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The discovery of the Ni(Cysteine-Glycine-Cysteine)2-, Ni(CGC)2-, in the A-cluster active site of Acetyl CoA Synthase has prompted the synthesis of many small molecule models which employ M(N2S2) complexes as metalloligands. In vitro studies have shown that nickel incorporates into the N2S2 binding pocket even when copper is in the enzyme growth medium, while copper is preferentially taken up in the proximal site, displacing the catalytically active nickel. (Darnault, C.; Volbeda, A.; Kim, E.J.; Legrand, P.; Vernede, X.; Lindahl, P.A.; Fontecilla-Camps, J.C. Nat. Struct. Biol. 2003, 10, 271-279.) The work herein has been designed to address the chemical viability of copper(II) within the tripeptide N2S2 ligand set. To this end, a series of CuN2S2 2- complexes, the resin-bound, O-Cu(CGC)2- (A) and free Cu(CGC)2- (B) complexes, as well as Cu(ema)2- (C) and Cu(emi)2- (D) dianions, have been characterized by UV-vis, EPR, and ESI-MS spectroscopies, cyclic voltammetry (CV), and, where appropriate, x-ray diffraction studies, and compared to the NiII congeners. EPR spectroscopic results have indicated that, in frozen DMF solution, the copper complexes are distorted square planar structures with nitrogen and sulfur donors. This is consistent with X-ray diffraction measurements which also show copper(II) in a distorted square planar environment that is bereft of CuN2S2 2- intermolecular interactions. DFT calculations resulted in optimized structures that are consistent with crystallographic data and indicated HOMO-SOMO gaps of 5.01 eV and 4.68 eV for C and D as respectively. Optimized structures of Ni(ema)2- and Ni(emi)2- share the same basic characteristics as for the copper(II) congeners. Electrochemical characterization of C and D resulted in a reversible CuIII/II couple at -1.20 V and - 1.40 V, respectively. Reactivity studies with Rh(CO)2+ show similar donor capabilities for complexes A-D. Analysis of A shows that transmetallation does not occur. From competitive metal uptake studies on immobilized tripeptide it is concluded that the N2S2 4- ligating unit has a slight preference for Cu2+ over Ni2+ and that the biosynthetic pathway responsible for constructing the distal site of ACS must be selective for nickel insertion or copper exclusion, or both.
The A-cluster active site of the bimetallic Ni enzyme, Acetyl CoA Synthase (ACS), shown in Figure 1, utilizes the metallopeptide Ni(Cysteine-Glycine-Cysteine)2-, Ni(CGC) 2-, as a metallodithiolate ligand to the catalytically active metal center, denoted as Mp (p = proximal, with respect to the 4Fe4S cluster, refer to Figure 1), which, in the case of Mp = Ni, performs the organometallic chemistry required to assemble acetyl coenzyme A.1,2 Salient findings are that (a) the presence of nickel in both sites is required for catalytic activity; (b) NiII at the proximal position may be selectively removed and reconstituted or replaced by other metals, retaining Ni in the distal site; and (c) the enzyme was inactive when CuI or CuII was added to the nickel-depleted enzyme (It has been reasonably assumed that the removable or labile nickel is Nip, with no evidence of depletion at Nid).3 In addition, ZnII substitution at the Nip position resulted in no catalytic activity.1,2 From these studies it was concluded that the active form of the ACS A-cluster contains a labile nickel center, Nip that is replaceable by other metals.3 Nevertheless, nickel was purported to be the distal site metal in all reported structures even when copper or zinc was present in the growth media. Thus, a preference for nickel at this site is indicated.3,4
Other metals (CoII and FeII) have been found in a strikingly similar N2S2 Cys-Ser-Cys motif of Nitrile Hydratase suggesting that the biological N2S2 ligation site is not restricted to nickel.5 Furthermore, many synthetic neutral MN2S2 small molecule models have been produced (where M = a variety of transition metals, most notably Ni, Fe, Co, Cu, Pt, and Zn) showing a considerable range in the binding capability of the N2S2 donor set.6 In particular, the nickel N2S2 complexes are known to be exceptional metalloligands, and there are several studies of their ability to nucleate bi, tri, and polymetallics through controlled thiolate sulfur aggregation.7 Nickel-containing small molecules have been prepared as mimics of the Nid site, Ni(CGC)2-, found in ACS.8 Riordan and coworkers have extensively studied the Ni(II) complex of the CGC4- ligand, showing its ability to bind to a second low valent, “organometallic-like” nickel through the thiolate sulfur of Ni(CGC)2-.9,10
Similar to biological studies, investigations using the neutral Ni(bme-daco) (bme-daco = bis(mercaptoethyl)diazacyclooctane) complex showed that the reactive thiolates had a greater specificity for copper over nickel and zinc, but there was no evidence of NiII ejection or Ni/Cu exchange from the N2S2 core.3,11 These solution studies in combination with the biological studies discussed above piqued interest as to the reason why the M(Cysteine-X-Cysteine) biological motif is yet to be discovered with M = CuII and how the distal Ni(Cysteine-Glycine-Cysteine)2- site of ACS is constructed sans CuII. In this regard, a baseline study of the stability and reactivity of biologically relevant CuN2S2 2- complexes via synthetic model complexes has been undertaken.
In nature, copper is found with diverse functions and binding sites reflecting its wide range of structural and redox utility in proteins.12,13 An additional aim of this study is to develop an understanding of how the unique electron transfer properties of biological copper in Cu/SR binding sites are achieved with no consequence for sulfur oxidation to disulfide. Frequently, the preparation of CuII monomeric complexes as synthetic analogues is challenged by a facile auto-oxidation process: 2 Cu2+ + 2 RS- → 2 CuI + RSSR.14 This is often followed by CuI binding to unreacted CuII-SR- moieties producing complex aggregates.15 A number of N3S pseudo-tetrahedral CuII-thiolate containing complexes have been produced utilizing scorpionate-type ligands as models of these systems.16 However, only a small number of structurally and spectroscopically characterized synthetic N2S2 square planar, copper(II)-thiolate complexes have been reported as monomers capable of withstanding copper reduction and aggregation (Figure 2).17,18 Although a number of these complexes model blue copper electron transfer sites, only the Krüger complex, Cu(phmi)2- (phmi = N,N'-1,2-phenylenebis(2-mercapto-2-methylpropionamide)), contains amido donors that mimic the metal donor groups found in the Md site of ACS.18 An advantage for the lack of aggregation of this model complex is that the ligand is composed of a bulky phenylene group on the N-to-N linker as well as gem-dimethyl groups on the aliphatic carbons adjacent to the thiolates.
