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The influence of N-alkylation and chain length in the coordination ability of diamines towards silver(I) in dimethylsulfoxide
Polyhedron 21 (2002) 1337 /1342 www.elsevier.com/locate/poly
The influence of N-alkylation and chain length in the coordination ability of diamines towards silver(I) in dimethylsulfoxide Clara Comuzzi *, Veronica Novelli, Roberto Portanova, Marilena Tolazzi * Dipartimento di Scienze e Tecnologie Chimiche, Universita` di Udine, Via del Cotonificio 108, I-33100 Udine, Italy Received 27 September 2001; accepted 16 October 2001
Abstract The thermodynamic functions for the complexation of Ag(I) by the following diamines: N ,N -dimethyldiethylenediamine (N ,N dmen), N ,N -dimethyl-1,3-propanediamine (N ,N -dmtn) and N ,N ,N ?,N ?-tetramethyl-1,3-propanediamine (tmtn) have been determined in dimethylsulfoxide (dmso) by potentiometric and calorimetric techniques at 298 K and 0.1 mol dm 3 ionic strength (NEt4ClO4). Only mononuclear complexes are formed (AgLj , j /1, 2) where the ligands act as monodentate or chelate agents. All the complexes are enthalpy stabilized whereas the entropy changes counteract the complexation. The different basicities and steric requirements of both the ligands and complexes formed together with the size of the chelate rings are taken into account to discuss the results presented here. # 2002 Elsevier Science Ltd. All rights reserved. Keywords: Silver(I); N -methylated polyamines; Complexes; Dimethylsulfoxide; Thermodynamics
1. Introduction In the last few years there has been a great deal of interest in the field of complex formation of metal ions with neutral N-donors in water [1 /8] as well as in nonaqueous or mixed solvents [9 /17]. The aim of these works has been to improve the knowledge of the coordination chemistry of soft and hard metal ions with these ligands, to investigate the strong influence of different basicities and steric properties of the ligands on the stability and nature of the complexes and to design ligands for selective complexation of metal ions. These studies have now also attracted renewed interest as the coordination properties of open chain polyamines towards different cations have been taken as models allowing direct comparison with corresponding macrocyclic systems [18]. In particular many papers have reported data concerning Group 11 metal complexation with these ligands [1,3,7,12,16,17]. These data have shown that the co-
* Corresponding authors. Tel.: /39-0432-558852/558882; fax: /390432-558803. E-mail addresses: [email protected] (C. Comuzzi), [email protected] (M. Tolazzi).
ordination properties of amines towards these ions may be modulated through the number and the basicity of nitrogen atoms present in the ligand, the different chelate ring sizes the ligands can form, the different Nfunctionalization and the nature of the solvent. In this context and as an extension of previous works, we report here the results of a thermodynamic study on the complex formation of Ag(I) with the following diamines in the aprotic solvent dimethylsulfoxide (dmso): N ,N -dimethyldiethylenediamine (N ,N -dmen), N ,N -dimethyl-1,3-propanediamine (N ,N -dmtn) and N ,N ,N ?,N ?-tetramethyl-1,3-propanediamine (tmtn). The ligands have been chosen in order to provide further information on how the different alkylation and/ or the different size of the potential chelate rings formed may affect their coordination ability toward metal ions. For this purpose also a comparison with data previously reported on Ag(I) complex formation with similar diamines possessing the same skeleton, but differently methylated, i.e. ethylenediamine (en) [12c], N ,N ?-dimethylethylenediamine (dmen) [12e], N ,N ,N ?N ?-tetramethylethylenediamine (tmen) [12e] and 1,3propanediamine (tn) [12c], will be carefully taken into account. Potentiometric and calorimetric measurements have been used to obtain the free energy and enthalpies of the
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reactions, respectively. All measurements have been performed at 298 K in an ionic medium adjusted to 0.1 mol dm 3 with NEt4ClO4 as neutral salt. 1H NMR spectroscopy has been used to provide insight into the coordination modes of the ligands.
periodically checked in dmso solutions containing no coordinating ligands. In the concentration range 106 B/[Ag ]B/102 mol dm 3, the emf values varied with the metal ion concentration according to Nernst’s law. The computer program HYPERQUAD [21] was used for the calculation of the stability constants.
