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Thermodynamic studies on the complexation of cobalt(II) with nitrogen donor ligands in dimethyl sulfoxide
www.elsevier.nl/locate/ica Inorganica Chimica Acta 321 (2001) 49 – 55
Thermodynamic studies on the complexation of cobalt(II) with nitrogen donor ligands in dimethyl sulfoxide Clara Comuzzi, Michela Grespan, Pierluigi Polese, Roberto Portanova, Marilena Tolazzi * Dipartimento di Scienze e Tecnologie Chimiche, Uni6ersita` di Udine, Via Cotonificio 108, I-33100 Udine, Italy Received 14 February 2001; accepted 25 May 2001
Abstract Co(II) complex formation with the N-donor ligands n-butylamine (n-but), diethylenetriamine (dien), N,N¦-dimethyldiethylenetriamine (dmdien) and N,N,N%,N¦,N¦-pentamethyldiethylenetriamine (pmdien) has been studied at 298 K in the aprotic solvent dimethyl sulfoxide (dmso) in an ionic medium set to 0.1 mol dm − 3 with NEt4ClO4 in anaerobic conditions. Potentiometric, spectrophotometric and calorimetric measurements have been carried out to obtain the thermodynamic parameters for the systems investigated. Only mononuclear CoLj2 + complexes are formed ( j = 1–3 for n-but; j= 1, 2 for dien and dmdien; j = 1 for pmdien) where the triamines act prevalently as terdentate agents. All the complexes are enthalpy stabilized whereas the entropy changes counteract the complex formation. The results are discussed in term of different basicities and steric requirements of both the ligands and the complexes formed. Preliminary information on the different affinities of the complexes for dioxygen is reported. © 2001 Elsevier Science B.V. All rights reserved. Keywords: Cobalt(II); Amines; Complexes; Thermodynamics; Dimethylsulfoxide
1. Introduction Complexes of a number of transition metals able to bind, transport and activate small molecules such as dioxygen, carbon monoxide, carbon dioxide, nitric oxide and sulfur dioxide, both natural (metallo-proteins and -enzymes) and synthetic, are very important in chemical, biochemical, biological, environmental and industrial fields [1– 15]. The activation of these small molecules usually involves a direct or indirect interaction with the transition metal centre, normally of the first row, situated in a very specific coordination environment [14]. In this context it is well known that Co(II) complexes both in the solid state and in solution, are able to bind dioxygen and nitric oxide and they find large applications in the catalytic oxidation of organic substrates in mild conditions and in several industrial processes * Corresponding author. Tel.: + 39-0432-558852; fax: + 39-0432558803. E-mail address: [email protected] (M. Tolazzi).
[7,10,15]. In addition Co(II) complexes are very important in biology because they represent one of the most successful classes of synthetic oxygen carriers, characterized by a broad range of dioxygen affinities [4,5,7,11]. Moreover, Co(II) and other metal complexes play an important role in the industrial cleaning of gaseous effluents from power plants through the suggested simultaneous absorption of nitric oxide and sulfur dioxide [8,14,15]. So far, great interest has been paid to complexation studies of these metal ions in aqueous solution [4,5,7,11] and minor attention has been devoted to the behaviour of such complexes in non aqueous solvents where they can be used extensively as oxidation catalysts [7,8,10]. As part of a research program on the study of molecular systems for oxygen transport and activation, our group started a study on the thermodynamics of complex formation of Co(II) with linear and cyclic polyaza ligands and of their oxygenation in non aqueous solutions with the aim to design complexes suitable for practical applications and to seek for the best solvents to obtain a good reversibility of the
0020-1693/01/$ - see front matter © 2001 Elsevier Science B.V. All rights reserved. PII: S 0 0 2 0 - 1 6 9 3 ( 0 1 ) 0 0 5 1 3 - 8
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uptake and release processes of dioxygen to be used particularly for catalytic applications. The first step of this work requires the determination of the coordination properties of Co(II) with these ligands in the solvents concerned in the absence of oxygen. In the present paper we report a study of the Co(II) complex formation with the following N-donor ligands: n-butylamine (n-but), diethylenetriamine (dien), N,N¦dimethyldiethylenetriamine (dmdien) and N,N,N%, N¦,N¦-pentamethyldiethylenetriamine (pmdien) in the aprotic solvent dimethyl sulfoxide (dmso) in anaerobic conditions. These studies hold also an academic interest for the understanding of the coordination properties of cobalt(II) and of the factors governing its complexation ability and selectivity in non aqueous solutions. In this context the triamines investigated offer the opportunity to investigate both the influence of different basicities and steric hindrances in the Co(II) complex formation, influence which has been shown to be important in many transition metals amino complexes [16]. In addition, as the dioxygen affinity of the cobalt complexes which have a quite similar structure is dependent on the electron donor ability of ligand [17] (the stronger the electron donor ability of the ligand, the greater the electron density on cobalt center, the easier the electron drift from cobalt to dioxygen and the higher the dioxygen affinity), the triamines here investigated have been chosen as they possess the same skeleton and differ mainly for the basicity of the amino groups in dmso, decreasing in the order – NH2 \ – NHR \ – NR2 [18].
The values of the stability constants of the various systems investigated were determined using different techniques. In particular for the Co(II)-pmdien system the determination of the stability constants was made via potentiometric measurements, using silver(I) as auxiliary metal ion. Silver(I) could not be used to obtain the stability constants for the other Co(II) amine systems because of redox problems: in fact voltammetric measurements recorded at a Pt working electrode on dmso solutions containing Co(II) and the other ligands in the molar ratio 1:1 show that the reduction potential of the Co(III),Co(II) couple was about 0.1 V versus a reference Ag/AgCl, Cl− (sat.) electrode. Therefore the stability constants for these systems were determined: for Co(II)–n-but system by direct spectrophotometric titrations; for Co(II)– dmdien and –dien complexes direct spectrophotometric titrations were not successful [19] in determining these equilibrium constants, due to the high stabilities of the complexes, as indicated by calorimetric results (see in Section 2). Therefore, Cd(II) was chosen as competitive ion and its stability constants with the same ligands were determined previously by potentiometric measurements using the ion selective cadmium electrode. The enthalpy values were obtained by direct calorimetric titrations. All measurements were performed at 298 K and in an ionic medium adjusted to 0.1 mol dm − 3 with NEt4ClO4 as neutral salt taking extreme care to obtain and maintain the lowest water and oxygen content in the systems.
Table 1 Overall stability constants and thermodynamic functions for the reaction Co2++jLlCoLj2+ at 298 K and I =0.1 mol dm−3 in dmso and in water [20]. The errors quoted correspond to three standard deviations
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Fig. 1. Typical electronic spectra of cobalt(II) –dien system in dmso solution. The numbers represent j within the [Co(dien)j ]2 + .
