Mechanism of Hydrogen Peroxide Production by Copper-Bound

Jan 5, 2009 - The amyloid beta peptide (Aβ) of Alzheimer's disease evolves hydrogen peroxide in vitro in the presence of Cu(II), external reducing ag...
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J. Phys. Chem. B 2009, 113, 1202–1209

Mechanism of Hydrogen Peroxide Production by Copper-Bound Amyloid Beta Peptide: A Theoretical Study Nadine Hewitt and Arvi Rauk* Department of Chemistry, Memorial UniVersity of Newfoundland, St. John’s, NL, Canada A1B 3X7, and Department of Chemistry, UniVersity of Calgary, Calgary, AB, Canada T2N 1N4 ReceiVed: August 15, 2008; ReVised Manuscript ReceiVed: October 22, 2008

The amyloid beta peptide (Aβ) of Alzheimer’s disease evolves hydrogen peroxide in vitro in the presence of Cu(II), external reducing agents, and molecular oxygen, without producing detectable amounts of the oneelectron reduced intermediate, superoxide, O2-•. The mechanism of this process was examined by ab initio computational chemistry techniques in systems that model the binding of Cu(II) to the His13His14 fragment of Aβ. The catalytic cycle begins with the reduction of the most stable Cu(II) complex to the most stable Cu(I) complex. This Cu(I) complex forms a Cu(II)-like adduct with 3O2 that cannot dissociate in water to yield O2-•. However, it can be reduced by proton-coupled electron transfer to an adduct between HOO- and the Cu(II)-like complex, which in turn can be protonated. The protonated complex decomposes to yield H2O2 by an associative-dissociative mechanism, thus completing the cycle. Introduction Alzheimer’s disease (AD) is a progressive disease which affects 2% of the population in industrialized countries.1 In the development of AD, β amyloid peptide plays a crucial role,2-5 and oxidative stress is a major factor in its neurotoxicity.6-8 The β amyloid peptide binds to the redox active transition metal copper with high affinity.9,10 Early reports indicated that the Cu(II)/Aβ complex has a relatively high reduction potential, E° ) +0.72-0.77 V, versus the standard hydrogen electrode (SHE),11 and is readily reduced to Cu(I) by Aβ itself or endogenous reducing agents such as ascorbate and thiols.12 More recent studies13,14 on Aβ and fragments of it clearly show that the reduction potential is 0.28-0.34 V versus SHE, and that the monomeric copper/Aβ complex is probably not reduced by Aβ itself,13 although its reduction by exogenous reducing agents was confirmed. Oxidation of the reduced metal Cu(I) back to Cu(II) by molecular oxygen is formally a single-electron process and should yield superoxide, O2-•, as a product. However, in cell-free environments, superoxide is not detected, but the two-electron reduction product, hydrogen peroxide is observed.11,15 The mechanism of this process is not clearly understood but is thought to follow superoxide dismutaselike or galactose oxidase-like chemistry.16 Three superoxide dismutases (SODs) exist in mammalian systems, a Mn-based SOD in mitochondria (SOD2) and two CuZn forms, SOD1 in the cytosol and SOD3 in extracellular spaces. SODs convert superoxide to oxygen and hydrogen peroxide in two single-electron redox steps, as shown in eqs 1 and 2:

Mn + O2-• f Mn-1 + O2

(1)

Mn-1 + O2-•+2H+ f Mn + H2O2

(2)

where n denotes the cation charge. In eq 1, the oxidized form of the metal is reduced by superoxide, yielding the reduced metal * Author to whom correspondence should be addressed. E-mail: [email protected].

and oxygen. In the second step, the reduced metal further reduces another molecule of superoxide, returning the metal to its oxidized form and generating hydrogen peroxide after addition of two protons. The standard reduction potential for reduction of oxygen to superoxide is E°(O2/O2-•) ) -0.33 V.17 This requires that the reduction potential of the metal, E°(Mn/Mn-1), be greater than -0.33 V in order to carry out the first step (eq 1). For the second step (eq 2), E°(O2-• + 2H+/H2O2) ) 0.89 V.17 The reduction potential of the metal should be less than this, E°(Mn/Mn-1) e 0.89 V. The reduction potential of human CuZn SOD falls in the middle of this range, E° ) 0.36 ( 0.01 V,18 as does Cu(II)/Aβ. Most endogenous antioxidants (ascorbate, glutathione, etc.) are able to reduce the copper to Cu(I). Thus, the first step in the production of H2O2 by Cu/Aβ(1-42) is not eq 1 but, rather, its reverse (eq 3):

