DALTON J. Chem. Soc. Dalton Trans. 1996 Pages 65–73 65 Metal complexes of a tetraaza macrocycle with N-carboxymethyl groups as pendant arms* Judite Costa,a,b Rita Delgado,a,c M. do Carmo Figueira,a Rui T. Henriques c,d and Miguel Teixeira a,e a Instituto de Tecnologia Química e Biológica Rua da Quinta Grande 6 2780 Oeiras Portugal b Faculdade de Farmácia de Lisboa Av. das Forças Armadas 1600 Lisboa Portugal c Instituto Superior Técnico Av. Rovisco Pais 1096 Lisboa codex Portugal d Instituto Tecnológico e Nuclear Dep. Química 2686 Sacavém codex Portugal e Universidade Nova de Lisboa Faculdade de Ciências e Tecnologia Monte da Caparica Portugal The protonation constants of H2L1 (7-methyl-3,7,11,17-tetraazabicyclo[11.3.1]heptadeca-1(17),13,15-triene-3,11- diacetic acid) and stability constants of its complexes with Mg2+ Ca2+ Co2+ Ni2+ Cu2+ Zn2+ Cd2+ Pb2+ Ga3+ Fe3+ and In3+ were determined by potentiometric methods.This macrocycle is selective exhibiting a very high stability constant for Cu2+ (log KML = 21.61) while for the remaining complexes of the first-row transition-metal ions the stability decreases sharply. The calcium complex has very low stability (log KML = 3.74) but that of Mg2+ is higher (log KML = 5.30) and those of trivalent metal ions (Ga3+ Fe3+ and In3+) have very low stability the value for Fe3+ (log KML = 20.64) being lower than that of Cu2+. Spectroscopic studies (electronic and EPR) in aqueous solution have shown that H2L1 upon complexation (with Co2+ Fe3+ and possibly Cu2+) exhibits a tendency to adopt a folded conformation with the nitrogen atom of the macrocycle opposed to the pyridine in the axial position the basal plane being formed by the three other nitrogen atoms of the ring and the other carboxylate in a square-pyramidal arrangement for the metal ion.The complex [Fe(L1)Br]?4H2O shows a temperature-dependent magnetic behaviour with meff ranging from 3.58 mB at 292 K to 1.70 mB at 3.1 K. The search for new compounds which may form selective complexes with the first-row transition-metal ions led us to study a N-acetate derivative of a pyridine-containing tetraazamacrocycle (H2L1) synthesized before 1 and for which the crystal structure of a copper(II) complex [Cu(H2L1)Cl]Cl was described.1 It is known that N-carboxymethyl derivatives of 14-membered macrocycles are more selective for the first-row transition divalent metal ions than are 12- or 13-membered rings,2–4 although the latter form complexes with higher stability constants.2,4 In general the complexes of those ligands obey the Irving–Williams order of stability,5 without surprising inversions.2,4 For alkaline-earth-metal ion complexes the stability constants decrease with increasing size of the cavity any metal ion being particularly favoured.The complexes of these ions with 14-membered macrocyclic ligands having contiguous propane chains exhibit very low stability constants.3 Although no surprising inversions in the usual trends of stability constants were observed some remarkably stable complexes (thermodynamic and kinetically) were found namely the 1,4,7,10-tetraazacyclododecane-1,4,7,10- tetraacetate (dota) complexes of Ca2+ (ref.2) and of some trivalent lanthanides.6,7 The compound studied in the present work although having two acetate arms has a complexing behaviour which is closer to that of linear polyamines than to cyclic polyaminocarboxylate ligands taking the advantage of some properties of the latter namely water solubility easy purification neutral complexes with divalent metal ions and faster kinetics of formation of metal complexes. The completely deprotonated form of the macrocycle is (L1)22 but for simplicity the charge is generally disregarded in the text. * Non-SI units employed mB ª 9.27 × 10224 J T21 G = 1024 T. Experimental Reagents The parent L2 was synthesized in our laboratory by previously reported procedures.8 Bromoacetic acid and Dowex 1 × 8 ionexchange resin (20–50 mesh Cl2 form) were obtained from Aldrich Chemical Co.H4egta (ethylenedioxydiethylenedinitrolotetraacetic acid) from Sigma and K2H2edta (edta = ethylenedinitrilotetraacetate) from Fluka. All the chemicals N N N N R3 R1 R1 R2 R2 HO2C N N CO2H CO2H N N HO2C H4trita N N N N HO2C CO2H HO2C CO2H D6/02948E/A1 H2L1 R1 = H; R2 = CH2CO2H; R3 = CH3 L2 R1 R2 = H; R3 = CH3 L3 R1 R2 R3 = H L4 R1 = H; R2 R3 = CH3 L5 R1 R3 = H; R2 = CH2Ph L6 R1 R2 R3 = CH3 L7 R1 R3 = CH3; R2 = H L8 R1 = CH3; R2 R3 = H H4dota H4teta N N N N HO2C CO2H HO2C CO2H 66 J. Chem. Soc. Dalton Trans. 1997 Pages 65–73 were of reagent grade and used as supplied without further purification (the resin was treated with 5% KOH before use). The organic solvents were purified by standard methods.9 Synthesis and characterization of the macrocycle The N-carboxymethyl groups were introduced by condensation of the parent amine (4.03 mmol 1.00 g) with potassium bromoacetate (obtained by addition of 3 mol dm23 KOH solution to concentrated aqueous bromoacetic acid 8.1 mmol 1.13 g at 5 8C) in aqueous basic solution (4 cm3).The temperature was kept at 25–30 8C and the pH < 9 by slow addition of 3 mol dm23 KOH. At the end of the reaction the mixture was cooled and adjusted to pH 1.9 with 3 mol dm23 hydrobromic acid. The solution was then concentrated and a small amount of methanol added. The inorganic matter formed was filtered off and the filtrate purified by chromatography using an anionic resin in the OH2 form washed with water and then eluted with a solution of 0.1 mol dm23 HBr. Yield 63%. M.p. 233–235 8C; 1H NMR (D2O) [sodium 4,4-dimethyl-4-silapentane-1-sulfonate (dss) as reference] d 7.76 (1 H t) 7.40 (2 H d) 4.52 (4 H s) 4.01 (4 H s) 3.22 (8 H m) 2.66 (3 H s) and 2.17 (4 H m); 13C (1,4- dioxane) d 168.36 149.47 140.41 126.53 56.44 55.90 51.28 49.99 41.89 and 17.67 (Found C 31.6; H 5.9; N 8.0.Calc. for C18H31Br3N4O4?4H2O C 31.8; H 5.8; N 8.3%). Synthesis of [Fe(L1)Br]?4H2O Iron(III) hydroxide was freshly prepared by addition of 0.1 mol dm23 KOH to Fe(NO3)3 (0.098 mmol). The precipitate formed was separated by centrifugation and added to an aqueous solution of [H5L1]Br3 (0.089 mmol). The mixture was heated for 1 h at 40 8C and stirred overnight at room temperature. The solution was then concentrated and