The Silver Complexes of Porphyrins, Corroles, and Carbaporphyrins

Research: Science and Education

The Silver Complexes of Porphyrins, Corroles, and Carbaporphyrins: Silver in the Oxidation States II and III Christian Brückner Department of Chemistry, University of Connecticut, Storrs, CT 06269-3060; [email protected]

Silver in the oxidation states II and III? If a student calculates an oxidation state of II or III for silver in one of its compounds, most instructors would reach for a red pen without hesitation. We teach that silver has an oxidation state of I in its compounds and we are familiar with AgICl, AgIBr, AgINO3 and, perhaps, the oxide AgI2O. Advanced inorganic texts, however, teach of higher oxidation states of silver (1). Most of the higher oxidation state silver compounds are prepared under extreme conditions. For instance, the reaction of elemental fluorine with finely divided silver produces AgIIF2. Further oxidation generates the square-planar complex [AgIIIF4]− or, using high pressure fluorine in the presence of CsF, even the silver(IV) complex Cs 2[Ag IVF 6].

H2N N

N AgIII

HN N H2N

3+

NH2 NH N

H

NH2

H

Figure 1. Structure of a (dibiguanidinium)AgIII complex.

Oxidation of AgI2O with peroxodisulfate (S2O82−) at 90 ⬚C in strongly alkaline solutions generates the solid oxide “AgIIO”, which was shown to be a mixed oxidation state species formulated as AgIAgIIIO2. High oxidation state species of uncomplexed silver are not stable in neutral aqueous solutions as they oxidize water to oxygen. Surprising in the context of the high oxidation potential of the Ag2+ and Ag3+ ions, organic complexes containing silver in these oxidation states have been known for quite some time. The oldest examples are the dibiguanidinium complexes containing silver in a III oxidation (2) although no structural proof is available (Figure 1). Other examples have appeared over the years (3). Recent work from a number of laboratories around the world, including our own, has shown that well-defined complexes of tetrapyrrolic ligands containing silver in II and III oxidation states are readily accessible, stable, and devoid of oxidizing characteristics (4– 8). They are structurally characterized and the determination of their central metal oxidation states is unambiguous. These complexes should not be regarded as exceptions or curiosities as their formation follows classic coordination chemistry rules. In fact, their existence supports one of the tenets of coordination chemistry, namely, that a given ligand environment largely determines the oxidation state of a coordinated metal ion. The following discussion will explain why stable AgII and AgIII complexes should not surprise, and why tetrapyrrolic ligands are ideally suited to provide these high oxidation state complexes. While the focus of the discussion is on the silver complexes, a discussion of the tetrapyrrolic metal complexes of the metals surrounding silver in the periodic table, group 10 (nickel group), 11 (coinage metal group), and 12 (zinc group), will help to put the silver chemistry into perspective and will highlight the generality of the findings (Figure 2). Porphyrins, Corroles, and Carbaporphyrins as Ligands

Figure 2. Detail of the periodic table.

N N

N

H HH

H N

N

N

N H N

N

porphyrin

HH H N

N

corrole

N

carbaporphyrin

Figure 3. Structures of porphyrins, corroles, and carbaporphyrins highlighting their aromatic 18 π-electron systems.

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Porphyrins are rigid tetrapyrrolic macrocycles that, after the loss of the inner protons, can serve as dianionic ligands (the so-called porphyrinato ligands), presenting four nitrogens in a square-planar arrangement to metal ions located in their center (Figure 3). Porphyrins form complexes with the vast majority of metals of the periodic table (9). Corroles are “contracted porphyrins”; that is, they are tetrapyrrolic fully conjugated macrocycles containing three methine bridges and one direct pyrrole–pyrrole linkage (Figure 3; ref 10). Both porphyrins and corroles contain aromatic 18 π-electron systems that are in conjugation with two or one peripheral double bonds, respectively. The 18 π-system in the contracted framework of corroles is maintained by the change of the oxidation state of one nitrogen. Three pyrroletype nitrogens and one imine-type nitrogen thus line the cen-