In order to investigate the prospect of a CuII in the Cys-Gly-Cys N2S2 4- binding site of ACS and concomitantly avoiding aggregate formation as a consequence of redox processes we have prepared the CuN2S2 2- (B, C and D) complexes shown in Chart 1, as well as a CuII-N2S2 2- system (O-Cu(CGC)2-, A, tethered to a polymer scaffold reported to provide site-site isolation.10 The resin-bound system is used for investigating Ni/Cu exchange processes while preventing or impeding the formation of large aggregates.
For solution studies, the N-acetylated and C-amidated Ac-CGC-NH2 tripeptide was obtained by the Fmoc method of solid phase peptide synthesis followed by cleavage from the NovaSyn TGR (TGR = TentaGel® Resin).10 The K2[CuII(CGC)] complex, represented in Chart 1 as B, was prepared by addition of CuII(OAc)2 to a basic methanol solution of Ac-CGC-NH2 in N,N'-dimethylformamide over a period of 15 min in order to minimize aggregate formation. The purple K2[Cu(CGC)] complex, [ESI-MS ([Cu(CGC)]- = 380.97 m/z (100%), Calc. = 380.98 m/z)], is air-sensitive but stable indefinitely when stored under N2 as a solid. In contrast, decomposition occurs in methanol or N,N'-dimethylformamide solutions after a few hours resulting in a highly insoluble yellow-brown solid. Its insolubility and mass spectral data suggest aggregate formation, likely resulting in S-oxidation and CuII reduction to CuI. 14-16
In previous studies of dianionic NiN2S2 2- complexes, we have shown Ni(ema)2- to model the Ni(CGC)2- complex in terms of additional metal binding through S-based ligating ability.10 Due to the increased instability in solution and decomposition pathways of Cu(CGC)2- versus the nickel analogue, synthetic analogues [Et4N]2[Cu(ema)] and [Et4N]2[Cu(emi)], (C and D, Chart 1), were prepared.
According to eq 1, and similarly to the synthesis of the Cu(CGC)2- complex, the air sensitive Et4N+ salts of C and D were prepared, recrystallized and analyzed by -ESI-MS. Complexes C and D showed parent ion peaks at 266.93 and 323.00 m/z respectively ([CuN2S2]- = 100% shown in Supporting Information, Figure S1). In the absence of air, the complexes are soluble and stable for at least 24 h in acetonitrile, methanol, N,N'-dimethylformamide, and water solutions as demonstrated by invariant UV-vis spectra. The Evans' method of determining magnetic moments19 revealed a μeff of 1.52 B.M. for Cu(ema)2- and 1.59 B.M. for Cu(emi)2- indicating a CuII ion with one unpaired electron at room temperature in acetonitrile solution for each case. As the cyclic voltammograms, vide infra, show completely reversible CuIII/II redox couples, and as the N2S2 4- type ligands have been shown by both Holm and Krüger to permit isolation of the MIII oxidation state in Ni and Cu complexes, we conclude that the deviation from the ideal magnetic moment of 1.73 B.M. is likely attributed to [CuIII(ema)]- or [CuIII(emi)]-.18,20
Dark purple crystals suitable for x-ray diffraction studies were obtained by layering concentrated acetonitrile solutions with diethyl ether. Despite attempts to maintain dry conditions, water molecules were found within the crystal forms. The molecular structures of C and D along with the nickel analogues [Et4N]2[Ni(ema)] and[Et4N]2[Ni(emi)], reported by Holm et al. are shown for comparison as ball and stick representations in Table 1.8, 20 The MN2S2 2- (M = Ni, Cu) complexes contain a symmetry-imposed mirror plane perpendicular to the MN2S2 plane that bisects the N-C-C-N linker backbone. The Cu-S and Cu-N distances are similar to those reported for the Cu(phmi)2- complex shown in Figure 2.18 Complex D has symmetry imposed, strict planarity while C has a mean deviation of ±0.0033Å in the CuN2S2 plane. Despite the symmetry present in the molecule, inducing the planarity indicated above, elongated thermal parameters in the crystal structures of C and D may be suggestive of a slight tetrahedral twist in the ligand set. An interesting structural feature is the elongation of the M-S bond resulting in the expansion of the S-M-S angle in the CuII complexes compared to their NiII analogues. This increase in bond length is attributed to population of an antibonding orbital within the σ bonding orbital regime, vide infra. The packing diagram of each structure (see Figure 3 and Figure S5) shows that the MN2S2 complexes are linked via hydrogen bonds of interstitial water molecules; for example, molecules of C are linked through water molecules from the amido oxygen of one CuN2S2 unit to the thiolate sulfur of the next. These linkages create sheets composed of MN2S2 complexes and water molecules with the Et4N+ cations separating each layer. Therefore, in the solid state no intermolecular CuN2S2 2- interactions are observed.
