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A biologically relevant iron(III) phenoxyl radical complex: A thermodynamic investigation on the structure-radical stability relationship
Journal of Molecular Structure 1022 (2012) 109–116
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A biologically relevant iron(III) phenoxyl radical complex: A thermodynamic investigation on the structure-radical stability relationship Iraj Saberikia a, Elham Safaei a,⇑, Mohammad Rafiee a, Patricia Coticˇ b, Giuseppe Bruno c, Hadi Amiri Rudbari c a b c
Institute for Advanced Studies in Basic Sciences (IASBS), 45195 Zanjan, Iran Institute of Mathematics, Physics and Mechanics, University of Ljubljana, Ljubljana, Slovenia Dipartimento di Chimica Inorganica, Vill. S. Agata, Salita Sperone 31, Università di Messina, 98166 Messina, Italy
h i g h l i g h t s " A binuclear Fe(III) complex of aminophenol ligand synthesized. " Thermodynamic parameters of electrochemically produced radicals were investigated. " Radical complex/Fe(III) center interaction make it stabilized such as RNR enzyme.
a r t i c l e
i n f o
Article history: Received 6 March 2012 Accepted 27 April 2012 Available online 9 May 2012 Keywords: Ribonucleotide reductase Aminophenol Model complex Iron complexes
a b s t r a c t A kind of new amine-chloro substituted phenol and its iron(III) complex has been prepared and characterized by spectroscopic, X-ray techniques and magnetic susceptibility studies. X-ray analysis revealed a binuclear complex, Fe2(LTHF)3 in which two Fe(III) centers were surrounded by three ligands. The magnetic susceptibility indicates antiferromagnetic coupling between two iron centers through phenolate bridges. Voltammetric study of ligand (H2LTHF) and complex Fe2(LTHF)3 revealed that the coordinated phenolate ligands undergo reversible two-electron oxidations with formation of coordinated phenoxyl radicals. It has seen that the oxidation of Fe2(LTHF)3 comparing to its analog, Fe2(LtBu)2 occurs at significantly higher potentials. It was attributed to the relationship between the radical stability with electronic and steric properties of ligands. Stability constants and thermodynamic parameters for the formation of Fe2(LTHF)3 complex and the role of ligand coordination to iron center in phenoxyl radical stabilization has been investigated using anodic peaks potentials. Ó 2012 Elsevier B.V. All rights reserved.
1. Introduction Radicals are highly reactive species due to their unpaired electrons. They have a key role in living organism’s physiology, biochemistry and some chemical processes, such as polymerization. Interest in the field of protein and amino acid radicals has been dramatically increasing because these active species serve as enzyme co-factors to lots of enzymatic reactions in living organisms. [1]. One of the most important amino acid radicals is tyrosyl radical which plays vital roles in the chemistry of living systems [1,2]. Enzymes that incorporate this kind of radical include active form of galactose oxidase (GO) [1,3,4], glyoxal oxidase (GLO) [2] and ribonucleotide reductase, RNR (Scheme 1) [5–9]. This last one is the only known enzyme that carries a stable free radical in its resting state and catalyzes DNA synthesis by ⇑ Corresponding author. Tel.: +98 241 4153200; fax: +98 241 4153232. E-mail address: [email protected] (E. Safaei). 0022-2860/$ - see front matter Ó 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.molstruc.2012.04.084
reducing ribonucleotides to their corresponding deoxyribonucleotides (Scheme 2). It has shown that the catalytic action of this enzyme involve formation of a tyrosyl radical close (tyrosyl 122 radical, Y122), to the redox active iron center which is responsible for the radical stability in RNR [8,9]. Stabilization of the mentioned radicals is very important due to their role in biological processes [1,10]. Scientists have found two mechanisms for tyrosyl radical stabilization: 1. Stabilization by hydrogen bonding to its neighbor amino acid residues (e.g. in GO). 2. Stabilization by interaction with a transition metal center (e.g. in RNR). The key role of the second mechanism in biological and catalytic radical processes is due to two major reasons [1,11]: (a) Redox reaction assisted radical generation. (b) Interaction of transition metals with their adjacent radicals.
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Scheme 1. The enzyme structure ribonucleotide reductase (RNR) [9].
Scheme 2. Enzymatic reaction catalyzed by ribonucleotide reductase [8].
On the other hand, the ‘‘metal–radical’’ approach has been noticed by scientists as one of the more attractive strategies for the design of biocatalysts and molecule-based materials. Therefore, nature and properties of d-transition metal bound to one or more radicals is much interest. Phenoxyl type radicals are one of these groups that their chemistry has been much interest because of key participation in biological redox processes [12–18]. Because of their close relationship with tyrosine-containing metalloenzymes, iron and copper complexes of chelating amine-phenolate ligands have been studied [19–27]. We report here the coordination chemistry, magnetic and redox properties of [6,60 -((((tetrahydrofuran-2-yl)methyl)azanediyl) bis(methylene))bis(2,4-dichlorophenol)] (H2LTHF) (Scheme 3) and its iron(III) complex with this ligand. In this way, the role of phenoxyl radical/iron(III) center interaction in radical stabilization has been investigated.
Scheme 3. Structure of [6,60 -((((tetrahydrofuran-2-yl)methyl) azanediyl)bis(methylene))bis(2,4-dichlorophenol)] (H2LTHF).
2. Experimental 2.1. Materials and physical measurements Reagents or analytical grade materials were obtained from commercial suppliers and used without further purification, except those for electrochemical measurements. Elemental analyses (C H N) were performed by the Research Institute of Petroleum Industry (RIPI). Fourier transform infrared spectroscopy on KBr pellets was performed on a FT-IR Bruker Vector 22 instrument. NMR measurements were performed on a Bruker 400 instrument. UV–Vis absorbance digitized spectra were collected using a CARY 100 spectrophotometer. Magnetic susceptibility were measured from powder samples of solid material in the temperature range 2–300 K by using a SQUID susceptometer (Quantum Design MPMS-XL-5) in a magnetic field of 1000 Oe. Voltammetric measurements were made with a computer controlled electrochemical system (ECO Chemie, Ultrecht, The Netherlands) equipped with a PGSTA 30 model and driven by GPES (ECO Chemie). The reference electrode was an Ag wire as the quasi reference electrode. A glassy carbon electrode with a surface area of 0.035 cm2 was used as a working electrode and a platinum wire served as the counter electrode. Prior to the measurement, the GC electrode was polished with 0.1 mm alumina powder and washed with distilled water. The voltage scan rate was set at 50 mV sec1. The solutions were deoxygenated by bubbling nitrogen gas through them. Ferrocene was added as an internal standard after completion of a set of experiments, and potentials are referenced vs. the ferrocenium/ferrocene couple (Fc+/Fc). Diffraction data for Fe2(LTHF)3 was measured on a Bruker–Nonius X8 ApexII diffractometer equipped with a CCD area detector by using graphite-monochromated Mo Ka radiation (k = 0.71073)