2. Experimental 2.2. Calorimetric measurements Anhydrous silver perchlorate was obtained from AgClO4 ×/H2O (Fluka puriss) as described previously [12a]. Dimethylsulfoxide (Fluka /99%) was purified by distillation according to the described procedures ˚ molecular sieves. [12a] and stored over activated 4 A The background salt NEt4ClO4 was recrystallized twice from MeOH and dried at 110 8C. The ligands N ,N dmen (Aldrich 95%), N ,N -dmtn (Aldrich 99%) and tmtn (Aldrich 98%) were purified by fractional distillation [19]. The Ag(I) solutions were prepared from anhydrous AgClO4 and freshly distilled dmso. (Warning ! Although we encountered no problems when handling AgClO4, metal perchlorates solvated by dmso are powerful explosives under certain conditions [20] and should be handled with caution!) The metal concentration in these solutions was checked by potentiometric titrations with chloride. Solutions of the ligands were prepared by dissolving weighted amounts in the anhydrous solvent. NEt4ClO4 was used to obtain the required ionic strength. All the solutions were prepared afresh before each experiment in a MB Braun 150 glove box under atmosphere of dry nitrogen. The water content in the solutions, typically 10 /20 ppm, was determined by a Metrohm 684 KF Coulometer. 2.1. Potentiometric measurements All measurements were carried out 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 complexes were the equilibrium concentrations of the Ag(I) ion, which were obtained from the emf of a galvanic cell previously described [12a]. The emf were measured by means of an Amel 338 pH meter equipped with a Metrohm 6.0328.000 Ag electrode as a working electrode and a Metrohm 6.0718.000 Ag electrode as a reference. An experimental run consisted in collecting many equilibrium data points when solutions of silver perchlorate (2.00 B/C 8M B/20.0 mmol dm 3) were titrated with solutions of the ligands (50 B/C 8L B/200 mmol dm 3). Titrations were performed with at least three different initial Ag(I) concentrations and some titrations were carried out in duplicate to verify the reproducibility of the system. The electrode couple was
A Tronac model 87-558 precision calorimeter was employed to measure the heat of reaction. The 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 DH8//47.61 kJ mol1, in good agreement with the accepted value of /47.539/0.13 kJ mol 1 [1]. The calorimetric titrations were performed at 298.009/0.02 K by adding known volumes of ligand solutions (50 B/C 8L B/200 mmol dm 3) to 20 ml of Ag(I) solution (2.00 B/[Ag ]B/20.0 mmol dm 3). The heats of dilution of the reactants, determined in separate runs, were found negligible. The least-squares computer program LETAGROP KALLE [22] was used for the calculation of the enthalpy changes. 2.3. NMR measurements 1
H NMR spectra were recorded at 298 K on a Brucker AC 200F QNP spectrometer. 1H chemical shifts refer to SiMe4. Measurements were performed on dmsod6 solutions containing the dissolved salt AgClO4 ×/ 4dmso (approximately 20 mmol dm 3 in Ag ) and the ligand N ,N -dmtn in the molar ratio Rc /CL/CM / 0.5. Analogous measurements were carried out on a solution containing only the ligand concerned.
3. Results The computer treatment of the potentiometric data shows that the best fit is obtained when the mononuclear species reported in Table 1 are taken into account. In the table, the overall stability constants and free energies of formation, with the limits of error indicated, are listed for the reactions: Ag /jL X/ AgLj (j /1, 2; L is the ligand concerned). In Fig. 1, the distribution of Ag(I) between the different complexes as a function of the ligand-to-metal ratio, Rc / CL/CM, is plotted for all the systems studied. The experimental data obtained from the calorimetric measurements are reported in Fig. 2 as Dhn , the total heats of reaction per mole of metal ion versus Rc. From an examination of the shape of these curves some conclusions can be drawn about the prevalent
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Table 1 Overall stability constants and thermodynamic functions for the reaction Ag j L X AgLj in dmso at 298 K and I 0.1 mol dm 3
The errors quoted correspond to three standard deviations. aRef. [12c]. bRef. [12e].
coordinated species formed, which are consistent with the information obtained from the analysis of the potentiometric data. For Ag(I) /tmtn system, the splitting of the curves is only indicative of formation of weak complex(es), whereas for Ag(I) /N ,N -dmen and N ,N -dmtn systems the enthalpy data which coincide up to Rc /1, slightly diverge at the inflection point (Rc /2) and superimpose at higher Rc values can be easily explained by the formation of only two successive, almost equally stable, mononuclear complexes. The enthalpy values and the overall stability constants in Table 1 were used to calculate the full lines in Fig. 2. The match of the experimental data is quite good indicating that all the systems are satisfactorily described. Also in Table 1, the stability constants and the thermodynamic functions for the complexation reac-
tions of Ag(I) by the ligands en [12c], dmen, tmen [12e] and tn [12c] previously determined in dmso are reported, for comparison. Fig. 3 shows the 1H NMR spectra of N ,N -dmtn and of its [Ag(N ,N -dmtn)] complex in dmso-d6.
4. Discussion As a consequence of the formation of strong covalent Ag(I) /N bonds and the relatively weak solvation of the species [9] favourable enthalpy values are obtained whereas the entropy terms oppose the complex formation. In a previous work [12e] the drop in the stability constant values between en, dmen and tmen, all behaving as chelating agents, was explained as due to: (i) the
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Fig. 3. 1H NMR spectra at 298 K of: (a) N ,N -dmtn/10.1 mmol dm 3; (b) C 8Ag /20.0, N ,N -dmtn/10.1 mmol dm 3 in dmso-d6.