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tral cavity of a corrole as compared to two nitrogens of each type in a porphyrin. Consequently, corroles are slightly smaller trianionic N4-donors as compared to the dianionic N4-donating porphyrins. Carbaporphyrins are porphyrin isomers in which one pyrrolic building block is “inverted” relative to a regular porphyrin. They are also known as “N-confused porphyrins”(11, 12). In their metal complexes they act as trianionic CN3-donors, providing a square-planar coordination environment similar in size to that of porphyrins. The coordination chemistry of corroles (13) and carbaporphyrins (5, 7, 14, 15) is still nascent but a rich coordination chemistry can be expected. Additional coordination of axial ligands to the metal center in all of the tetrapyrrolic ligands may lead to square-pyramidal or octahedral coordination spheres (16). Group 10, Copper, and Group 12 Complexes of Porphyrins The stable oxidation states of nickel, palladium, platinum, copper, zinc, cadmium, and mercury in their porphyrinato complexes, that is, the oxidation state the metal will adopt in aerated and wet solutions of its porphyrin complexes, are all II (9). Thus, these metals form neutral, squareplanar complexes with porphyrins according to equation 1: + porphyrin–H2 + M2 + 2 base (porphyrinato)MII + 2 baseH+ (1)

M = Ni, Pd, Pt, Cu, Zn, Cd, Hg The II oxidation states follow the general oxidation state preferences for these metals. The observed square-planar coordination geometries are as expected for the d8 ions and the d9 ion Cu2+. This coordination sphere for the d9 ion CuII can be viewed as the extreme of the Jahn–Teller-distorted (tetragonal elongated) octahedral coordination sphere typically associated with this ion. The size of the central cavity of porphyrins is a little too large for low-spin square-planar NiII (0.63 Å) (17), resulting in the observation of relatively long Ni⫺N bond distances (18) but is well suited to accommodate PdII, PtII, and CuII (0.78 Å, 0.74 Å, and 0.71 Å, respectively) (17). The good steric and electronic fit of these four metals in their porphyrin complexes is demonstrated by their extraordinary stability toward acid-induced demetallation, requiring extremely harsh conditions such as concentrated H2SO4 at elevated temperatures. The neutral, square-planar porphyrinato complexes of ZnII, CdII, and HgII also follow the general oxidation state preferences of these metals, which are governed by their particularly stable d10 electron configuration. Many of these metalloporphyrins display a square-planar coordination sphere but especially the Lewis acidic Zn2+ can also coordinate to additional axial ligands (19). As a result of their sizes and zero ligand-field stabilization, all zinc-group porphyrinato complexes are sensitive toward acid-induced demetallation, with their lability increasing with ion size (ionic radii of ZnII: 0.74 Å; CdII: 0.92 Å and HgII: 1.12 Å) (17). The HgII porphyrins are quantitatively and rapidly demetallated by dilute mineral acids. 1666


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Silver and Gold Porphyrins The d10 ion AgI has a considerable affinity for nitrogen donor atoms but prefers linear coordination geometries (1). Upon warming of a solution of porphyrin and two equivalents of AgI acetate in pyridine, metallation of the porphyrin takes place and one equivalent of elemental silver is deposited (20): porphyrin–H2 + 2AgI(acetate) (2) (porphyrinato)AgII + Ag0 + 2acetic acid Evidently, “squeezing” the large AgI ion (1.16 Å) (17) into the rigid porphyrin lowers its oxidation potential such that the second equivalent of AgI in the reaction mixture is strong enough to oxidize it, giving rise to the observed disproportionation reaction. The smaller oxidized d 9 ion (0.93 Å) (17) fits the porphyrin cavity well. The lowering of the oxidation potential of AgI can be understood considering that the electrostatic interaction between the dicationic metal and the dianionic metal ligand is much increased as compared to the initial monocation–dianion pair. Further, the resulting d9 ion is subject to crystal-field stabilization effects whereas the closed shell d10 ion AgI is not. The crystal structure of the AgII porphyrinato complex shows that the AgII ion is, like NiII and CuII in their porphyrinato complexes, coordinated in a near-ideal square-planar fashion (Figure 4). Like the CuII complex, the d9 complex is paramagnetic with a magnetic moment corresponding to one unpaired electron. The strong-field ligand environment of the porphyrin has further implications: the lone electron of the d9 ion occupying the highest occupied molecular orbital (HOMO) is relatively high in energy and therefore accessible toward oxidation (Figure 5). In other words, the potential to oxidize the Ag2+ ion located in the square-planar ligand field of a porphyrin is significantly reduced as compared to that of the free ion. Hence, AgII porphyrins can be chemically or electrochemically oxidized to form the corresponding AgIII porphyrins. The oxidation of the AgII porphyrin complex is accomplished at a lower potential (E0 = 0.59 V SCE) (22) than the

Figure 4. Stick representation of X-ray crystals structure of (mesotetraphenylporphyrinato)AgII (4, 21). The idealized square-planar coordination environment around the central metal is highlighted.


