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Heme ferrous–hydroperoxo complexes: some theoretical considerations
ABB Archives of Biochemistry and Biophysics 424 (2004) 137–140 www.elsevier.com/locate/yabbi
Heme ferrous–hydroperoxo complexes: some theoretical considerations Radu Silaghi-Dumitrescu* Department of Chemistry and Center for Metalloenzyme Studies, University of Georgia, Athens, GA 30602, USA Department of Chemistry, Babes-Bolyai University, Cluj-Napoca RO-3400, Romania Received 17 January 2004, and in revised form 17 February 2004
Abstract We report density functional calculations on complexes of ferrous hemes with hydroperoxide, where the axial ligand trans to OOH is imidazole, thiolate, or phenoxide. The geometrical parameters and charge distributions within the Fe–O–O–H moiety are identical between the ferrous complexes reported here and their ferric counterparts previously described, even though the latter contain one unpaired electron on iron as opposed to the former, which are diamagnetic. The extra negative charge upon going from a formally ferric state to formally ferrous appears to be distributed essentially on the porphyrin. These findings support recent experimental data showing that the ferrous state of certain hemoproteins can interact with peroxides in a catalytically competent fashion, cleaving the O–O bond heterolytically in a manner reminiscent of the ‘‘canonical’’ ferric–peroxo complexes, and contrary to any expectations based on the Fenton concept commonly invoked in non-heme chemistry. Ó 2004 Elsevier Inc. All rights reserved. Keywords: Heme; DFT; Peroxide; Peroxidase; Catalase; Alkylhydroperoxide reductase; Hemoprotein; Fenton; Free radical
Ferric heme active sites of various proteins are known to react readily with peroxides, in a process that results in immediate heterolytic cleavage of the O–O bond within a ferric–hydroperoxo (or ferric–alkylperoxo) complex [1–3]. This reaction is accepted to define the physiological function of peroxidases [2,4] and catalases [3,5], while for other hemoproteins (e.g., cytochromes P450) [1], this reaction has been observed in vitro but is generally assumed to have little physiological relevance. The product of O–O bond cleavage is a species denoted Compound I, currently described as iron(IV)–oxo þ porphyrin cation radical [6]. One electron reduction of Compound I further yields Compound II, formally an iron(IV)–oxo species. One electron reduction of Compound II then yields the resting ferric state of the heme [1,4]. By contrast, the ferrous form of the above-mentioned active sites was proposed to be less reactive with peroxides, and this reactivity was thought to involve liber* Fax: 1-706-542-9454. E-mail address: [email protected].
0003-9861/$ - see front matter Ó 2004 Elsevier Inc. All rights reserved. doi:10.1016/j.abb.2004.02.017
ation of hydroxyl radicals via homolytic cleavage of the O–O bond in Fe(II)–O–OH complexes—the so-called Fenton chemistry [7,8]. Protein damage has been reported to occur upon reaction of the ferrous form of hemoglobin with hydrogen peroxide and this damage has been associated with hydroxyl radicals [7]. It is still unclear as to what extent such protein damage can be attributed to hydroxyl radical-independent pathways— such as heterolytic cleavage of the O–O bond within a ferrous–hydroperoxo complex (yielding the high-valent intermediate Compound II). On the other hand, ferrous horseradish peroxidase was shown to induce heterolytic cleavage of the O–O bond in H2 O2 [9]. A similar (heterolytic) behavior was reported for myeloperoxidase [10]. Likewise, under anaerobic conditions, ferrous flavohemoglobin (the active site of which features a histidine-ligated heme similar to that of canonical peroxidases) efficiently catalyzes heterolytic cleavage of the O–O bond in alkylperoxides, forming alcohols and Compound II [11]. With these previous results in mind, we have investigated the origin of the unexpected non-Fenton reactivity of these ferrous hemes towards
138
R. Silaghi-Dumitrescu / Archives of Biochemistry and Biophysics 424 (2004) 137–140
peroxides. We report here density functional calculations on ferrous–hydroperoxo heme models, where the axial ligand trans to hydroperoxide is either imidazole (as seen in myeloperoxidase, flavohemoglobin, and most heme peroxidases), thiolate (as seen in chloroperoxidases) or phenoxide (as seen in catalases). We conclude that the Fe–O–OH moiety in ferrous–hydroperoxo complexes is essentially identical to ferric–hydroperoxo complexes. Physiological implications of this conclusion are discussed.
