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Thermodynamic stabilities and formation mechanisms of iron(III) complexes with Nα-carboxymethyl-Nα-salicyl-l -phenylalanine and Nα-salicyl-l -phenylalanine: Enhanced thermodynamic stability by additional methylcarboxylate group
Pergamon
0277-5387(94)EQO90-3
Polyhedron Vol. 13, No. 15116, pp. 2343-2351, 1994 Copyright 0 1994 Ekvier Science Ltd Printed in Great Britain. Au lights -ad 0277-5387/94 57.00+0.00
THERMODYNAMIC STABILITIES AND FORMATION MECHANISMS OF IRON(HI) COMPLEXES WITH N=CARBOXYMETHYL-N=SALICYL-L-PHENYLALANINE AND N=-SALICYL-L-PHENYLALANINE: ENHANCED THERMODYNAMIC STABILITY BY ADDITIONAL METHYLCARBOXYLATE GROUP TOMOHIRO
OZAWA, KOICHIRO JITSUKAWA, HISAHIKO EINAGA”
HIDEKI
MASUDA and
Department of Applied Chemistry, Faculty of Engineering, Nagoya Institute of Technology, Gokiso-cho, Showa-ku, Nagoya 466, Japan (Received 22 December 1993 ; accepted 23 February 1994)
Abstract-The thermodynamic stability and the formation reaction mechanism have been investigated for the iron(II1) complexes with a new tetradentate ligand, N”-carboxymethylN*-salicyl+phenylalanine (1, H,cslpa) and the related N”-salicyl-L-phenylalanine (2, Hzslpa). The mono(ligand)iron(III) complex of 1 formed under both conditions for the metal ion concentration (C,) being in large excess over the ligand 1 (CL) and Dice versa, while 2 formed a bis(ligand)iron(III) complex in consecutive stages under the conditions CL >>C,. The stability constant of the mono(ligand)iron(III) complex of 1 (log/$, = 21.61 fO.10 ; b,, = [Fe(cslpa)]/[Fe3+][cslpa3-1) was larger than the constant of the complex of 2 (log /IL1= 17.71 kO.09; /I,, = [Fe(slpa)+]/[Fe3+][slpa2-I), which was mainly caused by the chelate effect of an additional methylcarboxylate group. The complexes of 1 and 2 were formed through the composite pathways of Fe’+ +H,L-@’ ‘) (k32) and FeOH*+ + H3Lo(“’+)(k23) and another single pathway of FeOH’+ + H,L-@‘IO)(k22>[L3-(“’ ‘--) = cslpa3- (or slpa*-)]. The rate constants of each pathway for 1 and 2 are the same in magnitude as the corresponding constants of iron(III)-Na-pyridoxyl+aspartic acid (4, H,plaF) and-N”-salicyl-L-alanine systems, respectively. This good agreement in the rate constants suggests a mechanism that the complexes of 1 and 2 form through the slow coordination of the phenolato oxygen followed by the rapid donations of other coordinating atoms.
Microorganisms and plants synthesize and liberate some types of chelating compounds termed siderophores to incorporate effectively iron(III),’ which easily forms an insoluble hydroxide, Fe(OH), [(ldg Ksp = -39; Kp = [Fe3+][bH-]3)],2 in the physiological pH region. Investigations have been made for the thermodynamic stability and the formation reaction mechanisms of the iron(II1) complexes with some artificial chelating compounds that were composed of pyridoxal and some
*Author to whom correspondence should be addressed.
L-amino acids3,4with the aim to compare them with the natural siderophores and to design new chelating compounds as artificial siderophores. These ligands can form thermodynamically very stable iron(II1) complexes3*4 and their stability constants correspond to those of some ferrioxamines,’ which are natural siderophores. In a succeeding investigation, a new tetradentate ligand, N”-carboxymethyl-N”-salicyl+phenylalanine (1, H,cslpa) was synthesized by using salicylaldehyde instead of pyridoxal in order to compare the coordination behaviour to iron(II1) of the former with the pyridoxyl derivatives. The parent ligand, Na-
2343
2344
T. OZAWA et al.
salicyl-L-phenylalanine (2, H,slpa) was also investigated as a reference to 1.
N-
CHCOOH I
filtered off. The ethanol solution was evaporated to dryness and the resulting insoluble product was
NH -
/
CHCOOH I
CH,
NH -
/
CHCOOH I
NH / CH,
3 Hdw
EXPERIMENTAL Materials W-Salicyl-L-phenylalanine (2). L-Phenylalanine (5.0 g, 30.3 mmol) was dissolved in methanol (40 cm3) with an equimolar amount of potassium hydroxide. To this methanol solution was added salicylaldehyde (3.2 cm3, 30.6 mmol) and the resulting yellow solution was stirred for 10 min. Two equimolar amounts of sodium borohydride (2.34 g, 6.19 mmol) were then added with stirring for 20 min to give a clear solution. Precipitate was obtained by treating with ethanolic hydrochloride to the acidity of its isoelectric point. This precipitate was collected by filtration and washed twice with cold water to remove any inorganic salts. Analysis of 2 : Found : C, 68.3; H, 6.1; N, 5.0. Calc. for C,6H,7N03 *0.2H,O*O.lNaCl: C, 68.4; H, 6.2; N; 5.0%. N”-Carboxymethyl-W-salicyl-L-phenylalanine (1).
Ligand 2 (3.0 g, 11.1 mmol) was dissolved in water (40 cm3) by adjusting to pH 11 with an aqueous NaOH solution (1 mol dme3). To this solution was added sodium bromoacetate (2.0 g, 14.4 mmol) and the solution was stirred at 45-5O”C, maintaining the pH at 10-l 1. After more than 90% of the reaction was completed by checking with thin-layer chromatography, a precipitate was obtained at pH 2 by acidifying the reaction solution with an aqueous HCl solution (1 mol dm-‘). This solid compound was redissolved in a small amount of hot ethanol and any insoluble residue was immediately
CHCOOH I CH, - COOH
4 Was
recovered. Pure 1 was obtained by recrystallizing several times from water. Analysis of 1: Found : C, 65.0; H, 5.9; N, 4.2. Calc. for ClsHr9N05*0.1H20: C, 65.3; H, 5.8; N, 4.2%. All other materials were supplied by Wako Pure Chemicals Ind. Ltd, Osaka, and Tokyo Kasei Organic Chemicals, Co., Ltd, Tokyo and they were used without further purification. Measurements
Iron(II1) stock solution was prepared by dissolving an appropriate amount of (NH,) Fe(SO& * 12Hz0 in water to keep the solution slightly acidic (ca pH 2) by adding perchloric acid (l-5 cm3, 1 mol dme3). The ligands 1 and 2 were dissolved in water with the addition of an aqueous sodium hydroxide solution. Experiments were carried out at a constant ionic strength of sodium perchlorate (0.10 mol dme3) at 25.0 +O. 1°C. The hydrogen ion concentration was adjusted with a perchloric acid-sodium perchlorate system. The hydrogen ion concentration was calculated according to eq. (1). - log [H + 1 = PH,,,, + log fi+
(1)
Here, PH,,, and fn+ represent the measured pH value and the activity coefficient of H+, respectively ; logfn+ of -0.086 was adopted in this work. The apparatus used in this investigation has been described elsewhere.7
Thermodynamic stabilities and formation mechanisms of Fe”’ complexes RESULTS Electronic absorption spectra
The absorption spectra of 1,2 and their iron(II1) complexes are depicted in Fig. 1. The complexes as well as the ligands show similar spectra to each other. Intense absorption bands observed in the ultraviolet region of (35.M5.0) x lo3 cm-’ in all ligands and complexes can be assigned to the n* + R transition of the aromatic rings of the ligands. New absorption bands appear only in the spectra of the iron(II1) complexes in the lower energy region [(17.5-25.0) x lo3 cm-‘]. These absorption bands were assigned to a ligand-to-metal charge-transfer (CT) transition.3*4 Protonation
constants
The protonation constants of 1 and 2 were determined on the basis of the previously reported procedure* by using absorbance change vs -log [H+] data, which could be analysed to two protonation/deprotonation steps at the phenolate oxygen and amino nitrogen sites in the pH range 513. The wavelengths used here to determine the protonation constants were 285 and 284 nm for phenolate oxygen [K&,OHJ and 285 and 275 nm for amino nitrogen [Ka(amNHj] in 1 and 2, respectively. The protonation constant of the carboxylate oxygen of the amino acid moiety [K,,,,,] in 1 as well as 2 could not be determined spectrophotometrically because of little absorbance change
2345
with [H+]. Therefore, a potentiometric titration method was used to determine the K,(,cooHj with the SUPERQUAD program.’ Additionally, the protonation/deprotonation behaviour could not be observed for the newly inserted methylcarboxylate oxygen of 1; the constant, log K,~coonj, may be less than unity. The protonation constants thus determined are listed in Table 1.
