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4.11: Nature of the Metal-Dioxygen Linkage in Biological Systems

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    Many techniques have been used to probe the geometry and electronic structure of the metal-dioxygen moiety in biological systems and in synthetic models. The results form the basis of any understanding of the factors that determine and modulate oxygen affinity.

    Oxyhemocyanin and Oxyhemerythrin

    By resonance Raman techniques, the O—O stretch is observed at 744 cm-1 in oxyhemocycanin and at 844 cm-1 in oxyhemerythrin.140-142 Dioxygen is therefore coordinated as peroxo species. By use of unsymmetrically labeled dioxygen, 18O—16O, it was established that dioxygen coordinates symmetrically:141


    Carbon monoxide binds to only one of the CuI centers in deoxyhemocyanin, through the C atom,* apparently blocking the second CuI site.15 Similar behavior is also seen for the nitrosyl adduct.

    * A report143 that CO binds to the copper center through the O atom, an unprecedented mode, has been challenged.144

    On the other hand, in an experiment parallel to that just described for hemocyanin, dioxygen was found to bind asymmetrically in oxyhemerythrin:142

    Figure 4.13).45 This mode of dioxygen coordination remains unobserved to date in small-molecule synthetic systems. The O2-binding process is formally very similar to the combinations of Reactions (4.29a) and (4.29b), except that dioxygen attaches to only one Fe center (bridging carboxylato groups omitted in Reaction 4.45):


    The existence and location of the proton, which cannot be proven in the crystal structure of deoxyhemerythrin or of oxyhemerythrin, are inferred from a model system for the former that contains a hydroxo group bridging two high-spin, weakly antiferromagnetically coupled Fell centers (J = -10 cm-1).88 For oxyhemerythrin (and for one conformation of methydroxohemerythrin), a small change in the position of the symmetric Fe—O—Fe mode is observed when H2O is replaced by D2O.146 The strong antiferromagnetic coupling observed for methemerythrin and oxyhemerythrin (-J ~ 100 cm-1)147 is uniquely consistent with a bridging oxo moiety between a pair of FeIII centers.80 Finally, a Bohr effect (release or uptake of protons) is absent in oxygen binding to hemerythrin.16,18 These observations are consistent with a \(\mu\)-oxo group slightly perturbed by hydrogen bonding to a coordinated hydroperoxo species. An important role for the protein in hemerythrin is to assemble an asymmetric diiron(II) species. Only a few of the myriad of known \(\mu\)-oxodiiron(III) complexes are asymmetric,81 and the synthesis of realistic asymmetric models remains a challenge.

    Deoxyhemocyanin (CuI d10) and deoxyhemerythrin (Fell d6) are colorless. In the oxygenated derivatives there is considerable charge transfer between the coordinated peroxo groups and the metal centers. This phenomenon makes the essentially d-d metal transitions more intense than those for the simple aquated Fe3+ or Cu2+ ions, and permits facile measurement of oxygen-binding curves. The spectral changes accompanying oxygenation are shown in Figures 4.20 and 4.21.5

    Nitric oxide binds to deoxyhemocyanin, to deoxyhemerythrin, and to the mixed valence FeIII • • • Fell semimethemerythrin.148 Carbon monoxide binds to neither form of hemerythrin: apparently the other ligands have insufficiently strong fields to stabilize the low-spin state for which electron density would be available for back donation into the CO \(\pi\)* orbitals.

    Figure 4.20 - Spectral changes accompanying oxygenation of hemocyanin: deoxyhemocyanin (- - -); oxyhemocyanin (——).5 Reproduced with permission from A. G. Sykes, in A. G. Sykes, ed., Advances in Irorganic and Bioinorganic Mechanisms, Academic Press, 1 (1985), 121-178.
    Figure 4.21 - Spectral ehanges accompanying oxygenation of hemerythrin.5 Reprodueed with permission from A. G. Sykes, in A. G. Sykes, ed., Advances in Inorganic and Bioinorganic Mechanisms, Academic Press, 1 (1985), 121-178.


    The O—O stretch that is observed by difference infrared techniques at around 1105 cm-1 for oxyhemoglobin and oxymyoglobin149 clearly categorizes the dioxygen moiety as a superoxo species; that is, the order of the O—O bond is about 1.5. Considerable ink has been spilled about the nature of the Fe—O2 fragment since Pauling's original suggestion150 in 1948 that dioxygen binds to iron in an end-on bent fashion:


    He subsequently reaffirmed this geometry, and proposed that hydrogen bonding between the coordinated dioxygen and the distal imidazole H—N group was important in stabilizing the Fe—O2 species.137 In an alternative model Weiss proposed that a low-spin FeIII center (S = \(\frac{1}{2}\)) was very strongly antiferromagnetically coupled to a superoxide anion radical (S = \(\frac{1}{2}\)).151 A triangular peroxo mode has also been advanced.152,153 The problem has been how to resolve the observed diamagnetism of oxyhemoglobin92,98 with UV-visible, x-ray absorption, and resonance Raman spectroscopic characteristics154 that are distinctly different from those of Fell systems (such as carbonmonoxyhemoglobin and low-spin six-coordinated hemochromes, such as Fe(Porph)(Py)2) and from unambiguously FeIII systems (such as chloromethemoglobin or cyanomethemoglobin).

