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New insight into the oxidation of Fe(II) by desferrioxamine B (DFB): spectrophotometric and capillary electrophoresis (CE) study
Inorganic Chemistry Communications 6 (2003) 131–134 www.elsevier.com/locate/inoche
New insight into the oxidation of Fe(II) by desferrioxamine B (DFB): spectrophotometric and capillary electrophoresis (CE) study .A. Enyedy, I. F E. Farkas *, E abi an Department of Inorganic and Analytical Chemistry, University of Debrecen, H-4010 Debrecen, Hungary Received 26 July 2002; accepted 18 October 2002
Abstract Results on the previously described irreversible redox reaction taking place between iron(II) and desferrioxamine B (DFB) under anaerobic conditions have been complemented by additional capillary electrophoresis (CE) and kinetic studies in the present work. Reduction of the oxidizing agent, DFB to monoamide derivative, was confirmed by CE technique and suggestion for the most probable kinetically active species and mechanism of the initial step is discussed in the paper. Ó 2002 Elsevier Science B.V. All rights reserved. Keywords: Fe(II); Desferrioxamine B; Oxidation; Kinetic studies; Capillary electrophoresis
1. Introduction Microbial siderophores solubilize and transport iron(III) into the cells in the required concentrations. In iron(III) complexes formed with hydroxamate-based siderophores the reduction of the metal centre within the cells plays a crucial role in the mechanism of iron release [1]. Consequently, investigation of the interaction between iron(II) and naturally occurring hydroxamatebased siderophores (e.g., DFB, see Scheme 1 for its formulae) or model hydroxamic acids can contribute to a better understanding of the above-mentioned biological process. In our previous work [2] complexation of iron(II) with monohydroxamic acids, dihydroxamic acids and DFB was studied. It was found that all the studied ligands with the exception of DFB bind to iron(II) as usual, via their hydroxamate oxygens. However, unexpected oxidation of iron(II) to iron(III) by DFB under strictly oxygen-free conditions was observed and this reaction was assumed to have importance in microbial iron uptake under anaerobic conditions. According to our preceding results, the interaction between iron(II) ion and DFB, that
occurs above pH 4, is accompanied by a continuous pH decrease and by the parallel development of the characteristic charge-transfer band of the Fe(III)–tris(hydroxamate) complex (½FeðIIIÞðHDFBÞþ ) in the visible spectrum (kmax 430, emax 2600 M1 cm1 ) [2]. The pH-stat measurements at pH 5.5 confirmed the fact that two moles of protons per one mole of Fe(II) are released when the ligand to metal ratio is at least 3 to 2. Oxidation of iron(II) was excluded both by the atmospheric oxygen and solvent water. The final conclusion was that DFB is the oxidizing agent and one of its hydroxamic acid groups is reduced to amide in the reaction. The oxidized metal ion is complexed by the excess of DFB while the reduced ligand does not coordinate. Similar reactions are known in the literature in which Mo(V), V(III), V(IV), U(IV) and Rh(I) are oxidized by simple monohydroxamic acids or dihydroxamic acids [3–5]. Osmium(II) was also found to be a reducing agent of N-phenylbenzohydroxamic acids, but in these systems the reduced amide remained C,N-coordinated to Os(III) [6]. The following stoichiometry was proposed for the Fe(II)–DFB reaction: 2Fe2þ þ 3H4 DFBþ ¼ 2½FeðIIIÞðHDFBÞþ
*
Corresponding author. Tel.: +36-52-512-900-2306; fax: +36-52489-667. E-mail address: efarkas@delfin.klte.hu (E. Farkas). 1387-7003/02/$ - see front matter Ó 2002 Elsevier Science B.V. All rights reserved. PII: S 1 3 8 7 - 7 0 0 3 ( 0 2 ) 0 0 7 0 3 - 7
þ H3 DFB-monoamideþ þ H2 O þ 4Hþ
ð1Þ
132
E. Farkas et al. / Inorganic Chemistry Communications 6 (2003) 131–134
Scheme 1. Formulae of desferrioxamine B (H4 DFBþ ).
