Reversible uptake of molecular oxygen by heteroligand Co(II)–l-α-amino acid–imidazole systems: equilibrium models at full mass balance

The capability of compounds called natural respiratory pigments to reversibly absorb molecular oxygen has been the subject of intensive research since the end of the 19th Century and has been inspiring the creation of artificial systems to imitate their activity [1–14]. Example models of synthetic oxygen carriers include mixed complexes of the type Co(II)–auxiliary ligand–imidazole, in which imidazole coordinates in trans position against the bound O₂ molecule, alike imidazole of the proximal histidine in myoglobin and hemoglobin [15]. In contrast to classical methods of preparing such compounds by mixing separate solutions of Co(II) salts, appropriate amino acids and imidazole [16–18], an original method has been applied, in which cobalt(II) and imidazole were introduced in the form of a polymeric, pseudo-tetrahedral, semi-conductive complex [Co(imid)₂]_n. This results in the formation of definite, unique structures with an imidazole molecule in an axial position opposite the O₂ molecule [19–26].

[Co(imid)₂]_n is a coordination compound crystallizing in an infinite polymeric net, in which each cobalt(II) ion is joined via imidazole bridges with four adjacent ions of the metal [27, 28]. Each Co(II) ion forms two dative bonds with the nitrogen atoms of two deprotonated imidazole moieties and two ionic bonds with the nitrogen atoms of two other imidazoles (Fig. 1). Therefore, this alternative method of obtaining dioxygen complexes with a strictly defined structure by starting from the [Co(imid)₂]_n polymer is much more effective than the method in which appropriate so-called “active” complexes capable of reversible dioxygen uptake are formed in an equilibrium mixture during titration of a solution containing Co(II) ions, the suitable auxiliary ligand (e.g. amino acid) and imidazole [16, 17].

The peculiar property of O₂ transport in such Co(II)–amac–Himid systems, as with the natural dioxygen carriers, results from the rapidly stabilizing equilibrium present in solution between the “active” form and the dioxygen-containing form. The “active” form, responsible for the dioxygen transport, is usually a paramagnetic, high-spin, hexacoordinate Co(II) complex of Co^II(amac)₂(Himid)(H₂O) composition, containing two chelate–like connected amino acid molecules forming an equatorial plane, as well as two axial ligands–imidazole and water. After substitution of the dioxygen molecule for water, a dimeric, diamagnetic [Co^III(amac)₂(Himid)]₂O₂ ²⁻ complex is formed with the O₂ molecule coordinated in peroxide order i.e. with a O₂ ²⁻ (μ-peroxy) bridge between two cobalt ions formally oxygenated to Co(III). This complex, because of the eventual partial irreversible oxidation of Co(II) to mononuclear Co(III) products, is frequently denoted as an intermediate oxygen adduct. Owing to the elongation of the dioxygen bond from 120.7 pm for the triplet O₂ to 149.0 pm for the peroxide O₂ ²⁻ anion, the oxygen adducts may be used as intermediate complexes in catalytic processes [29–34].

The O₂ ²⁻ bridge (μ-peroxy) exists within pH = 3–9, but upon a rise in basicity above pH 10, this is transformed into a poorly reversible dibridged Co(III)O₂ ²⁻OH⁻Co(III) (μ-peroxy–μ-hydroxy) form. This double-bridge appears in place of the two carboxyl groups, which easily undergo dissociation and which are found in cis position towards the coordinated dioxygen molecule. Such a complex is a much less effective O₂ carrier due to its higher affinity for autoxidation. An alternative known description of the oxygen bridges is the form type η, corresponding to “side on” bridge μ-peroxy structures [35]. In turn, acidification of the solution at a low temperature (−3 to 0 °C) leads to protonation of the μ-peroxy bridge, whereas the forming intermediate Co(III)O₂ ²⁻H⁺Co(III) product undergoes a rapid decay accompanied by Co(III) ion formation. In addition, at a temperature around 0 °C and in acidic medium, the O₂ ²⁻ (μ-peroxy) bridge may be subsequently oxidized by means of strong oxidizers, e.g. Ce⁴⁺, MnO₄ ⁻ or Cl₂ ions. As a result, a paramagnetic, stable {[Co^III(amac)₂(Himid)]₂O₂ ⁻}⁺ complex is formed, with an irreversibly bound dioxygen moiety in the Co(III)–O₂ ⁻–Co(III) (μ-superoxy) bridge.

