Polarographic study of mercury complexes with some recently synthesized benzo-substituted macrocyclic diamides in binary acetonitrile + water mixtures

J~R~L

ELSEVIER

OF

Journal of Electroanalytical Chemistry 405 (1996) 177-181

Polarographic study of mercury complexes with some recently synthesized benzo-substituted macrocyclic diamides in binary acetonitrile + water mixtures Mohammad Reza Ganjali a, Hossein Eshghi b, Hashem Sharghi b, Mojtaba Shamsipur c,. a Department of Chemistry, Tehran Unioersity, Tehran, lran b Department of Chemistry, Shiraz Unioersity, Shiraz, lran c Department of Chemistry, Razi University, Kermanshah, lran Received 6 September 1995; in revised form 17 October 1995

Abstract The complex formation of Hg 2+ ion with some benzo-substituted macrocyclic diamides in binary acetonitrile + water mixtures was studied by differential pulse polarography at 25°C. The stoichiometry and stability of the complexes were determined by monitoring the shift in the Hg 2+ differential pulse peak potential against the macrocycle concentration. In all cases studied, it was found that the stability of the resulting 1 : 1 complex decreases drastically by increasing the amount of water in the binary mixtures. The observed stability order in a given solvent mixture is discussed in terms of the cavity size, structural flexibility and nature of the substituents on the macrocyclic ring. Keywords: Mercury; Polarography; Macrocyclic diamides; Mixed solvent

1. Introduction The design and synthesis of new functionalized macrocyclic ligands for some specific applications is of increasing interest [1-3]. Among these macrocyclic ligands, derivatives of crown ethers and azacrown ethers are of great importance owing to their applications in separation, chemical and biological sciences [4-7]. The efficient synthesis of some interesting benzo-substituted macrocyclic diamides has been recently reported by this research group [8] and others [9]. Among a variety of physicochemical methods used for the study of cation-macrocycle interactions [10], polarographic methods are well known techniques for the study of the electrochemical behavior of macrocyclic ligands and the stability and selectivity of their complexes with various cations in both aqueous and non-aqueous media [11-21]. Despite the importance of heavy metal ion complexes of macrocyclic ligands in chemistry and industry [5,6], little attention has been focused on the study of these complexes, compared with the extensive amount of re-

* Corresponding author. 0022-0728/96/$15.00 © 1996 Elsevier Science S.A. All rights reserved SSDI 0 0 2 2 - 0 7 2 8 ( 9 5 ) 0 4 4 1 3-2

search work on the corresponding complexes with alkali and alkaline earth complexes [10]. Thus, we were recently involved in the study of transition and heavy metal complexes with different macrocyclic and macrobicyclic ligands in non-aqueous and mixed solvents [18-24]. In this paper we report a differential pulse polarographic study of Hg 2+ complexes with eight recently synthesized benzosubstituted macrocyclic diamides [8] in acetonitrile + water mixtures at 25°C. The structures of the ligands are shown in Fig. 1.

2. Experimental

2.1. Reagents Benzo-substituted macrocyclic diamides 1,15-diaza3,4; 12,13-dibenzo-5,8,11 -trioxacycloheptadecane-2,14-dione (L1), 1,15-diaza-3,4;12,13-dibenzo-17-methyl-5,8,11trioxacycloheptadecane-2,14-dione (L2), 1,15-diaza3,14; 12,13-dibenzo-5,8,11-trioxacyclouneicosane-2,14dione (L3), 1,15-diaza-3,4;12,13;16,17-tribenzo-5,8,11trioxacycloheptadecane-2,14-dione (L4), 1,15-diaza-

M.R. Ganjali et al. / Journal of Electroanalytical Chemistry 405 (1996) 177-181

178

(g8)

0.025 M and the concentration range of the macrocyclic ligands added was 1.0 to 10 mM. It is well known that mercury from the DME can be oxidized anodically to the bivalent state leading to the overall electrode process [26]

(LI)

X = -,(CH2)2-

(L2)

X = -CHICHCH3

(L3)

x = -(CH2),-

Hg = Hg2++ 2e-

(L4)

X =~

(Ls)

x

(L6)

X = ~v-~-U~

In the presence of most complexing agents, this oxidation reaction can be facilitated by chemical complexation following the electrochemical oxidation of mercury [26]:

(L7)

X = -CH2CH2NHCH2CI-I~-

Fig. 1. Structures of benzo-substituted macrocyclic diamides used.