Reaction (1) reveals that two moles of Fe(II) are oxidized by one of the three functional groups of one DFB, while the other two DFB molecules are required to bind Fe(III) in the high stability tris-hydroxamato complex. However, we were not successful either detecting the H3 DFB-monoamideþ formed in the reduction as byproduct or in characterizing the kinetically active species. In the present work, we complement our previous work, studying the system by capillary electrophoresis and time-resolved spectrophotometric methods.
mol=dm3 . Under such conditions, strictly first-order kinetic behaviour was observed and kinetic traces were fitted to an experimental function. Altogether ca. 30 kinetic runs were made. The pH ð log½HÞ was adjusted by MES (2-morphine-ethanesulphonic acid, pK 6.06) or HEPES(2-[4-(2-hydroxyethyl)-1-piperazinyl]-ethanesulphonic acid, pK 7.50) buffers. Absorbances were recorded at 429 nm. During the calculations the absorbance–time curves were fitted by the use of the software of the HP Instrument. Capillary electrophoresis measurements were made on a HP 3D Instrument and fused-silica (Polymicro Technology, 64:5 cm 50 lm i.d.) capillary was used. The applied voltage was +25 kV, while injections were made at 100 mbar. The iron was removed from the samples by using 100-hold excess of quinoline-8-olate according to method described in literature [9]. The concentration of individual DFB and MES samples was 1:67 g=dm3 , while in the case of the mixture (DFB, MES and the monoamide derivative) 6:67 g=dm3 . 25 mmol=dm3 Na2 HPO4 (pH 9.10) solution was used as a buffer. Electropherograms were evaluated by ChemStation 7.01 HP program.
2. Experimental DFB was obtained from CIBA Geigy. The purity of the ligand and the concentration of the solution were determined by GranÕs method [7]. FeCl2 stock solution was prepared by the dissolution of iron (Reanal) in excess of HCl, then the solution was filtered and stored under strictly oxygen-free Ar atmosphere. The concentrations of FeCl2 and HCl were determined by titrimetry with KMnO4 and by pH-potentiometry, respectively. Argon overpressure was applied when the stock solution was added to the samples. The pH-potentiometric measurements were carried out at an ionic strength of 0.2 M (KCl). Carbonate-free KOH solutions of known concentrations (ca. 0.2 M) were used as titrant. Radiometer pHM 84 instrument equipped with a Metrohm 62104130 combined electrode was used to collect the data. The electrode system was calibrated according to Irving et al. [8]. Kinetic measurements were carried out on a HP 8453 Instrument in the region of 300–980 nm. Path length was 1 cm. A special, tightly closed tandem cuvette (Tandem Cell 236 HELLMA) was used in which both isolated pockets were completely deoxygenated by bubbling a stream of argon for 10–10 min before FeCl2 was added to the samples. The spectra were recorded immediately after mixing the reactants. In the samples the metal to ligand ratio was changed in the range of 1/10–1/180 in the case of the concentration dependence studies while the concentration of Fe(II) was in the range of 1:5 104 3 104 mol=dm3 . The pH dependence of the reaction rate was studied at 1/10 metal to ligand ratio and the concentration of Fe(II) was 3 104
3. Results and discussion 3.1. Indication of the by-product, DFB-monoamide by CE method Since the unexpected redox reaction (1) between Fe(II) and DFB was discovered, numerous trials for the detection of the by-product have been performed. However, the presence of the iron(III) complex, ½FeðIIIÞðHDFBÞþ , in the equilibrium mixture interferes with these studies, thus Fe(III) was removed from the complex (for details see Section 2). The Fe(III)-free samples obtained this way contained DFB, its monoamide derivative, MES buffer and traces of the extracting agent, quinoline-8-olate. This sample was used to perform CE measurements. Independent samples of DFB, MES and quinoline-8-olate were also prepared and their UV-spectra and electropherograms were recorded (representative examples are shown in Figs. 1 and 2). Fortunately, neither the MES (see Fig. 1, inset b) nor the quinoline-8-olate has measurable absorbance at 220 nm where DFB exhibits a characteristic band (see Fig. 1, inset a). As Fig. 2 shows, in addition to the characteristic peak of DFB a new one appears in the first part of the electropherogram of the mixture. The new peak belongs to a species specified by somewhat higher charge/mass value compared to DFB. Considering the fact that the molecular weight of the presumed monoamide derivative is less by 17 than that of DFB, but the overall charges of the two molecules should be about the same, the charge/
E. Farkas et al. / Inorganic Chemistry Communications 6 (2003) 131–134
Fig. 1. Electropherograms of DFB and MES at 220 nm with the inserted UV spectra of DFB (a) and MES (b) (cDFB;MES ¼ 1:67 g=dm3 , pH 9.10).
