Equilibrium studies of schiff bases and their complexes with Ni(II), Cu(II) and Zn(II) derived from salicylaldehyde and some α-amino acids

Article abstract of DOI:10.1007/s10953-005-5592-9

The acid-base equilibria of Schiff bases derived from salicylaldehyde, glycine, alanine, serine, tyrosine, and phenylalanine, and their Ni(II), Cu(II) and Zn(II) complex formation equilibria were investigated by a potentiometric method in aqueous solution (t = 25°C, μ = 0.1 M, KCl). The data from the potentiometric titrations were evaluated by means of the BEST computer program. The order of the formation constant values of the Schiff bases was Sal-Ala > Sal-Gly > Sal-Ser > Sal-Phe > Sal-Tyr, which is the same order as the increasing log K ₁ values of amino acids (and the log K ₂ values of tyrosine) with the exception of an inversion between serine and phenylalanine. Also, it was seen that the stability constants, log β₁ and log β₂, of Schiff base-metal complexes vary for all the metal ions investigated, viz., Sal-Gly > Sal-Ala > Sal-Ser > Sal-Tyr > Sal-Phe with the exception of Sal-Gly in the copper complex. The effect of the nature of the amino acids on their formation, protonation and stability constants was also discussed.

Full text of DOI:10.1007/s10953-005-5592-9

ꢀC
Journal of Solution Chemistry, Vol. 34, No. 5, May 2005 ( 2005)
DOI: 10.1007/s10953-005-5592-9
Equilibrium Studies of Schiff Bases and Their
Complexes with Ni(II), Cu(II) and Zn(II) derived
from Salicylaldehyde and Some α-Amino Acids
Havva Demirelli1,∗ and Fitnat Ko¨seog˘lu1
Received June 5, 2004; revised December 25, 2004
The acid-base equilibria of Schiff bases derived from salicylaldehyde, glycine, ala-
nine, serine, tyrosine, and phenylalanine, and their Ni(II), Cu(II) and Zn(II) complex
formation equilibria were investigated by a potentiometric method in aqueous solution
(t = 25◦C, µ = 0.1 M, KCl). The data from the potentiometric titrations were evaluated
by means of the BEST computer program. The order of the formation constant values
of the Schiff bases was Sal-Ala > Sal-Gly > Sal-Ser > Sal-Phe > Sal-Tyr, which is
the same order as the increasing log K1 values of amino acids (and the log K2 values of
tyrosine) with the exception of an inversion between serine and phenylalanine. Also, it
was seen that the stability constants, log β1 and log β2, of Schiff base–metal complexes
vary for all the metal ions investigated, viz., Sal-Gly > Sal-Ala > Sal-Ser > Sal-Tyr >
Sal-Phe with the exception of Sal-Gly in the copper complex. The effect of the nature
of the amino acids on their formation, protonation and stability constants was also
discussed.
KEY WORDS: Schiff bases; salicylaldehyde; α-amino acid; formation constant; pro-
tonation constant; stability constant; potentiometric method.
1. INTRODUCTION
A large number of Schiff bases and their complexes have been studied for their
interesting and important properties, e.g., their ability to reversibly bind oxygen,(1)
their catalytic activity in the hydrogenation of olefins,(2) the transfer of an amino
group,(3) their photochromic properties,(4) their complexing ability towards some
toxic metals.(5) On the other hand, the interactions between metal ions and amino
acids is of considerable interest as models for metal–protein reactions and models
in a variety of biological systems.(6) The ternary complexes of the Schiff bases and
1Department of Chemistry, Faculty of Gazi Education, University of Gazi, Ankara, Turkey; e-mail;
havva@gazi.edu.tr.
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0095-9782/05/0500-0561/0 2005 Springer Science+Business Media, Inc.