Materials and methods Geometries for all models (cf. Fig. 1) were optimized without constraints (except where specified) at the DFT level in the Spartan package [12]. The BP86 functional, which uses the gradient-corrected exchange functional proposed by Becke [13], the correlation functional by Perdew [14], and the DN** numerical basis set (comparable in size to 6-31G**) were used as implemented in Spartan. For the SCF calculations, a fine grid was used, and the convergence criteria were set to 106 (for the root-mean square of electron density) and 108 (energy), respectively. For geometry optimization, convergence criteria were set to 0.001 a.u. (maximum gradient criterion) and 0.0003 (maximum displacement criterion). Charges were derived from Mulliken population analyses after DFT geometry optimization. Lowdin bond orders were derived from similar BP86/6-31G** calculations performed within the Spartan 02 package [15].
Results and discussion Fig. 1 shows the ferrous– and ferric–hydroperoxo models employed in the present study, and Table 1 shows geometry optimization results for these models. Table 1 shows only negligible differences in O–O and Fe–OOH bond distances between Fe(II) and Fe(III)
models, with the Fe(II) models featuring slightly longer distances as expected. Lowdin and Mulliken bond orders (not shown) for the O–O bond are also identical between ferrous and ferric models. Further arguing for a similarity in bond strength between Fe(II)–OOH and Fe(III)–OOH, we found that, in histidine-ligated mod required els, elongation of the O–O bond by 0.2 A 6 kcal/mol for both Fe(II) and Fe(III) models, while similarly required 15 kcal/mol for elongation by 0.4 A both ferrous and ferric. No significant spin density developed on the OH group upon elongating the O–O bond for either ferrous or ferric models. Rather, a progressive increase of negative charge (0.1–0.5 units) on the OH is seen in both cases. A full comparative investigation of the O–O bond cleavage process in ferrous– and ferric–hydroperoxo complexes would have to at least take into account a potential proton donor for the ‘‘leaving’’ OH group, as well as medium polarization effects. Ideally, the entire protein scaffold would have to be included, leading to a QM/MM approach. However, at this early stage our data do indicate a strong resemblance between the O–O bonds in ferrous and ferric heme hydroperoxo complexes. This resemblance includes identical bond lengths, identical calculated bond orders (with either the Lowdin or the Mulliken analysis) and identical resistance to O–O bond elongation. By contrast, the O–O bond in most non-heme ferrous–hydroperoxo complexes that we have examined (unpublished results), was much weaker than in heme ferrous– and ferric–hydroperoxo models examined here. This O–O bond in non-heme Fe(II)–OOH complexes in fact cleaves simply upon geometry optimization—as expected based on common knowledge of Fenton chemistry [16]. Partial atomic charges on iron are essentially identical between Fe(II) and Fe(III) models (cf. Table 1). Spin densities for the ferric models were shown elsewhere [17] to be entirely consistent with a low-spin ferric state. As expected, no unpaired spins are found on the iron in the ferrous state. The Fe–X (non-peroxo axial ligand) distances are identical between Fe(II) and Fe(III) in the
Fig. 1. Models examined in the present work (ferrous and ferric, for each structure).