Thermodynamic
stability
The pH dependence of the absorbance at the wavelength of the CT region of the iron(III)-1 and -2 systems are depicted in Fig. 2. The same pHdependent absorbance-change behaviour can be found in the iron(III)-1 system under both conditions of CM >>CL and CL >>CM [CM and CL represent the total concentrations of iron(II1) and ligand, respectively], indicating that only a mono(ligand)iron(III)-type complex can form in the acidity region ca pH l&3.0. This result was supported by another study using the method of continuous variations ;I0 the maximum absorbance showed at 0.5 on the molar ratio of C,/(C,+ CM). Therefore, the stability constant of the iron(II1) complex of 1 was determined spectrophotometrically under the conditions of CL >>CM in order to avoid the influence of the hydrolysis of iron(II1). The complex formation equilibrium of the iron(III)-1 system can be defined by equilibrium (I) : K&) Fe3+ + H,cslpa(“-‘I+ _ A Fe(cslpa) + nH+ (I)
CM = [Fe3+] + [Fe(cslpa)] CL = [H,cslpa] + [H,cslpa-] where the value of n is 2 or 3. In an analogous way as that of a previous literature,” eq. (2) log {(&lax -A1)/(A’
-&in)}
= log W+l”-‘(1 +f&ncm~,W+I) -hGeq(1)cL
Fig. 1. Electronic absorption spectra of 1 (3 ; - - -) and 2 (4; --) and mono(ligand)iron(III) complexes of 1 (1; -) and 2 (2 ; ---). 1 : [Fe(cslpa)] (PH 1.68) ; 2 : lFe(slpa)] + (pH 2.72) ; 3 : H&pa @H 1.35) ; 4 : H3slpa+ (PH 1.48).
(2)
can be derived on the basis of a materials balance and the protonation constants of 1. Here, A;,, A,i, and A’ represent the absorbances of solutions in which the complex forms quantitatively, free ligand alone is present and the complex and the free ligand coexist at [H+], respectively. Plotting of the left-hand term Lf(A’)] vs the hydrogen ion concentration term (g[H+]) of eq. (2) by using the data shown in Fig. 2 gave a straight line of the regression
2346
T. OZAWA et al. Table 1. Protonation
and stability constants
Ligand
log &phOH) (mol-’ dm’)
log &lM) (mol-’ dm3)
log &amcoo~j (mol-’ dm3)
Remarks
Ref.
1 (H+zslpa) 2 W&a)
11.09* 0.03 10.52+0.02
7.30+ 0.03 8.09f 0.02
2.09 +0.02= 2.58*0.02”
b 6
This work This work
Complex
log 811 (mol-’ dm3)
1% 812 (mol-* dm6)
5 [Fe(cslpa)] 6 lWW41 + ]Fe(IWpa)J+ ’ [Fe(hbida)ld
21.61fO.10 17.71+0.09 18.77f0.14 22.4
39.06kO.26
Remarks
Ref.
fill” = /L(1) B,,” =*/L(2)
This work This work 3
e
15
“The value was estimated by a potentiometric titration method ;0.10 mol dmP3 (KN03), 25°C. *O.lOmol dmW3(NaC10,),25°C. “[Fe(Hplpa)d + : bis(ligand)iron(III) complex of 3, with the pyridyl nitrogen protonating. d[Fe(hbida)] : mono(ligand)iron(III) complex of N-(o-hydroxybenzyl)iminodiacetic acid. ‘0.10 mol dm-‘(KNO,), 25°C.
equation off(A’) = ug[H+] +b with a being equal to 1.08, b to -0.45 and n- 3 and allowed us to estimate the equilibrium constant, I&(l), from its intercept. The number of (n- 1) in eq. (2), which represents the liberated proton number on the complex formation from the diprotonated ligand species of 1 at the amino nitrogen and phenolato oxygen, was 2, namely, the complex formation equi-
o”*w
0 0
*A
0
AA
d 8 A
0 A
0
A
0
A
A AA
0
Af+
0 1
3
2
PH
Fig. 2. Relation between absorbance and pH for the mono(ligand)iron(III) complexes of 1 and 2. 0 : 1; C‘>> c,; C,: 1.50x 10e3 mol dm-‘; CM: 1.00x 10V4 moldm-3;l:525mn;~:1;CM>>C,;CM:1.47x10-3 mol dme3; CL: 9.87 x lo-‘mol dmp3; 1: 525 nm; A: 2; C,>> CM; C,: 1.00x lo-’ mol dmp3; CM: 1.00x 10e4 mol dmp3 ; 1: 525 nm ; 0.10 mol dm-’ (NaClO& 25°C.
librium under the experimental conditions can be represented as equilibrium (I)’ : K&)
Fe3+ +H,cslpa-
+
Fe(cslpa) + 2H+.
(I)’
This result is in agreement with the proton liberations from the amino nitrogen and the phenolato oxygen of the diprotonated ligand species by the formation of the mono(ligand)iron(III) complex of 1. The stability constant of the mono(ligand) iron(II1) complex of 1 is calculated on the basis of eq. (3). /I 11(l)
=
Ka(phOH)~&mNH)-kkql
(1)
(3)
A remarkable spectral change in the CT region could not be observed in the region ca pH 2.0-5.0, indicating that the ligand 1 forms a mono(ligand)iron(III) complex in the region ca pH 3.G 5.0 as well as in that of ca pH 1.0-2.0. For the iron(III)-2 system, it was found (cf. Fig. 3) that the absorption peak in the CT region firstly increases from ca pH 2.0 to 3.0, followed by no change from cu pH 3.0 to 3.5 and again an increase from cu pH 3.5 to 5.0. This result is in contrast to that of the iron(II1) - 1 system and indicates that 2 forms firstly a mono(ligand)iron(III) complex in the pH region cu 2.0-3.0 and secondly a bis(ligand)iron(III) complex in the region cu 3.5-5.0 in a consecutive stage: the compositions of these two complexes were ascertained by a continuous variations method.‘O
Thermodynamic
stabilities and formation mechanisms of Fe”’ complexes
2347
2. The complex formation equilibrium (II) can be rewritten in equilibrium (II)‘. K&1 Fe3+ +H,slpae
600
Wavelengthlnm
B11m
Fig. 3. Dependence of absorption spectra in the CT region on pH for the iron(III)-2 system. a : pH 2.11; b : pH 2.60; c: pH 3.56; d: pH 4.08; e: pH 4.60; CL: 1.01 x lo-’ mol drne3 ; CM : 1.00 x lop4 mol dm-‘.
For the mono(ligand)iron(III) equilibrium (II)
(II)’
The value of 2 corresponds to the protons of phenolato oxygen and amino nitrogen, where the protonation/deprotonation behaviour of the carboxylate oxygen has been renormalized into lyeq(2) in the same way as in the iron(III)-1 system. The stability constant, b11 (2), can be calculated according to eq. (5) by the use of the equilibrium constant, Keq (2), estimated from the intercept in the plots of the absorbance term and the hydrogen ion concentration term in eq. (4).
I
so0
Fe(slpa)+ + 2H+
complex of 2,
=
GJdlOH~~a(amNH)~q
(2)
(5)
Any attempts to determine reliably the stability constant of the bis(ligand)iron(III) complex was unsuccessful because free iron(II1) species in equilibrium with the complex even under the conditions of C, >>CM tended to undergo complex hydrolysis reactions including the precipitation of Fe(OH), above pH 3.5.13 All the stability constants thus obtained are listed in Table 1.
