8006
J. Phys. Chem. B 1999, 103, 8006-8015
Ab Initio Quantum Mechanical Study of Metal Substitution in Analogues of Rubredoxin: Implications for Redox Potential Control Brian W. Beck,† John B. Koerner, and Toshiko Ichiye* Department of Biochemistry/Biophysics, Washington State UniVersity, Pullman, Washington 99164-4660 ReceiVed: April 20, 1999; In Final Form: July 26, 1999
Substitution of different metals into the redox sites of metalloproteins is a means of studying the structure of the native protein and of varying the redox properties of the protein. The implicit assumption is often made that metal substitution changes only intrinsic properties of the redox site such as the ionization potential without altering the surrounding protein or solvent. However, if this is not true, structural studies of metalsubstituted proteins will not reflect the native protein and the differences in redox potential upon metal substitution will not be simply the differences in ionization potential of the redox sites because of perturbations in the extrinsic electric field. Here, we present an ab initio unrestricted Hartree-Fock quantum mechanical study of metal substitution in the [M(SCH3)4)]2-/1- analogue, where M ) Fe, Co, Ni, and Zn, of the protein rubredoxin. Variations in several physical properties were determined and compared to experimental data. Upon metal substitution, only minor variations in geometry, atomic spin, and atom-centered partial charges of the redox site are observed. However, significant variation is found in the energies of reduction, on the order of 100-1000 mV. This indicates that when such substitutions are made into an Fe-S metalloprotein, little change will occur in the interactions between the metal site and the surrounding protein and thus the surrounding protein structure and the resultant electric field will not change. Thus, the structure is relevant to the native protein and the redox properties are mainly determined by the variations in the intrinsic ionization potential of the metal site and not the extrinsic field of the surrounding protein and solvent.
Introduction A crucial question in the study of biological electron transfer is identification of the molecular determinants of the redox potentials of electron-transfer proteins, both to gain a more fundamental understanding of the protein function and to guide efforts in engineering proteins with different redox properties. Metal substitution into the redox sites of electron transfer metalloproteins has become a means of studying the structure and function of the native protein. For instance, zinc substitution for the paramagnetic iron of the Fe-S proteins greatly simplifies nuclear magnetic resonance studies.1 In addition, metal substitution may be a means of engineering the redox potential of a protein since mutations in amino acids often give unpredictable results.2,3 Electronic structure calculations of analogues of metalloprotein redox sites are an important means of studying their structure and energetics.4-7 These studies also play an important role in determining parameters for force fields for molecular mechanics calculations, which have generally lagged behind for metalloproteins although there are studies in this direction.8-10 The focus here is on metallosulfur redox sites since sulfurs, usually from cysteinyl side chains, are among the most common ligands in metalloproteins. In particular, they are the ligands in the iron-sulfur (Fe-S) proteins, which are ubiquitous to living systems. The metal tetrathiolate [M(SR)4]1-,2- system studied here has a single metal in what is referred to here as a [1M] core, where M is the metal, with four tetrahedral sulfur ligands. * To whom correspondence should be addressed. E-mail:
[email protected]. Fax: 509-335-9688. † Current address: Institute for Molecular Design, University of Houston, Houston, TX 77204-5641.
The [1Fe] core is the simplest of the Fe-S redox sites and is found in the small proteins rubredoxin11 and desulforedoxin12 as well as in larger proteins such as desulfoferrodoxin, rubrerythrin, and nigerythrin13,14 A large number of synthetic analogues have been constructed,15-23 and a growing body of experiments on transition metal substitution in the rubredoxins and desulforedoxins is being formed.24-31 The [1Ni] core is of particular interest since Ni2+-substituted rubredoxins and desulforedoxins mimic some of the activities of Ni-containing hydrogenases32 and are structurally homologous to part of the Ni-Fe site in bacterial hydrogenases.33 Thus, there are numerous recent studies of the Ni2+-substituted rubredoxins and desulforedoxins.24,27-29,31 Although metal substitutions into the cubane-like [Fe4S4(SR)4]1-,2-,3- systems, referred to here as the [4Fe-4S] cores, have also been studied extensively in experiments of the ferredoxins and their analogues,34-42 studies of metal substitution in the [1Fe] core have certain advantages. For instance, electronic structure calculations of the [1Fe] core are much less computationally intensive than those of the [4Fe-4S] core. Perhaps more important is that substitution into the [1Fe] core is simpler to interpret since the single metal is substituted completely, whereas in the [4Fe-4S] core, generally only one of the four irons is replaced to yield a [M3Fe-S] core. Thus, for [1Fe] substitution, changes in redox properties can be linked directly with the metal being substituted, while for [4Fe-4S] substitution it is less clear how much the heterometal perturbs the redox reactions of the remaining irons. For example, the [Zn] core in substituted rubredoxin is not redox active,26,43 whereas the [Zn3Fe-4S] core in DesulfoVibrio africanus (Da) ferredoxin III and Pyrococcus furiosus (Pf) ferredoxin is redox
10.1021/jp9912814 CCC: $18.00 © 1999 American Chemical Society Published on Web 08/31/1999
Metal Substitution in Analogues of Rubredoxin active.37-40 In the latter, presumably at least two of the irons take part in the reduction, as is the case for the [4Fe-4S] core, and evidence indicates that generally the [3Fe-4S] fragment and not the heterometal is the primary locus of electron density change in the substituted cores.34 However, the redox potentials of the [Zn3Fe-4S] proteins differ by 80 to 100 mV relative to the [4Fe-4S] proteins, thus indicating the influence of the Zn. Thus, in an arbitrary [M3Fe-4S] core, it must be established if the substitution is changing redox properties because a metal that is being reduced is changed, as in the case of substitutions in [1Fe] cores, or because of indirect influences, as in the case of the [Zn3Fe-4S] core. To provide a framework in which to understand the molecular determinants of the redox potential, the redox potential of a redox-active protein can be divided into two major contributing factors: (1) the intrinsic ionization potential of the redox site independent of the protein (i.e., the difference in energy of the oxidized and reduced state of the redox site in a vacuum) and (2) the extrinsic electric field produced by the surrounding protein and solvent at the redox site.3,18,44,45 Of course, the protein may also influence the ionization potential by perturbing the nuclear and electronic structure of the redox site, and the redox site may also influence the protein and solvent structure by changes in the electrostatic distribution or the geometry of the site. Careful studies of the structures of these proteins as well as how changes in the protein (i.e., mutations) or the redox site perturb the redox potential are both important steps in understanding the molecular determinants of the redox potential. Altering the identity of the metal in a redox site affects the physical properties of the site due to variations in both the nuclear structure and the orbital occupancy, which can affect the relative energies of the oxidized and reduced states and thus the redox potentials. Among these properties are the internal energies, geometries, electron density, and atomic spin populations. The most obvious effect of metal substitution on the redox potential is alteration of the intrinsic ionization potential of the redox site. If this were the only change, then a metal-substituted protein could be considered to give reliable structural information about the native protein. Moreover, metal substitution could be a reliable means of altering the redox potential since the change in redox potential should simply be the difference in intrinsic ionization potentials of redox site analogues with the different metals. However, changes in the redox site geometry and electron density distribution may cause perturbations to the protein structure, which would also alter the extrinsic electric field produced by the protein and solvent. In addition, the protein may alter the redox site geometry, thus affecting the intrinsic ionization potential. Thus, the change in the redox potential in a protein would no longer be simply the difference in intrinsic ionization potentials of redox site analogues. In the Fe-S proteins, changes in the geometry upon metal substitution could potentially change the protein structure because the redox site is attached to the rest of the protein by covalent bonds between the cysteinyl sulfurs and the redox site irons. For example, the root-mean-square (RMS) difference between the crystallographic structures of Clostridium pasteurianum (Cp) rubredoxin in the native state (1IRO, 1.1 Å resolution)43 and with Zn substitution (1IRN, 1.2 Å resolution)43 from the Brookhaven Protein Data Bank46 is 0.23 Å for backbone atoms and 0.31 Å for all atoms.47 While these are relatively small differences, they are on the same order as the difference between the oxidized and reduced forms of Pf rubredoxin, which is 0.25 Å for backbone atoms,47 and could contribute significantly to the redox energy if the differences