Fig. 2. Electropherogram of the Fe(III)-free mixture at 220 nm with the inserted UV spectrum of the monoamide (cmixture ¼ 6:67 g=dm3 , pH 9.10).
mass value of the former molecule should be somewhat higher than that of the latter one. It is also worth mentioning that the UV-spectrum of the Fe(III)-free mixture is very similar to the spectrum of DFB (cf. inset a in Fig. 1 and inset in Fig. 2). Consequently, the new peak in Fig. 2 most likely belongs to the reduced form of DFB.
133
In order to obtain further information on the overall process, a systematic kinetic study was performed. Because of the hydrolytic processes of Fe(II), the reaction mixtures containing metal excess were very difficult to study. Therefore, kinetic runs were always made at high ligands excess. At least 1 to 10 or higher metal to ligand ratios were applied in order to provide pseudo-firstorder condition and the concentration of the Fe(II) was constant. The well-known characteristic charge-transfer band of the tris-chelated ½FeðIIIÞðHDFBÞþ was used to follow the reaction (see inset in Fig. 3). As an example, typical experimental traces together with the fitted curve are shown in Fig. 3. All kinetic runs were made at 429 nm. The concentration and pH dependencies of k obs indicate somewhat complex kinetic patterns. Thus, the redox reaction was observed only within a limited pH region and k obs goes through a maximum as a function of pH. The highest reaction rate was observed at physiological pH. At relatively low pH, the reaction rate increased almost linearly by increasing the concentration of DFB and the corresponding plot of the data showed a slight saturation character at higher ligand concentrations. In contrast, at higher pH kobs was practically independent of the ligand concentration in the experimentally accessible concentration range. While an explicit expression for the pH dependency of kobs cannot be derived the corresponding plot of kobs clearly indicates that the concentration of the reactive species goes though a maximum as a function of pH. This can be interpreted by assuming that the variation of the reaction rate reflects the concentration and pH dependencies of the concentration distribution of the kinetically active Fe(II) complex. Presumably, the ligand is partly protonated in this species the concentration of which first increases by increasing the pH. Eventually the complex undergoes deprotonation in more alkaline solution and vanishes.
3.2. Kinetic studies Although the stoichiomery of the redox reaction does not provide any information about the mechanism, it is an important observation that the pH where the interaction between Fe(II) and DFB starts coincides with the pH where the redox reaction starts. Moreover, since the redox reaction is not fast enough, the formation of a Fe(II) complex could be detected spectrophotometrically (kmax ¼ 950 nm, emax ¼ 1 M1 cm1 ) immediately after mixing the sample. The spectrum agrees well with that detected for the Fe(II)–acetohydroxamic acid complex ðkmax ¼ 900 nm, emax ¼ 1 M1 cm1 Þ. Conclusively, the above findings strongly suggest that the formation of a Fe(II) complex is the initial step of the redox reaction studied.
Fig. 3. A typical kinetic curve for Fe(II)–DFB system registered at 429 nm (d) with the fitted curve (solid line) (metal to ligand ratio ¼ 1/10, cFeðIIÞ ¼ 3:0 104 M, pH 6.75). Inset shows the absorbance spectra as a function of time.