Fe(OH)‘:_ i + H,slpa(“-2)+ e Fe(slpa)+ + (i+n - 3)H+
(II)
CM = [Fe’+] + [FeOH’+] + [Fe(slpa) ‘1 C, = [H,slpa+] + [H,slpa] can be defined in a way analogous to a previous paper,” where the values of i as well as n are 2 and 3. Equation (4) can be derived on the basis of equilibrium (II) :
-A2)l(A2 -4rin)} log { bGax = log[H+]“+i-3(1 +K,(,,,,[H+]) x (1 +%rI[H+I)-logK,(WL
(4)
by the same procedure as that of the iron(III)-1 system. Here, A:,, Amin and A2 correspond to those of A,!_, and Ati,,, and A’ in the iron(III)-1 system, respectively, and K& is the hydrolysis constant of iron(II1) [K&L = [FeOH2+][H+]/[Fe3+] ; log K& = -2.78 (0.1 mol dme3 (NaCIO,)), 25“C].” Plotting the relation between the absorbance term [AA’)] and the hydrogen ion concentration term (g[H+]) of eq. (4) gave a straight line of a regression equation off(A2) = ag[H+] + b, with a being equal to 1.02, b to - 3.95, and n + i of 5, indicating that the net proton number liberated by the complex formation, (n+i- 3), is equal to
Kinetics The kinetics of iron(II1) complex formation was investigated by a stopped-flow spectrophotometric method under the pseudo-first-order conditions C, >>C, at the wavelength 525 nm for the Fe(III)1 system and 514 nm for the Fe(III)-2 system. The relation of the absorbance change with time showed a single exponential curve for more than 3.5 halflife periods. This result indicates that the complex formation reaction proceeds through a single ratelimiting stage. The observed rate constant, kobs,was determined at the acidity region pH 1.2-2.5 for the iron(III)-1 system and pH 2.43.2 for the iron(III)2 system, where the species of Fe3+ and Fe(OH)‘+ for iron(II1) and H,L’(“’ +) and H,L+(“’ 2+) (L = cslpa3- or slpa2-) are considered to contribute to the kinetics based on the hydrolysis constant of iron(II1) and protonation constants of the ligands. Equation (6) can be derived for the iron(II1) complex formation with 1. Fe(OH);ti
$yU) + H,cslpac-3)+ --P Fe(cslpa) + (i+j-
3)H+
(6)
The kobawill be newly defined here for the iron(III)-
2348
T.
OZAWA et al.
1 system according to a previous paper ;4 it can be represented as given in eqs (7) and (8) Kbs = k,,,M)lC, [H+13
+k23
&(~~oH)&~NH)
x
(1
[H
[H+13
+ 1’ )
+~&IW+I)
(14)
CM = [Fe3+] + [Fe(OH)2+] + [Fe(slpa)+]
+ [k32 (1) + k23 (~)~%&MXIOH)] [H
+
(7)
kLbs= k33(1)K a(phOH) Ka(amNH) Ka(amCOOH)
X &ptto~)%(mN~)
a(2) = (Ka@hOH) Ka(amNH) Ko(amCOOH)
C, = [H,slpa+] + [H,slpa]
+ 12
(~)K~~%@~oH~K~(~~NH)[H+I
(8)
+Ka(pho~)%(am~~)[H+l~) x (1 +J%I/[H+I)
(9)
CM = [Fe3+] + [Fe(OH)2+] + [Fe(cslpa)]
by using k,(2) in place of k,(l) in eqs (7) and (8). A plot of the relation between k& calculated according to eq. (12) and [H+] indicated that the same form of the quadratic function [eq. (lo)] holds for the iron(III)-2 system after a least-squares treatment with a constant a of 6.03 x 1O22and that of b of 5.32 x 1O’9under the conditions C, = 2.50 x lop4 and CM = 2.50 x lo-’ mol dmw3. The pathways and rate constants thus determined are listed in Table 2 together with those of our concern.
C, = [H,cslpa] + [H,cslpa-] by using14kObswhen it is taken into consideration that the protonation/deprotonation of the ligand is much faster than the coordination process.” A plot of the relation between kobscalculated from eq. (7) and [H+] revealed that kobs could be represented by a quadratic function of [H+] passing through the origin using a least-squares treatment, namely eq. (lo) Y=aP+bX
(10)
with a representing a constant containing k32(1) and k,,(l) of 1.78 x 102’ and b representing one containing k22(1) of 8.48 x 10” gave a best-fit to the experimental data obtained under the conditions C, = 6.50 x 10e4 and CM = 5.00 x lo-’ mol dme3. Here, X and Y are [H+] and k&, respectively. The pathways and rate constants thus determined are listed in Table 2. Likewise, eq. (11) can be derived for the iron(III)2 system : equations (7) and (8) can also be obtained Fe(OH)\ti
$0) + HJ.slpaGV2)+Fe(slpa)+ + (i+jKbs = k,+(NCi_
3)H+
(11) (12)
k:bs = k33(~)Ka@hOH)&unNH) x
&amCOOH)
+
k23
W+13
+
W32(2)
(2)~~%z,amcOoH,1
x &(pho~)&(amN~)
[H
+k23(2)~~K~(ph~~)Ka(am~~)[H+l
+ 12
(13)
DISCUSSION Thermodynamic stability The thermodynamic stability constant of the mono(ligand)iron(III) complex of 1 [log fill(l) = 21.61+0.10] is comparable to that of the complex of the related tetradentate ligand, N-(0hydroxybenzyl)iminodiacetic acid [H,hbida : log /?,, = 22.4; fill = [Fe(hbida)]/[Fe3+][hbida3-] ; 0.1 mol dme3(KN03), 25”C]15and it is larger than the constants of the mono(ligand)iron(III) complexes of 2 and 3 that are terdentate. This result supports the coordination of the methylcarboxylate group that caused the chelate effect (cf. Fig. 4). This chelate effect is further reflected in the equilibria (I)’ and (II)’ ; a mono(ligand)iron(III) complex forms quantitatively as low as pH 1.2 for 1, whereas it forms just as pH 2.7 for 2 (cf. Fig. 2). Such a remarkable difference in complex formation behaviour has not been observed between the iron(III)3 and -4 systems3 that also have different numbers of chelate rings. This difference may probably be caused by the different structures of these two tetradentate ligands through a strain due to the six-membered chelate ring. It is interesting to discuss here the difference in the complex formation behaviour between 1 and 4. Ligand 1 forms only a mono(ligand)iron(III) complex even under the conditions CL >>CM, while ligand 4 can form a bis(ligand)iron(III) complex under the same conditions. This difference in the behaviour can be assumed to be responsible for the ligand structures of 1 (a salicylaldehyde derivative) and 4 (a pyridoxal derivative) : ligand 4 can form a CH-n intramolecular interaction in its bis(lig-
2349
Thermodynamic stabilities and formation mechanisms of Fe”’ complexes Table 2. Pathways and formation rate constants Pathways
Ligand 1 (H,cslpa)
2 (H+lpa)
3 (H,plpa)
4 (H,plas)
H,slal’
H,mugnY Cd&OH (Hph)
Fe3++ Hgslpa- (ks2) and Fe(OH)‘+ + H,cslpa (kz3) Fe(OH)‘+ + Hsslpa (k& Fe’+ + H,slpa (kx2) and Fe(OH)*+ + H,slpa+ (kz3) Fe(OH)‘+ + H,slpa (k2J Fe3++ Hsplpa+(k,,) and Fe(OH)‘+ + H,plpa*+ (ku) Fe(OH)‘+ + H,plpa (k,,) Fe3++ H,plas (k,,) and Fe(OH)‘+ + H,plas+ (kza) Fe(OH)‘+ + H,plas (k,,) Fe3++ H,slal (k3J and Fe(OH)‘+ +H3slal+ (k,,) Fe(OH)‘+ + H,slal (k& Fe3++ H,mugm (k33 Fe(OH)2+ + H,mugm- (kz2) Fe’+ +Hph (k,,) Fe(OH)2+ +Hph (k2,)
k,
(mol-’ dm3 s-r)
Remarks
Ref.