J. Phys. Chem. B, Vol. 103, No. 37, 1999 8007 are localized near the redox site, as in the Pf rubredoxin case.48 Such structural changes may result in a change in the extrinsic electric field. A further consideration is that the geometric variations due to metal substitution observed in analogues may differ in proteins because the protein may restrain the geometry of the redox site. However, in terms of redox properties, our previous ab initio studies of the geometry optimized [Fe(SCH3)4]2-/1analogue indicate small variations in the bonds and bond angles and changes in dihedral angles controlling orientation of ligands do not significantly affect the intrinsic ionization potentials.7 However, an experimental study in which a serine was substituted for each of the cysteinyl ligands in turn showed a decrease in redox potential of 100 mV for the interior ligand mutants and of 200 mV for the exterior ligand mutants.49 The difference was attributed mainly to the 0.04 Å increase in bond length of the interior over the surface ligand metal-ligand bonds, which is seen both in the wild-type and the serine mutants, which would reduce ligand σ-donor strength, thereby increasing Z′eff, although other differences such as the orientation of the amide dipoles, number of NH‚‚‚S-Cys hydrogen bonds, and solvent accessibility were noted. A variation in charge distribution upon metal substitution with no change in protein structure will also change the electrostatic interaction energy and thus the redox potential, since the charge distribution determines the classical electrostatic interaction energy between the redox site and the electric field of the protein. More subtly, a variation in charge distribution may perturb the protein structure and thus the electric field due to the protein, which will also affect the redox potential. In fact, classical molecular dynamics (MD) simulations of rubredoxin and active site analogues with different partial charge distributions have significant variations in the solvation energy as well as the protein and solvent structure.47 Here, we present a quantum mechanical study of metal substitution by Co, Ni, and Zn in the [Fe(SCH3)4)]2-/1- analogue of the protein rubredoxin. Previous studies indicate that this analogue is a good model for the rubredoxin site. For instance, electronic structure calculations indicate that it is necessary and sufficient to model the β positions with methyl groups rather than to terminate with hydrogens.50 Moreover, our calculations indicate that different conformations with unstrained S-FeS-C dihedral angles have total energies that are within 1 kcal/ mol of each other and have very similar charge distributions so that variations due to the protein environment may be modeled classically.7 Various physical properties are compared to experimental data for metal-substituted sites and predictions as to the relative importance of these factors in the intrinsic versus extrinsic contributions to redox potentials are made. The first property examined is the geometry of the redox site. The next property examined is the electron density distribution of the redox site, which is evaluated by the Mulliken and electrostatic potential (ESP) derived atom-centered charges. Finally, explicit electronic properties associated with the metal are also examined. For instance, the ionization potentials are examined because they can be directly linked to the redox potential exhibited by the protein. In addition, the electronic properties such as the structures of the orbitals occupied by electrons provide insights into some of the observed geometric and electrostatic parameters. The overall goals are three-fold. The first is to evaluate the changes of the analogue with metal substitution that are likely to affect the protein structure, which are the geometry and the partial charges. The second is to evaluate the changes in the analogue with metal substitution
8008 J. Phys. Chem. B, Vol. 103, No. 37, 1999
Figure 1. Schematic representation of the D2d symmetry metallothiolate model.
that may contribute to the differences in redox potential of the analogue with metal substitution, which include the ionization potential and possibly the geometry and partial charges. The final goal is to develop a set of equilibrium geometries and partial charges for the different metal sites that can be used to study metal substitution in a given protein or to study other redox sites such as that of the Ni-Fe hydrogenases. Methods Electronic Structure Calculations. The protein redox site analogues in this study have a single 3d transition metal (Fe, Co, Ni, or Zn) tetrahedrally ligated to four methiolate ([SCH3]-) ligands arranged in the highly symmetric D2d conformation (Figure 1). In this conformation, all four thiolate ligands are equivalent by symmetry, thus increasing computational efficiency and simplifying analysis. The orientation of these ligands can be described as the equatorial conformation with respect to the principal axis of rotation of the molecule. The initial structure for each metal-substituted analogue, irrespective of oxidation state or metal identity, was based on the oxidized Cp rubredoxin crystal structure51 but was modified slightly to conform to ideal D2d symmetry: M-S, 2.29 Å; S-C, 1.80 Å, S-M-S, 109.5°; M-S-C, 105.0°; S1-M-S2-C2, 180.0°; S1M-S3/4-C3/4, (60.0°; where M ) Fe, Co, Ni, or Zn. Here, Si,j-M-Sk,l refers to all four possible S-M-S angles between either atoms i or j and either atoms k or l, while Si/j-M-Sk/l refers to only angles Si-M-Sk or Sj-M-Sl. Similarly, SiM-Sj/k-Cj/k refers to either dihedral angle Si-M-Sj-Cj or SiM-Sk-Ck. Methyl group bond lengths and angles were set to idealized geometries: C-H, 1.09 Å; X-C-H, 109.5°; M-SC-Ha, 180.0°; where X ) S or H. Each model was then geometry optimized as part of the electronic structure calculation. Both oxidized and reduced states of each model system were examined for all metals except Zn, in which only the “reduced” state was determined. Metals were formally +3 in the oxidized and +2 in the reduced states, which correspond to net charges of -1 and -2, respectively, of the analogue. The formal d counts for Fe(III), Co(III), and Ni(III) are d5, d6, and d7, respectively, while Fe(II), Co(II), and Ni(II) are d6, d7, and d8. For Zn(II), all of the d orbitals are full. The energy levels of the metal-centered d orbitals are roughly equivalent to that expected for tetrahedral coordination, which means that the z2 and x2 - y2 orbitals (e set) are lower in energy than the xy, xz, and yz orbitals (t2 set). Because of the D2d symmetry of the structure, inequivalent π donation of the sulfurs lifts the tetrahedral degeneracy such that the energies of the d orbitals are z2 < x2 - y2 < xy < xz ) yz.
Beck et al. The spin state of the metal is an important consideration since, although the coordination of the [1Fe] site is tetrahedral with a high-spin state, the possiblity of distortion leading to a lowspin state should be considered. For all of the metals, the highspin state was examined to be consistent with most experimental data. While tetrahedral transition metal complexes are generally high spin, in this case there is experimental evidence that there may be considerable distortion of the tetrahedron and thus the possibility of a low-spin state. Experiments on rubredoxin and substituted rubredoxins unequivocally assign the high-spin state in the iron52 and cobalt24 systems. Magnetic circular dichroism (MCD) spectra also show Ni(II)-rubredoxin as a high-spin triplet (S ) 1).24,29 The near-infrared/visible/UV absorption24 and MCD29 spectra of Ni(II)-rubredoxin are both consistent with the tetragonally elongated tetrahedral thiolate coordination for Ni(II) found in crystal structures of [Ni(SPh)4]2-.23 However, EPR indicates oxidized Ni-rubredoxin has a low spin (S ) 1/ ) Ni(III), probably due to a tetrahedral environment that is 2 highly distorted toward a square planar, trigonal bipyramidal, or tetragonally compressed octahedral geometry, since a pure tetrahedral symmetry of the Ni(III) ion would lead to a highspin (S ) 3/2) state.27,28 The spin state of the oxidized analogue is not known, and there may be differences between the redox sites in analogues and in proteins. For instance, recent experimental work on metal substitution in [MFe3S4(SR)4]2-/3analogues shows that, although the trends in spin states and ionization potentials of analogues are consistent upon substitution of eight different metals (not including Ni(III)), they do not always agree with reported values of substituted proteins.34 Thus, given experimental evidence that most analogues are in the high-spin state, Ni(III) is also examined in the high-spin (S ) 3/2) state. This use of all high-spin states is also in agreement with recent density functional theory (DFT) calculations of transition metal chlorides, [MCl4]2-, for the same metals as studied here.53 The interaction diagram (Figure 2) for the initial oxidized structure was constructed from extended Huckel calculations,54 which used a modified Wolfsberg-Helmholtz formula55 with standard parameters taken from the literature.56 The ab initio electronic structure calculations are similar to those reported for the [1Fe] core7 and were performed using the program GAMESS.57,58 The high-spin state necessitates the use of unrestricted Hartree-Fock (UHF) calculations. The basis sets and inclusion of electron correlation were examined in a previous study;7 the set used (referred to as basis set II in that work) at the UHF level was deemed to be sufficiently accurate for examining ionization potentials and better than the use of an effective core potential (ECP) on the metal with electron correlation included by second-order perturbation theory (MP2). The metal basis sets were the full core primitive sets of Wachters,59 which are approximately triple-ζ in the valence region, as modified and implemented within GAMESS.57,60 The basis set on the sulfurs utilized an ECP, which removes the 1s, 2s, and 2p electrons from explicit consideration. In addition, the associated bases were used for the sulfur, which were contracted to give a double-ζ basis in the valence region (21/ 21). This sulfur basis was then further augmented with an additional d polarization function (ζ ) 0.6). The C atoms used the standard 3-21G basis,61 while the methyl hydrogens were represented with the smaller STO-3G basis.62,63 Redox Potentials and Ionization Potentials. Although attempts are being made to measure ionization potentials of Fe-S sites,64 the currently available experimental data that are most relevant to the ionization potentials are the redox potentials,
Metal Substitution in Analogues of Rubredoxin
J. Phys. Chem. B, Vol. 103, No. 37, 1999 8009
Figure 2. Orbital interaction diagram for Fe(SCH3)4- illustrating the electronic origins of rhombic distortions. The d-centered orbitals are shown with only the σ donor orbitals from the ligands without the π effects. The tetrahedral e set orbitals (left-hand side) are shown with no σ donors since the ligands are directed toward the angular nodes. Left-hand side shows an idealized tetrahedron, and the right-hand side shows the results of tetragonal elongation.