R. Silaghi-Dumitrescu / Archives of Biochemistry and Biophysics 424 (2004) 137–140
139
Table 1 and partial atomic charges in ferrous–hydroperoxo and ferric–hydroperoxo heme models shown in Fig. 1 Relevant bond distances (A) Model
O–O
Fe–O
Fe–X
Fe
O1a
O2 b
O2 Hc
Xd
Pe
SCH3 –Fe(III)–OOH SCH3 –Fe(II)–OOH PhO–Fe(III)–OH PhO–Fe(II)–OH ImH–Fe(III)–OOH ImH–Fe(II)–OOH
1.46 1.49 1.48 1.49 1.46 1.47
1.89 1.93 1.84 1.89 1.80 1.88
2.31 2.41 1.96 2.08 2.07 2.09
0.53 0.57 0.69 0.68 0.65 0.63
)0.29 )0.36 )0.30 )0.35 )0.25 )0.35
)0.28 )0.34 )0.29 )0.34 )0.23 )0.37
)0.26 )0.40 )0.31 )0.39 )0.20 )0.37
)0.32 )0.40 )0.48 )0.68 0.37 0.18
)0.92 )1.77 )0.90 )1.61 )0.72 )1.44
a
Iron-bound oxygen atom. Protonated oxygen atom. c Sum over the hydroperoxo ligand. d Sum over all axial ligand atoms (i.e., thiolate, phenoxide, and imidazole). e Sum over all porphyrin atoms. b
for the imidazole-ligated models, but differ by 0.1 A anionic thiolate and phenoxide ligands, not unexpectedly. Only minimal differences are seen in the overall charge of the hydroperoxo ligand, between Fe(II) and Fe(III) models. Noticeable differences (0.7–0.9 units) are however seen in porphyrin charges. Thus, the extra electron in Fe(II) models compared to their Fe(III) counterparts is localized mainly on the porphyrin ring. Others [18] and we [17] have already shown that a similar situation, with the porphyrin serving as a ‘‘charge buffer,’’ occurs when comparing ferrous vs. ferric heme complexes of various axial ligation (e.g., bis-histidine, histidine–dioxygen, histidine–oxo, histidine–hydroxo, thiolate–dioxygen, thiolate–oxo, thiolate–hydroxo, and others). We have thus shown that, beyond any concern regarding the formal oxidation state of the iron, the Fe–O–O–H moiety in heme ferrous–hydroperoxo complexes appears to be essentially identical to that in ferric–hydroperoxo complexes. It is then not surprising that, at least in certain hemoproteins [10,11], heterolytic cleavage of the O–O bond in peroxides has been shown experimentally to be catalyzed by ferrous and ferric hemes alike. Peroxidases and catalases feature redox potentials between +50 and )300 mV, a range that more or less parallels the varying in vivo redox potentials that these hemoproteins will encounter (from strongly reducing anaerobic bacteria to human phagocytes) [19–22]. This implies that non-negligible amounts of these hemoproteins will always exist in the ferrous state in vivo. At such a point, expectations based on the classical Fenton concept would lead one to believe that encounters between the hemoprotein and its substrate(s) would be fatal for the protein if not for the entire cell. Our data, however, suggest that such encounters between ferrous hemoproteins and peroxides are in fact catalytically competent and hold physiological relevance. Experimental results [9–11] support this prediction. Peroxides may face competition from O2 , NO or CO when reacting with ferrous hemoproteins in vivo; how-
ever, at least one case is clear: flavohemoglobin features Km values for alkylhydroperoxides sufficiently low to justify its recent description as alkylhydroperoxide reductase, with a ferrous resting state [11]. For hemoproteins involved in oxidative stress defense [23], it now appears that a ferric or a ferrous resting state would serve the purpose of removing toxic hydroperoxides equally well. On the other hand, hemoproteins involved in metabolic pathways will often specifically require formation of Compound I (starting from a ferric resting state), rather than just Compound II (starting from a ferrous resting state), as Compound I is a much stronger oxidant than Compound II [2,4,21,24,25]. Also, hydrogen peroxide has been shown to react somewhat slower with ferrous, compared with ferric, hemes [10]. This difference was ascribed to a contraction of the active site upon reduction (imposing extra steric constraints for peroxide binding) [26] as well as to electrostatic repulsion between HOO and the extra electron present on the heme in the ferrous state [10]. Notably, O–O bond cleavage in the ferrous–hydroperoxo adduct appears to be very efficient, in that (similar to ferric–hydroperoxo adducts) it is much faster than initial binding of peroxide to the heme: an Fe(II)–OOH reaction intermediate prior to Compound II formation could not be detected in any of the stopped-flow experiments reported so far [9–11]. To conclude, we add that our data on heme models appears to support the ‘‘no-radical’’ interpretation of Fenton (ferrous–peroxo) chemistry [27]. Efforts are under way to similarly investigate non-heme ferrous–peroxo reactivity.
Acknowledgments Dr. D.M. Kurtz, Jr. (University of Georgia) is thanked for support and helpful discussions. Dr. I. Silaghi-Dumitrescu (Babes-Bolyai University) and Dr. V. Huang (University of Georgia) are thanked for helpful discussions.