k326 = k,,(l), ku = k,(l)
This work
kzb = k,,(I) k326= k&2), kz = k,,(2)
This work This work
k,P = k&2)
This work 3
(7.26f 0.20) x lo2 0 k2,:(2.08
f0.30) x 10”
(1.48f0.20) x 104’ k2,:(7.87f0.30)
x
103d
b
(1.08f0.20) x 1o*=f k,,:(1.27&0.30) x 10” (5.22f0.20) x 10’g*h
b
k2,:(1.31kO.30) x lo39 (5.43kO.20) x 103”
b
k2,:(3.18f0.30) x 10’ k,,:(6.17&0.20) x lo2 k,,:(1.92*0.30) x 102 h : -25 k2, : 7.2 x lo2 1.1 x 102 1.5 x 102
b
b
3 3
3 18
18 14 14 19 19 20 21
%2+ h&mco,,K& bO.10mol dm-’ (NaClO,), 25°C. ‘k3,+ k&,,,,Kb:, = (6.90f0.20) x 10’ mol-’ dm3 s-l. dk3r+ kz&
and)iron(III) complex, as has been suggested previously3 by the use of the methyl group of pyridoxal moiety [cf. Fig. 4 (8)], whereas 1 cannot form this type of interaction as is easily recognized by a molecular model consideration. In fact, the X-ray structurer6 of the bis(ligand)cobalt(III) complexes of N”-pyridoxyl+alanine revealed that the pyridine rings of pyridoxal moieties were placed close to each other through the CH-x intramolecular interaction of the methyl group of the pyridoxal moiety with the aromatic ring. This intramolecular interaction is expected to probably increase the
L-
thermodynamic stability and the formation of bis(ligand)iron(III) complex with 4. Additionally, the less strain of the chelate ring structure of 5 compared with 7 (de supra; cf. Fig. 4) may also enhance the thermodynamic stability of the mono(ligand)iron(III) complex of 1. The intramolecular interactions that have been found in the X-ray structun? are also expected in the bis(ligand)iron(III) complex of 3. Therefore, the difference of the complex formation behaviour that 3 can form a bis(ligand)iron(III) complex in a single step in contrast to 2 in consecutive stages
2350
T. OZAWA et al.
B
[Wcslpa)l
II
F-C", ,P
Fig. 4. The proposed structures of mono(ligand)iron(III)
under the conditions of CL >. CM may reasonably be explained. Kinetics and mechanisms
For the tetradentate ligand 1, the complex formation rate constant of the composite pathways of Fe3+ + H,cslpaa[k,, (l)] and Fe(OH)2+ + H3cslpa ]k23 W1[k32(1)+k23 ~~~~~~~~~~~~~ K& = (7.26f0.20) x lo2 mol-’ dm3 s-i] is a reasonable magnitude as a rate constant of iron(II1) (cf. Table 2). The rate constant is in the same magnitude as that of the pyridoxal-derived ligand, 3. Furthermore, the constant of the single pathway of Fe(OH)2+ + H,cslpa- [k,,(l)] [k,,(l) = (2.08 f0.30) x lo3 mol-’ dm3 s-‘1 corresponds in magnitude to that of the same pathway of 33 (cf. Table 2). These facts indicate that the complex formation mechanism of 1 is the same as that of the iron(III)-3 system. For the terdentate ligand 2, on the other hand, the complex formation rate constant of the composite pathways of Fe3+ + H,slpa [k32(2)] and Fe(OH)2+ +H,slpa+ [k23(2)] (k32(2)+k23(2) K .CamCOOH&,~ = (1.48 + 0.20) x lo4 mol- ’ dm3 s-l] is considerably larger than the corresponding constants of 1 and 33 (cf. Table 2). Incidentally, much of the literature3*4~‘4*‘7-z’ has reported the rate constants of Fe3+- and Fe(OH)‘+-related pathways, indicating that the Fe(OH)2+-related
P
\
,COOH
C”
complexes of 1,2and 4.
pathways are generally one or two orders of magnitude larger in the constants than the Fe3+related pathways for the same ligand,17 so long as there are no restrictions such as steric hindrance due to the ligand structure, etc. Accordingly, the Fe(OH)‘+-related pathways are mainly expected to contribute to the composite pathways for the system with 2. The fact that the rate constants of the single pathways [k,,(l) and k,,(2)] of 1 and 2 related only to Fe(OH)2f are close in magnitude to the composite constants of k32(2) + k,,(2) K aCamCOOH&,h of the system with 2 allows us to substantiate the discussion given above. The comparison of the rate constants of k,,(l) for 1 and k,,(2) for 2 with those of the corresponding pathways of the iron(III)-N”-pyridoxyl-type ligand systems3s4suffices for the deduction of complex formation mechanism. The fact that the values of k,,(l) for 1 and k22(2) for 2 are the same in magnitude allows us to propose the resemblance in the complex formation mechanism for the iron(II1) - 1 and -2 systems. Therefore, both of the complex formation reactions of 1 and of 2 proceed equally through the slow coordination of the phenolato oxygen to [Fe(H20)J3+ or [Fe(OH)(H20),]2+ species, followed by the rapid donations of other coordinating atoms as shown in Scheme I. In conclusion, a new tetradentate ligand (1) exhibits an increased thermodynamic stability of the mono(ligand)iron(III) complex compared with its
Thermodynamic
stabilities and formation mechanisms of Fe”’ complexes
CHz-N+-CHCOOH
2351
CH2-N+-CHCOOH
+H+ +H+
5 (R=CH&Oo-
1
B (R=H) /
-0.
+
FeOHat
/
r.d.8
Scheme I parent terdentate ligand 2 ; this increased stability is mainly due to the chelate effect by a newly introduced methylcarboxylate group. The complex formation mechanism for 1, however, remains unchanged in the case of the mechanism for 2. Thus, the introduction of the methylcarboxylate group to a phenol-amino acid-type ligand is recommended to increase the thermodynamic stability towards iron(II1) without changing the formation reaction mechanism. Acknowledgements-We wish to express our gratitude to Professor 0. Yamauchi and Mr T. Yajima of Nagoya University for their help to determine the protonation constants by the potentiometric titration method.
REFERENCES 1. C. Chimiak, R. C. Hider, A. Lui, J. B. Neilands, K. Nomoto and Y. Sugiura, Siderophores from Microorganisms and Plants. Springer, Berlin (1984). W. M. Latimer, Oxidation Potential. Prentice-Hall,
Englewood, NJ (1974). T. Ozawa, K. Jitsukawa, H. Masuda and H. Einaga, Ber. Buns. Phys. Chem. 1994,98,66. T. Ozawa, K. Jitsukawa and H. Einaga, $er. Buns. Phys. Chem. 1994,98,59. S. Konetchny-Rapp, G. Jung, K. N. Raymond, J. Meiwes and H. Z&ner, J. Am. Chem. Sot. 1992,114, 2224. 6. J. Kielland, J. Am. Chem. Sot. 1937,59, 1675.
7. T. Ozawa, K. Jitsukawa, H. Masuda and H. Einaga, Polyhedron, submitted. 8. Y. Hara, K. Okamoto, J. Hidaka and H. Einaga, Bull. Chem. Sot. Japan 1984,63, 1211. 9. P. Gans, A. Sabatini and A. Vacca, J. Chem. Sot., Dalton Trans. 1985, 1195. 10. H. L. Schlafer, Komplexbildung in Liisung, Ch. 8.
Springer, Berlin (1960). 11. M. Noritake, K. Okamoto, J. Hidaka and H. Einaga, Bull. Chem. Sot. Japan 1990,63, 353. 12. L. G. Sill&n and A. E. Martell, Stability Constants of Metal-Zon Complexes, Sp. Pub. No. 17, p. 53. Chemical Society, London (1964). 13. C. F. Baes Jr and R. E. Mesmer, The Hydrolysis of Cations, p. 237. John Wiley & Sons, New York (1976). 14. H. Yoshida, T. Ozawa, K. Jitsukawa and H. Einaga, Polyhedron 1993, 12, 1319. 15. W. R. Harris, R. J. Motekaitis and A. E. Martell, Znorg. Chem. 1975, 14,974. 16. K. Jitsukawa, K. Iwai, H. Masuda, H. Ogoshi and H. Einaga, Chem. Lett. 1994,303. 17. D. W. Margerum, G. R. Cayley, D. C. Weatherbum and G. K. Pagenkopf, Coordination Chemistry (Edited by A. E. Martell), Vol. 2, Ch. 1. American Chemical Society, Washington, DC (1978). 18. T. Ozawa, K. Iwai, K. Jitsukawa, H. Masuda and H. Einaga, Polyhedron, in press. 19. S. Gouder and J. Stuehr, Znorg. Chem. 1974,13,379. 20. K. Nakamura, T. Tsucbida, A. Yamagishi and M. Fujimoto, BUN. Chem. Sot. Japan 1973,46,456. 21. F. P. Cavacino and E. D. Dio, J. Chem. Sot. A 1970, 1151.
