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Spectrophotometric study of the alkali metal—murexide complexes in some non-aqueous solutions
Talanta,Vol. 36, No. 7, pp. 773-716, 1989 Printed in GreatBritain. All rights reserved
0039-9140/89 53.00+ 0.00 Copyright Q 1989Pcrgamon Prm plc
SPECTROPHOTOMETRIC STUDY OF THE ALKALI METAL-MUREXIDE COMPLEXES IN SOME NON-AQUEOUS SOLUTIONS MOJTABA SHAMSIPLJR+
Department
of Chemistry, Shiraz University, Shirax, Iran
SIAVASH MADAENI
and S~I-IEILAKASELWAN
Department of Chemistry, Raxi University, Bakhtaran, Iran (Received 9 July 1988. Revtked 1 January 1989. Accepted 15 January 1989) Summary-The complexes of murexide with alkali-metal cations have been studied spectrophotometrically in methanol, dimethylformamide and dimethylsulphoxide media at 25”. The stoichiometry of the complexes was found to be I : 1. The formation constants of the complexes were determined, and found to decrease in the order Na+ > K+ > Rb+ m Li+ for all solvents studied. The complex formation constants varied inversely with the Gutmann donicity of the solvents.
chemistry of the alkali-metal ions was very largely unknown. The discovery of crown ethers’ and cryptands’ opened a new era in co-ordination of metal ions, and during the past 15 years thermodynamic and kinetic studies of the complexation reactions between these ligands and alkali-metal ions have become a very important field of research.3 In the past, however, the studies of alkali-metal complexes with more conventional lig ands have been much less popular, mostly due to the weak interactions between these ligands and cations (especially in aqueous solutions), which are often undetectable by most physicochemical techniques.’
ExPERIMBNTAL
Two decades ago the co-ordination
Reagent grade murexide (Merck) and the bromides of lithium (BDH), sodium (Merck), potassium @I & B) and rubidium (Merck) were used without further treatment except drying. The methanol (Baker, MeOH), dimethylformamide (Fisher. DMF) and dimethvlsulnhoxide (Fisher. DMSO) used were purified as repo&d~elsewhere.‘O All spectra were obtained with a Beckman 34 spectrometer at 25 f 2”. The formation constants of the 1: 1 complexes __
were determined by measuring the absorbance changes in the complex formation reactions. The concentration of the ligand was kept constant at 2.0 x lO_rM and the concentration of the salts was varied from 1.0 x lo-’ to 1.0 x lo-*M. 4 was determined from a linear plot of l/(aA-eL) vs. l/C,+,” since 1 -=CA-CL
L
I
-!
Murexide, the ammonium salt of purpuric acid (I), is well known as a metallochromic indicator for determinatiot?” and dynamic studies’ of alkalineearth metal cations in aqueous solution. The use of murexide in spectrophotometric studies of the kinetics of alkali-metal complexation by a variety of ligand molecules in methanol solution has also been reported.* We have recently reported the results of a study of alkaline-earth metal complexes with murexide in some non-aqueous solvents,9 and have now extended the study to the alkali-metal complexes formed with murexide in methanol, dimethylformamide and dimethylsulphoxide solutions. *Author to whom correspondence
should be addressed.
[ML1
1 (ML -
EL ( l+x;c,
1 (1) >
where EA = A/C,, A is the absorbance of the solution, CL is the initial concentration of mumxide, C,,, is the total concentration of alkali-metal ion and eL and tlrlL am the molar absorptivities of the ligand and complex, respectively.
RESULTS AND DISCUSSION
The spectra of murexide and its complexes with Li+, Na+, K+ and Rb+ were obtained in MeOH, DMF and DMSO solutions. The spectra in DMSO are shown in Fig. 1. In all three solvents, the alkalimetal complexes are distinguished by a strong and ion-specific spectral shift towards shorter wavelengths. As we have noted in the case of the alkaline-earth metal murexide complexes,9 such a pronounced shift is possible only if the two rings of the murexide molecule are twisted around the central 773
MOJTABASELAMSIPUR et al.