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I. Saberikia et al. / Journal of Molecular Structure 1022 (2012) 109–116 Table 1 Crystal data and structure refinement for Fe2(LTHF)3. Empirical formula Formula weight Temperature (K) Wavelength (Å) Crystal system Space group Unit cell dimensions Volume (Å3) Z Calculated density (Mg/m3) Absorption coefficient (mm1) F(0 0 0) Crystal size (mm) Limiting indices Theta range for data collection Reflections collected/unique Completeness Extinction coefficient Refinement method Data/restraints/parameters Goodness-of-fit on F2 Final R indices [I > 2sigma(I)] R indices (all data) Largest diff. peak and hole (e.A3)
C57H48Cl2Fe2N3O9 1456.08 296(2) K 0.71073 A Triclinic, P-1 12.5896(19) Å alpha = 84.493(9)° 15.533(2) Å beta = 77.649(9)° 19.816(3) Å gamma = 71.867(9)° 3595.7(9) 1, 1.377 1.345 0.901 1520 0.9 0.7 0.4 mm 16 h 16, 19 k 19, 25 l 25 1.05–25.00° 111059/15550 [R(int) = 02782] 99.5% (theta = 25.00) 0.036(4) Full-matrix least-squareson F2 12592/0/755 1.555 R1 = 0.1738, wR2 = 0.4444 R1 = 0.3006, wR2 = 0.5088 1.370 and 1.324
Table 2 Selected bond lengths (Å) and angles (°) for Fe2(LTHF)3. Coordination sphere of Fe1
Coordination sphere of Fe2
Fe(1)AO(3) Fe(1)AO(4) Fe(1)AO(5) Fe(1)AN(1) O(5)AC(8) O(3)AFe(1)AO(4) O(3)AFe(1)AO(2) O(4)AFe(1)AO(2) O(3)AFe(1)AO(1) O(4)AFe(1)AO(1) O(2)AFe(1)AO(1) O(3)AFe(1)AO(5) O(4)AFe(1)AO(5) O(2)AFe(1)AO(5) O(1)AFe(1)AO(5) O(3)AFe(1)AN(1)
Fe(2)AO(6) Fe(2)AO(7) Fe(2)AN(3) Fe(2)AN(2) O(7)AC(34) O(6)AFe(2)AO(7) O(6)AFe(2)AO(1) O(7)AFe(2)AO(1) O(6)AFe(2)AO(2) O(7)AFe(2)AO(2) O(1)AFe(2)AO(2) O(6)AFe(2)AN(3) O(7)AFe(2)AN(3) O(1)AFe(2)AN(3) O(2)AFe(2)AN(3) O(6)AFe(2)AN(2)
1.943(11) 1.938(10) 2.124(10) 2.239(13) 1.40(2) 168.7(4) 93.9(4) 94.9(4) 95.5(4) 93.6(4) 74.9(4) 87.1(4) 84.9(4) 173.7(4) 98.8(4) 86.3(5)
1.899(10) 1.897(9) 2.246(12) 2.260(13) 1.294(15) 106.7(4) 90.9(4) 160.4(4) 158.6(4) 91.6(4) 73.1(3) 89.4(4) 85.0(4) 86.8(4) 103.5(4) 83.8(5)
537, 473. UV–vis in CH2Cl2: kmax, nm (e, M1 cm1): 290 (20, 140), 540 (4088). 2.3. Spectrophotometric investigation of H2LTHF and Fe complexation
generated from a sealed tube source. Data were collected and reduced by smart and saint software [28] in the Bruker package. The structure was solved by direct methods [29], than developed by least squares refinement on F2 [30,31]. All non-H atoms were placed in calculated positions and refined as isotropic with the ‘‘riding-model technique’’. Details concerning collection and analysis are reported in Tables 1 and 2. 2.2. Preparations 2.2.1. Synthesis of ligand Ligand was synthesized according to the modified literature procedure [32]. 2.2.1.1. Synthesis of H2LTHF [6,60 -((((tetrahydrofuran-2-yl)methyl)azanediyl)bis(methylene))bis(2,4-dichlorophenol)]. A solution of 2,4-dichloro phenol (4.8 g, 29.00 mmol), (tetrahydrofuran-2-yl)methanamine (1.46 ml, 14.50 mmol), and 37% aqueous formaldehyde (4.31 mL, 58 mmol) was refluxed for 48 h. Upon cooling, a large quantity of beige solid was formed. The solvent was decanted, and the remaining solid residue was washed with cold methanol to give a pure, white powder (10.40 g, 80% yield) (Table 2). 1 H NMR (400 MHz, CDCl3, 298 K): d 1.47 (m, 1H); 1.94 (m, 2H); 2.01 (m, 1H); 2.58 (m, 1H); 2.68 (m, 1H); 3.75 (d, 2H); 3.95 (d, 2H);3.97 (m, 2H); 4.23 (m, 1H); 6.98 (d, 2H); 7.30 (d, 2H). 13C NMR (100.6 MHz, CDCl3, 298 K): d 151.1, 129.2, 128.5, 124.2, 124.1, 122.0, 68.6, 56.6, 56.15, 29.8, 25.1. IR (cm1): 3405 (OH); 2965 (CAH); 1564 (C@C, phenyl ring). m.p. 150 °C. 2.2.2. Synthesis of Fe2(LTHF)3 Complex Sodium methoxide (0.20 g, 2 mmol) was added to a solution of H2LTHF (0.448 g, 1 mmol) in 50 ml acetonitrile. FeCl36H2O (0.270 g, 1.00 mmol) was added to this solution and the resulting mixture was refluxed for 1 h, resulting in an intense purple solution. Solvent was removed and the solid material crystallized in 1:1 dichloromethane/hexane mixture. Yield: 0.35 g (65%). Anal. calcd for C38H34Cl8Fe2N2O6 (1005.9 g/mol): C, 45.0; H, 4.40; N, 2.80. Found: C, 44.80; H, 4.20; N, 2.60%. IR (KBr, cm1): 3433, 2955, 1630, 1458, 1385, 1308, 1271, 1200, 1171, 1135, 1022, 922, 874, 836, 746, 599,
In an experiment, 2 ml of Fe(NO3)34H2O solution in methanol (2.5 104 M) was transferred into a cuvette. UV–vis spectra were recorded in the range of 300–800 nm about 5 min after each addition of 10 lL of H2LTHF (5 103 M) solution. Changes in the absorbance of iron nitrate complex upon addition of H2LTHF solution were monitored at the LMCT maximum wavelength of complexes. 3. Results and discussion [6,60 -((((tetrahydrofuran-2-yl)methyl)azanediyl)bis(methylene)) bis(2,4-dichlorophenol)] H2LTHF) was prepared from (tetrahydrofuran-2-yl)methanamine, formaldehyde and 2,4-di-chloro phenol, in two steps Mannich condensation. Aminophenol ligands are usually prepared in methanol solutions but we describe herein green synthesis of these reactions when water is used as the reaction medium. In fact, since aqueous formaldehyde is one of the reagents, no additional solvent is required beyond that is already present in the reagent solution. Iron complex was formed in good yields by refluxing methanolic solution of the ligand with iron(III) chloride and base in suitable ratio as follows.
H2 L þ FeCl2 4H2 O þ Base
Reflux
!
MeOH=1 h
Fe2 ðLTHF Þ3
ð1Þ
In IR spectra of complex, the strong and sharp OH stretching band of the phenols around 3300–3500 cm1 for the mOH stretch of ligands were replaced by a broad band, proving the coordination of phenol groups to the metal. The electronic absorption spectra of complex have been measured in dichloromethane in the 300–800 nm range, exhibit one band in the near UV and one more in the visible region. For all the present, the lower energy band (400–600 nm) is assigned to a charge-transfer transition from the phenolate (pp) to the halffilled dp orbital of iron atom, and thus its band position falls in the range of other phenolato compounds (Section 2.2.2). The higher energy band (300 nm) is associated with p/p intera-ligand charge transfer, ILCT [33–35]. Comparing the transition energies with ones for similar complex with amin-tert-butyl substituted phenol ligands [36] shows a blue shift for both transitions. It can attribute to the
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Fig. 1. ORTEP diagram and atom labeling scheme for complex Fe2(LTHF)3.