562
Demirelli and Ko¨seog˘lu
amino acids with metal ions bear the importance of both Schiff bases and amino
acids.(7)
Metal ions greatly enhance the formation of pyridoxal-amino acid Schiff
bases, which play an important role in many biological processes involving amino
acids and keto acids, such as transamination, decarboxylation and deamination
in aqueous solution.(8,9) The metal complexes of Schiff bases lend themselves
ideally to a study of the effect of metal ions on the equilibrium between two
small molecules and a large molecule, which may be produced from them. The
chemistry of the metal complexes of amino acid Schiff bases are of interest, be-
cause of their possible biological importance.(10–24) Transition metal complexes of
salicylaldehyde–amino-acid Schiff bases are nonenzymatic models for pyridoxal-
amino acid systems, which are of considerable importance as key intermediates
in many metabolic reactions of amino acids catalyzed by enzymes, which require
pyridoxal as a cofactor.
Although much work has been reported on the synthesis and structural char-
actarization of the metal complexes of salicylidene amino acids, a literature survey
reveals that the formation and protonation constants of Schiff bases prepared from
salicylaldehyde and amino acid, and the stability constants of their complexes
with Ni(II), Cu(II) and Zn(II), have not been determined systematically under the
conditions studied here. Particilarly, some of the equilibrium constants for sal-
icylidenephenylalanine, salicylideneserine and salicylidenetyrosine, which were
investigated in this study, have not been reported previously. This article, there-
fore, deals with the determination of the stoichiometric formation and protona-
tion constants of salicylideneglycine (Sal-Gly2−), salicylidenealanine (Sal-Ala2−),
salicylidenephenylalanine (Sal-Phe2−), salicylideneserine (Sal-Ser2−) and sali-
cylidenetyrosine (Sal-Tyr2−), and the stability constants of their complexes with
Ni(II), Cu(II) and Zn(II) in aqueous medium. Also, in this study, the stoichiomet-
ric protonation constants of saliyclaldehyde (Sal), glycine (Gly), l-alanine (Ala),
l-phenylalanine (Phe), l-serine (Ser) and l-tyrosine (Tyr), and the stability constants
of their complexes with Ni(II), Cu(II) and Zn(II) have been reported, and the rela-
tionship between the stoichiometric protonation constants of the amino acids and
the formation constants of salicylideneamino acids has been investigated.
The stoichiometric equilibrium constants were measured by a potentiometric
titration method at (25.0 0.1) ◦C under a nitrogen atmosphere in aqueous solu-
tions, in which the ionic strength was adjusted with 0.10 M KCl and calculations
were carried out using the BEST computer software.(25,26)
2. EXPERIMENTAL
Salicylaldehyde, α-amino acids and the other reagents were purchased from
Merck and used as received. Stock solutions (0.030 M) were prepared in water
and their purities were checked by potentiometric titration.(27,28) A 0.1 M HCl
solution was prepared in water and standardized by titration against primary

