Uptake and Release of Iron by Ferritin Adsorbed at Tin-Doped Indium

Department of Chemistry, University of Colorado, Denver, Colorado 80217. Received April 7, 1999. In this work, the uptake and release of iron by ferri...
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Uptake and Release of Iron by Ferritin Adsorbed at Tin-Doped Indium Oxide Electrodes Moon-Son Pyon, Ryan J. Cherry, A. Jason Bjornsen, and Donald C. Zapien* Department of Chemistry, University of Colorado, Denver, Colorado 80217 Received April 7, 1999 In this work, the uptake and release of iron by ferritin using direct electrochemical techniques have been investigated for the first time. Adsorption of ferritin from phosphate solution onto tin-doped indium oxide (ITO) at open circuit potential gives about one monolayer of ferritin. The ferritin layer is electroactive and therefore lends itself to electrochemical analysis; cyclic voltammetry was the principal method used in this study. In the presence of EDTA, iron is not removed at open circuit potential. However, when the ITO/ ferritin electrode is subjected to -0.70 V, the anodic branch is no longer present in the current-potential curve, indicating that the reduction of core iron had induced the removal of iron from the protein shell. Adsorbed ferritin, emptied in this fashion, was exposed to ferrous ion at 0.20 V. The ensuing currentpotential curve showed the original peak pattern, suggesting that iron had been reincorporated into the apoferritin shell. An authentic sample of adsorbed apoferritin exposed to ferrous ion, exhibited the same current-potential response as electrochemically emptied ferritin, supporting the conclusion that apoferritin results from the reduction of ITO/ferritin in the presence of an iron chelator. These studies show that ferritin adsorbed at an ITO electrode is a promising venue to study not only ferritin’s electrochemistry but also its functions.

Introduction Cellular ferritin is a protein whose principal functions are to sequester excess iron, to store this iron, and to supply iron when it is needed, such as in the synthesis of ironcontaining proteins. Ferritin is roughly spherical (12.0 nm diameter) protein which has a molecular weight of 450 000 and has the capacity to contain as many as 4500 iron atoms.1 The protein shell is composed of 24 subunits, associated in 4:3:2 symmetry, surrounding a mineral core whose average composition is (FeOOH)8‚FeOPO3H2. The arrangement of the subunits results in the formation of eight hydrophilic and six hydrophobic channels, which link the exterior with the interior of the shell.1 Since the earliest in vitro studies of iron release, it has been known that iron chelators can remove iron only at extremely slow rates, attesting to the stability of the ferritin complex. However, when treated with reducing agents, ferritin readily releases iron at approximately 1 × 106 times the rate when no reduction was performed.2 It was once believed that small reductants, such as dithionite and thioglycolic acid, might enter the protein shell and effect the iron release. However, large reductants such as FMNH2,3 xanthine,4 and NADH5 have been used to reduce ferritin, molecules too large to pass through the channels. Tunneling has been suggested as a possible mechanism.6 Alternatively, the electrons might pass through the shell mediated by iron ions residing in the * To whom correspondence should be addressed. (1) (a) Thiel, E. C. In Advances in Inorganic Biochemistry; Thiel, E. C., Eichhorn, G. L., Marzilli, L. G., Eds.; Elsevier: New York, 1983, Vol. 5, pp 1-38. (b) Thiel, E. C. Annu. Rev. Biochem. 1987, 56, 289-315. (2) Harrison, P. M.; Clegg, G. A.; May, K. Ferritin Structure and Function. In Iron in Biochemistry and Medicine II; Jacobs, A., Worwood, M., Eds.; Academic Press: London, 1980; pp 131-172. (3) Ulvik, R. J.; Romslo, I.; Roland, F.; Crichton, R. R. Biochim. Biophys. Acta 1981, 351, 224-229. (4) Topham, R.; Goger, M.; Pearce, K.; Schultz, P. Biochem. J. 1989, 261, 137-143. (5) Sirivech, S.; Frieden, E.; Osaki, S. Biochem. J. 1974, 143, 311315. (6) Watt, G. D.; Jacobs, D.; Frankel, R. B. Proc. Natl. Acad. Sci. U.S.A. 1988, 85, 7457-7461.