electron-withdrawing effect of chlorine substituent on stabilizing phenolate p (HOMO) orbital. 3.1. Description of crystal structure Red brown crystals of Fe2(LTHF)3 for X-ray analysis were obtained from methanol/dichloromethane (1:2) solution. Crystal data and details of data collection and refinement are summarized in Table 1. A complete listing of bond lengths and angles for these complexes can be found in the supplementary material. The structural data for Fe2(LTHF)3 have been deposited with the Cambridge Crystallographic Data Centre, the deposition number being CCDC 832084. The structure of complex Fe2(LTHF)3 consists of an unsymmetrical binuclear unit which passes through the bridging phenoxo oxygen atoms O(1) and O(2). Selected bond distances and angles are listed in Table 1. An ORTEP view of the dimeric unit is shown in Fig. 1. The iron ion, Fe(1) is in deformed octahedral coordination sphere in which the phenolate O(3), O(4), THF O(5), the bridging O(1), O(2) and the amine N(1) of one aminophenol ligand occupy six positions. Fe(2) is in distorted octahedral coordination sphere. O(6), O(7) and the bridging phenolates O(1), O(2) occupy the basal positions whilst the and amine N(2) and N(3) of two aminophenol ligands lie at the apical position of the trigonal bipyramid. The O(8) and O(9) of THF moieties have not been coordinated to Fe(2). Therefore, the structure is unsymmetrical with Fe2O2 – plane formed by the phenolate oxygen atoms. 3.2. Spectral data analysis of H2LTHFAFe complexation The titration of ligand (H2LTHF) solution at fixed concentration of iron and varying concentration of H2LTHF have been conducted. The titration spectra of iron nitrate upon increasing addition of H2LTHF are shown in Fig. 2. During the titration, the hypochromicity was
Fig. 2. The titration absorption spectra of Fe(NO3)34H2O (2.5 104 M) by H2LTHF.
observed without any shift in kmax band of 540 nm, which presents the existence of interaction between H2LTHF and the iron ion. The appearance of two isosbestic points in iron nitrate spectraclearly indicates the existence of simple equilibrium between H2LTHF and H2LTHF AFe complex. The complex composition of 3:2 was determined by plotting the absorption changes vs. ligand to metal mole ratio (nL/nFe). Fig. 3 shows a mole ratio plot of iron nitrate upon increasing addition of H2LTHF. 3.3. Magnetic susceptibility measurement Magnetic measurements were performed on a Quantum Design MPMS XL-5 SQUID magnetometer. The dc magnetic susceptibility
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Fig. 3. Determination of Fe2(LTHF)3 complex composition by the mole ratio plot. Fig. 5. Cyclic voltammograms of Fe2(LTHF)3 in CH2Cl2. The scan rate is 200 mV/s and T = 40 °C.
v = M/H was determined in the temperature range 2–300 K in a magnetic field H of 1000 Oe. The measured susceptibility monotonically increases with decreasing temperature as is shown in Fig. 4. In the inset of Fig. 4, we show a temperature dependence of effective magnetic moment calculated per binuclear unit Fe2(LTHF)3. At room temperature the effective magnetic moment leff is 7.7 lB which is less than the calculated (8.3 lB) for two uncoupled high-spin (S = 5/2) Fe(III) ions. The effective magnetic moment decreases upon cooling showing an antiferromagnetic coupling between Fe(III) ions in a unit. According to the binuclear structure of the complex we have described the measured susceptibility with the v(T) dependence that follows from the isotropic interaction Hamiltonian H = JSA SB [37], where SA = SB = 5/2 for high-spin Fe(III) ions. We add a temperature constant term to the function v(T) derived in Ref. [37] that accounts for diamagnetic contribution of inner shell electrons and temperature independent paramagnetism. The result of a fitting procedure is shown as a full line in inset to Fig. 4 giving us the interaction J = 0.3 cm1. The interaction J is negative as expected for an antiferromagnetic coupling. It is rather small compared to the interactions in similar bridged compounds collected in [38]. However, a relatively small interaction J is in agreement with our experimental data where no maximum in temperature dependent susceptibility, which is a fingerprint of strong antiferromagnetic coupling, has been observed.
3.4. Electrochemistry Cyclic voltammogram (CV) of complex Fe2(LTHF)3 have been recorded in CH2Cl2 solutions containing 0.1 M [(nBu)4N]ClO4 as a supporting electrolyte. The CV voltammograms (CV) of Fe2(LTHF)3 complex revealed oxidations and reductions peaks (Fig. 5). Complex Fe2(LTHF)3 exhibits two reversible oxidation processes at 0.840 V and 1.3 V, corresponds to ligand-centered peaks as seen in Fig. 5. The metal-centered voltammogram shows two quasireversible peaks in the negative potential range, which corresponds to the FeIIIFeIII/FeIIIFeII and FeIIIFeII/FeIIFeII reduction of Fe2(LTHF)3. The anodic to cathodic peak separation is 120 mV and increases with scan rates (Fig. 6), indicating that the process deviates from a kinetic reversibility and the electron transfer at the electrode surface is relatively slow. Fig. 7 shows the cyclic voltammograms of H2LTHF ligand and Fe2(LTHF)3 complex in positive scan region at 40 °C. In both voltammograms two anodic peaks appear; A1 and A2 for complex and Á1 and Á2 for ligand. These two anodic peaks are related to the oxidation of phenolic moieties of ligands (known as ligand-centered potentials) to their semiquinone and quinine species. There are two considerable differences between voltammograms of H2LTHF ligand and Fe2(LTHF)3 complex. (a) Both oxidation steps of Fe2(LTHF)3 complex are reversible at 40 °C with two well-defined
8 6
μeff (BM)
χ (emu/mol)
1.5
1.0
4 2
0.5
0.0
0
0
H = 1 kOe 0
100
100
T (K) 200
200
300
300
T (K) Fig. 4. Temperature dependent susceptibility and effective magnetic moment (inset) of Fe2(LTHF)3 in magnetic field of 1 kOe. The full line in inset is a fit with theoretical susceptibility of binuclear S = 5/2 units as described in text.
Fig. 6. Cyclic voltammograms of Fe2(LTHF)3 (3 103 M) in CH2Cl2 with M [(nBu)4N]ClO4 as supporting electrolyte. Potentials are referenced vs. Fc, Scan rates are 50, 100, 200, 300, 400, 500 and 600 mV/s and T = 40 °C.
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Fig. 7. Cyclic voltammograms of Fe2(LTHF)3 and ligand in CH2Cl2 The scan rate is 200 mV/s and T = 40 °C.
counter parts, C´1 and C´2 which are related to the reduction of semiquinone and quinone forms. It has seen that the cyclic voltammograms of ligand do not show any reduction peaks. The reversibility of cyclic voltammograms in Fe2(LTHF)3 comparing to its free ligand can attributed to the role of Lewis acidity of iron center on stabilization of phenoxyl radical bound to it, as seen in ribonucleotide reductase, RNR enzyme. (b) comparing the oxidation potential of Fe2(LTHF)3 with H2LTHF show a positive shift in the anodic peak potentials of complexes, 460 mV for A1–Á1 and 450 mV for A2–Á2. It may be related to stability of complex and lower density of negative charge on phenolate group which coordinated to iron (III) center. In order to investigate the effect of electronic and steric effects of substituents peak potentials of Fe2(LTHF)3 were compared with previously reported tert-butyl substituted complex Fe2(LtBu)2 [36]. The half wave potentials for both oxidation peaks (Fig. 8) of Fe2(LTHF)3 shift to more positive values comparing to Fe2(LtBu)2. This criterion can attribute to more highly positive charge density on iron center due to electron-withdrawing effect of chlorine substituent’s on phenolate moities in Fe2(LTHF)3. It can lead us to this fact that despite of the more steric effect around radical center in Fe2(LtBu)2 on radical stabilization, Fe2(LTHF)3 show a good phenoxyl radical stability. In other word, one can say that sometimes electronic parameters are more important than steric parameters for radical stabilization.
Fig. 9. Cyclic voltammograms of Fe2(LTHF)3 in CH2Cl2 at different temperatures. The scan rate is 200 mV/s.
3.5. Investigation of thermodynamic stability of Fe2(LTHF)3 radical complex We shed light into the stability constants and thermodynamic parameters of the electrochemically production of Fe2(LTHF)3 radical complex, Fe2 ðLTHF Þ3 through characterization of the energetic governing radical complex formation. In this way, the voltammetric study of complex was performed at different temperatures. There is a regular shift in the peak potentials of complex, A1. This peak shifts to more positive potentials by decreasing temperature (Fig. 9). The height of cathodic peak decreases by increasing the temperature which is in good accordance with instability of produced Fe2(LTHF)3 radical complex at higher temperatures. Based on the shift in peak potentials, temperature dependence formation constants were obtained by Nernst Equation as follows.