134
E. Farkas et al. / Inorganic Chemistry Communications 6 (2003) 131–134
Because of the noted experimental difficulties, direct determination of the composition and the stability constant of the redox active complex were not feasible. However, it is reasonable to assume that the equilibrium features of the Fe(II)–DFB system are very similar to those of related metal ion–DFB systems, i.e., the ligand may coordinate to iron(II) via one or more hydroxamate groups depending on the pH and the actual geometry of the complexes formed. The stability constants determined for hydroxamic acid complexes of 3d5 –3d10 metal(II) ions [2,10] demonstrate unambiguously the validity of the ‘‘Irving-Williams order’’. Therefore, the logb values of the Fe(II)–DFB complexes were estimated on the basis of the stability constants of the corresponding complexes with Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) [10,11] (Fig. 4). The stability constant calculated for the Fe(II)–trishydroxamato complex and shown in the Fig. 4 agrees well with the value, which can
be calculated by the use of E1=2 and the stability constant of [Fe(III)(HDFB)] complex [1]. The concentration distribution curves for the Fe(II)– DFB system were calculated by using the estimated stability constants for the various complexes. As shown in Fig. 5, kobs shows very similar pH dependency to that of the concentration of the monochelated complex. Considering the uncertainties associated with the estimation of the stability constants, the correlation seems to be reasonable between these curves leading to the conclusion that the redox active species is the ½FeðDFBH3 Þ2þ complex. In accordance with these considerations, we propose that fast formation of this species precedes the rate determining intramolecular electron transfer step. Further studies to provide a quantitative kinetic description of this reaction are in progress in our laboratories. Acknowledgements The authors are grateful for the financial support from OTKA T 034674, OTKA TS 040685 and M 028244. References
Fig. 4. Stability constants (log b) for complexes formed between DFB and bivalent 3d5–10 metal ions with the estimated values for Fe(II); (b ¼ ð½MðDFBHq ÞÞ=ð½M½DFB½Hq Þ); t ¼ 25 °C, I ¼ 0:2 M(KCl).
Fig. 5. The observed reaction rates measured in the Fe(II)–DFB system at different pH values (d) and the supposed concentration distribution curves for the same system (metal to ligand ratio ¼ 1/10, cFeðIIÞ ¼ 3:0 104 M).
[1] A.M. Albrecht-Gary, A.L. Crumbliss, in: H. Sigel, A. Sigel (Eds.), Metal Ions in Biological Systems, vol. 35, Marcel-Dekker, New York, 1998. .A. Enyedy, L. Zekany, Gy. Deak, J. Inorg. Biochem. [2] E. Farkas, E 83 (2001) 107. [3] D.A. Brown, H. B€ ogge, R. Cooga, D. Doocey, T.J. Kemp, A. M€ uller, B. Neumann, Inorg. Chem. 35 (1996) 1674. [4] W.L. Smith, K.N. Raymond, J. Inorg. Nucl. Chem. 41 (1979) 1431. [5] A. Das, F. Basuli, S.M. Peng, S. Bhattacharya, Inorg. Chem. 41 (2002) 440. [6] A. Das, F. Basuli, I.R. Falvello, S. Battacharya, Inorg. Chem. 40 (2001) 4085. [7] G. Gran, Acta Chem. Scand. 4 (1950) 559. [8] H. Irving, M.G. Miles, L.D. Pettit, Anal. Chim. Acta 38 (1967) 475. [9] M.A.F. Jalal, D. van derHelm, in: G. Winkelmann (Ed.), CRC Handbook of Microbial Iron Chelates, vol. 7, CRC Press, Boca Raton, 1990. [10] E. Farkas, H. Cs oka, G. Micera, A. Dessi, J. Inorg. Biochem. 65 (1997) 281. .A. Enyedy, PhD Thesis, Debrecen (2002) (logb values for [11] E complexes of Mn(II) and Co(II): ½MnðH3 DFBÞ2þ : 31.8, ½MnðH2 DFBÞþ : 25.73, [Mn(HDFB)]: 17.37, ½MnðDFBÞ : 6.37; ½CoðH3 DFBÞ2þ : 32.8, ½CoðH2 DFBÞþ : 27.44, [Co(HDFB)]: 19.57, ½CoðDFBÞ : 8.55).