All known O₂ carriers (both natural and synthetic) form complexes of two types: monomeric, with an M:O₂ stoichiometry of 1:1, and dimeric, with an M:O₂ stoichiometry of 2:1. An analysis of the theoretically estimated values of the free standard Gibbs energy of the O₂ reactions with metal ions and their complexes could be expected to favor the dimeric structures. In fact, the ΔG° value for the dimer formation reaction attains negative values for a much higher number of metals than is the case for monomer formation. This effect refers to the displacement of complex-formation decidedly to the right [36]. The data find a practical confirmation because among all the known dioxygen carriers, in aqueous solution we observe formation of stable dimeric complexes.

Previous investigations of the Co(II)–amac–Himid systems have not included the key aspect, i.e. accurate calculations of the Co(II), amac and Himid concentrations at equilibrium, by using the formation constants reported in our work for analogous oxygen-free systems [37]. These calculations may allow the equilibrium concentrations of all equilibrium forms present in solution to be determined, and for the \(K_{{{\text{O}}_{2} }}\) equilibrium constants to be evaluated without using any simplified approximations, which for instance take into account only the “active” complex and the oxygen adduct within the mass balance system [19, 38]. Moreover, the advantage of the experimental methods used in the present work, i.e. a direct gas–volumetric experiment with simultaneous pH measurement, is that it allows the degree of reversibility of O₂ uptake to be taken into account. As for many other complexes, including a majority of complexes with amino acids and peptides, the irreversible part of the reaction is quite rapid (e.g. t_1/2 <5 min for glygly), which excludes the use of the most commonly applied method based only on potentiometric titration [39–41].

The optimum amac to Co(II) ratio equaled 2:1. Above this value, the amount of dioxygen taken up did not change (see Additional file 1: Figure S1). The amount of imidazole released from the [Co(imid)₂]_n moiety as a result of the mixed Co(II)–amac–Himid–O₂ complex formation (0.3 mol Himid per 0.3 mol Co) indicates that the stoichiometric Co(II): imidazole ratio was 1:1, which confirms that one of the two [Co(imid)₂]_n imidazole moieties remains in the coordination sphere of cobalt(II) of the final complex (see Additional file 2: Figure S2). In other words the structure of the forming dioxygen adducts are unified by the presence of one imidazole in the coordination sphere.

Investigations of oxygenation of the Co(II)–amac–Himid systems indicated that dioxygen uptake is accompanied by a rise in pH to 9–10. An example of the time dependence between pH and the number of mmoles of bound dioxygen is shown for l-α-histidine in Fig. 2. The percentage of reversibility noted after acidification of the solution to the initial pH ranged within ca 30–60% for l–α–alanine, 40–70% for l-α-asparagine and 50–90% for l-α-histidine; this rose as the share of the amino acid in the Co(II): amac: Himid ratio increased (Table 1). The results confirm that the axial imidazole plays a role in enhancing reversibility of the O₂ uptake as opposed to the systems with the same amino acids but lacking imidazole [42]. Imidazole is in fact an important complement of the coordination sphere of the central ion as a donor of a free-electron pair of the N3 nitrogen.