2RT ðLnK f Þ DE1=2 ¼ 3nF
ð2Þ
The energetic of Fe2(LTHF)3 equilibrium can be conveniently determined by thermodynamic parameters such as Gibbs free energy (DG). The standard Gibbs energy change is usually calculated from the equilibrium constant, K, of the electrochemical process, by the following relationship:
DG0 ¼ RTLnK f
ð3Þ
where R and T are the gas constant and the absolute temperature, respectively. Since the activity coefficients of the reactions are not known, the usual procedure is to assume them unity and to use the equilibrium concentrations instead of the activity. Therefore, it would be appropriate to adjust the terminology of apparent equilibrium constant K0 , and Gibbs energy DG°0 . Apparent standard enthalpies per mole of cooperative unit can be obtained from the dependence on temperature of the apparent formation constant K0 , by van’t Hoff equation:
0 DH 1 @ @LnK f ¼ T R
ð4Þ
This is the so-called van’t Hoff enthalpy. The apparent standard entropy change, DS°0 , can be derived from the following equation:
DS0 ¼ ðDH0 DG0 Þ=T Fig. 8. Cyclic voltammograms of Fe2(LTHF)3 and Fe2(LtBu)2 in CH2Cl2. The scan rate is 200 mV/s and T = 40 °C.
ð5Þ
The van’t Hoff plot for electrochemically produced radical complex of Fe2 ðLTHF Þ3 is shown in Fig. 10.
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Acknowledgments Authors are grateful to the Institute for Advanced Studies in Basic Sciences (IASBSs), University of Ljubljana and Università di Messina. E. Safaei gratefully acknowledges the support by the Institute for Advanced Studies in Basic Sciences (IASBS) Research Council under Grant No. G2012IASBS127. Appendix A The structural data for Fe2(LTHF)3 has been deposited with the Cambridge Crystallographic Data Centre, the deposition number being CCDC 832084. These data can be obtained free of charge via http://www.ccdc.cam.ac.uk/conts/retrieving.html, or from the Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336 033; or e-mail: [email protected]. Fig. 10. plot of LnKf vs. 1/T according to the van’t Hoff plot.
Appendix B. Supplementary material Table 3 Formation constants and related thermodynamic parameters of Fe2(LTHF)3. T/K
log K0 (M1)
DG°0 (kJ mol1)
DH°0 (kJ mol1)
DS°0 (J mol1 K1)
298 283 248 233
3.4 3.9 4.5 5.0
19.3 ± 0.310 21.2 ± 0.100 21.5 ± 0.020 22.4 ± 0.008
21.3 ± 0.6 21.3 ± 0.6 21.3 ± 0.6 21.3 ± 0.6
6.8 ± 0.0031 0.3 ± 0.0025 1.0 ± 0.0025 4.6 ± 0.0026
The calculated thermodynamic parameters are listed in Table 3. The negatively large standard free energy changes (DG°0 ) implies that Fe2 ðLTHF Þ3 product is favored over reactant. The more negative values of DG°0 at lower temperatures can be related to the higher stability of Fe2 ðLTHF Þ3 because of decreasing radical side reactions. It has also been indicated that formation of Fe2 ðLTHF Þ3 is an exothermic process. In terms of driving forces, the present case implies the stabilization of phenoxyl radicals (LTHF) due to their interaction with iron centers. The considerable negative values of DH°0 and also small positive (or negative) values of DS°0 in mentioned electrochemical process indicate that the contribution of the negative enthalpy changes (DH°0 ) results in more negative DG°0 and favoring the radical complex production process (Table 3), it seems that the major contributing factor in radical stabilizing is enthalpic in origin. It can be concluded that the negative enthalpy changes are the driving forces in interaction between iron center and ligand phenoxyl radical (LTHF). 4. Conclusions A kind of new [NAO]-donor tripodal amine chloro-substituted phenol ligand and its iron(III) complex of the type Fe2 LTHF 3 , have been synthesized and characterized. X-ray crystal structure of Fe2 LTHF reveal that this complex has unsymmetrical dimeric struc3 ture, in which two Fe(III) centers are surrounded by amine nitrogen’s, phenolate atoms of three ligands. Magnetostructural studies of this complex displayed an antiferromagnetic coupling between Fe(III) ions through phenolate bridges. It has been shown that electrochemical oxidations are ligand-centered, i.e. formation of the phenoxyl radicals from the coordinated phenolates, while reductions are metal-centered related to FeIII/FeII couple. Our electrochemically generated Fe2(LTHF)3 phenoxyl radical complex (Fe2 ðLTHF Þ3 ) shows a good stability and this process is exothermic due to the predominant role of phenoxyl radical (LTHF)/iron interaction. This means that metal ions can be good candidates as radical stabilizers as seen in enzymes for application in new molecular switches and/or data storage devices [11].
Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.molstruc.2012. 04.084. References [1] L. Benisvy, A.J. Blake, D. Collison, E.S. Davies, C.D. Garner, E.J.L. McInnes, J. McMaster, G. Whittaker, C. Wilson, Chem. Commun. (2001) 1824–1825. [2] S. Kimura, E. Bill, E. Bothe, T. Weyhermüller, K. Wieghardt, J. Am. Chem. Soc. 123 (2001) 6025–6039. [3] J.W. Whittaker, Chem. Rev. 103 (2003) 2347–2363. [4] C. Wright, A.G. Sykes, J. Inorg. Biochem. 85 (2001) 237–243. [5] J. Stubbe, W.A. van der Donk, Chem. Rev. 98 (1998) 705–762. [6] J. Stubbe, D.G. Nocera, C.S. Yee, M.C.Y. Chang, Chem. Rev. 103 (2003) 2167– 2202. [7] K.E. Silva, T.E. Elgren, L. Que, M.T. Stankovich, Biochemistry 34 (1995) 14093– 14103. [8] A. Graslund, M. Sahlin, B.-M. Sjobergt, Environ. Health Perspect. 64 (1985) 139–149. [9] M. Gbom, M. Galander, M. Andersson, M. Kolberg, W. Hofbauer, G. Lassmann, P. Nordlund, F. Lendzian, PNAS 100 (2003) 3209–3214. [10] R.G. Hicks, Stable Radicals: Fundamentals and Applied Aspects of OddElectron Compounds, in: R.G. Hicks (Ed.), John Wiley & Sons Ltd., West Sussex, 2010. [11] L. Benisvy, E. Bill, A.J. Blake, D. Collison, E.S. Davies, C.D. Garner, E.J.L. McInnes, J. McMaster, S. Ross, C. Wilson, Dalton. Trans. (2006) 258–267. [12] K.U. Ingold, Adv. Chem. Ser. 75 (1968) 296–305. [13] L.R. Mahoney, Angew. Chem. Int. Ed. Engl. 8 (1969) 547–555. [14] J.A. Howard, Adv. Free Radical Chem. 4 (1972) 49–173. [15] E.B. Burlakova, Russ. Chem. Rev. 44 (1975) 871–880. [16] I.V. Khudyakov, V.A. Kuzmin, Russ. Chem. Rev. 44 (1975) 801–815. [17] G.W. Burton, K. Ingold, Acc. Chem. Res. 19 (1986) 194–201. [18] S. Steenken, P. Neta, In the Chemistry of Phenols, in: Z. Rappoport (Ed.), Wiley, New York, 2003. [19] R.H. Holm, P. Kennepohl, E. Solomon, Chem. Rev. 96 (1996) 2239–2314. [20] M. Velusamy, M. Palaniandavar, R.S. Gopalan, G.U. Kulkarni, Inorg. Chem. 42 (2003) 8283–8293. [21] R. Viswanathan, M. Palaniandavar, T. Balasubramanian, T.P. Muthiah, Inorg. Chem. 37 (1998) 2943–2951. [22] L. Que, W.B. Tolman, Angew. Chem. Int. Ed. 41 (2002) 1114–1137. [23] P.J. Cappillino, P.C. Tarves, G.T. Rowe, A.J. Lewis, M. Harvey, C. Rogge, A. Stassinopoulos, W. Lo, W.H. Armstrong, J.P. Caradonna, Inorg. Chim. Acta 362 (2009) 2136–2150. [24] O.C. Belle, I.G. Luneau, J.L. Pierre, C. Scheer, Inorg. Chem. 35 (1996) 3706–3708. [25] M. Suzuki, S. Fujinami, T. Hibino, H. Hori, Y. Maeda, A. Uehara, M. Suzuki, Inorg. Chim. Acta 283 (1998) 124–135. [26] A. Neves, S.M.D. Erthal, V. Drago, K. Griesar, W. Haase, Inorg. Chim. Acta 197 (1992) 121–124. [27] A. Neves, M.A. de Brito, V. Drago, K. Griesar, W. Haase, Y.P. Mascarenhas, Inorg. Chim. Acta 214 (1993) 5–8. [28] Bruker AXS Inc., SMART, Version 5.060 and SAINT, Version 6.02, Wisconsin, USA, 1999. [29] M.C. Burla, R. Caliandro, M. Camalli, B. Carrozzini, G.L. Cascarano, L. De Caro, C. Giacovazzo, G. Polidori, R. Spagna, J. Appl. Crystallogr. 38 (2005) 381–388. [30] G.M. Sheldrick, SHELXL97, Program for Crystal Structure Refinement, University of Gttingen, Germany, 1997. [31] SHELXT LN, Version 5.10, Bruker Analytical X-ray Inc., Madison, WI, USA, 1998.