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oxidation of the corresponding CuII porphyrinato complex (E0 = 1.00 V SCE) (23). Furthermore, the oxidation product of the CuII porphyrinato complex is a CuII complex with an oxidized porphyrin ligand. This falls within the general trends observed within the periodic table. An electron is more readily removed from a larger ion of identical charge and valence electron configuration. In case of the porphyrinato ligand presenting the donor atoms in a rigidly fixed arrangement, the larger AgIII ion (0.81 Å) (17) presumably also fits better than the hypothetical very small CuIII ion (estimated to be < 0.6 Å). What if an even larger metal such as gold is inserted into a porphyrin? AuI, inserted into a porphyrin, is oxidized spontaneously and forms a cationic AuIII porphyrin complex (24). The resulting d8 complex features gold in its “regular” high oxidation state (1). The oxidative metallation is facilitated by the fact that the ligand-field splitting for the heavier elements is significantly larger than for the lighter elements. The oxidation to the +3 ion generates a small (0.82 Å) (17) and, in the square-planar coordination environment of the porphyrin, very stable d8 ion that is isoelectronic to the equally stable (porphyrinato)PtII. Silver Corroles The structural differences of corroles as compared to porphyrins are reflected in their altered metal-coordination properties. Following purely electrostatic arguments, one can predict that the smaller trianionic corrolato ligand has a greater ability to stabilize higher central metal oxidation states than the larger dianionic porphyrinato ligand. This is indeed observed (8, 25–32). Warming a free base corrole with three equivalents of silver(I) acetate in pyridine, smooth metallation of the macrocycle can be observed. Again, the AgIII complex forms in a disproportionation reaction (8): corrole–H3 + 3AgI(acetate) + 3pyridine (corrolato)AgIII + 2Ag0 + 3acetate− + 3pyH +

(3)

A single-crystal X-ray structure of the complex proves the connectivity of the complex and its square coordination sphere (Figure 6). The coordination sphere is not perfectly

planar because the corrole is slightly saddled (alternating pyrrole units are tilted up and down). This is an indication that the Ag3+ central ion is slightly too big to fit into the central cavity of a planar corrole as this saddling enlarges the macrocycle core. In turn, this implies that the even larger Ag2+ ion would have been much too big to fit, supporting the size argument brought forward to explain the observed oxidation states. How does one demonstrate the presence of a particular oxidation state? Some spectroscopic and analytical data in concert lend support to the notion of a III oxidation state of the central metal in (meso-tritolylcorrolato)AgIII (8): the resulting silver complex of the trianionic corrolato ligand is neutral. Sharp signals in the NMR spectra of this complex prove it to be diamagnetic, strongly supporting the d8 electron configuration of AgIII. An X-ray photoelectron spectroscopy investigation of the complex showed that higher energies were required to remove a particular inner-shell electron from the silver ion in the corrolato complex as compared to the AgII porphyrinato complex. The central metal in the AgIII corrolato complex is isoelelectronic with PdII for which the observation of square-planar complexes is the norm. Although “free” Ag3+ is a potent oxidant, the (corrolato)AgIII complex formed “spontaneously” with only Ag+ present as an oxidant. Thus, one might expect the complex not to show any pronounced oxidation capabilities. This is also experimentally found (8). The complex is very stable and shows no signs of being an oxidant. Even treatment of the (corrolato)AgIII complex with gaseous H2S does not lead to any reaction. This highlights the extraordinary stability the corrolato ligand imparts onto the III oxidation state of the coordinated metal ion. However, treatment of the (corrolato)AgIII complex with weaker reducing acids of stronger acidity such as HBr and HI quickly leads to a reductive demetallation reaction according to eq 4. The reaction is presumably also driven by the extraordinary low solubility of the silver(I) halogenides formed. (corrolato)AgIII + 4 HI (4) corrole–H3 + AgI(s) + H + + I3−

dx 2−y 2

dxy dxz

dxy

dyz

dz 2 dx 2−y 2 dz 2 dyz

dxz 9

free d ion

9

d ion in a square-planar ligand field

Figure 5. Ligand-field splitting diagram for a d9 ion in a squareplanar geometry.