3,4; 12,13; 16,19-tribenzo-5,8,11 -trioxacyclononadecane2,14-dione (L5), 1,15-diaza-3,4;12,3-dibenzo-16,18naphthalino-5,8,11 -trioxacyclooctadecane-2,14-dione (L6), 1,15,18-triaza-3,4; 12,13-dibenzo-5,8,1 t -trioxa-cycloeicosane-2,14-dione (L7) and 1,15-diaza-3,4;12,13-dibenzo5,8,11-trioxa-bicyclo[ 13,2,2]heptadecane-2,14-dione (L8) were prepared, purified and dried as described before [8] (Fig. 1). Reagent grade tetraethylammonium perchlorate (TEAP, Merck) was used without any further purification except for vacuum drying over P2Os. HPLC grade actonitrile (Fluka) and triply distilled deionized water were used for the preparation of the solvent mixtures. 2.2. Apparatus The polarographic measurements were carried out with a dropping mercury electrode (DME) in a three-electrode arrangement. The counter electrode was a platinum wire with a considerably larger surface area than that of the DME. A silver Isilver chloride Isaturated KC1 reference electrode was placed in a 0.1 M TEAP solution in nonaqueous solvents and connected to the electrolyzed solution by means of a bridge containing the base electrolyte. All the solutions were deaerated for 10 min with pure argon and an inert atmosphere was maintained over the solution during the oxidation. A Polarecord E-506 Metrohm Herisau instrument Was used for the polarographic measurements. The usual instrumental parameters were: constant drop time, 0.8 s; mercury height, 50 cm; scan rate, 5 mV s - l ; pulse duration, 0.02 s; pulse magnitude, 40 mV. All experiments were carried out at 25 _ 0.1°C using a model FK2 Haake thermostat with water bath. 2.3. Procedure The determination of the stability constants of Hg 2+crown ether complexes, studied in different solvents, was based on measurements of the half-wave potentials El~ 2 brought about by the addition of an increasing amount of the ligands [25]. In all experiments the concentration of TEAP, as supporting electrolyte, was kept constant at

(1)

Hg + e L = [HgL~] 2* + 2e-

(2)

The half-wave potential Ell 2 of this process is related to the ligand concentration [LJand the stability constant /3p, by the well-known equation [25,27] RT 2 (p- ])D 1/2 El~ 2 = E~g + 2 F In P ( D ' ) 1/2/3p[L](,-l)

(3)

where D and D' are the diffusion coefficients of the ligand and complex in solution respectively; the ratio of these two values is reasonably assumed to be unity [25]. It has been reported that the peak potential Ep in differential pulse polarography is related to the half-wave potential El~ 2 of the corresponding d.c. polarogram by the expression [28,29] Ep = E l / 2 - - A E / 2 , where A E is the magnitude of the differential pulse. In cases where A E is small, Ep lies very close to the value of El~ 2 [29]. It is clear that Ep and El~ 2 are linearly related and thus it is possible to replace El~ z with Ep in Eq. (3). The free ligand concentration [L] in equilibrium with Hg 2+ depends on the total concentration c L of ligand and pH according to the equation [L] = CL/OtL(n),

(4)

where at4H) is the proton side reaction coefficient. Considering the fact that Ep-.~ El~2, substitution of [L] from Eq. (4) into Eq. (3) and rearrangement gives Ep=E~g-

RT RT 2---~ln /3p-- 2 F l n c(LP-l)

RT 2(p - 1) + 2Fln ~ ( p - 1) P

(5)