774
Li*
I
I
I
I
I
I
450
500
550
600
650
WAVELENGTH
hm)
Fig. 1. Visible region spectra of murexide and its alkalimetal complexes in DMSO. nitrogen-bridge axis in the complexation. Such an assumption was proposed earlier.’ The stoichiometry of the complexes was determined by the continuous variations method,‘* and found to be 1: 1. The plot for the Na+ complex in DMSO is shown in Fig. 2. The existence of well defined isosbestic points in the spectra of murexide recorded during its titration with alkali-metal solution is also further evidence for a simple 1: 1 complexation equilibrium (Fig. 3). To determine the formation constants of the complexes, the spectra of solutions containing a constant amount of murexide and varying amounts of the alkali metal were obtained (c$ Fig. 3). The plots of l/(eA - Ed) us. l/C, gave straight lines, in accordance with equation (l), and the formation constants were calculated from the slopes and intercepts. All the values obtained are given in Table 1. The relationships between the stability constants and the crystal radii of the alkali-metal cations” are shown in Fig. 4. It is seen (Table 1) that in methanol solutions the stability of the Nat complex is not affected by a
WAVELEIJGTH
Fig. 3. Visible
MOLE FRACTION
Fig. 2. Continuous variation plot for Nat-murexide in DMSO. change in the counter-ion from chloride to iodide. It is evident, therefore, that in the relatively high dielectric constant16 solvents we used and at the low salt concentrations studied, the formation of the complexes is unaffected by ion-pairing. Comparison of our values with those reported from kinetic data for the complexation of Lit, Na+ and K+ ions by murexide in methanol, measured by the spectrophotometric electric field-jump relaxation method,’ shows a satisfactory agreement. As shown in Fig. 1, the spectral behaviour of the alkali-metal complexes of murexide consists of strong cation-specific shifts towards shorter wavelengths. Whereas, in all solvents used, the displacement and intensity of the absorption band of the complex both increase with decreasing radius of the alkali-metal ion, the formation constants of the complexes vary in the order Nat > K+ > Rb+ _ Lit. The same kind of
(nm)
region spectra for titration of 2.0 x 10v5M murexide with Li+ in DMSO at 25”. [LPI: 1, 1.0 x 10-2M; 2, 8.0 x lo-“M; 3, 6.0 x lo-‘M; 4, 4.0 x lo-‘M; 5, 2.0 x lo-‘M; 6, 1.0 x lo-‘M; 7, murexide alone.
Spectrophotometric
study of the alkali metal-murexide
complexes
775
Table 1. Log X, of different alkali-metal cation complexes with murexide in various solvents at 25”, with chloride as counter-ion Solvent
Dielectric constant’6
Gutma~ donor numberI
Methanol
32.7
19.7
DMF
36.7
26.6
DMSO
46.7
29.8
Cation L.i+ Li+ Na+ Na+ Na+ ::
log x,
Rb+ Li+ Na+ K+ Rb+ Li+ Na+ K+
2.85 f 0.07 2.9* 3.38 f 0.06 3.43 f 0.087 3.4. 3.08 f 0.08 3.1’ 2.95 f 0.06 2.28 f 0.07 2.90 f 0.06 2.50 f 0.06 2.37 rt 0.05 1.99 f 0.07 2.58 + 0.07 2.18 &0.06
Rb+
2.02 f 0.05
*From Diebler et ~1.~ tWith iodide as counter-ion.