diethyl ether was added to the residue. An orange precipitate was obtained which was dried under vacuum.Yield 70% (Found C 38.1; H 5.75; N 9.45. Calc. for C18H26BrFeN4O4?4H2O C 37.9; H 6.0; N 9.85%). Potentiometric measurements Reagents and solutions. Metal-ion solutions were prepared at about 0.025 mol dm23 from the nitrate salts (analytical grade) with demineralized water (obtained by a Millipore/Milli-Q system) and standardized by titration with K2H2edta.10 A back titration with a standard solution of ZnSO4 was necessary for Ga3+. The solutions of the trivalent metal ions were kept in an excess of nitric acid to prevent hydrolysis. Carbonate-free solutions of the titrant NMe4OH were prepared as described.11 Solutions were discarded when the percentage of carbonate was about 0.5% of the total amount of base. Equipment and working conditions. The equipment used was as described.11 All the experiments were monitored by computer.The temperature was kept at 25.0 ± 0.1 8C; atmospheric CO2 was excluded from the cell during the titration by passing purified N2 across the top of the experimental solution in the reaction cell. The ionic strength of the solutions was kept at 0.10 mol dm23 with NMe4NO3. Measurements. The [H+] of the solutions was determined by measurement of the electromotive force of the cell E = E98 + Q log [H+] + Ej where E98 and Q were obtained by previous calibration titrating a standard solution of known hydrogen-ion concentration at the same ionic strength using the values of the acid range. The term pH is defined as 2log [H+]. The liquidjunction potential Ej was found to be negligible under the conditions used. The value of KW = [H+][OH2] was determined from the alkaline region of the calibration considering E98 and Q valid for the entire pH range and found equal to 10213.80 mol2 dm26.The potentiometric equilibrium measurements were made on macrocycle solutions (ª2.50 × 1023 mol dm23 20.00 cm3) diluted to a final volume of 30.00 cm3 first in the absence of metal ions and then in the presence of each metal ion the cL cM ratios being 1 1 1 2 and in several cases 2 1. A minimum of two replicates were made. The E data were taken after additions of 0.025 or 0.050 cm3 increments of standard NMe4OH solution and after stabilization in this direction equilibrium was then approached from the other direction by adding standard nitric acid. In the cases of Cu2+ Ga3+ Fe3+ and In3+ the extent of formation of the metal complexes at the beginning of the titration was too high for the use of the direct potentiometric method and so ligand–ligand competition titrations were performed to determine the constants; K2H2edta was used as the second ligand in the cases of Cu2+ and In3+ and H4egta for Fe3+.Ratios cL 1 cL9 cM 1 1 1 1 0.7 1 and 1.5 1 1 were used respectively; L9 is the reference ligand K2H2edta or H4egta for which it is necessary to know accurately the stability constants of complexes with the same metal ion.12 A competition reaction which can be written in terms of equilibrium (1) was considered [ML9]n2 + [H5L1]3+ [ML1]m+ + 3 H+ + [H2L9]22 (1) appropriate when all complexed species exist in solution at least 30% concentration with respect to the total metal ion; n = 1 (for trivalent metal ion) or 2 (for divalent metal ion) and m = 0 (for divalent metal ion) or 1 (for trivalent metal ion).Stability constants for the gallium complexes were calculated by relying on the competition or displacement reaction (2).13 [GaL1(OH)] + 3OH2 [Ga(OH)4]2 + (L1)22 (2) Constants corresponding to the formation of [GaL1]+ and [GaL1(OH)] can be determined at pH > 6 when [Ga(OH)4]2 starts to be formed and used in other parts of the titration curves (at low pH) to obtain the constants for the other equilibrium reactions. As the value of the stability constant for the iron(III) complex was lower than expected other techniques were used to confirm the value. A competition with another metal ion Cu2+ in the ratio cL 1 cM cM9 1 0.5 0.5 and also a direct redox method using a couple of platinum and reference electrodes to follow the Fe3+–Fe2+ equilibrium at pH 2.14 In the competition reactions the equilibria were slow to attain but even so automated titrations were possible; 10–15 min were necessary at each point of the titration in the pH range where the competition reaction took place.The same values of stability constants were obtained either using the direct or the back titration curves. Calculation of equilibrium constants. Protonation constants Ki H = [HiL]/[Hi21L][H] were calculated by fitting the potentiometric data obtained for the free macrocycle using the SUPERQUAD program.15 Stability constants of the various species formed in solution were obtained from the experimental data corresponding to the titration of solutions of different ratios of the macrocycle and metal ions also with the aid of the SUPERQUAD program.The initial computations gave overall stability constants or bMmHhLl = [MmHhLl]/[M]m[L]l[H]h. Only mononuclear species ML MHL MLH21 (bMLH21 = bML(OH)KW) were found. Differences in log units between bM(HL) (or bMLH21) and bML provide the stepwise protonation reaction constants. The species considered were limited to those which can be justified by established principles of co-ordination chemistry. The errors quoted are the standard deviations of the overall stability constants given directly by the program. In the case of the stepwise constants the standard deviations were determined by normal propagation rules and do not represent the total experimental errors. J. Chem. Soc. Dalton Trans. 