2. Results The best fit of the potentiometric and spectrophotometric data was obtained when the species reported in Table 1 were taken into account. In Table 1, the overall stability constants and the free energies of formation, with the limits of error indicated, are listed for the reactions: Co2 + +jL X CoLj2 + (where L is the amine concerned and j= 1 – 3). No polynuclear or mixed Co– L – Ag or Co – L –Cd complexes have been evidenced in the range of concentration investigated. As an example in Fig. 1 typical UV– Vis spectra of Co –dien system, as given by the Hyperquad output using the stability constants in Table 1, is reported in term of molar absorbances vs. wavelengths. The absorption maximum for Co(ClO4)2 in dmso occurs at 535 nm with an extinction molar coefficients m = 11.9 mol − 1 dm3 cm − 1. The absorption maximum of CoL and CoL2 are shifted to lower wavelengths, in line with an increase in the ligand field strength. The experimental data obtained from the calorimetric measurements for all the systems investigated are reported in Fig. 2 as Dhv, the total heats of reaction per mole of metal ion, as a function of R = CL/CM, the ratio between the moles of ligand and the moles of cobalt(II) in the calorimetric vessel. The enthalpy curves are fully consistent with the information obtained from the analysis of potentiometric and spectrophotometric data. In fact, the curves representing the complexation of Co(II) with n-but indicate formation of complex(es) of low stability(es) whereas the Co(II)–pmdien data are consistent with the formation of only one complex of relatively high stability. For
both Co(II)–dmdien and – dien systems the superimposing of the curves up to R= 2 indicates formation of two complexes of high stabilities. In the case of dmdien the small change of slope at R= 1 indicates a slightly lower stability for the second complex, while for Co(II)–dien system the absence of any change of slope clearly indicates that two complexes of comparable stabilities are formed. The experimental heats of reaction and the overall stability constants in Table 1 were used to calculate the full lines in Fig. 2. The fit of the experimental data is quite good indicating that all the
Fig. 2. The total molar enthalpy changes, Dhv, as a function of R =CL/CM for cobalt(II) – amine systems in dmso. (a) n-but: a= 0, (2) 10.20, (") 30.42 mmol dm − 3 in Co2 + ; (b) dien: a = −50, () 9.98, () 29.96 mmol dm − 3 in Co2 + ; (c) dmdien: a = − 50, () 10.12, ( ) 30.36 mmol dm − 3 in Co2 + ; (d) pmdien: a = −50, ( ) 10.20, ( ) 30.42 mmol dm − 3 in Co2 + . The solid lines have been calculated from the values of ij and DH°ij in Table 1.
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systems are satisfactorily described. In Table 1 are also reported, for comparison, the available stability constants and thermodynamic functions for the complexation reactions of Co(II) by the same ligands in aqueous solutions [20].
3. Discussion The negative enthalpy values found here (Table 1) are typical of reactions involving complexation of metal ions by neutral ligands in aprotic solvents [21]. The entropy terms, also negative, oppose the complex formation. The release of solvent molecules from the coordination sphere of the metal cation accompanying complexation does not compensate the decrease in internal entropy of the ligand and the negative entropy change due to the loss of translational entropy of the reagents when the complexes are formed. The electronic spectrum of [Co(dmso)6]2 + (Fig. 1) is very similar to that arising from the octahedral [Co(H2O)6]2 + complex in water [22a]. The CoL and CoL2 spectra for all the systems concerned, similar in trends to those reported in Fig. 1 for the Co(II)–dien system, do not exhibit any characteristics of the tetrahedral coordination, i.e. much higher extinction coefficients and absorptions at higher wavelengths [22]: therefore it seems reasonable to suppose that the octahedral structure is retained after the coordination of the amines. In the case of Co(II)– n-but complex formation a feature of the system is that the stepwise −DH°2 \ − DH°1, due to a large desolvation occurring in the first step of complexation. Evidently the enthalpy needed for the desolvation is so large that DH°1 becomes less exothermic in spite of the fact that the heat evolved in the Co(II) –amine bond formation is certainly larger in the first step. This hypothesis is confirmed by the more unfavorable entropy term associated with the second complexation step. The thermodynamic values associated to the third complexation step reflect the lowering of the strength of Co –N bonds as more ligands are added [21a]. Cobalt(II) has been classified as a borderline acid in water [23] and its hard character is expected to be enhanced in a more basic solvent like dmso [23b]: therefore, the affinity for the hard N-atom is expected to be quite high. Nevertheless, and despite the Co(II) maximum coordination number is six, no more complexes beyond the third one were detected in our experimental conditions with n-but ligand. The reason for this is most likely to be found in the large activity of the solvent which evidently prevents the monodentate n-but to compete for the complex formation when three ligands are already coordinated and the net charge of the central ion is lowered. It is not to be excluded that
also steric hindrances may be responsable for this evidence. The stability constants associated to the first complexation step for the Co(II)– dien and –dmdien systems, indeed very high if compared with the Co(II)–n-but system, indicate that all the N atoms coordinate the central ion, forming two fused five-membered chelate rings probably disposed on an octahedral face, as found in water for the same system [20]. The very exothermic value of the DH°i 1 and the unfavorable entropy terms are in line with the hypothesis of terdentation, further confirmed by: (i) a comparison with the thermodynamic data available in water for the Co(II)– dien system (see Table 1) where the ligand behaves as tridentate [20]; (ii) FT-IR spectra run on solutions containing Co(II)– dien in the molar ratio 1:0.5. In fact while the asymmetrical and symmetrical NH stretching modes of free dien in dmso is observed at 3360 and 3300 cm − 1 respectively, [24,25] in the presence of the Co(II) ion a bathocromic shift of these NH stretching modes is observed towards 3268 and 3180 cm − 1, respectively. No bands typical of free NH groups were detected. For Co(II)–dmdien in the same molar ratio, the single band of the secondary NH stretching [24] at 3311 cm − 1 disappears and only the NH stretching mode ascribed to the bonded amine is observed at 3240 cm − 1. The reason for the decrease in i1 and for the less exothermic DH°i 1 values going from dien to the N,N¦dimethylated dmdien might be a minor strenght in the M–N bonds due both to the decreased basicity of the N-atoms and to the steric crowding produced by N-alkyl groups which causes an elongation of M–N bonds [16]. The less unfavorable entropy