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R. Silaghi-Dumitrescu / Archives of Biochemistry and Biophysics 424 (2004) 137–140
References [1] M. Sono, M.P. Roach, E.D. Coulter, J.H. Dawson, Chem. Rev. 96 (1996) 2841. [2] I.S. Isaac, J.H. Dawson, Essays Biochem. 34 (1999) 51. [3] T.L. Poulos, Adv. Inorg. Chem. 7 (1988) 1. [4] H.B. Dunford, in: J. Everse, K.E. Everse, M.B. Grisham (Eds.), Peroxidases in Chemistry and Biology, vol. II, CRC Press, Boca Raton, 1991, pp. 2–17. [5] O.M. Lardinois, Free Radic. Res. 22 (1995) 251. [6] D.L. Harris, Curr. Opin. Chem. Biol. 5 (2001) 724. [7] S.M. Sadrzadeh, E. Graf, S.S. Panter, P.E. Hallaway, J.W. Eaton, J. Biol. Chem. 259 (1984) 14354. [8] E. Nagababu, J.M. Rifkind, Biochem. Biophys. Res. Commun. 247 (1998) 592. [9] R.W. Noble, Q.H. Gibson, J. Biol. Chem. 245 (1970) 2409. [10] W. Jantschko, P. Furtmuller, M. Zederbauer, M. Lanz, C. Jakopitsch, C. Obinger, Biochem. Biophys. Res. Commun. 312 (2003) 292. [11] A. Bonamore, P. Gentili, A. Ilari, M.E. Schinina, A. Boffi, J. Biol. Chem. 278 (2003) 22272. [12] Spartan 5.0, Wavefunction, Inc., 18401 Von Karman Avenue Suite 370, Irvine, CA 92612, USA.
[13] A.D. Becke, Phys. Rev. (1988) 3098. [14] J.P. Perdew, Phys. Rev. B 33 (1986) 8822. [15] SPARTAN Õ02 for Windows, Wavefunction Inc., 18401 Von Karman Avenue, Suite 370 Irvine, CA 92612, USA. [16] Fenton chemistry is commonly associated with non-heme chemistry. [17] R. Silaghi-Dumitrescu, I. Silaghi-Dumitrescu, Rev. Roum. Chim. (2003) in press. [18] M.P. Johansson, M.R.A. Blomberg, D. Sundholm, M. Wikstrom, Biochim. Biophys. Acta 1553 (2002) 183. [19] R. Makino, R. Chiang, L.P. Hager, Biochemistry 15 (1976) 4748. [20] C.D. Millis, D. Cai, M.T. Stankovich, T. Ming, Biochemistry 28 (1989) 8484. [21] J. Arnhold, P. Furtmuller, C. Obinger, Redox Rep. 8 (2003) 179. [22] N.L. Wengenack, H. Lopes, M.J. Kenndey, P. Tavares, A.S. Pereira, I. Moura, J.J.G. Moura, F. Rusnak, Biochemistry 39 (2000) 11508. [23] E.D. Coulter, D.M. Kurtz Jr., Chemtracts Inorg. Chem. 14 (2001) 407. [24] A.M. English, G. Tsaprailis, Adv. Inorg. Chem. 43 (1995) 79. [25] M. Gajhede, Biochem. Soc. Trans. 29 (2001) 91. [26] H.M. Abu-Soud, S.L. Hazen, Biochemistry 40 (2001) 10747. [27] H.B. Dunford, Coord. Chem. Rev. 233–234 (2002) 311.
Heme ferrous–hydroperoxo complexes: some theoretical considerations Radu Silaghi-Dumitrescu* Department of Chemistry and Center for Metalloenzyme Studies, University of Georgia, Athens, GA 30602, USA Department of Chemistry, Babes-Bolyai University, Cluj-Napoca RO-3400, Romania Received 17 January 2004, and in revised form 17 February 2004
Abstract We report density functional calculations on complexes of ferrous hemes with hydroperoxide, where the axial ligand trans to OOH is imidazole, thiolate, or phenoxide. The geometrical parameters and charge distributions within the Fe–O–O–H moiety are identical between the ferrous complexes reported here and their ferric counterparts previously described, even though the latter contain one unpaired electron on iron as opposed to the former, which are diamagnetic. The extra negative charge upon going from a formally ferric state to formally ferrous appears to be distributed essentially on the porphyrin. These findings support recent experimental data showing that the ferrous state of certain hemoproteins can interact with peroxides in a catalytically competent fashion, cleaving the O–O bond heterolytically in a manner reminiscent of the ‘‘canonical’’ ferric–peroxo complexes, and contrary to any expectations based on the Fenton concept commonly invoked in non-heme chemistry. Ó 2004 Elsevier Inc. All rights reserved. Keywords: Heme; DFT; Peroxide; Peroxidase; Catalase; Alkylhydroperoxide reductase; Hemoprotein; Fenton; Free radical
Ferric heme active sites of various proteins are known to react readily with peroxides, in a process that results in immediate heterolytic cleavage of the O–O bond within a ferric–hydroperoxo (or ferric–alkylperoxo) complex [1–3]. This reaction is accepted to define the physiological function of peroxidases [2,4] and catalases [3,5], while for other hemoproteins (e.g., cytochromes P450) [1], this reaction has been observed in vitro but is generally assumed to have little physiological relevance. The product of O–O bond cleavage is a species denoted Compound I, currently described as iron(IV)–oxo þ porphyrin cation radical [6]. One electron reduction of Compound I further yields Compound II, formally an iron(IV)–oxo species. One electron reduction of Compound II then yields the resting ferric state of the heme [1,4]. By contrast, the ferrous form of the above-mentioned active sites was proposed to be less reactive with peroxides, and this reactivity was thought to involve liber* Fax: 1-706-542-9454. E-mail address: [email protected].