1997 Pages 65–73 67 Hydrolysis of the trivalent metal ions.The trivalent metal ions studied easily form hydrolytic species in aqueous solution the constants of which show some discrepancies in the literature. We have used the values reported before considered more reliable.11 Spectroscopic studies Proton NMR spectra were recorded with a Bruker CXP-300 spectrometer. Solutions of the macrocycle for the measurements (ª 0.01 mol dm23) were made up in the D2O and the pD was adjusted by adding DCl or CO2-free KOD using an Orion 420A instrument fitted with a combined Ingold microelectrode. The 2log [H+] was measured directly in the NMR tube after the calibration of the microelectrode with buffered aqueous solutions; dss was used as an internal reference. The 13C NMR spectra were recorded with the same spectrometer and 1,4-dioxane as internal reference.The metal complexes were prepared in water by addition of the metal ion in the form of Co(NO3)2 or Fe(NO3)3 to an equivalent amount of the macrocycle and enough KOH to give a final pH of 5.8 and 2.8 respectively and after evaporation of water were taken up in D2O. The magnetic moment of the complexes were determined by the Evans method16 in solution at room temperature. Electronic spectra were measured with a Shimadzu model UV-3100 spectrometer for UV/VIS/near IR using aqueous solutions of the complexes prepared by addition of the metal ion (in the form of its nitrate salt) to the macrocycle at the appropriate pH value (corresponding to total formation of the metal complex). The X-band EPR spectra were recorded with a Bruker ESP 380 spectrometer equipped with a continuous-flow cryostat for liquid helium or for liquid nitrogen.The spectra of the complexes of Cu2+ and Fe3+ (9.0 × 1024 and 7.55 × 1023 mol dm23 respectively in 1.0 mol dm23 NaClO4) were recorded at 86 K for the first complex and in the range 4.6–200 K for the second. The quantification of signals of the iron complex was obtained by comparison of the second integral of its EPR spectra with that of an external standard corrected by the gp factor for the sample and the standard.17 As a standard we used a solution of hydrated copper(II) ion (5 mmol CuSO4?5H2O + 5 mmol HCl in 2.0 mol dm23 NaClO4). The simulation of the EPR spectrum of the copper complex was carried out with a program for a microcomputer.18 Magnetic susceptibility studies The magnetic susceptibility measurements for the iron(III) complex in the solid state were performed in the range 3–292 K using a longitudinal Faraday system (Oxford Instruments) and a 7 T superconducting magnet.A polycrystalline sample of [Fe(L1)Br]?4H2O (24 mg) was placed into a thin-walled Teflon bucket previously measured. The magnetic field used was 1 T and force was measured with a microbalance (Sartorius S3D-V) applying forward and reverse gradients of 5 T m21. The paramagnetic susceptibility was calculated from the raw susceptibility data correcting for diamagnetism estimated from Pascal’s constants 19 as 3.2 × 1024 cm3 mol21. Results Protonation and stability constants Titration of the macrocycle in the form [H5L1]Br3 showed two inflection points at a = 3 and 4 respectively (a being a number of equivalents of base added per mol of macrocycle).The protonation constants obtained are summarized in Table 1 together with the values for other tetraazamacrocycles for comparison. The titration curves obtained for mixtures of the macrocycle and metal ions (1 1) showed one inflection at a = 5 with the exception of those of Ga3+ and Fe3+ due to the formation of stable ML1(OH) species at low pH which had an inflection at a = 6. The curves for Mg2+ Ca2+ and Pb2+ exhibit another inflection at a = 3 at which formation of complexes starts. Titration curves corresponding to other mixtures like 2 1 or 1 2 were not significantly different from those of ratio 1 1. The values of the stability constants for the metal complexes of L1 studied in this work determined in water are also compiled in Table 1.In most cases only ML and M(HL) species are formed; but hydroxo complexes are found for some metal ions especially Ga3+ Fe3+ In3+ and also Zn2+. For Fe3+ and In3+ a precipitate is formed at pH ª 3.6 and 8 respectively; it is impossible to obtain reliable constants for the ML1(OH) species. In the case of iron(III) probably a m-oxo dimer[(FeL1)2O] will be formed. We have checked the possibility of formation of other species like protonated MHiL (i > 2) or polynuclear M2L but they do not appear to be formed under our conditions. As the overall basicity of L1 is not very high when compared with that of dota or teta the complexes of Cu2+ and trivalent metal ions (Ga3+ Fe3+ and In3+) were completely or almost completely formed even at low pH and it was impossible to determine the values of stability constants by direct potentiometry.However competition reactions (1) and (2) with another ligand (edta egta or OH2) enabled the determination of the constants. The values of the protonation and stability constants of edta and egta were taken from the literature 12a and are compiled in Table 2. Spectroscopic studies The UV/VIS/near-IR data for complexes of Co2+ Ni2+ Cu2+ and Fe3+ and EPR data for those of Cu2+ and Fe3+ with L1 in water solution are collected in Tables 3–5. The electronic spectrum of the cobalt complex exhibits a broad band at 484 nm (e = 267.9 dm3 mol21 cm21) with shoulders at 493 and 591 nm another at 342.5 nm (e = 362.8 dm3 mol21 cm21) and three smaller bands at 1400 1520 and 1670 nm ( = 10.0 9.0 and 18.0 dm3 mol21 cm21 respectively).The magnetic moment of the complex is 4.3 mB. This orange-pink complex is slow to form (4.5 h were needed to attain the maximum absorbance at pH 5.8) but did not suffer any degradation with time. The violet solution of the nickel complex exhibits three main bands at 940 570 and 355 nm (e = 19.4 18.1 and 86.1 dm3 mol21 cm21 respectively) and a shoulder at 810 nm (e = 18.1 dm3 mol21 cm21). The copper complex is blue exhibiting a broad band in the visible region at 630 nm with a shoulder at lower energy (at 780 nm) due to the d–d transitions and an intense band in the ultraviolet region. The spectrum of the yellow iron complex shows two intense peaks at 276 and 328 nm a small band at 534 nm with a shoulder at higher energy. The magnetic moment of this complex in aqueous solution is 3.4 mB at 301 K.The EPR spectra of the complexes of Cu2+ and Fe3+ are shown in Figs. 1 and 2. The spectrum of [CuL1] exhibits three well resolved lines of the four expected at low field due to the interaction of the unpaired electron spin with the copper nucleus and no superhyperfine splitting due to coupling with the four nitrogen atoms of the macrocycle. The fourth copper line is completely overlapped by the much stronger and unresolved band of the high-field part of the spectrum. The computational simulation of the spectrum 18 leads to three different principal values of g which indicates that the Cu2+ ion of this complex is in a rhombically distorted ligand field. The hyperfine coupling constants and g values are presented in Table 4. The EPR spectra of [FeL1]+ were recorded at several temperatures into the range 4.6–200 K.All the spectra exhibit two types of signals the intensity of which decreases with increasing temperature in the same proportion for both signals. One of the signals is typical of a rhombic iron(III) complex in the low-spin state with signals at 2.683 2.337 and 1.721 similar to those of iron(III) porphyrin complexes. The other signals can be ascribed to two high-spin iron(III) species. The concentration 68 J. Chem. Soc. Dalton Trans. 