behaviour was observed for the alkaline-earth metal murexide complexes,9 where Ca*+ (which has about the same ionic size as Na+) forms the most stable of the alkaline-earth metal complexes. Whilst any simple model must correspond to the monotonic decrease in the stability with increasing radius of the cation, the question may arise of why the complexes of metal ions of a particular size have the highest stability constant. Since in the case of the alkali metals and alkaline-earth metals we are dealing with ions with “noble gas” electron coniigurations, the selectivity of complex formation cannot be a consequence of rearrangement of the electron configuration as in the case of transition metal ions. Therefore, it seems more likely to be a special
I
I
I
I
0.8
1.0
1.2
1.4
IONIC RADIUS th Fig. 4. Stability constants of the alkali-metal complexes of murexide, in various solvents at 25” us. ionic radii of the cations. TAL %/l-E
property of the complexing ligand and the reaction medium. First, it should be noted that the thermodynamic stability constant is not just a measure of the solutesolute interaction, but is a measure of the relative strength of this in comparison with that of the solute4olvent interactions. Thus for a given group of metal ions the stability constant will be afIected mainly by the differences in the cation-ligand binding strengths and the solvation energies of the metal ions and their complexes. Another factor which could affect the stability constant is the conformational geometry of the ligand in solution. As mentioned earlier,s” murexide has a relatively flexible structure in solution, in which the two rings of the dye molecule can twist relative to each other around the central nitrogen-bridge axis, so that the donor atoms (bridg ing nitrogen atom and neighbouring oxygen atoms) can form a variable geometry. Clearly the highest binding energy would be associated with a particular cation size favouring a suitable spatial fit. Cations with smaller or larger radius would fail to achieve the maximum stability. The reported crystalline structures for the lithium” and potassium’s complexes of murexide support this discussion. In both cases, the two approximately planar barbiturate rings of each anion (I) are. not co-planar, the interaction between neighbouring carbonyl groups causing torsion about the central nitrogen+arbon bonds. It is obvious that in solution the ligand, with much more flexibility, has a better opportunity to achieve the most convenient structure. From Table 1 it can be seen that for a given cation the stabilities of the complexes are very dependent on the nature of the solvent, but with all the solvents tested there is an inverse relationship between the stabilities of the complexes and the donicity of the solvents, as expressed by the Gutmann donor numbers.i6 Methanol is the solvent with lowest donicity
MOJTABASIU~~SIPURer al.
176
and, therefore, is the least competitive with the ligand for binding the cations, which in turn results in higher stability for all the murexide complexes in a given metal series. There are several earlier articles which clearly show the same type of the solvent effect on the stabilities of various alkali-metal and alkaline-earth metal complexes.9*‘7-20
8. H. Diebler, M. Eigen, G. Ilgenfritz, G. Maass and
10. 11. 12. 13. 14.
REFERENCES
1. C. J. Pedersen, J. Am. Chem. Sot., 1967, 89, 7017. 2. B. Dietrich, J. M. Lehn and J. P. Sauvage, Tetrahedron Lea, 1959, 2885. 3. R. M. Izatt. J. S. Bradshaw. S. A. Nielsen. J. D. Lamb and J. J. Christensen, Chek Rev., 1985, 85, 271. 4. N. S. Poonia and A. V. Bajaj, ibid., 1979, 79, 389. 5. A. Scarna. Methodr Enzvmol.. 1972, 24B. 343. 6. D. S. Russell, J. B. Carnpbell~and S: S. Berman, Anal. Chim. Acta, 1961, 25, 81. 7. V. M. Loyola, R. Pizer and R. G. Wilkins, J. Am. Chem. Sot., 1977, 99, 7185.
R. Winkler. Pure ADDS. Chem.. 1969. 20. 93. M.’ ‘B. Gholivand, S. -Madaeni, A. Nikrahi and M. Shamsipur, Polyhedron, 1988,7, 1227. M. S. Greenberg and A. I. Popov, Spectrochim. Acra, 1975, 31A, 697. J. P. Birk, P. B. Chock and J. Halpern, J. Am. Chem. SOL, 1968, 90, 6959. W. Likussar and D. F. Boltz, Anal. Chem., 1971, 43, 1262. J. M. Lehn, Sfrnct. Bonding (Berlin), 1973, 16, 1. H. B. Burgi, S. Diuric, M. Dobler and J. D. Dunitz. Helv. Ch&Acta,w197i, 56, 1771. R. L. Martin. A. H. White and A. C. Willis. J. Chem. Sot. Dalton i%ans., 1977, 1336. V. Gutmann, Coordination Chemistry in Nonaqueous Solvents,Springer-Verlag. Vienna, 1968. E. Schmidt, A. Hourdakis and A. I. Popov, Inorg. Chim. Acta, 1981, 52, 91. M. Shainsiour. G. Rounaahi J. Soln. - and A. 1. Ponov. _ Chem., 1980, 9, 701. M. B. Gholivand and M. Shamsipur, Inorg. Chim. Acta, 1986, 121, 53. M. B. Gholivand, S. Kashanian and M. Shamsipur, Polyhedron, 1987, 6, 535.