Cu(I)/Aβ + O2 f Cu(II)/Aβ + O2-•

(3)

In principle, the one-electron oxidation of the reduced copper complex with Aβ should produce superoxide. However, superoxide is not detected;11 rather, hydrogen peroxide is seen, presumably by the equivalent of eq 2, namely, eq 4:

Cu(I)/Aβ + O2-• + 2H+ f Cu(II)/Aβ + H2O2

(4)

Thus, the reduced copper complex must be able to reduce both molecular oxygen and the product superoxide. In this study, we examine by ab initio techniques the detailed mechanism of both events and offer an explanation of the lack of observation of the intermediate superoxide. We assume that endogenous reducing agents can maintain a steady state concentration of the reduced copper/Aβ complex. As in previous studies,5,19,20 we adopt 3-(5-imidazolyl)propionylhistamine (1) as a model of the His13His14 region of Aβ and examine the chemistry of its complexes with Cu(I) and Cu(II). Computational Methods All calculations were carried out by Gaussian 0321 using the hybrid density functional method, B3LYP.22 Geometry optimi-

10.1021/jp807327a CCC: $40.75  2009 American Chemical Society Published on Web 01/05/2009

Mechanism of H2O2 Production by Copper-Bound Aβ

zation, harmonic frequency calculation, and thermochemical parameters were determined using the 6-31+G(d) basis set, which we will refer to as the small basis set (SB). The frequency calculation confirmed that the optimized structure was at a local minimum on the potential energy hypersurface. The zero point energies were scaled by 0.9806.23 However, this was not done for the thermal correction of enthalpy or entropies. For more accurate enthalpies, single point energies were calculated using B3LYP with a large basis (LB), 6-311+(2df,2p). Atomic net charges and spin distributions were determined by natural atomic and bond orbital (NBO) analysis.24 Details of all computed quantities, and structural information, are provided in Supporting Information Tables S1 and S2, respectively. Molden 4.0 was used as an visualization tool.25 Free Energies of Solvation, ∆Gsolv, and Empirical Corrections. In order to calculate the free energy change in water, ∆G(aq), the change in the free energy of solvation, ∆∆Gsolv, was added to the free energy change in the gaseous phase, ∆G(g), corrected for a standard state of 1 M. ∆Gsolv was determined using the integral equation following the polarizable continuum model (IEFPCM)26 and the B3LYP/6-31G(d) density. The basis set without diffuse functions was selected to reflect the more compact nature of the wave function in solution. In our experience, charged species are undersolvated by the IEFPCM with standard scaling of the united atom Hartree-Fock (UAHF) radii, so selective scaling was applied as follows: the radii of the metal ion, atoms directly attached to it, and any atoms bearing a formal charge, whether positive or negative, were scaled by a factor of 1.1; all other atoms were scaled by the default value, 1.2. Experimental rather than calculated relative free energies of solvation were applied where available in order to reduce errors further. For the proton, ∆Gsolv(H+) ) -1107 kJ/mol was adopted.27 In the case of ∆Gsolv(O2-•), the equation 3 O2 + e f O2-• was employed. The standard reduction potential, E°(O2/O2-•) ) -0.33 V vs the standard hydrogen electrode (SHE), was used to derive ∆G(aq) ) -387 kJ/mol (ignoring the electron and using ∆G(aq) ) 418 kJ/mol for the SHE half-cell reaction).28 At the CCSD(T)/LB//B3LYP/SB level, ∆G(g) ) -20.8 kJ/mol. If one assumes that there is no error in the calculated ∆Gsolv(3O2) ) 5.8 kJ/mol, then one can derive ∆Gsolv(O2-•) ) -360 kJ/mol, which compares well with the calculated value, -354 kJ/mol. The free energy of solvation of protonated superoxide, HO2•, may be similarly derived. For the reaction HOO•(aq) f O2-•(aq) + H+(aq) and the experimental pKa ) 4.8, ∆G (aq) ) 28.8 kJ/mol. At the CCSD(T)/LB//B3LYP/ SB level, ∆G (g) ) 1469.2 kJ/mol. If one adopts the experimental values for ∆Gsolv(O2-•) ) -360 kJ/mol and ∆Gsolv(H+) ) -1107 kJ/mol, then one can derive ∆Gsolv(HOO•) ) -25.2 kJ/ mol (again in reasonable agreement with the calculated value, -31 kJ/mol). Lastly, from the experimental acid dissociation constant for H2O2, pKa ) 11.8, from which ∆G(aq) ) 66.3 kJ/ mol, the CCSD(T)/LB//B3LYP/SB level gaseous phase free energy change, ∆G(g) ) 1562.8 kJ/mol, and assuming that the calculated value for ∆Gsolv(H2O2) ) -44.0 kJ/mol is overestimated by 6 kJ/mol as for HOO•, one adopts ∆Gsolv(H2O2) ) -38.0 kJ/mol and derives ∆Gsolv(HOO-) ) -427.5 kJ/mol (calculated -416 kJ/mol). In summary, experimental (or best estimate) values of ∆Gsolv were adopted as follows (in kJ/mol):