The copper complexes B, C, and D react with [Rh(CO)2Cl]2 similarly to NiN2S2 x- complexes (x = 0, 2)10. Examples of crystallographically characterized NiN2S2M(CO)x complexes are shown in Figure 4 as structures E, F, and G. Evidence for adduct formation with the CuN2S2 x- complexes as metalloligands is obtained from the two band ν(CO) pattern observed in the FTIR spectra of solutions following mixing of the Cu and Rh reagents (Figures S4 and S5). Table 2 compares the ν(CO) values of the Rh(CO)2 + adducts of the nickel complexes with the copper congeners. The Cotton-Kraihanzel force constants (k) are very similar, however, the interaction force constants (ki) are slightly higher for Cu vs. Ni,21 which could be related to the difference in the S-M-S bite angle noted in the crystallographic metric parameters (vide supra). Nevertheless the two ν(CO) bands are of similar intensity in all cases, indicating carbonyls at ca. 90° angles. In the absence of crystals of the copper derivatives IR measurements suggest that the copper-rhodium complex is analogous to the dimeric nickel derivative (G) shown in Figure 4,10 although the CuRh heterobimetallic, which would be analogous to F as shown below, cannot be ruled out.
The gas-phase DFT computations of the Ni(N2S2)2- complexes (B3LYP functional and the 6-311G(d,p) basis set) and of the Cu(N2S2)2- complexes (utilizing both unrestricted and restricted open shell B3LYP calculations and the 6-311G(d,p) basis set) resulted in optimized structures and metric parameters fully consistent with the crystallographic experimental results (see Supporting Information, Table S3).
The optimized structure of Ni(emi)2- has nearly identical geometric parameters to the Ni(ema)2- complex22,23 therefore the following discussion is appropriate to both complexes. Figure 5 illustrates that the HOMO and HOMO-1 molecular orbitals are close in energy, and are comprised of metal-d and sulfur-p π-antibonding interactions, involving anti-symmetric and symmetric combinations of the sulfur pz orbitals in the following linear combinations:
The slight energy difference could be explained in terms of the contribution from both the amido nitrogen and oxygen atoms in the HOMO-1, whereas the oxygen atom interaction is absent in the HOMO.
Table 3 gives the atomic orbital contributions to the frontier molecular orbitals for the nickel and copper complexes. The values presented for sulfur and nitrogen refer to each of the two atoms. In both Ni(ema)2- and Ni(emi)2-, the metal contribution to the HOMO is essentially the same; however, a slight difference is observed for the sulfur contributions. The HOMO-1 of both Ni(ema)2- and Ni(emi)2- complexes is comprised of Ni dyz character and pz character from each sulfur and also a symmetric combination of the nitrogen pz orbitals. The lowest unoccupied molecular orbital (LUMO) of Ni(emi)2- consists of an antibonding σ-orbital set, in the x-y plane. The HOMO-LUMO gap of Ni(emi)2- is calculated to be 4.20 eV which is slightly higher than that of Ni(ema)2-, found to be 4.12 eV.22.23
The frontier orbitals of Cu(ema)2- and Cu(emi)2- share the same fundamental characteristics with differences in orbital contributions from the copper and the ligand set. The orbital composition of the SOMO, similar to that of the LUMO of the corresponding Ni complexes, involves a σ-antibonding overlap of the dxy orbital of copper with the pσ orbitals of sulfur and nitrogen. The HOMO has a minor (4 %) copper contribution to the antisymmetric M-S dπ-pπ overlap and a thiolate sulfur contribution of 41-43% each. These orbital compositions are different for the corresponding nickel complexes where a larger metal contribution is observed. For both Ni and Cu, the HOMO-1 is more delocalized relative to the HOMO. The HOMO-SOMO gaps of Cu(ema)2- and Cu(emi)2- are 5.01 eV and 4.68 eV, respectively, a trend analogous to the HOMO-LUMO gaps of the Ni(ema)2- and Ni(emi)2- complexes (vide supra).
The difference in metal orbital contributions to the HOMO and HOMO-1 between the nickel and copper complexes can be attributed to intrinsic metal properties, in that the bonding d orbitals of the copper(II) are stabilized in energy relative to nickel(II), while the anti-bonding dxy, destabilized as a result of the ligand field symmetry, is indicated by the increase of copper character in the SOMO. Overall, the calculations indicate a greater degree of covalency in the metal-sulfur bond of the nickel complexes as compared to the copper most directly observed in the HOMO and HOMO-1 orbitals.
In general, computational structural parameters (described above) are corroborated by experiment. The most significant differences between the Cu and Ni complexes appears in the M-S and M-N bond lengths as well as the S-M-S and N-M-N angles. Due to the presence of a populated σ* orbital, which is unpopulated in the Ni cases, bond lengths in the copper complexes are ca. 0.1 Å greater than that found in the analogous nickel complexes and the S-M-S angle expands by approximately 5° for copper complexes. The distortions from planarity as indicated by the S1N1N2S2 dihedral angle are very small for the Ni complexes (3 to 4°), whereas for the Cu complexes a distortion in the form of a Td twist is between 8.5 and 9.5°, indicating somewhat of a shift towards a pseudo-tetrahedral geometry which is likely for a d9 metal ion. The fact that this is not observed in the experimental structures probably reflects crystallographically imposed symmetry. Nevertheless, the amido-thiolate N2S2 4- ligands are much more rigid binding sites as contrasted to the N2S2 2- ligands.
The electrostatic potential plots as well as the Mulliken charges for the M(N2S2)2- [M = NiII, CuII; N2S2 = ema4-, emi4-] series offer insight into the probable sites for charge-controlled, electrophilic and nucleophilic reactivity of the four complexes, see Figure 6 and Table 4. The overall charge delocalization imparted by the carboxamido unit on these dianionic NiN2S2 complexes as compared to the neutral NiN2S2 analogues was earlier noted. The negative charge of the carboxamido group is shared between both the nitrogen and the oxygen atoms, creating an iminolate contribution to the metal center.22,23 The Cu(ema)2- and Cu(emi)2- complexes display a greater degree of charge polarization as compared to the NiII congeners, implying that the CuII derivatives have less of a covalent interaction with the N2S2 4- ligand set. This is seen by inspection of the electrostatic potential maps of Figure 6 where there are similar areas of positive and negative character on the sulfur, nitrogen, and oxygen atoms of the N2S2O2 backbone for all four complexes. However, a major change in these four systems is seen by the amount of positive character on the copper in Cu(ema)2- and Cu(emi)2- versus the NiII analogues, indicating the greater ionicity in the former complexes.