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[32] E. Safaei, H. Sheykhi, A. Wojtczak, Z. Jaglicic, A. Kozakiewicz, Polyhedron 30 (2011) 1143–1148. [33] M.I. Davies, A.M. Orville, F. Neese, J.M. Zaleski, J.D. Lipscomb, E.I. Solomon, J. Am. Chem. Soc. 124 (2002) 602–614. [34] R. Mayilmurugan, E. Suresh, M. Palaniandavar, Inorg. Chem. 46 (2007) 6038– 6049.
[35] R. Mayilmurugan, K. Visvaganesan, E. Suresh, M. Palaniandavar, Inorg. Chem. 48 (2009) 8771–8783. [36] E. Safaei, T. Weyhermüller, E. Bothe, K. Wieghardt, P. Chaudhuri, Eur. J. Inorg. Chem. (2007) 2334–2344. [37] O. Khan, Molecular Magnetism, VCH Publishers Inc., New York, 1993. [38] R. Werner, S. Ostrovsky, K. Griesar, W. Haase, Inorg. Chim. Acta 326 (2001) 78–88.
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A biologically relevant iron(III) phenoxyl radical complex: A thermodynamic investigation on the structure-radical stability relationship Iraj Saberikia a, Elham Safaei a,⇑, Mohammad Rafiee a, Patricia Coticˇ b, Giuseppe Bruno c, Hadi Amiri Rudbari c a b c
Institute for Advanced Studies in Basic Sciences (IASBS), 45195 Zanjan, Iran Institute of Mathematics, Physics and Mechanics, University of Ljubljana, Ljubljana, Slovenia Dipartimento di Chimica Inorganica, Vill. S. Agata, Salita Sperone 31, Università di Messina, 98166 Messina, Italy
h i g h l i g h t s " A binuclear Fe(III) complex of aminophenol ligand synthesized. " Thermodynamic parameters of electrochemically produced radicals were investigated. " Radical complex/Fe(III) center interaction make it stabilized such as RNR enzyme.
a r t i c l e
i n f o
Article history: Received 6 March 2012 Accepted 27 April 2012 Available online 9 May 2012 Keywords: Ribonucleotide reductase Aminophenol Model complex Iron complexes
a b s t r a c t A kind of new amine-chloro substituted phenol and its iron(III) complex has been prepared and characterized by spectroscopic, X-ray techniques and magnetic susceptibility studies. X-ray analysis revealed a binuclear complex, Fe2(LTHF)3 in which two Fe(III) centers were surrounded by three ligands. The magnetic susceptibility indicates antiferromagnetic coupling between two iron centers through phenolate bridges. Voltammetric study of ligand (H2LTHF) and complex Fe2(LTHF)3 revealed that the coordinated phenolate ligands undergo reversible two-electron oxidations with formation of coordinated phenoxyl radicals. It has seen that the oxidation of Fe2(LTHF)3 comparing to its analog, Fe2(LtBu)2 occurs at significantly higher potentials. It was attributed to the relationship between the radical stability with electronic and steric properties of ligands. Stability constants and thermodynamic parameters for the formation of Fe2(LTHF)3 complex and the role of ligand coordination to iron center in phenoxyl radical stabilization has been investigated using anodic peaks potentials. Ó 2012 Elsevier B.V. All rights reserved.
1. Introduction Radicals are highly reactive species due to their unpaired electrons. They have a key role in living organism’s physiology, biochemistry and some chemical processes, such as polymerization. Interest in the field of protein and amino acid radicals has been dramatically increasing because these active species serve as enzyme co-factors to lots of enzymatic reactions in living organisms. [1]. One of the most important amino acid radicals is tyrosyl radical which plays vital roles in the chemistry of living systems [1,2]. Enzymes that incorporate this kind of radical include active form of galactose oxidase (GO) [1,3,4], glyoxal oxidase (GLO) [2] and ribonucleotide reductase, RNR (Scheme 1) [5–9]. This last one is the only known enzyme that carries a stable free radical in its resting state and catalyzes DNA synthesis by ⇑ Corresponding author. Tel.: +98 241 4153200; fax: +98 241 4153232. E-mail address: [email protected] (E. Safaei). 0022-2860/$ - see front matter Ó 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.molstruc.2012.04.084
reducing ribonucleotides to their corresponding deoxyribonucleotides (Scheme 2). It has shown that the catalytic action of this enzyme involve formation of a tyrosyl radical close (tyrosyl 122 radical, Y122), to the redox active iron center which is responsible for the radical stability in RNR [8,9]. Stabilization of the mentioned radicals is very important due to their role in biological processes [1,10]. Scientists have found two mechanisms for tyrosyl radical stabilization: 1. Stabilization by hydrogen bonding to its neighbor amino acid residues (e.g. in GO). 2. Stabilization by interaction with a transition metal center (e.g. in RNR). The key role of the second mechanism in biological and catalytic radical processes is due to two major reasons [1,11]: (a) Redox reaction assisted radical generation. (b) Interaction of transition metals with their adjacent radicals.
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Scheme 1. The enzyme structure ribonucleotide reductase (RNR) [9].
Scheme 2. Enzymatic reaction catalyzed by ribonucleotide reductase [8].
On the other hand, the ‘‘metal–radical’’ approach has been noticed by scientists as one of the more attractive strategies for the design of biocatalysts and molecule-based materials. Therefore, nature and properties of d-transition metal bound to one or more radicals is much interest. Phenoxyl type radicals are one of these groups that their chemistry has been much interest because of key participation in biological redox processes [12–18]. Because of their close relationship with tyrosine-containing metalloenzymes, iron and copper complexes of chelating amine-phenolate ligands have been studied [19–27]. We report here the coordination chemistry, magnetic and redox properties of [6,60 -((((tetrahydrofuran-2-yl)methyl)azanediyl) bis(methylene))bis(2,4-dichlorophenol)] (H2LTHF) (Scheme 3) and its iron(III) complex with this ligand. In this way, the role of phenoxyl radical/iron(III) center interaction in radical stabilization has been investigated.