0277-5387(94)EQO90-3
Polyhedron Vol. 13, No. 15116, pp. 2343-2351, 1994 Copyright 0 1994 Ekvier Science Ltd Printed in Great Britain. Au lights -ad 0277-5387/94 57.00+0.00
THERMODYNAMIC STABILITIES AND FORMATION MECHANISMS OF IRON(HI) COMPLEXES WITH N=CARBOXYMETHYL-N=SALICYL-L-PHENYLALANINE AND N=-SALICYL-L-PHENYLALANINE: ENHANCED THERMODYNAMIC STABILITY BY ADDITIONAL METHYLCARBOXYLATE GROUP TOMOHIRO
OZAWA, KOICHIRO JITSUKAWA, HISAHIKO EINAGA”
HIDEKI
MASUDA and
Department of Applied Chemistry, Faculty of Engineering, Nagoya Institute of Technology, Gokiso-cho, Showa-ku, Nagoya 466, Japan (Received 22 December 1993 ; accepted 23 February 1994)
Abstract-The thermodynamic stability and the formation reaction mechanism have been investigated for the iron(II1) complexes with a new tetradentate ligand, N”-carboxymethylN*-salicyl+phenylalanine (1, H,cslpa) and the related N”-salicyl-L-phenylalanine (2, Hzslpa). The mono(ligand)iron(III) complex of 1 formed under both conditions for the metal ion concentration (C,) being in large excess over the ligand 1 (CL) and Dice versa, while 2 formed a bis(ligand)iron(III) complex in consecutive stages under the conditions CL >>C,. The stability constant of the mono(ligand)iron(III) complex of 1 (log/$, = 21.61 fO.10 ; b,, = [Fe(cslpa)]/[Fe3+][cslpa3-1) was larger than the constant of the complex of 2 (log /IL1= 17.71 kO.09; /I,, = [Fe(slpa)+]/[Fe3+][slpa2-I), which was mainly caused by the chelate effect of an additional methylcarboxylate group. The complexes of 1 and 2 were formed through the composite pathways of Fe’+ +H,L-@’ ‘) (k32) and FeOH*+ + H3Lo(“’+)(k23) and another single pathway of FeOH’+ + H,L-@‘IO)(k22>[L3-(“’ ‘--) = cslpa3- (or slpa*-)]. The rate constants of each pathway for 1 and 2 are the same in magnitude as the corresponding constants of iron(III)-Na-pyridoxyl+aspartic acid (4, H,plaF) and-N”-salicyl-L-alanine systems, respectively. This good agreement in the rate constants suggests a mechanism that the complexes of 1 and 2 form through the slow coordination of the phenolato oxygen followed by the rapid donations of other coordinating atoms.
Microorganisms and plants synthesize and liberate some types of chelating compounds termed siderophores to incorporate effectively iron(III),’ which easily forms an insoluble hydroxide, Fe(OH), [(ldg Ksp = -39; Kp = [Fe3+][bH-]3)],2 in the physiological pH region. Investigations have been made for the thermodynamic stability and the formation reaction mechanisms of the iron(II1) complexes with some artificial chelating compounds that were composed of pyridoxal and some
*Author to whom correspondence should be addressed.
L-amino acids3,4with the aim to compare them with the natural siderophores and to design new chelating compounds as artificial siderophores. These ligands can form thermodynamically very stable iron(II1) complexes3*4 and their stability constants correspond to those of some ferrioxamines,’ which are natural siderophores. In a succeeding investigation, a new tetradentate ligand, N”-carboxymethyl-N”-salicyl+phenylalanine (1, H,cslpa) was synthesized by using salicylaldehyde instead of pyridoxal in order to compare the coordination behaviour to iron(II1) of the former with the pyridoxyl derivatives. The parent ligand, Na-
2343
2344
T. OZAWA et al.
salicyl-L-phenylalanine (2, H,slpa) was also investigated as a reference to 1.
N-
CHCOOH I
filtered off. The ethanol solution was evaporated to dryness and the resulting insoluble product was
NH -
/
CHCOOH I
CH,
NH -
/
CHCOOH I
NH / CH,
3 Hdw
EXPERIMENTAL Materials W-Salicyl-L-phenylalanine (2). L-Phenylalanine (5.0 g, 30.3 mmol) was dissolved in methanol (40 cm3) with an equimolar amount of potassium hydroxide. To this methanol solution was added salicylaldehyde (3.2 cm3, 30.6 mmol) and the resulting yellow solution was stirred for 10 min. Two equimolar amounts of sodium borohydride (2.34 g, 6.19 mmol) were then added with stirring for 20 min to give a clear solution. Precipitate was obtained by treating with ethanolic hydrochloride to the acidity of its isoelectric point. This precipitate was collected by filtration and washed twice with cold water to remove any inorganic salts. Analysis of 2 : Found : C, 68.3; H, 6.1; N, 5.0. Calc. for C,6H,7N03 *0.2H,O*O.lNaCl: C, 68.4; H, 6.2; N; 5.0%. N”-Carboxymethyl-W-salicyl-L-phenylalanine (1).
Ligand 2 (3.0 g, 11.1 mmol) was dissolved in water (40 cm3) by adjusting to pH 11 with an aqueous NaOH solution (1 mol dme3). To this solution was added sodium bromoacetate (2.0 g, 14.4 mmol) and the solution was stirred at 45-5O”C, maintaining the pH at 10-l 1. After more than 90% of the reaction was completed by checking with thin-layer chromatography, a precipitate was obtained at pH 2 by acidifying the reaction solution with an aqueous HCl solution (1 mol dm-‘). This solid compound was redissolved in a small amount of hot ethanol and any insoluble residue was immediately
CHCOOH I CH, - COOH
4 Was
recovered. Pure 1 was obtained by recrystallizing several times from water. Analysis of 1: Found : C, 65.0; H, 5.9; N, 4.2. Calc. for ClsHr9N05*0.1H20: C, 65.3; H, 5.8; N, 4.2%. All other materials were supplied by Wako Pure Chemicals Ind. Ltd, Osaka, and Tokyo Kasei Organic Chemicals, Co., Ltd, Tokyo and they were used without further purification. Measurements
Iron(II1) stock solution was prepared by dissolving an appropriate amount of (NH,) Fe(SO& * 12Hz0 in water to keep the solution slightly acidic (ca pH 2) by adding perchloric acid (l-5 cm3, 1 mol dme3). The ligands 1 and 2 were dissolved in water with the addition of an aqueous sodium hydroxide solution. Experiments were carried out at a constant ionic strength of sodium perchlorate (0.10 mol dme3) at 25.0 +O. 1°C. The hydrogen ion concentration was adjusted with a perchloric acid-sodium perchlorate system. The hydrogen ion concentration was calculated according to eq. (1). - log [H + 1 = PH,,,, + log fi+
(1)
Here, PH,,, and fn+ represent the measured pH value and the activity coefficient of H+, respectively ; logfn+ of -0.086 was adopted in this work. The apparatus used in this investigation has been described elsewhere.7
Thermodynamic stabilities and formation mechanisms of Fe”’ complexes RESULTS Electronic absorption spectra
The absorption spectra of 1,2 and their iron(II1) complexes are depicted in Fig. 1. The complexes as well as the ligands show similar spectra to each other. Intense absorption bands observed in the ultraviolet region of (35.M5.0) x lo3 cm-’ in all ligands and complexes can be assigned to the n* + R transition of the aromatic rings of the ligands. New absorption bands appear only in the spectra of the iron(II1) complexes in the lower energy region [(17.5-25.0) x lo3 cm-‘]. These absorption bands were assigned to a ligand-to-metal charge-transfer (CT) transition.3*4 Protonation
constants
The protonation constants of 1 and 2 were determined on the basis of the previously reported procedure* by using absorbance change vs -log [H+] data, which could be analysed to two protonation/deprotonation steps at the phenolate oxygen and amino nitrogen sites in the pH range 513. The wavelengths used here to determine the protonation constants were 285 and 284 nm for phenolate oxygen [K&,OHJ and 285 and 275 nm for amino nitrogen [Ka(amNHj] in 1 and 2, respectively. The protonation constant of the carboxylate oxygen of the amino acid moiety [K,,,,,] in 1 as well as 2 could not be determined spectrophotometrically because of little absorbance change
2345
with [H+]. Therefore, a potentiometric titration method was used to determine the K,(,cooHj with the SUPERQUAD program.’ Additionally, the protonation/deprotonation behaviour could not be observed for the newly inserted methylcarboxylate oxygen of 1; the constant, log K,~coonj, may be less than unity. The protonation constants thus determined are listed in Table 1.