terms going from dien to dmdien reflect the minor solvation of N-methylated dmdien and the decrease in the outer sphere solvation energy of the complexes. The latter effect is a consequence of the fact that N-alkylation inhibits the formation of hydrogen bonds of the type M–N–H–O with the solvent and increase the radii of the complexes [16]. The minor solvation also contributes to a decreased exothermicity. The same general effects are certainly responsible for the decreased stability of the Co(II)–pmdien complex. In this case the huge drop of stability and of exothermicity for the complex formation might exclude a terdentate behaviour of the ligand. As monodentation would result in lower i1 value, as indicated by the i1 value for the Co(II) complex formation with the primary n-but ligand, the thermodynamic values for the Co(II)–pmdien system may better fit with a bidentate behaviour of the ligand: the values of the enthalpy and entropy terms seem to confirm this hypothesis. The steric requirements of the bulky pmdien is likely the reason why the formation of a second complex is prevented. That the steric requirements of the ligands
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play an important role on the behaviour of the complexes formed is clearly demonstrated by preliminary experiments carried out in aerated solutions. While dioxygen clearly binds to Co(II)– dien and – dmdien complexes, even if different abilities seem to be shown by the two systems, the visible spectra of deareated solutions of the Co(II)– pmdien complex are identical to the aereated ones in the range 650– 300 nm showing that the cobalt(II) complexes with polyamine ligands do not bind dioxygen and are not oxidized when all the nitrogen donor groups are tertiary, which was already found in water with other linear tertiary polyamines [26]. This finding agrees with previous reports that N-alkylation of amine ligands shifts the redox potentials of the couples Mn + 1/nLm anodically relative to those of the corresponding non-alkylated [16]. The low value of the ratio K1/K2 as well as the similarity of the thermodynamic functions for the 1:1 and 1:2 complexes are consistent with the second ligand in [Co(dien)2]2 + being terdentate also. FT-IR spectra run on Co:dien solutions in the molar ratio 1:2 confirm this hypothesis as only the NH stretching modes ascribing to bonded amine are observed at 3267 and 3176 cm − 1, respectively. No bands typical of free NH groups were detected. It is to note that for the Co(II)– dien system the K2 value is of the same order of magnitude of K1 differently from what occurs in water where K2 K1 [20]. This feature may be explained admitting that desolvation is primarily important when the first complexation step is considered in dmso: the higher exothermic enthalpy value and the higher unfavourable entropy term associated with the second complexation step (Table 1) confirm this hypothesis. As for the [Co(dmdien)2]2 + complex formation, the thermodynamic functions seem to indicate that the second ligand coordinated as terdentate which is confirmed by the FT-IR spectra of a 1:2 Co(II)– dmdien solution where only the band at 3240 cm − 1 appears: the steric requirements of the methyl groups should be mainly reasonably responsible for the drop in K2 and for the less favorable enthalpy value associated with the second step of complexation [16]. Higher stabilities are shown by the Co(II)– dien complexes in dmso than in water. Unfortunately no data concerning the heat of transfer of this ion from water to any other solvent are known at the best of our knowledge [27]. Nevertheless considering the similarity in the charge density of Zn2 + and Co2 + ions and the fact that Zn2 + and other divalent cations are more favorably solvated in dmso than in water, as reported by Kalidas et al. [27b], it is reasonable to expect that also Co(II) ion would be stronger solvated by dmso than by water. This effect alone should bring to a minor stability of the Co(II)– dien complexes in the former solvent, which
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is not the case. As no remarkable differencies between the structure of the starting salts and/or of the complexes are evident in the two solvent concerned [20,22], the results obtained here are probably connected with a larger solvation of the amine in water, via strong hydrogen bonds, which should compensate for the stronger solvation of Co(II) in dmso: this seems to be confirmed by more unfavorable entropy change in dmso. In conclusion, it is shown that little differencies in basicities and steric hindrances of the ligands influence strongly Co(II) complex formation not only in terms of stability of the species formed but also in term of the stoichiometry of the compounds present in solution. Moreover, it seems that different methylation of the triamines concerned causes quite important effects on dioxygen affinity of Co(II) complexes and this will be object of further investigations.
4. Experimental
4.1. Chemicals Co(ClO4)2·6dmso and Cd(ClO4)2·6dmso were prepared by precipitation from almost saturated solutions of the hydrated salts (Aldrich 98%) in dimethyl sulfoxide with diethylether. After filtration they were dried under vacuum for several days at 50 °C. The compounds were characterized as M(ClO4)2·6dmso (M= Co, Cd) by elemental analysis and by titration of the Co(II) or Cd(II) with EDTA [28]. The molar conductivities of 1.00 mmol dm − 3 solutions of Co(ClO4)2. and Cd(ClO4)2. (respectively, about 82 and 75 V − 1 mol − 1 cm2) were measured with a Methrohm 712 conductometer and were found in good agreement with the accepted value for 1:2 elecrolytes in dmso [29]. Anhydrous silver perchlorate was obtained from AgClO4·H2O (Fluka puriss) as described previously [30]. Dimethylsulfoxide (Fluka \ 99%) was purified by distillation according to the described procedures [30], degassed by a pumping–freezing procedure and stored over 4 A, molecular sieves. The ligands n-but, dien, pmdien (Aldrich \ 97%) were purified by fractional distillation [31]. The dmdien ligand was synthesized and purified as described earlier [32,33]. Its purity (\99%) was checked by 1H and 13C NMR techniques. Perchlorate stock solutions of Cd(II) and Co(II) ions were prepared by dissolving in anhydrous degassed dmso weighed amounts of the adducts and their concentrations were checked by titration with EDTA [28]. The solutions of AgClO4 were prepared and standardized as before [30]. The background salt NEt4ClO4 was recrystallized twice from methanol and dried at 110 °C. Solutions of the ligands were prepared by dissolving weighted
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amounts in dmso and standardized by thermometric titration with standard HClO4 solutions. All standard solutions were prepared and stored in a MB Braun 150 glove box under a controlled atmosphere containing less than 1 ppm of water and less than 1 ppm of oxygen. The water content in the solutions, typically 10} 20 ppm, was determined by a Metrohm 684 KF coulometer.