Fig. 1. The percent distribution of the metal ion in the Ag(I) /diamine systems in dmso at C 8Ag /10 mmol dm 3. (a) N ,N -dmen; (b) N ,N dmtn; (c) tmtn.
Fig. 2. The total molar enthalpy changes, Dhn , as a function of Rc / CL/CM for Ag(I) /diamine systems in dmso. (a) N ,N -dmen: (k) 5.10, (m) 19.85 mmol dm 3 in Ag ; (b) N ,N -dmtn: (I) 4.85, (j) 19.98 mmol dm 3 in Ag ; (c) tmtn: (2) 4.91, (") 19.55 mmol dm3 in Ag . Only some of the experimental points, chosen at random, have been plotted. The solid lines have been calculated from the values of bj and DH 8b in Table 1. j
different basicity of amino group, which decreases in the order /NH2 / /NHR / /NR2 in dmso [23]; (ii) the elongation of M /N bonds, as a consequence of the steric hindrance of the ligands [7]; and (iii) the decreased solvation of the complexes as a consequence both of their increased radii and of the minor extent of hydrogen bonding brought about by the N -alkylated amines [7].
The log b1 value here reported for the Ag(I) /N ,N dmen system is about of the same order of magnitude as that of dmen: this indicates that the ligand behaves as bidentate as confirmed by the enthalpy and entropy values. It is of interest to note that the major difference in the thermodynamic values between dmen and N ,N -dmen lies in the entropy term which is less unfavourable for the latter system. This points out that no great differences in the medium M /N bond strength are found when changing two secondary nitrogens with one primary and one tertiary ones whereas the asymmetric N -methylation in N ,N -dmen influences the first step of complexation from an entropic point of view. Evidently, both the ligand and the complex formed should be seen as less hydrogen bonded by the dmso molecules with respect to dmen system where the solvent may well interact with alongside /NHR groups, due to the formation of less crowded solvates. The thermodynamic data associated with the second complexation step indicate that in the AgL2 complex the ligand behaves also as bidentate. The higher exothermic enthalpy value of the second complexation step, DH 8K2, as compared with DH 8K1 is probably due to a particular desolvation occurring in the first step of complexation: the enthalpy needed for desolvating causes the DH8K1 to be less exothermic than DH8K2 in spite of the larger exothermicity associated to the Ag(I) interaction in the first complexation step. This hypothesis is confirmed by the unfavourable high stepwise entropy term which is also in line with the above mentioned relatively scanty outer sphere solvation of the [Ag(N ,N -dmen)] complex. As far as N ,N -dmtn is concerned, the thermodynamic data are somewhat ambigous in indicating the ligand coordination mode, i.e. if it acts as mono- or bi-dentate. A comparison of the data previously reported for monoamines [12a,12b] and tn [12c] in dmso shows that: (i) the values of log b1 and DH 81 for N ,N -dmtn are
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slightly more favourable and the DS 81 value much more unfavourable than those found previously for the monodentate n -butylamine (n-but) (log b1 /3.58; /DH 81 /31.4 kJ mol 1; /DS81 /37.3 J mol 1 K 1 [12b]; (ii) log b1 for N ,N -dmtn is about two orders of magnitude lower than that found for tn and DH81 about half the enthalpy value for tn, which has been found to be linearly chelated to Ag(I) [12c], whereas the same differences between the two chelated en and dmen are much less important (see Table 1). Point (i) seems to indicate that N ,N -dmtn is able to act as a chelating agent, whereas point (ii) seems to imply a different conclusion. Nevertheless the trend in thermodynamic functions discussed in (ii) should also be due to different geometric requirements in Ag(I) /N ,N dmtn and Ag(I) /tn systems. In fact, as mentioned above, tn was found to behave as linear bidentate in dmso according with the observed fact that large metal ions such as Ag(I) and Hg(II) will tend to have low coordination numbers and therefore large N /M /N angles [24]: this is an exception to the general behaviour of metal ions whose coordination numbers tend to become higher as metal ions become larger, behaviour which generally favours five-membered chelate ring formation more than six-membered ones [24]. As a consequence of this exception, the decreased complex stability expected for large metal ions when chelate ring size is changed from five- to six-membered, is not observed when Ag(I) is considered (see also data reported in Table 1 for en and tn systems) but, on the contrary, an increase in complex stability occurs. In conclusion, the trend in the thermodynamic data here reported seems to indicate that, if the ligand acts as bidentate, the tetrahedral configuration of the Ag(I) solvate [9] is retained: likewise, when N -methyl groups are present in N ,N -dmtn other steric requirements are added to the six-membered chelate ring formed which prevents Ag(I) from being coordinated linearly. The 1H NMR spectra of N ,N -dmtn and of its Ag(I) complex may be of some help in investigating the coordination mode of the ligand concerned. In Fig. 3, a Dd (Dd/dcomplex/dligand) of about 0.1 is observed for both N -methyl groups (from 2.18 to 2.08) and the CH2 (C) (see Scheme 1) (from 2.32 to 2.22) whereas the Dd of CH2 (A) is about 0.2 ppm (from 2.72 to 2.52). For comparison also the 1H NMR spectra of the monodentate n -but and of its Ag(I) complex were