Equilibrium Studies of Schiff Bases and Their Complexes
563
standard sodium carbonate. Double-distilled deionized water was used throughout
the experiments. Chemically pure KCl was used to maintain a constant ionic
strength. Alkali solutions containing 0.10 M KCl were prepared using KOH and
standardizedpotentiometricallyagainstanHClsolutionusingGranplottechniques
and also allowing a determination of the dissolved carbonate impurity.(29,30) Stock
solutions of Ni(II), Zn(II) and Cu(II) ions were prepared by dissolving the required
quantities of the respective chloride compound in water containing the calculated
quantity of hydrochloric acid to prevent hydrolysis. This solution was standardized
by both the atomic absorption spectroscopic method and titration with EDTA.(31)
Potentiometric titrations were performed in a 65 mL glass vessel equipped
with a combined pH electrode (Ingold), nitrogen inlet and outlet tubes, a magnetic
stirrer, and titrant inlet. The electrode was modified by substituting its aqueous
potassium chloride solution for a mixture of 0.10 M KCl saturated with AgCl.
An Orion Model 940A pH Ionmeter was used to measure the cell e.m.f. (the
uncertainity of the e.m.f. measurements was 0.1 mV). The KOH solution was
added with an Orion 960 automatic Titermeter. The temperature was maintained
at (25.0 0.1) ◦C.
The potentiometric cell was calibrated before each experiment to obtain pH
(= −log [H+]) values for the titration medium.(25) The potentiometric titrations
were carried out at constant temperature and in an inert atmoshere of nitrogen with
CO2-free standardized 0.10 M KOH in a 50.0 mL solution containing 0.10 M KCl:
(i) (2.5 × 10−3) M HCl (for cell calibration); (ii) (1.5 × 10−3) M salicylaldehyde
(for protonation constants); (iii) (1.50 × 10−3) M HCl + (1.50 × 10−3) M amino
acid (for protonation constants); (iv) (1.50 × 10−3) M HCl + (1.5 × 10−3) M
salicylaldehyde + (1.5 × 10−3) M amino acid (for formation and protonation
constants of Schiff bases); (v) (1.5 × 10−3) M HCl + (1.5 × 10−3) M salicy-
laldehyde or amino acid + (5.0 × 10−4) M NiCl2, CuCl2 or ZnCl2 (for stability
constants of salicylaldehyde or amino acids with Ni(II), Cu(II) or Zn(II)); and
(vi) (1.5 × 10−3) M HCl + (1.5 × 10−3) M salicylaldehyde + (1.5 × 10−3) M
amino acids + (7.5 × 10−3) M NiCl2, CuCl2 or ZnCl2 {for stability constants of
the Schiff base complexes with Ni(II), Cu(II), or Zn(II)}. The cell potential was
read after waiting to establish the equilibrium throughout the titration, which was
performed at a constant ionic strength of 0.1 M with KCl. The BEST computer
software was used to determine the formation, protonation and stability constants
from the potentiometric data.(25,26)
3. RESULTS AND DISCUSSION
3.1. Binary Systems
The stoichiometric protonation constants of salicylaldehyde, glycine, alanine,
phenylalanine, serine and tyrosine, and the stoichiometric stability constants of
these amino acids and the salicylaldehyde complexes with Ni(II), Cu(II) and
Zn(II) that were obtained here and elsewhere(32–36) are presented in Table I.

564
Demirelli and Ko¨seog˘lu

Equilibrium Studies of Schiff Bases and Their Complexes
565
The values of the protonation and stability constants are in reasonable agreement
with data available in the literature.
3.2. Ternary Systems
The equilibrium pH profiles were plotted as a function of added base for
the systems containing equimolar quantities of amino acids and salicylaldehyde
to determine whether the Schiff bases form under the conditions studied. An
example plot is given in Fig. 1, where the experimental values labeled as Gly-
Sal SB is different from the pH profile calculated assuming that there is no
Fig. 1. Potentiometric titration curves of salicylaldehyde, glycine, and salicylideneglycine
as a function of added KOH (a = moles of base added per mole of ligand).