channels.7 Factors such as pH, core phosphate content, and buffer ion are all known to affect the rates of iron release.8 The mechanistic aspects of iron uptake have been studied more extensively. It appears that iron enters the protein sphere via the hydrophilic 3-fold channels and is chelated in regions rich in carboxylate-terminated residues at subunit trimer interfaces.9 Recent studies involving recombinant H- and L-chain ferritin have shown that the H-chain homopolymer was able to catalyze the oxidation of iron, whereas L-chain did not, leading to the hypothesis that the former possessed “ferroxidase activity”.10 One of the leading proposed mechanisms suggests that two Fe2+ are chelated to glutamate residues in the ferroxidase center in close proximity to each other.11 The two iron atoms then hydrolyze water forming an Fe2+-OH-Fe2+ dimer. Due to their close proximity, the iron(II) ions simultaneously transfer two electrons to molecular oxygen, generating the Fe3+-O-Fe3+ and hydrogen peroxide. The proposed mechanism is supported by the observation that peroxide was produced in the early stages of core develop(7) Ford, G. C.; Harrison, P. M.; Rice, D. W.; Smith, J. M. A.; Treffery, A.; White, J. L.; Yariv, J. J. Philos. Trans. R. Soc. London, Sect. B 1984, 304, 551-565. (8) (a) Watt, G. D.; Frankel, R. B.; Papaefthymiou, G. C. Proc. Natl. Acad. Sci. U.S.A. 1985, 82, 3640-43. (b) Watt, G. D.; Frankel, R. B.; Jacobs, D.; Huang, H.; Papaefthymiou, G. C. Biochemistry 1992, 31, 5672-5679. (9) (a) Rice, D. W.; Ford, G. C.; White, J. L.; Smith, J. M. A.; Harrison, P. M. Adv. Inorg. Biochem. 1983, 5, 39-50. (b) Harrison, P. M.; Lilley, T. H. Physical Bio-Inorganic Chemistry; Loher, T. M., Ed.; VCH: New York, 1989; Vol. 5, pp 123-238. (10) (a) Levi, S.; Luzzago, A.; Cesareni, G.; Cozzi, A.; Franceschinelli, F.; Albertini, A.; Arosio, P. J. Biol. Chem. 1988, 263, 18086-18092. (b) Levi, S.; Salfeld, J.; Franceschinelli, F.; Cozzi, A.; Dorner, M. H.; Arosio, P. Biochemistry 1989, 28, 5179-5184. (c) Treffry, A.; Harrison, P. M.; Luzzago, A.; Cesareni, G. FEBS 1989, 247, 268-272. (d) Lawson, D. M.; Treffry, A.; Artymiuk, P. J.; Harrison, P. M.; Yewdall, S. J.; Luzzago, A.; Cesareni, G.; Levi, S.; Arosio, P. FEBS 1989, 254, 207-210. (e) Levi, S.; Yewdall, S. J.; Harrison, P. M.; Santambrogio, P.; Cozzi, A.; Rovida, E.; Albertini, A.; Arioso, P. Biochem. J. 1992, 288, 591-596. (f) Treffery, A.; Bauminger, E. R.; Hechel, D.; Hodson, N. W.; Nowik, I.; Yewdall, S. J.; Harrison, P. M. Biochem. J. 1993, 296, 721-728. (11) Treffry, A.; Hirzmann, J.; Yewdall, S. J.; Harrison, P. M. FEBS 1992, 302, 108-112.

10.1021/la990403g CCC: $15.00 © 1999 American Chemical Society Published on Web 07/30/1999

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ment (less than 24 iron atoms/ferritin molecule).12 Studies using Mossbauer and extended X-ray absorption fine structure (EXAFS) spectroscopy, and other methods have revealed that larger clusters form, probably at nucleation sites on the inner surface of the protein, followed by continued iron oxidation on the developing mineral core.13 It is not clear what constitutes the ferroxidase center; however there is compelling data which suggests that tyrosinate residues are probably involved.14 In short, these studies have revealed that electron-transfer steps are involved in the loading and unloading of iron by ferritin. The formation of ferritin involves an oxidation process, while the mobilization of iron from the mineral core requires that iron is first reduced. However, the details of the electron-transfer steps remain unclear. Learning more about the role of electron transfer, therefore, is key to understanding the mechanisms of iron uptake and release. Electrochemical methods have been used extensively to explore the redox properties of proteins. Effecting the direct electron transfer of the protein with the electrode allows one to observe how the protein responds directly to applied potential and to measure electron transfer kinetics. The availability of a variety of electrode materials and electrode modification techniques can greatly facilitate these investigations. In our initial work, the direct electron transfer of ferritin was studied on gold electrodes modified with 3-mercaptopropionic acid (MPA).15 Subsequently, it was discovered that ferritin can be immobilized by adsorption at submonolayer coverages on tin-doped indium oxide (ITO) electrodes in an electroactive state. Significant hydrophobic interaction between ferritin and ITO appears to be important in the adsorption process.16 This work has been an extension of many studies which have used ITO to characterize the electrochemical properties of other proteins such as cytochromes,17-23 myoglobins,24-28 ferredoxins,29-33 and hemoglobin.34-35 Though most of the (12) Xu, B.; Chasteen, N. D. J. Biol. Chem. 1991, 266, 19965-19970. (13) Yang, C. Y.; Meagher, A.; Huynh, B. H.; Sayers, D. E.; Thiel, E, C. Biochemistry 1987, 26, 497-503. (14) Waldo, G. S.; Ling, J.; Sanders-Loehr, J.; Thiel, E. C. Science 1993, 259, 796-798. (15) Martin, T. D.; Monheit, S. A.; Niichel, R. J.; Peterson, S. C.; Campbell, C. H.; Zapien, D. C. J. Electroanal. Chem. 1997, 420, 279290. (16) Cherry, R. J.; Bjornsen, A. J.; Zapien, D. C. Langmuir 1998, 14, 1971-1973. (17) Yeh, P.; Kuwana, T. Chem. Lett. 1977, 1145-8. (18) Cohen, D. J.; Hawkridge, F. M.; Blount, H. N.; Hartzell, C. R. Charge Field Eff. Biosyst., [Pap. - Int. Symp. Bioelectrochem. Bioenerg.]; Allen, Milton Joel, Usherwood, Peter Norman Russell, Eds.; Abacus Press: Tunbridge Wells, U.K., 1984. (19) Koller, K. B.; Hawkridge, F. M.; Fauque, G.; LeGall, J. Biochem. Biophys. Res. Commun. 1987, 145, 619-24. (20) Taniguchi, I.; Kurihara, H.; Yoshida, K.; Tominaga, M.; Hawkridge, F. M. Denki Kagaku oyobi Kogyo Butsuri Kagaku 1992, 60, 1043-9. (21) Bowden, E. F.; Hawkridge, F. M.; Chlebowski, J. F.; Bancroft, E. E.; Thorpe, C.; Blount, H. N. J. Am. Chem. Soc. 1982, 104, 7641-4. (22) Yuan, X.; Hawkridge, F. M.; Chlebowski, J. F. J. Electroanal. Chem. 1993, 350, 29-42. (23) Salamon, Z.; Tollin, G. Photochem. Photobiol. 1993, 58, 730-6. (24) King, B. C.; Hawkridge, F. M. J. Electroanal. Chem. Interfacial Electrochem. 1987, 237, 81-92. (25) Taniguchi, I.; Watanabe, K.; Tominaga, M.; Hawkridge, F. M. J. Electroanal. Chem. 1992, 333, 331-8. (26) King, B.; Hawkridge, F. M.; Hoffman, B. M. J. Am. Chem. Soc. 1992, 114, 10603-10608. (27) Tominaga, M.; Kumagai, T.; Takita, S.; Taniguchi, I. Chem. Lett. 1993, 1771-4. (28) Nassar, A.-E. F.; Willis, W. S.; Rusling, J. F. Anal. Chem. 1995, 67, 2386-92. (29) Taniguchi, I. Proc.-Electrochem. Soc. 1993, 93-11(Proceedings of the Fifth International Symposium on Redox Mechanisms and Interfacial Properties of Molecules of Biological Importance, 1993), 9-20. (30) Taniguchi, I.; Hayashi, K.; Tominaga, M.; Muraguchi, R.; Hirose, A. Denki Kagaku oyobi Kogyo Butsuri Kagaku 1993, 61, 774-5.