1997 Pages 65–73 Table 1 Protonation (log Ki H) constants of L1–L3 dota and teta and stability constants (log KMmHhLl) of some of their metal complexes (25.0 8C I = 0.10 mol dm23) Ion Equilibrium quotient L1 L2a L3a dota teta H+ [HL]/[H][L] [H2L]/[HL][H] [H3L]/[H2L][H] [H4L]/[H3L][H] [H4L]/[L][H]4 10.72(2) 7.74(4) 4.05(7) 1.8(1) 24.31 9.74 8.67 4.67 <1 <24 9.92 8.56 4.66 <1 <24 12.09 b 9.76 b 4.56 b 4.09 b 30.50 10.52 c 10.18 c 4.09 c 3.35 c 28.14 Mg2+ Ca2+ Co2+ Ni2+ Cu2+ Zn2+ Cd2+ Pb2+ Ga3+ Fe3+ In3+ [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[ML(OH)][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[ML(OH)][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[ML(OH)][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[ML(OH)][H] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[ML(OH)][H] [ML]/[M][L] [M(HL)]/[ML][H] 5.30(7) — 8.74(6) 3.74(2) — 14.4(1) 4.1(1) 16.59(1) 2.94(2) 21.61(4) 2.28(4) 14.01(2) 4.05(2) 7.66(6) 14.56(3) 3.84(6) 10.89(2) 4.96(6) 9.9(2) 18.02(5) 3.09(9) 3.75(5) 20.64(6) 2.84(7) (5.9) g 18.94(2) 2.38(5) ————————— 20.23 — 11.91 — 8.06 8.77 — 9.03 —————————— ——————— 16.27 — 19.76 — 12.82 — 8.48 9.76 — 9.72 — 10.95 ———————— 11.92 d 4.09 d — 17.23 d 3.54 d 20.27 b 4.08 b 20.03 d 3.51 d 22.25 b 3.78 b 21.099 b 4.178 b — 21.31 b 4.39 b 22.69 b 3.86 b — 21.33 f 4.00 f — 29.4 f 3.23 f — 23.9 f 3.44 f 1.97 e —— 8.32 e — 16.38 c 4.04 c 19.83 c 4.14 c 20.49 c 3.77 c 16.40c 4.10 c — 18.02 c 4.04 c 14.32 c 4.75 c — 19.91 f 3.66 f — 27.46 f 2.64 f — 23.00 f 3.33 f aI = 0.10 mol dm23 KNO3.8 bI = 0.10 mol dm23 NMe4NO3.2b cI = 0.10 mol dm23 KNO3.2b dI = 0.10 mol dm23 NMe4NO3.2a eI = 0.10 mol dm23 KNO3.2a fI = 0.10 mol dm23 KCl.20 gApproximate value only as precipitation occurs.of the low-spin species was evaluated at 9.0 149.0 and 200.0 K and 35 ± 5% was obtained for the three temperatures.The low-spin signal is typical of d5 systems in distortedoctahedral environments.27–29 In general in this configuration the g values vary widely and are sensitive to small changes in structure. They can be related to the parameters which describe the electronic ground state of the complex by coefficients A B and C which are related to the axial distortion m the rhombic distortion parameter R the energy of the Kramers doublet E1 and the spin–orbit coupling constant l (see equations in Table 5). The experimental spectra give only the magnitude of the g values the problem being to assign the values obtained to gx gy and gz and to determine the sign of each. Substitution of the three g values (having positive or negative sign) into the equations for A B and C gives 48 combinations but only six satisfy the normalization condition A2 + B2 + C2 = 1.From these six combinations only one satisfies the conditions that |m/l| is a maximum and R/l is positive. The values calculated for our Table 2 Protonation (log Ki H) constants of edta and egta and their stability constants (log KMmHhLl) with metal ions used in competition reactions 12a (25 8C I = 0.10 mol dm23) Ion Equilibrium quotient edta egta H+ [HL]/[H][L] [H2L]/[HL][H] [H3L]/[H2L][H] [H4L]/[H3L][H] 10.19 6.13 2.69 2.00 9.40 8.78 2.66 2.0 Cu2+ Fe3+ In3+ [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[ML(OH)][H] [ML]/[M][L] [ML]/[M][L] [M(HL)]/[ML][H] [ML]/[ML(OH)][H] 18.78 3.1 11.4 — 24.9 1.5 8.49 ——— 20.5 ——— complex together with those of other similar iron(III) complexes from the literature are compiled in Table 5.The signal of lower field is typical of a high-spin d5 state of Fe3+. Using the spin-Hamiltonian formalism for high-spin iron(III),30 the low-field resonances can be assigned to two species with different rhombic (E/D) distortions species I with E/D ª 0.275 and II with E/D ª 0.15. The effective g values expected for each Kramers doublet are indicated in Table 6 (those observed in the spectra Fig. 2 are underlined). Since the overall lineshape does not change over the entire temperature range studied it is not possible to determine the zero-field splitting accurately; the fact that for both species even at 4.6 K resonances from the |±3/2Ò and |±1/2Ò doublets are observed suggests a very small value for D. The presence of the two types of signals in the EPR spectra can be interpreted by an equilibrium of two spin states for the iron(III) complex or by the presence of different isomers.To understand this we have studied the temperature-dependent static magnetic susceptibility behaviour over the range 3.1– 291.7 K in the solid state. Plots of the magnetic susceptibility and magnetic moment as a function of temperatures are given in Fig. 3. The magnetic moment is 3.58 mB at 292 K and below ca. 100 K decreases gradually to 1.70 mB at 3.1 K. Table 3 Spectroscopic UV/VIS/near IR data for the complexes of Co2+ Ni2+ Cu2+ and Fe3+ with L1 Complex pH colour lmax/nm (e/dm3 mol21 cm21) [CoL1] 5.8 orangepink 1670 (18) 1520 (9) 1400 (10) 591 (sh) (9) 493 (sh) (37) 484 (267.9) 342.5 (362.8) [NiL1] 6.4 violet 940 (19.4) 810 (sh) (18.1) 570 (18.1) 355 (86.1) [CuL1] 3.95 blue 780 (sh) (25.1) 630 (97.1) 265 (2238) [FeL1]+ 2.77 yellow 534 (55.4) 458 (sh) (92.4) 328 (1305.4) 276 (1822.7) J.Chem. Soc. Dalton Trans. 