Amac	Co(II): amac:Himid (mmole ratio)	Time (min)	pH	mmoles O2	% revers
Alanine	0.3: 0.15:0.3	90	9.882	0.0515	30.33
	0.3:0.225:0.3	105	9.557	0.0559	31.25
	0.3:0.3:0.3	210	9.751	0.0918	31.80
	0.3:0.375:0.3	215	9.629	0.0936	46.76
	0.3:0.5:0.3	215	9.562	0.1193	57.71
	0.3:0.6:0.3	270	9.603	0.1343	49.68
	0.3:0.75:0.3	195	9.418	0.1457	45.56
	0.3:0.9:0.3	270	9.422	0.1661	46.65
	0.3:1.05:0.3	200	9.276	0.1690	48.97
	0.3:1.2:0.3	120	9.247	0.1704	57.72
Asparagine	0.3:0.15:0.3	210	9.498	0.0352	43.37
	0.3:0.225:0.3	150	9.752	0.0567	45.04
	0.3:0.3:0.3	270	9.634	0.0600	44.37
	0.3:0.375:0.3	190	9.656	0.0936	49.77
	0.3:0.45:0.3	150	9.543	0.0901	53.55
	0.3:0.6:0.3	120	9.152	0.1219	60.78
	0.3:0.75:0.3	120	8.919	0.1237	57.84
	0.3:0.9:0.3	200	9.022	0.1304	65.27
	0.3:1.2:0.3	180	8.616	0.1275	67.67
Histidine	0.3:0.15:0.3	90	9.668	0.0373	51.16
	0.3:0.225:0.3	70	9.687	0.0565	52.85
	0.3:0.3:0.3	150	9.535	0.0755	72.99
	0.3:0.375:0.3	70	9.412	0.0932	68.98
	0.3:0.45:0.3	150	9.732	0.1123	69.81
	0.3:0.6:0.3	90	9.193	0.1496	82.63
	0.3:0.75:0.3	105	9.526	0.1513	86.69
	0.3:0.9:0.3	90	9.345	0.1502	85.06
	0.3:1.2:0.3	90	9.234	0.1507	84.81

In comparison with the system with histidine, the systems with alanine and asparagine demonstrated some higher values for the share of the free Co(II) ion and the particular complex species existing in solution apart from the oxygen adduct (regarded as the entire amount of cobalt engaged in both reversible and irreversible oxygenation) Fig. 3. For the two amino acids, the oxygenation reaction is right shifted to a relatively lower extent. The competitive binary complexes CoL₃ (or CoL₂ for histidine), which are able to reversibly take up dioxygen, are also present in solution in relatively low concentrations below 0.2% (Fig. 3); according to Fallab’s rule, they have a sufficient number of 3N in the coordination sphere [40]. However, the experimental results indicate that the only active complexes taking up dioxygen in practice are heteroligand species with imidazole as the second ligand (Additional file 2: Figure S2). Therefore, the \(K_{{{\text{O}}_{2} }}\) equilibrium constants can be calculated using formula (16), where the “active” complex is an appropriate heteroligand species with a concentration directly following the full mass balance equation. The fact that the value of log  \(K_{{{\text{O}}_{2} }}\) remained constant between different initial total Co(II), amino acid and imidazole concentrations, within limits of error (Table 2), indicates that the “active” complex was a heteroligand CoL₂L′complex for alanine and asparagine (Fig. 4), but also for histidine although the structure differs in participation of the N3 nitrogens of the additional imidazole side group (Fig. 5). The imidazole N1–H side-group does not dissociate in the measurable pH range due to it having a pK of 14.4 [43]. In addition, for histidine, as in the case of alanine and asparagine, the dioxygen substitutes a relatively weak donor, i.e. the deprotonated carboxyl oxygen, instead of the water molecule.

Amac	Co(II):amac:Himid (mmole ratio)	Log \(K_{{{\text{O}}_{2} }}\)
Alanine	0.3:0.5:0.3	9.442 10.957 10.216
	0.3:0.6:0.3	8.348 9.200 8.650
	0.3:0.75:0.3	10.061 10.107 10.100
Mean: 9.7 ± 0.8
Asparagine	0.3:0.45:0.3	7.377 8.739 7.996
	0.3:0.6:0.3	7.520 8.246 7.959
	0.3:0.75:0.3	7.609 7.831 7.743
	0.3:0.9:0.3	7.371 7.593 7.439
	0.3:1.2:0.3	7.952 7.997 7.978
Mean: 7.8 ± 0.4
Histidine	0.3:0.15:0.3	15.050 15.422 15.425
	0.3:0.375:0.3	14.002 14.175 14.179
	0.3:0.45:0.3	15.585 15.712 15.714
	0.3:0.6:0.3	14.537 14.471 14.475
Mean: 14.9 ± 0.7