9. S. Kashanian,
15. 16. 17. 18. 19. 20.
0039-9140/89 53.00+ 0.00 Copyright Q 1989Pcrgamon Prm plc
SPECTROPHOTOMETRIC STUDY OF THE ALKALI METAL-MUREXIDE COMPLEXES IN SOME NON-AQUEOUS SOLUTIONS MOJTABA SHAMSIPLJR+
Department
of Chemistry, Shiraz University, Shirax, Iran
SIAVASH MADAENI
and S~I-IEILAKASELWAN
Department of Chemistry, Raxi University, Bakhtaran, Iran (Received 9 July 1988. Revtked 1 January 1989. Accepted 15 January 1989) Summary-The complexes of murexide with alkali-metal cations have been studied spectrophotometrically in methanol, dimethylformamide and dimethylsulphoxide media at 25”. The stoichiometry of the complexes was found to be I : 1. The formation constants of the complexes were determined, and found to decrease in the order Na+ > K+ > Rb+ m Li+ for all solvents studied. The complex formation constants varied inversely with the Gutmann donicity of the solvents.
chemistry of the alkali-metal ions was very largely unknown. The discovery of crown ethers’ and cryptands’ opened a new era in co-ordination of metal ions, and during the past 15 years thermodynamic and kinetic studies of the complexation reactions between these ligands and alkali-metal ions have become a very important field of research.3 In the past, however, the studies of alkali-metal complexes with more conventional lig ands have been much less popular, mostly due to the weak interactions between these ligands and cations (especially in aqueous solutions), which are often undetectable by most physicochemical techniques.’
ExPERIMBNTAL
Two decades ago the co-ordination
Reagent grade murexide (Merck) and the bromides of lithium (BDH), sodium (Merck), potassium @I & B) and rubidium (Merck) were used without further treatment except drying. The methanol (Baker, MeOH), dimethylformamide (Fisher. DMF) and dimethvlsulnhoxide (Fisher. DMSO) used were purified as repo&d~elsewhere.‘O All spectra were obtained with a Beckman 34 spectrometer at 25 f 2”. The formation constants of the 1: 1 complexes __
were determined by measuring the absorbance changes in the complex formation reactions. The concentration of the ligand was kept constant at 2.0 x lO_rM and the concentration of the salts was varied from 1.0 x lo-’ to 1.0 x lo-*M. 4 was determined from a linear plot of l/(aA-eL) vs. l/C,+,” since 1 -=CA-CL
L
I
-!
Murexide, the ammonium salt of purpuric acid (I), is well known as a metallochromic indicator for determinatiot?” and dynamic studies’ of alkalineearth metal cations in aqueous solution. The use of murexide in spectrophotometric studies of the kinetics of alkali-metal complexation by a variety of ligand molecules in methanol solution has also been reported.* We have recently reported the results of a study of alkaline-earth metal complexes with murexide in some non-aqueous solvents,9 and have now extended the study to the alkali-metal complexes formed with murexide in methanol, dimethylformamide and dimethylsulphoxide solutions. *Author to whom correspondence
should be addressed.
[ML1
1 (ML -
EL ( l+x;c,
1 (1) >
where EA = A/C,, A is the absorbance of the solution, CL is the initial concentration of mumxide, C,,, is the total concentration of alkali-metal ion and eL and tlrlL am the molar absorptivities of the ligand and complex, respectively.