which are related as follows. The total free energy of a reduction reaction is related to the redox potential via the Nernst equation
∆G ) -nFE
(1)
This free energy can be decomposed approximately into an intrinsic part, ∆Gint, due to the redox site and an environmental contribution, ∆Genv, due to the surrounding protein and/or solvent
∆G ) ∆Genv + ∆Gint
(2)
The ∆Gint is composed of ∆Eint, the energy difference between the oxidized and reduced states with their respective optimized geometries or adiabatic ionization energy and an entropy term
∆Gint ) ∆Eint - T∆Sint
(3)
Furthermore, ∆Eint is simply
∆Eint ) -IP + ∆Erlx
(4)
where IP is the vertical ionization potential (IP) and ∆Erlx is the internal structural relaxation energy of the redox site upon reduction. As a first approximation, it is assumed that the intrinsic and environmental parts can be calculated separately.
The purpose here was to try to calculate differences in redox potential that would occur if different metals were substituted into the site of a given protein. Thus, the environmental contribution should remain approximately constant and the differences in redox potential upon metal substitution should correlate directly with the differences in ∆Eint. The ∆Eint is calculated here simply by the energy difference between the absolute energies of the oxidized and reduced states. In previous work,7 the ionization potential was approximated by the ∆SCF method,65 which is simply the difference between the absolute self-consistent field (SCF) energy of the oxidized and reduced states. This has been shown to be reasonably accurate when compared to experimental measurements for high-spin porphyrins.66 When used for the open-shell ironsulfur site in rubredoxin,7 it was shown that the ∆SCF of the oxidized geometry gave an underestimate of ∆Eint while the reduced geometry gave an overestimate of ∆Eint. An average of these two energies was very close to ∆Eint.7 Partial Charges. Atom-centered partial charges at the optimized geometries of all complexes in both oxidation states were calculated by both Mulliken population analysis and by fitting charges to the electrostatic potential (ESP) using the CHELPG method,67 as implemented in GAMESS. Mulliken partial charges are calculated by subtracting the number of core electrons and the Mulliken valence population from the nuclear charge. For the ESP calculations, the potential at points on a cubic grid between an outer radius of 5 Å from each atom and the van der Waals radii of each atom was fit. Breneman and Wiberg values67 were utilized for the van der Waals radii of all the atoms except the metals. For the latter, the GAMESS value of 1.8 Å for any transition metal was used because of the general lack of information available on transition metal van der Waals radii. A cubic grid point spacing of 0.25 Å had previously been found to give stable charges for the [1Fe] analog7 and therefore was utilized here. Such radii and grid spacing yielded a number of points ranging from 112 000 to 110 000 for oxidized species and 115 000 to 113 000 for reduced. The least-squares fit of the nuclear charges to the potentials constrained the total charge, the dipole moment, and the quadrupole moment to those from the electronic wave function. Results The optimized geometries of [M(SCH3)4]1-/2- with each of the four metals were calculated (Table 1). The changes in metal-sulfur (M-S) and sulfur-carbon (S-C) bond lengths with change in metal and with change in redox state compare favorably with those in known crystal structures;16,21-23,51 however, the UHF calculations predictably give absolute bond lengths that are 0.08-0.19 Å larger than known crystal structure values.7,68,69 The calculated M-S bond lengths are all found to increase by 0.17-0.19 Å upon reduction due to increased antibonding between the metal and the ligand. Among the various metals, the M-S bond lengths differ by less than 0.08 Å within either the oxidized or reduced states. The S-C bond lengths are nearly identical in both oxidation states (1.84 ( 0.01 Å) regardless of the metal and differ from crystal structures by no more than 0.08 Å. The S-M-S angles also agree well with known experimental geometries.16,21-23,51 While the average S-M-S angle is always 109.5° for both oxidation states regardless of the metal, the S1-M-S2 and, by symmetry, the S3-M-S4 angles, are significantly smaller than the tetrahedral value of 109.5°, while the other four angles are significantly larger than 109.5°. This tetragonal elongation away from the pure tetrahedron is more pronounced in the reduced state and
8010 J. Phys. Chem. B, Vol. 103, No. 37, 1999
Beck et al.
TABLE 1: Comparison of Optimized Oxidized and Reduced Geometries with Crystal Structures Fe(SCH3)4 geometric valuea
Co(SCH3)4
Ni(SCH3)4
oxidation state
QM
exptb
QM
M-S
-1
2.370
2.319
2.299
S-C
-1
1.844
1.841
1.841
S1,2-M-S3,4d
-1
112.10
2.29 2.267(2) 1.83 1.840(6) 112.3 110.7(1) 103.9 107.0(1) 109.5 109.5(1) 105.3 101.5(2)
111.40
113.06
105.68
102.51
109.49
109.54
102.08
100.75
S1/3-M-S2/4d
104.34
S-M-S average M-S-C
-1
109.58
-1
100.42
M-S
-2
2.543
S-C
-2
1.851
S1,2-M-S3,4d
-2
110.27
S1/3-M-S2/4d S-M-S average M-S-C
107.89 -2
109.47
-2
100.16
2.35(1) 2.356(5) 1.77(1) 1.835(15) 114.6(1) 108.3(2) 99.6(1) 111.7(2) 109.6(1) 109.4(2) 110.9(3) 114.6(10)
expte
2.509
QM
2.33(1)
1.851
2.481
Zn(SCH3)4
expte
QM
exptc
2.29(1)
2.467
2.35(1) 2.34(2)
1.850
1.849
111.60
116.4(2)
113.90
118.6(2)
111.64
105.29
96.3(2)
100.94
92.4(2)
105.22
109.50
109.7(2)
109.58
109.9(2)
109.50
98.99
109.9(2)
98.13
108.5(2)
99.20
1.82(3) 115.4(2) 112.1(5) 98.2(2) 104.3(2) 109.7(2) 109.5(5) 109.6(2) 102.7(30)
a Bond lengths in Å, bond angles in degrees. Listed experimental data are average values. b The upper number is from oxidized rubredoxin values51 and from reduced [Fe(SPh)4] analogue values,22 respectively. The lower number is from the oxidized and reduced [Fe(S2-o-xyl)2] analogue.16 c The upper number is from reduced [M(SPh) ] analogue values,23 while the lower number is from Zn-rubredoxin.43 d See text for definitions of 4 these angles.
TABLE 2: Mulliken Atomic Spin Populations for [M(SCH3)4]1-/2atom M3+ S C Ha Hb % ionicity of M M2+ S C Ha Hb % ionicity of M
Fe
Co
Oxidation State -1 4.6252 3.4435 0.0885 0.1413 0.0019 -0.0076 0.0015 0.0018 0.0009 0.0018 92.5 86.1 Oxidation State -2 3.9652 2.9436 0.0094 0.0144 -0.0020 -0.0016 0.0005 0.0006 0.0004 0.0003 99.1 98.1
Ni 2.3600 0.1611 -0.0072 0.0031 0.0015 78.7
TABLE 3: Total Energies (and Differences from the Energy of Fe) in Atomic Units of [M(SCH3)4]1-/2- and [MCl4]2Complexes metal
[M(SCH3)4]1-
[M(SCH3)4]2-
[MCl4]2-a
Fe
-1459.77080 (0) -1578.66014 (-118.88934) -1704.08054 (-244.30974)
-1459.69240 (0) -1578.64305 (-118.95065) -1704.07241 (-244.38001) -1975.00273 (-515.31033)
-3104.79494 (0) -3223.84676 (-119.05182) -3349.35734 (-244.56240) -3620.42194 (-515.62700)
Co Ni Zn
1.9147 0.0225 -0.0031 0.0010 0.0005 95.7
exhibits a weak tendency to be larger with increasing metal atomic number. The atomic spin populations of the Fe, Co, and Ni methiolates (Zn has no unpaired spins) approach ideal values of +5, +4, and +3, respectively, for the oxidized state of the metals, and +4, +3, and +2, respectively, for the reduced state (Table 2). The bulk of the unpaired electron density is resident on the metal. The ratio of the actual to the ideal spin population or % ionicity shows strong ionic character, which is greater for the reduced state. In addition, the % ionicity decreases in both oxidation states with increasing atomic number. However, there is a greater decrease in % ionicity with increasing atomic number in the oxidized state. The absolute energies of [M(SCH3)4]1-/2- with each of the four metals were calculated and the [M(SCH3)4]2- energies compared to absolute energy results for a DFT quantum mechanical study of [MCl4]2- metal tetrahalides53 (Table 3). The corresponding reduction energy differences, ∆Eint, for [M(SCH3)4]1-/2- as well as the redox potentials for heterometal-
a
From ref 53.