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Figure 6. Stick representation of X-ray crystal structure of (mesotritolylcorrolato)AgIII (8,21). The slightly distorted square planar coordination environment around the central metal is highlighted.


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Silver Carbaporphyrins

3−

Carbaporphyrins were also shown to stabilize AgIII (5– 7). This is convincingly supported by their diamagnetic properties and by X-ray crystallography. Silver in the AgIII state is stabilized by the trianionic carbaporphyrinato ligand despite the fact that the ligand contains a central cavity size comparable to that of the dianionic porphyrinato ligand. As outlined, the dianionic porphyrin is capable of forming the III oxidation state of silver only under strongly oxidizing conditions. Hence, the charge component of the electrostatic stabilization of high oxidation states of metal ions coordinated to carbaporphyrins appears to be crucial. To what extent the superior σ-donor capabilities of the carbaporphyrins influence their ability to stabilize AgIII remains to be determined. Metal–Nitrogen Bond Distances as a Measure of the Metal Oxidation States In general, the higher the oxidation state of a metal, the smaller its ion size and the shorter the expected metal–ligand bond distances. However, this analysis fails in the cases of the largely rigid tetrapyrrolic ligands because the metal–N bond distances are determined by the metrical parameters of the ligand alone. For instance, an average Ag⫺N bond distance of 2.09 Å is found in the (porphyrinato)AgII complex and an average of 2.06 Å is found for the Ag⫺N bond distances in the similarly dimensioned (carba-porphyrinato)AgIII complexes (5–7). By comparison, the Ag⫺N bond distances measured in the (corrolato)AgIII complex are significantly shorter and average 1.95 Å (8). Thus, the oxidation states of the central metal do not define the metal–donor atom bond distances and cannot be used as a measure for the oxidation state of the metal. High Oxidation State Stabilization by Corroles and Carbaporphyrins and the “Non-Innocent” Nature of Tetrapyrrolic Ligands Is the stabilization of higher oxidation states by corroles and carbaporphyrins general? This question can be answered in the affirmative. The stable oxidation state of nickel in corroles and carbaporphyrins is also III (7, 12, 26). A number of high oxidation state corrolato complexes of iron (27, 30, 33, 34) and chromium (32, 35) containing FeIV and CrIV, and CrV, respectively, have been reported in the past few years. However, one important feature of tetrapyrrolic ligands is their “non-innocent” nature. This means that the large electron-rich π-systems of the tetrapyrrolic ligand are readily capable of providing an electron to the central metal. This reduces the formal oxidation state of the metal and generates a ligand radical. However, in case of the CuIII complexes of corroles, for instance, this then gives rise to the following ambiguity: are the complexes observed “true” CuIII complexes with a trianionic corrolato ligand [(corrolato)3− CuIII] or are they CuII complexes of a dianionic radical corrolato ligand [(corrolato)2− • CuII] (Figure 7) (26, 36, 37)? Calculations have shown that the ground state of the copper complex of corroles is a diamagnetic CuIII state, with CuII ligand π-radical states slightly higher in energy (36). We found that the equilibrium between these two states is sol1668


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N

N

N

∆

III

Cu N

2−

N

N II

Cu N

N

Figure 7. Redox equilibrium of (corrolato)CuIII complexes.