The Mulliken charges listed in Table 4 for the ema4- and emi4- metal complexes are entirely consistent with the electrostatic potential maps but they provide a more quantitative analysis of the charge distribution. For example, the polarization of charge in the copper systems is observed in the thiolate sulfur character which is in the range of -0.584 and -0.638 while the thiolate sulfurs of the nickel complexes are significantly more positive, -0.509 and -0.571. This corresponds to a higher positive charge on copper of ca. 0.20 units as compared to nickel. There are no significant differences in the negative charge of the carboxyamido oxygen atoms throughout the series of complexes (between -0.511 and -0.520, see Table 4). For complexes containing the same metal ion, the differences in the charges of the sulfurs and the metal are small and probably insignificant. Due to the asymmetry of the ligand set, the negative charges of the nitrogen and the oxygen atoms in the CGC4- complexes are unequal, see Table 4.
The copper coordination environments for B, C, and D were assessed by electronic absorption spectroscopy and EPR spectroscopy. Table 5 shows the λmax and extinction coefficients characteristic of the electronic absorption spectrum of each complex in the UV-visible region in acetonitrile solvent. The position and intensity of the high energy bands at ~320 nm and 290 nm in the Cu(ema)2- and the Cu(emi)2- complexes suggest charge transfer transitions whereas the moderate intensity (ε ~400-800) of the low energy bands at ~ 400 nm and ~500 nm are likely Sσ→Cu and Sπ→Cu transitions respectively.12,13 Solomon and others noted that spectroscopic differences can be used to highlight structural dissimilarities between copper binding sites such as those in plastocyanin, nitrocyanin and synthetic metalloprotein active site models.12,17,18 For complexes B, C and D the Sσ → Cu transitions are red-shifted relative to nitrosocyanin and are of a much lesser intensity.12 The contour plots for the ground state wavefunctions of Cu(ema)2- and Cu(emi)2- illustrate that the S-Cu bonding is dominated by a dπ-pσ sigma interaction which is considered to be weak owing to a poor dπ-pσ orbital energy match and is a likely explanation for the weak band intensity.
The Sσ → Cu transition of the CuII complexes around 400 nm is lower in energy than the corresponding transition in NiII complexes.20 The relatively weak sigma interaction in the CuII complexes, when compared with NiII complexes, is indicated in the electrostatic potential diagram which shows a significant degree of charge separation and also in the longer M-S bond lengths in the CuII complexes. Whereas the flat nature of the molecule would normally facilitate both sigma and pi overlap, in this case, the frontier molecular orbital diagram shows that there is a poor energy match between the p orbitals on sulfur and the d orbitals on copper.
The nature of the copper environment in B, C, and D was assessed by EPR spectroscopy. The X-band EPR spectra resulting from frozen N,N'-dimethylformamide solutions (9 K) of the CuN2S2 2- complexes show axial signals with g > g > 2.003 which is the signature pattern expected for CuII (S = ½, I = 3/2) in a square planar environment with a dx2-y2 or dxy ground state (Figures S2 and S3). The anisotropic spin-Hamiltonian parameters exhibit values in the region expected for N2SX complexes where X = N, O or S. When A‖ and g‖ values are compared with the 2N2S region in Peisach-Blumberg correlation diagrams24 the observation is that values for B, C and D complexes lie in the region between the 4S and 4N complexes but outside of the region delineated for 2N2S complexes of CuII.24 This deviation is a consequence of larger g‖ values for B, C and D and is appreciable when compared to natural and artificial proteins although the degree of fit is improved when compared with the latter.24 There are several factors that influence the values of A‖ and g‖ and they may operate in concerted or opposing ways. Such factors include ligand charge, covalency and geometry of the metal binding domain.
EPR spectra were recorded for B, C and D in the coordinating solvent DMF; therefore, there is a strong likelihood that axial solvent coordination is influencing the g‖ values. It has been reported that axial coordination decreases the covalency in the axial bonds, resulting in the increased g‖ value observed for B, C and D.25 Assuming that the electric field around the CuII is unchanged going from liquid to frozen solution then the increased g‖ value, which places the complexes outside of the delineated region for 2N2S complexes in the Peisach-Blumberg correlations, 24 could be rationalized by considering axial solvent coordination. This argument is supported by the calculated Mulliken charges (vide supra) which suggest that there is a high degree of charge polarization in the Cu complexes. The implication is that electron delocalization, reduced in the plane of the ligand atoms, is manifested as increased g‖ values. The magnitude of A‖ and g‖ depends not only on the on the nature of the coordinating ligands but also the type of adjacent atoms. Since the axial EPR spectrum displayed by B, C and D is characteristic of CuII in a tetragonal environment and not representative of a sulfur-based radical it is reasonable to suggest that the magnetic parameters indicate both N and S binding to copper. There is no significant variation in the EPR parameters for all three complexes (B, C and D) indicating the structural similarities in copper binding geometries and ligand donor sets.
Cyclic voltammograms of C and D, measured in DMF, display two redox events (a full scan of C is in Figure S8). The response in the range of -800 to -1000 mV is assigned to the CuIII/II redox couple and its reversibility is verified by Ipa/Ipc values close to 1 (C = 0.957, D = 0.983), and by small ΔE values measured at 82 and 98 mV for C and D, respectively (Table S2). An irreversible oxidation attributed to ligand decomposition is observed at more positive potentials as well (Figure S8, Table S2).