RESULTS AND DISCUSSION
The spectra of murexide and its complexes with Li+, Na+, K+ and Rb+ were obtained in MeOH, DMF and DMSO solutions. The spectra in DMSO are shown in Fig. 1. In all three solvents, the alkalimetal complexes are distinguished by a strong and ion-specific spectral shift towards shorter wavelengths. As we have noted in the case of the alkaline-earth metal murexide complexes,9 such a pronounced shift is possible only if the two rings of the murexide molecule are twisted around the central 773
MOJTABASELAMSIPUR et al.
774
Li*
I
I
I
I
I
I
450
500
550
600
650
WAVELENGTH
hm)
Fig. 1. Visible region spectra of murexide and its alkalimetal complexes in DMSO. nitrogen-bridge axis in the complexation. Such an assumption was proposed earlier.’ The stoichiometry of the complexes was determined by the continuous variations method,‘* and found to be 1: 1. The plot for the Na+ complex in DMSO is shown in Fig. 2. The existence of well defined isosbestic points in the spectra of murexide recorded during its titration with alkali-metal solution is also further evidence for a simple 1: 1 complexation equilibrium (Fig. 3). To determine the formation constants of the complexes, the spectra of solutions containing a constant amount of murexide and varying amounts of the alkali metal were obtained (c$ Fig. 3). The plots of l/(eA - Ed) us. l/C, gave straight lines, in accordance with equation (l), and the formation constants were calculated from the slopes and intercepts. All the values obtained are given in Table 1. The relationships between the stability constants and the crystal radii of the alkali-metal cations” are shown in Fig. 4. It is seen (Table 1) that in methanol solutions the stability of the Nat complex is not affected by a
WAVELEIJGTH
Fig. 3. Visible
MOLE FRACTION
Fig. 2. Continuous variation plot for Nat-murexide in DMSO. change in the counter-ion from chloride to iodide. It is evident, therefore, that in the relatively high dielectric constant16 solvents we used and at the low salt concentrations studied, the formation of the complexes is unaffected by ion-pairing. Comparison of our values with those reported from kinetic data for the complexation of Lit, Na+ and K+ ions by murexide in methanol, measured by the spectrophotometric electric field-jump relaxation method,’ shows a satisfactory agreement. As shown in Fig. 1, the spectral behaviour of the alkali-metal complexes of murexide consists of strong cation-specific shifts towards shorter wavelengths. Whereas, in all solvents used, the displacement and intensity of the absorption band of the complex both increase with decreasing radius of the alkali-metal ion, the formation constants of the complexes vary in the order Nat > K+ > Rb+ _ Lit. The same kind of
(nm)
region spectra for titration of 2.0 x 10v5M murexide with Li+ in DMSO at 25”. [LPI: 1, 1.0 x 10-2M; 2, 8.0 x lo-“M; 3, 6.0 x lo-‘M; 4, 4.0 x lo-‘M; 5, 2.0 x lo-‘M; 6, 1.0 x lo-‘M; 7, murexide alone.
Spectrophotometric
study of the alkali metal-murexide
complexes
775
Table 1. Log X, of different alkali-metal cation complexes with murexide in various solvents at 25”, with chloride as counter-ion Solvent
Dielectric constant’6
Gutma~ donor numberI
Methanol
32.7
19.7
DMF
36.7
26.6
DMSO
46.7
29.8
Cation L.i+ Li+ Na+ Na+ Na+ ::
log x,
Rb+ Li+ Na+ K+ Rb+ Li+ Na+ K+
2.85 f 0.07 2.9* 3.38 f 0.06 3.43 f 0.087 3.4. 3.08 f 0.08 3.1’ 2.95 f 0.06 2.28 f 0.07 2.90 f 0.06 2.50 f 0.06 2.37 rt 0.05 1.99 f 0.07 2.58 + 0.07 2.18 &0.06
Rb+
2.02 f 0.05
*From Diebler et ~1.~ tWith iodide as counter-ion.