566
Demirelli and Ko¨seog˘lu
interaction between the two Schiff base-forming components, indicated by the
dotted curve. Although the differences seem small in the acidic and neutral regions,
the computational analysis indicated the formation of a substantial amount of HSB,
H2SB, H3SB species and relatively smaller amounts of the Schiff bases.
The stoichiometric formation constants (log β0), which were evaluated for the
formation of the unprotonated Schiff bases from the unprotonated forms of various
amino acids and salicylaldehyde, and the stoichiometric protonation constants of
Schiff bases were calculated by using the computer program mentioned above
and are listed in Table II together with the stoichiometric stability constants of the
Ni(II), Cu(II), and Zn(II) complexes of these Schiff bases where the values are the
cumulative constants, β, for the reactions,
n−2
i Xn+ + j Sal− + j AA−
X (Sal − AA)
+ H2O
(1)
ꢀ
ꢁ
i
j
where X is H+, Ni(II), Cu(II), or Zn(II) and AA− are amino acids.
An inspection of the log β0 values for the formation of unprotonated Schiff
bases (Table II) reveals that the order is Sal-Ala > Sal-Gly> Sal-Ser > Sal-Phe >
Sal-Tyr, which is the same order as the increasing log K1 values of amino acids (and
the log K2 values of tyrosine) with the exception of an inversion between serine
and phenylalanine. The inversion of these two amino acids may be explained by
steric effects. The steric interference between the phenyl ring in phenylalanine and
the planar Sal− is higher than between the hydroxyl group in serine and the planar
Sal−. The log β0 values for various unprotonated Schiff bases reported(8,13,21)
for salicylaldehyde range between 0.14 and 1.34 in agreement with the constants
reported here, which are in the range from 0.15 to 0.83.
The log K1 values of the Schiff bases that can be calculated from the
data in Table II using Eq. (2) are 11.13 (Sal-Gly), 10.96 (Sal-Ala), 11.08 (Sal-
Ser), 11.77 (Sal-Phe), and 12.15 (Sal-Tyr). The data obtained by Leach and
Leussing(21) for (Sal-Gly) and (Sal-Ala) are in good agreement with our finding
(see Table II).
log K1 = log β1 − log β0
(2)
A comparison between the log K1 values for Schiff bases and those for log K1
for amino acids (and the log K2 values of tyrosine) suggest that the azomethine
nitrogen has a higher basicity than those of amino acids from which it is derived,
because the proton bound to the azomethine nitrogen is stabilized by the H-bonding
interaction with phenoxide and carboxylate groups in all the Schiff bases.(8)
When the log K1 values for Schiff bases are compared, the following order
can be seen.
Sal − Tyr > Sal − Phe > Sal − Gly > Sal − Ser > Sal − Ala

Equilibrium Studies of Schiff Bases and Their Complexes
567

568
Demirelli and Ko¨seog˘lu
Scheme 1.
This order can be used as an indication to the extent of the contributions of the
parent amino acids to the basicity of the azomethine nitrogen. If the log K1 value
of Sal-Ser is neglected it is observed that there is a somewhat negative correlation
between the log K1 values of Schiff bases and the log K1 values of amino acids
(and the log K2 values of tyrosine) (r = −0.988). This negative correlation can be
explained by the formation of the carbonion(37) (Scheme I).
The resonance forms a and b of the α-deprotonated Schiff base, which is
----
monoprotonated at the azomethine nitrogen, are stabilized by the covalent N
H
bond as well as by delocalization of the negative charge. In contrast to the other
Schiff bases, salicylidenephenylalanine and salicylidenetyrosine, because they
contain a phenyl ring, which is electron withdrawing due to conjugation, the
negative charge is delocalized and the protons, which are bound to the azomethine
nitrogen, are stabilized. Also, the conjugation effect reduces the log K2 value of
tyrosine that is related to the phenolate protonation of the phenyl ring.
The other protonation sequence of the Schiff base is assigned as follows: the
protonation at the phenolic oxygen and protonation of the carboxylate group. The
constant for phenoxide protonation of the salicylidene Schiff bases is smaller than
the value observed for phenolate protonation of salicylaldehyde. This low affinity
of the phenolate group for a proton arises from the repulsive influence of the nearby
imminium ion coupled with the energy required to break the imminium-phenolate
hydrogen bond.(8)
Determination of the equilibrium constants permits the calculation of the
concentration of all species that are formed in appreciable concentrations as a
function of p[H] or other solution conditions. The distribution curve of each
Schiff base concentration vs. p[H] (Fig. 2) shows that the concentration of the