In the mixed solvents used there is no proton side reaction; ~t~H) Can be taken as unity. Thus, Eq. (5) can be simplified to RT Ep = E~lg - - - I n 2F

RT /3p -

RT 2(p - 1) In c(Le- o + - - l n ~ 2F 2F p

(6) where E~g is the formal potential of the HgZ+/Hg couple. The reported value for E~g in aqueous solution is 0.589 V vs. SCE at an ionic strength of 1.0 M [30]. Unfortunately, reports on E~g in mixed and non-aqueous solutions are quite sparse [31]. But, as in aqueous solution, the formal

M.R. Ganjali et al. / Journal of Electroanalytical Chemistry 405 (1996) 177-181

I

.~25

0.400

0.~75

0.~50 E (v)

0.525|

0.~00

0.~75

Fig. 2. Anodic differential pulse polarograms of mercury electrode in 0.025 M TEAP in different acetonitrile + water mixtures in the presence of 3 mM of ligand L2 at 25°C. Wt.% acetonitrile in the solvent mixture is: I (90%), 2 (80%), 3 (70%), 4 (60%).

sponding to a reversible two-electron oxidation ( 3 0 _ 2 mV). The observation of relatively narrow differential pulse peaks in all cases studied further supports the reversible behavior of the H g ° / H g 2÷ couple [28,29]. This reversible behavior indicates that the exchange between the Hg2+-crown species is rapid on the measurement time scale. The anodic oxidation potentials of the mercury electrode in various solvent mixtures were measured and are given in Table 1. It is interesting to note that there is a good linear correlation between the oxidation potential of mercury and the molar fraction X w of water in the mixed solvent. The higher the molar fraction of water, the easier the oxidation and dissolution of Hg ° in the solution. The formation constants of the resulting 1 : 1 complexes were determined from the following simplified equation: gp - g~g = - ( RT/2F)ln

potentials should be very close to the mercury electrode potentials. Therefore, in the mixed solvents used in this study, the oxidation potentials of the mercury electrode were used for E~g in Eq. (6) [32]. The oxidation potentials of the mercury electrode in the acetonitrile + water mixtures used were obtained from the intersection of the anodic pulse polarograms of the corresponding 0.025 M TEAP solutions in the absence of any ligand with their base lines (see, for example, polarogram 1 in Fig. 2). Errors associated with the potentials were reported as + SD from at least three replicate measurements. However, it should be noted that, owing to some differences in the properties of various solvents used, the existence of some different contributions from possible liquid junction potentials (although at a low relative level) to the measured E~g values cannot be neglected.

3. R e s u l t s a n d d i s c u s s i o n

In all four acetonitrile + water mixtures used, the electrochemical process for the oxidation of the mercury electrode was reversible and diffusion controlled. The same kind of electrochemical behavior was recently reported in pure acetonitrile solution [20,32]. As an example, the anodic differential pulse polarograms of the mercury electrode in the presence of ligand L2 in different solvent mixtures are shown in Fig. 2. It is interesting to note that, in all cases studied, the peak potentials for electrochemical oxidation of mercury in the presence of different macrocyclic diamides were found to be independent of the ligand concentration, while only an increase in the peak current was observed with increasing ligand concentration. This behavior, which is in accordance with p = 1 in Eq. (6), indicates the formation of 1 : 1 complexes of Hg 2+ with the ligands in solution. The Eappl vs. l o g [ I / ( I d - I)] for the corresponding d.c. polarograms gave straight lines of Nernstian slope corre-