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collected and it was observed that the multiplet signal of CH2 (at 1.30) in g position with respect to /NH2 is unaffected by complexation. The small downfield shift of CH2 (C) (and of methyl protons) observed in Ag(I) / N ,N -dmtn system should be therefore taken as reflecting an interaction between the /NR2 moiety and the metal ion. Nevertheless, this M /NR2 interaction is weak, given that the formation of a six-membered chelate ring, normally more strained than that of a five-membered ring in a tetrahedral arrangement [24], is assisted by a tertiary, low basic, amino group [23]. The weakness of this interaction is supported not only by the values of the thermodynamic functions associated to [Ag(N ,N dmtn)] formation, but also by the absence of polynuclear species which have usually been found to be formed in dmso when chelating ligands having primary nitrogen atoms separated by more than two methylene groups were considered [12c,12d]. The thermodynamic parameters associated to the second complexation step are very close to those corresponding to the formation of [Ag(n-but)2] complex (log b2 /7.34; /DH 82 /71.5 and /DS 82 /99.6) [12b], thus indicating that N ,N -dmtn behaves as monodentate in this complexation step. For tmtn system, the data clearly indicate that the ligand behaves as monodentate in both the complexation steps: the log b1 value is even lower than that found for Ag(I) /tributylamine (tribut) system (log b1 tribut / 2.2) [17]. The enthalpy value for 1:1 Ag(I) /tribut complex has not been reported: therefore the most reliable comparison with literature data is with the enthalpy values relative to the secondary monodentate dibutylamine (dibut) which are, respectively DH 81 / /31.80 and DH 82 //61.60 kJ mol 1 for the 1:1 and 1:2 Ag(I) /dibut complexes [12b]. These values are certainly much more favourable than those found for tmtn system further confirming monodentation. In conclusion, it is of interest to note that symmetric dimethylation of an amino group in a potential sixmembered N ,N -dmtn chelate ligand reverses the peculiar feature of silver ion to form more stable sixmembered chelate rings with respect to five-membered chelate ones, strongly influencing the efficiency of the ligands concerned toward Ag(I). In the case of the tetramethylated 1,3-propanediamine the low basicity of the tertiary amino groups and the important steric requirement of the ligand even prevent chelation.
5. Supplementary material Scheme 1.
The material is available from the authors on request.
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Acknowledgements This work has been supported by the Ministero dell’Universita` e della Ricerca Scientifica e Tecnologica (MURST, Rome) within the program COFIN 2000. We thank Pierluigi Polese for technical assistance.
References [1] (a) R.M. Smith, A.E. Martell, Critical Stability Constants, vol. 6 (2nd Supplement), Plenum Press, New York, 1989; (b) A.E. Martell, R.M. Smith, R.J. Motekaitis, Critically Selected Stability Constants of Metal Complexes Database, version 5.0, 1998. [2] R.M. Clay, S. Corr, M. Micheloni, P. Paoletti, Inorg. Chem. 24 (1985) 3330. [3] R. Hancock, A.E. Martell, Chem. Rev. 89 (1989) 1875. [4] A. Bencini, A. Bianchi, P. Paoletti, Coord. Chem. Rev. 120 (1992) 51. [5] F.C.J.M. van Veggel, W. Verboom, D.N. Reinboudt, Chem. Rev. 94 (2) (1994) 279. [6] C. Bazzicalupi, A. Bencini, A. Bianchi, V. Fusi, E. GarciaEspana, P. Paoletti, P. Paoli, B. Voltancoli, Inorg. Chem. 32 (1995) 4900. [7] D. Meyerstein, Coord. Chem. Rev. 185 /186 (1999) 141 (and references therein). [8] G. Golub, H. Cohen, P. Paoletti, A. Bencini, D. Meyerstein, J. Chem. Soc., Dalton Trans. (1996) 2055. [9] S. Ahrland, in: J.J. Lagowsky (Ed.), The Chemistry of Nonaqueous Solvents, vol. 5a (and references therein), Academic Press, New York, 1978. [10] B.G. Cox, J. Garcia-Rosas, H. Schneider, J. Am. Chem. Soc. 103 (1981) 1384. [11] H.K. Chantooni, I.M. Kolthoff, J. Solution Chem. 14 (1985) 1.
