Complexes 1–6 were prepared in high yield as detailed in Section 3.3. Elemental analyses establish the formula MnL¹(H₂O)₂(ClO₄) for 1, MnL¹(H₂O)₂(AcAc) for 2, MnL²(H₂O)₂(NO₃)(CH₃OH) for 3, MnL²(H₂O)(DCA) for 4, MnL³(H₂O)₂(ClO₄) for 5, and MnL⁴(H₂O)₂(NO₃)(H₂O)₃ for 6. Molar conductivity measurements in 10⁻³ M dimethylformamide solutions, in the range 68–115 µS, show the electrolyte behavior for 1, 3, 5 and 6, while 2 and 4 show non-electrolyte behavior.

ESI (electrospray-ionization) mass spectra registered in methanol show peaks corresponding to the fragment [MnL]⁺ for all the complexes, indicating the coordination of the Schiff base ligand to the metal centre. Other minor signals could be assigned to [Mn₂L₂(X)]⁺ units (being X the anion), which could be attributed to the presence of dimeric species.

Values for the magnetic moments at room temperature are very close to the spin-only value of 4.9 B.M., as expected for a high-spin magnetically diluted d⁴ manganese(III) ion.

Infrared spectroscopy (IR) verifies the coordination of the Schiff base ligand to the manganese ion, since in all cases we observe a negative shift (4–10 cm⁻¹) of the ν(CN_imine) band and also a positive shift (10–30 cm⁻¹) of the ν(CO_phenol) mode with respect to the free ligand. These data suggest the coordination of the Schiff base in its dianionic form through the inner phenol oxygen and the imine nitrogen atoms. Some other IR spectroscopical features allow understanding how the former manganese salt anions are bound to the central ion or even indicating their coordination behavior:

The broad unsplit band at 1120 cm⁻¹ in 1 and 5 is indicative of the presence of the uncoordinated perchlorate anion [12], its characteristic ν₄ stretching mode at ca. 635 cm⁻¹ can also easily be identified;

Single crystals of Complex 4, suitable for X-ray diffraction studies, were obtained by slow evaporation of the methanolic solution at room temperature. Intensity data were collected on a Bruker X8 APEXII difractometer employing graphite-monochromated Mo-Kα radiation (λ = 0.71073 Å) at 100 K.

The geometry around the manganese ion can be described as distorted octahedral. Main bond distances and angles are collected in Table 1. An ORTEP view of 4 with the atomic numbering scheme is shown in Figure 2.

The coordination sphere around each manganese center comprises the planar Schiff base ligand, tightly bound to the metal ion through the inner N₂O₂ compartment by the N_imine and O_phenol atoms (Mn–N_imine bond lengths of 1.98–1.99 Å and Mn–O_phenol of 1.88 Å which are typical of such complexes and corroborate the bisdeprotonation of the ligands), occupying the equatorial positions and giving rise to two six-membered chelate rings (which are nearly planar) and an additional five-membered chelate ring.

The axial positions of the octahedron are occupied by a capping water molecule and a dicyanamide molecule. The Jahn-Teller elongation expected for d⁴ high-spin manganese(III) appears outstanding in the axis orthogonal to the plane of the Schiff base ligand, with distances ranging from 2.26 to 2.30 Å, considerably longer than the equatorial Mn–O bond lengths quoted earlier [12,13,14,15].

The superstructure of 4 involves associations via a combination of π-aryl offset interactions [20] and hydrogen bonds between capping water molecules and both phenoxy and methoxy oxygen atoms of the neighboring Schiff base ligand, forming μ-aquo dimeric units. These hydrogen bonds are charge-assisted, that is, the hydrogen bond donor and/or acceptor carry positive and negative ionic charges, respectively, and hence are rather strong and short. As result of these supramolecular interactions, the Mn···Mn distances of about 4.84 Å are short for monomeric compounds.

The electrochemical behavior of complexes 1, 3, 4 and 5 has been studied. Cyclic voltammograms of the complexes exhibit a quasi-reversible one-electron reduction-oxidation wave, which at slow scan rates of about 0.02 V s⁻¹ is reversible (Figure 3, Table 2).

The shape of the cyclic voltammograms for all the complexes is quite similar, with a peak-to-peak separation from 0.080 to 0.108. The rate of reversibility of these compounds suggests that they could act as catalysts in different processes. The redox potentials of 1, 4 and 5 appear in the same range, the values of 5 being slightly more positive. The reduction wave of 3 is observed at more negative potentials (E_red = −0.237 V).