TABLE 4: Energies of Reduction, ∆Eint, for [M(SCH3)4]1-/2and NegatiWe of Redox Potentials, E, for [MFe3S4(SR)4]2-/3metal Fe Co Ni
[M(SCH3)4]1-/2- [MFe3S4(SR)4]2-/3- a [MFe3S4(SR)4]2-/3∆Eint (meV) -E ( 10 (mV vs NHE) -E (mV vs SCE) 2140 470 220
450 245 360
b
1200 1020 900
a Ferredoxin redox potential values from ref 36. b Analogue redox potential values from ref 34.
substituted thiocubanes, [MFe3S4(SR)4]2-/3-, from proteins and analogues34,36 are also shown (Table 4). The ∆Eint decreases from 2140 mV for Fe to 470 mV for Co to 220 mV for Ni. Zn has only one relevant oxidation state in this complex so no ∆SCF energies were evaluated. Both Mulliken (Table 5) and ESP fitted (Table 6) atomcentered partial charges were calculated for the analogue with each of the four metals for both oxidation states except for Zn(II). Irrespective of the method of charge determination or metal, there is an increase in positive charge on the metal upon reduction with the exception of the ESP charges of Fe. Upon reduction, the metal charges change by at most 0.17e for Mulliken charges and at most 0.03e for ESP fitted charges.
Metal Substitution in Analogues of Rubredoxin
J. Phys. Chem. B, Vol. 103, No. 37, 1999 8011
TABLE 5: Mulliken Population Chargesa for [M(SCH3)4]1-/2atom
Fe
Co
Ni
Zn
M3+ S C Ha Hb
0.723 -0.2754 -0.5364 0.11445 0.1333
oxidation state -1 0.63305 0.528 -0.25175 -0.2254 -0.53673 -0.5363 0.11387 0.1141 0.13355 0.1328
M2+ S C Ha Hb
0.735 -0.4651 -0.5505 0.08925 0.1213
oxidation state -2 0.725 0.693 -0.4644 -0.4540 -0.5503 -0.5504 0.08885 0.08845 0.1223 0.12135
0.671 -0.4562 -0.5461 0.08885 0.12285
a Reported to three decimal places of accuracy; figures beyond are to ensure an integer net charge.
TABLE 6: ESP (CHELPG) Chargesa for [M(SCH3)4]1-/2atom
Fe
Co
Ni
M3+ S C Ha 2*Hb number of points
oxidation state -1 1.328 1.298 -0.6781 -0.6755 0.0879 0.1036 0.0575 0.0519 -0.02465 -0.02725 111912 111492
1.244 -0.6514 0.0718 0.0608 -0.0211 110524
M2+ S C Ha 2*Hb number of points
oxidation state -2 1.297 1.315 -0.8671 -0.8557 0.0426 -0.0141 0.03269 0.04885 -0.01622 -0.00390 115390 114054
1.251 -0.8255 -0.0854 0.06491 0.01662 113032
Zn
1.258 -0.838 -0.034 0.0521 0.0027 113080
a Reported to three decimal places of accuracy; figures beyond are to ensure an integer net charge.
Instead, the bulk of the electron density that is added to the system upon reduction resides on the more electronegative sulfurs, which become further separated in space as the ironsulfur bond lengths expand upon reduction. There is also a counterintuitive increase in positive charge at the metal, which is discussed further below. Perhaps the most surprising result is that neither method shows significant variation in the partial charges among the metals. The largest variation of charge on any atom with metal substitution occurs for the metal when the analogue is in the oxidized state, with much less variation occurring in the reduced state. This is especially true for the Mulliken charges, which vary 0.2e in the oxidized state but only 0.06e in the reduced. In contrast, the ESP fitted partial charges are barely changed at all with regard to the metal present, varying by 0.08e in the oxidized state and 0.06e in the reduced. Discussion The effects of metal substitutions on the properties of metal tetramethiolate complexes, [M(SCH3)4]2-/1-, where M ) Fe, Co, Ni, Zn, are discussed here. First, the general properties of the analogues are discussed and compared with experiment. Next, variations in the calculated properties observed upon metal substitution are discussed. Finally, the implications of the observed perturbations in analogue properties with metal substitution to the structure, redox potentials, and simulations of these proteins are discussed. General Properties of the [1M] Sites. A comparison of the optimized geometries to the experimental geometries of known compounds indicates these calculations are generally accurate
(Table 1). The geometries of [M(SCH3)4]1-/2- from our ab initio optimizations show generally good agreement with both rubredoxin analogues16,22,23 and oxidized C. pasteurianum (Cp) rubredoxin,51 although the differences in the sulfur ligand substituent groups between the experimental structures and the compounds studied here will lead to certain systematic deviations. The optimized geometries show an increase in bond length upon reduction as seen in experiment, which primarily serve to reduce the metal-ligand antibonding that occurs upon reduction but secondarily to reduce the ligand-ligand charge repulsion due to the increased charges on the sulfurs upon reduction. The ab initio M-S bond lengths are too long by 0.1 to 0.2 Å relative to experiment as is common with UHF calculations.7,68,69 The ab initio S-C bond lengths in the Fe complexes are less than 0.02 Å longer than those from the oxidized rubredoxin and [Fe(S2-o-xyl)2]1-/2- structures and are about 0.08 Å longer than those from the [Fe(SPh)4]2- structures. The shorter bond lengths in the [Fe(SPh)4]2- analogue are in part explained by the increased π character of the aryl C-S bonds relative to alkyl C-S bonds;22 thus, it is important to reconsider the ligand independence of the electronic structure suggested by Noodleman and co-workers.50 Perhaps the most significant deviation from experiment is that while an increase in bond length is seen in both, the ab initio M-S bonds expand 0.17-0.19 Å upon reduction, whereas those of the [Fe(S2-o-xyl)2]1-/2- analogue only expand by 0.09 Å. This is due to the lack of electron correlation, an important factor when evaluating metal-ligand charge-transfer excitations, since the Hartree-Fock method underestimates the covalent character of transition metal-ligand bonds.68 Another important geometric feature is the tetragonal (D2d) elongation of the sulfur ligands away from an ideal tetrahedral coordination. A measure of this is the six S-M-S angles for [M(SCH3)4]2-, which would have a value of 109.5° for ideal tetrahedral coordination but are distorted so that the four angles S1,2-M-S3,4 are larger and the two angles S1/3-M-S2/4 are smaller. The ab initio S-M-S angles are ∼4° closer to tetrahedral than experimental values for [M(SPh)4]2-; the larger values in the latter can in part be explained by a steric clash between the bulky phenyl substituent groups. The ab initio [Fe(SCH3)4]1- S-Fe-S values are even closer to those from oxidized Cp rubredoxin, showing less than 0.5° variance. On the other hand, experimental values for the [Fe(S2-o-xyl)2]1-/2analogue S-M-S angles show much less distortion away from tetrahedral relative to the ab initio values for [Fe(SCH3)4]1-/2-. In fact, what angular distortion is shown appears to be away from S4 along the Fe-S4 axis. The apparent lack of S-M-S distortion with such bidentate ligands may be an artifact, as [Fe(S2C4O2)2]2- demonstrates larger distortions than the monodentate ligands (data not shown).22 Regardless, both experimental and theoretical geometries exhibit tetragonal elongation away from pure tetrahedral symmetry, regardless of oxidation state or ligand substituent groups. The electronic origins of the tetragonal elongation can be interpreted by examining the orbital interaction diagram (Figure 2). The most important interactions involve the d orbitals of the metal with the pseudo-σ donor orbitals from the ligands.70 As the two angles bisected by the z axis become more acute, the degeneracy of the tetrahedral orbitals is lifted. Thus, the tetrahedral t2 set (left-hand side of Figure 2) is split into a b2 xy orbital and an e set composed of xz and yz. The e set is highest in energy, presumably because the σ donors move away from the angular nodes of the d orbitals as the molecule distorts. In contrast, the b2 metal-centered orbital is lowered in energy,
8012 J. Phys. Chem. B, Vol. 103, No. 37, 1999 presumably because the metal-σ antibonding is reduced. The tetrahedral e set (left-hand side of Figure 2) is also split as the molecule distorts to D2d symmetry. The x2 - y2 orbital remains rigorously nonbonding throughout the distortion and its energy is unaffected. On the other hand, the energy of the z2 orbital of a1 symmetry is perturbed with the geometric distortion. Overlap between the z2 orbital and the ligand-centered donors increases as the angles become more acute. While the metal-ligand interaction is primarily antibonding for the z2 orbital, metal s is also mixed into the metal-centered orbital in a bonding fashion to the ligand lone pairs; this three-orbital mixing (shown at the bottom of Figure 2) results in an a1 molecular orbital that is more pronounced in the torus. The degree of distortion will be different based on electron counts. A high-spin d1 or d6 system is likely to be distorted since the a1 orbital will tend to be stabilized in D2d symmetry. Also, a high-spin d7 system, in which 1a1 and 1b1 are both doubly occupied, is likely to be distorted since b1 is rigorously nonbonding and its energy would be unaffected by the distortion. However, for a d7 system, there is a possibility for a low-spin state. If the splitting is large, b2 and e have a large energy gap, which may lead to a favorable spin-paired interaction in the b2 orbital. Thus, a larger distortion in D2d symmetry favors a lowspin system as seen for oxidized Ni. This could serve as an alternate explanation for the EPR signal for the Ni(III) as opposed to the flattening of the tetrahedron toward square planar. Adding an extra electron would then favor the high-spin system seen for reduced Ni, which would tend to be less distorted since the e set is now doubly occupied. For the spherically symmetric metal in the high-spin d5 systems of Fe(III), there is still distortion toward D2d symmetry. This indicates that π effects may also be involved in the distortion, although this would be a secondary effect. Variations with Metal Substitution. The calculated geometries of the different [M(SCH3)4] analogues are similar regardless of metal. The M-S bond