1997 Pages 65–73 69 Table 4 Spectroscopic EPR data for the copper(II) complex of L1 and similar complexes l/nm (e/dm3 EPR g|| 2 2 Complex mol21 cm21) g1 g2 g3 104A1/cm21 104A2/cm21 104A3/cm21 Geometry g^ 2 2 Ref. [CuL1] [CuL3]2+a [CuL9]2+ [CuL10] [CuL11] [CuL12]2+ [CuL13]2+ [CuL14]2+ [Cu(dota)]22 [Cu(trita)]22 [Cu(teta)]22 630 (97.1) 560 (187) 695 (161) 599 (250) 513 (100) 622 (147) 626.4 (160) 690.6 (161) 734 (100) 719 (370) 646 (70) 2.027 2.084 2.034 2.060 2.033 2.084 2.057 2.049 2.050 2.059 2.027 2.082 2.037 2.077 2.062 2.047 2.050 2.221 2.188 2.210 2.198 2.186 2.224 2.216 2.226 2.300 2.202 2.249 14.9 21.3 0.5 3.4 26.6 38.9 24.1 38.7 10.9 20.5 26.9 15.1 23.8 21.6 — — — — — — 165.4 192.9 161.0 184.2 205.0 183.1 160.2 162.8 150.3 190.6 168.0 SP SPY SPY SP SPY c SPY SPY d OC — OC 3.98 4 3.59 3.47 3.79 4.11 3.96 3.96 4.84 4.29 4.98 This work 88 21 24 22 22 22 23 23 23 a This work.The values published 8 were redetermined because the EPR spectrum then obtained exhibited evidence of aggregation. A spectrum considerably better resolved was obtained when the solution was diluted 1 1 with 1 mol dm23 NaClO4. b SP = square planar SPY = square pyramid OC = octahedron. c The equatorial plane is formed by the four nitrogen atoms of the ring. d The equatorial plane with a degree of tetrahedral distortion. L9 = 3,6,9,15-Tetraazabicyclo[9.3.1]pentadeca-1(15),11,13-triene; L10 = 1,4,7,10-tetraazacyclododecane(cyclen); L11 = 1,4,8,11-tetraazacyclotetradecane( cyclam); L12 = 1-oxa-4,8,12-triazacyclotetradecane; L13 = 1-oxa-4,8,11-triazacyclotridecane; L14 = 1-oxa-4,7,10-triazacyclododecane.Table 5 Spectroscopic EPR data for the iron(III) complex of L1 and similar complexes Complex gz gy gx m/l R/Î E1/l R/m Ref. [FeL1]+a [FeL15]2+b [FeL11]3+ [FeL16]3+e 1.721 1.910 21.15 1.631 2.337 2.237 2.23 2.463 2.683 2.511 23.26 2.841 23.735 26.103 2.595 23.033 1.925 3.957 1.266 1.379 22.647 24.166 21.789 22.214 20.515 20.648 (0.019)c 0.488d 20.455 This work 26 26 25 A = (gx + gy 2 2gz)/4(gx + gy 2 gz)��� ; B = (gx + gy)/2[2(gx + gy 2 gz)]��� ; C = (gy 2 gx)/4(gx + gy 2 gz)��� ; m/l = {2A2 + B2 + C2 + 22��� [AB + (BC2/A)]}/ 2��� [2AB + (BC2/A)]; R/Î = (2AC + 2��� BC)/(C2 2 A2); E1/l = (2m/3l) 2 (A/2��� B). a Spectrum obtained at 12.4 K main signal.b HL15 = 1,4,8,11- Tetraazacyclotetradecane-1-acetic acid. c Value presented by Szulbinski and Busch.26 d Our calculation for the experimental values given in ref. 26. e L16 = 6,13-Dimethyl-1,4,8,11-tetraazacyclotetradecane-6,13-diamine. Discussion The macrocycle L1 has six basic centres nevertheless only four protonation constants were obtained; the last two are very low to be determined by potentiometric measurements. All the compounds listed in Table 1 have two high (or fairly high) and two low values of the protonation constants. Taking into account the sequence of protonation of L2 and L3 (ref. 8) and other similar N-carboxymethyl derivatives of tetraaza macrocycles 31,32 studied by 1H NMR spectroscopy the interpretation of the values of the protonation constants of L1 based on the following sequence is straightforward the first two protons added to the basic form of the macrocycle are spread over all the nitrogen atoms 50% of them remaining non-protonated; the third protonation will occur at carboxylate groups linked to Fig.1 The X-band EPR spectrum of the copper complex of L1 in 1.0 mol dm23 NaClO4 recorded at 86 K n = 9.41 GHz microwave power 2.4 mW and modulation amplitude 1 mT and its simulation non-protonated nitrogen atoms (the value is similar to that found for teta lower than that corresponding to protonation of a third nitrogen of the ring of the parent amine); the very low value of the fourth constant is typical of protonation of a carboxylate group linked to a protonated nitrogen atom31 and is certainly what occurs in this case.The last two nitrogen atoms of the ring will be only partially protonated at very low pH values. The second and the fourth protonation constants of L1 are lower than those of dota and teta due to the presence of the pyridine ring and the existence of only two carboxylate groups which contribute strongly to a lower overall basicity of L1 compared to that of teta or dota (a decrease of about 4 and 6 log units respectively) a situation which will have important repercussions in metal complex formation. Consideration of the stability constants in Table 1 and their variations in Fig. 4 allows some interesting conclusions. (1) Compound L1 is very selective exhibiting a very high stability constant for Cu2+ while for the remaining complexes of the first-row transition divalent metal ions the stability decreases sharply Zn2+ and Ni2+ having ML constants which are 7.6 and 5.02 log units lower than that of Cu2+.The alkaline-earth-metal ion complexes present very surprising aspects namely the very low value of the stability constant of the calcium complex and the value for Mg2+ being higher than that of Ca2+. Fig. 4 shows that the trend of stability constants for the various complexes of L1 with the first-row transition divalent metal ions and also of Cd2+ and Pb2+ is similar to that of other N-acetate derivatives of 14-membered macrocycles such as teta 2a,b or H2L18 (1-oxa- 4,8,12-triazacyclotetradecane-4,12-diacetic acid),3 all of them presenting pronounced differences in stability on progressing along the series of metal ions. This behaviour is in complete contrast with that of the 12-membered complexes such as those of 3,6,9,15-tetraazabicyclo[9.3.1]pentadeca-1(15),11,13-triene- 3,6,9-triacetic acid),11 dota,2a,b or H2L18 (1-oxa-4,7,10-triazacyclododecane- 4,10-diacetic acid),4 which are unable to differ- 70 J.Chem. Soc. Dalton Trans. 