J. Phys. Chem. B, Vol. 113, No. 4, 2009 1203 H+, -1107; H2O, -16.2; O2-•, -360.0; HOO•, -25.2; HOO-, -427.5; H2O2, -38.0. For all other species, free energies of solvation were taken as calculated by the procedure described above. Calculation of Reduction Potentials for “Cu(II)”/“Cu(I)” Redox Couples. The mechanisms proposed below involve both inter- and intramolecular single-electron transfer processes. The standard reduction potential of a “Cu(II)”/“Cu(I)” couple, relative to the standard hydrogen electrode (SHE), E°(“Cu(II)”/ “Cu(I)”), is defined by eq 5

E°(“Cu(II)”/“Cu(I)”) ) -(∆G(aq)Cu - ∆G(aq)SHE)/F (5) where F is the Faraday constant, F ) 96.485 kJ mol-1 V-1, ∆G(aq)SHE is the free energy change for the standard hydrogen cell half-reaction, 1/2H2(g) + e- f H+(aq) (∆G(aq)SHE ) -418 kJ mol-1, ignoring the electron),28 and ∆G(aq)Cu is the calculated free energy change for reaction 6, again ignoring the electron.

“Cu(II)”(aq) + e- f “Cu(I)”(aq) + aL

(6)

In eq 6, “Cu(II)” and “Cu(I)” represent species containing oxidized and reduced copper, respectively. The symbol aL recognizes the fact that a number of ligands may be shed in the reduction process, and that the associated entropy change may be an important component of the free energy change. The actual potential, E, of the half-reaction under ambient conditions is related to the standard potential, E°, through the Nernst equation (eq 7)

E ) E° - (RT/F) ln Q

(7)

where Q is the reaction quotient specifying concentrations of oxidized and reduced components and other species associated with the chemical change. In the special case that n protons are consumed in solution buffered at pH 7 under otherwise standard conditions, the reaction quotient reduces to Q ) 107n, and the symbol E°′ denotes the potential at pH 7 (E°′ ) E° - (RT/ F) ln Q ) E° - 0.41n V). For most of the energy differences calculated in the following sections, errors inherent in the calculation of absolute values could be expected to cancel, yielding reliable relative energies. However, this is less likely to be the case for the calculation of aqueous free energy changes for reactions such as eq 6. Since a transition element is involved and the number of electrons changes, the enthalpy change will be less accurately described at this theoretical level than expected for lighter elements. An extreme case is illustrated by the difference between the calculated and experimental second ionization potential of atomic copper (IE2 (calc) ) 2008 kJ mol-1; IE2 (exp) ) 1958 kJ mol-1 29). The discrepancy is most likely due to the unequal treatment of electron correlation (an enthalpic term). In a previous publication,5 we assumed that the error in the ionization potential of Cu+ will be present in the reduction potentials, E°(“Cu(II)”(aq)/“Cu(I)”(aq)), irrespective of the metal environment, since they all involved the change in copper oxidation state from +2 to +1. This approach is not appropriate in the present oxygen substituted systems, since reduction involves partial electron transfer both to the metal and the coordinated oxy species. In the present work, we adopt a softer approach, by correcting the enthalpy change for each reaction involving open shell Cu species by -50∆R kJ/mol, where ∆R is the change in the spin