Thermodynamic
stability
The pH dependence of the absorbance at the wavelength of the CT region of the iron(III)-1 and -2 systems are depicted in Fig. 2. The same pHdependent absorbance-change behaviour can be found in the iron(III)-1 system under both conditions of CM >>CL and CL >>CM [CM and CL represent the total concentrations of iron(II1) and ligand, respectively], indicating that only a mono(ligand)iron(III)-type complex can form in the acidity region ca pH l&3.0. This result was supported by another study using the method of continuous variations ;I0 the maximum absorbance showed at 0.5 on the molar ratio of C,/(C,+ CM). Therefore, the stability constant of the iron(II1) complex of 1 was determined spectrophotometrically under the conditions of CL >>CM in order to avoid the influence of the hydrolysis of iron(II1). The complex formation equilibrium of the iron(III)-1 system can be defined by equilibrium (I) : K&) Fe3+ + H,cslpa(“-‘I+ _ A Fe(cslpa) + nH+ (I)
CM = [Fe3+] + [Fe(cslpa)] CL = [H,cslpa] + [H,cslpa-] where the value of n is 2 or 3. In an analogous way as that of a previous literature,” eq. (2) log {(&lax -A1)/(A’
-&in)}
= log W+l”-‘(1 +f&ncm~,W+I) -hGeq(1)cL
Fig. 1. Electronic absorption spectra of 1 (3 ; - - -) and 2 (4; --) and mono(ligand)iron(III) complexes of 1 (1; -) and 2 (2 ; ---). 1 : [Fe(cslpa)] (PH 1.68) ; 2 : lFe(slpa)] + (pH 2.72) ; 3 : H&pa @H 1.35) ; 4 : H3slpa+ (PH 1.48).
(2)
can be derived on the basis of a materials balance and the protonation constants of 1. Here, A;,, A,i, and A’ represent the absorbances of solutions in which the complex forms quantitatively, free ligand alone is present and the complex and the free ligand coexist at [H+], respectively. Plotting of the left-hand term Lf(A’)] vs the hydrogen ion concentration term (g[H+]) of eq. (2) by using the data shown in Fig. 2 gave a straight line of the regression
2346
T. OZAWA et al. Table 1. Protonation
and stability constants
Ligand
log &phOH) (mol-’ dm’)
log &lM) (mol-’ dm3)
log &amcoo~j (mol-’ dm3)
Remarks
Ref.
1 (H+zslpa) 2 W&a)
11.09* 0.03 10.52+0.02
7.30+ 0.03 8.09f 0.02
2.09 +0.02= 2.58*0.02”
b 6
This work This work
Complex
log 811 (mol-’ dm3)
1% 812 (mol-* dm6)
5 [Fe(cslpa)] 6 lWW41 + ]Fe(IWpa)J+ ’ [Fe(hbida)ld
21.61fO.10 17.71+0.09 18.77f0.14 22.4
39.06kO.26
Remarks
Ref.
fill” = /L(1) B,,” =*/L(2)
This work This work 3
e
15
“The value was estimated by a potentiometric titration method ;0.10 mol dmP3 (KN03), 25°C. *O.lOmol dmW3(NaC10,),25°C. “[Fe(Hplpa)d + : bis(ligand)iron(III) complex of 3, with the pyridyl nitrogen protonating. d[Fe(hbida)] : mono(ligand)iron(III) complex of N-(o-hydroxybenzyl)iminodiacetic acid. ‘0.10 mol dm-‘(KNO,), 25°C.
equation off(A’) = ug[H+] +b with a being equal to 1.08, b to -0.45 and n- 3 and allowed us to estimate the equilibrium constant, I&(l), from its intercept. The number of (n- 1) in eq. (2), which represents the liberated proton number on the complex formation from the diprotonated ligand species of 1 at the amino nitrogen and phenolato oxygen, was 2, namely, the complex formation equi-
o”*w
0 0
*A
0
AA
d 8 A
0 A
0
A
0
A
A AA
0
Af+
0 1
3
2
PH
Fig. 2. Relation between absorbance and pH for the mono(ligand)iron(III) complexes of 1 and 2. 0 : 1; C‘>> c,; C,: 1.50x 10e3 mol dm-‘; CM: 1.00x 10V4 moldm-3;l:525mn;~:1;CM>>C,;CM:1.47x10-3 mol dme3; CL: 9.87 x lo-‘mol dmp3; 1: 525 nm; A: 2; C,>> CM; C,: 1.00x lo-’ mol dmp3; CM: 1.00x 10e4 mol dmp3 ; 1: 525 nm ; 0.10 mol dm-’ (NaClO& 25°C.
librium under the experimental conditions can be represented as equilibrium (I)’ : K&)
Fe3+ +H,cslpa-
+
Fe(cslpa) + 2H+.
(I)’
This result is in agreement with the proton liberations from the amino nitrogen and the phenolato oxygen of the diprotonated ligand species by the formation of the mono(ligand)iron(III) complex of 1. The stability constant of the mono(ligand) iron(II1) complex of 1 is calculated on the basis of eq. (3). /I 11(l)
=
Ka(phOH)~&mNH)-kkql
(1)
(3)
A remarkable spectral change in the CT region could not be observed in the region ca pH 2.0-5.0, indicating that the ligand 1 forms a mono(ligand)iron(III) complex in the region ca pH 3.G 5.0 as well as in that of ca pH 1.0-2.0. For the iron(III)-2 system, it was found (cf. Fig. 3) that the absorption peak in the CT region firstly increases from ca pH 2.0 to 3.0, followed by no change from cu pH 3.0 to 3.5 and again an increase from cu pH 3.5 to 5.0. This result is in contrast to that of the iron(II1) - 1 system and indicates that 2 forms firstly a mono(ligand)iron(III) complex in the pH region cu 2.0-3.0 and secondly a bis(ligand)iron(III) complex in the region cu 3.5-5.0 in a consecutive stage: the compositions of these two complexes were ascertained by a continuous variations method.‘O
Thermodynamic
stabilities and formation mechanisms of Fe”’ complexes
2347
2. The complex formation equilibrium (II) can be rewritten in equilibrium (II)‘. K&1 Fe3+ +H,slpae
600
Wavelengthlnm
B11m
Fig. 3. Dependence of absorption spectra in the CT region on pH for the iron(III)-2 system. a : pH 2.11; b : pH 2.60; c: pH 3.56; d: pH 4.08; e: pH 4.60; CL: 1.01 x lo-’ mol drne3 ; CM : 1.00 x lop4 mol dm-‘.
For the mono(ligand)iron(III) equilibrium (II)
(II)’
The value of 2 corresponds to the protons of phenolato oxygen and amino nitrogen, where the protonation/deprotonation behaviour of the carboxylate oxygen has been renormalized into lyeq(2) in the same way as in the iron(III)-1 system. The stability constant, b11 (2), can be calculated according to eq. (5) by the use of the equilibrium constant, Keq (2), estimated from the intercept in the plots of the absorbance term and the hydrogen ion concentration term in eq. (4).
I
so0
Fe(slpa)+ + 2H+
complex of 2,
=
GJdlOH~~a(amNH)~q
(2)
(5)
Any attempts to determine reliably the stability constant of the bis(ligand)iron(III) complex was unsuccessful because free iron(II1) species in equilibrium with the complex even under the conditions of C, >>CM tended to undergo complex hydrolysis reactions including the precipitation of Fe(OH), above pH 3.5.13 All the stability constants thus obtained are listed in Table 1.