Normal pulse voltammetry can be applied to such systems and provides an additional proof of the oxidation state of the manganese in the former complexes. Anodic and cathodic currents were observed when the initial potentials were more negative than the wage range. However, when the initial potentials were more positive than the wave voltage range, only cathodic currents were observed. These data indicate that in solution only the oxidized forms of the redox systems exist, i.e., manganese(III).

ABTS is colorless and it reacts readily with H₂O₂ in the presence of a peroxidase catalyst to yield a stable green colored radical cation ABTS^·+ [21], which presents characteristic absorption bands at 415, 650, 735 and 815 nm. The extent of the reaction can be measured quantitatively at λ = 650 nm since ε = 12000 M⁻¹ cm⁻¹ has been determined. The oxidation potential of ABTS to provide ABTS^·+ is invariable over a wide range of pH. In the absence of the complex, a solution of ABTS and H₂O₂ is stable for several hours without showing any formation of ABTS^·+.

The reaction of ABTS with H₂O₂ in the presence of 1–4 generates ABTS^·+ and the characteristic absorption bands of this species could be established (Figure 4). The rate of formation of ABTS·⁺ of about 40–62% indicates a relevant peroxidase activity by Complexes 1–4. Complexes 5–6 do not show significant peroxidase-like activity since the UV absorbances corresponding to the ABTS^·+ radical cation are negligible in the assays with these complexes.

The extent of the conversion of ABTS increases to 70–76% using a larger amount of the oxidant (250 μL 10 M). However the higher turnovers numbers (TON; calculated as the number of moles of substrate that a mole of catalyst can convert before becoming inactivacted) of catalytic cycles is achieved by reducing the concentration of the catalyst in solution (Figure 5). This behavior can be explained by the reaction completion so that more turnovers are observed with lower catalyst load. The percentages of conversions in these conditions range from 53 to 61%.

The exact mechanism of oxidation process is not clear at present. High-valent intermediates such as Mn(IV) or Mn(V) have been proposed in different redox processes catalyzed by other synthetic manganese complexes [22,23,24,25]. Accordingly, Complexes 1–4 could follow a pathway involving the interconversion between the Mn^III₂ and [Mn^IV=O]₂ forms. The resolution of the crystal structure of 4 reveals a short distance between the manganese ions. Both the short intermetallic separation and the occurrence of two labile coordination sites could facilitate the µ-bridging mode of peroxide leading to O-O cleavage and formation of the [Mn^IV=O]₂ form of catalyst.

The way to achieve the dimeric nature of the complexes is decisive on their peroxidase activity since the oxidation of the hydroperoxide ion involves an intramolecular two-electron transfer reaction which is forbidden for a monomeric Mn(III) complex. In this sense, the self-assembly of the manganese complexes through hydrogen bonding arises as a key issue to enhance the peroxidase activity for this type of complex [13]. The methoxy groups of the outer compartments of the Schiff bases H₂L¹-H₂L⁴ are essential to establish a supramolecular network that approximates the manganese ions to distances below 5 Å, which is considerably shorter than the 10 Å for monomeric compounds [26].

Veratryl alcohol (3,4-dimethoxybenzyl alcohol) is a commonly used lignin model [27]. This substrate is optimal for our experiments since Complexes 1–6 have no absorbance at 310 nm, whereas its oxidized product veratraldehyde absorbs strongly at 310 nm (ε = 9300 M⁻¹ cm⁻¹) [28]. The oxidation studies have been done both in presence and absence of air flow bubbling in order to check for the effect of a saturated atmosphere of molecular oxygen in the process. Dioxygen is an attractive choice as a nonpoisonous and inexpensive oxidant for industrial processes. The processes that are based on molecular oxygen can also be considered to be ecologically benign processes, because water and hydrogen peroxide are the main side products derived from the oxidant [29].

Veratraldehyde product were isolated and purified following the procedure described in Section 3.6. ¹H nuclear magnetic resonance (NMR) spectroscopy confirms the absence of the methylene proton signal at 4.61 ppm from the original veratryl alcohol as well as the appearance of a new peak at 9.62 ppm corresponding to the aldehyde. A control reaction under the same conditions without catalyst did not give any fraction of veratraldehyde.