lengths differ by less than 0.08 Å within either oxidation state, while the S-C bond lengths show less than 0.01 Å variation regardless of oxidation state. The changes observed, though small, have implications. Shorter metal-ligand bonds as the metal is changed from Fe to Co to Ni suggest better donating ability of the ligand and thus a decreased effective nuclear charge for the metal.70 The change upon reduction is roughly the same for all metals; however, the relative change is greatest for nickel and smallest for iron. Metal substitution seems to primarily affect the geometries by altering the bond angles and, in particular, the S-M-S angle distortion away from tetrahedral symmetry. Of the three transition metals examined in both states (Zn having full d-orbitals), the tetragonal elongation not only is more pronounced in the reduced state, but generally increases with increasing d orbital occupancy. Such a finding supports our model of the origins of the distortion (Figure 2). Somewhat surprising is that the atom-centered partial charges show remarkably little variation. However, this lack of atomic charge variation does agree with the homogeneity of the atomic electronegativities of the metals used in this study (1.83 for Fe, 1.88 for Co, 1.91 for Ni, 1.65 for Zn71). This lack of variation in the charge distribution is less surprising when one considers that the exponents used on the metal to describe the diffusivity in the Wachter’s set are all similar.59 These results are important because classical molecular dynamics simulations indicate that the rubredoxin structure is sensitive to the values of the partial
Beck et al. charges when they are in the range of those found for the [1Fe] site.47 These issues will be discussed further at the end of this section. The two methods for calculating atomic partial charges, Mulliken population analysis and the CHELPG method, provide different pictures of the partitioning of charge over the atoms even though the same wave functions are used. Since the CHELPG method fits a set of partial charges located at the atoms centers to reproduce the electrostatic potential calculated from the wave functions in a shell around each molecule, it samples points far enough away from the metal such that significant variations in the electrostatic potentials due to electrons occupying the various metal-centered orbitals may simply be averaged out. Moreover, while ESP fitted charges have been found to have reasonable magnitudes, the values for buried atoms such as the metals in the [M(SCH3)4] site are more difficult to calculate accurately.67,72-74 This effect should be more pronounced for the oxidized states since the bond lengths are shorter, which would tend to support the idea that the greater difference in the CHELPG charges upon metal substitution for the oxidized states than for the reduced states is significant. However, the relative change in the bond lengths is greater in the oxidized than the reduced state, so the greater variations in charges may simply reflect the greater variations in bond lengths. On the other hand, the Mulliken charges are a direct measure of the electron density since the Mulliken population analysis partitions half of the overlap probability density to each of the atomic orbitals in a given bond, whereas the values of the individual ESP charges are less directly related to the electron density since the ESP charges are fits to the potentials due to the electron density constrained to the total charge, dipole moment, and quadrupole moment of the electronic wave function. The Mulliken charges show greater variation as the metal in the redox site changes than the ESP charges, but the changes are still small. Also, although both have the same general trends, the Mulliken charges correlate better with increasing number of electrons. However, in terms of determining charges for classical mechanics calculations, the magnitudes of the ESP charges are probably more accurate than Mulliken charges since the latter show a strong basis set dependency and fail to reproduce higher order moments.72,75 Thus, we assume that the Mulliken charges better reflect changes in the electron density upon metal substitution, even though the magnitudes of the charges may not be as accurate. Therefore, physical changes will be described in terms of Mulliken charges Although the variation is small, the metal charge decreases upon substituting the Fe with Co and Ni, most noticeably in the oxidized state. This reflects the increasing electronegativity of the metal, except for Zn (1.83 for Fe, 1.88 for Co, 1.91 for Ni, 1.65 for Zn71). For a given metal site, the charges on the constituent atoms of the redox site are expected to become more negative upon reduction. The sulfur donors become more negative with the Mulliken charge decreasing from ∼-0.25 to ∼-0.45. However, for the metal in the reduced state, the M-S bonds are longer, have more antibonding character, and are less covalent, which would tend to increase the positive charge on the metal so that there is a balance of these two effects. The charge on the metal increases upon reduction for Fe and increases even more upon substitution of the Fe by Co and Ni, which is consistent with the relative change in the bond lengths for the oxidized systems being greater than for the reduced upon substitution. The near ideal atomic spins of the metals in these complexes (Table 2) indicates that the systems are very ionic, with the
Metal Substitution in Analogues of Rubredoxin bulk of the unpaired electron density resident in the metal. The increase in % ionicity upon reduction correlates with the increase in metal-ligand antibonding and longer M-S bond lengths. Thus, the electron density is dispersed from the internuclear zone, thereby resulting in a more localized bonding picture with the electrons residing on the atoms. There is a small decrease in the % ionicity of the reduced states as the metal is changed from Fe to Ni and, like the total density discussed in terms of partial charges, the oxidized states show greater variation. This in part reflects the increasing electronegativity of the metal (1.83 for Fe, 1.88 for Co, 1.91 for Ni71). In addition, the shorter the M-S bond length, the greater the tendency for the unpaired density to spread out over the molecular framework. This means there is an increase of unpaired density on the ligands upon metal substitution, the effects being more pronounced in the oxidized state (Table 2). As a result, the change in % ionicity between oxidized and reduced states (∆ ionicity) increases markedly from Fe to Ni. Of course, these can be altered by the protein, and the recent hybrid DFT study of hyperfine NMR chemical shifts showed that the protein environment affected the Fermi contact spin densities of the atoms of the redox site of rubredoxins.76 For a ligated metal site, Holm has described three possible ways the ligands may affect the ionization potential.70 The first way is that the energy of the redox-active orbital will be strongly affected by the ligand field geometry. As the energy of this orbital increases, the site becomes easier to oxidize. The second way is that the effective nuclear charge (Z′eff) will be affected by the donor propensity of the ligands. The energy of the d orbital manifold would tend to increase as Z′eff is lowered, which makes the site easier to oxidize. The third way is electronic relaxation, which is manifested by changes in electronic structure. A greater change in electronic structure upon oxidation makes the site easier to oxidize. In a somewhat different approach to altering the ionization potential, a recent study found that the substitution of serines for cysteinyl ligands in rubredoxin changes the metal-ligand bond length by 0.4 Å and the redox potential by as much as 200 mV.49 The [M(SCH3)4]2- analogues show the same relative variation in absolute energies with metal identity as do those from a DFT calculation of [MCl4]2-, which used the Gaussian-type basis sets: (63321/531/41) and (6321/521/1) for transition metals and chlorine with the uncontracted basis sets of (10/5/5) and (9/4/ 4), respectively,53 differing by less than 0.13 au (Table 3). The smaller decrease in energy in going from Fe to Zn with [M(SCH3)4]2- relative to [MCl4]2- may be due to the better charge-donating ability of the thiolate relative to chloride ligands (with relative Pauling electronegativities of 2.58 and 3.16,71 respectively). Since this would tend to lower Z′eff, the energy of the HOMO would tend to increase. Unfortunately, no calculations were made of [MCl4]- in that study so no comparisons of ionization potentials can be made. The calculated results show that ∆Eint is large and positive (Table 4) as the overall charge changes from -1 to -2, as might be expected for adding an electron to an already negatively charged system. Moreover, ∆Eint becomes more favorable (i.e., smaller) as the atomic number increases as the metal is changed from Fe to Ni. This correlates with experimental results in which the d-d ligand field stabilization energies for both Co and Ni tetrathiolates are smaller than those for similar Fe compounds,16,77 which implies that the redox-active orbital is higher in energy. Also, the trends in ∆Eint are consistent with the trends in the third IP of the elemental ions (Table 7) in that even though the energy necessary to reduce M(III) (i.e., the negative of the
J. Phys. Chem. B, Vol. 103, No. 37, 1999 8013 TABLE 7: Energies of Reduction, ∆Eint, and NegatiWe of Third Ionization Potentials, IP, of Metal Ions ∆Eint (meV) -first IPa (meV) -second IPa (meV) -third IPa (meV) (∆Eint (-[first IP (-[second IP (-[third IP first IP(Fe)]) second IP(Fe)]) third IP(Fe)]) metal ∆Eint(Fe)) Fe Co Ni
2140 (0) 470 (-1670) 220 (-1920)
Zn a
-7870 (0) -7860 (70) -7635 (235) -9394 (-1524)
-16180 (0) -17050 (-870) -18150 (-1970) -17960 (-1780)
-30651 (0) -33500 (-2849) -35170 (-4519) -39722 (-9071)
From ref 83.