New insight into the oxidation of Fe(II) by desferrioxamine B (DFB): spectrophotometric and capillary electrophoresis (CE) study .A. Enyedy, I. F E. Farkas *, E abi an Department of Inorganic and Analytical Chemistry, University of Debrecen, H-4010 Debrecen, Hungary Received 26 July 2002; accepted 18 October 2002
Abstract Results on the previously described irreversible redox reaction taking place between iron(II) and desferrioxamine B (DFB) under anaerobic conditions have been complemented by additional capillary electrophoresis (CE) and kinetic studies in the present work. Reduction of the oxidizing agent, DFB to monoamide derivative, was confirmed by CE technique and suggestion for the most probable kinetically active species and mechanism of the initial step is discussed in the paper. Ó 2002 Elsevier Science B.V. All rights reserved. Keywords: Fe(II); Desferrioxamine B; Oxidation; Kinetic studies; Capillary electrophoresis
1. Introduction Microbial siderophores solubilize and transport iron(III) into the cells in the required concentrations. In iron(III) complexes formed with hydroxamate-based siderophores the reduction of the metal centre within the cells plays a crucial role in the mechanism of iron release [1]. Consequently, investigation of the interaction between iron(II) and naturally occurring hydroxamatebased siderophores (e.g., DFB, see Scheme 1 for its formulae) or model hydroxamic acids can contribute to a better understanding of the above-mentioned biological process. In our previous work [2] complexation of iron(II) with monohydroxamic acids, dihydroxamic acids and DFB was studied. It was found that all the studied ligands with the exception of DFB bind to iron(II) as usual, via their hydroxamate oxygens. However, unexpected oxidation of iron(II) to iron(III) by DFB under strictly oxygen-free conditions was observed and this reaction was assumed to have importance in microbial iron uptake under anaerobic conditions. According to our preceding results, the interaction between iron(II) ion and DFB, that
occurs above pH 4, is accompanied by a continuous pH decrease and by the parallel development of the characteristic charge-transfer band of the Fe(III)–tris(hydroxamate) complex (½FeðIIIÞðHDFBÞþ ) in the visible spectrum (kmax 430, emax 2600 M1 cm1 ) [2]. The pH-stat measurements at pH 5.5 confirmed the fact that two moles of protons per one mole of Fe(II) are released when the ligand to metal ratio is at least 3 to 2. Oxidation of iron(II) was excluded both by the atmospheric oxygen and solvent water. The final conclusion was that DFB is the oxidizing agent and one of its hydroxamic acid groups is reduced to amide in the reaction. The oxidized metal ion is complexed by the excess of DFB while the reduced ligand does not coordinate. Similar reactions are known in the literature in which Mo(V), V(III), V(IV), U(IV) and Rh(I) are oxidized by simple monohydroxamic acids or dihydroxamic acids [3–5]. Osmium(II) was also found to be a reducing agent of N-phenylbenzohydroxamic acids, but in these systems the reduced amide remained C,N-coordinated to Os(III) [6]. The following stoichiometry was proposed for the Fe(II)–DFB reaction: 2Fe2þ þ 3H4 DFBþ ¼ 2½FeðIIIÞðHDFBÞþ
*
Corresponding author. Tel.: +36-52-512-900-2306; fax: +36-52489-667. E-mail address: efarkas@delfin.klte.hu (E. Farkas). 1387-7003/02/$ - see front matter Ó 2002 Elsevier Science B.V. All rights reserved. PII: S 1 3 8 7 - 7 0 0 3 ( 0 2 ) 0 0 7 0 3 - 7
þ H3 DFB-monoamideþ þ H2 O þ 4Hþ
ð1Þ
132
E. Farkas et al. / Inorganic Chemistry Communications 6 (2003) 131–134
Scheme 1. Formulae of desferrioxamine B (H4 DFBþ ).
