the Expected Pattern for

Rate-Determining Association

60

56

52

48

44

40

36

32

28

24

2016

1284

0

59

Rate-Determining Dissociation. When the dissociation step is

rate determining, the regeneration of the catalyst becomes the impor

tant step. The rate will be rapid until the catalyst becomes saturated

with decarboxylated substrate, and then decrease as it becomes limited

by catalyst regeneration.

LM + S " IMS - f-aS-t-> LMP + C02 W > LM + P + C02

This would occur only when the concentration of the catalyst is smaller

than the concentration of the substrate and only after an amount of

substrate about equal to the catalyst concentration had reacted. How

ever, an observed kinetic plot which shows two straight line portions

could also be an indication of a change in the mechanism of the reac

tion. This type of behavior was found for a few catalysts including

the aqueous copper, 4,7-dimethoxy-l,10-phenanthroline--manganese and

zinc, and 5,6-dimethyl-1,10-phenanthroline--manganese and zinc cata

lysts.

Lowest Lying Phenanthroline TT— »TT'f Transitions

The transitions listed in Table IV were identified as 'JT— >7f *

transitions by comparing the frequencies in methanol and cyclohexane.

This absorption consistently shifted to higher energy when the solvent

was changed from methanol to cyclohexane which is characteristic of

this transition (12). Extinction coefficients were similar for both

the complexed and uncomplexed phenanthrolines. In several cases, the

transition appeared as an indistinct shoulder on the strong adjacent

60

Table IV. Wave Lengths of the Lowest Lying 7T— > 7T* Transition of the Phenanthrolinesa

Derivative Uncomplexed 1:1 Zinc Complex1:1 Mn

Comp1ex1:1 Fe

Complex

Unsubs tituted 308 311 312 (318)

4,7-Dichloro 302 306 308 307

4,7-Dibromo 302 309 308 307

4,7-Dimethyl 300 (305) (310) 303

5-Nitro (306) 312 311 —

4,7-Diphenyl 310 312 313 (311)

3,4,7,8-Tetrame thy1 312 (315) (314) 322

5,6-Dimethyl — (299) (302) - -

3,5,6,8-Tetramethyl - - 307 307 (308)

a. Wave lengths in millimicrons. All metals in +2 valence state and all spectra taken in methanol.

61

absorption, and its energy could not be accurately determined. This

was especially true when the phenanthroline was complexed. The fre

quencies in Table IV which were estimated from shoulders are shown in

parentheses. The consistently fast 4,7-dimethoxy-1,10-phenanthroline

catalysts have anomalous spectra due to strong new absorptions which

appear at the position predicted for the 7T >TT* transition.

Oxidation Potentials and Polarography

The oxidation potentials that were measured showed no signifi

cant difference from the potentials of the aqueous Fe(II) ion when an

equimolar amount of phenanthroline was added. This indicates that, at

a 1:1 ratio, the phenanthroline had no effect on the oxidation poten

tial of iron.

The half wave potentials which were taken from the polarograms

of 1:1 complexes of Mn(II) and Zn(II) with the phenanthrolines showed

no definite change from the half wave potentials of the aqueous metal

ions. Insolubility was a large problem in these measurements. Some

changes which did not result from the uncomplexed phenanthroline were

noticed in the polarograms of the complexes with respect to those of

the aqueous metal ions, but were not interpretable.

Proton Magnetic Resonance

The chemical shifts of the protons on 1,10-phenanthrolines are

listed in Table V. The shifts indicate how substituents on aromatic

systems affect the 7T electron density at the various carbons. The

proton chemical shift of a proton bound to a carbon is dependent on

62

Table V. N.M.R. Chemical Shifts of 1,10-Phenanthrolinesa

DerivativeP o s i t i o n

2,9 3,8 4,7 5,6

Unsubstituted -9.28 -7.68 -8.28 -7.82

4,7-Dichloro 40.09 -0.19 - - -0.63

5-Nitro -0.22(2) -0.29(3) -0.87(4) -0.96(6)

-0.18(9) -0.25(8) -0.29(7) - -

5,6-Dimethyl 40.03 -0.02 -0.22 - -

4,7-Diphenyl -0.01 -0.07 - - -0.06

4,7-Dimethyl 40.17 40.18 -0.22

3,4,7,8-Tetramethyl 40.27 — — -0.22

4,7-Dimethoxy 40.11 40.55 — -0.51

a. 1,10-Phenanthroline figures are relative to tetramethyl- silane as an external standard. Other numbers are relative to1,10-phenanthroline. All shifts in parts per million.

63

the polarization of the C-H bond, which is in turn related to the

electron density at the carbon to which the proton is attached. This

relationship is shown (14, 16) by the equation

proton on the reference system and jOL is the change in electron density

at the ith atom. The value chosen for the proportionality constant, K,

had no effect on the results since relative electron densities were

calculated. The need to correct for ring current effects was elimi

nated by using 1,10-phenanthroline as the reference. The reliability

of the calculated electron densities at the 2 and 9 positions was

limited, however, by the magnetic anisotropy of the neighboring nitro

gens. Magnetic anisotropy has an effect of unknown size on chemical

shifts (50). The calculated electron densities are listed in Table VI.

where chemical shift of the ith proton from the corresponding

64

Table VI. Relative Shifts

Carbon Electron Densities From N.M.R. Chemica1

Phenanthroline P o s i t i o nDerivative 2,9 3,8 4,7 5,6

4,7-Dichloro 1.008 0.982 — 0.952

5-Nitro 0.982* 0.975* 0.946* 0.907b

5,6-Dimethyl 1.004 1.002 0.979 - -

4,7-Diphenyl 0.999 1.004 - - 0.994

Unsubstituted 1.000 1.000 1.000 .1.000

4,7-Dimethyl 1.016 1.017 - - 0.979

3,4,7,8-Te trame thy1 1.026 - " - - 0.979

4,7-Dimethoxy 1.011 1.051 — ' 0.952

a. Average of the two positions.

b. Position six only.

DISCUSSION

Aqueous Metal Ion Catalysts

The catalytic effect of metal ions on the decarboxylation of

oxaloacetic acid is related to the degree of interaction of the metal

ions with oxaloacetic acid in the decarboxylation step of the reaction

(22). If the interaction of metal ion and substrate in the transition

state of this step is similar to the keto complex, a linear relation

ship between the rate constants and the corresponding association con

stants for oxaloacetates would be expected. If this transition state is

tween the rates and oxalate association constants. This argument

should also hold for dime thyloxaloacetic acid, because the methyl

groups do not interfere to a significant extent with chelation. Oxalo

acetic acid decarboxylation rates follow the oxalate association con

stants much better than the oxaloacetate constants, indicating that the

transition state is more like the enol complex (22).

Reaction Intermediate OxalateComplex

Reaction Intermediate(enol) (keto)

more like the enol complex, a linear relationship would be expected be-

Figure 9 shows the association constants for oxalate and oxalo-

acetate with various metals plotted against the appropriate corrected

rate constants for dimethyloxaloacetic acid0 The oxalate constants

show a good linear relationship with the rates9 although the oxaloace-

tate constants also show a fair linear relationship0 However, oxalo-

acetate complexes can be either in the keto or enol form, whereas only

the keto form is possible for dimethyloxaloacetic acido The relation

ship between the rates and oxalate constants suggests that the transi

tion state is similar to the oxaloacetic acid case* However, the

association constants for dimethy1oxaloacetate are not available, and

so a definite conclusion cannot be drawn concerning the transition

state*

One metal deserves further comment* The observed kinetic plots

for Cu(ll) show two very definite linear portions* This is an indica

tion of tw