third IP) is large and negative, the energy becomes more favorable (i.e., more negative) as the atomic number increases. For both ∆Eint and the third IP of the elemental ions, this can be interpreted in terms of the interaction of the added electron with the effective nuclear charge. As a rough estimate, the effective nuclear charges using Slater’s rules78 for a 3d electron are 6.25, 6.90, and 7.55 for Fe, Co, and Ni, respectively, which indicates an increase in effective nuclear charge with atomic number. The increasing electronegativity with atomic number for Fe to Ni (Table 7) also indicates the increasing attraction for the electron. However, the ligands obviously play a role since the energy does not become as favorable with increasing atomic number for ∆Eint of the ligated metal sites as for the third IP of the elemental ions and becomes even less so when relaxation effects (eq 4) are excluded from ∆Eint.7 In addition, the relative change in ∆Eint between Fe and Co is much larger than between Co and Ni, unlike the relative changes in the third IP. Considering Holm’s analysis, the general decrease in favorability of reduction as the atomic number increases for the liganded metal sites over the elemental ions may be attributed to the ligand effects on Z′eff. Extensive charge donation of the thiolate ligand will tend to reduce Z′eff relative to the actual nuclear charge, which would make the site harder to reduce. There are several indications that the ligand donation is becoming greater and thus Z′eff is increasing more slowly with increasing atomic number for the ligated versus elemental metal. From the Mulliken population analysis, the valence populations are about 7.28, 8.37, and 9.47 electrons for Fe, Co, and Ni, respectively, in the oxidized analogues and about 7.27, 8.28, and 9.31 electrons for Fe, Co, and Ni, respectively, in the reduced analogues (Table 5) and the % ionicity decreases with atomic number (Table 2); furthermore, the metal-sulfur bond lengths decrease with atomic number (Table 1). Moreover, considering that the valence population of the reduced state is closer to a metal ion with a formal charge of +1, the change in ∆Eint is closer that of the second IP of the elemental ions (Table 7). However, these are relatively uniform changes with atomic number and provide an interpretation only for the general decrease in favorability for the ligated versus elemental metal. The source of the larger change in ∆Eint for the Co relative to Fe site than for the Ni relative to Co site may lie in the oxidized Fe analogue. Whereas the Fe, Co, and Ni sites in the reduced state and the Co and Ni sites in the oxidized state all show increasing tetragonal distortion with atomic number, the Fe is more distorted than the Co site in the oxidized state. Moreover, a large degree of electronic relaxation has been noted for these sites with an Fe, since while the high-spin Fe(II) complex is normal, the high-spin Fe(III) complex has an inverted bonding description.70 This would tend to make the site harder to reduce and thus to increase ∆Eint of the Fe over the Co and Ni sites, as has been found in this study.
8014 J. Phys. Chem. B, Vol. 103, No. 37, 1999 The ∆Eint for the different metals indicates that the redox potential of Fe should be much more negative than Co and Ni. A direct comparison of the ∆Eint in this study to the experimental redox potentials of other metal tetrathiolates is problematic as the contribution of solvent and/or protein matrix has not been well established nor are solvent effects included in this work. Our calculations also ignore possible contributions from lowspin states also discussed in the Introduction. The sensitivity of ∆Eint to the R group of the ligands also bears re-examination beyond the early studies of Noodleman et al.50 given the recent studies of Stavrev and Zerner.79 Moreover, no experimental data for redox potentials of different metals in the [1M] core are available; therefore, trends in the ab initio data were compared to values for the [4Fe-4S] core. Substitution of Co and Ni for one Fe in the [4Fe-4S] thiocubanes of D. gigas ferredoxin36 and [MFe3S4(SR)4]2-/3- analogues34 increases the potential (Table 4) but by inconsistent amounts, probably due to the protein environment. Recalling that the redox potential E is proportional to -∆Eint plus an additive term (eq 1-4), the analogue data is correlated with our findings, although it is not clear that the [4Fe-4S] core data should be consistent with the [1M] core. Implications for Protein Structure, Redox Potentials, and Simulations. The calculated trends in geometry, partial charge, and atomic spin overall indicate that the identity of the metal in the [1M] site will not affect the protein structure. The small variation of M-S bond lengths, S-M-S angles, and M-S-C angles with metal substitution in either oxidation state, as well as the negligible variation in S-C bond length, suggests that one should expect little distortion of the surrounding protein structure upon metal substitution by changes in the redox site geometry. Moreover, the relatively small variation in the partial charges with metal substitution indicates that the protein will not change much in response to changes in the partial charges. This indicates that metal substitution is unlikely to produce the large differences in surrounding protein and solvation energy that were observed by Yelle et al. in simulations with varying partial charges for the redox site.47 Such a prediction is consistent with crystallographic data showing that the RMS difference between Fe(III) and Zn(II) rubredoxins43 is only 0.23 Å for backbone atoms and only 0.31 Å for all atoms47 despite the difference in oxidation state, although the perturbations are larger nearer the redox site and differences in surface charged side chains occurred. In addition, a very recent X-ray crystallographic study shows that the RMS differences for the CR atoms between Fe(III), In(III), and Ga(III) desulforedoxins are from 0.12 and 0.15 Å, between Fe(III), In(III), or Ga(III) and Cd(II) or Hg(II) desulforedoxins are from 0.29 and 0.32 Å, and between Cd(II) and Hg(II) is 0.20 Å.31 Again, for the larger metals ions in the 2+ oxidation state, perturbations near the redox site were more significant and differences in surface charged side chains occurred. Thus, our results are consistent with the general assumption that studies of metal-substituted [1M] proteins can be interpreted assuming little change in the protein structure. Our results also predict trends for the reduction energies of the [1M] sites that indicate that there should be significant differences in the redox potentials, as to be expected, but that these differences may not follow those predicted for the [4Fe4S] sites, which again may be expected. The lack of structural perturbation described above also indicates that the differences in redox potential will most likely be due mainly to the differences in the reduction energies of the [1M] site alone, with little contribution from changes in the electrostatic environment of the protein. Thus, our trends in the reduction energies of the [1M] analogue should correlate directly with the redox potentials of the [1M] proteins.