The mean values are followed by standard deviations

The optical absorption spectra for the Co(II)–amac–Himid system with histidine indicate a significant increase of the molar absorption coefficients resulting from the O₂ uptake (Fig. 6); similar observations have been reported for analogous systems with alanine and asparagine [38]. The low energetic asymmetric d–d band in curve (a) can be attributed to the asymmetric, quasi-octahedral T _1g→T _1g(P) transition of the Co(H₂O) ₆ ²⁺ aquo-ion. Curve (b) is a spectral curve mainly characterizing the formed heteroligand CoL₂L′ active complex, predominating at pH ~9 under oxygen-free conditions, with a blue-shifted d–d band at λ _max 485 nm (ε _max ~20). Curve (c) corresponds to a μ-peroxo-type dioxygen adduct with two components of the LMCT band from the split antibonding π*(O₂) orbital of dioxygen to the unfilled dσ*(Co) orbital: π*_h → dσ* (in-plane) and π*_v→dσ* (out-of-plane). It can be seen that the molar absorption coefficient of both the bands (ε _max ~5 × 10⁴) is much higher than that of the “active” complex. The intensity of the two LMCT components was relatively comparable, this being typical of monobridged peroxo complexes, which are usually non-planar [38, 44].

As can be seen in Table 2, the highest value of log \(K_{{{\text{O}}_{2} }}\) occurs for the histidine system, most likely due to participation of the aforementioned additional effective N3 donor of the histidine imidazole side group. This is not surprising as it is already known that for histidine, the “active” complex is the most thermodynamically stable complex also under oxygen-free conditions [37]. On the other hand, the lower value of log  \(K_{{{\text{O}}_{2} }}\) for the asparagine system in comparison with the alanine system is most likely due to steric hindrance, which arises from one of the asparagine amide side groups during formation of the dimer. In this case, a greater share of the amino acid in the Co(II): amac: Himid molar ratio favorably displaces the equilibrium towards oxygen adduct formation. This also makes it possible to obtain chemically reasonable (positive) solutions of equation system (1) at higher excesses of the amino acid (cf. Table 2). For the two remaining amino acids alanine and histidine, particularly histidine, the oxygen adduct (for both the reversible and irreversible parts together) almost entirely uses up the accessible cobalt when the share of the amino acid greatly exceeds the stoichiometric ratio Co(II): amac:Himid = 1:2:1; the concentrations of the other complex species, including the “active” complex, fall to such low levels that it is impossible for the equation system (1) to converge in the form of three positive solutions.

At a decreased temperature close to 0 °C, the Co(II)–amac–Himid systems demonstrate enhanced reversible uptake of molecular oxygen. Coordination of the dioxygen molecule by the “active” complex occurs as exchange of the axial H₂O or carboxyl oxygen to O₂, occurring together with simultaneous formal intramolecular redox oxidation of Co(II) to Co(III) and the reduction of the charge of the dioxygen molecule to a bridging peroxide O₂ ²⁻ ion.

The log  \(K_{{{\text{O}}_{2} }}\) values are highest for the oxygenated forms of the heteroligand complexes with histidine, as their coordination sphere is formed by a chelating tridentate ligand (with imidazole, NH₂, COO⁻ donors). The essential impact on the electron structure of the dioxygen bridge, and by that on reversibility of O₂ uptake, is due to the first of the groups mentioned above. The two remaining amac ligands engaged in the mixed complexes (i.e. alanine and asparagine) were bidentate ligands. Even the potentially tridentate l-α-asparagine behaves as a bidentate ligand in the attainable pH range of around 9–10, illustrated in Table 1, which follows also from the previous reports concerning oxygen–free conditions. However, the reversibility of O₂ uptake in the latter systems containing an axial imidazole, unequivocally indicates a much higher reversibility than that previously reported for Co(II)–amac systems in the absence of imidazole.

Chemically reasonable (positive) values of both [Co(II)], [amac], [Himid] equilibrium concentrations and hence, appropriate log \(K_{{{\text{O}}_{2} }}\) values, could be attained only for limited Co(II):amac:Himid molar ratios, irrespective of the degree of equilibrium displacement towards oxygen adduct formation.