The cyclic voltammograms of the isolated reversible CuIII/II redox couple of C and D are shown in Figure 10. To our knowledge, complex D, which is considered to have the most negative CuIII/II couple reported to date at E½ = -1.40 V vs. Fc+/Fc, is 0.20 V more difficult to access compared to that of C. Although CuIII is considered rare, access to the high valent cuprate state can be attributed to stabilization afforded by the highly negative and polarizable carboxamido nitrogens from the ligand set. This was first noted by Krüger and co-workers with the Cu(phmi)2- complex (E½ = -1.16 V vs. Fc+/Fc in CH3CN).18 Figure 7 shows that the potentials of the CuIII/II couples of C and D are >350 mV more negative than the reversible NiIII/II couples of the Ni(ema)2- and Ni(emi)2- complexes. This is consistent with the greater effective nuclear charge of CuII over NiII and further demonstrates, as discussed by Kitagawa, et al., how the choice of metal can be used to tune the properties of a biological ligand.26
The N-acetylated, resin-bound O-CGC4- tripeptide was obtained by the Fmoc method of solid phase peptide synthesis as described above for complex B (vide supra).10 The potassium salt of the resin-bound copper complex O-Cu(CGC)2- was prepared in a fritted syringe by the addition of 0.142 mmol/g of CuII(OAc)2 (assuming a similar O-CGC4- uptake capacity as was found in nickel studies)10b to a N,N'-dimethylformamide (DMF) slurry of O-CGC4- loaded beads producing a deep purple coloration within 10 minutes, Scheme 2. A schematic representation of the composition of the resin is given in Figure 8. The CGC loaded beads/CuII(OAc)2 slurry was agitated for an additional 50 min, then extensively washed with DMF, dichloromethane, and methanol solvents, and dried in vacuo. An EPR spectrum of the 1:1 copper/peptide reaction product was recorded (Figure 9a) and an axial spectrum, similar to B, C, and D was obtained. The measured values of the anisotropic spin-Hamiltonian parameters (g = 2.230, g = 2.100 and A = 195 G) correlate well with those of complexes B, C and D (vide supra) suggesting an N2S2 coordination for Cu.II The fact that the resin-bound complex possesses EPR characteristics similar to B, C and D indicates that site isolation, rather than polynuclear products, has been achieved within the resin. A recent investigation involving resin-bound cysteine as part of the ligand system demonstrated the formation of a stable Ni2+ complex as its Ph2PCH2CH2Ph2 derivative, O-(N-S)2- Ni(dppe).29
When excess copper was added to the O-CGC/Cu2+ loaded beads, an additional interaction was indicated by line broadening in the X-band EPR frozen solution spectrum (10 K) Figure 9b. An EPR spectrum of the resin and Cu2+ in the absence of peptide is also shown in Figure 9c. The higher g-values (g = 2.110 and g = 2.304) correspond to an oxygen environment and are indicative of a PEG-CuII interaction.24 A number of reports have shown that chelation of various transition metal ions such as RhIII by the PEG can occur.27 Uptake of transition metals with an affinity for O-donors is not surprising in view of the fact that the TentaGel® resin-beads are mainly comprised of oxygen-rich PEG units.28 Since 1 equivalent of Cu(OAc)2 results in O-Cu(CGC)2-, it is reasonable to conclude that binding of CuII by the CGC moiety occurs preferentially and that excess CuII results in stable Cu-PEG interactions.
In order to further identify the presumed resin-bound O-Cu(CGC)2-, attempts were made to derivatize the CuN2S2 site with Rh(CO)2 + since this procedure was effective in studies with NiN2S2 2-.10 The υ(CO) infrared spectrum (ATR-FTIR) obtained from the resulting purple-red beads on addition of a solution of [Rh(CO)2Cl]2 followed by DMF, methanol, and dichloromethane washes and drying in vacuo, showed two bands of equal intensity at 2066 and 1989 cm-1. This indicates cis-CO moieties ca. 90° to one another (Figure S4b) and donor capabilities that are similar to the Cu(ema)2- and Cu(emi)2- complexes. An indication of the air-stability of the O-Cu(CGC)2- beads was noted by their ability to react with Rh(CO)2 + after 2 weeks of exposure to air. In contrast, the resin-free analogues, B, C, and D, decompose within minutes.
Since the thiolates of Ni(CGC)2- found in Acetyl CoA Synthase display a high affinity for both CuII and NiII and model studies of the synthetic, neutral NiN2S2 complexes gave similar results, exogenous metal binding and possible transmetallation reactions of CuII within the N2S2 core were investigated.
As a preliminary investigation to compare the affinity of CuII versus NiII for the O-CGC4- site, a rudimentary competition study was performed. [The word “rudimentary” is used because for solubility reasons, two different sources of Cu and Ni are used and the nature of the metal ion species present in solution is not defined.]30 Since aggregates may form in solution, the resin-bound analogue O-CGC4- was used for this study. To a suspension of O-CGC4- in N,N'-dimethylformamide equimolar solutions of Cu(OAc)2 and Ni(acac)2 were added simultaneously. The contents of the flask containing the resin-bound peptide suspension were agitated for 1 h, collected via filtration, washed and dried. An EPR spectrum of the orange-brown beads showed a signal corresponding to O-Cu(CGC)2-, but was difficult to quantify. On exposure of the beads to solutions of [Rh(CO)2Cl]2, the broad bands in the IR spectrum centered at 2065 and 1991 cm-1 indicate a mixture of two or more species with the O-Cu(CGC)2- adduct of Rh(CO)2 + predominating.
The Cu/Ni competition experiment for O-CGC4- experiment was repeated and the suspension of beads and reagents was agitated for 6 h, collected via filtration, washed and dried. The resulting concentration of nickel and copper within the beads was determined via Neutron Activation Analysis (Sample 1) and compared to samples of O-Ni(CGC)2- (Sample 2) and O-Cu(CGC)2- (Sample 3) produced from the same batch of O-CGC4-. As shown in Table 7, when presented with both copper(II) and nickel (II) as the salts described above, both metals are bound and the overall uptake of copper is greater than that of nickel resulting in a Cu:Ni ratio of 1.7:1. (This overall Cu:Ni uptake ratio however does not consider the issue of two sites within the resin beads that bind copper, vida infra.) The EPR spectrum of the orange-brown beads from Sample 1 showed a signal corresponding to O-Cu(CGC)2-.