0003-9861/$ - see front matter Ó 2004 Elsevier Inc. All rights reserved. doi:10.1016/j.abb.2004.02.017
ation of hydroxyl radicals via homolytic cleavage of the O–O bond in Fe(II)–O–OH complexes—the so-called Fenton chemistry [7,8]. Protein damage has been reported to occur upon reaction of the ferrous form of hemoglobin with hydrogen peroxide and this damage has been associated with hydroxyl radicals [7]. It is still unclear as to what extent such protein damage can be attributed to hydroxyl radical-independent pathways— such as heterolytic cleavage of the O–O bond within a ferrous–hydroperoxo complex (yielding the high-valent intermediate Compound II). On the other hand, ferrous horseradish peroxidase was shown to induce heterolytic cleavage of the O–O bond in H2 O2 [9]. A similar (heterolytic) behavior was reported for myeloperoxidase [10]. Likewise, under anaerobic conditions, ferrous flavohemoglobin (the active site of which features a histidine-ligated heme similar to that of canonical peroxidases) efficiently catalyzes heterolytic cleavage of the O–O bond in alkylperoxides, forming alcohols and Compound II [11]. With these previous results in mind, we have investigated the origin of the unexpected non-Fenton reactivity of these ferrous hemes towards
138
R. Silaghi-Dumitrescu / Archives of Biochemistry and Biophysics 424 (2004) 137–140
peroxides. We report here density functional calculations on ferrous–hydroperoxo heme models, where the axial ligand trans to hydroperoxide is either imidazole (as seen in myeloperoxidase, flavohemoglobin, and most heme peroxidases), thiolate (as seen in chloroperoxidases) or phenoxide (as seen in catalases). We conclude that the Fe–O–OH moiety in ferrous–hydroperoxo complexes is essentially identical to ferric–hydroperoxo complexes. Physiological implications of this conclusion are discussed.
Materials and methods Geometries for all models (cf. Fig. 1) were optimized without constraints (except where specified) at the DFT level in the Spartan package [12]. The BP86 functional, which uses the gradient-corrected exchange functional proposed by Becke [13], the correlation functional by Perdew [14], and the DN** numerical basis set (comparable in size to 6-31G**) were used as implemented in Spartan. For the SCF calculations, a fine grid was used, and the convergence criteria were set to 106 (for the root-mean square of electron density) and 108 (energy), respectively. For geometry optimization, convergence criteria were set to 0.001 a.u. (maximum gradient criterion) and 0.0003 (maximum displacement criterion). Charges were derived from Mulliken population analyses after DFT geometry optimization. Lowdin bond orders were derived from similar BP86/6-31G** calculations performed within the Spartan 02 package [15].