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Figure 1. Structures 3B(aq), 1B(aq), and 2B(aq) optimized in solution at the B3LYP/6-31+G(d) level. All others were optimized in the gaseous phase. The large green balls are copper, the smallest balls are hydrogens, and the intermediate balls are the following: red, oxygen; blue, nitrogen; black, carbon. The numbers are lengths of bonds in Å.

on the Cu site for the reaction considered. Thus, for a pure ionization, Cu+ f Cu2+ + e-, ∆R ) 1.0, and the correction to the calculated ∆H is -50 kJ/mol. With the procedures described above, we expect aqueous free energy changes, ∆G(aq), will be accurate to (15 kJ/mol for all of the reactions considered. Results and Discussion Both steps in the mechanism of H2O2 production (eqs 3 and 4) begin with the reduced Cu(I)/Aβ complex, modeled here as a complex of Cu(I) with structure 1. The His13 and H14 residues are firmly established as ligands for Cu(II) in Cu(II)/Aβ. Previous modeling of Cu(II) binding to 1 has shown that OC (the oxygen of the amide bridge) is almost certainly a third ligand.20 According to spectroscopic studies on Cu/Aβ(1-28),30 the N-terminal amino group and His6 are involved in the coordination to the copper but not Tyr10. X-ray absorption spectroscopic data on Cu(II)/Aβ solutions, supported by DFT calculations, suggested an octahedrally distorted Cu site with 3N3O stoichiometry.31 The three N ligands were taken as

unconnected His residues, and the oxygen ligands were modeled as carboxylate groups and a water. At the present level of theory, axial ligands dissociate in water. For simplicity, we employ a water molecule as a surrogate fourth neutral ligand for Cu(II). The Cu(I) structure does not have the water. The most stable Cu(I) complex with 1 is A (Figure 1), with a predicted binding affinity of 93 kJ/mol at the B3LYP/6-31G(d) level.20 The structure of A was reoptimized at the B3LYP/SB level. The Cu(I)-N bonds increased slightly from 1.88 to 1.89 Å, and the tenuous coordination to Oc (the carbonyl oxygen) weakened further with an increase of the Cu-Oc distance from 2.13 to 3.23 Å, forming a macrocycle with the Cu(I) bridging the two imidazole groups. The N-Cu-N angle became more linear, changing from 156 to 177°. Overall, the Cu(I) species, A, has become a more open structure at the B3LYP/SB level, and the Cu(I) binding affinity is reduced to 83 kJ/mol. Possible pathways in the reaction of A with oxygen (eq 3) are shown in Figures 2 and 3. Figure 4 shows possible pathways in the reaction of A with superoxide (eq 4). In the figures, typical Cu(II) species are shown as bicyclic, as exemplified by structure 2E, where

Mechanism of H2O2 Production by Copper-Bound Aβ

J. Phys. Chem. B, Vol. 113, No. 4, 2009 1205

Figure 2. Scheme showing the interaction of 3O2 with A (Cu(I)/Aβ model) according to eq 3. Left angle bracket denotes spin multiplicity; right angle bracket denotes net charge. Numbers on arrows are ∆G(aq) in kJ/mol.

Figure 3. Scheme showing the interaction of 1O2 with A (Cu(I)/Aβ model) according to eq 3. Left angle bracket denotes spin multiplicity; right angle bracket denotes net charge. Numbers on arrows are ∆G(aq) in kJ/mol.