4.2. Potentiometric measurements All measurements were carried out under dry, oxygen free nitrogen in the MB Braun 150 glove box in a thermostated cell maintained at 298.09 0.1 K. The experimental data required for the determination of the stability constants of the Co(II)–pmdien system complexes were the equilibrium concentrations of the silver ion, which were obtained from the e.m.f. data of a galvanic cell similar to that reported previously [25]. The emfs were measured by means of an Amel 338 pH meter equipped with a Metrohm 6.0328.000 silver electrode as a working electrode and a Metrohm 6.0718.000 silver electrode as a reference. Aliquots of ligand solutions of known concentrations were first added to the cell containing solutions of silver(I) perchlorate only (2.00 BC°Ag B 10.00 mmol dm − 3) of exact known concentration and the free silver(I) concentration was measured after each titrant addition. Once the stability constants of the Ag(I) complexes with pmdien [34] had been determined, similar titrations were carried out on solutions containing, in addition to silver(I), also cobalt perchlorate in the concentration range 2.00B C°Co B 40.00 mmol dm − 3. In order to check that no formation of mixed complexes of the type Ag– L – Co occurred in solution, all titrations were performed with at least two different initial silver(I) concentrations, ranging from 2.00 to 10.00 mmol dm − 3. The electrode couple was periodically checked in the range the 10 − 6 B[Ag+]B 10 − 2 mol dm − 3 in the presence of Co(II) ion: the emf values varied with the silver ion concentration according to Nernst’s law. The stability constants of the Cd(II)– dmdien and –dien complexes were obtained using a Weiss WCD1001 Cd ion selective electrode as working electrode and a Methrom 6.0718.000 silver electrode as reference one. To Cd(II) solutions in the concentration range 2.00BC°Cd B40.00 mmol dm − 3, ligand solutions of known concentration were added and the free cadmium(II) concentration was measured after each titrant addition. Some titrations were carried out in duplicate to verify the reproducibility of the system. Equilibrium was reached tipically in 2– 5 min. The Nernstian response of the Cd electrode was obtained in the range 10 − 2 BCd B 10 − 6 mmol dm − 3. The computer program Hyperquad [35] was used for the calculation of the stability constants.
4.3. Spectrophotometric measurements UV –Vis spectra were recorded with a Varian Cary 2300 Spectrophotometer equipped with a thermostatted compartment and with a computer which recorded data in 5 nm intervals over the wavelength range 300–650 nm. A quartz cuvette with a pathlength of 1 cm (117.100 Bracco cell for anaerobic applications) was used which was filled and ermetically closed in glove box under rigorous absence of oxygen.
4.3.1. Co(II) –n-but system Co(ClO4)2 dmso solutions (2 ml) (20.00BC°Co B 40.00 mmoldm − 3) in cell were titrated directly with n-but solutions 100B C°L B 500 mmol dm − 3. 4.3.2. Co –dmdien and –dien systems The spectrophotometric determination of the stability constants was made by using Cd(II) as competive ion [36]. Different types of titrations were carried out in order to obtain the best competition between Cd(II) and Co(II) for the same ligand and/or to evidence possible formation of mixed or polynuclear species: (i) to 2 ml solutions in the cell containing 10.00BC°Co B 40.00 mmol dm − 3 and 10.00B C°L B 40.00 mmol dm − 3, Cd(ClO4)2 solutions of concentrations 25.00B C°Cd B 40.00 mmol dm − 3 were added; (ii) to 2 ml solutions in the cell containing 20.00B C°Cd B30.00 mmol dm − 3, 20.00B C°L B 30.00 mmol dm − 3 and 10.00B C°Co B 20.00 mmol dm − 3, a solution of Co(ClO4)2 about 50 mmol dm − 3 were added; (iii) to 2 ml solutions in cell 10.00B C°Co B 20.00 mmol dm − 3, 20.00B C°L B 40.00 mmol dm − 3, solutions of 50.00 mmol dm − 3 in C°L and 50.00 mmol dm − 3 in C°Cd were added. The fact that equilibrium was attained was supported by the observation that addition of titrant solution produce an immediate spectral change with no further change noted after several hours. The absorbance data at about 20 different wavelengths in the range 450–550 nm were analyzed and the formation constants of the CoLj complexes ( j=1, 2) were determined, together with the molar extinction coefficients of the complexes at each relevant wavelength, using the Hyperquad program [35]. 4.4. Calorimetric measurements A Tronac model 87–558 precision calorimeter was employed to measure the heats of reaction. The cover of the titration vessel and its connection to the calorimeter were modified in order to make a gasproof closure. To be sure of operating in absence of O2, both the vessel and the burette were filled inside the glove box, joined together, taken out of the glove box and connected to the calorimeter for the measurements. The
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calorimeter was checked by titration of tris(hydroxymethyl)aminomethane (tham) with a standard solution of HCl in water. The experimental value of the heat of neutralization of tham was found to be DH°= −47.48 kJ mol − 1, in good agreement with the accepted value of − 47.539 0.13 kJ mol − 1 [37]. The procedure of calorimetric titration was as described before [25]. The heats of dilution of the reactants, determined in separate runs, were found negligible. The least squares computer program Letagrop Kalle [38] was used for the calculation of the enthalpy changes.
4.5. FT-IR measurements The FT-IR spectra were obtained using a Bio-Rad FTS 40 spectrometer (maximum resolution 4 cm − 1; 16 scans). A cell with barium fluoride windows (thickness of 25 mm) was used. The cells were filled and closed in glove box and transferred quickly to the spectrometer. The C°Co in the dmso solutions was about 50 mmol dm − 3.
Acknowledgements This work has been supported by the Ministero dell’Universita` e della Ricerca Scientifica e Tecnologica (MURST, Rome) within the program COFIN 98.