1997 Pages 65–73 entiate between metal ions exhibiting almost the same value of the stability constant for the entire series although higher than those of the corresponding 14-membered complexes. In spite of the intermediate values of stability constants exhibited by L1 between those of teta and L17 (cf. Fig. 4) apart from the copper complex which shows exceptional high stability L1 is the most selective of those macrocycles shown in Table 1 and in Fig. 4 even of all N-carboxymethyl macrocyclic compounds studied up to now.33 (2) In general complexes formed by L1 have lower stability constants than those of the corresponding teta complexes this being expected as teta has two more acetate groups for co-ordination.Nevertheless it was shown that all donor Fig. 2 The X-band EPR spectra of the iron complex of L1 in 1.0 mol dm23 NaClO4 at different temperatures (a) (4.6) (b) (12.0) and (c) (25.0 K) were recorded at n = 9.64 GHz (d) (121.9) and (e) (200.0 K) at n = 9.61 GHz and are expanded four times. Microwave power 2.4 mW modulation amplitude 1 mT. Resonances 1–3 S = ��� species; 4–7 one high-spin species (E/D ª 0.275); and 8 and 9 another high-spin species (E/D ª 0.15); x is an impurity of the instrument cavity. g values for lowspin species 2.683(1) 2.337(2) and 1.721(3); for both high-spin species; 9.50(4) 4.61(5) 3.93(6) 4.20(7) 8.68(8) and 5.39(9) Fig.3 Reciprocal of the paramagnetic susceptibility 1/cp (d) and magnetic moment meff (e) of [Fe(L1)Br]?4H2O in the temperature range 3–292 K atomskalineearth or lanthanide ions i.e. in complexes which have mainly electrostatic interactions and in these cases the macrocycle is in a very strained conformation.2a,b,34 However teta complexes of the first-row transition-metal ions have stabilities lower than expected a fact only explained if some of the donor atoms do not co-ordinate for structural reasons.2a,b Spectroscopic data in solution and X-ray diffraction analysis of crystals have shown that the ML complexes (M = Cu2+ or Ni2+) have two free acetate groups linked to opposite nitrogen atoms,35 and these groups may be involved in the formation of bi- or poly-nuclear species 36 making teta less interesting for applications involving the first-row transition metals.Compound L1 however has only the exact number of donor atoms to form octahedral complexes the polynuclear species being less probable as confirmed by the present thermodynamic studies. The lower stability of L1 complexes compared with those of teta is mainly due to the lower overall basicity of L1 (a difference of 3.83 log units 2a,b) and in general the stabilities of corresponding complexes of both macrocycles differ by the same amount. Exceptions to this statement would only be explained by a special structural arrangement of the ligand on complexation the complexes of Mg2+ and Cu2+ with L1 are more stable than those of teta (3.33 Fig.4 Variation of the stability constants (log KML) of the metal complexes of L1 (d) L2 (m),8 dota (h),2b,20 teta (s),2b,20 L17 (n)4 and L18 (j)3 with the atomic number of the metal ion Table 6 Rhombic distortions (E/D) and effective g values expected from the high-spin species of the iron(II) complexes of L1. (The values observed in the experimental spectra are underlined) E/D Doublet g1 g2 g3 0.15 |±5/2Ò |±3/2Ò |±1/2Ò 9.95 5.39 1.44 0.14 2.82 8.68 0.16 3.13 2.71 0.275 |±5/2Ò |±3/2Ò |±1/2Ò 9.79 4.61 0.82 0.43 3.93 9.50 0.57 4.2 1.24 J. Chem. Soc. Dalton Trans. 1997 Pages 65–73 71 and 1.12 log units in inverse position of the overall basicity) 2a,b and the opposite happens with Ca2+ In3+ and especially Fe3+ the complexes of which are destabilized by L1 when compared with teta (by 4.58 4.06 and 6.82 log units respectively).2a,20 (3) When the stabilities of the metal complexes of L1 are compared with those of the parent (L2 or L3) 8 the values for L1 are always higher which is an indication of the co-ordination of at least one acetate group.Also the cobalt(II) complexes of the parent amines are not stable.8 (4) The low stabilities of the complexes of trivalent metal ions with L1 (Ga3+ Fe3+ and In3+) are very intriguing. The value for the iron complex is even lower than that of the copper complex which is very unusual for a polyaminopolycarboxylic ligand. The extraordinary selectivity for the copper complex renders L1 an interesting ligand for medical applications in nuclear medicine using 64Cu (a b+ emitter of potential use in positron emission tomography),37 in radioimmunotherapy using 67Cu38 or for the treatment of copper intoxication in cases of metal poisoning in patients suffering from Wilson’s disease.39 The problem here involves removing the Cu2+ without perturbation of the other essential ions present in biological systems namely Ca2+ Zn2+ or Fe3+.The compounds trien (triethylenetetramine) 3,7-diazanonane-1,9-diamine (L19) are the most used ligands in chelation therapy the last forming a very stable complex with Cu2+ (23.2 in log units 12a). The complex [CuL1] has a lower stability constant than that of the nonanediamine however L1 has a lower overall basicity than that of linear tetraamines and at physiological pH the competition between protons and Cu2+ is less important for L1 as can be observed from the values of pM determined at pH 7.4 (cf.Table 7);40 also the linear tetraamines form charged complexes while [CuL1] is neutral and therefore most suited for diffusion into tissues.39 So L1 seems to be a useful ligand for this medical application better than L19 if no important toxic aspects occur. Structural data Some of the specificities of L1 upon complexation may be found in the configuration that the macrocycle seems to show a tendency to adopt when one or more substituents are linked to the nitrogen atoms. Some derivatives of L2 frequently exhibit a five-co-ordinated arrangement around the metal ion the macrocycle being folded in such a way that the nitrogen atom opposed to the pyridine binds at an axial position the basal plane being formed by the three other nitrogen atoms of the ring and another ligand or by a donor atom of one of the substituents at N (in a trans position to the pyridine N atom) in a square-pyramidal arrangement for the metal ion.Folding the macrocycle leads to a larger cavity and subsequently to longer M]N bond lengths and probably also to less stable complexes. Examples are [Ni(L4)X][ClO4]n (n = 1 X = Cl; n = 2 X = Me2SO);41 [Ni(L4)X]2+ (X = H2O or N3);42 [IrL2(Cl)H]PF6 in which the Ir3+ is octahedrally co-ordinated to all four nitrogen donors of L2 three of them equatorially and the fourth bent away to an axial position with the chloride and hydride ligands bound cis to each other;43 [Ni(L5)Cl]ClO4?H2O;44 Table 7 The pCua values for copper(II) complexes of L1 trien and L19 at pH 7.4 Ligand L1 trien b L19c pCu 17.79 15.97 18.54 a Values calculated for 100% excess of free ligand under physiological conditions pH 7.4; cCu = 1.0 × 1025 mol dm23 cL = 2.0 × 1025 mol dm23.b log K1 = 9.74 log K2 = 9.07 log K3 = 6.59 log K4 = 3.27 log KCuL = 20.05 log KM(HL) = 3.7 log KML(OH) = 10.7.12a c log K1 = 10.08 log K2 = 9.26 log K3 = 6.88 log K4 = 5.45 log KCuL = 23.2.12a [Co(L6)Cl][ClO4];45 [Ni(L8)(en)][ClO4]2 (en = ethane-1,2- diamine);46 [RuL7(Cl)(CO)][BPh4] shows octahedral Ru2+ coordinated to three nitrogen donors of the macrocycle bound equatorially the fourth N being bent away to bind at an axial position with the carbonyl and Cl completing the coordination. 