[12] (a) A. Cassol, P. Di Bernardo, P.L. Zanonato, R. Portanova, M. Tolazzi, J. Chem. Soc., Dalton Trans. (1987) 657; (b) A. Cassol, P. Di Bernardo, P.L. Zanonato, R. Portanova, M. Tolazzi, G. Tomat, J. Chem. Soc., Dalton Trans. (1988) 1781; (c) A. Cassol, P. Di Bernardo, P.L. Zanonato, R. Portanova, M. Tolazzi, G. Tomat, V. Cucinotta, D. Sciotto, J. Chem. Soc., Faraday Trans. I 85 (8) (1989) 2445; (d) A. Cassol, P. Di Bernardo, P.L. Zanonato, R. Portanova, M. Tolazzi, G. Tomat, J. Chem. Soc., Faraday Trans. 86 (16) (1990) 2841; (e) C. Comuzzi, V. Novelli, R. Portanova, M. Tolazzi, Supramol. Chem. 13 (3) (2001) 455. [13] A.F. De Namor, J. Chem. Soc., Faraday Trans. 1 84 (7) (1988) 2441. [14] A. Lewandowski, A. Szukalska, M. Galinski, J. Chem. Soc., Faraday Trans. 91 (7) (1995) 1097 (and references therein). [15] F. Arnaud-Neu, Chem. Soc. Rev. (1994) 235. [16] N. Heidari, A. Thaler, H. Schneider, B.G. Cox, Inorg. Chim. Acta 279 (1998) 186. [17] A. Thaler, N. Heidari, B.G. Cox, H. Schneider, Inorg. Chim. Acta 286 (1999) 160. [18] I. Bertini, L. Messori, G. Galub, H. Cohen, D. Meyerstein, Inorg. Chim. Acta 235 (1995) 5. [19] D.D. Perrin, W.L.F. Armarego, D.R. Perrin, Purification of Laboratory Chemicals, 2nd ed., Pergamon, Oxford, 1980. [20] (a) S. Ahrland, I. Persson, Acta Chem. Scand., Sect. A 34 (1980) 645; (b) D. Martin, H.G. Hauthal, Dimethyl Sulfoxide, van Nostrand Reinhold, Wokingham, Berkshire, 1975, pp. 448 /449. [21] P. Gans, A. Sabatini, A. Vacca, Talanta 43 (1996) 1739. [22] R. Arnek, Ark. Kemi 32 (1970) 81. [23] (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. [24] A.E. Martell, R.D. Hancock, in: J.P. Fackler, Jr. (Ed.), Metal Complexes in Aqueous Solutions, Plenum, New York, 1996.
The influence of N-alkylation and chain length in the coordination ability of diamines towards silver(I) in dimethylsulfoxide Clara Comuzzi *, Veronica Novelli, Roberto Portanova, Marilena Tolazzi * Dipartimento di Scienze e Tecnologie Chimiche, Universita` di Udine, Via del Cotonificio 108, I-33100 Udine, Italy Received 27 September 2001; accepted 16 October 2001
Abstract The thermodynamic functions for the complexation of Ag(I) by the following diamines: N ,N -dimethyldiethylenediamine (N ,N dmen), N ,N -dimethyl-1,3-propanediamine (N ,N -dmtn) and N ,N ,N ?,N ?-tetramethyl-1,3-propanediamine (tmtn) have been determined in dimethylsulfoxide (dmso) by potentiometric and calorimetric techniques at 298 K and 0.1 mol dm 3 ionic strength (NEt4ClO4). Only mononuclear complexes are formed (AgLj , j /1, 2) where the ligands act as monodentate or chelate agents. All the complexes are enthalpy stabilized whereas the entropy changes counteract the complexation. The different basicities and steric requirements of both the ligands and complexes formed together with the size of the chelate rings are taken into account to discuss the results presented here. # 2002 Elsevier Science Ltd. All rights reserved. Keywords: Silver(I); N -methylated polyamines; Complexes; Dimethylsulfoxide; Thermodynamics
1. Introduction In the last few years there has been a great deal of interest in the field of complex formation of metal ions with neutral N-donors in water [1 /8] as well as in nonaqueous or mixed solvents [9 /17]. The aim of these works has been to improve the knowledge of the coordination chemistry of soft and hard metal ions with these ligands, to investigate the strong influence of different basicities and steric properties of the ligands on the stability and nature of the complexes and to design ligands for selective complexation of metal ions. These studies have now also attracted renewed interest as the coordination properties of open chain polyamines towards different cations have been taken as models allowing direct comparison with corresponding macrocyclic systems [18]. In particular many papers have reported data concerning Group 11 metal complexation with these ligands [1,3,7,12,16,17]. These data have shown that the co-
* Corresponding authors. Tel.: /39-0432-558852/558882; fax: /390432-558803. E-mail addresses: [email protected] (C. Comuzzi), [email protected] (M. Tolazzi).