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studies have investigated the electron-transfer properties of dissolved proteins, the electrochemistry of proteins adsorbed at ITO electrodes has also been studied.36-40 In this work, the ITO surface will be used as a venue where both the release and uptake of iron by ferritin can be studied directly. The direct electron transfer of ferritin adsorbed at ITO electrodes makes it possible to induce the mobilization of iron from the mineral core matrix by applying a reducing potential to the ferritin layer. Moreover, adsorbed apoferritin can be reconstituted through its exposure to ferrous ion at a potential which sustains ferritin in the oxidized form. For the first time, the primary functions of ferritin can be manipulated by controlling the electrochemical potential. Experimental Section The electrochemical cells used in this study have been described elsewhere.39 Voltammetric scans were performed with a Cypress model Omni 90 potentiostat (Lawrence, KS) and a BioAnalytical systems Model RXY recorder (West Lafayette, IN). A PerkinElmer (Norwalk, CT) model 552 UV-vis spectrophotometer was used to measure the concentration of ferritin following sizeexclusion chromatography. The purification of commercially available ferritin utilized a Flex-Column gravity liquid chromatography column purchased from Kontes Glass Co.(Vineland, NJ). Water was purified using a Milli-Q water treatment system manufactured by Millipore Corporation (Bedford, MA). Sodium hydroxide (Analytical Reagent) and sodium phosphate, monobasic (Analytical Reagent), were purchased from Mallinckrodt Specialty Chemicals Co. (Paris, KY), and EDTA, disodium salt, dihydrate (Analyzed Reagent) was obtained from J. T. Baker Chemical Co. (Phillipsburg, NJ). Bio-Rad Protein Assay G-250 Dye was purchased from Bio-Rad Laboratories (Hercules, CA). Sigma Chemical Co. (St. Louis, MO) was the source of the following compounds: bovine albumin (fraction 5 powder), horse spleen ferritin (108 mg/mL), horse spleen apoferritin (50 mg/ mL), sodium azide (>99.0%), and phenylmethyl sulfonyl fluoride (PMSF) (>99%). All of the chemicals except for ferritin and apoferritin were used as obtained without further treatment. Tin-doped indium oxide on glass was cut into 6 × 10 mm pieces, followed by sonication in saturated Alconox (in 95% ethanol) for 5 min, then sonicated twice for 5 min in Milli-Q water. Finally, the electrodes were allowed to hydrate for 24 h. In the experiments described below, the electrode was clipped to an alligator type connector and suspended in solution by the electrical lead wire so that only a 6 × 6 mm area was immersed in solution. In cases in which an electrode was transferred between different cells under an oxygen-free environment, the electrode was suspended in a tube (Supporting Information Figure 1). Opening the Teflon stopcock allows a stream of pressurized nitrogen to blanket the electrode. A column of dimensions 25 cm × 2.5 cm was packed with a gel filtration medium of Sephadex G-200 (protein fractionation range, 5000-600000 Da). The column was equilibrated with 500 mL of G-200 buffer (20 mM pH 7.0 phosphate buffer, 0.9% NaCl, (31) Taniguchi, I.; Hirakawa, Y.; Kwakiri, K.; Tominaga, M.; Nishiyama, K. J. Chem. Soc., Chem. Commun. 1994, 953-4. (32) Taniguchi, I.; Muraguchi, R.; Nishiyama, K. Denki Kagaku oyobi Kogyo Butsuri Kagaku 1994, 62, 985-6. (33) Nishiyama, K.; Ishida, H.; Taniguchi, I. J. Electroanal. Chem. 1994, 373, 255-58. (34) Grubbs, William T.; Rickard, Lyman H. Charge Field Eff. Biosyst.-2, [Proc. Int. Symp.]; Allen, Milton Joel, Cleary, Stephen F., Hawkridge, Fred M., Eds.; Plenum: New York, NY, 1989; pp 129-136. (35) Detrich, J. L.; Erb, G. A.; Beres, D. A.; Rickard, L. H. Charge Field Eff. Biosyst.-3 [Int. Symp.] Editor(s): Allen, Milton J., Ed.; Birkhaeuser: Boston, MA, 1992; pp 41-52. (36) Bowden, E. F.; Hawkridge, F. M.; Blount, H. N. J. Electroanal. Chem. Interfacial Electrochem. 1984, 161, 355-76. (37) Koller, K. B.; Hawkridge, F. M. J. Am. Chem. Soc. 1985, 107, 7412-17. (38) Koller, K. B.; Hawkridge, F. M. J. Electroanal. Chem. Interfacial Electrochem. 1988, 239, 291-306. (39) Willit, J.; Bowden, E. F. J. Electroanal. Chem. 1987, 221, 265274. (40) Daido, T.; Akaike, T. J. Electroanal. Chem. 1993, 344, 91-106.