The results of the oxidation of veratryl alcohol to veratraldehyde are shown in Figure 6. The activity of the catalysts 1–6 decreases as the length of the alkyl bridge between the coordinating amines increases.

Complexes 1, 2 and 4, behave as better catalysts in this process. An explanation arises from the short two-carbon chain between imine groups in these complexes; the short chain length constricts the chelate ring once nitrogens coordinate to the metal and it leads to a tetragonally elongated octahedral geometry. An axial water molecule in this class of distorted geometries constitutes a quite labile ligand, which would generate a vacant position in the coordination sphere to accommodate the substrate molecule. On the other hand, the flexible three-membered alkyl chain between the imine groups in the Mn compounds 5–6 favors a better stabilization of a high-symmetry octahedral geometry, which subsequently makes the generation of a vacancy difficult. The activity of 3, intermediate between these two groups of complexes, may be derived from its more negative potentials observed in electrochemical studies.

Conversions up to 30% were obtained at 22 °C with an air flow bubbling for 10 h, using the Catalyst 4 and three equivalents of H₂O₂ respective to the substrate. These are mild conditions compared with the experimental setup of other authors [30,31], where high temperatures and pressures are needed to reach similar yields with different catalysts. Moreover, most of the oxygen-activating homogeneous transition-metal catalysts are investigated in organic solvents [32,33]; however, only a few are known to be active in water [34].

No conversion has been shown in absence of catalyst or in absence of hydrogen peroxide. Air flow providing oxygen as oxidant is crucial to increase the yield, but a degradation of about 7% of veratryl alcohol is observed when the experiment was done without air flow. However, presence of both catalyst and H₂O₂ is necessary to find some degradation of veratryl alcohol. This behavior suggests that hydrogen peroxide could induce the oxidation of the catalyst to a more catalytically active species, such as [Mn^IV=O]₂ dimers, which subsequently would degrade the veratryl alcohol using dioxygen as oxidant.

All the starting materials (Aldrich) and solvents (Probus) used for the synthesis were of commercially available reagent grade and were used without further purification.

Elemental analyses were performed on a Carlo Erba Model 1108 CHNS-O elemental analyzer. The IR spectra were recorded as KBr pellets on a Bio-Rad FTS 135 spectrophotometer in the range of 4000–400 cm⁻¹. ¹H NMR spectra were recorded on a Bruker AC-300 spectrometer using CD₃OD (296 K) as solvent and SiMe₄ as an internal reference. The electro-spray mass spectra of the compounds were obtained on a Hewlett-Packard model LC-MSD 1100 instrument (positive ion mode, 98:2 CH₃OH-HCOOH as mobile phase, 30 to 100 V). Room-temperature magnetic susceptibilities were measured using a digital measurement system MSB-MKI, calibrated using mercury tetrakis(isothiocyanato)cobaltate(II) Hg[Co(NCS)₄] as a susceptibility standard. Conductivities of 10⁻³ M solutions in DMF were measured on a Crison microCM 2200 conductivity meter.

Electrochemical experiments were performed using an EG&G PAR model 273 potentiostat, controlled by EG&G PAR model 270 software. A Metrohm model 6.1204.000 graphite disc coupled to a Metrohm model 628-10 rotating electrolyte device was used as a working electrode. A saturated calomel electrode was used as a reference and a platinum wire used as an auxiliary electrode. All measurements were made with ca. 10⁻³ mol dm⁻³ solutions of the complexes in dimethylformamide using 0.2 mol dm⁻³ NBu₄PF₆ as a supporting electrode. Cyclic voltammetry measurements were performed with a static graphite electrode, whilst direct-current and pulse voltammograms were recorded with the graphite disc rotating at 2000 revolutions per minute.