Scheme 3. Structure of [6,60 -((((tetrahydrofuran-2-yl)methyl) azanediyl)bis(methylene))bis(2,4-dichlorophenol)] (H2LTHF).
2. Experimental 2.1. Materials and physical measurements Reagents or analytical grade materials were obtained from commercial suppliers and used without further purification, except those for electrochemical measurements. Elemental analyses (C H N) were performed by the Research Institute of Petroleum Industry (RIPI). Fourier transform infrared spectroscopy on KBr pellets was performed on a FT-IR Bruker Vector 22 instrument. NMR measurements were performed on a Bruker 400 instrument. UV–Vis absorbance digitized spectra were collected using a CARY 100 spectrophotometer. Magnetic susceptibility were measured from powder samples of solid material in the temperature range 2–300 K by using a SQUID susceptometer (Quantum Design MPMS-XL-5) in a magnetic field of 1000 Oe. Voltammetric measurements were made with a computer controlled electrochemical system (ECO Chemie, Ultrecht, The Netherlands) equipped with a PGSTA 30 model and driven by GPES (ECO Chemie). The reference electrode was an Ag wire as the quasi reference electrode. A glassy carbon electrode with a surface area of 0.035 cm2 was used as a working electrode and a platinum wire served as the counter electrode. Prior to the measurement, the GC electrode was polished with 0.1 mm alumina powder and washed with distilled water. The voltage scan rate was set at 50 mV sec1. The solutions were deoxygenated by bubbling nitrogen gas through them. Ferrocene was added as an internal standard after completion of a set of experiments, and potentials are referenced vs. the ferrocenium/ferrocene couple (Fc+/Fc). Diffraction data for Fe2(LTHF)3 was measured on a Bruker–Nonius X8 ApexII diffractometer equipped with a CCD area detector by using graphite-monochromated Mo Ka radiation (k = 0.71073)
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I. Saberikia et al. / Journal of Molecular Structure 1022 (2012) 109–116 Table 1 Crystal data and structure refinement for Fe2(LTHF)3. Empirical formula Formula weight Temperature (K) Wavelength (Å) Crystal system Space group Unit cell dimensions Volume (Å3) Z Calculated density (Mg/m3) Absorption coefficient (mm1) F(0 0 0) Crystal size (mm) Limiting indices Theta range for data collection Reflections collected/unique Completeness Extinction coefficient Refinement method Data/restraints/parameters Goodness-of-fit on F2 Final R indices [I > 2sigma(I)] R indices (all data) Largest diff. peak and hole (e.A3)
C57H48Cl2Fe2N3O9 1456.08 296(2) K 0.71073 A Triclinic, P-1 12.5896(19) Å alpha = 84.493(9)° 15.533(2) Å beta = 77.649(9)° 19.816(3) Å gamma = 71.867(9)° 3595.7(9) 1, 1.377 1.345 0.901 1520 0.9 0.7 0.4 mm 16 h 16, 19 k 19, 25 l 25 1.05–25.00° 111059/15550 [R(int) = 02782] 99.5% (theta = 25.00) 0.036(4) Full-matrix least-squareson F2 12592/0/755 1.555 R1 = 0.1738, wR2 = 0.4444 R1 = 0.3006, wR2 = 0.5088 1.370 and 1.324
Table 2 Selected bond lengths (Å) and angles (°) for Fe2(LTHF)3. Coordination sphere of Fe1
Coordination sphere of Fe2
Fe(1)AO(3) Fe(1)AO(4) Fe(1)AO(5) Fe(1)AN(1) O(5)AC(8) O(3)AFe(1)AO(4) O(3)AFe(1)AO(2) O(4)AFe(1)AO(2) O(3)AFe(1)AO(1) O(4)AFe(1)AO(1) O(2)AFe(1)AO(1) O(3)AFe(1)AO(5) O(4)AFe(1)AO(5) O(2)AFe(1)AO(5) O(1)AFe(1)AO(5) O(3)AFe(1)AN(1)
Fe(2)AO(6) Fe(2)AO(7) Fe(2)AN(3) Fe(2)AN(2) O(7)AC(34) O(6)AFe(2)AO(7) O(6)AFe(2)AO(1) O(7)AFe(2)AO(1) O(6)AFe(2)AO(2) O(7)AFe(2)AO(2) O(1)AFe(2)AO(2) O(6)AFe(2)AN(3) O(7)AFe(2)AN(3) O(1)AFe(2)AN(3) O(2)AFe(2)AN(3) O(6)AFe(2)AN(2)
1.943(11) 1.938(10) 2.124(10) 2.239(13) 1.40(2) 168.7(4) 93.9(4) 94.9(4) 95.5(4) 93.6(4) 74.9(4) 87.1(4) 84.9(4) 173.7(4) 98.8(4) 86.3(5)
1.899(10) 1.897(9) 2.246(12) 2.260(13) 1.294(15) 106.7(4) 90.9(4) 160.4(4) 158.6(4) 91.6(4) 73.1(3) 89.4(4) 85.0(4) 86.8(4) 103.5(4) 83.8(5)
537, 473. UV–vis in CH2Cl2: kmax, nm (e, M1 cm1): 290 (20, 140), 540 (4088). 2.3. Spectrophotometric investigation of H2LTHF and Fe complexation
generated from a sealed tube source. Data were collected and reduced by smart and saint software [28] in the Bruker package. The structure was solved by direct methods [29], than developed by least squares refinement on F2 [30,31]. All non-H atoms were placed in calculated positions and refined as isotropic with the ‘‘riding-model technique’’. Details concerning collection and analysis are reported in Tables 1 and 2. 2.2. Preparations 2.2.1. Synthesis of ligand Ligand was synthesized according to the modified literature procedure [32]. 2.2.1.1. Synthesis of H2LTHF [6,60 -((((tetrahydrofuran-2-yl)methyl)azanediyl)bis(methylene))bis(2,4-dichlorophenol)]. A solution of 2,4-dichloro phenol (4.8 g, 29.00 mmol), (tetrahydrofuran-2-yl)methanamine (1.46 ml, 14.50 mmol), and 37% aqueous formaldehyde (4.31 mL, 58 mmol) was refluxed for 48 h. Upon cooling, a large quantity of beige solid was formed. The solvent was decanted, and the remaining solid residue was washed with cold methanol to give a pure, white powder (10.40 g, 80% yield) (Table 2). 1 H NMR (400 MHz, CDCl3, 298 K): d 1.47 (m, 1H); 1.94 (m, 2H); 2.01 (m, 1H); 2.58 (m, 1H); 2.68 (m, 1H); 3.75 (d, 2H); 3.95 (d, 2H);3.97 (m, 2H); 4.23 (m, 1H); 6.98 (d, 2H); 7.30 (d, 2H). 13C NMR (100.6 MHz, CDCl3, 298 K): d 151.1, 129.2, 128.5, 124.2, 124.1, 122.0, 68.6, 56.6, 56.15, 29.8, 25.1. IR (cm1): 3405 (OH); 2965 (CAH); 1564 (C@C, phenyl ring). m.p. 150 °C. 2.2.2. Synthesis of Fe2(LTHF)3 Complex Sodium methoxide (0.20 g, 2 mmol) was added to a solution of H2LTHF (0.448 g, 1 mmol) in 50 ml acetonitrile. FeCl36H2O (0.270 g, 1.00 mmol) was added to this solution and the resulting mixture was refluxed for 1 h, resulting in an intense purple solution. Solvent was removed and the solid material crystallized in 1:1 dichloromethane/hexane mixture. Yield: 0.35 g (65%). Anal. calcd for C38H34Cl8Fe2N2O6 (1005.9 g/mol): C, 45.0; H, 4.40; N, 2.80. Found: C, 44.80; H, 4.20; N, 2.60%. IR (KBr, cm1): 3433, 2955, 1630, 1458, 1385, 1308, 1271, 1200, 1171, 1135, 1022, 922, 874, 836, 746, 599,
In an experiment, 2 ml of Fe(NO3)34H2O solution in methanol (2.5 104 M) was transferred into a cuvette. UV–vis spectra were recorded in the range of 300–800 nm about 5 min after each addition of 10 lL of H2LTHF (5 103 M) solution. Changes in the absorbance of iron nitrate complex upon addition of H2LTHF solution were monitored at the LMCT maximum wavelength of complexes. 3. Results and discussion [6,60 -((((tetrahydrofuran-2-yl)methyl)azanediyl)bis(methylene)) bis(2,4-dichlorophenol)] H2LTHF) was prepared from (tetrahydrofuran-2-yl)methanamine, formaldehyde and 2,4-di-chloro phenol, in two steps Mannich condensation. Aminophenol ligands are usually prepared in methanol solutions but we describe herein green synthesis of these reactions when water is used as the reaction medium. In fact, since aqueous formaldehyde is one of the reagents, no additional solvent is required beyond that is already present in the reagent solution. Iron complex was formed in good yields by refluxing methanolic solution of the ligand with iron(III) chloride and base in suitable ratio as follows.