Results and discussion Fig. 1 shows the ferrous– and ferric–hydroperoxo models employed in the present study, and Table 1 shows geometry optimization results for these models. Table 1 shows only negligible differences in O–O and Fe–OOH bond distances between Fe(II) and Fe(III)
models, with the Fe(II) models featuring slightly longer distances as expected. Lowdin and Mulliken bond orders (not shown) for the O–O bond are also identical between ferrous and ferric models. Further arguing for a similarity in bond strength between Fe(II)–OOH and Fe(III)–OOH, we found that, in histidine-ligated mod required els, elongation of the O–O bond by 0.2 A 6 kcal/mol for both Fe(II) and Fe(III) models, while similarly required 15 kcal/mol for elongation by 0.4 A both ferrous and ferric. No significant spin density developed on the OH group upon elongating the O–O bond for either ferrous or ferric models. Rather, a progressive increase of negative charge (0.1–0.5 units) on the OH is seen in both cases. A full comparative investigation of the O–O bond cleavage process in ferrous– and ferric–hydroperoxo complexes would have to at least take into account a potential proton donor for the ‘‘leaving’’ OH group, as well as medium polarization effects. Ideally, the entire protein scaffold would have to be included, leading to a QM/MM approach. However, at this early stage our data do indicate a strong resemblance between the O–O bonds in ferrous and ferric heme hydroperoxo complexes. This resemblance includes identical bond lengths, identical calculated bond orders (with either the Lowdin or the Mulliken analysis) and identical resistance to O–O bond elongation. By contrast, the O–O bond in most non-heme ferrous–hydroperoxo complexes that we have examined (unpublished results), was much weaker than in heme ferrous– and ferric–hydroperoxo models examined here. This O–O bond in non-heme Fe(II)–OOH complexes in fact cleaves simply upon geometry optimization—as expected based on common knowledge of Fenton chemistry [16]. Partial atomic charges on iron are essentially identical between Fe(II) and Fe(III) models (cf. Table 1). Spin densities for the ferric models were shown elsewhere [17] to be entirely consistent with a low-spin ferric state. As expected, no unpaired spins are found on the iron in the ferrous state. The Fe–X (non-peroxo axial ligand) distances are identical between Fe(II) and Fe(III) in the
Fig. 1. Models examined in the present work (ferrous and ferric, for each structure).
R. Silaghi-Dumitrescu / Archives of Biochemistry and Biophysics 424 (2004) 137–140
139
Table 1 and partial atomic charges in ferrous–hydroperoxo and ferric–hydroperoxo heme models shown in Fig. 1 Relevant bond distances (A) Model
O–O
Fe–O
Fe–X
Fe
O1a
O2 b
O2 Hc
Xd
Pe
SCH3 –Fe(III)–OOH SCH3 –Fe(II)–OOH PhO–Fe(III)–OH PhO–Fe(II)–OH ImH–Fe(III)–OOH ImH–Fe(II)–OOH
1.46 1.49 1.48 1.49 1.46 1.47
1.89 1.93 1.84 1.89 1.80 1.88
2.31 2.41 1.96 2.08 2.07 2.09
0.53 0.57 0.69 0.68 0.65 0.63
)0.29 )0.36 )0.30 )0.35 )0.25 )0.35
)0.28 )0.34 )0.29 )0.34 )0.23 )0.37
)0.26 )0.40 )0.31 )0.39 )0.20 )0.37
)0.32 )0.40 )0.48 )0.68 0.37 0.18
)0.92 )1.77 )0.90 )1.61 )0.72 )1.44
a
Iron-bound oxygen atom. Protonated oxygen atom. c Sum over the hydroperoxo ligand. d Sum over all axial ligand atoms (i.e., thiolate, phenoxide, and imidazole). e Sum over all porphyrin atoms. b
for the imidazole-ligated models, but differ by 0.1 A anionic thiolate and phenoxide ligands, not unexpectedly. Only minimal differences are seen in the overall charge of the hydroperoxo ligand, between Fe(II) and Fe(III) models. Noticeable differences (0.7–0.9 units) are however seen in porphyrin charges. Thus, the extra electron in Fe(II) models compared to their Fe(III) counterparts is localized mainly on the porphyrin ring. Others [18] and we [17] have already shown that a similar situation, with the porphyrin serving as a ‘‘charge buffer,’’ occurs when comparing ferrous vs. ferric heme complexes of various axial ligation (e.g., bis-histidine, histidine–dioxygen, histidine–oxo, histidine–hydroxo, thiolate–dioxygen, thiolate–oxo, thiolate–hydroxo, and others). We have thus shown that, beyond any concern regarding the formal oxidation state of the iron, the Fe–O–O–H moiety in heme