Mn₂L¹₂(H₂O)₄(ClO₄)₂ (1), 1.99 mmol (1.00 g) of H₂L¹ was dissolved in 100 mL of a 1:1 methanol-ethanol mixture and 1.99 mmol (0.72 g) of Mn(ClO₄)₂.6H₂O was added to the initial yellow solution which changed to green (CAUTION: Although no problems were encountered in this work, perchlorates are potentially explosive and should be handled in only small quantities and with care!). After stirring for 10 min, 3.98 mmol (0.16 g) of NaOH, dissolved in a small quantity of water, was added, and the mixture turned dark. The progress of the reaction was followed by thin-layer chromatography (TLC) for three days and the mixture was then filtered. The complex was obtained from the filtrate as a brown solid after crystallization. It was isolated by filtration and washed with diethyl ether and dried in air. Anal. Calcd. for C₃₈H₄₄Br₄Cl₂Mn₂N₄O₁₆ (1312.4): C, 34.7; H, 3.4; N, 4.3. Found: C, 34.0; H, 3.4; N, 4.0 %. MS ES (m/z): 553 [MnL]⁺. IR (KBr, cm⁻¹): ν(O-H) 3438 (m), ν(C=N) 1625 (vs), ν(C-O) 1295 (s), ν₃(ClO₄⁻) 1120 (vs). μ = 4.7 BM. Conductivity (in DMF) Λ_M = 82 μS.

Mn₂L¹₂(H₂O)₄(AcAc)₂ (2), H₂L¹ (1.4 mmol, 0.70 g) was dissolved in methanol (40 mL) and Mn(AcAc)₃·4H₂O (1.4 mmol, 0.40 g), dissolved in methanol was added to the initial yellow solution, which changed to green. After stirring for 10 min, NaOH (2.8 mmol, 0.11 g), dissolved in a small quantity of water, was added and the mixture turned black. The reaction mixture was refluxed for 2 hours, and then concentrated in vacuo to half its volume. The complex was obtained as black crystals, which were isolated by filtration, washed with diethyl ether and dried in air. Anal. Calcd. for C₄₈H₅₈Mn₂Br₄N₄O₁₆ (1375.8): C, 41.8; H, 4.2; N, 4.1. Found: C, 40.4; H, 3.9; N, 4.3 %. MS ES (m/z): 553 [MnL]⁺. IR (KBr, cm⁻¹): ν(O-H) 3432 (m), ν(C=N) 1626 (vs), ν(C-O) 1298 (s), ν_asym(CO₂) 1546 (s), ν_sym(CO₂) 1438 (m). μ = 5.0 BM. Conductivity (in DMF) Λ_M = 10 μS.

Mn₂L²₂(H₂O)₄ (NO₃)₂(CH₃OH)₂ (3). 1.00 mmol (0.33 g) of H₂L² was dissolved in methanol (50 mL) and 1.00 mmol (0.26 g) of Mn(NO₃)₂·4H₂O was added to the initial yellow solution which changed to brown. After 3 h of stirring at room temperature slow evaporation of solvent lead to deposition of a brown compound. The product was collected by filtration, washed with diethyl ether and then dried in vacuo. Anal. Calcd. for C₃₈H₅₂Mn₂N₆O₂₀ (511.4): C, 44.6; H, 5.1; N, 8.2. Found: C, 44.5; H, 4.9; N, 8.3%. MS ESI (m/z): 381 [MnL¹]⁺, 824 [Mn₂L¹₂(NO₃)]⁺. IR (KBr, cm⁻¹): ν(O-H) 3420 (m), ν(C=N) 1619 (vs), ν(C-O) 1257 (s), ν(NO₃) 1384 (vs), 830 (m), 738 (s). ¹H NMR (DMSO-d₆, ppm): δ −30.52 (H4), −21.50 (H5). μ = 5.1 BM. Λ_M = 81 μS.

Mn₂L²₂(H₂O)₂(DCA)₂ (4), 0.46 mmol (0.15 g) of H₂L² was dissolved in methanol (40 mL) and 0.46 mmol (0.11 g) of Mn(CH₃COO)₂ was added to the initial yellow solution which changed to brown. The mixture was gently heated and stirred for 30 min, and then 0.46 mmol (0.04 g) of NaN(CN)₂ in 10 mL of methanol was added. Anal. Calcd. for C₄₀H₄₀Mn₂N₁₀O₁₀ (930.8): C, 51.6; H, 4.3; N, 15.0. Found: C, 50.8; H, 4.4; N, 14.8 %. MS ES (m/z): 381 [MnL]⁺; 828 [MnL(DCA)]⁺. IR (KBr, cm⁻¹): ν(O-H) 3421 (m), ν(C=N) 1624 (vs), ν(C-O) 1259 (s), ν_sym(C≡N) 2148 (vs), ν_asym(C≡N) 2266 (m). μ = 5.0 BM. Conductivity (in MeOH) Λ_M = 33 μS.