H2 L þ FeCl2 4H2 O þ Base
Reflux
!
MeOH=1 h
Fe2 ðLTHF Þ3
ð1Þ
In IR spectra of complex, the strong and sharp OH stretching band of the phenols around 3300–3500 cm1 for the mOH stretch of ligands were replaced by a broad band, proving the coordination of phenol groups to the metal. The electronic absorption spectra of complex have been measured in dichloromethane in the 300–800 nm range, exhibit one band in the near UV and one more in the visible region. For all the present, the lower energy band (400–600 nm) is assigned to a charge-transfer transition from the phenolate (pp) to the halffilled dp orbital of iron atom, and thus its band position falls in the range of other phenolato compounds (Section 2.2.2). The higher energy band (300 nm) is associated with p/p intera-ligand charge transfer, ILCT [33–35]. Comparing the transition energies with ones for similar complex with amin-tert-butyl substituted phenol ligands [36] shows a blue shift for both transitions. It can attribute to the
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Fig. 1. ORTEP diagram and atom labeling scheme for complex Fe2(LTHF)3.
electron-withdrawing effect of chlorine substituent on stabilizing phenolate p (HOMO) orbital. 3.1. Description of crystal structure Red brown crystals of Fe2(LTHF)3 for X-ray analysis were obtained from methanol/dichloromethane (1:2) solution. Crystal data and details of data collection and refinement are summarized in Table 1. A complete listing of bond lengths and angles for these complexes can be found in the supplementary material. The structural data for Fe2(LTHF)3 have been deposited with the Cambridge Crystallographic Data Centre, the deposition number being CCDC 832084. The structure of complex Fe2(LTHF)3 consists of an unsymmetrical binuclear unit which passes through the bridging phenoxo oxygen atoms O(1) and O(2). Selected bond distances and angles are listed in Table 1. An ORTEP view of the dimeric unit is shown in Fig. 1. The iron ion, Fe(1) is in deformed octahedral coordination sphere in which the phenolate O(3), O(4), THF O(5), the bridging O(1), O(2) and the amine N(1) of one aminophenol ligand occupy six positions. Fe(2) is in distorted octahedral coordination sphere. O(6), O(7) and the bridging phenolates O(1), O(2) occupy the basal positions whilst the and amine N(2) and N(3) of two aminophenol ligands lie at the apical position of the trigonal bipyramid. The O(8) and O(9) of THF moieties have not been coordinated to Fe(2). Therefore, the structure is unsymmetrical with Fe2O2 – plane formed by the phenolate oxygen atoms. 3.2. Spectral data analysis of H2LTHFAFe complexation The titration of ligand (H2LTHF) solution at fixed concentration of iron and varying concentration of H2LTHF have been conducted. The titration spectra of iron nitrate upon increasing addition of H2LTHF are shown in Fig. 2. During the titration, the hypochromicity was
Fig. 2. The titration absorption spectra of Fe(NO3)34H2O (2.5 104 M) by H2LTHF.
observed without any shift in kmax band of 540 nm, which presents the existence of interaction between H2LTHF and the iron ion. The appearance of two isosbestic points in iron nitrate spectraclearly indicates the existence of simple equilibrium between H2LTHF and H2LTHF AFe complex. The complex composition of 3:2 was determined by plotting the absorption changes vs. ligand to metal mole ratio (nL/nFe). Fig. 3 shows a mole ratio plot of iron nitrate upon increasing addition of H2LTHF. 3.3. Magnetic susceptibility measurement Magnetic measurements were performed on a Quantum Design MPMS XL-5 SQUID magnetometer. The dc magnetic susceptibility
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Fig. 3. Determination of Fe2(LTHF)3 complex composition by the mole ratio plot. Fig. 5. Cyclic voltammograms of Fe2(LTHF)3 in CH2Cl2. The scan rate is 200 mV/s and T = 40 °C.
v = M/H was determined in the temperature range 2–300 K in a magnetic field H of 1000 Oe. The measured susceptibility monotonically increases with decreasing temperature as is shown in Fig. 4. In the inset of Fig. 4, we show a temperature dependence of effective magnetic moment calculated per binuclear unit Fe2(LTHF)3. At room temperature the effective magnetic moment leff is 7.7 lB which is less than the calculated (8.3 lB) for two uncoupled high-spin (S = 5/2) Fe(III) ions. The effective magnetic moment decreases upon cooling showing an antiferromagnetic coupling between Fe(III) ions in a unit. According to the binuclear structure of the complex we have described the measured susceptibility with the v(T) dependence that follows from the isotropic interaction Hamiltonian H = JSA SB [37], where SA = SB = 5/2 for high-spin Fe(III) ions. We add a temperature constant term to the function v(T) derived in Ref. [37] that accounts for diamagnetic contribution of inner shell electrons and temperature independent paramagnetism. The result of a fitting procedure is shown as a full line in inset to Fig. 4 giving us the interaction J = 0.3 cm1. The interaction J is negative as expected for an antiferromagnetic coupling. It is rather small compared to the interactions in similar bridged compounds collected in [38]. However, a relatively small interaction J is in agreement with our experimental data where no maximum in temperature dependent susceptibility, which is a fingerprint of strong antiferromagnetic coupling, has been observed.
3.4. Electrochemistry Cyclic voltammogram (CV) of complex Fe2(LTHF)3 have been recorded in CH2Cl2 solutions containing 0.1 M [(nBu)4N]ClO4 as a supporting electrolyte. The CV voltammograms (CV) of Fe2(LTHF)3 complex revealed oxidations and reductions peaks (Fig. 5). Complex Fe2(LTHF)3 exhibits two reversible oxidation processes at 0.840 V and 1.3 V, corresponds to ligand-centered peaks as seen in Fig. 5. The metal-centered voltammogram shows two quasireversible peaks in the negative potential range, which corresponds to the FeIIIFeIII/FeIIIFeII and FeIIIFeII/FeIIFeII reduction of Fe2(LTHF)3. The anodic to cathodic peak separation is 120 mV and increases with scan rates (Fig. 6), indicating that the process deviates from a kinetic reversibility and the electron transfer at the electrode surface is relatively slow. Fig. 7 shows the cyclic voltammograms of H2LTHF ligand and Fe2(LTHF)3 complex in positive scan region at 40 °C. In both voltammograms two anodic peaks appear; A1 and A2 for complex and Á1 and Á2 for ligand. These two anodic peaks are related to the oxidation of phenolic moieties of ligands (known as ligand-centered potentials) to their semiquinone and quinine species. There are two considerable differences between voltammograms of H2LTHF ligand and Fe2(LTHF)3 complex. (a) Both oxidation steps of Fe2(LTHF)3 complex are reversible at 40 °C with two well-defined
8 6
μeff (BM)
χ (emu/mol)
1.5
1.0
4 2
0.5
0.0
0
0
H = 1 kOe 0
100
100
T (K) 200
200
300
300
T (K) Fig. 4. Temperature dependent susceptibility and effective magnetic moment (inset) of Fe2(LTHF)3 in magnetic field of 1 kOe. The full line in inset is a fit with theoretical susceptibility of binuclear S = 5/2 units as described in text.