the third coordination site is occupied by Oc and the fourth by water. 2E (Figure 1) is the most stable Cu(II) species and is the simplest model for Cu(II)/Aβ. Reduction of 2E yields A: E°(2E/ A+H2O) ) 0.52 V vs SHE,32 somewhat higher than a literature report for Cu(II)/Aβ(1-42), E ) 0.28-0.34 V.13 The scheme shown in Figure 5 explains the observation that reduction of Cu(II)/Aβ (as modeled by 2E) in the presence of oxygen yields hydrogen peroxide without the intermediary production of detectable amounts of free superoxide. Step 1: Oxidation of A with Oxygen (eq 3). In the reaction of A with molecular oxygen, both triplet and singlet electronic states must be considered. Initially, the triplet surface is lower than the singlet by 94.2 kJ/mol, the difference in energy between 3 O2 and 1O2.33 The B3LYP/LB enthalpy difference, 160.9 kJ/ mol, is in poor agreement with the experimental value; the CCSD(T)/LB value is little better, 125.6 kJ/mol. We will treat the singlet and triplet surfaces separately. The Triplet Potential Energy Surface for Step 1 (eq 3, Figure 2). On the triplet state gaseous phase potential energy surface, the direct interaction of 3O2 with A yields only a very weak charge-dipole complex, 3B(g), in which the shortest

distance between the oxygen and A is 3.3 Å. The net charge on the Cu site is +0.72, and no charge has been transferred to the O2 moiety. A scan of the Cu-O separation yields only a steadily rising potential energy curve. At a typical Cu-O distance of 1.95 Å, the energy is only 17 kJ/mol higher than 3 B(g). 3B(g) is unstable both in the gaseous phase, ∆G(g) ) 15 kJ/mol, and in aqueous solution, ∆G(aq) ) 32 kJ/mol, largely due to loss of entropy and solvation. In short, no stable species corresponding to a Cu(II)-superoxide complex is found in the gaseous phase on the triplet surface. However, the solutionoptimized structure 3B(aq) (Figure 1) has a geometry typical of Cu(II) coordination. The Cu-O (of O2) separation is 2.04 Å, and the carbonyl O atom has moved to within 2.16 Å of the copper. NBO analysis of the gaseous phase wave function at the solution-optimized geometry shows a charge of +1.04 at the copper site and -0.29 on the O2 moiety. 3B(aq) is predicted to be more stable than 3B(g) by 22 kJ/mol but less stable than A + 3O2 by 10 kJ/mol. Since this is within our expected average error, 3B(aq) should be in equilibrium with A + 3O2. Protonation of the O2 moiety of 3B(g) or 3B(aq) to yield 3C (Figure 1) should promote further electron transfer from the

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Figure 4. Scheme showing the interaction of O2-• with A (Cu(I)/Aβ model) according to eq 4. Left angle bracket denotes spin multiplicity; right angle bracket denotes net charge. Numbers on arrows are ∆G(aq) in kJ/mol.

Figure 5. Scheme showing the proposed mechanism for H2O2 production from 2E (model of Cu(II)/Aβ) and 3O2. Left angle bracket denotes spin multiplicity; right angle bracket denotes net charge. Numbers on arrows are ∆G(aq) in kJ/mol. Standard reduction potentials, E° in V vs SHE, are explicitly shown.

copper. Indeed, 3C already has a typical Cu(II) coordination geometry in the gaseous phase and was not reoptimized in solution. Both the oxygen of the carbonyl group and one of the oxygens of the O2 moiety are within bonding distance, 1.93 and 2.11 Å, respectively. However, 3C is very unstable in aqueous solution, ∆G(aq) ) 87 kJ/mol relative to A + 3O2 + H+(aq). It is isoenergetic with the dissociated products, 2D (the tricoordinated Cu(II) complex with 1, Figure 1) and protonated superoxide, HO2•, ∆G(aq) ) 2 kJ/mol. As the pKa of HO2• is 4.8,34 proton loss to water will be spontaneous at physiological pH. Addition of H2O to 2D to complete the coordination environment of the Cu(II), yielding 2E, is exergonic by 50 kJ/ mol. The overall reaction which yields free superoxide, namely, A + 3O2 + H2O f 2E + O2-•, is endergonic, ∆G(aq) ) 51 kJ/ mol. We investigated whether the decomposition of 3C, yielding the stable Cu(II) complex, 2E, might follow an eliminationaddition mechanism as described above, or it may follow an addition-elimination pathway. This would entail prior addition of water to give a pentacoordinated intermediate, 3F (Figure 1), followed by loss of HOO• to yield 2E directly. In the gaseous phase, 3F is stable. It optimizes to a complex in which the OOH