Equilibrium Studies of Schiff Bases and Their Complexes
569
Fig 2. Species’ distribution curves for the system, salicylaldehyde-serine, at a 1:1 molar ratio.
Schiff base is highest at p[H] = 9.0 and it is monoprotonated. (Sal-Ser2− is given
as an example Schiff base in Fig. 2). Extensive dissociation usually occurs in more
acidic regions where both phenolate and carboxylate groups are protonated.
3.3. Metal Complexes of Schiff Bases
The potentiometric titration curves of Schiff bases and their M(II) complexes,
together with the titration curve of the binary system M(II)/amino acid (taken as
a representative of salicylideneserine) are shown in Figs. 3–5. By comparing the
titration curves of the systems of M(II)/amino acid and M(II)/Schiff bases, it
can be seen that the beginning of each curve of the system M(II)/Schiff base is
nearly coincident with the beginning of the titration curve of the binary system
M(II)/amino acid. This behavior can be explained by the fact that amino acids lose
one proton at a lower pH, besides the fact that they bind more strongly to the metal
ion. Beyond a = 2 for the system Cu(II)/Schiff base and a = 3 for the systems
Ni(II)/Schiff base and Zn(II)/Schiff base {a = number of moles of base added
per mole of Cu(II)}, the curves of the M(II)/Schiff base were found to deviate
considerably from the curves of M(II)/amino acid, due to the formation of a Schiff

570
Demirelli and Ko¨seog˘lu
Fig 3. Potentiometric titration curves of the Cu(II)/serine and Cu(II)/salicylideneserine sys-
tems as a function of added KOH (a = moles of base added per moles of salicylideneserin
or Zn(II)).
base metal complex. For the metal complex curves, there is a significant lowering
from that of the free Schiff base, indicating formation of metal complexes by
release of protons. The Schiff bases formed in this research serve as tridentate
ligands by the coordination of the imine, phenoxide and carboxylate donor groups
with the metal ion. It has been suggested that for such complexes the metal ion

Equilibrium Studies of Schiff Bases and Their Complexes
571
Fig 4. Potentiometric titration curves of the Ni(II)/serine and Ni(II)/salicylideneserine
systems as a function of added KOH (a = moles of base added per moles of salicylideneserin
or Zn(II)).
is more strongly coordinated to the imine nitrogen and phenoxide oxygen than to
the carboxylate oxygen.(38)
The pH measurements were carried out for systems containing 1:1:1 and
1:1:0.5 molar ratios of the amino acid derivatives, salicylaldehyde and metal ion,
respectively. The appearance of a precipitate even in weakly acidic solutions was a
characteristic feature of the majority of the 1:1:1 systems. More useful information
was obtained from the metal ions whereby the complex species formed were
soluble over a large range of pH. The potentiometric data for the Cu(II)-Schiff
base systems indicate that an insignificant tendency exists toward the formation

572
Demirelli and Ko¨seog˘lu
Fig 5. Potentiometric titration curves of the Zn(II)-serine and Zn(II)-salicylideneserine systems as
a function of added KOH (a = moles of base added per moles of salicylideneserin or Zn(II)).
of M(SB)2 species, whereas the data for the Zn(II) and Ni(II) systems suggested
the formation of considerable amounts of these species in the case of Schiff bases
derived from glycine, alanine, serine, phenylalanine and tyrosine.
The proposed species’ model, which was used for each Schiff base system, is
based on other similar Schiff base systems.(8,17,21) The binary species presented in
Table I were also included and fixed in the model of Schiff base systems, but only
the ternary species were refined by the BEST program. The stability constants