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0.2 mM PMSF, and 0.05% NaN3). Ferritin (or apoferritin) fractions separated on the column were combined, and the protein concentration was determined by the Bradford method.41 Purified ferritin was diluted with pH 7 phosphate buffer, then the electrodes were immersed into ferritin solution at open circuit potential for at least 60 h to attain saturation coverage. Ionic strength was modified by varying the concentration of the phosphate buffer while maintaining a constant neutral pH. For the current-potential curves of adsorbed ferritin, the treated electrode was rinsed in pure electrolyte and then placed in pure electrolyte solution (pH 7 phosphate or HEPES buffer), and the potential was cycled between 0.10 and -0.80 V. The baseline of the anodic peak was estimated by drawing a straight line from the leading and tailing limits of the peak. The charge of the cathodic peak was determined by the cut-and-weigh method. Experimental packing densities Γ were determined by substituting the electrolytic charge into the Faraday law. The electrode surface area occupied by each ferritin molecule was calculated from the reported diameter of the protein (12.0 nm).1 The theoretical packing density of ferritin was estimated from the “footprint” area of ferritin. Experiments examining iron release by ferritin first involved adsorbing a layer of ferritin as described above. Following the rinsing away of soluble species, the electrode was immersed into pH 7 phosphate buffer containing 10 mM EDTA. The currentpotential curves were then scanned negatively from 0.10 V at a scan rate of 100 mV/s. Iron uptake experiments began with the removal of iron from an adsorbed layer of ferritin, as discussed above. The resulting ITO/apoferritin modified electrode was rinsed and then immersed into deaerated, pH 7 HEPES buffer (0.5 M in HEPES and 1.0 M in NaCl, adjusted to pH 7.0) containing 1 × 10-4 M ferrous ammonium sulfate solution for 20 min. Immersion is made at a controlled potential of 0.20 V, sustaining the adsorbed ferritin in the oxidized form. The electrode is rinsed and then immersed into pure HEPES buffer, and the potential is cycled from 0.20 and -0.70 V at 100 mV/s. When dioxygen was used as the oxidant, the ITO/apoferritin modified electrode was immersed into the HEPES/Fe2+ solution at open circuit potential for 20 min, followed by rinsing and transferring under nitrogen to a cell containing pure HEPES buffer. There, the potential was cycled between 0.20 and -0.70 V at 100 mV/s. To determine whether apoferritin in fact, resulted from electrochemical reduction in the presence of EDTA, apoferritin was applied directly onto the ITO electrode in a control experiment. ITO electrodes were exposed to apoferritin in phosphate buffer for 60 h. The electrodes were rinsed and then immersed in HEPES/Fe2+ for 20 min at a controlled potential of 0.20 V. This was followed by rinsing and immersing the electrode into pH 7 HEPES buffer, where the potential was cycled between 0.20 and -0.70 V.

Results and Discussion Characterization of the ITO/Ferritin Electrode. A clean tin-doped indium oxide electrode (ITO) was immersed into a ferritin sample for 60 h, followed by rinsing and immersion into pure pH 7 phosphate buffer (µ ) 1.0 M), and the potential was cycled between 0.10 and -0.80 V. The current-potential curve shown in Figure 1 clearly indicates the presence of an electroactive immobilized species, which we attribute to an adsorbed layer of ferritin.16 Three features are prominent: an initial cathodic peak at -0.62 V, an anodic peak at -0.19 V in the return scan, and a new cathodic peak at -0.36 V in the second cycle. The first cathodic peak is attributed to the reduction of adsorbed ferritin. The midpoint potential of dissolved ferritin, measured by mediated microcoulometry, is reported to be -0.19 (vs NHE), or about -0.41 V (vs Ag/Ag/Cl).8a The cathodic peak potential reported in this work is substantially more negative, indicating that adsorbed ferritin undergoes significant conformational (41) Bradford, M. M. Anal. Biochem. 1976, 72, 248-54.