moiety occupies an axial position and has a greatly elongated CuO bond, rCuO ) 3.34 Å. In water, the formation of 3F from 3 C is exergonic, ∆G(aq) ) -25 kJ/mol, and subsequent decomposition of 3F to 2E + HOO• is also exergonic, ∆G(aq) ) -24 kJ/mol. Because all steps in the addition-elimination pathway are exergonic, it implies that ligand exchange in a Cu(II) complex should occur by addition followed by elimination, as found for other complexes.35,36 The Singlet Potential Energy Surface for Step 1 (eq 3, Figure 3). On the singlet potential energy surface of A + 1O2, a complex, 1B(g), with typical Cu(II) coordination geometry, is formed. 1B(g) is lower in energy than A + 1O2, ∆H298 ) -38 kJ/mol, ∆G(aq) ) -18 kJ/mol. The Cu-O (of O2) separation is rCuO ) 1.95 Å, the Cu site bears a charge of +1.10, and the O2 moiety has a total charge of -0.42. In contrast to the triplet surface, optimization of 1B(g) in solution yields a structure, 1B(aq) (Figure 1), that is little different from the gaseous phase structure; the two CuO distances are rCuO ) 1.92 Å and rCuO ) 2.08 Å, respectively, and the charges on the Cu site and the O2 moiety are 1.1 and -0.46, respectively. Relative to A + 1O2, ∆H298 ) -29 kJ/mol, ∆G(aq) ) -19 kJ/mol. Thus, in both gaseous and aqueous phases, 1O2 forms a complex with A that

Mechanism of H2O2 Production by Copper-Bound Aβ resembles more closely a complex between superoxide and the tricoordinated Cu(II) complex, 2D. Protonation of 1B(aq) yields 1C (Figure 1), which also has the normal geometry of a Cu(II) complex. The OOH moiety is within normal bonding distance of Cu, rCuO ) 1.85 Å, and the oxygen of the bridging amide group completes the tetracoordination of the copper. 1C is less stable than 1B(aq) + H+(aq) by ∆G(aq) ) 14 kJ/mol. The decomposition of 1C yields the same products as 3C. The overall reaction, A + 1O2 + H2O f 2E + O2-•, is exergonic, ∆G(aq) ) -111 kJ/mol. In summary, the reaction of the reduced copper species, A, with triplet oxygen yields 3B, a complex with the characteristics of Cu(II) bound to superoxide. However, no free superoxide is produced, since the actual release of superoxide is strongly endergonic. On the other hand, singlet oxygen reacts exergonically with A to form 1B in which the copper is substantially oxidized, and 1B can dissociate exergonically to yield superoxide. Step 2: Reduction of Superoxide by A (eq 4, Figure 4). Step 2 is equivalent to the second step of the SOD mechanism (eq 2), namely, the production of hydrogen peroxide by the reduction of superoxide. The primary reaction of the most stable Cu(I) species, A, with O2-• corresponds to the interaction of two oppositely charged ions, and is strongly exothermic in the gaseous phase, ∆H298 ) -407 kJ/mol. However, the reaction is not accompanied by significant spin transfer from O2-• to A. The product complex, 2B(g), still has a Cu(I) coordination pattern, although the Cu-O (of O2-•) separation is relatively short, rCuO ) 2.17 Å. The high exothermicity in the gaseous phase is overwhelmed by the loss of solvation of the two ions, with the result that formation of 2B(g) is predicted to be endergonic in the aqueous phase, ∆G(aq) ) 101 kJ/mol. Reoptimization of 2B(g) with the solvent reaction field yields 2B(aq) (Figure 1) which has a structure very similar to 2B(g). For the formation of 2B(aq), ∆G(aq) ) 94 kJ/mol. Protonation of 2B(aq) at the terminal oxygen atom to form 2C is accompanied by electron transfer to the OOH moiety and oxidation of the copper to Cu(II). The copper binding environment rearranges to one typical of Cu(II), with both the oxygen of the carbonyl group, and of the OOH moiety within bonding distance, 2.11 and 1.89 Å, respectively. Net charges on the Cu site and the OOH fragment are +1.22 and -0.52, respectively, and the majority of the spin is on the copper and neighboring ligands, with 0.40 remaining on OOH. Formation of 2C by protonation of 2B(aq) in solution is also strongly exergonic, ∆G(aq) ) -141 kJ/mol. 2C can be formed directly from A + HOO• (∆G(aq) ) -47 kJ/mol), although the reaction may be slow at physiological pH because of the low concentration of HOO• (pKa ) 4.8). 2 C may dissociate directly into the coordinatively deficient Cu(II) complex, 2D and HOO-. While this reaction is implausible in the gaseous phase, the solvation of a +2 ion and a -1 ion in solution may permit the dissociation in solution. In fact, the dissociation in solution is also found to be endergonic, ∆G(aq) ) 57 kJ/mol. An associative pathway for the expulsion of HOOwas examined, but a pentacoordinated species was not found. The added water preferred to hydrogen bond to the OOH moiety instead. While the endothermicity of the dissociative pathway is essentially compensated by the additional 50 kJ/mol from addition of H2O to 2D to yield 2E, as mentioned above, and 27 kJ/mol from the protonation of HOO- to yield H2O2 at pH 7, it poses a barrier to the decomposition of 2C. An alternate pathway, involving prior protonation of 2C, was examined. Protonation of 2C to yield 2G (Figure 1) in solution is predicted to be isoergonic, ∆G(aq) ) -11 kJ/mol. Dissociative

