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Ligand effects on the metal ion catalyzeddecarboxylation of dimethyloxaloacetic acid
Item Type text; Dissertation-Reproduction (electronic)
Authors Claus, Kenneth Granger, 1941-
Publisher The University of Arizona.
Rights Copyright © is held by the author. Digital access to this materialis made possible by the University Libraries, University of Arizona.Further transmission, reproduction or presentation (such aspublic display or performance) of protected items is prohibitedexcept with permission of the author.
Download date 31/03/2021 06:04:48
Link to Item http://hdl.handle.net/10150/565187
http://hdl.handle.net/10150/565187
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
Recommended