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Figure 1. Current-potential curve of ferritin adsorbed at an ITO electrode. An ITO electrode was immersed in pH 7.0 phosphate buffer containing ferritin for 60 h, rinsed, and reimmersed in pure buffer: ferritin concentration, 0.10 mg/mL; electrode area, 0.36 cm2; scan rate, 100 mV/s.

deformation upon adsorption onto the ITO surface. It has been reported previously that at lower ionic strengths, the reduction process (of ITO/ferritin) is represented by two peaks, rather than one.16 In addition, the original layer reconstructs into a new layer which exhibits quasireversible kinetics, indicated by the anodic and new cathodic peaks. The reconstruction may entail a change to an orientation, which is electroinactive as evidenced by the disappearance of the peaks after five cycles. Alternatively, reduction followed by desorption may cause the peaks to become smaller; this point will be addressed below. Since the focus of this work is the uptake and release of iron by adsorbed ferritin, the cause of the peak disappearance will be a subject of future study. By use of the reported diameter of tissue ferritin of 12.0 nm,1 and given that the quaternary structure of ferritin is nearly spherical, the theoretical area occupied per ferritin molecule is calculated to be 1.1 × 10-12 cm2. With this estimate, a theoretical packing density for ferritin is calculated to be 1.5 pmol/cm2. An experimental packing density of 1.3 pmol/cm2 was estimated from the charge under the initial cathodic peak, an electrode area of 0.36 cm2, and 1500 electrons transferred per molecule (based on 1500 iron atoms per ferritin molecule).15 Comparison of the packing densities suggests that roughly a monolayer of ferritin is formed. The pI of horse spleen ferritin is 4.5,1 while that of indium oxide is about 6.40 Hence, at neutral pH the ionic interaction between ferritin and the electrode surface is not expected to be as favorable as is the case for cytochrome c (pI ) 10) and ITO. However, in prior work we showed that the packing density of ferritin increases with ionic strength, suggesting that the adsorption of ferritin to ITO probably involves a significant hydrophobic interaction.16 As adsorption time increases, the packing density also increases slowly until it levels off at about 60 h.16 This time-dependent adsorption behavior is different than that of adsorbed cytochrome c. In the latter case, packing density is essentially independent of exposure time.40 Ferritin’s adsorption kinetics are such that the adsorption-desorption equilibrium takes place over longer periods of time. It is not clear why more time is needed for the system to reach a steady-state coverage, but in general, when long periods of time are required, surface

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Figure 2. Current-potential curve of ferritin adsorbed at an ITO electrode. An ITO electrode was immersed in pH 7.0 HEPES buffer containing ferritin for 60 h, rinsed, and re-immersed in pure buffer. Other conditions are the same as in Figure 1.

attachment is accompanied by reorientation or conformational changes.42 The current-potential curves of adsorbed ferritin in pH 7.0 phosphate buffer (µ ) 1.0 M) were measured at scan rates of 50, 100, 200, and 400 mV/s. Each currentpotential curve shows the same voltammetric features. The currents of all the peaks increase linearly with scan rate (correlation coefficient ) 0.986), a trend not only predicted by theory but observed with electroactive proteins in the adsorbed state.43 The initial cathodic peak obtained in the first scan does not have a return wave, indicative of irreversible processes. Its peak potential becomes more negative, shifting from -0.59 to -0.71 V as the scan rate is increased. The ∆Ep of the reconstructed layer increases gradually over the same scan rate range, indicating an electron-transfer couple exhibiting quasireversible electron-transfer kinetics. The current-potential curve of the ITO/ferritin electrode in pH 7.0 HEPES buffer is shown in Figure 2. The general voltammetric features acquired in phosphate buffer can also be found in this curve. However, the initial cathodic peak appears broader in HEPES, and its potential is shifted from -0.62 V (in phosphate) to -0.67 V. The quasi-reversible couple has undergone changes in potential as well, the midpoint potential shifting from -0.275 to -0.235 V. In addition, the peak potential difference has increased from 170 to 220 mV. As the pH increases, the potential of the initial cathodic peak becomes more negative, suggesting that the reduction is thermodynamically favored at lower pH (Supporting Information Figure 2). In microcoulometric studies in which dissolved ferritin was reduced using electrochemical mediators, it was also determined that the reduction potential of dissolved ferritin became more negative with an increase in pH.8a More specifically, it was found that approximately 2H+ are transferred to the core for every Fe(III) reduced since the potential became more negative with pH at a rate of approximately 120 mV/pH. In another study in which mediators were used, the change was (42) Norde, W. Adv. Colloid Interface Sci. 1986, 25, 267-340. (43) Tachikawa, H.; Zeng, K.; Zhu, Z.; Davidson, V. L. Meeting Abstracts of The 193rd Meeting of The Electrochemical Society; Electrochemical Society, Inc.: Pennington, NJ, 1998; Vol. 9801, no. 915. (44) (a) Clegg, G. A.; Fitton, P. M.; Harrison, P. M.; Treffery, A. Prog. Biophys. Mol. Biol. 1980, 36, 56-86. (b) Funk, F.; Lenders, J. P.; Crichton, R. R.; Schneider, W. Eur. J. Biochem. 1985, 152, 167.