179

fll

(7)

by the measurement of Ep and E~g in the presence and absence of ligands respectively. Since Ep and E~g are measured under the same experimental conditions, except for the presence of neutral crown ethers in the case of Ep, the effects of possible liquid junction potentials would cancel out in Eq. (7). All calculated formation constants for the resulting 1 • 1 complexes between Hg 2+ ion and ligands L I - L 8 in different acetonitrile + water mixtures at 25°C are summarized in Table 2. It is immediately obvious that the solvent properties have a very fundamental effect on the stability and selectivity of the resulting complexes. In all cases, the stability of the complexes decreases drastically with increasing weight percent of water in the mixed solvent. It has been shown that the solvating ability of the solvent, as expressed by the Gutmann donicity number [33], plays an important role in the complexation reactions [10,12,13,1824]. Water is a solvent of high solvating ability (DN = 33) [34], which can compete strongly with the ligands for Hg 2+ ion. Thus, it is not surprising that the addition of water to acetonitrile as a relatively low donicity solvent ( D N - - 14.1) [33] will decrease the extent of interaction between the ligand donating atoms and Hg 2+ ion. In addition, the lower dielectric constant of acetonitrile (38.0) compared with that of water (78.5) would also cause

Table 1 Oxidation potentials of the mercury electrode in different acetonitrile+ water mixtures containing 0.025 M TEAP at 25°C Solvent composition Potential/mV /wt.% acetonitrile in water a 90 625+5 80 610+5 70 600+5 60 585+5 a The solutions have the following Xw: 90% (0.20), 80% (0.36), 70% (0.49), 60% (0.60).

M.R. Ganjali et al./ Journal of Electroai alytical Chemistry 405 (1996) 177-181

180

Table 2 Stability constants of Hg 2+ complexes with different macrocyclic diamide in various acetonitrile + water mixtures (wt% acetonitrile in water) at 25°C Ligand

log fit a 90

LI L2 L3 L4 L5 L6 L7 L8

4.20 4.33 3.10 3.25 3.05 2.83 4.50 4.83

80 + 0.03 + 0.03 5:0.04 -1- 0.05 + 0.04 + 0.05 + 0.02 + 0.03

3.05 3.00 2.05 2.33 2.15 2.50 2.83 3.73

70 + + + + + + + +

0.03 0.04 0.05 0.05 0.05 0.05 0.02 0.03

2.10 2.00 1.15 1.83 1.43 2.33 1.66 2.83

60 + 0.05 + 0.05 ::k 0.06 + 0.05 + 0.05 + 0.05 + 0.05 + 0.04

1.15 + 0.05 0.83 + 0.06 0.83 + 0.06 1.15 + 0.06 0.83:5:0.06 1.50 + 0.06 0.66 + 0.07 2.16 + 0.05

a The errors associated with stability constants are given as + SD.

the electrostatic contributions to the bond formation to decrease with increasing percentage of water in the solvent mixtures. It is interesting to note that, in all cases studied, there is actually a linear relationship between log fll and the molar fraction X w of water in the mixed solvents. The same trend has already been reported for a variety of complexes in different solvent mixtures [18,21,23,24,3540]. It seems reasonable to assume that the preferential hydration of Hg 2÷ ion is mainly responsible for such a monotonic dependence of the stability constants on the solvent mixture composition. From Table 2 it is seen that the stability of the resulting 1 : 1 complexes of Hg 2+ ion with different macrocyclic diamides, in 90% actonitrile mixture, decreases in the order L 8 > L 7 > L 2 > L I > L 4 > L 3 = L 5 > L 6 . There are at least five factors which can make significant contributions to the stability of the metal ion complexes with macrocyclic ligands: (1) the cavity size-cation diameter ratio, (2) the number and nature of donating atoms participating in cation binding, (3) the number and nature of substituents on the macrocyclic ring, (4) the conformations of the free and complexed macrocyclic ligands and (5) the extent of solvation of the species involved in the complexation reaction. As is seen from Fig. 1, the total number of ring atoms in the macrocyclic diamides used is 17 (for ligands L1, L2, L4 and L8), 18 (for L6), 19 (for L5), 20 (for L7) and 21 (for L3). The Hg 2÷ ion with an ionic radius of 1.19 ~, [41] seems to have the best fitting condition inside the 17crowns used, resulting in the most stable complexes in the series. However, the ligand L7, with 20 atoms in the ring, exceptionally forms the second most stable complex with Hg 2÷ ion in the series. This is most probably due to the existence of three soft - N H - groups in its ring (all other ligands have only two - N H - groups in their structures) which promotes the tendency of the soft cation for the interaction with this ligand. Comparison of the data given in Table 2 indicates that among different 17-crown diamides used, where the ring