Reaction (1) reveals that two moles of Fe(II) are oxidized by one of the three functional groups of one DFB, while the other two DFB molecules are required to bind Fe(III) in the high stability tris-hydroxamato complex. However, we were not successful either detecting the H3 DFB-monoamideþ formed in the reduction as byproduct or in characterizing the kinetically active species. In the present work, we complement our previous work, studying the system by capillary electrophoresis and time-resolved spectrophotometric methods.
mol=dm3 . Under such conditions, strictly first-order kinetic behaviour was observed and kinetic traces were fitted to an experimental function. Altogether ca. 30 kinetic runs were made. The pH ð log½HÞ was adjusted by MES (2-morphine-ethanesulphonic acid, pK 6.06) or HEPES(2-[4-(2-hydroxyethyl)-1-piperazinyl]-ethanesulphonic acid, pK 7.50) buffers. Absorbances were recorded at 429 nm. During the calculations the absorbance–time curves were fitted by the use of the software of the HP Instrument. Capillary electrophoresis measurements were made on a HP 3D Instrument and fused-silica (Polymicro Technology, 64:5 cm 50 lm i.d.) capillary was used. The applied voltage was +25 kV, while injections were made at 100 mbar. The iron was removed from the samples by using 100-hold excess of quinoline-8-olate according to method described in literature [9]. The concentration of individual DFB and MES samples was 1:67 g=dm3 , while in the case of the mixture (DFB, MES and the monoamide derivative) 6:67 g=dm3 . 25 mmol=dm3 Na2 HPO4 (pH 9.10) solution was used as a buffer. Electropherograms were evaluated by ChemStation 7.01 HP program.
2. Experimental DFB was obtained from CIBA Geigy. The purity of the ligand and the concentration of the solution were determined by GranÕs method [7]. FeCl2 stock solution was prepared by the dissolution of iron (Reanal) in excess of HCl, then the solution was filtered and stored under strictly oxygen-free Ar atmosphere. The concentrations of FeCl2 and HCl were determined by titrimetry with KMnO4 and by pH-potentiometry, respectively. Argon overpressure was applied when the stock solution was added to the samples. The pH-potentiometric measurements were carried out at an ionic strength of 0.2 M (KCl). Carbonate-free KOH solutions of known concentrations (ca. 0.2 M) were used as titrant. Radiometer pHM 84 instrument equipped with a Metrohm 62104130 combined electrode was used to collect the data. The electrode system was calibrated according to Irving et al. [8]. Kinetic measurements were carried out on a HP 8453 Instrument in the region of 300–980 nm. Path length was 1 cm. A special, tightly closed tandem cuvette (Tandem Cell 236 HELLMA) was used in which both isolated pockets were completely deoxygenated by bubbling a stream of argon for 10–10 min before FeCl2 was added to the samples. The spectra were recorded immediately after mixing the reactants. In the samples the metal to ligand ratio was changed in the range of 1/10–1/180 in the case of the concentration dependence studies while the concentration of Fe(II) was in the range of 1:5 104 3 104 mol=dm3 . The pH dependence of the reaction rate was studied at 1/10 metal to ligand ratio and the concentration of Fe(II) was 3 104
3. Results and discussion 3.1. Indication of the by-product, DFB-monoamide by CE method Since the unexpected redox reaction (1) between Fe(II) and DFB was discovered, numerous trials for the detection of the by-product have been performed. However, the presence of the iron(III) complex, ½FeðIIIÞðHDFBÞþ , in the equilibrium mixture interferes with these studies, thus Fe(III) was removed from the complex (for details see Section 2). The Fe(III)-free samples obtained this way contained DFB, its monoamide derivative, MES buffer and traces of the extracting agent, quinoline-8-olate. This sample was used to perform CE measurements. Independent samples of DFB, MES and quinoline-8-olate were also prepared and their UV-spectra and electropherograms were recorded (representative examples are shown in Figs. 1 and 2). Fortunately, neither the MES (see Fig. 1, inset b) nor the quinoline-8-olate has measurable absorbance at 220 nm where DFB exhibits a characteristic band (see Fig. 1, inset a). As Fig. 2 shows, in addition to the characteristic peak of DFB a new one appears in the first part of the electropherogram of the mixture. The new peak belongs to a species specified by somewhat higher charge/mass value compared to DFB. Considering the fact that the molecular weight of the presumed monoamide derivative is less by 17 than that of DFB, but the overall charges of the two molecules should be about the same, the charge/
E. Farkas et al. / Inorganic Chemistry Communications 6 (2003) 131–134
Fig. 1. Electropherograms of DFB and MES at 220 nm with the inserted UV spectra of DFB (a) and MES (b) (cDFB;MES ¼ 1:67 g=dm3 , pH 9.10).