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[15] K.J. Franz, S.J. Lippard, J. Am. Chem. Soc. 121 (1999) 10504. [16] D. Meyerstain, Coord. Chem. Rev. 185 – 186 (1999) 141. [17] H. Sugimoto, T. Nagayama, S. Maruyama, S. Fujinami, Y. Yasuda, M. Suzuki, A. Uehara, Bull. Chem. Soc. Jpn. 71 (1998) 2267. [18] (a) A. Mucci, R. Domain, R.L. Benoit, Can. J. Chem. 58 (1980) 953; (b) R.L. Benoit, M.J. Mackinnon, L. Bergeron, Can. J. Chem. 59 (1981) 1501. [19] F.J.C. Rossotti, H. Rossotti, The determination of Stability Constants, McGraw-Hill, New York, 1961. [20] M. Ciampolini, P. Paoletti, L. Sacconi, J. Chem. Soc. (1961) 2994. [21] (a) S. Ahrland, in: J.J. Lagowsky (Ed.), The Chemistry of Non-Aqueous Solvents, vol. 5a, Academic Press, New York, 1978; (b) F. Arnaud-Neu, Chem. Soc. Rev. (1994) 235; (c) A. Cassol, P. Di Bernardo, R. Portanova, M. Tolazzi, P.L. Zanonato, Inorg. Chim. Acta 262 (1997) 1 and references cited therein. [22] (a) F.A. Cotton, G. Wilkinson, Advanced Inorganic Chemistry, 4th ed., Wiley, New York, 1980, p. 794; (b) H. Suzuki, S.-I. Ishiguro, H. Hohtaki, J. Chem. Soc., Faraday Trans. 86 (1990) 2179. [23] (a) R.G. Pearson, J. Am. Chem. Soc. 85 (1963) 3533; (b) R.D. Hanckock, A.E. Martell, J. Chem. Ed. 73 (1996) 654. [24] R.M. Silverstein, G.C. Bassler, T.C. Morrill, Spectrometric Identification of Organic Compounds, 3rd ed., Wiley, New York, 1974. [25] A. Cassol, P. Di Bernardo, R. Portanova, M. Tolazzi, G. Tomat, P.L. Zanonato, J. Chem. Soc., Dalton Trans. (1992) 469. [26] I. Bertini, L. Messori, G. Golub, H. Cohen, D. Meyerstein, Inorg. Chim. Acta 235 (1995) 5. [27] (a) G. Gritzner, A. Lewandowski, J. Chem. Soc., Faraday Trans. 87 (1991) 2599; (b) C. Halidas, G. Heffer, Y. Marcus, Chem. Rev. 100 (2000) 819. [28] A.I. Vogel, in: Textbook of Quantitative Chemical Analysis, 5th ed., Longman Group, UK, 1989. [29] W.J. Geary, Coord. Chem. Rev. 7 (1971) 81. [30] A. Cassol, P. Di Bernardo, P.L. Zanonato, R. Portanova, M. Tolazzi, J. Chem. Soc., Dalton Trans. (1987) 657. [31] D.D. Perrin, W.L.F. Armarego, D.R. Perrin, Purification of Laboratory Chemicals, 2nd ed., Pergamon, Oxford, 1980. [32] R.M. Clay, S. Corr, M. Micheloni, P. Paoletti, Inorg. Chem. 24 (1985) 3330. [33] D.W. White, B.A. Karcher, R.A. Jacobson, J.K. Verkade, J. Am. Chem. Soc. (1979) 4921. [34] Ag+ +jL lAgLj + ( j=1, 2). Log i1 =5.44(0.01), log i2 = 6.97(0.06): the errors in parentheses correspond to three standard deviations. [35] P. Gans, A. Sabatini, A. Vacca, Talanta 43 (1996) 1739. [36] Cd2 + +jL lCdLj2 + ( j=1, 2). For L =dien: log i1 = 8.46(0.09), log i2 =17.15(0.03); for L =dmdien: log i1 = 7.51(0.08), log i2 =13.42(0.09): the errors in parentheses correspond to three standard deviations. [37] R.M. Smith, A.E. Martell, Critical Stability Constants, vol. 6, Plenum Press, New York, 1989 2nd Supplement. [38] R. Arnek, Ark. Kemi 32 (1970) 81.
Thermodynamic studies on the complexation of cobalt(II) with nitrogen donor ligands in dimethyl sulfoxide Clara Comuzzi, Michela Grespan, Pierluigi Polese, Roberto Portanova, Marilena Tolazzi * Dipartimento di Scienze e Tecnologie Chimiche, Uni6ersita` di Udine, Via Cotonificio 108, I-33100 Udine, Italy Received 14 February 2001; accepted 25 May 2001
Abstract Co(II) complex formation with the N-donor ligands n-butylamine (n-but), diethylenetriamine (dien), N,N¦-dimethyldiethylenetriamine (dmdien) and N,N,N%,N¦,N¦-pentamethyldiethylenetriamine (pmdien) has been studied at 298 K in the aprotic solvent dimethyl sulfoxide (dmso) in an ionic medium set to 0.1 mol dm − 3 with NEt4ClO4 in anaerobic conditions. Potentiometric, spectrophotometric and calorimetric measurements have been carried out to obtain the thermodynamic parameters for the systems investigated. Only mononuclear CoLj2 + complexes are formed ( j = 1–3 for n-but; j= 1, 2 for dien and dmdien; j = 1 for pmdien) where the triamines act prevalently as terdentate agents. All the complexes are enthalpy stabilized whereas the entropy changes counteract the complex formation. The results are discussed in term of different basicities and steric requirements of both the ligands and the complexes formed. Preliminary information on the different affinities of the complexes for dioxygen is reported. © 2001 Elsevier Science B.V. All rights reserved. Keywords: Cobalt(II); Amines; Complexes; Thermodynamics; Dimethylsulfoxide
1. Introduction Complexes of a number of transition metals able to bind, transport and activate small molecules such as dioxygen, carbon monoxide, carbon dioxide, nitric oxide and sulfur dioxide, both natural (metallo-proteins and -enzymes) and synthetic, are very important in chemical, biochemical, biological, environmental and industrial fields [1– 15]. The activation of these small molecules usually involves a direct or indirect interaction with the transition metal centre, normally of the first row, situated in a very specific coordination environment [14]. In this context it is well known that Co(II) complexes both in the solid state and in solution, are able to bind dioxygen and nitric oxide and they find large applications in the catalytic oxidation of organic substrates in mild conditions and in several industrial processes * Corresponding author. Tel.: + 39-0432-558852; fax: + 39-0432558803. E-mail address: [email protected] (M. Tolazzi).
[7,10,15]. In addition Co(II) complexes are very important in biology because they represent one of the most successful classes of synthetic oxygen carriers, characterized by a broad range of dioxygen affinities [4,5,7,11]. Moreover, Co(II) and other metal complexes play an important role in the industrial cleaning of gaseous effluents from power plants through the suggested simultaneous absorption of nitric oxide and sulfur dioxide [8,14,15]. So far, great interest has been paid to complexation studies of these metal ions in aqueous solution [4,5,7,11] and minor attention has been devoted to the behaviour of such complexes in non aqueous solvents where they can be used extensively as oxidation catalysts [7,8,10]. As part of a research program on the study of molecular systems for oxygen transport and activation, our group started a study on the thermodynamics of complex formation of Co(II) with linear and cyclic polyaza ligands and of their oxygenation in non aqueous solutions with the aim to design complexes suitable for practical applications and to seek for the best solvents to obtain a good reversibility of the
0020-1693/01/$ - see front matter © 2001 Elsevier Science B.V. All rights reserved. PII: S 0 0 2 0 - 1 6 9 3 ( 0 1 ) 0 0 5 1 3 - 8
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uptake and release processes of dioxygen to be used particularly for catalytic applications. The first step of this work requires the determination of the coordination properties of Co(II) with these ligands in the solvents concerned in the absence of oxygen. In the present paper we report a study of the Co(II) complex formation with the following N-donor ligands: n-butylamine (n-but), diethylenetriamine (dien), N,N¦dimethyldiethylenetriamine (dmdien) and N,N,N%, N¦,N¦-pentamethyldiethylenetriamine (pmdien) in the aprotic solvent dimethyl sulfoxide (dmso) in anaerobic conditions. These studies hold also an academic interest for the understanding of the coordination properties of cobalt(II) and of the factors governing its complexation ability and selectivity in non aqueous solutions. In this context the triamines investigated offer the opportunity to investigate both the influence of different basicities and steric hindrances in the Co(II) complex formation, influence which has been shown to be important in many transition metals amino complexes [16]. In addition, as the dioxygen affinity of the cobalt complexes which have a quite similar structure is dependent on the electron donor ability of ligand [17] (the stronger the electron donor ability of the ligand, the greater the electron density on cobalt center, the easier the electron drift from cobalt to dioxygen and the higher the dioxygen affinity), the triamines here investigated have been chosen as they possess the same skeleton and differ mainly for the basicity of the amino groups in dmso, decreasing in the order – NH2 \ – NHR \ – NR2 [18].