frame remains the same, the stabilities of the resulting mercury complexes vary in the order L8 > L2 > L1 > L4. The presence of an extra - C H 2 - C H 2- group on the ring of the ligand L1 can pump electrons into the ligand ring and thus increase the basicity of the donating nitrogen atoms of the ring, while the flexibility of the ligand remains more or less the same as L1. Thus it is not unexpected to observe the highest stability for HgE+-L8 complex. The presence of a - C H 3 group on the ring of L1 can also perform the same function, but to a lesser extent and, therefore, the HgE+-L2 complex is somewhat more stable than Hg2+-L1. However, the addition of another benzo group to the ring of L1 lowers the stability of the resulting mercury complex markedly. This behavior, observed for the H g 2 ÷ - L 4 complex, may be attributed to some combination of the electron withdrawing property of the benzo group, which weakens the electron-donor ability of the nitrogen atoms of the ring, and increased rigidity of the ligand which prevents the macrocyclic molecule wrapping itself around the cation. Among the three remaining ligands, i.e. L3, L5 and L6, L3 forms the most stable complex, although its cavity size is too large for the cation. This behavior seems to be due to the electron donating property of the - ( C H 2 ) 6 - bridge, which can increase the basicity of the nitrogen atoms of the ring, as well as the resulting increased flexibility of the macrocyclic ring of L3. In comparison, lower stability of the ligands L5 and, especially, L6 would be a consequence of the electron withdrawing effect of - C 6 H 4 - and - C i o H 6- groups respectively, as well as the very high degree of rigidity of the macrocyclic diamides. However, as the data given in Table 2 indicate, an increase in the fraction of water in the acetonitrile + water mixtures would result in some considerable changes in the stability order of Hg 2+ complexes with the macrocyclic diamides used. Such behavior can most probably be related to the variations in the specific interactions between the solvent components and macrocyclic ligands [42-44]. Unfortunately, the data about such solvent-ligand interactions are too meager to allow useful speculations.

M.R. Ganjali et al. /Journal of E lectroanalytical Chemistry 405 (1996) 177-181

References [1] K. Naemura, H. Iwasaka and H. Chikamatsu, Bull. Chem. Soc. Jpn., 60 (1987) 4181. [2] D. Baldwin, P.A. Buckworth, G.R. Erickson, L.F. Lindoy, M. McPartlin, G.M. Mockler, W.E. Moody and P.A. Tasker, Aust. J. Chem., 40 (1987) 1861. [3] K.E. Krakowiak and J.S. Bradshaw, lsr. J. Chem., 32 (1992) 3. [4] R.M. Izatt, G.A. Clark, J.S. Bradshaw, J.D. Lamb and J.J, Christensen, Sep. Purif. Meth., 15 (1986) 21. [5] L.F. Lindoy and P.S. Baldwin, Pure Appl. Chem., 61 (1989) 909. [6] M. Dozol, in L. Cecille, M. Casaraci and L. Pietrelli (Eds.), New Separation Chemistry Techniques for Radioactive Waste and Other Specific Applications, Elsevier, Amsterdam, 1991. [7] J.S. Shih, J. Chin. Chem. Soc., 39 (1992) 551. [8] H. Sharghi and H. Eshghi, Tetrahedron, 51 (1995) 1993. [9] Y.A. lbrahim and A.H.M. Elwahy, Synthesis, (1993) 503. [10 R.M. Izatt, K. Pawlak, J.S. Bradshaw and R.L. Bmening, Chem. Rev., 91 (1991) 1721. [11 F. Peter, M. Gross, L. Pospisil and J. Kuta, J. Eiectroanal. Chem., 90 (1978) 239. [12 A. Hofmanova, J. Koryta, M. Brezina and L. Mittal, Inorg. Chim. Acta, 28 (1978) 73. [13 E.L. Yee, J. Tabib and M.J. Weaver, J. Electroanal. Chem., 96 (1979) 241. [14 C. Boudon, F. Peter and M. Gross, J. Electroanal Chem., 135 (1982) 93. [15 C. Luca, H.A. Azab and I. Tanase, Anal. Lett., 18 (1985) 449. [16 E. Lada, S. Filipek and M.K. Kalinowski, Aust. J. Chem., 41 (1988) 437. [17] Z. Samec and P. Papoff, Anal. Chem., 62 (1990) 1010. [18] H. Parham and M. Shamsipur, J. Electroanal. Chem., 314(1991) 71. [19] A. Semnani and M. Shamsipur, J. Electroanal. Chem., 315 (1991) 95. [20] A. Rouhollahi, M. Shamsipur and M.K. Amini, Talanta, 41 (1994) 1465. [21] M.R. Ganjali and M. Shamsipur, J. lnclus. Phenom., in press. [22] N. Alizadeh and M. Shamsipur, Talanta, 40 (1993) 503.