As discussed earlier, in addition to the O-CGC4- copper-binding sites the PEG moiety can interact with copper which could account for the slightly higher metal concentration in Sample 3 relative to Sample 2. Earlier studies with good reproducibility demonstrated a 1:1 correlation of O-CGC4- and Ni2+ uptake,10b hence we can assume that the copper-only sample has ca. 0.02 mmol/g Cu2+ taken up in the PEG. If this assumption is correct, then the mixed-metal sample contains ca. 0.10 mmol/g Cu2+ and a Cu:Ni ratio of 1.4:1. We cannot account for the discrepancy of the total metal uptake in Sample 1 and the amount of Cu taken up in Sample 3, and we do not know if this discrepancy indicates the error in the experiment. To this end, further analytical studies are underway. Nevertheless it is clear that both Cu and Ni compete for the O-CGC4- site.
The addition of 20 equivalents of Ni(acac)2 to purple beads of O-Cu(CGC)2- resulted in no color change over the course of 6 h. Following washes and vacuum drying the resulting EPR spectrum displays the axial signal previously assigned to O-Cu(CGC)2- further indicating that there is no significant exchange of Ni2+ for Cu2+. The reverse experiments (i.e. exchange of Cu2+ for Ni2+) attempting opposite exchange were also unsuccessful. Furthermore, under the conditions described above the thiolates of the O--Ni(CGC)2- and O-Cu(CGC)2- do not produce stable paramagnetic, multimetallic species when exposed to exogenous CuII or NiII in the form of CuCl2, Cu(OAc)2, NiCl2, Ni(acac)2 or Ni(OAc)2.
The synthesis of new and stable dianionic [CuIIN2S2]2- complexes has been described along with characterization including -ESI-MS, UV-vis, EPR, and x-ray diffraction studies. The UV-vis and EPR data show that the complexes have a CuII center with two amido nitrogen and two thiolate sulfur donors provided by the ligands emi4-, ema4-, and CGC4-. The conclusions drawn from solution spectroscopy are corroborated by two solid state structures which show the CuII is in a square planar N2S2 arrangement. The metal carbonyl derivatives show that the thiolate reactivity with Rh(CO)2 + of these CuN2S2 2- complexes parallels the well-studied NiN2S2 analogues. The CuRh bimetallic series also reveals that Cu(ema)2-, Cu(emi)2-, and Cu(CGC)2- are equal in donor ability. These results are reflected in the electrostatic potential maps derived from DFT calculations. As the EP's and Mulliken charges indicate a greater M-S charge separation when M = Cu, the ionicity within M(N2S2)2- is greater than when M = Ni. Nevertheless, the greater negative charge on sulfur does not correspond to a greater donor ability as indicated by the M(N2S2)Rh(CO)2 - complexes.
Synthesis of O-Cu(CGC)2- using previously reported methodology for the O-Ni(CGC)2- species produces dark purple beads that are stabilized against decomposition with O2.10 The immobilized O-Cu(CGC)2- complex can be spectroscopically identified through EPR signals. Electron paramagnetic spectroscopy also showed that CuII binds to the underivatized TentaGel® beads in the absence of the CGC unit, presumably to the ether oxygen donors of PEG. Nevertheless, the capture of CuII by the CGC N2S2 units surpasses the Cu-PEG interaction. Derivatization with the Rh(CO)2 + unit qualitatively shows that the same reactivity as observed in solution may occur resin-bound, further confirming that the Cu(CGC)2- is immobilized.
Metal uptake studies showed that CuII and NiII are stable once in the N2S2 core, i.e., transmetallation or metal ion exchange with Ni(acac)2 and Cu(OAc)2 as source of Cu and NiII, respectively, does not occur. Furthermore, multimetallic species formed by thiolates bridging to NiII or CuII were not observed using the acac or acetato sources for exogenous Ni or Cu. Thus, the reactivity with Rh(CO)2+ indicates the requirement for stabilizing ligands on the exogenous metal ion for producing bimetallic species such as O-Cu(CGC)Rh(CO)21-.
Despite obvious caveats to the competitive metal uptake studies described above, we conclude that the N2S24- ligating unit, either free in solution or immobilized in a matrix, has a slight preference for CuII over NiII. A reasonable conclusion with regards to the ACS active site is that the biosynthetic pathway responsible for constructing the distal site of ACS must be selective for nickel insertion or copper exclusion, or both. We note that no biological CuII sites composed of the Cys-X-Cys ligand set have been reported to date. However, examples of Ni, Fe, and Co are known for enzymes with strikingly similar N2S2 ligands albeit with quite different biological roles. As our studies indicate a strong similarity between NiII and CuII in a N2S24- donor environment, and, in fact, a slight preference for CuII binding, an eventual sighting of a Cys-X-Cys-CuII moiety will not be surprising.
Solvents were purified according to standard procedures and were freshly distilled under N2 prior to use or purified and degassed via a Bruker solvent system.31 Other reagents were purchased from commercial sources and used as received unless noted.