Fig. 2. Electropherogram of the Fe(III)-free mixture at 220 nm with the inserted UV spectrum of the monoamide (cmixture ¼ 6:67 g=dm3 , pH 9.10).
mass value of the former molecule should be somewhat higher than that of the latter one. It is also worth mentioning that the UV-spectrum of the Fe(III)-free mixture is very similar to the spectrum of DFB (cf. inset a in Fig. 1 and inset in Fig. 2). Consequently, the new peak in Fig. 2 most likely belongs to the reduced form of DFB.
133
In order to obtain further information on the overall process, a systematic kinetic study was performed. Because of the hydrolytic processes of Fe(II), the reaction mixtures containing metal excess were very difficult to study. Therefore, kinetic runs were always made at high ligands excess. At least 1 to 10 or higher metal to ligand ratios were applied in order to provide pseudo-firstorder condition and the concentration of the Fe(II) was constant. The well-known characteristic charge-transfer band of the tris-chelated ½FeðIIIÞðHDFBÞþ was used to follow the reaction (see inset in Fig. 3). As an example, typical experimental traces together with the fitted curve are shown in Fig. 3. All kinetic runs were made at 429 nm. The concentration and pH dependencies of k obs indicate somewhat complex kinetic patterns. Thus, the redox reaction was observed only within a limited pH region and k obs goes through a maximum as a function of pH. The highest reaction rate was observed at physiological pH. At relatively low pH, the reaction rate increased almost linearly by increasing the concentration of DFB and the corresponding plot of the data showed a slight saturation character at higher ligand concentrations. In contrast, at higher pH kobs was practically independent of the ligand concentration in the experimentally accessible concentration range. While an explicit expression for the pH dependency of kobs cannot be derived the corresponding plot of kobs clearly indicates that the concentration of the reactive species goes though a maximum as a function of pH. This can be interpreted by assuming that the variation of the reaction rate reflects the concentration and pH dependencies of the concentration distribution of the kinetically active Fe(II) complex. Presumably, the ligand is partly protonated in this species the concentration of which first increases by increasing the pH. Eventually the complex undergoes deprotonation in more alkaline solution and vanishes.
3.2. Kinetic studies Although the stoichiomery of the redox reaction does not provide any information about the mechanism, it is an important observation that the pH where the interaction between Fe(II) and DFB starts coincides with the pH where the redox reaction starts. Moreover, since the redox reaction is not fast enough, the formation of a Fe(II) complex could be detected spectrophotometrically (kmax ¼ 950 nm, emax ¼ 1 M1 cm1 ) immediately after mixing the sample. The spectrum agrees well with that detected for the Fe(II)–acetohydroxamic acid complex ðkmax ¼ 900 nm, emax ¼ 1 M1 cm1 Þ. Conclusively, the above findings strongly suggest that the formation of a Fe(II) complex is the initial step of the redox reaction studied.
Fig. 3. A typical kinetic curve for Fe(II)–DFB system registered at 429 nm (d) with the fitted curve (solid line) (metal to ligand ratio ¼ 1/10, cFeðIIÞ ¼ 3:0 104 M, pH 6.75). Inset shows the absorbance spectra as a function of time.