47 Molecular mechanics theoretical studies also show this tendency of the macrocycle L4 to fold the nitrogen donor of the macrocycle opposed to the pyridine being in apical position.48 We could not obtain crystals of the complexes with appropriate size for X-ray diffraction analysis but some spectroscopic studies in solution were undertaken which also show that L1 seems to adopt the conformation described above at least in the complexes of Co2+ and Fe3+.The electronic spectrum of the cobalt complex exhibiting near-infrared and visible absorption together with a low value of the magnetic moment indicates a five-co-ordinate symmetry of a high-spin species.49,50 It is difficult to distinguish between spectra of high-spin squarepyramidal or trigonal-bipyramidal complexes but in general the former have lower intensities (often <100 dm3 mol21 cm21).51–53 The intense band at 342.5 nm is probably a chargetransfer band.The electronic spectrum of the nickel complex is characteristic of a tetragonal (D4h) symmetry. Following the conclusions of Busch and co-workers 54 on some tetraazamacrocycles we tentatively assigned the bands of our spectrum (3B1g æÆ 3Eg a 940; 3B1g æÆ 3B2g 810; 3B1g æÆ 3Eg b 570; and 3B1g æÆ 3Eg c 355 nm) and values of Dqxy and Dqz were calculated 1235 and 893 cm21; Dqz is strongly influenced by the in-plane ligand field decreasing as Dqxy increases. The complex [NiL1] has spectral parameters similar to those of the teta complex Dqxy = 1220 and Dqz = 740 cm21 or to those of [Ni([15]aneN4)(NCS)2] Dqxy = 1202 and Dqz = 908 cm21 ([15]aneN4 = 1,4,8,12- tetraazacyclopentadecane).54 There is no published structure for the nickel complex of teta but only one for copper Ba[Cu- (teta)]?6H2O.35 The geometry of the latter is a distorted octahedron with four amino nitrogens in a plane and two apical acetate oxygen donors.This geometry seems also be adopted by the nickel complex according to spectroscopic measurements.36 The low value obtained for Dqz is also an indication of this. Complexes of 14-membered tetraazamacrocycles (examples can be seen in ref. 54) have stronger equatorial fields. Two hypotheses can be advanced to explain our values if the equatorial plane is formed by the four nitrogen atoms of the ring the size of the cavity so formed should not be appropriate for Ni2+ and the macrocycle should fold for the co-ordination as happens with the [15]aneN4 complex or the structure adopted by the complex is such that it has a N3O set of donor atoms in the equatorial plane (three nitrogen atoms of the macrocycle and one oxygen from a carboxylate) the last nitrogen atom of the ring and the last carboxylate oxygen being in axial positions.The first hypothesis is less probable as L8 with the same set of donor atoms exactly in the same position behaves similarly to other 14-membered ligands.54 The second hypothesis is supported by the mentioned theoretical molecular mechanics study 48 and the crystal structures.41–47 This hypothesis also explains the high value of Dqz impossible to understand if the two carboxylate oxygens were in axial positions. It is well known that electronic spectra of copper(II) complexes are not especially good indicators of geometry.55 However some comparisons are possible with similar complexes such as those compiled in Table 4.The EPR spectrum of [CuL1] shows g3 > (g1 + g2)/2 which is typical of rhombic symmetry for the copper(II) ion where the distortion takes the form of elongation of the axial bonds and where a dx22y2 ground state is present and would be consistent with elongated rhombicoctahedral rhombic square-coplanar or distorted squarebased- pyramidal stereochemistries. It appears that it is possible to exclude a trigonal-bipyramidal geometry or a tetragonal structure involving compression of axial bonds.21,55 72 J. Chem. Soc. Dalton Trans. 1997 Pages 65–73 The electronic properties of the copper(II) complexes can be explained by the usual factors taken from the equations of the EPR parameters derived from ligand-field theory.56–58 The addition of a fifth ligand to a square-planar arrangement has the effect of decreasing Az while increasing gz with a simultaneous red shift in the electronic spectra.59,60 Comparing the data of Table 4 for the copper complex of the parent L3 with those of L1 we can see that there is an increase of gz a decrease of Az and a red shift of the electronic spectra.As the structure of the copper parent complex was considered square planar it is possible to infer that the structure of [CuL1] will be square pyramidal or even tetragonal. The ratio (g|| 2 2) (g^ 2 2) is nearly equal to 4 1 which is characteristic for the Cu2+ ion in a square-based pyramidal co-ordination.61,62 Also the electronic data for [CuL1] are very similar to those obtained for other complexes considered square pyramidal such as those of L14 or L13.22 Both have a geometry around the copper which can be described as a distorted square pyramid the first one with an equatorial plane formed by three nitrogen atoms of the ring and a halogen atom the oxygen of the macrocycle being in the apical position while in the second the four nitrogen atoms of the macrocycle form the basal plane and the halogen atom completes the coordination in the apical position.The spectral data for our complex are also very similar to those of the [Cu(teta)]22 complex 23 which is a distorted octahedron as already mentioned. The crystal structure of the copper complex of the diprotonated form of L1 was obtained in very acidic media.1 In this complex [Cu(H2L1)Cl]Cl the copper(II) has a distorted trigonal-bipyramidal geometry with apical co-ordination to Cl2 and the pyridine nitrogen of the ring and equatorial coordination to the other three nitrogen atoms of the macrocycle and only one of the two acetate groups the MeN group and one of the two acetate arms of the ligand being protonated and