Figure 3. (a) Current-potential curve of an ITO/ferritin electrode in the presence of 10 mM EDTA. Other conditions are same as in Figure 1. (b) Current-potential curve of an ITO/ ferritin in pure electrolyte following immersion in EDTA solution. An ITO/ferritin electrode was immersed in 5 mM EDTA in pH 7.0 phosphate buffer for 6 h, followed by rinsing and immersion in pure buffer at 0.20 V. Other conditions are as in Figure 1.

measured at 100 mV/pH.45 The rate observed here for the direct electrolysis of adsorbed ferritin is -125 ( 3 mV/ pH, which agrees very closely to the rates measured for dissolved ferritin. This agreement suggests that adsorbed ferritin behaves electrochemically very similar to dissolved ferritin. One interpretation of the data is that changes occurring in the protein structure upon adsorption have not significantly altered its ability to behave as ferritin in the dissolved state. Curiously, the midpoint potential of the quasi-reversible couple changes at a rate of -15 ( 0.5 mV/pH. This result suggests protons are not involved in the electrochemical reactions of the core following its initial reduction; however, the details of the reaction are not clear. Studies of Iron Release by Adsorbed Ferritin. It is known that iron can be induced to exit the mineral core following the reduction of ferritin iron.3-6 Reductioninduced iron transport from adsorbed ferritin in low ionic strength electrolyte has been reported previously.16 Figure 3a shows the i-E curve of an ITO/ferritin electrode immersed in pH 7.0 phosphate buffer (µ ) 1.0 M) containing 10 mM EDTA. The electrode was removed from solution, rinsed, and placed back into pure buffer. Scanning to -0.70 V reduces the iron core as evidenced by the cathodic peak at -0.60 V. However, what is striking is the absence of the anodic branch of the reconstructed layer, (45) Chasteen, N. Dennis; Ritchie, I. M.; Webb, J. Anal. Biochem. 1991, 195, 296-302.

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indicating that following reduction of the core, the iron has exited the protein shell and has complexed with the surrounding EDTA. Moreover, the missing cathodic peak on the second scan, indicating the absence of iron(III) in the ferritin core, further supports this conclusion. In the present system, no reducing agents were used. Hence, the unloading mechanism could not have involved reducing agents entering the protein through the channels as suggested in previous studies.44 Electrons are transferred directly from the electrode to the protein core. These data suggest that electrons pass through the protein coat by a tunneling mechanism, or perhaps electron transfer is mediated, for example, by iron ions residing in the channels.7 It was important to determine whether the iron release, discussed above, was induced by reduction of ferritin iron or simply caused by direct chelation of Fe(III). An ITO/ ferritin electrode was immersed in pH 7.0 phosphate buffer containing 10 mM EDTA for 6 h, providing ample time for EDTA to remove iron from the complex. The electrode was then rinsed free of residual EDTA and immersed in a pure phosphate buffer. The current-potential curve, shown in Figure 3b, displays all the peaks given in Figure 1. The presence of the anodic and cathodic peaks of the reconstructed layer suggests that iron still remains in the protein shell even after prolonged exposure of ferritin (oxidized form) to EDTA. The significance of this result for the present work is that the apparent loss of iron observed in Figure 3a is not due to EDTA acting on the iron core apart from reduction. Rather, iron appears to be mobilized only after it is reduced. The consistency of this result with those of other in vitro studies referenced above implies that the ferritin in the adsorbed state possesses similar stability to ligand exchange as the native protein. Iron Release Dependence on Buffer Anion. Referring once again to Figure 3a, the anodic and cathodic peaks of the reconstructed layer are absent following reduction in EDTA in phosphate buffer. On the other hand, when HEPES buffer is used, the anodic and cathodic peaks of the reconstructed layer are still present, though small. Though in each buffer the curve shows that the reduction of ferritin iron has been effected, there remains some unreacted Fe(II) in the core when the experiment is done in HEPES buffer. This result suggests that the driving force for the formation of the Fe2+-EDTA complex is smaller in HEPES buffer than in phosphate buffer. The attenuation of this driving force is not well understood. It has been suggested that phosphate, since it is a constituent of the mineral core, may be involved in buffering the pH changes within the cores resulting from iron redox reactions.8a The small size of phosphate ion allows it to enter and leave the protein shell, whereas the relatively large size of HEPES precludes its transfer through the channels. Perhaps the reduction of ferritin is facilitated by the buffering effect of phosphate. Studies of Iron Uptake by Adsorbed Ferritin. In vitro studies have shown that ferritin can be reconstituted from apoferritin in the presence of Fe(II), and not Fe(III).1 An ITO/ferritin electrode was immersed in EDTA/buffer solution, and the potential scanned to -0.70 V, emptying adsorbed ferritin of its iron. The electrode was then incubated in HEPES buffer containing 1 × 10-4 M ferrous ammonium sulfate for 20 min at 0.20 V. The controlled potential of 0.20 V is negative enough to maintain dissolved iron essentially in its reduced form, while sustaining ferritin iron in its oxidized form. The electrochemical system is kept oxygen free by deaerating the buffer with nitrogen from the bottom of the cell. Anaerobic conditions are essential in order to eliminate any possibility of

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Figure 4. Current-potential curves of an ITO/“reconstituted ferritin” electrode following uptake of iron at 0.20 V. An ITO/ ferritin electrode was immersed in 1 × 10-4 M ferrous ammonium sulfate in pH 7 HEPES buffer at 0.20 V for 20 min, rinsed, and immersed in pure HEPES buffer at 0.20 V. Other conditions as in Figure 2.