Equilibrium Studies of Schiff Bases and Their Complexes
573
for the Schiff base/M(II) complexes are reported in Table II. The data obtained in
aqueous medium are in good agreement with data available in the literature.(8,20,21)
Comparison of the data in Tables I and II indicates that Schiff base/metal
complexes are significantly more stable than binary complexes having the corre-
sponding stoichiometry. This may be caused by the fact that Schiff bases behave
as tridentate ligands, each of which forms five- and six-membered fused two-ring
systems.(39,40)
According to Table II, when log β1Ni, log β2Ni, log β1Zn, and log β2Zn values
vary in the order, Sal-Gly > Sal-Ala > Sal-Ser > Sal-Tyr ≥ Sal-Phe, whereas the
log βCu values vary as, Sal-Ala > Sal-Ser > Sal-Gly > Sal-Tyr ≥ Sal-Phe. These
orders can explain the strength of the C=N bond, the basicity of the Schiff base
and the steric effect.(21) If the rank of the stability constants and the formation
and protonation constants is compared in order to investigate the influence of
the first two factors, it is seen that there is correlation between them. The linear
correlation coefficients, r, between the logarithms of the stability constants of the
complexes and the formation and protonation constants of the ligands are set out
in Table III. An inspection of the correlation coefficients (Table III) reveals that
none of the stability constants correlates significantly with log K2H (log K3H for
Sal-Tyr) and log K3H (log K4H for Sal-Tyr) and the values of the stability constants
depend primarily on the basicity of the imine group of the Schiff bases. The
negative correlation between the stability constants and the protonation constants
of the imine group indicates that the complexation equilibria of metal ions are
competitive. When a Schiff base releases a proton from the imine nitrogen, the
metal ion can bond to it.
The order of the stability constants may be due to steric interactions caused
by the bulky alkyl group attached to the amino acids. The presence of the bulky
alkyl group on the α-carbon atom of the amino acids is seen to cause a small,
Table III. The Linear Correlation Coefficients Between Pairs of Variablesa,b
Stepwise stability constants of Schiff base complexes
Ni(II)
Cu(II)
Zn(II)
Formation and protonation
constants of Schiff bases
log KML log KML2 log KML log KML log KML2
log β0
0.833∗
0.856∗
0.898∗
0.793
−0.981∗∗
0.936∗
0.545
−0.478
log K1H
log K2H
log K3H
−0.909∗ −0.781
−0.882∗ −0.833∗
−0.442
0.431
−0.531
0.133
−0.377 −0.513
0.226
0.532
aThe correlation coefficients were computed by SPSS.
bIn each case, it was used five pairs of variables in calculation of correlation coefficients.
∗Correlation is significant at the 0.05 level.
∗∗Correlation is significant at the 0.01 level.

574
Demirelli and Ko¨seog˘lu
but measurable, decrease in the stability of their Schiff base complexes. This
effect does not appear to arise from interactions of the alkyl group with other
coordinated ligands, because the presence of bulkier ligands has little effect on the
ꢀ
ꢁ
relative stability; the logarithms of the constants for reaction, M(SB) + HSB
M(SB)2 + H are −3.32, −3.29, -3.07, −3.87, −4.06 (M is Ni) and −4.56, −4.86,
−4.33, −4.55, −4.87 (M is Zn) for Sal-Gly, Sal-Ala, Sal-Ser, Sal-Phe and Sal-Tyr,
respectively. Probably the influence of the alkyl groups arises from its effects on
the solvation of the complex and on the =NCH(R)COO− bond angle.(21)
Replacement of the proton bound to the azomethine nitrogen of a Schiff base
by a divalent metal ion, except for Cu(II), does not seem to make a remarkable
difference in the stability of the complex; the M(SB) chelates investigated have
even lower formation constants than HSB. The stability sequences observed for
the species M(SB) and M(SB)2 are Cu(II) > Ni(II) > Zn(II). Thus, the Irving–
Williams order of complexation, which is related to the ligand-field stabilization
energies, is fulfilled.(41,42)
Estimation of the equilibrium concentrations of metal(II) complexes as a
function of pH provides a useful picture of metal ion binding in biological systems.
In all of the species’ distributions, the concentration of the Schiff base metal
complexes increases with increasing pH, thus making complex formation more
favored in the physiological pH range (see Figs. 6–8). The species’ distribution
pattern for the Ni-Sal-Ser system indicates that Ni(Sal-Ser) starts to form at a pH
∼5 and reaches a maximum concentration (∼60%) at pH ∼ 7.5–9.5. In this pH
range, Ni(Sal-Ser)2 begins to form and is found with Ni(Sal-Ser) in solution. The
predominant form is Ni(Sal-Ser). The distribution curve of the Zn-Sal-Ser system,
in which the Schiff base complex starts to form at pH ∼6 and its concentration
reaches a maximum (∼60%) at pH ∼ 8.0–9.0, is similar to the curve of Ni-Sal-Ser
system. On the other hand, the distribution curve of the Cu-Sal-Ser system is quite
different from the other two in that the formation of the Cu(Sal-Ser) complex
begins at pH ∼ 2.5 and Cu(Sal-Ser) is the predominant species (∼98%) in the
pH range 4–10. This difference, which results from the higher affinity of Cu(II)
for nitrogen and oxygen donors, may be evidence that Cu(II) is most effective
in catalyzing the deamination reactions, because these reactions proceed at an
optimum rate between pH 4 and 5 and are negligible above pH 9.(43)
4. CONCLUSION
The present results may have important biological implications. To interpret
properly the biological processes involving amino acids it is necessary to have
quantitative data on all pertinent equilibrium constants [i.e., the formation con-
stants of the Schiff base, the protonation constants of the Schiff base, the stability
constants of Schiff base metal complexes, and the protonation constants of the
amino acids and salicylaldehyde from which the Schiff bases are derived as well