Fig. 6. Cyclic voltammograms of Fe2(LTHF)3 (3 103 M) in CH2Cl2 with M [(nBu)4N]ClO4 as supporting electrolyte. Potentials are referenced vs. Fc, Scan rates are 50, 100, 200, 300, 400, 500 and 600 mV/s and T = 40 °C.
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Fig. 7. Cyclic voltammograms of Fe2(LTHF)3 and ligand in CH2Cl2 The scan rate is 200 mV/s and T = 40 °C.
counter parts, C´1 and C´2 which are related to the reduction of semiquinone and quinone forms. It has seen that the cyclic voltammograms of ligand do not show any reduction peaks. The reversibility of cyclic voltammograms in Fe2(LTHF)3 comparing to its free ligand can attributed to the role of Lewis acidity of iron center on stabilization of phenoxyl radical bound to it, as seen in ribonucleotide reductase, RNR enzyme. (b) comparing the oxidation potential of Fe2(LTHF)3 with H2LTHF show a positive shift in the anodic peak potentials of complexes, 460 mV for A1–Á1 and 450 mV for A2–Á2. It may be related to stability of complex and lower density of negative charge on phenolate group which coordinated to iron (III) center. In order to investigate the effect of electronic and steric effects of substituents peak potentials of Fe2(LTHF)3 were compared with previously reported tert-butyl substituted complex Fe2(LtBu)2 [36]. The half wave potentials for both oxidation peaks (Fig. 8) of Fe2(LTHF)3 shift to more positive values comparing to Fe2(LtBu)2. This criterion can attribute to more highly positive charge density on iron center due to electron-withdrawing effect of chlorine substituent’s on phenolate moities in Fe2(LTHF)3. It can lead us to this fact that despite of the more steric effect around radical center in Fe2(LtBu)2 on radical stabilization, Fe2(LTHF)3 show a good phenoxyl radical stability. In other word, one can say that sometimes electronic parameters are more important than steric parameters for radical stabilization.
Fig. 9. Cyclic voltammograms of Fe2(LTHF)3 in CH2Cl2 at different temperatures. The scan rate is 200 mV/s.
3.5. Investigation of thermodynamic stability of Fe2(LTHF)3 radical complex We shed light into the stability constants and thermodynamic parameters of the electrochemically production of Fe2(LTHF)3 radical complex, Fe2 ðLTHF Þ3 through characterization of the energetic governing radical complex formation. In this way, the voltammetric study of complex was performed at different temperatures. There is a regular shift in the peak potentials of complex, A1. This peak shifts to more positive potentials by decreasing temperature (Fig. 9). The height of cathodic peak decreases by increasing the temperature which is in good accordance with instability of produced Fe2(LTHF)3 radical complex at higher temperatures. Based on the shift in peak potentials, temperature dependence formation constants were obtained by Nernst Equation as follows.
2RT ðLnK f Þ DE1=2 ¼ 3nF
ð2Þ
The energetic of Fe2(LTHF)3 equilibrium can be conveniently determined by thermodynamic parameters such as Gibbs free energy (DG). The standard Gibbs energy change is usually calculated from the equilibrium constant, K, of the electrochemical process, by the following relationship:
DG0 ¼ RTLnK f
ð3Þ
where R and T are the gas constant and the absolute temperature, respectively. Since the activity coefficients of the reactions are not known, the usual procedure is to assume them unity and to use the equilibrium concentrations instead of the activity. Therefore, it would be appropriate to adjust the terminology of apparent equilibrium constant K0 , and Gibbs energy DG°0 . Apparent standard enthalpies per mole of cooperative unit can be obtained from the dependence on temperature of the apparent formation constant K0 , by van’t Hoff equation:
0 DH 1 @ @LnK f ¼ T R
ð4Þ
This is the so-called van’t Hoff enthalpy. The apparent standard entropy change, DS°0 , can be derived from the following equation:
DS0 ¼ ðDH0 DG0 Þ=T Fig. 8. Cyclic voltammograms of Fe2(LTHF)3 and Fe2(LtBu)2 in CH2Cl2. The scan rate is 200 mV/s and T = 40 °C.
ð5Þ
The van’t Hoff plot for electrochemically produced radical complex of Fe2 ðLTHF Þ3 is shown in Fig. 10.
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Acknowledgments Authors are grateful to the Institute for Advanced Studies in Basic Sciences (IASBSs), University of Ljubljana and Università di Messina. E. Safaei gratefully acknowledges the support by the Institute for Advanced Studies in Basic Sciences (IASBS) Research Council under Grant No. G2012IASBS127. Appendix A The structural data for Fe2(LTHF)3 has been deposited with the Cambridge Crystallographic Data Centre, the deposition number being CCDC 832084. These data can be obtained free of charge via http://www.ccdc.cam.ac.uk/conts/retrieving.html, or from the Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336 033; or e-mail: [email protected]. Fig. 10. plot of LnKf vs. 1/T according to the van’t Hoff plot.
Appendix B. Supplementary material Table 3 Formation constants and related thermodynamic parameters of Fe2(LTHF)3. T/K
log K0 (M1)
DG°0 (kJ mol1)
DH°0 (kJ mol1)
DS°0 (J mol1 K1)
298 283 248 233
3.4 3.9 4.5 5.0
19.3 ± 0.310 21.2 ± 0.100 21.5 ± 0.020 22.4 ± 0.008
21.3 ± 0.6 21.3 ± 0.6 21.3 ± 0.6 21.3 ± 0.6
6.8 ± 0.0031 0.3 ± 0.0025 1.0 ± 0.0025 4.6 ± 0.0026
The calculated thermodynamic parameters are listed in Table 3. The negatively large standard free energy changes (DG°0 ) implies that Fe2 ðLTHF Þ3 product is favored over reactant. The more negative values of DG°0 at lower temperatures can be related to the higher stability of Fe2 ðLTHF Þ3 because of decreasing radical side reactions. It has also been indicated that formation of Fe2 ðLTHF Þ3 is an exothermic process. In terms of driving forces, the present case implies the stabilization of phenoxyl radicals (LTHF) due to their interaction with iron centers. The considerable negative values of DH°0 and also small positive (or negative) values of DS°0 in mentioned electrochemical process indicate that the contribution of the negative enthalpy changes (DH°0 ) results in more negative DG°0 and favoring the radical complex production process (Table 3), it seems that the major contributing factor in radical stabilizing is enthalpic in origin. It can be concluded that the negative enthalpy changes are the driving forces in interaction between iron center and ligand phenoxyl radical (LTHF). 4. Conclusions A kind of new [NAO]-donor tripodal amine chloro-substituted phenol ligand and its iron(III) complex of the type Fe2 LTHF 3 , have been synthesized and characterized. X-ray crystal structure of Fe2 LTHF reveal that this complex has unsymmetrical dimeric struc3 ture, in which two Fe(III) centers are surrounded by amine nitrogen’s, phenolate atoms of three ligands. Magnetostructural studies of this complex displayed an antiferromagnetic coupling between Fe(III) ions through phenolate bridges. It has been shown that electrochemical oxidations are ligand-centered, i.e. formation of the phenoxyl radicals from the coordinated phenolates, while reductions are metal-centered related to FeIII/FeII couple. Our electrochemically generated Fe2(LTHF)3 phenoxyl radical complex (Fe2 ðLTHF Þ3 ) shows a good stability and this process is exothermic due to the predominant role of phenoxyl radical (LTHF)/iron interaction. This means that metal ions can be good candidates as radical stabilizers as seen in enzymes for application in new molecular switches and/or data storage devices [11].
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