oxidation of ferritin by dissolved oxygen. In addition, though a HEPES radical has been shown to be produced during iron oxidation by dissolved ferritin,46 the absence of oxygen should prevent the formation of the radical. Following a rinsing away of dissolved species after the incubation, the electrode was placed in pure buffer and the potential cycled between 0.20 and -0.70 V. The current-potential curve, shown in Figure 4, exhibits a negatively shifted cathodic wave on the first scan and a quasi-reversible anodic/cathodic couple on subsequent scans as displayed in Figure 2. The exposure of the adsorbed layer to ferrous ion at a potential of 0.20 V has thus resulted in the appearance of a voltammetric response similar to that in Figure 2, suggesting that iron has indeed entered the protein shell. Given that the peak potentials are similar to those of adsorbed ferritin, it is probable that the iron oxidation leads to the ferrihydrite structure that has been found with reconstituted ferritin in other studies.7 Moreover, in the absence of molecular oxygen, Fe(II) is still converted to Fe(III) by the applied electrode potential, suggesting that during, or after, entry into the protein sphere, the iron was oxidized. However, the potentials exhibited by adsorbed reconstituted ferritin are not exactly the same as those for adsorbed natural ferritin, the initial cathodic peak has shifted from -0.67 to -0.46 V, and the midpoint potential of the quasi-reversible couple is essentially unchanged. Moreover, the peak potential difference has changed from 220 to 280 mV. Since phosphate is largely adventitious, ferritin can be reconstituted without phosphate.7,47 In addition, it was found through Mo¨ssbauer spectroscopy that the iron core of reconstituted ferritin was indeed that of the ferrihydrite structure and that the structure was more ordered than that of native ferritin.48 It should be mentioned that phosphate is no longer in the iron core when ferritin is reconstituted electrochemically, and the absence of phosphate may be responsible for the lower redox potentials. The reduction potential of bacterioferritin from Azobacter vinlandii is lower than that of mammalian ferritin, (46) Grady, J. K.; Chasteen, N. D.; Harris, D. C. Anal. Biochem. 173, 111-115. (47) Macara, I. G.; Hoy, T. G.; Harrison, P. M. Biochem. J. 1972, 135, 151-162. (48) Mann, S.; Williams, J. M.; Treffry, A.; Harrison, P. M. J. Mol. Biol. 1987, 198, 405-416.

Uptake and Release of Iron

presumably due to its higher concentration of phosphate.8b It can be seen that the peaks of the i-E curve in Figure 4 are more than twice the size of those shown in Figure 2 (i-E curve of adsorbed, natural ferritin). The increase is probably due to a greater number of iron atoms in reconstituted ferritin than was present originally. In vitro studies have reported reconstituted ferritin containing three times the iron as native ferritin.48 In general, the current-potential curves of reconstituted ferritin and natural ferritin are more similar than they are different. Ferritin from horse spleen consists of >90% L-type subunit (22 of 24). Studies which compare the iron uptake behavior of recombinant L-chain and H-chain ferritin have revealed that the respective subunits have cooperative roles in iron uptake.10e Regions on the H-chain subunit have been found to possess ferroxidase activity, affording the ability to catalyze iron oxidation. While L-chain subunits do not possess such ferroxidase centers, ferritins composed purely of L-chain have a greater ability to nucleate iron. The presence of two H-subunits in horse spleen ferritin is sufficient to catalyze the iron oxidation. It is clear that the reconstitution of ferritin at an applied potential negative of the oxidation of ferrous ion strongly supports the theory that the protein catalyzes the oxidation of Fe(II). Iron uptake results in all of the adsorbed ferritin becoming electroactive again as indicated by the reappearance of the initial cathodic peak. It had been previously reported that the reduction-induced reconstruction of natural ferritin may be accompanied by desorption.16 However, these new data on iron uptake indicate that ferritin, which was thought to have desorbed, had actually undergone a reorientation, conformational change, or some other reconstruction rendering it electroinactive. What the reconstruction entails will be investigated in future work. Apoferritin Control Studies. Presumably, it is apoferritin which remains on the ITO surface after the reduction of adsorbed ferritin in the presence of EDTA. A control experiment has been designed which entailed forming a layer from an authentic apoferritin sample and to observe how apoferritin behaves under similar iron uptake conditions. First, however, it was of interest to ascertain the electrochemical response of adsorbed apoferritin. After purification, a layer of commercially prepared apoferritin was formed on an ITO electrode, as was done for ferritin. No peaks appear in the current-potential curve for the ITO/apoferritin electrode (Supporting Information Figure 3a) due to the absence of iron in adsorbed apoferritin. In the reconstitution experiment, an ITO/ apoferritin electrode was immersed into pH 7.0 HEPES buffer containing ferrous ammonium sulfate for 20 min at 0.20 V in order to induce the loading of iron. The electrode was rinsed again, immersed in pure HEPES buffer, and the potential cycled between 0.20 and -0.70 V (Supporting Information Figure 3b). The same peaks are displayed as those in Figure 2 indicating that the ferritin, emptied electrochemically, behaves similarly to apoferritin when re-exposed to ferrous ion. This result strongly supports the conclusion that adsorbed ferritin is converted to adsorbed apoferritin following reduction in the presence of a complexing agent. Iron Uptake Dependence on Electrode Potential. To better define the role of electrode potential on the uptake of iron, the immersion of ITO/apoferritin electrodes in ferrous ion solution was made for the same duration, but at different potentials. Figure 5 shows the initial cathodic peak area plotted against electrode potential under which ferritin was reconstituted. At 0.10 V, iron