Equilibrium Studies of Schiff Bases and Their Complexes
575
Fig 6. Species’ distribution curves of the system, Ni(II)-salicylaldehyde-serine, at a
1:1:1 molar ratio.
Fig 7. Species’ distribution curves for the system, Zn(II)-salicylaldehyde-serine,
at a 1:1:1 molar ratio.

576
Demirelli and Ko¨seog˘lu
Fig 8. Species’ distribution curves for the system, Cu(II)-salicylaldehyde-serine, at a 1:1:1
molar ratio.
as their metal binding constants]. By suitably varying the nature of the metal ion
and the experimental parameters it should be possible to increase considerably
our understanding of the chemical factors which are involved in more complicated
biological processes.
REFERENCES
1. R. D. Jones, D. A. Summerville, and F. Basolo, Chem. Rev. 79, 139 (1979).
2. G. Henrici-Olive and S. Olive, The Chemistry of the Catalyzed Hydrogenation of Carbon Monoxide
(Springer, Berlin, 1984).
3. H. Dugas and C. Penney, Bioorganic Chemistry (Springer, New York, 1981).
4. J. D. Margerum and L. J. Miller, Photochronism (Wiley, New York, 1971).
5. W. J. Sawodny and M. Riederer, Angew Chem., Int. Edn. Engl. 16, 214 (1977).
6. M. M. Shoukry and R. Koussini, Bull. Soc. Chim. Fr. 128, 465 (1991).
7. B. E. Fischer and H. Sigel, J. Am. Chem. Soc. 1029, 2998 (1980).
8. L. Leussing, in Metal Ions in Biological Systems, Vol. 5, Chap. 1, H. Sigel, ed. (Marcell Dekker,
New York, 1976).
9. R. G. Kallen, T. Korpela, A. E. Martell, Y. Matsushima, C. M. Metzler, and D. E. Metzler, in
Transaminases, Vol. 2, Chap. 2, P. Christen and D. E. Metzler, eds. (Willey, New York, 1985).
10. I. I. Mathews, P. Ajoy, S. Vasudevan, and H. Monahar, Inorg. Chem. 30, 2181 (1991).
11. J. W. Pyrz, A. I. Roe, L. J. Strern, and L. Que, Jr., J. Am. Chem. Soc. 107, 614 (1985).
12. A. E. Martell and B. Szpoganiez, Inorg. Chem. 28, 4199 (1989).
13. C. M. Metzler, A. Cahill, and D. Metzler, J. Am. Chem. Soc. 102, 6075 (1980).
14. D. L. Leussing and E. Hanna, J. Am. Chem. Soc. 88, 693 (1966).