Fe(OH)‘:_ i + H,slpa(“-2)+ e Fe(slpa)+ + (i+n - 3)H+
(II)
CM = [Fe’+] + [FeOH’+] + [Fe(slpa) ‘1 C, = [H,slpa+] + [H,slpa] can be defined in a way analogous to a previous paper,” where the values of i as well as n are 2 and 3. Equation (4) can be derived on the basis of equilibrium (II) :
-A2)l(A2 -4rin)} log { bGax = log[H+]“+i-3(1 +K,(,,,,[H+]) x (1 +%rI[H+I)-logK,(WL
(4)
by the same procedure as that of the iron(III)-1 system. Here, A:,, Amin and A2 correspond to those of A,!_, and Ati,,, and A’ in the iron(III)-1 system, respectively, and K& is the hydrolysis constant of iron(II1) [K&L = [FeOH2+][H+]/[Fe3+] ; log K& = -2.78 (0.1 mol dme3 (NaCIO,)), 25“C].” Plotting the relation between the absorbance term [AA’)] and the hydrogen ion concentration term (g[H+]) of eq. (4) gave a straight line of a regression equation off(A2) = ag[H+] + b, with a being equal to 1.02, b to - 3.95, and n + i of 5, indicating that the net proton number liberated by the complex formation, (n+i- 3), is equal to
Kinetics The kinetics of iron(II1) complex formation was investigated by a stopped-flow spectrophotometric method under the pseudo-first-order conditions C, >>C, at the wavelength 525 nm for the Fe(III)1 system and 514 nm for the Fe(III)-2 system. The relation of the absorbance change with time showed a single exponential curve for more than 3.5 halflife periods. This result indicates that the complex formation reaction proceeds through a single ratelimiting stage. The observed rate constant, kobs,was determined at the acidity region pH 1.2-2.5 for the iron(III)-1 system and pH 2.43.2 for the iron(III)2 system, where the species of Fe3+ and Fe(OH)‘+ for iron(II1) and H,L’(“’ +) and H,L+(“’ 2+) (L = cslpa3- or slpa2-) are considered to contribute to the kinetics based on the hydrolysis constant of iron(II1) and protonation constants of the ligands. Equation (6) can be derived for the iron(II1) complex formation with 1. Fe(OH);ti
$yU) + H,cslpac-3)+ --P Fe(cslpa) + (i+j-
3)H+
(6)
The kobawill be newly defined here for the iron(III)-
2348
T.
OZAWA et al.
1 system according to a previous paper ;4 it can be represented as given in eqs (7) and (8) Kbs = k,,,M)lC, [H+13
+k23
&(~~oH)&~NH)
x
(1
[H
[H+13
+ 1’ )
+~&IW+I)
(14)
CM = [Fe3+] + [Fe(OH)2+] + [Fe(slpa)+]
+ [k32 (1) + k23 (~)~%&MXIOH)] [H
+
(7)
kLbs= k33(1)K a(phOH) Ka(amNH) Ka(amCOOH)
X &ptto~)%(mN~)
a(2) = (Ka@hOH) Ka(amNH) Ko(amCOOH)
C, = [H,slpa+] + [H,slpa]
+ 12
(~)K~~%@~oH~K~(~~NH)[H+I
(8)
+Ka(pho~)%(am~~)[H+l~) x (1 +J%I/[H+I)
(9)
CM = [Fe3+] + [Fe(OH)2+] + [Fe(cslpa)]
by using k,(2) in place of k,(l) in eqs (7) and (8). A plot of the relation between k& calculated according to eq. (12) and [H+] indicated that the same form of the quadratic function [eq. (lo)] holds for the iron(III)-2 system after a least-squares treatment with a constant a of 6.03 x 1O22and that of b of 5.32 x 1O’9under the conditions C, = 2.50 x lop4 and CM = 2.50 x lo-’ mol dmw3. The pathways and rate constants thus determined are listed in Table 2 together with those of our concern.
C, = [H,cslpa] + [H,cslpa-] by using14kObswhen it is taken into consideration that the protonation/deprotonation of the ligand is much faster than the coordination process.” A plot of the relation between kobscalculated from eq. (7) and [H+] revealed that kobs could be represented by a quadratic function of [H+] passing through the origin using a least-squares treatment, namely eq. (lo) Y=aP+bX
(10)
with a representing a constant containing k32(1) and k,,(l) of 1.78 x 102’ and b representing one containing k22(1) of 8.48 x 10” gave a best-fit to the experimental data obtained under the conditions C, = 6.50 x 10e4 and CM = 5.00 x lo-’ mol dme3. Here, X and Y are [H+] and k&, respectively. The pathways and rate constants thus determined are listed in Table 2. Likewise, eq. (11) can be derived for the iron(III)2 system : equations (7) and (8) can also be obtained Fe(OH)\ti
$0) + HJ.slpaGV2)+Fe(slpa)+ + (i+jKbs = k,+(NCi_
3)H+
(11) (12)
k:bs = k33(~)Ka@hOH)&unNH) x
&amCOOH)
+
k23
W+13
+
W32(2)
(2)~~%z,amcOoH,1
x &(pho~)&(amN~)
[H
+k23(2)~~K~(ph~~)Ka(am~~)[H+l
+ 12
(13)
DISCUSSION Thermodynamic stability The thermodynamic stability constant of the mono(ligand)iron(III) complex of 1 [log fill(l) = 21.61+0.10] is comparable to that of the complex of the related tetradentate ligand, N-(0hydroxybenzyl)iminodiacetic acid [H,hbida : log /?,, = 22.4; fill = [Fe(hbida)]/[Fe3+][hbida3-] ; 0.1 mol dme3(KN03), 25”C]15and it is larger than the constants of the mono(ligand)iron(III) complexes of 2 and 3 that are terdentate. This result supports the coordination of the methylcarboxylate group that caused the chelate effect (cf. Fig. 4). This chelate effect is further reflected in the equilibria (I)’ and (II)’ ; a mono(ligand)iron(III) complex forms quantitatively as low as pH 1.2 for 1, whereas it forms just as pH 2.7 for 2 (cf. Fig. 2). Such a remarkable difference in complex formation behaviour has not been observed between the iron(III)3 and -4 systems3 that also have different numbers of chelate rings. This difference may probably be caused by the different structures of these two tetradentate ligands through a strain due to the six-membered chelate ring. It is interesting to discuss here the difference in the complex formation behaviour between 1 and 4. Ligand 1 forms only a mono(ligand)iron(III) complex even under the conditions CL >>CM, while ligand 4 can form a bis(ligand)iron(III) complex under the same conditions. This difference in the behaviour can be assumed to be responsible for the ligand structures of 1 (a salicylaldehyde derivative) and 4 (a pyridoxal derivative) : ligand 4 can form a CH-n intramolecular interaction in its bis(lig-
2349
Thermodynamic stabilities and formation mechanisms of Fe”’ complexes Table 2. Pathways and formation rate constants Pathways
Ligand 1 (H,cslpa)
2 (H+lpa)
3 (H,plpa)
4 (H,plas)
H,slal’
H,mugnY Cd&OH (Hph)
Fe3++ Hgslpa- (ks2) and Fe(OH)‘+ + H,cslpa (kz3) Fe(OH)‘+ + Hsslpa (k& Fe’+ + H,slpa (kx2) and Fe(OH)*+ + H,slpa+ (kz3) Fe(OH)‘+ + H,slpa (k2J Fe3++ Hsplpa+(k,,) and Fe(OH)‘+ + H,plpa*+ (ku) Fe(OH)‘+ + H,plpa (k,,) Fe3++ H,plas (k,,) and Fe(OH)‘+ + H,plas+ (kza) Fe(OH)‘+ + H,plas (k,,) Fe3++ H,slal (k3J and Fe(OH)‘+ +H3slal+ (k,,) Fe(OH)‘+ + H,slal (k& Fe3++ H,mugm (k33 Fe(OH)2+ + H,mugm- (kz2) Fe’+ +Hph (k,,) Fe(OH)2+ +Hph (k2,)
k,
(mol-’ dm3 s-r)
Remarks
Ref.