    Any adequate theoretical treatment must also explain how iron-porphyrin systems can bind not only O2, but also CO, NO, alkyl isocyanides, and alkyl-nitroso moieties. A simple qualitative model presented by Wayland and coworkers129,155 conveniently summarizes ligand-binding geometries of cobalt and iron porphyrins. Although a reasonable quantitative theoretical consensus exists for 1:1 cobalt-dioxygen species, the same cannot be said yet for irondioxygen systems.

    A Simple Model for the Electronic Structure of Liganded Hemoglobins

    Why does dioxygen bind to iron and cobalt porphyrins in an end-on bentbond fashion as in (4.37) and (4.46)? Why does carbon monoxide bind in a linear manner (Equation 4.40)? Why are six-coordinate dioxygen and carbonmonoxide adducts more stable than five-coordinate ones? A unified picture of ligand binding that addresses these questions is important in understanding properly the specific case of dioxygen binding to hemoglobin and related systems.

    The splitting of the metal d orbitals for a four-coordinate metalloporphyrin is shown in the center of Figure 4.22. These orbitals contain some porphyrin character and are antibonding with respect to metal-porphyrin bonds. As shown in Figure 4.16, the primary effect of a single \(\sigma\)-donor axial ligand, such as pyridine or 1-methylimidazole, is to elevate the energy of the antibonding dz2 and lower the energy of the dx2-y2 orbital and hence lead to a high-spin species in place of the intermediate-spin four-coordinate one. Thus, for simplicity in highlighting interaction of the metal center with the diatomic \(\sigma\)-donor: \(\pi\)-acid ligands CO, NO, and O2, the perturbations wrought by primarily \(\sigma\)-donor ligands, such as 1-methylimidazole, are omitted. For the corresponding cobalt(II) compound, there is an additional electron. The diatomic ligands of interest share a qualitatively similar molecular orbital scheme. The filling of orbitals for CO is shown on the left-hand side. Dioxygen, which is shown on the right-hand side, has two more electrons than CO; these occupy the doubly degenerate \(\pi\)* orbitals. Quantitative calculations show that the energy of the \(\pi\)* orbitals decreases monotonically from CO to NO to O2, indicating increasing ease of reduction of the coordinated molecule, a feature that has not been included in the diagram. Only those interactions of molecular orbitals that have appropriate symmetry and energy to interact significantly with the metal d orbitals are shown.

    Two extremes are shown in Figure 4.22 for the interaction of a diatomic molecule A-B with the metal center: a linear geometry on the left and a bent geometry on the right. A side-on geometry is omitted for the binding of O2 to a CoIl or Fell porphyrin, since this would lead to either an MIll side-on superoxo or an MIV peroxo species; both these modes of coordination to these metals are currently without precedent.

    Figure 4.22 - A simple general model for ligand binding to metalloporphyrins:129 AB = CO(\(\pi\)*0), NO (\(\pi\)*1), O2(\(\pi\)*2); M=Co Il(d7), Fell(d6).

    Linear diatomic metal bonding maximizes the metal-d\(\pi\) to ligand-p\(\pi\)* bonding. When a ligand coordinates in a bent manner, axial symmetry is destroyed, and the degeneracy of the ligand p\(\pi\)* orbitals is lifted. One p\(\pi\)* orbital is now oriented to combine with the metal dz2 orbital to form a \(\sigma\) bond, and the other is oriented to combine with dxz and dyz orbitals to form a \(\pi\) bond. A bent geometry for the diatomic molecule will result when either or both of the metal dz2 or the ligand p\(\pi\)* orbitals are occupied, since this geometry stabilizes the occupied dz2 orbital in the five-coordinate complex. Thus O2 binds in a strongly bent manner to CoIl and Fell porphyrins; NO binds in a strongly bent manner to CoIl porphyrins; CO binds in a linear fashion to Fell porphyrins.

    The interaction of NO with Fell porphyrins and CO with CoIl porphyrins—the resultant species are formally isoelectronic—is more complicated. The degree of bending seen in Fell(TPP)(NO) is midway between the two extremes.111For CO the higher-energy p\(\pi\)* orbitals lead to a greater mismatch in energy between the dz2 and p\(\pi\)* orbitals, and less effective \(\sigma\) bonding. In EPR experiments the odd electron is found to be localized in a molecular orbital with about 0.87 metal dz2 character for the five-coordinate Co—CO adduct, as expected for a nearly linear geometry.129 On the other hand, for the Fe—NO adduct the metal dz2 character of the odd electron is about 0.4 to 0.5;155 a somewhat bent geometry (140°) is observed in the crystal structure of Fe(TPP)(NO). Because the CO ligand is a very weak \(\sigma\) donor, the Co—CO species exists only at low temperatures.

    Only qualitative deductions can be made from this model about the extent of electron transfer, if any, from the metal onto the diatomic ligand, especially for dioxygen. The higher in energy the metal dz2 orbital is with respect to the dioxygen p\(\pi\)* orbitals, the closer the superoxo ligand comes to being effectively a coordinated superoxide anion. With an additional electron, the dioxygen ligand in Co—O2 complexes can acquire greater electron density than it can in Fe—O2 complexes.

    From the diagram it may be inferred that a ligand with very strong \(\pi\)-acceptor properties will lower the energy of the dxz and dyz orbitals through strong (dxz, dyz)-\(\pi\)* interaction. The resultant energy gap between these two orbitals and the other three metal d orbitals may be sufficient to overcome the energy involved in spin-pairing, and hence lead to five-coordinate low-spin species, as happens for complexes containing phosphines and carbon monosulfide.99,132

    4.11: Nature of the Metal-Dioxygen Linkage in Biological Systems is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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