Standard Fmoc Merrifield techniques were employed for all peptide syntheses. Fmoc-Cys(Mmt)-OH and Fmoc-Gly-OH were obtained from NovaBioChem. The solid supports, NovaSyn TGR resin-beads (Novabiochem) averaging 90 μm in diameter, are composed of polystyrene with crosslinking via divinylbenzene and grafted with polyethyleneglycol bearing a Rink linker (Rink = trialkoxybenzhydrylamine) as the free amine termini (0.4 mmol/g loading). Plastic-fritted syringes, 10 mL, were used as reaction vessels to facilitate the multiple additions and removal of reagents and wash solvents. Mixing of beads and reagents was accomplished by an automated shaker. Synthesis of resin-bound Cys-Gly-Cys (O-CGC4-) was performed as described previously.10
Syntheses of air sensitive Ni(II) complexes were performed under anaerobic conditions using distilled/degassed solvents and standard Schlenk line techniques under an argon atmosphere. The [Et4N]2[Ni(ema)] and [Et4N]2[Ni(emi)] complexes were prepared following published procedures and used in CV analysis to be consistent with conditions used for the new complexes reported herein.20
Solution infrared spectra were recorded on a Bruker Tensor 27 FTIR spectrometer using 0.1 mm NaCl sealed cells. The Pike MIRacle™ attachment from Pike Technologies was used for Attenuated Total Reflectance Infrared Spectra for solid state samples. UV-Vis spectra were recorded on a Hewlett Packard HP8452A diode array spectrometer. Mass spectrometry (ESI-MS) was performed by the Laboratory for Biological Mass Spectrometry at Texas A&M University. Elemental analyses were performed by the Canadian Microanalytical Servies, Ltd., Delta, British Columbia, Canada. Cyclic Voltammograms were obtained under an Ar atmosphere at 22°C using a BAS100W potentiostat equipped with a 3.0 mm glassy carbon working electrode, a platinum wire auxiliary electrode, and Ag/AgNO3 reference electrode. Measurements were performed in a DMF solution with 0.1 M [Bu4N][BF4] as supporting electrolyte. Ferrocene was used as an internal standard and reported relative to NHE (Fc/Fc+ = +692 mV vs. NHE).32 EPR spectra were obtained with a Bruker ESP 300 equipped with an Oxford ER910 cryostat operating at 10 K. Samples were 1 mM in analyte in DMF and frozen in liquid N2 prior to recording the EPR spectra at 10 K. The g values reported were as simulated using the WinEPR Simfonia program.33 Magnetic moments were determined utilizing the Evans method using a mixed solvent system of CD3CN and CH3CN.19 Measurements were taken by 1H NMR using an Inova 300 MHz NMR spectrometer
For C and D, a BRUKER D8 GADDS general-purpose three-circle X-ray diffractometer was employed for sample screening and data collection. The goniometer was controlled using the GADDS software suite (Microsoft Win 2000 operating system). A 0.7 mm nylon loop was dipped in the sample. The mounted tube was then placed on the diffractometer and in a cold nitrogen stream (Oxford) maintained at 280K. The sample was optically centered with the aid of a video camera such that no translations were observed as the sample was rotated through all positions. The detector was set at 12cm from the sample (MWPC Hi-Star Detector, 1024 × 1024 pixel). The X-ray radiation employed was generated from a Cu sealed X-ray tube (Kα = 1.54184Å with a potential of 40 kV and a current of 40 mA) and filtered with a graphite monochromator in the parallel mode (175 mm collimator with 0.5 mm pinholes).
The beam intersection coordinates were compared to the corundum (Al2O3) standard coordinates and changes were made accordingly. A single data frames were taken at widths of -179° in θ (continuous rotation on ϕ: Gandolfi Scan) with an exposure time of 5 seconds. Three frame sets at 20, 55, and 90° 2-θ settings were collected. The data was reduced by area integration methods to produce a single powder diffraction pattern for each frame. Merging the individual powder diffraction patterns in the program EVA produced a single one-dimensional pattern, which is reported. The parameters are listed in Table S1.
For C and D, the integrated intensity information for each reflection was obtained by reduction of the data frames with the program SAINT V6.63.34 The integration method employed a three-dimensional profiling algorithm and all data were corrected for Lorentz and polarization factors, as well as for crystal decay effects. Finally the data was merged and scaled to produce a suitable data set. The absorption correction program SADABS was employed to correct the data for absorption effects. X-Seed was employed for the final data presentation and structure plots.35 The tetraethylammonium cation and water molecules in each structure were found to be disordered between two positions. The anions were modeled by employing local bond distance restraints and included in the final refinement. CCDC reference numbers: 691889 and 691890.
DFT calculations were performed using a hybrid functional (the three-parameter exchange functional of Becke (B3)36 and the correlation functional of Lee, Yang, and Parr (LYP)37) (B3LYP) as implemented in Gaussian 03.38 For each calculation, all atoms were optimized via the use of the 6-311g(d,p) basis set. Due to the paramagnetic nature of the model copper complexes, an unrestricted open-shell calculation was initially performed; however, in order to prevent spin contamination inherent in the unrestricted open-shell calculations, restricted open-shell calculations were also investigated. Optimized bond lengths and angles in these two methods were in all cases nearly identical. Cartesian coordinates for the starting input geometry of Ni(ema)2- was extracted from the crystallographic structure, and the models of Ni(emi)2- and the copper analogues were modified by adding the gem-dimethyl groups or changing the metal center, respectively. A frequency calculation was performed alongside each geometry optimization to ensure the stability of the ground state as ascertained by the absence of imaginary frequencies. Graphical visualizations of the electron density of the individual molecular orbitals and the electrostatic potentials were first generated as cube files from Gaussian, and these images were implemented into and viewed by the Cerius2 software package.39 Optimized geometries of each of the complexes were imported into and visualized with JIMP2 and can be located in the Supporting Information.40 Geometric parameters were extracted from the optimized structures by use of the GaussView program.41 For each complex in this text, the energies have been converted from values in Hartrees to eV.
Deprotection of the TentaGel®-bound, N-acylated, Mmt Cys-S-protected Ac-CysGlyCys-NH2 tripeptide was accomplished via a 1:94:5 mixture of trifluoroacetic acid:dichloromethane:triisopropyl silane as described for the synthesis of O-Ni(CGC)2-.10 Following washes (3 × 5 mL each) with pure dichloromethane and methanol solvent, a basic solution of Cu(OAc)2 in 5 mL of 50:50 methanol:dichloromethane was introduced to the resin bed. The light-yellow TentaGel® beads changed to a deep purple after 10 minutes. After 50 min, the CuII supernatant solution was expelled and the beads were again washed with methanol, dichloromethane, and diethyl ether solvent until the residual washes were completely colorless. The purple beads were dried in vacuo and stored in a vacuum desiccator.