unco-ordinated.However the complex of the completely deprotonated species for which there is no crystal structure is thermodynamically very stable as verified in the present work the stability constant being of the same order of magnitude as that of the dota complex. It could be predicted for the former complex that all the donor atoms will be involved in coordination forming a tetragonal complex or as the spectroscopic data seem to indicate a five-co-ordinated complex probably with an acetate arm not involved in the co-ordination (non-co-ordination of a nitrogen of the ring would lead to a less stable complex).Note that the copper ion in a five-co-ordinated geometry (square pyramidal or trigonal bipyramidal) is also stabilized by crystal-field stabilization energy.53 The complex of L1 with Fe3+ has a surprisingly low stability constant as do the other trivalent metal-ion complexes in this work. The electronic spectrum as expected for a d5 configuration gives no information because the tail of the intense charge-transfer absorptions overlaps the weak forbidden bands of the visible region producing the yellow colour. The band at 328 nm is assigned to a metal-to-ligand charge-transfer transition because of the relatively high absorption coefficient. The EPR spectra of this complex in the range of temperature studied show two types of signals one for a typical low-spin system in a rhombic field and a high-spin species (S = 5�2 ).The low-spin species has a large rhombic distortion |R/m| = 0.515 (the maximum rhombicity value is equal to 0.667),28a although a value which is intermediate between that of the 1,4,8,11-tetraazacyclotetradecane (cyclam L11) complex and that of the cyclam derivative with only one acetate arm L15 (ref. 26) cf. Table 5. The magnetic moment at room-temperature measured in solution by the Evans method is 3.40 mB. This agrees quite well with the value obtained at room temperature from static magnetic susceptibility measurements meff (292 K) = 3.58 mB. Two hypotheses can be advanced for the interpretation of this value it may be due to three unpaired electrons with no orbital contribution (S = ��2 spin state) or it results from a weighted average of isomers with different spin states (S = 5�2 and ��� ).In the latter hypothesis the magnetic susceptibility could be described as the sum of the contributions of the low-spin (ls) and high-spin (hs) species c = xcls + (1 2 x)chs where x is the mole fraction of the low-spin isomer. The calculation using the above equation yields values in very large excess of those found by the EPR experiment in solution (for instance ª98% by calculation against 35% from the EPR experiment at 9 K). The first hypothesis implies that in the solid state the ligand geometry around the FeIII is slightly different from that in solution giving rise to an energy splitting that allows the S = ��2 spin configuration.The decrease in the magnetic susceptibility with temperature may then be due to a spin transition or to antiferromagnetic interactions. Several examples are known of FeIII in the S = ��2 spin state mainly five-co-ordinated in a distorted square-pyramidal geometry which exhibit a S = ��2 �Æ S = ��� spin crossover such as [FeL(NO)] with L = N,N9- bis(salicylidene)ethane-1,2-diaminate,63,64 1,4,8,11-tetramethyl- 1,4,8,11-tetraazacyclotetradecane,65 or N,N9-bis(salicylidene)- o-phenylenediaminate.66 The comparison with [Fe(L1)Br]?4H2O reveals some common features but also some differences. As with [Fe(L1)Br]?4H2O at least at room temperature all the complexes mentioned above have five-co-ordinated iron and the ligand is somewhat bent. These complexes show an intermediate spin �Æ low spin equilibrium over a narrow range of temperature.It is known however that spin transitions may extend over a wide range of temperature or even be incomplete.67 It is also worth mentioning that theoretical studies although performed for the high spin �Æ intermediate spin transitions of trigonal five-co-ordinated iron(III) species show that the temperature dependence of the magnetic moment from the state of higher spin (at high temperature) to the lower (at low temperature) extends over a wide range.68 The temperature profile those found in these theoretical studies. To clarify whether the magnetic behaviour at low temperature results from a spin crossover or is due to the onset of antiferromagnetic interactions requires further studies particularly Mössbauer experiments.Conclusions In spite of forming less stable complexes than those of similar tetraazamacrocyclic ligands without pyridine in the ring such as dota or teta L1 is more selective for divalent metal ions its complexes of trivalent metal ions such as Ga3+ Fe3+ or In3+ being comparatively less stable. The copper(II) complex is very stable rendering L1 an interesting ligand for medical applications in nuclear medicine or for the treatment of patients suffering from Wilson’s disease. Besides the high selectivity for Cu2+ which some linear tetraamines also have L1 exhibits a relatively low overall basicity and forms a neutral complex which makes it the best candidate for these medical applications if no toxic effects are found. The special selectivity of L1 in complex formation is probably due to its tendency to exhibit a folded conformation the metal ion being in a fiveco- ordinated environment.The observation of a low-spin iron(III) complex with g values such as those found in this work is of particular relevance to the interpretation of EPR spectra of iron(III) porphyrins in biological systems.69–71 So far only a limited number of haem iron axial co-ordinations has been unambiguously found [histidine (N) methionine (S) cysteine (S2) and amine (N)]; species with gmax values of ª2.6 have been generally assigned to hydroxylbound species.69 The assignment of haem iron co-ordination solely on the basis of EPR g values has recently become less and less reliable in part due to the lack of suitable model compounds. 70,71 The present study shows that carboxylate ligands can yield the same kind of EPR spectra opening a large range of new possible assignments of haem iron axial coordination.J. Chem. Soc. 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