The values of the stability constants of the various systems investigated were determined using different techniques. In particular for the Co(II)-pmdien system the determination of the stability constants was made via potentiometric measurements, using silver(I) as auxiliary metal ion. Silver(I) could not be used to obtain the stability constants for the other Co(II) amine systems because of redox problems: in fact voltammetric measurements recorded at a Pt working electrode on dmso solutions containing Co(II) and the other ligands in the molar ratio 1:1 show that the reduction potential of the Co(III),Co(II) couple was about 0.1 V versus a reference Ag/AgCl, Cl− (sat.) electrode. Therefore the stability constants for these systems were determined: for Co(II)–n-but system by direct spectrophotometric titrations; for Co(II)– dmdien and –dien complexes direct spectrophotometric titrations were not successful [19] in determining these equilibrium constants, due to the high stabilities of the complexes, as indicated by calorimetric results (see in Section 2). Therefore, Cd(II) was chosen as competitive ion and its stability constants with the same ligands were determined previously by potentiometric measurements using the ion selective cadmium electrode. The enthalpy values were obtained by direct calorimetric titrations. All measurements were performed at 298 K and in an ionic medium adjusted to 0.1 mol dm − 3 with NEt4ClO4 as neutral salt taking extreme care to obtain and maintain the lowest water and oxygen content in the systems.
Table 1 Overall stability constants and thermodynamic functions for the reaction Co2++jLlCoLj2+ at 298 K and I =0.1 mol dm−3 in dmso and in water [20]. The errors quoted correspond to three standard deviations
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51
Fig. 1. Typical electronic spectra of cobalt(II) –dien system in dmso solution. The numbers represent j within the [Co(dien)j ]2 + .
2. Results The best fit of the potentiometric and spectrophotometric data was obtained when the species reported in Table 1 were taken into account. In Table 1, the overall stability constants and the free energies of formation, with the limits of error indicated, are listed for the reactions: Co2 + +jL X CoLj2 + (where L is the amine concerned and j= 1 – 3). No polynuclear or mixed Co– L – Ag or Co – L –Cd complexes have been evidenced in the range of concentration investigated. As an example in Fig. 1 typical UV– Vis spectra of Co –dien system, as given by the Hyperquad output using the stability constants in Table 1, is reported in term of molar absorbances vs. wavelengths. The absorption maximum for Co(ClO4)2 in dmso occurs at 535 nm with an extinction molar coefficients m = 11.9 mol − 1 dm3 cm − 1. The absorption maximum of CoL and CoL2 are shifted to lower wavelengths, in line with an increase in the ligand field strength. The experimental data obtained from the calorimetric measurements for all the systems investigated are reported in Fig. 2 as Dhv, the total heats of reaction per mole of metal ion, as a function of R = CL/CM, the ratio between the moles of ligand and the moles of cobalt(II) in the calorimetric vessel. The enthalpy curves are fully consistent with the information obtained from the analysis of potentiometric and spectrophotometric data. In fact, the curves representing the complexation of Co(II) with n-but indicate formation of complex(es) of low stability(es) whereas the Co(II)–pmdien data are consistent with the formation of only one complex of relatively high stability. For
both Co(II)–dmdien and – dien systems the superimposing of the curves up to R= 2 indicates formation of two complexes of high stabilities. In the case of dmdien the small change of slope at R= 1 indicates a slightly lower stability for the second complex, while for Co(II)–dien system the absence of any change of slope clearly indicates that two complexes of comparable stabilities are formed. The experimental heats of reaction and the overall stability constants in Table 1 were used to calculate the full lines in Fig. 2. The fit of the experimental data is quite good indicating that all the
Fig. 2. The total molar enthalpy changes, Dhv, as a function of R =CL/CM for cobalt(II) – amine systems in dmso. (a) n-but: a= 0, (2) 10.20, (") 30.42 mmol dm − 3 in Co2 + ; (b) dien: a = −50, () 9.98, () 29.96 mmol dm − 3 in Co2 + ; (c) dmdien: a = − 50, () 10.12, ( ) 30.36 mmol dm − 3 in Co2 + ; (d) pmdien: a = −50, ( ) 10.20, ( ) 30.42 mmol dm − 3 in Co2 + . The solid lines have been calculated from the values of ij and DH°ij in Table 1.
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systems are satisfactorily described. In Table 1 are also reported, for comparison, the available stability constants and thermodynamic functions for the complexation reactions of Co(II) by the same ligands in aqueous solutions [20].