k326 = k,,(l), ku = k,(l)
This work
kzb = k,,(I) k326= k&2), kz = k,,(2)
This work This work
k,P = k&2)
This work 3
(7.26f 0.20) x lo2 0 k2,:(2.08
f0.30) x 10”
(1.48f0.20) x 104’ k2,:(7.87f0.30)
x
103d
b
(1.08f0.20) x 1o*=f k,,:(1.27&0.30) x 10” (5.22f0.20) x 10’g*h
b
k2,:(1.31kO.30) x lo39 (5.43kO.20) x 103”
b
k2,:(3.18f0.30) x 10’ k,,:(6.17&0.20) x lo2 k,,:(1.92*0.30) x 102 h : -25 k2, : 7.2 x lo2 1.1 x 102 1.5 x 102
b
b
3 3
3 18
18 14 14 19 19 20 21
%2+ h&mco,,K& bO.10mol dm-’ (NaClO,), 25°C. ‘k3,+ k&,,,,Kb:, = (6.90f0.20) x 10’ mol-’ dm3 s-l. dk3r+ kz&
and)iron(III) complex, as has been suggested previously3 by the use of the methyl group of pyridoxal moiety [cf. Fig. 4 (8)], whereas 1 cannot form this type of interaction as is easily recognized by a molecular model consideration. In fact, the X-ray structurer6 of the bis(ligand)cobalt(III) complexes of N”-pyridoxyl+alanine revealed that the pyridine rings of pyridoxal moieties were placed close to each other through the CH-x intramolecular interaction of the methyl group of the pyridoxal moiety with the aromatic ring. This intramolecular interaction is expected to probably increase the
L-
thermodynamic stability and the formation of bis(ligand)iron(III) complex with 4. Additionally, the less strain of the chelate ring structure of 5 compared with 7 (de supra; cf. Fig. 4) may also enhance the thermodynamic stability of the mono(ligand)iron(III) complex of 1. The intramolecular interactions that have been found in the X-ray structun? are also expected in the bis(ligand)iron(III) complex of 3. Therefore, the difference of the complex formation behaviour that 3 can form a bis(ligand)iron(III) complex in a single step in contrast to 2 in consecutive stages
2350
T. OZAWA et al.
B
[Wcslpa)l
II
F-C", ,P
Fig. 4. The proposed structures of mono(ligand)iron(III)
under the conditions of CL >. CM may reasonably be explained. Kinetics and mechanisms
For the tetradentate ligand 1, the complex formation rate constant of the composite pathways of Fe3+ + H,cslpaa[k,, (l)] and Fe(OH)2+ + H3cslpa ]k23 W1[k32(1)+k23 ~~~~~~~~~~~~~ K& = (7.26f0.20) x lo2 mol-’ dm3 s-i] is a reasonable magnitude as a rate constant of iron(II1) (cf. Table 2). The rate constant is in the same magnitude as that of the pyridoxal-derived ligand, 3. Furthermore, the constant of the single pathway of Fe(OH)2+ + H,cslpa- [k,,(l)] [k,,(l) = (2.08 f0.30) x lo3 mol-’ dm3 s-‘1 corresponds in magnitude to that of the same pathway of 33 (cf. Table 2). These facts indicate that the complex formation mechanism of 1 is the same as that of the iron(III)-3 system. For the terdentate ligand 2, on the other hand, the complex formation rate constant of the composite pathways of Fe3+ + H,slpa [k32(2)] and Fe(OH)2+ +H,slpa+ [k23(2)] (k32(2)+k23(2) K .CamCOOH&,~ = (1.48 + 0.20) x lo4 mol- ’ dm3 s-l] is considerably larger than the corresponding constants of 1 and 33 (cf. Table 2). Incidentally, much of the literature3*4~‘4*‘7-z’ has reported the rate constants of Fe3+- and Fe(OH)‘+-related pathways, indicating that the Fe(OH)2+-related
P
\
,COOH
C”
complexes of 1,2and 4.
pathways are generally one or two orders of magnitude larger in the constants than the Fe3+related pathways for the same ligand,17 so long as there are no restrictions such as steric hindrance due to the ligand structure, etc. Accordingly, the Fe(OH)‘+-related pathways are mainly expected to contribute to the composite pathways for the system with 2. The fact that the rate constants of the single pathways [k,,(l) and k,,(2)] of 1 and 2 related only to Fe(OH)2f are close in magnitude to the composite constants of k32(2) + k,,(2) K aCamCOOH&,h of the system with 2 allows us to substantiate the discussion given above. The comparison of the rate constants of k,,(l) for 1 and k,,(2) for 2 with those of the corresponding pathways of the iron(III)-N”-pyridoxyl-type ligand systems3s4suffices for the deduction of complex formation mechanism. The fact that the values of k,,(l) for 1 and k22(2) for 2 are the same in magnitude allows us to propose the resemblance in the complex formation mechanism for the iron(II1) - 1 and -2 systems. Therefore, both of the complex formation reactions of 1 and of 2 proceed equally through the slow coordination of the phenolato oxygen to [Fe(H20)J3+ or [Fe(OH)(H20),]2+ species, followed by the rapid donations of other coordinating atoms as shown in Scheme I. In conclusion, a new tetradentate ligand (1) exhibits an increased thermodynamic stability of the mono(ligand)iron(III) complex compared with its
Thermodynamic
stabilities and formation mechanisms of Fe”’ complexes
CHz-N+-CHCOOH
2351
CH2-N+-CHCOOH
+H+ +H+
5 (R=CH&Oo-
1
B (R=H) /
-0.
+
FeOHat
/
r.d.8
Scheme I parent terdentate ligand 2 ; this increased stability is mainly due to the chelate effect by a newly introduced methylcarboxylate group. The complex formation mechanism for 1, however, remains unchanged in the case of the mechanism for 2. Thus, the introduction of the methylcarboxylate group to a phenol-amino acid-type ligand is recommended to increase the thermodynamic stability towards iron(II1) without changing the formation reaction mechanism. Acknowledgements-We wish to express our gratitude to Professor 0. Yamauchi and Mr T. Yajima of Nagoya University for their help to determine the protonation constants by the potentiometric titration method.
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Englewood, NJ (1974). T. Ozawa, K. Jitsukawa, H. Masuda and H. Einaga, Ber. Buns. Phys. Chem. 1994,98,66. T. Ozawa, K. Jitsukawa and H. Einaga, $er. Buns. Phys. Chem. 1994,98,59. S. Konetchny-Rapp, G. Jung, K. N. Raymond, J. Meiwes and H. Z&ner, J. Am. Chem. Sot. 1992,114, 2224. 6. J. Kielland, J. Am. Chem. Sot. 1937,59, 1675.
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Springer, Berlin (1960). 11. M. Noritake, K. Okamoto, J. Hidaka and H. Einaga, Bull. Chem. Sot. Japan 1990,63, 353. 12. L. G. Sill&n and A. E. Martell, Stability Constants of Metal-Zon Complexes, Sp. Pub. No. 17, p. 53. Chemical Society, London (1964). 13. C. F. Baes Jr and R. E. Mesmer, The Hydrolysis of Cations, p. 237. John Wiley & Sons, New York (1976). 14. H. Yoshida, T. Ozawa, K. Jitsukawa and H. Einaga, Polyhedron 1993, 12, 1319. 15. W. R. Harris, R. J. Motekaitis and A. E. Martell, Znorg. Chem. 1975, 14,974. 16. K. Jitsukawa, K. Iwai, H. Masuda, H. Ogoshi and H. Einaga, Chem. Lett. 1994,303. 17. D. W. Margerum, G. R. Cayley, D. C. Weatherbum and G. K. Pagenkopf, Coordination Chemistry (Edited by A. E. Martell), Vol. 2, Ch. 1. American Chemical Society, Washington, DC (1978). 18. T. Ozawa, K. Iwai, K. Jitsukawa, H. Masuda and H. Einaga, Polyhedron, in press. 19. S. Gouder and J. Stuehr, Znorg. Chem. 1974,13,379. 20. K. Nakamura, T. Tsucbida, A. Yamagishi and M. Fujimoto, BUN. Chem. Sot. Japan 1973,46,456. 21. F. P. Cavacino and E. D. Dio, J. Chem. Sot. A 1970, 1151.