ordination properties of amines towards these ions may be modulated through the number and the basicity of nitrogen atoms present in the ligand, the different chelate ring sizes the ligands can form, the different Nfunctionalization and the nature of the solvent. In this context and as an extension of previous works, we report here the results of a thermodynamic study on the complex formation of Ag(I) with the following diamines in the aprotic solvent dimethylsulfoxide (dmso): N ,N -dimethyldiethylenediamine (N ,N -dmen), N ,N -dimethyl-1,3-propanediamine (N ,N -dmtn) and N ,N ,N ?,N ?-tetramethyl-1,3-propanediamine (tmtn). The ligands have been chosen in order to provide further information on how the different alkylation and/ or the different size of the potential chelate rings formed may affect their coordination ability toward metal ions. For this purpose also a comparison with data previously reported on Ag(I) complex formation with similar diamines possessing the same skeleton, but differently methylated, i.e. ethylenediamine (en) [12c], N ,N ?-dimethylethylenediamine (dmen) [12e], N ,N ,N ?N ?-tetramethylethylenediamine (tmen) [12e] and 1,3propanediamine (tn) [12c], will be carefully taken into account. Potentiometric and calorimetric measurements have been used to obtain the free energy and enthalpies of the
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reactions, respectively. All measurements have been performed at 298 K in an ionic medium adjusted to 0.1 mol dm 3 with NEt4ClO4 as neutral salt. 1H NMR spectroscopy has been used to provide insight into the coordination modes of the ligands.
periodically checked in dmso solutions containing no coordinating ligands. In the concentration range 106 B/[Ag ]B/102 mol dm 3, the emf values varied with the metal ion concentration according to Nernst’s law. The computer program HYPERQUAD [21] was used for the calculation of the stability constants.
2. Experimental 2.2. Calorimetric measurements Anhydrous silver perchlorate was obtained from AgClO4 ×/H2O (Fluka puriss) as described previously [12a]. Dimethylsulfoxide (Fluka /99%) was purified by distillation according to the described procedures ˚ molecular sieves. [12a] and stored over activated 4 A The background salt NEt4ClO4 was recrystallized twice from MeOH and dried at 110 8C. The ligands N ,N dmen (Aldrich 95%), N ,N -dmtn (Aldrich 99%) and tmtn (Aldrich 98%) were purified by fractional distillation [19]. The Ag(I) solutions were prepared from anhydrous AgClO4 and freshly distilled dmso. (Warning ! Although we encountered no problems when handling AgClO4, metal perchlorates solvated by dmso are powerful explosives under certain conditions [20] and should be handled with caution!) The metal concentration in these solutions was checked by potentiometric titrations with chloride. Solutions of the ligands were prepared by dissolving weighted amounts in the anhydrous solvent. NEt4ClO4 was used to obtain the required ionic strength. All the solutions were prepared afresh before each experiment in a MB Braun 150 glove box under atmosphere of dry nitrogen. The water content in the solutions, typically 10 /20 ppm, was determined by a Metrohm 684 KF Coulometer. 2.1. Potentiometric measurements All measurements were carried out 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 complexes were the equilibrium concentrations of the Ag(I) ion, which were obtained from the emf of a galvanic cell previously described [12a]. The emf were measured by means of an Amel 338 pH meter equipped with a Metrohm 6.0328.000 Ag electrode as a working electrode and a Metrohm 6.0718.000 Ag electrode as a reference. An experimental run consisted in collecting many equilibrium data points when solutions of silver perchlorate (2.00 B/C 8M B/20.0 mmol dm 3) were titrated with solutions of the ligands (50 B/C 8L B/200 mmol dm 3). Titrations were performed with at least three different initial Ag(I) concentrations and some titrations were carried out in duplicate to verify the reproducibility of the system. The electrode couple was
A Tronac model 87-558 precision calorimeter was employed to measure the heat of reaction. The 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 DH8//47.61 kJ mol1, in good agreement with the accepted value of /47.539/0.13 kJ mol 1 [1]. The calorimetric titrations were performed at 298.009/0.02 K by adding known volumes of ligand solutions (50 B/C 8L B/200 mmol dm 3) to 20 ml of Ag(I) solution (2.00 B/[Ag ]B/20.0 mmol dm 3). The heats of dilution of the reactants, determined in separate runs, were found negligible. The least-squares computer program LETAGROP KALLE [22] was used for the calculation of the enthalpy changes. 2.3. NMR measurements 1
H NMR spectra were recorded at 298 K on a Brucker AC 200F QNP spectrometer. 1H chemical shifts refer to SiMe4. Measurements were performed on dmsod6 solutions containing the dissolved salt AgClO4 ×/ 4dmso (approximately 20 mmol dm 3 in Ag ) and the ligand N ,N -dmtn in the molar ratio Rc /CL/CM / 0.5. Analogous measurements were carried out on a solution containing only the ligand concerned.