3. Discussion The negative enthalpy values found here (Table 1) are typical of reactions involving complexation of metal ions by neutral ligands in aprotic solvents [21]. The entropy terms, also negative, oppose the complex formation. The release of solvent molecules from the coordination sphere of the metal cation accompanying complexation does not compensate the decrease in internal entropy of the ligand and the negative entropy change due to the loss of translational entropy of the reagents when the complexes are formed. The electronic spectrum of [Co(dmso)6]2 + (Fig. 1) is very similar to that arising from the octahedral [Co(H2O)6]2 + complex in water [22a]. The CoL and CoL2 spectra for all the systems concerned, similar in trends to those reported in Fig. 1 for the Co(II)–dien system, do not exhibit any characteristics of the tetrahedral coordination, i.e. much higher extinction coefficients and absorptions at higher wavelengths [22]: therefore it seems reasonable to suppose that the octahedral structure is retained after the coordination of the amines. In the case of Co(II)– n-but complex formation a feature of the system is that the stepwise −DH°2 \ − DH°1, due to a large desolvation occurring in the first step of complexation. Evidently the enthalpy needed for the desolvation is so large that DH°1 becomes less exothermic in spite of the fact that the heat evolved in the Co(II) –amine bond formation is certainly larger in the first step. This hypothesis is confirmed by the more unfavorable entropy term associated with the second complexation step. The thermodynamic values associated to the third complexation step reflect the lowering of the strength of Co –N bonds as more ligands are added [21a]. Cobalt(II) has been classified as a borderline acid in water [23] and its hard character is expected to be enhanced in a more basic solvent like dmso [23b]: therefore, the affinity for the hard N-atom is expected to be quite high. Nevertheless, and despite the Co(II) maximum coordination number is six, no more complexes beyond the third one were detected in our experimental conditions with n-but ligand. The reason for this is most likely to be found in the large activity of the solvent which evidently prevents the monodentate n-but to compete for the complex formation when three ligands are already coordinated and the net charge of the central ion is lowered. It is not to be excluded that
also steric hindrances may be responsable for this evidence. The stability constants associated to the first complexation step for the Co(II)– dien and –dmdien systems, indeed very high if compared with the Co(II)–n-but system, indicate that all the N atoms coordinate the central ion, forming two fused five-membered chelate rings probably disposed on an octahedral face, as found in water for the same system [20]. The very exothermic value of the DH°i 1 and the unfavorable entropy terms are in line with the hypothesis of terdentation, further confirmed by: (i) a comparison with the thermodynamic data available in water for the Co(II)– dien system (see Table 1) where the ligand behaves as tridentate [20]; (ii) FT-IR spectra run on solutions containing Co(II)– dien in the molar ratio 1:0.5. In fact while the asymmetrical and symmetrical NH stretching modes of free dien in dmso is observed at 3360 and 3300 cm − 1 respectively, [24,25] in the presence of the Co(II) ion a bathocromic shift of these NH stretching modes is observed towards 3268 and 3180 cm − 1, respectively. No bands typical of free NH groups were detected. For Co(II)–dmdien in the same molar ratio, the single band of the secondary NH stretching [24] at 3311 cm − 1 disappears and only the NH stretching mode ascribed to the bonded amine is observed at 3240 cm − 1. The reason for the decrease in i1 and for the less exothermic DH°i 1 values going from dien to the N,N¦dimethylated dmdien might be a minor strenght in the M–N bonds due both to the decreased basicity of the N-atoms and to the steric crowding produced by N-alkyl groups which causes an elongation of M–N bonds [16]. The less unfavorable entropy terms going from dien to dmdien reflect the minor solvation of N-methylated dmdien and the decrease in the outer sphere solvation energy of the complexes. The latter effect is a consequence of the fact that N-alkylation inhibits the formation of hydrogen bonds of the type M–N–H–O with the solvent and increase the radii of the complexes [16]. The minor solvation also contributes to a decreased exothermicity. The same general effects are certainly responsible for the decreased stability of the Co(II)–pmdien complex. In this case the huge drop of stability and of exothermicity for the complex formation might exclude a terdentate behaviour of the ligand. As monodentation would result in lower i1 value, as indicated by the i1 value for the Co(II) complex formation with the primary n-but ligand, the thermodynamic values for the Co(II)–pmdien system may better fit with a bidentate behaviour of the ligand: the values of the enthalpy and entropy terms seem to confirm this hypothesis. The steric requirements of the bulky pmdien is likely the reason why the formation of a second complex is prevented. That the steric requirements of the ligands
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play an important role on the behaviour of the complexes formed is clearly demonstrated by preliminary experiments carried out in aerated solutions. While dioxygen clearly binds to Co(II)– dien and – dmdien complexes, even if different abilities seem to be shown by the two systems, the visible spectra of deareated solutions of the Co(II)– pmdien complex are identical to the aereated ones in the range 650– 300 nm showing that the cobalt(II) complexes with polyamine ligands do not bind dioxygen and are not oxidized when all the nitrogen donor groups are tertiary, which was already found in water with other linear tertiary polyamines [26]. This finding agrees with previous reports that N-alkylation of amine ligands shifts the redox potentials of the couples Mn + 1/nLm anodically relative to those of the corresponding non-alkylated [16]. The low value of the ratio K1/K2 as well as the similarity of the thermodynamic functions for the 1:1 and 1:2 complexes are consistent with the second ligand in [Co(dien)2]2 + being terdentate also. FT-IR spectra run on Co:dien solutions in the molar ratio 1:2 confirm this hypothesis as only the NH stretching modes ascribing to bonded amine are observed at 3267 and 3176 cm − 1, respectively. No bands typical of free NH groups were detected. It is to note that for the Co(II)– dien system the K2 value is of the same order of magnitude of K1 differently from what occurs in water where K2 K1 [20]. This feature may be explained admitting that desolvation is primarily important when the first complexation step is considered in dmso: the higher exothermic enthalpy value and the higher unfavourable entropy term associated with the second complexation step (Table 1) confirm this hypothesis. As for the [Co(dmdien)2]2 + complex formation, the thermodynamic functions seem to indicate that the second ligand coordinated as terdentate which is confirmed by the FT-IR spectra of a 1:2 Co(II)– dmdien solution where only the band at 3240 cm − 1 appears: the steric requirements of the methyl groups should be mainly reasonably responsible for the drop in K2 and for the less favorable enthalpy value associated with the second step of complexation [16]. Higher stabilities are shown by the Co(II)– dien complexes in dmso than in water. Unfortunately no data concerning the heat of transfer of this ion from water to any other solvent are known at the best of our knowledge [27]. Nevertheless considering the similarity in the charge density of Zn2 + and Co2 + ions and the fact that Zn2 + and other divalent cations are more favorably solvated in dmso than in water, as reported by Kalidas et al. [27b], it is reasonable to expect that also Co(II) ion would be stronger solvated by dmso than by water. This effect alone should bring to a minor stability of the Co(II)– dien complexes in the former solvent, which
53
is not the case. As no remarkable differencies between the structure of the starting salts and/or of the complexes are evident in the two solvent concerned [20,22], the results obtained here are probably connected with a larger solvation of the amine in water, via strong hydrogen bonds, which should compensate for the stronger solvation of Co(II) in dmso: this seems to be confirmed by more unfavorable entropy change in dmso. In conclusion, it is shown that little differencies in basicities and steric hindrances of the ligands influence strongly Co(II) complex formation not only in terms of stability of the species formed but also in term of the stoichiometry of the compounds present in solution. Moreover, it seems that different methylation of the triamines concerned causes quite important effects on dioxygen affinity of Co(II) complexes and this will be object of further investigations.
4. Experimental
4.1. Chemicals Co(ClO4)2·6dmso and Cd(ClO4)2·6dmso were prepared by precipitation from almost saturated solutions of the hydrated salts (Aldrich 98%) in dimethyl sulfoxide with diethylether. After filtration they were dried under vacuum for several days at 50 °C. The compounds were characterized as M(ClO4)2·6dmso (M= Co, Cd) by elemental analysis and by titration of the Co(II) or Cd(II) with EDTA [28]. The molar conductivities of 1.00 mmol dm − 3 solutions of Co(ClO4)2. and Cd(ClO4)2. (respectively, about 82 and 75 V − 1 mol − 1 cm2) were measured with a Methrohm 712 conductometer and were found in good agreement with the accepted value for 1:2 elecrolytes in dmso [29]. Anhydrous silver perchlorate was obtained from AgClO4·H2O (Fluka puriss) as described previously [30]. Dimethylsulfoxide (Fluka \ 99%) was purified by distillation according to the described procedures [30], degassed by a pumping–freezing procedure and stored over 4 A, molecular sieves. The ligands n-but, dien, pmdien (Aldrich \ 97%) were purified by fractional distillation [31]. The dmdien ligand was synthesized and purified as described earlier [32,33]. Its purity (\99%) was checked by 1H and 13C NMR techniques. Perchlorate stock solutions of Cd(II) and Co(II) ions were prepared by dissolving in anhydrous degassed dmso weighed amounts of the adducts and their concentrations were checked by titration with EDTA [28]. The solutions of AgClO4 were prepared and standardized as before [30]. The background salt NEt4ClO4 was recrystallized twice from methanol and dried at 110 °C. Solutions of the ligands were prepared by dissolving weighted
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amounts in dmso and standardized by thermometric titration with standard HClO4 solutions. All standard solutions were prepared and stored in a MB Braun 150 glove box under a controlled atmosphere containing less than 1 ppm of water and less than 1 ppm of oxygen. The water content in the solutions, typically 10} 20 ppm, was determined by a Metrohm 684 KF coulometer.