J. Phys. Chem. B, Vol. 113, No. 4, 2009 1207 decomposition of 2G directly to 2D + H2O2 is slightly endergonic, ∆G(aq) ) 19 kJ/mol. However, it is likely that formation of H2O2 is by an associative route. Addition of water to 2G did not yield a pentacoordinated complex, but rather, the water inserted to become the fourth ligand, simultaneously expelling the H2O2. Displacement of H2O2 from 2G by added water to yield 2E is exergonic, ∆G(aq) ) -31 kJ/mol. Catalytic Production of H2O2 without Formation of Superoxide. Numerous observations have documented the catalytic production of H2O2 by Cu(II) complexes of Aβ in vitro in the presence of air and excess ascorbate or other reducing agents.16,37 The one-electron reduction product, superoxide, was not detected.11 The proposed mechanism in terms of the present models of Cu(II)/Aβ and Cu(I)/Aβ, namely, 2E and A, respectively, is shown in Figure 5. The Cu(II) species, 2E, is reduced to the Cu(I) species, A, by an external reducing agent, such as ascorbate or glutathione: E°(2E/A+H2O) ) 0.52 V. A reacts with oxygen to form a weak Cu(II)-like complex, 3B. 3B is subsequently reduced directly to 2C by proton-coupled electron transfer (PCET) from an external reducing agent: E°(3B+H+/ 2 C) ) 0.58 V. Direct reduction of 3B to 2B is not possible due to its very low reduction potential: E°(3B/2B) ) -0.88 V. 2C then is protonated further to 2G, a complex with hydrogen peroxide weakly coordinated to Cu(II). 2G then undergoes an SN2-like substitution reaction with water to complete the cycle. The cycle requires two reduction steps and two protonation steps. If the protons must be provided by the aqueous environment at physiological pH, then the free energy of these steps would increase by about 40 kJ/mol. However, the proton source may be the reducing agent, ascorbate or glutathione, or even Aβ itself,38 thus avoiding the penalty. The mechanism of catalytic production of H2O2 by Cu(II)/ Aβ has previously been studied by computational means by Barnham, Bush, and co-workers (BB&c).16 In their model system, the Cu(II) was assumed to be coordinated to the three His residues and to Tyr10 (anion) of Aβ, as well as an ascorbate moiety. The initial reduction to Cu(I)/Aβ was effected by electron transfer from the coordinated ascorbate to Cu(II), accompanied by proton transfer from ascorbate to tyrosinate and expulsion of the Tyr10 ligand, yielding a tetracoordinated Cu(I) species. Addition of 3O2 to this complex yielded an intermediate pentacoordinated Cu(I)-O2 species that underwent internal oxidation to Cu(II)-O2-• with the assistance of Hbonding to the nearby Tyr10. Subsequent H-atom transfer and a second proton from the aqueous medium completed the formation of H2O2 and the generation of a tyrosyl radical. Reduction of the tyrosyl radical by a second ascorbate reestablished the starting Cu(II)/Aβ system, completing the cycle. Evidence for the intermediacy of the tyrosyl radical is the observation of some dityrosine. The mechanism proposed by BB&c is similar to that proposed here, with some important differences. We think it is unlikely that Tyr10 would be so intimately involved on the production of H2O2. There is no direct evidence that Tyr10 is a ligand for the Cu(II) in Cu(II)/Aβ; it has been explicitly excluded spectroscopically in Cu/Aβ(1-28). 30 Second, at the higher level of theory employed here, tetra- and pentacoordinated Cu(I) complexes are unstable in water, as are penta- and hexacoordinated Cu(II) species. In BB&c’s mechanism, Tyr10 plays three roles. The first role is to accept a proton to facilitate electron transfer to the copper from ascorbate. We have assumed that this role could be played by the aqueous solvent or any other nucleophilic species that may be nearby. The second role is to “activate” the bound oxygen toward reduction by hydrogen