134
E. Farkas et al. / Inorganic Chemistry Communications 6 (2003) 131–134
Because of the noted experimental difficulties, direct determination of the composition and the stability constant of the redox active complex were not feasible. However, it is reasonable to assume that the equilibrium features of the Fe(II)–DFB system are very similar to those of related metal ion–DFB systems, i.e., the ligand may coordinate to iron(II) via one or more hydroxamate groups depending on the pH and the actual geometry of the complexes formed. The stability constants determined for hydroxamic acid complexes of 3d5 –3d10 metal(II) ions [2,10] demonstrate unambiguously the validity of the ‘‘Irving-Williams order’’. Therefore, the logb values of the Fe(II)–DFB complexes were estimated on the basis of the stability constants of the corresponding complexes with Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) [10,11] (Fig. 4). The stability constant calculated for the Fe(II)–trishydroxamato complex and shown in the Fig. 4 agrees well with the value, which can
be calculated by the use of E1=2 and the stability constant of [Fe(III)(HDFB)] complex [1]. The concentration distribution curves for the Fe(II)– DFB system were calculated by using the estimated stability constants for the various complexes. As shown in Fig. 5, kobs shows very similar pH dependency to that of the concentration of the monochelated complex. Considering the uncertainties associated with the estimation of the stability constants, the correlation seems to be reasonable between these curves leading to the conclusion that the redox active species is the ½FeðDFBH3 Þ2þ complex. In accordance with these considerations, we propose that fast formation of this species precedes the rate determining intramolecular electron transfer step. Further studies to provide a quantitative kinetic description of this reaction are in progress in our laboratories. Acknowledgements The authors are grateful for the financial support from OTKA T 034674, OTKA TS 040685 and M 028244. References
Fig. 4. Stability constants (log b) for complexes formed between DFB and bivalent 3d5–10 metal ions with the estimated values for Fe(II); (b ¼ ð½MðDFBHq ÞÞ=ð½M½DFB½Hq Þ); t ¼ 25 °C, I ¼ 0:2 M(KCl).
Fig. 5. The observed reaction rates measured in the Fe(II)–DFB system at different pH values (d) and the supposed concentration distribution curves for the same system (metal to ligand ratio ¼ 1/10, cFeðIIÞ ¼ 3:0 104 M).
[1] A.M. Albrecht-Gary, A.L. Crumbliss, in: H. Sigel, A. Sigel (Eds.), Metal Ions in Biological Systems, vol. 35, Marcel-Dekker, New York, 1998. .A. Enyedy, L. Zekany, Gy. Deak, J. Inorg. Biochem. [2] E. Farkas, E 83 (2001) 107. [3] D.A. Brown, H. B€ ogge, R. Cooga, D. Doocey, T.J. Kemp, A. M€ uller, B. Neumann, Inorg. Chem. 35 (1996) 1674. [4] W.L. Smith, K.N. Raymond, J. Inorg. Nucl. Chem. 41 (1979) 1431. [5] A. Das, F. Basuli, S.M. Peng, S. Bhattacharya, Inorg. Chem. 41 (2002) 440. [6] A. Das, F. Basuli, I.R. Falvello, S. Battacharya, Inorg. Chem. 40 (2001) 4085. [7] G. Gran, Acta Chem. Scand. 4 (1950) 559. [8] H. Irving, M.G. Miles, L.D. Pettit, Anal. Chim. Acta 38 (1967) 475. [9] M.A.F. Jalal, D. van derHelm, in: G. Winkelmann (Ed.), CRC Handbook of Microbial Iron Chelates, vol. 7, CRC Press, Boca Raton, 1990. [10] E. Farkas, H. Cs oka, G. Micera, A. Dessi, J. Inorg. Biochem. 65 (1997) 281. .A. Enyedy, PhD Thesis, Debrecen (2002) (logb values for [11] E complexes of Mn(II) and Co(II): ½MnðH3 DFBÞ2þ : 31.8, ½MnðH2 DFBÞþ : 25.73, [Mn(HDFB)]: 17.37, ½MnðDFBÞ : 6.37; ½CoðH3 DFBÞ2þ : 32.8, ½CoðH2 DFBÞþ : 27.44, [Co(HDFB)]: 19.57, ½CoðDFBÞ : 8.55).