3. Results The computer treatment of the potentiometric data shows that the best fit is obtained when the mononuclear species reported in Table 1 are taken into account. In the table, the overall stability constants and free energies of formation, with the limits of error indicated, are listed for the reactions: Ag /jL X/ AgLj (j /1, 2; L is the ligand concerned). In Fig. 1, the distribution of Ag(I) between the different complexes as a function of the ligand-to-metal ratio, Rc / CL/CM, is plotted for all the systems studied. The experimental data obtained from the calorimetric measurements are reported in Fig. 2 as Dhn , the total heats of reaction per mole of metal ion versus Rc. From an examination of the shape of these curves some conclusions can be drawn about the prevalent
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Table 1 Overall stability constants and thermodynamic functions for the reaction Ag j L X AgLj in dmso at 298 K and I 0.1 mol dm 3
The errors quoted correspond to three standard deviations. aRef. [12c]. bRef. [12e].
coordinated species formed, which are consistent with the information obtained from the analysis of the potentiometric data. For Ag(I) /tmtn system, the splitting of the curves is only indicative of formation of weak complex(es), whereas for Ag(I) /N ,N -dmen and N ,N -dmtn systems the enthalpy data which coincide up to Rc /1, slightly diverge at the inflection point (Rc /2) and superimpose at higher Rc values can be easily explained by the formation of only two successive, almost equally stable, mononuclear complexes. The enthalpy values and the overall stability constants in Table 1 were used to calculate the full lines in Fig. 2. The match of the experimental data is quite good indicating that all the systems are satisfactorily described. Also in Table 1, the stability constants and the thermodynamic functions for the complexation reac-
tions of Ag(I) by the ligands en [12c], dmen, tmen [12e] and tn [12c] previously determined in dmso are reported, for comparison. Fig. 3 shows the 1H NMR spectra of N ,N -dmtn and of its [Ag(N ,N -dmtn)] complex in dmso-d6.
4. Discussion As a consequence of the formation of strong covalent Ag(I) /N bonds and the relatively weak solvation of the species [9] favourable enthalpy values are obtained whereas the entropy terms oppose the complex formation. In a previous work [12e] the drop in the stability constant values between en, dmen and tmen, all behaving as chelating agents, was explained as due to: (i) the
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Fig. 3. 1H NMR spectra at 298 K of: (a) N ,N -dmtn/10.1 mmol dm 3; (b) C 8Ag /20.0, N ,N -dmtn/10.1 mmol dm 3 in dmso-d6.
Fig. 1. The percent distribution of the metal ion in the Ag(I) /diamine systems in dmso at C 8Ag /10 mmol dm 3. (a) N ,N -dmen; (b) N ,N dmtn; (c) tmtn.
Fig. 2. The total molar enthalpy changes, Dhn , as a function of Rc / CL/CM for Ag(I) /diamine systems in dmso. (a) N ,N -dmen: (k) 5.10, (m) 19.85 mmol dm 3 in Ag ; (b) N ,N -dmtn: (I) 4.85, (j) 19.98 mmol dm 3 in Ag ; (c) tmtn: (2) 4.91, (") 19.55 mmol dm3 in Ag . Only some of the experimental points, chosen at random, have been plotted. The solid lines have been calculated from the values of bj and DH 8b in Table 1. j
different basicity of amino group, which decreases in the order /NH2 / /NHR / /NR2 in dmso [23]; (ii) the elongation of M /N bonds, as a consequence of the steric hindrance of the ligands [7]; and (iii) the decreased solvation of the complexes as a consequence both of their increased radii and of the minor extent of hydrogen bonding brought about by the N -alkylated amines [7].
The log b1 value here reported for the Ag(I) /N ,N dmen system is about of the same order of magnitude as that of dmen: this indicates that the ligand behaves as bidentate as confirmed by the enthalpy and entropy values. It is of interest to note that the major difference in the thermodynamic values between dmen and N ,N -dmen lies in the entropy term which is less unfavourable for the latter system. This points out that no great differences in the medium M /N bond strength are found when changing two secondary nitrogens with one primary and one tertiary ones whereas the asymmetric N -methylation in N ,N -dmen influences the first step of complexation from an entropic point of view. Evidently, both the ligand and the complex formed should be seen as less hydrogen bonded by the dmso molecules with respect to dmen system where the solvent may well interact with alongside /NHR groups, due to the formation of less crowded solvates. The thermodynamic data associated with the second complexation step indicate that in the AgL2 complex the ligand behaves also as bidentate. The higher exothermic enthalpy value of the second complexation step, DH 8K2, as compared with DH 8K1 is probably due to a particular desolvation occurring in the first step of complexation: the enthalpy needed for desolvating causes the DH8K1 to be less exothermic than DH8K2 in spite of the larger exothermicity associated to the Ag(I) interaction in the first complexation step. This hypothesis is confirmed by the unfavourable high stepwise entropy term which is also in line with the above mentioned relatively scanty outer sphere solvation of the [Ag(N ,N -dmen)] complex. As far as N ,N -dmtn is concerned, the thermodynamic data are somewhat ambigous in indicating the ligand coordination mode, i.e. if it acts as mono- or bi-dentate. A comparison of the data previously reported for monoamines [12a,12b] and tn [12c] in dmso shows that: (i) the values of log b1 and DH 81 for N ,N -dmtn are