1208 J. Phys. Chem. B, Vol. 113, No. 4, 2009 bonding. They propose interaction with the vicinal oxygen of the O2 moiety. In our model system, the equivalent step is protonation of 3B(aq) at the distal oxygen to form 3C, a very endergonic process on the triplet potential energy surface (Figure 2). The third role of Tyr10 in BB&c’s scheme is the actual reduction of the oxygen by H-atom transfer, thereby generating the tyrosyl radical. In our model, the second reduction parallels the first by PCET from the reducing agent to generate the reduced species, 2C (Figure 5). In fact, the second reduction requires proton transfer with the proton ending on the distal oxygen. No involvement of Tyr10 is required. This is why, as BB&c themselves have shown, catalytic production of H2O2 continues even if the Tyr10 is mutated to Ala, albeit at a lower rate.16 The role of the water ligand in our model could conceivably be filled in Cu/Aβ by tyrosinate, but this is unlikely, since the addition of a negatively charged ligand to Cu(II) considerably lowers its reduction potential.5,39,40 We favor involvement of the N-terminal amino group or a third His residue, both of which are excellent ligands for Cu(II) and Cu(I). Conclusions We have examined the oxidation of Cu(I)/Aβ by singlet and triplet oxygen using a model, A, that encompasses the minimum Aβ coordination environment of the copper, namely, two adjacent His residues and the connecting amide bridge. Triplet oxygen forms a weak adduct with the copper in water, 3B(aq) (Figures 1 and 2) that has the bonding characteristics of Cu(II) but does not release superoxide. On the other hand, singlet oxygen is predicted to generate superoxide in the presence of A. We also examined the possibility that A could further reduce superoxide to hydrogen peroxide, by analogy with superoxide dismutase. In water, the adduct between A and O2-•, 2B(aq) (Figures 1 and 4), is predicted to be unstable by almost 100 kJ/mol, but reaction of A with protonated superoxide, HOO•, to form 2C is exergonic which is accompanied by internal oxidation of the copper to Cu(II). Hydrogen peroxide is readily released upon protonation of 2 C and exchange of the H2O2 by another ligand (water in our model). The reactions of A with 3O2 and with O2-• suggest a mechanism for the observed catalytic generation of H2O2 by Cu(II)/Aβ, oxygen, and reducing agents (Figure 5). Cu(II)/ Aβ, modeled by 2E (Figure 1), is reduced to Cu(I)/Aβ (modeled by A). A forms 3B, a loose adduct with 3O2. 3B is reduced by proton-coupled electron transfer from a reducing agent, to form 2C. 2C has the characteristics of a Cu(II) complex attached to HOO-. Upon protonation, 2C releases H2O2 by an associative substitution reaction, reforming the starting Cu(II) complex, 2E, in the process. This completes the catalytic cycle. Acknowledgment. The financial support of the Natural Sciences and Engineering Council of Canada (NSERC) is gratefully acknowledged. N.H. thanks NSERC for the award of an Undergraduate Research Fellowship. A generous amount of computer resources provided by Westgrid is greatly appreciated. Supporting Information Available: Tables showing the primary computed properties of all compounds discussed in the manuscript and the Gaussian archive entries of all B3LYP/631+G(d)-optimized structures. This material is available free of charge via the Internet at http://pubs.acs.org.

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