vent- and temperature-dependent (37). Heating of a diamagnetic sample in a NMR magnet causes line shifts and line broadening. As expected for the narrow singlet–triplet state energy gap, heating shifts the equilibrium toward the paramagnetic species. While the distinction between these two states is relatively straightforward in the copper case, the current discussion of the nature of the iron corroles illustrates the difficulty to distinguish the limiting paramagnetic species [(corrolato)3– FeIV] and [(corrolato)2–• FeIII] (33, 34). Conclusion Considering this discussion, are the findings of stable AgIII complexes altogether surprising? Hardly, although only most recently has their detailed description appeared in the literature. Will these high oxidation state complexes be of greater practical use than providing textbook examples for principles of coordination chemistry? It is too early to provide a definitive answer, but the high oxidation state iron, manganese, and chromium complexes of porphyrins and corroles have already been demonstrated to possess unique catalytic properties in oxidation, epoxidation, and aziridination reactions (25, 28, 30). The (corrolato)AgIII complex is electrochemically readily and reversibly reduced to the corresponding AgII complex, suggesting this complex might be a competent catalyst in electron-transfer reactions. Ongoing research in this intriguing area of inorganic chemistry will certainly provide more clues regarding uses of high oxidation state complexes. Literature Cited 1. (a) Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Butterworth–Heinenmann: Oxford, 1997; pp 1180– 1193. (b) Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M. Advanced Inorganic Chemistry, 6th ed.; John Wiley & Sons: New York, 1999; pp 1084–1097. (c) Hollemann, A. F.; Wiberg, E.; Wiberg, N. Inorganic Chemistry, 1st English ed.; Academic Press: San Diego, CA, 2001; pp 1272–1273. 2. For a representative overview see: Chatterjee, B.; Syamal, A. J. Indian Chem. Soc. 1970, 47, 1021–1022. 3. See for example: (a) Barefield, E. K.; Mocella, M. T. Inorg. Chem. 1973, 12, 2829–2832. (b) Kirschenbaum, L. J.; Rush, J. D. J. Am. Chem. Soc. 1984, 106, 1003–1010. (d) Eujen, R.; Hoge, B.; Brauer, D. J. Inorg. Chem. 1997, 36, 1464–1475. 4. Scheidt, W. R.; Mondal, J. U.; Eigenbrot, C. W.; Adler, A.; Radonovich, L. J.; Hoard, J. L. Inorg. Chem. 1986, 25, 795– 799.


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Research: Science and Education 5. Furuta, H.; Ogawa, T.; Uwatoko, Y.; Araki, K. Inorg. Chem. 1999, 38, 2676–2682. 6. Furuta, H.; Maeda, H.; Osuka, A. J. Am. Chem. Soc. 2000, 122, 803–807. 7. (a) Muckey, M. A.; Szczepura, L. F.; Ferrence, G. M.; Lash, T. D. Inorg. Chem. 2002, 41, 4840–4842. (b) Lash, T. D.; Rasmussen, J. M.; Bergman, K. M.; Colby, D. A. Org. Lett. 2004, 6, 549–552. 8. Brückner, C.; Barta, C. A.; Briñas, R. P.; Krause Bauer, J. A. Inorg. Chem. 2003, 42, 1673–1680. 9. Buchler, J. W. In The Porphyrins; Dolphin, D., Ed.; Academic Press: New York, 1978; Vol. 1, pp 389–483. 10. Paolesse, R. In The Porphyrin Handbook; Kadish, K. M., Smith, K. M., Guilard, R., Eds.; Academic Press: San Diego, CA, 2000; Vol. 2, pp 201–300. 11. Furuta, H.; Asano, T.; Ogawa, T. J. Am. Chem. Soc. 1994, 116, 767–768. 12. Chmielewski, P. J.; Latos-Grazynski, L.; Rachlewicz, K.; Glowiak, T. Angew. Chem., Int. Ed. Engl. 1994, 33, 779–781. 13. Licoccia, S.; Paolesse, R. Struct. Bonding (Berlin) 1995, 84, 71–133. 14. Chmielewski, P. J.; Latos-Grazynski Inorg. Chem. 1997, 36, 840–845. 15. Furuta, H.; Kubo, N.; Maeda, H.; Ishizuka, T.; Osuka, A.; Nanami, H.; Ogawa, T. Inorg. Chem. 2000, 39, 5424–5425. 16. Scheidt, R. W. In The Porphyrin Handbook; Kadish, K. M., Smith, K. M., Guilard, R., Eds.; Academic Press: San Diego, CA, 2000; Vol. 3, pp 49–112. 17. Shannon, R. D. Acta Crystallogr. 1976, A32, 751–767. 18. The average Ni⫺N bond distances in meso-tetraphenylporphyrin are 1.96 Å (Fleischer, E. B.; Miller, C. K.; Webb, L. E. J. Am. Chem. Soc. 1964, 86, 2342–2347.) whereas the average Ni⫺N bond distance for square-planar, diamagnetic NiII⫺aromatic N donor atoms is closer to 1.88 Å (see for example Brückner, C.; Karunaratne, V.; Rettig, S. J.; Dolphin, D. Can. J. Chem. 1996, 74, 2182–2193.) 19. Sanders, J. K. M.; Bampos, N.; Clyde-Watson, Z.; Darling, S. L.; Hawley, J. C.; Kim, H.-J.; Mak, C. C.; Webb, S. In The Porphyrin Handbook; Kadish, K. M., Smith, K. M., Guilard, R., Eds.; Academic Press: San Diego, CA, 2000; Vol. 3, pp 1–48. 20. Adler, A. D.; Longo, F. R.; Kampas, F.; Kim, J. J. Inorg. Nucl. Chem. 1970, 32, 2443–2445.