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Figure 5. Iron uptake dependence on electrode potential. An ITO/“apoferritin” electrode was immersed in 1 × 10-4 M ferrous ammonium sulfate in pH 7.0 HEPES buffer for 20 min at the indicated potentials, rinsed, and immersed in pure HEPES buffer at 0.20 V. Other conditions are as in Figure 2.

uptake has reached its saturation level, indicating that this potential was sufficiently positive to oxidize iron entering the protein shell to Fe(III); potentials above 0.20 V do not increase the level of iron loading. At -0.40 V ferritin is maintained in its reduced form according to cyclic voltammetry. The experiment was repeated at -0.20 and -0.40 V; however, no peaks were observed in the ensuing scan, suggesting that at these potentials iron had either not entered the protein or entered but had not been oxidized. In a control experiment to show whether iron enters the apoferritin shell at -0.40 V, the ITO/apoferritin was exposed to ferrous ion at -0.40 V for 20 min, rinsed, and finally immersed in pure buffer for 20 min at 0.20 V in order to oxidize any iron which may have entered the protein. No appreciable peaks were displayed on the current-potential curve indicating that iron does not enter the apoferritin cavity at -0.40 V. Post-uptake Oxidation by Dissolved Oxygen. Most in vitro iron uptake studies have subjected apoferritin to ferrous ion in the presence of dissolved oxygen. To show whether ferritin, in the electrochemical system we are presenting in this report, behaves similarly to that in other in vitro studies, the applied oxidizing potential was replaced by exposure to dissolved oxygen. Ideally, exposure to oxygen should be done simultaneously with ferrous ion exposure; however, oxidation of ferrous ion was observed when this condition was attempted. Therefore, the exposure to oxygen was made after exposure to ferrous ion. An ITO/apoferritin (emptied electrochemically) was prepared as described above and then immersed into deaerated ferrous ammonium sulfate in HEPES buffer at open circuit potential. The electrode was rinsed under nitrogen to maintain an anaerobic condition and immersed in HEPES buffer saturated with oxygen. After rinsing under nitrogen and re-immersion in pure HEPES buffer, the potential was cycled between 0.20 and -0.70 V (Supporting Information Figure 4a). The curve displays similar electroactivity as shown in Figure 2 suggesting that oxygen has effected the oxidation of iron taken up by ferritin. The peaks however are 10 times smaller than when an oxidizing potential is applied during exposure to an anaerobic Fe(II) solution. The open circuit potential is approximately -0.17 V, a potential which, according to the results above, does not allow for saturated reconstitution. These current-potential data also show that oxidation can take place after iron uptake, suggesting that some Fe(II) can be loaded into the protein. These results are

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consistent with EXAFS experiments, which reveal that Fe(II) can load into ferritin without oxidizing.49 Though it appears that oxygen was responsible for converting Fe(II) to Fe(III), it was necessary to determine whether the application of 0.20 V at the start of the scan was responsible for oxidation of the iron. The above experiment was repeated, however omitting the exposure to molecular oxygen. The resulting current-potential curve (Supporting Information Figure 4b) shows essentially no electroactivity, indicating that oxidation was not effected by the applied potential during the scan. Clearly, omitting the exposure to oxygen allows the iron to remain in the reduced state, supporting the conclusion that oxygen effects the oxidation of iron(II) already present in ferritin. It was also of interest to determine whether iron can be oxidized by an applied potential after reconstitution at open circuit potential (OCP). Following immersion of an ITO/apoferritin electrode at OCP in ferrous ion for 20 min, the electrode was placed into pure buffer at 0.20 V for 20 min. The current-potential curve shows identical currents as that when molecular oxygen was used as the oxidizer, suggesting that oxidation by oxygen or applied potential is equally effective following the loading of iron(II). Conclusions Adsorbed ferritin, in the oxidized form, does not readily yield iron to chelating agents in solution. However, when iron in adsorbed ferritin is reduced by sufficient negative (49) Rohrer, J. S.; Joo, M.-S.; Dartyge, E.; Sayers, D. E.; Fontaine, A.; Thiel, E. C. J. Biol Chem. 1987, 262, 13385-13387.

Pyon et al.

potential, the mobilization of iron is induced. Electrochemically emptied ferritin, when exposed to ferrous ion, takes up iron at a potential positive enough to sustain ferritin in the oxidized form; at negative potentials (-0.20 to -0.40) V, iron does not accumulate in the emptied shell. Adsorbed apoferritin takes up iron similarly to electrochemically emptied ferritin, strongly suggesting that apoferritin is what remains as a result of electrochemically reducing ferritin in the presence of EDTA. Though immobilized, ferritin adsorbed at an ITO surface exhibits iron uptake and release in a fashion similar to ferritin in solution. Clearly, the results presented in this report show ITO/ferritin to be a system by which the primary functions of ferritin, the loading and unloading of iron, can be examined by controlling the potential of the electrode. This system thus provides an attractive means to probe the role of electron transfer in the mechanisms of ferritin’s functions. Acknowledgment. The authors gratefully acknowledge Research Corporation for support of this work (Cottrell Research Grant Number C-3735) and Applied Films Incorporated (Longmont, CO) for donating the tindoped indium oxide. The authors also thank Professor Edmond F. Bowden at North Carolina State University for his many helpful suggestions. Supporting Information Available: A drawing of the ITO electrode, cathodic peak potential dependence on pH, and current-potential curves of a commercial ITO electrode and modified ITO electrode. This material is available free of charge via the Internet at http://pubs.acs.org. LA990403G