Beck et al. It is also important to note that while the metal may not significantly perturb the protein, the protein may hold the cysteinyl ligands in a nonideal geometry thus favoring certain spin states in certain metals such as appears to be the case in the spectroscopy of Ni-substituted rubredoxin.24,27 The idea of the protein straining the metal into an energized or entatic state has been proposed for the zinc in carbonic anhydrase or copper in the blue-copper proteins80 and more recently disputed in the blue copper proteins.81,82 However, the crystallographic results on metal substitution in desulforedoxin, which did not include Ni(II), indicates that the geometry of the metal site is being enforced by the protein.31 In Ni-rubredoxin, the protein appears to accentuate the rhombic distortion observed in other [M(SR)4] systems such that an equilibrium between high- and low-spin states may exist.27-29 Such a shift in the Ni(III) spin state in the protein should also coincide with a raised redox potential with respect to the other metals27 relative to that of Ni-substituted analogues with respect to the other metals. Our results for the geometries and partial charges also will be used for molecular dynamics simulation studies of metal substitution in the [1Fe] proteins and of proteins such as the Ni-Fe hydrogenase that have other [1M] sites. However, for studies of metal substitution, our examination of the magnitude of these differences indicates that the protein structures will not change much so that the emphasis of such studies may focus on questions such as the entatic nature of the substitutions. Conclusions Ab initio quantum mechanical calculations were performed on metal tetramethiolate complexes, [M(SCH3)4]2-/1-, where M ) Fe, Co, Ni, Zn. Metal substitution leads to only minor variations in geometry, atomic spin, and atom-centered partial charges with significant variation only in the energies of reduction. Given the difficulties of engineering protein contributions to give predictable redox potential changes, our findings indicate that variation of the metal identity in Fe-S proteins may be a less problematic method for engineering redox potential properties. Moreover, our findings support the idea that experiments involving metal substitution may be interpreted under the assumption that the protein structure will be relatively invariant. Finally, our geometries and partial charges can now be used for computer simulation studies of metal substitution in the [1Fe] site and for proteins with other metals in the [1M] site. Acknowledgment. The authors thank Dr. Thomas A. Albright and Dr. James K. Hurst for their helpful comments. This work was supported by a grant from the National Institutes of Health (GM45303). Computation was performed on a 16processor IBM SP2 (9076) funded by the National Science Foundation (BIR-9512538) and Washington State University. Additional computer time was provided by the Pittsburgh Supercomputing Center (sponsored by the National Science Foundation, MCB950005P), and the Maui High Performance Computing Center (sponsored in part by the Phillips Laboratory, Air Force Materiel Command, USAF, under cooperative agreement number F29601-93-2-0001). The views and conclusions contained in this document are those of the authors and should not be interpreted as necessarily representing the official policies or endorsements, either expressed or implied, of Phillips Laboratory or the U.S. Government. We also thank the VADMS Laboratory at WSU for additional computational resources. References and Notes (1) Blake, P. R.; Park, J.-B.; Zhou, Z. H.; Hare, D. R.; Adams, M. W. W.; Summers, M. F. Protein Sci. 1992, 1, 1508.
Metal Substitution in Analogues of Rubredoxin (2) Zeng, Q.; Smith, E. T.; Kurtz, D. M.; Scott, R. A. Inorg. Chim. Acta 1996, 242, 245. (3) Shen, B.; Jollie, D. R.; Stout, C. D.; Diller, T. C.; Armstrong, F. A.; Gorst, C. M.; La Mar, G. N.; Stephens, P. J.; Burgess, B. K. J. Biol. Chem. 1994, 269, 8564. (4) Guckert, J. A.; Lowery, M. D.; Solomon, E. I. J. Am. Chem. Soc. 1995, 117, 2817. (5) Olsson, M. H. M.; Ryde, U.; Roos, B. O.; Pierloot, K. J. Biol. Inorg. Chem. 1998, 3, 109. (6) Mouesca, J. M.; Chen, J. L.; Noodleman, L.; Bashford, D.; Case, D. A. J. Am. Chem. Soc. 1994, 116, 11898. (7) Koerner, J. B.; Ichiye, T. J. Phys. Chem. B 1997, 101, 3633. (8) Hoops, S. C.; Anderson, K. W.; Merz, K. M., Jr. J. Am. Chem. Soc. 1991, 113, 8262. (9) Vedani, A.; Huhta, D. W. J. Am. Chem. Soc. 1990, 112, 4759. (10) Shen, J.; Wong, C. F.; Subramaniam, S.; Albright, T. A.; McCammon, J. A. J. Comput. Chem. 1990, 11, 346. (11) Lovenberg, W.; Sobel, B. Proc. Natl. Acad. Sci. U.S.A. 1965, 54, 193. (12) Moura, I.; Brushci, M.; LeGall, J.; Moura, J. J. G.; Xavier, A. V. Biochem. Biophys. Res. Commun. 1977, 75, 1037. (13) Moura, I.; Tavares, P.; Ravi, N. Methods Enzymol. 1994, 243, 216. (14) Moura, I.; Pereira, A.; Tavares, P.; Moura, J. J. G. In AdVances in Inorganic Chemistry; Cammack, R., Ed.; Academic Press: San Diego, 1999; Vol. 47, p 361. (15) Hagen, K. S.; Holm, R. H. J. Am. Chem. Soc. 1982, 104, 5496. (16) Lane, R. W.; Ibers, J. A.; Frankel, R. B.; Papaefthymiou, G. C.; Holm, R. H. J. Am. Chem. Soc. 1977, 99, 84. (17) Laskowski, E. J.; Frankel, R. B.; Gillum, W. O.; Papaefthymiou, G. C.; Renaud, J.; Ibers, J. A.; Holm, R. H. J. Am. Chem. Soc. 1978, 100, 5322. (18) Holm, R. H.; Ibers, J. A. In Iron-Sulfur Proteins; Lovenberg, W., Ed.; Academic Press: New York, 1977; Vol. 3, p 205. (19) Holm, R. H. In AdVances in Inorganic Chemistry; Cammack, R., Ed.; Academic Press: San Diego, 1992; Vol. 38, p 1. (20) Werth, M. T.; Kurtz, D. M.; Howes, B. D.; Huynh, B. H. Inorg. Chem. 1989, 28, 1357. (21) Coucouvanis, D.; Swenson, D.; Baenziger, N. C.; Holah, D. G.; Kostikas, A.; Simopoulos, A.; Petrouleas, V. J. Am. Chem. Soc. 1976, 98, 5721. (22) Coucouvanis, D.; Swenson, D.; Baenzinger, N. C.; Murphy, C.; Holah, D. G.; Sfarnas, N.; Simopoulos, A.; Kostikas, A. J. Am. Chem. Soc. 1981, 103, 3350. (23) Swenson, D.; Baenziger, N. C.; Coucouvanis, D. J. Am. Chem. Soc. 1978, 100, 1932. (24) Moura, I.; Teixeira, M.; LeGall, J.; Moura, J. J. G. J. Inorg. Biochem. 1991, 44, 127. (25) Poutney, D. L.; Henehan, C. J.; Vasak, M. Protein Sci. 1995, 4, 1571. (26) Blake, P. R.; Lee, B.; Summers, M. F.; Park, J.-B.; Zhou, Z. H.; Adams, M. W. W. New J. Chem. 1994, 18, 387. (27) Mus-Veteau, I.; Diaz, D.; Gracia-Mora, J.; Guigliarelli, B.; Chottard, G.; Bruschi, M. Biochim. Biophys. Acta 1991, 1060, 159. (28) Huang, Y.-H.; Park, J.-B.; Adams, M. W. W.; Johnson, M. K. Inorg. Chem. 1993, 32, 375. (29) Kowal, A. T.; Zambrano, I. C.; Moura, I.; Moura, J. J. G.; LeGall, J.; Johnson, M. K. Inorg. Chem. 1988, 27, 1162. (30) Archer, M.; Huber, R.; Tavares, P.; Moura, I.; Moura, J. J. G.; Carrondo, M. A.; Sieker, L. C.; LeGall, J.; Roma˜o, M. J. J. Mol. Biol. 1995, 251, 690. (31) Archer, M.; Carvalho, A. L.; Teixeira, S.; Moura, I.; Moura, J. J. G.; Rusnak, F.; Roma˜o, M. J. Protein Sci. 1999, 8, 1536. (32) Saint-Martin, P.; Lespinat, P. A.; Fauque, G.; Berlier, Y.; LeGall, J.; Moura, I.; Teixeira, M.; Xavier, A. V.; Moura, J. J. G. Proc. Natl. Acad. Sci. U.S.A. 1988, 85, 9378. (33) Volbeda, A.; Charon, M. H.; Piras, C.; Hatchikian, E. C.; Frey, M.; Fontecilla-Camps, J. C. Nature 1995, 373, 580. (34) Zhou, J.; Raebiger, J. W.; Crawford, C. A.; Holm, R. H. J. Am. Chem. Soc. 1997, 119, 6242. (35) Moura, I.; Moura, J. J. G.; Mu¨nck, E.; Papaefthymiou, V.; LeGall, J. J. Am. Chem. Soc. 1986, 108, 349. (36) Moreno, C.; Macedo, A. L.; Moura, I.; LeGall, J.; Moura, J. J. G. J. Inorg. Chem. 1994, 53, 219. (37) Finnegan, M. G.; Conover, R. C.; Park, J.-B.; Zhou, Z. H.; Adams, M. W. W.; Johnson, M. K. Inorg. Chem. 1995, 34, 5358. (38) Srivastava, K. K. P.; Surerus, K. K.; Conover, R. C.; Johnson, M. K.; Park, J.-B.; Adams, M. W. W.; Mu¨nck, E. Inorg. Chem. 1993, 32, 927. (39) Surerus, K. K.; Mu¨nck, E.; Moura, I.; Moura, J. J. G.; LeGall, J. J. Am. Chem. Soc. 1987, 109, 3805. (40) Butt, J. N.; Armstrong, F. A.; Breton, J.; George, S. J.; Thomson, A. J.; Hatchikian, E. C. J. Am. Chem. Soc. 1991, 113, 6663. (41) Butt, J. N.; Sucheta, A.; Armstrong, F. A.; Breton, J.; Thomson, A. J.; Hatchikian, E. C. J. Am. Chem. Soc. 1991, 113, 8948.