181

[23] A. Rouhollahi, M.K. Amini and M. Shamsipur, J. Solut. Chem., 23 (1994) 63. [24] M. Shamsipur and J. Ghasemi, J. Inclus. Phenom., 20 (1995) 157. [25] D.R. Crow, Polarography of Metal Complexes, Academic Press, New York, 1969. [26] I.M. Kolthoff and C.S. Miller, J. Am. Chem. Soc., 63 (1941) 1405. [27] J. Heyrovky and J. Kuta, Principles of Polarography, Academic Press, New York, 1969. [28] E.P. Parry and R.A. Osteryoung, Anal. Chem., 37 (1965) 1634. [29] A.J. Bard and L.R. Eaulkner, Electrochemical Methods, Fundamentals and Applications, Wiley, New York, 1980, p. 194. [30] C.J. Nyman and E.P. Parry, Anal. Chem., 30 (1958) 1255. [31] P.K. Wrona and Z. Galus, in A.J. Bard (Ed.), Encyclopedia of Electrochemistry of the Elements, Vol. IX, Marcel Dekker, New York, 1982. [32] M. Hojo, M. Hajiwara, H. Nagai and Y. lmai, J. Electroanal. Chem., 234 (1987) 251. [33] V. Gutman, Coordination Chemistry in Nonaqueous Solutions, Springer, New York, 1960. [34] R.H. Erlish and A.I. Popov, J. Am. Chem. Soc., 93 (1971) 5620. [35] R.M. Izatt, R.E. Terry, D.P. Nelson, Y. Chan, D.J. Eatough, J.S. Bradshaw, L.D. Hansen and J.J. Christensen, J. Am. Chem. Soc., 98 (1976) 7626. [36] B.G. Cox, G. Gnminski and H. Schneider, J. Am. Chem. Soc., 104 (1982) 3789. [37] B.G. Cox, P. Fiman, D. Gudin and H. Schneider, J. Phys. Chem., 86 (1982) 4986. [38] M.B. Gholivand and M. Shamsipur, Inorg. Chim. Acta, 121 (1986) 53. [39] M. Sharnsipur, A. Esmaeili and M.K. Amini, Talanta, 36 (1989) 1300. [40] M. Saeidi and M. Shamsipur, J. Coord. Chem., 22 (1991) 131. [41] R.D. Shannon, Acta Crystallogr., 32A (1976) 751. [42] J.A.A. DeBoer, D.M. Reinhoudt, S. Harkema, GJ. van Hummel and F. de,long, J. Am. Chem. Soc., 104 (1982) 4073. [43] P.A. Mosier-Boss and A.J. Popov, J. Am. Chem. Soc., 107 (1985) 6168. [44] R.M. Izatt, J.S. Bradshaw, K. Pawlak, R.L. Bruening and BJ. Tarbet, J. Chem. Rev., 92 (1992) 1261.