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21. Structural data from Cambridge Structural Data Base (CSD) version 5.24, Nov. 2002, processed with CAChe v.4.9, Fujitsu, 2002. 22. Kadish, K. M.; Lin, X. Q.; Ding, J. Q.; Wu, Y. T.; Araullo, C. Inorg. Chem. 1986, 25, 3236–3242. 23. Kadish, K. M.; Morrison, M. M. Bioinorg. Chem. 1977, 7, 107–115. 24. Chambron, J.-C.; Heitz, V.; Sauvage, J.-P. New. J. Chem. 1997, 21, 237–240. 25. Gross, Z.; Golubkov, G.; Simkhovich, L. Angew. Chem., Int. Ed. 2000, 39, 4045–4047. 26. Will, S.; Lex, J.; Vogel, E.; Schmickler, H.; Gisselbrecht, J.-P.; Haubtmann, C.; Bernard, M.; Gross, M. Angew. Chem., Int. Ed. Engl. 1997, 36, 357–361. 27. Vogel, E.; Will, S.; Tilling, A. S.; Neumann, L.; Lex, J.; Bill, E.; Trautwein, A. X.; Wieghardt, K. Angew. Chem., Int. Ed. Engl. 1994, 33, 731–735. 28. Gross, Z.; Simkhovich, L.; Galili, N. Chem. Commun. 1999, 599–600. 29. Gross, Z. J. Biol. Inorg. Chem. 2001, 6, 733–738. 30. Simkhovich, L.; Gross, Z. Tetrahedron Lett. 2001, 42, 8089– 8092. 31. Golubkov, G.; Bendix, J.; Gray, H. B.; Mahammed, A.; Goldberg, I.; DiBilio, A. J.; Gross, Z. Angew. Chem., Int. Ed. Engl. 2001, 40, 2132–2134. 32. Meier-Callahan, A. E.; Di Bilio, A. J.; Simkhovich, L.; Mahammed, A.; Goldberg, I.; Gray, H. B.; Gross, Z. Inorg. Chem. 2001, 40, 6788–6793. 33. Steene, E.; Wondimagegn, T.; Ghosh, A. J. Phys. Chem. B 2001, 105, 11406–11413 (Correction and addition: J. Phys. Chem. B 2002, 106, 15312.). 34. Cai, S.; Licocca, S.; D’Ottavi, C.; Paolesse, R.; Nardis, S.; Bulach, V.; Zimmer, B.; Shokhivera, T. K.; Walker, F. A. Inorg. Chim. Acta 2002, 339, 171–178. 35. Meier-Callahan, A. E.; Gray, H. B.; Gross, Z. Inorg. Chem. 2000, 39, 3605–3607. 36. (a) Ghosh, A.; Wondimagegn, T.; Parusel, A. B. J. J. Am. Chem. Soc. 2000, 122, 5100–5104. (b) Wasbotten, I. H.; Wondimagegn, T.; Ghosh, A. J. Am. Chem. Soc. 2002, 124, 8104–8116. (c) Luobeznova, I.; Simkhovich, L.; Goldberg, I.; Gross, Z. Eur. J. Inorg. Chem. 2004, 1724–1732. 37. Brückner, C.; Briñas, R. P.; Krause Bauer, J. A. Inorg. Chem. 2003, 42, 4495–4497.


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The Silver Complexes of Porphyrins, Corroles, and Carbaporphyrins

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