J. Phys. Chem. B, Vol. 103, No. 37, 1999 8015 (42) Butt, J. N.; Niles, J.; Armstrong, F. A.; Breton, J.; Thomson, A. J. Nature Struct. Biol 1994, 1, 427. (43) Dauter, Z.; Wilson, K. S.; Sieker, L. C.; Moulis, J. M.; Meyer, J. Proc. Natl. Acad. Sci. U.S.A. 1996, 93, 8836. (44) Kassner, R. J. Proc. Natl. Acad. Sci. U.S.A. 1972, 69, 2263. (45) Yang, Y.; Beck, B. W.; Shenoy, V. S.; Ichiye, T. J. Am. Chem. Soc. 1993, 115, 7439. (46) Bernstein, F. C.; Koetzle, T. F.; Williams, G. J. B.; Meyer, E. F.; Brice, M. D.; Rodgers, J. R.; Kennard, O.; Shimanouchi, T.; Tasumi, M. J. Mol. Biol. 1977, 112, 535. (47) Yelle, R. B.; Beck, B. W.; Koerner, J. B.; Sacksteder, C. A.; Ichiye, T. Proteins, in press. (48) Swartz, P. D.; Ichiye, T. Biophys. J. 1997, 73, 2733. (49) Xiao, Z.; Lavery, M. J.; Ayhan, M.; Scrofani, S. D. B.; Wilce, M. C. J.; Guss, J. M.; Tregloan, P. A.; George, G. N.; Wedd, A. G. J. Am. Chem. Soc. 1998, 120, 4135. (50) Noodleman, L.; Norman, J. G.; Osborne, J. H.; Aizman, A.; Case, D. A. J. Am. Chem. Soc. 1985, 107, 3418. (51) Watenpaugh, K. D.; Sieker, L. C.; Jensen, L. H. J. Mol. Biol. 1979, 131, 509. (52) Phillips, W. D.; Poe, M.; Weiher, J. F.; McDonald, C. C. Nature 1970, 227, 574. (53) Waizumi, K.; Masuda, H.; Einaga, H.; Fukushima, N. Bull. Chem. Soc. Jpn. 1993, 66, 3648. (54) Hoffmann, R.; Lipscomb, W. N. J. Chem. Phys. 1962, 37, 2782. (55) Ammeter, J. H.; Burgi, H.-B.; Thibeault, J. C.; Hoffmann, R. J. Am. Chem. Soc. 1978, 100, 3686. (56) Hoffman, D. M.; Hoffmann, R.; Fisel, C. R. J. Am. Chem. Soc. 1982, 104, 3858. (57) Schmidt, M. W.; Baldridge, K. K.; Boatz, J. A.; Elbert, S. T.; Gordon, M. S.; Jensen, J. H.; Koseki, S.; Matsunaga, N.; Nguyen, K. A.; Su, S.; Windus, T. L.; Dupuis, M.; Montgomery, J. A. J. Comput. Chem. 1993, 14, 1347. (58) Schmidt, M. W.; Baldridge, K. K.; Boatz, J. A.; Jensen, J. H.; Koseki, S.; Gordon, M. W.; Nguyen, K. A.; Windus, T. L.; Elbert, S. T. QCPE Bull. 1990, 10, 52. (59) Wachters, A. J. H. J. Phys. Chem. 1970, 52, 1033. (60) Rappe, A. K.; Smedley, T. A.; Goddard, W. A., III J. Phys. Chem. 1981, 85, 2607. (61) Binkley, J. S.; Pople, J. A.; Hehre, W. J. J. Am. Chem. Soc. 1980, 102, 939. (62) Collins, J. B.; Schleyer, P. v. R.; Binkley, J. S.; Pople, J. A. J. Chem. Phys. 1976, 64, 5142. (63) Hehre, W. J.; Stewart, R. F.; Pople, J. A. J. Chem. Phys. 1969, 51, 2657. (64) Wang, L.-S. Personal Communication. (65) Ohno, K. In Horizons of Quantum Chemistry; Fukui, K., Pullman, B., Eds.; D. Reidel: Dortrecht, 1980; p 245. (66) Dedieu, A.; Rohmer, M. M.; Veillard, A. In AdVances in Quantum Chemistry; Lowdin, P. O., Ed.; Academic Press: New York, 1982; Vol. 16, p 43. (67) Breneman, C. M.; Wiberg, K. B. J. Comput. Chem. 1990, 11, 361. (68) Luthi, H. P.; Siegbahn, P. E. M.; Almlof, J.; Aegri, K. F., Jr.; Heiberg, A. Chem. Phys. Lett. 1984, 111, 1. (69) Taylor, T. E.; Hall, M. B. Chem. Phys. Lett. 1985, 114, 338. (70) Holm, R. H.; Kennepohl, P.; Solomon, E. I. Chem. ReV. 1996, 96, 2239. (71) Allred, A. L. J. Inorg. Nucl. Chem. 1961, 17, 215. (72) Wiberg, K. B.; Rablen, P. R. J. Comput. Chem. 1993, 14, 1504. (73) Cornell, W. D.; Cieplak, P.; Bayly, C. I.; Gould, I. R.; Merz, K. M., Jr.; Ferguson, D. M.; Spellmeyer, D. C.; Fox, T.; Caldwell, J. W.; Kollman, P. A. J. Am. Chem. Soc. 1995, 117, 5179. (74) Francl, M.; Carey, C.; Chirlian, L. E.; Gange, D. M. J. Comput. Chem. 1996, 17, 367. (75) Chirlian, L. E.; Francl, M. M. J. Comput. Chem. 1987, 8, 894. (76) Wilkens, S. J.; Xia, B.; Weinhold, F.; Markley, J. L.; Westler, W. M. J. Am. Chem. Soc. 1998, 120, 4806. (77) Davison, A.; Switkes, E. S. Inorg. Chem. 1971, 10, 837. (78) Slater, J. C. Phys. ReV. 1930, 36. (79) Stavrev; Zerner, M. Int. J. Quantum Chem. 1995, QB Symp. 22, 155. (80) Williams, R. J. P. Eur. J. Biochem. 1995, 234, 363. (81) Ryde, U.; Olsson, M. H. M.; Pierloot, K.; Roos, B. O. J. Mol. Biol. 1996, 261, 586. (82) De Kerpel, J. O. A.; Ryde, U. Proteins 1999, 36, 157. (83) CRC Handbook of Chemistry and Physics, 68th ed.; Weast, R. C., Ed.; CRC Press: Boca Raton, FL, 1987.