MnL³(H₂O)₂(ClO₄) (5), Synthetic procedure similar to 1, using 2.00 mmol (1.0 g) of H₂L³, 2.00 mmol (0.72 g) of Mn(ClO₄)₂.6H₂O and 4.00 mmol (0.16 g) of NaOH. Anal. Calcd. for C₁₉H₂₂Br₂ClMnN₂O₁₀ (688.2): C, 33.1; H, 3.2; N, 4.1. Found: C, 33.5; H, 3.2; N, 4.0 %. MS ES (m/z): 554 [MnL]⁺. IR (KBr, cm⁻¹): ν(O-H) 3431 (m), ν(C=N) 1624 (vs), ν(C-O) 1294 (s), ν₃(ClO₄⁻) 1119 (vs). μ = 4.6 BM. Conductivity (in DMF) Λ_M = 115 μS.

MnL⁴(H₂O)₂(NO₃)(H₂O)₃, (6). Synthetic procedure similar to 3, using 0.68 mmol (0.25 g) of H₂L⁴, and 0.68 mmol (0.18 g) of Mn(NO₃)₂.4H₂O. Anal. Calcd. for C₂₁H₃₂MnN₃O₁₁ (557.4): C, 45.2; H, 5.8; N, 7.5. Found: C, 45.3; H, 5.7; N, 7.5%. MS ESI (m/z): 423 [MnL⁵]⁺, 908 [Mn₂L⁵₂(NO₃)]⁺. IR (KBr, cm⁻¹): ν(O-H) 3427 (m), ν(C=N) 1609 (vs), ν(C-O) 1254 (s), ν(NO₃) 1384 (vs), 850 (m), 739 (s). ¹H NMR (DMSO-d₆, ppm): δ -21.97 (H4), -18.00 (H5). μ = 5.3 BM. Λ_M = 68 μS.

Single crystals of Complex 4 suitable for X-ray diffraction studies were obtained by slow evaporation of the methanolic solution at room temperature. Detailed crystal data collection and refinement are summarized in Table 3. The structures were solved by direct methods [35] and finally refined by full-matrix least-squares base on F². An empirical absorption correction was applied using sadabs [36]. All non-hydrogen atoms were included in the model at geometrically calculated positions.

The Cambridge crystallographic data centre (CCDC) reference number 916337 contains the supplementary crystallographic data for 4. These data can be obtained free of charge at www.cdcc.cam.ac.uk/conts/retrieving.html (or from the Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK [Fax: + 44-1223/336-033; E-Mail: deposit@ccdc.cam.ac.uk].

Oxidation of ABTS with H₂O₂ at ca. pH 7 in the presence of the complexes was tested in the following manner. An aqueous solution of ABTS (50 μL; 0.009 M; 4.5 × 10⁻⁷ mol) and a methanolic solution of the complex (10 μL; 10⁻³ M; 10⁻⁸ mol) were added to water (3 mL). The intensity of the UV absorption bands of ABTS started to increase immediately after addition of an aqueous solution of H₂O₂ (50 μL; 10 M; 5 × 10⁻⁴ mol). The disproportionation of hydrogen peroxide by the Complexes 1–6 (catalase activity) was not observed for the concentrations of this peroxidase assay.

Veratryl alcohol (200 mmol, 0.29 mL), previously dissolved in 20 mL of dihydrogenphosphate/NaOH buffer pH 8, was mixed with a methanolic solution (15 mL) of the catalytic complex (0.5 mmol, ca. 6 mg). The oxidant (600 mmol) was added in three portions at 20 min intervals. The reaction mixture was stirred for 10 hours at 22 °C with an air flow bubbling through the solution and then filtered through a short silica gel plug to remove the catalyst and excess oxidant. The reaction mixture was extracted in CH₂Cl₂ in the presence of saturated aqueous NaCl solution. The organic layer was dried over MgSO₄ and evaporated under reduced pressure. The mixture was then purified by column chromatography using a 1:1 mixture of ethyl acetate:hexane as eluent. The veratraldehyde fraction was characterized by NMR spectrometry. ¹H NMR (300 MHz) for the veratraldehyde fraction: 9.62 (s, 1H); 7.44 (d, 1H); 7.33 (s, 1H); 7.18 (d, 1H); 3.83 (s, 6H); ¹H NMR (300 MHz) for original commercial veratryl alcohol: 7.03 (s, 1H); 6.81 (d, 1H); 6.55 (d, 1H); 4.61 (s, 2H); 3.83 (s, 6H).