In an analogous manner to the preparation of O-Ni(CGC)Rh(CO)21-, a sample of O-Cu(CGC)2-, purple in color, was placed in a fritted syringe and swelled in N,N'-dimethylformamide. A yellow dichloromethane solution of [Rh(CO)2Cl]2 was added resulting in dark red beads after 30 min. The beads were washed thoroughly with dichloromethane, methanol, and diethyl ether (3 × each) and then dried in vacuo. ATR-FTIR on solid, dried beads: ν(CO) = 2061s, 1983s cm-1.
The H4CGC (20 mg, 0.062 mmol) ligand and KOH (13.92 mg, 0.0248 mmol) were combined in 20 mL of N,N'-dimethylformamide and allowed to stir for 20 minutes. To this a blue solution of Cu(OAc)2 (15.6 mg, 0.0610mmol in MeOH) was added drop wise. The reaction mixture developed into a mauve purple after 15 minutes and was allowed to stir overnight. A light purple solid was obtained after precipitation with Et2O. Yield 16.3 mg (70.0% yield). Absorption spectrum (CH3CN): λmax (ε, M-1 cm-1) 232 (26,185), 294 (18,962), 322sh (13,450) and 367(930) nm.- ESI-MS: m/z 380.9314 [M]-
In a degassed Schlenk flask N, N'-ethylenebis(2-mercaptoacetamide) (2.09 g, 4.61 mmol), KOH (1.5 g, 27.8 mmol) and tetraethyl ammonium chloride (1.65 g, 10.0 mmol) were combined with 100 mL of methanol and magnetically stirred for 30 min. To this a bright blue methanol (100 mL) solution of Cu(OAc)2 (0.8 g, 4.3 mmol) was added drop wise over the course of 5 minutes. The mixture was stirred for 15 min during which time the light yellow solution changed to maroon and then purple. After 1.5 h additional reaction time, the solvent was removed in vacuo and the residue was redissolved in ca. 40 mL of CH3CN. The solution was filtered through a Celite pad and reduced to 20 mL. A second filtration using filter paper wrapped around a cannula and secured with copper wire produced a deep purple solution. A solid was obtained from the filtrate by addition of 250 mL of diethyl ether. (1.74 g, 83% yield) Absorption spectrum (CH3CN): λmax (ε, M-1 cm-1) 233 (24,880), 293 (28,700), 320sh (22,100), 407 (1740) nm.- ESI-MS: m/z 266.9275 [M]-. Anal. Calcd (found) for C22H48N4CuO2S2: C 50.01 (49.75), N 10.60 (10.57), H 9.16 (9.30).
The solids N,N'-ethylenebis(2-benzylthio)isobutyramide) (1.61 g, 4.62 mmol), KOH (1.5 g, 27.8 mmol) and tetraethyl ammonium chloride (1.65 g, 10.0 mmol) were dissolved in 200 mL CH3OH and stirred for 30 min under a blanket of Ar. The slow addition of a MeOH solution of Cu(OAc)2 (0.8 g, 4.3 mmol) to this solution followed by stirring at room temperature for 30 min produced a red solution. The solution was allowed to stir for an additional 30 min and filtered through a Celite pad; the solvent was removed in vacuo and the residue taken up in ca. 40 mL of CH3CN. The solution was filtered through Celite pad and reduced to 20 mL. A second filtration using a paper filter wrapped around a cannula and secured with copper wire produced a deep purple solution. (1.36 g, 74% yield). Crystals suitable for x-ray diffraction were obtained at -10° by layering a concentrated CH3CN solution with hexane and then ether. Absorption spectrum (CH3CN): λmax (ε, M-1 cm-1) 231 (23,799), 294 (25,250), 319sh (21,300), and 413 (2120) nm.- ESI-MS: m/z 323.0042 [M]-. Anal. Calcd (found) for C26H56N4CuO2S2•H2O: C 50.33 (50.24), N 9.75 (9.72), H 9.03 (9.02).
Under an argon blanket, the purple solid, [Et4N]2[Cu(ema)] (20 mg, 0.037 mmol), was suspended in 20 mL N,N'-dimethylformamide. After 30 minutes stirring the solid was completely dissolved to give a yellow-red solution. A yellow solution of [Rh(CO)2Cl]2 (7.373 mg, 0.0189 mmol) was then added drop wise. After stirring for 2 h the resulting red-brown solution showed ν(CO) = 2061, 1982 cm-1. A brown solid was obtained following precipitation with diethyl ether. Absorption spectrum (N,N'-dimethylformamide): λmax (ε, M-1 cm-1) 270(27,610) and 330(18, 610) nm.
To a red solution of [Et4N]2[Cu(emi)] (20 mg, 0.034 mmol) in 20 mL N,N'-dimethylformamide, a yellow solution of [Rh(CO)2Cl]2 (7.4 mg, 0.019 mmol) was added drop wise. The resulting brown-red solution provided ν(CO) = 2062, 1984 cm-1. A brown solid was obtained following precipitation with diethyl ether. Absorption spectrum (N,N'-dimethylformamide): λmax (ε, M-1 cm-1) 289(25,260) and 390(18,540) nm.
We acknowledge financial support from the National Science Foundation (CHE-0616695 to M.Y.D, MCB 0542291 DAR) with contributions from the R.A Welch Foundation (A-0924 to MYD) and the National Institutes of Health (Chemistry-Biology Interface Training Grant to K.G. T32GM008523). We sincerely thank the TAMU x-ray crystallography facility (Drs. Nattamai Bhuvanesh and Joseph H. Reibenspies), The Laboratory for Molecular Simulation at Texas A&M University (Dr. Lisa Perez), and the Laboratory for Biological Mass Spectrometry (Dr. Shane Tichy) for their expert assistance in this work.
Supporting Information Available: Crystallographic data for C and D in both CIF and table format and a packing diagram for D, electrochemical data as both cyclic voltammograms and in table format, additional EPR and IR spectra of CuII complexes, and computational data including metric parameters and optimized structures. This material is available free of charge via the Internet at http://pubs.acs.org.