ferrous–hydroperoxo complexes appears to be essentially identical to that in ferric–hydroperoxo complexes. It is then not surprising that, at least in certain hemoproteins [10,11], heterolytic cleavage of the O–O bond in peroxides has been shown experimentally to be catalyzed by ferrous and ferric hemes alike. Peroxidases and catalases feature redox potentials between +50 and )300 mV, a range that more or less parallels the varying in vivo redox potentials that these hemoproteins will encounter (from strongly reducing anaerobic bacteria to human phagocytes) [19–22]. This implies that non-negligible amounts of these hemoproteins will always exist in the ferrous state in vivo. At such a point, expectations based on the classical Fenton concept would lead one to believe that encounters between the hemoprotein and its substrate(s) would be fatal for the protein if not for the entire cell. Our data, however, suggest that such encounters between ferrous hemoproteins and peroxides are in fact catalytically competent and hold physiological relevance. Experimental results [9–11] support this prediction. Peroxides may face competition from O2 , NO or CO when reacting with ferrous hemoproteins in vivo; how-
ever, at least one case is clear: flavohemoglobin features Km values for alkylhydroperoxides sufficiently low to justify its recent description as alkylhydroperoxide reductase, with a ferrous resting state [11]. For hemoproteins involved in oxidative stress defense [23], it now appears that a ferric or a ferrous resting state would serve the purpose of removing toxic hydroperoxides equally well. On the other hand, hemoproteins involved in metabolic pathways will often specifically require formation of Compound I (starting from a ferric resting state), rather than just Compound II (starting from a ferrous resting state), as Compound I is a much stronger oxidant than Compound II [2,4,21,24,25]. Also, hydrogen peroxide has been shown to react somewhat slower with ferrous, compared with ferric, hemes [10]. This difference was ascribed to a contraction of the active site upon reduction (imposing extra steric constraints for peroxide binding) [26] as well as to electrostatic repulsion between HOO and the extra electron present on the heme in the ferrous state [10]. Notably, O–O bond cleavage in the ferrous–hydroperoxo adduct appears to be very efficient, in that (similar to ferric–hydroperoxo adducts) it is much faster than initial binding of peroxide to the heme: an Fe(II)–OOH reaction intermediate prior to Compound II formation could not be detected in any of the stopped-flow experiments reported so far [9–11]. To conclude, we add that our data on heme models appears to support the ‘‘no-radical’’ interpretation of Fenton (ferrous–peroxo) chemistry [27]. Efforts are under way to similarly investigate non-heme ferrous–peroxo reactivity.
Acknowledgments Dr. D.M. Kurtz, Jr. (University of Georgia) is thanked for support and helpful discussions. Dr. I. Silaghi-Dumitrescu (Babes-Bolyai University) and Dr. V. Huang (University of Georgia) are thanked for helpful discussions.
140
R. Silaghi-Dumitrescu / Archives of Biochemistry and Biophysics 424 (2004) 137–140
References [1] M. Sono, M.P. Roach, E.D. Coulter, J.H. Dawson, Chem. Rev. 96 (1996) 2841. [2] I.S. Isaac, J.H. Dawson, Essays Biochem. 34 (1999) 51. [3] T.L. Poulos, Adv. Inorg. Chem. 7 (1988) 1. [4] H.B. Dunford, in: J. Everse, K.E. Everse, M.B. Grisham (Eds.), Peroxidases in Chemistry and Biology, vol. II, CRC Press, Boca Raton, 1991, pp. 2–17. [5] O.M. Lardinois, Free Radic. Res. 22 (1995) 251. [6] D.L. Harris, Curr. Opin. Chem. Biol. 5 (2001) 724. [7] S.M. Sadrzadeh, E. Graf, S.S. Panter, P.E. Hallaway, J.W. Eaton, J. Biol. Chem. 259 (1984) 14354. [8] E. Nagababu, J.M. Rifkind, Biochem. Biophys. Res. Commun. 247 (1998) 592. [9] R.W. Noble, Q.H. Gibson, J. Biol. Chem. 245 (1970) 2409. [10] W. Jantschko, P. Furtmuller, M. Zederbauer, M. Lanz, C. Jakopitsch, C. Obinger, Biochem. Biophys. Res. Commun. 312 (2003) 292. [11] A. Bonamore, P. Gentili, A. Ilari, M.E. Schinina, A. Boffi, J. Biol. Chem. 278 (2003) 22272. [12] Spartan 5.0, Wavefunction, Inc., 18401 Von Karman Avenue Suite 370, Irvine, CA 92612, USA.
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