4.2. Potentiometric measurements All measurements were carried out under dry, oxygen free nitrogen in the MB Braun 150 glove box in a thermostated cell maintained at 298.09 0.1 K. The experimental data required for the determination of the stability constants of the Co(II)–pmdien system complexes were the equilibrium concentrations of the silver ion, which were obtained from the e.m.f. data of a galvanic cell similar to that reported previously [25]. The emfs were measured by means of an Amel 338 pH meter equipped with a Metrohm 6.0328.000 silver electrode as a working electrode and a Metrohm 6.0718.000 silver electrode as a reference. Aliquots of ligand solutions of known concentrations were first added to the cell containing solutions of silver(I) perchlorate only (2.00 BC°Ag B 10.00 mmol dm − 3) of exact known concentration and the free silver(I) concentration was measured after each titrant addition. Once the stability constants of the Ag(I) complexes with pmdien [34] had been determined, similar titrations were carried out on solutions containing, in addition to silver(I), also cobalt perchlorate in the concentration range 2.00B C°Co B 40.00 mmol dm − 3. In order to check that no formation of mixed complexes of the type Ag– L – Co occurred in solution, all titrations were performed with at least two different initial silver(I) concentrations, ranging from 2.00 to 10.00 mmol dm − 3. The electrode couple was periodically checked in the range the 10 − 6 B[Ag+]B 10 − 2 mol dm − 3 in the presence of Co(II) ion: the emf values varied with the silver ion concentration according to Nernst’s law. The stability constants of the Cd(II)– dmdien and –dien complexes were obtained using a Weiss WCD1001 Cd ion selective electrode as working electrode and a Methrom 6.0718.000 silver electrode as reference one. To Cd(II) solutions in the concentration range 2.00BC°Cd B40.00 mmol dm − 3, ligand solutions of known concentration were added and the free cadmium(II) concentration was measured after each titrant addition. Some titrations were carried out in duplicate to verify the reproducibility of the system. Equilibrium was reached tipically in 2– 5 min. The Nernstian response of the Cd electrode was obtained in the range 10 − 2 BCd B 10 − 6 mmol dm − 3. The computer program Hyperquad [35] was used for the calculation of the stability constants.
4.3. Spectrophotometric measurements UV –Vis spectra were recorded with a Varian Cary 2300 Spectrophotometer equipped with a thermostatted compartment and with a computer which recorded data in 5 nm intervals over the wavelength range 300–650 nm. A quartz cuvette with a pathlength of 1 cm (117.100 Bracco cell for anaerobic applications) was used which was filled and ermetically closed in glove box under rigorous absence of oxygen.
4.3.1. Co(II) –n-but system Co(ClO4)2 dmso solutions (2 ml) (20.00BC°Co B 40.00 mmoldm − 3) in cell were titrated directly with n-but solutions 100B C°L B 500 mmol dm − 3. 4.3.2. Co –dmdien and –dien systems The spectrophotometric determination of the stability constants was made by using Cd(II) as competive ion [36]. Different types of titrations were carried out in order to obtain the best competition between Cd(II) and Co(II) for the same ligand and/or to evidence possible formation of mixed or polynuclear species: (i) to 2 ml solutions in the cell containing 10.00BC°Co B 40.00 mmol dm − 3 and 10.00B C°L B 40.00 mmol dm − 3, Cd(ClO4)2 solutions of concentrations 25.00B C°Cd B 40.00 mmol dm − 3 were added; (ii) to 2 ml solutions in the cell containing 20.00B C°Cd B30.00 mmol dm − 3, 20.00B C°L B 30.00 mmol dm − 3 and 10.00B C°Co B 20.00 mmol dm − 3, a solution of Co(ClO4)2 about 50 mmol dm − 3 were added; (iii) to 2 ml solutions in cell 10.00B C°Co B 20.00 mmol dm − 3, 20.00B C°L B 40.00 mmol dm − 3, solutions of 50.00 mmol dm − 3 in C°L and 50.00 mmol dm − 3 in C°Cd were added. The fact that equilibrium was attained was supported by the observation that addition of titrant solution produce an immediate spectral change with no further change noted after several hours. The absorbance data at about 20 different wavelengths in the range 450–550 nm were analyzed and the formation constants of the CoLj complexes ( j=1, 2) were determined, together with the molar extinction coefficients of the complexes at each relevant wavelength, using the Hyperquad program [35]. 4.4. Calorimetric measurements A Tronac model 87–558 precision calorimeter was employed to measure the heats of reaction. The cover of the titration vessel and its connection to the calorimeter were modified in order to make a gasproof closure. To be sure of operating in absence of O2, both the vessel and the burette were filled inside the glove box, joined together, taken out of the glove box and connected to the calorimeter for the measurements. The
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calorimeter was checked by titration of tris(hydroxymethyl)aminomethane (tham) with a standard solution of HCl in water. The experimental value of the heat of neutralization of tham was found to be DH°= −47.48 kJ mol − 1, in good agreement with the accepted value of − 47.539 0.13 kJ mol − 1 [37]. The procedure of calorimetric titration was as described before [25]. The heats of dilution of the reactants, determined in separate runs, were found negligible. The least squares computer program Letagrop Kalle [38] was used for the calculation of the enthalpy changes.
4.5. FT-IR measurements The FT-IR spectra were obtained using a Bio-Rad FTS 40 spectrometer (maximum resolution 4 cm − 1; 16 scans). A cell with barium fluoride windows (thickness of 25 mm) was used. The cells were filled and closed in glove box and transferred quickly to the spectrometer. The C°Co in the dmso solutions was about 50 mmol dm − 3.
Acknowledgements This work has been supported by the Ministero dell’Universita` e della Ricerca Scientifica e Tecnologica (MURST, Rome) within the program COFIN 98.
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