Equilibrium Studies of Schiff Bases and Their Complexes
577
15. L. J. Nunez and G. L. Eichhorn, J. Am. Chem. Soc. 84, 901 (1962).
16. G. L. Eichhorn and J. W. Dawes, J. Am. Chem. Soc. 76, 5663 (1954).
17. K. S. Bai and D. L. Leussing, J. Am. Chem. Soc. 89, 6126 (1967).
18. D. Heinert and A. E. Martell, J. Am. Chem. Soc. 85, 188 (1963).
19. D. Hopgood and D. L Leussing, J. Am. Chem. Soc. 91, 3740 (1969).
20. D. L. Leussing, and K. S. Bai, Anal. Chem. 40, 575 (1968).
21. B. E. Leach and D. L. Leussing, J. Am. Chem. Soc. 93, 3377 (1971).
22. E. S. Jayadevappa and S. C. Galgali, J. Indian Chem. Soc. 60, 1098 (1983).
23. G. L. Eichhorn and N. D. Marchand, J. Am. Chem. Soc. 78, 2688 (1956).
24. C. M. Metzler, A. Cahill, and D. E. Metzler, J. Am. Chem. Soc. 102, 6075 (1980).
25. A. E. Martell and R. J. Motekaitis, The Determination and Use of Stability Constants (VCH, New
York, 1988).
26. R. J. Motekaitis and A. E. Martell, Can. J. Chem. 60, 2403 (1982).
27. D. D. Perrin and W. L. F. Armarego, Purification of Laboratory Chemicals, 3rd edn. (Pergamon,
New York, 1992).
28. H. Rossotti, Chemical Applications of Potentiometry (Van Mostrand, London, 1968).
29. G. Gran, Acta Chem. Scand. 4, 559 (1950).
30. G. Gran, Analyst 77, 661 (1952).
31. G. Schwarzenbach and M. Flascha, Complexometric Titrations (Methuen, London, 1957).
32. I. Sovago, T. Kiss, and A. Gergely, Pure and Appl. Chem. 65, 1030 (1993).
33. O. Yamauchi and A. Odani, Pure and Appl. Chem. 68, 469 (1996).
34. A. E. Martell and R. M. Simith, Critical Stability Constants, Amino Acids (Plenum, New York,
1974).
35. A. Gergely and T. Kiss, J. Inorg. Nucl. Chem. 39, 109 (1978).
36. F. Dallavalle, G. Folesoni, A. Sabatini, M. Tegoni, and A. Vacula, Polyhedron 20, 103 (2001).
37. A. E. Martell, in Chemical and Biological Aspects of Pyridoxal Catalysis, I.U.B. Symp. Ser. No.
30 (Pergamon, London, 1963).
38. S. Capasso, F. Giordano, C. Mattia, L. Mazzarella, and A. Ripamonti, J. Chem. Soc., Dalton Trans.
2228 (1974).
39. I. Bkouche-Waksman, J. M. Barbe, and A. Kvıck, Acta Cryst. Sect. B. 44, 595 (1988).
40. A. G. Raso, J. J. Fiol, A. L. Zafra, I. Mata, E. Espinosa, and E. Molins, Polyhedron 19, 673 (2000).
41. H. Irving and J. P. Williams, Nature 162, 746 (1948).
42. H. Irving and R. J. P. Williams, J. Chem. Soc. 3192 (1953).
43. V. B. Shanbhag, Metal Ion and Vitamin B6 Catalyzed Oxidative Determination of Some Amino
Acids in the Presence of Molecular Oxygen, unpublished Doctoral Thesis (Texas A&M University,
Texas, 1989).