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slightly more favourable and the DS 81 value much more unfavourable than those found previously for the monodentate n -butylamine (n-but) (log b1 /3.58; /DH 81 /31.4 kJ mol 1; /DS81 /37.3 J mol 1 K 1 [12b]; (ii) log b1 for N ,N -dmtn is about two orders of magnitude lower than that found for tn and DH81 about half the enthalpy value for tn, which has been found to be linearly chelated to Ag(I) [12c], whereas the same differences between the two chelated en and dmen are much less important (see Table 1). Point (i) seems to indicate that N ,N -dmtn is able to act as a chelating agent, whereas point (ii) seems to imply a different conclusion. Nevertheless the trend in thermodynamic functions discussed in (ii) should also be due to different geometric requirements in Ag(I) /N ,N dmtn and Ag(I) /tn systems. In fact, as mentioned above, tn was found to behave as linear bidentate in dmso according with the observed fact that large metal ions such as Ag(I) and Hg(II) will tend to have low coordination numbers and therefore large N /M /N angles [24]: this is an exception to the general behaviour of metal ions whose coordination numbers tend to become higher as metal ions become larger, behaviour which generally favours five-membered chelate ring formation more than six-membered ones [24]. As a consequence of this exception, the decreased complex stability expected for large metal ions when chelate ring size is changed from five- to six-membered, is not observed when Ag(I) is considered (see also data reported in Table 1 for en and tn systems) but, on the contrary, an increase in complex stability occurs. In conclusion, the trend in the thermodynamic data here reported seems to indicate that, if the ligand acts as bidentate, the tetrahedral configuration of the Ag(I) solvate [9] is retained: likewise, when N -methyl groups are present in N ,N -dmtn other steric requirements are added to the six-membered chelate ring formed which prevents Ag(I) from being coordinated linearly. The 1H NMR spectra of N ,N -dmtn and of its Ag(I) complex may be of some help in investigating the coordination mode of the ligand concerned. In Fig. 3, a Dd (Dd/dcomplex/dligand) of about 0.1 is observed for both N -methyl groups (from 2.18 to 2.08) and the CH2 (C) (see Scheme 1) (from 2.32 to 2.22) whereas the Dd of CH2 (A) is about 0.2 ppm (from 2.72 to 2.52). For comparison also the 1H NMR spectra of the monodentate n -but and of its Ag(I) complex were
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collected and it was observed that the multiplet signal of CH2 (at 1.30) in g position with respect to /NH2 is unaffected by complexation. The small downfield shift of CH2 (C) (and of methyl protons) observed in Ag(I) / N ,N -dmtn system should be therefore taken as reflecting an interaction between the /NR2 moiety and the metal ion. Nevertheless, this M /NR2 interaction is weak, given that the formation of a six-membered chelate ring, normally more strained than that of a five-membered ring in a tetrahedral arrangement [24], is assisted by a tertiary, low basic, amino group [23]. The weakness of this interaction is supported not only by the values of the thermodynamic functions associated to [Ag(N ,N dmtn)] formation, but also by the absence of polynuclear species which have usually been found to be formed in dmso when chelating ligands having primary nitrogen atoms separated by more than two methylene groups were considered [12c,12d]. The thermodynamic parameters associated to the second complexation step are very close to those corresponding to the formation of [Ag(n-but)2] complex (log b2 /7.34; /DH 82 /71.5 and /DS 82 /99.6) [12b], thus indicating that N ,N -dmtn behaves as monodentate in this complexation step. For tmtn system, the data clearly indicate that the ligand behaves as monodentate in both the complexation steps: the log b1 value is even lower than that found for Ag(I) /tributylamine (tribut) system (log b1 tribut / 2.2) [17]. The enthalpy value for 1:1 Ag(I) /tribut complex has not been reported: therefore the most reliable comparison with literature data is with the enthalpy values relative to the secondary monodentate dibutylamine (dibut) which are, respectively DH 81 / /31.80 and DH 82 //61.60 kJ mol 1 for the 1:1 and 1:2 Ag(I) /dibut complexes [12b]. These values are certainly much more favourable than those found for tmtn system further confirming monodentation. In conclusion, it is of interest to note that symmetric dimethylation of an amino group in a potential sixmembered N ,N -dmtn chelate ligand reverses the peculiar feature of silver ion to form more stable sixmembered chelate rings with respect to five-membered chelate ones, strongly influencing the efficiency of the ligands concerned toward Ag(I). In the case of the tetramethylated 1,3-propanediamine the low basicity of the tertiary amino groups and the important steric requirement of the ligand even prevent chelation.
5. Supplementary material Scheme 1.
The material is available from the authors on request.
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Acknowledgements This work has been supported by the Ministero dell’Universita` e della Ricerca Scientifica e Tecnologica (MURST, Rome) within the program COFIN 2000. We thank Pierluigi Polese for technical assistance.
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