behaviour was observed for the alkaline-earth metal murexide complexes,9 where Ca*+ (which has about the same ionic size as Na+) forms the most stable of the alkaline-earth metal complexes. Whilst any simple model must correspond to the monotonic decrease in the stability with increasing radius of the cation, the question may arise of why the complexes of metal ions of a particular size have the highest stability constant. Since in the case of the alkali metals and alkaline-earth metals we are dealing with ions with “noble gas” electron coniigurations, the selectivity of complex formation cannot be a consequence of rearrangement of the electron configuration as in the case of transition metal ions. Therefore, it seems more likely to be a special
I
I
I
I
0.8
1.0
1.2
1.4
IONIC RADIUS th Fig. 4. Stability constants of the alkali-metal complexes of murexide, in various solvents at 25” us. ionic radii of the cations. TAL %/l-E
property of the complexing ligand and the reaction medium. First, it should be noted that the thermodynamic stability constant is not just a measure of the solutesolute interaction, but is a measure of the relative strength of this in comparison with that of the solute4olvent interactions. Thus for a given group of metal ions the stability constant will be afIected mainly by the differences in the cation-ligand binding strengths and the solvation energies of the metal ions and their complexes. Another factor which could affect the stability constant is the conformational geometry of the ligand in solution. As mentioned earlier,s” murexide has a relatively flexible structure in solution, in which the two rings of the dye molecule can twist relative to each other around the central nitrogen-bridge axis, so that the donor atoms (bridg ing nitrogen atom and neighbouring oxygen atoms) can form a variable geometry. Clearly the highest binding energy would be associated with a particular cation size favouring a suitable spatial fit. Cations with smaller or larger radius would fail to achieve the maximum stability. The reported crystalline structures for the lithium” and potassium’s complexes of murexide support this discussion. In both cases, the two approximately planar barbiturate rings of each anion (I) are. not co-planar, the interaction between neighbouring carbonyl groups causing torsion about the central nitrogen+arbon bonds. It is obvious that in solution the ligand, with much more flexibility, has a better opportunity to achieve the most convenient structure. From Table 1 it can be seen that for a given cation the stabilities of the complexes are very dependent on the nature of the solvent, but with all the solvents tested there is an inverse relationship between the stabilities of the complexes and the donicity of the solvents, as expressed by the Gutmann donor numbers.i6 Methanol is the solvent with lowest donicity
MOJTABASIU~~SIPURer al.
176
and, therefore, is the least competitive with the ligand for binding the cations, which in turn results in higher stability for all the murexide complexes in a given metal series. There are several earlier articles which clearly show the same type of the solvent effect on the stabilities of various alkali-metal and alkaline-earth metal complexes.9*‘7-20
8. H. Diebler, M. Eigen, G. Ilgenfritz, G. Maass and
10. 11. 12. 13. 14.
REFERENCES
1. C. J. Pedersen, J. Am. Chem. Sot., 1967, 89, 7017. 2. B. Dietrich, J. M. Lehn and J. P. Sauvage, Tetrahedron Lea, 1959, 2885. 3. R. M. Izatt. J. S. Bradshaw. S. A. Nielsen. J. D. Lamb and J. J. Christensen, Chek Rev., 1985, 85, 271. 4. N. S. Poonia and A. V. Bajaj, ibid., 1979, 79, 389. 5. A. Scarna. Methodr Enzvmol.. 1972, 24B. 343. 6. D. S. Russell, J. B. Carnpbell~and S: S. Berman, Anal. Chim. Acta, 1961, 25, 81. 7. V. M. Loyola, R. Pizer and R. G. Wilkins, J. Am. Chem. Sot., 1977, 99, 7185.
R. Winkler. Pure ADDS. Chem.. 1969. 20. 93. M.’ ‘B. Gholivand, S. -Madaeni, A. Nikrahi and M. Shamsipur, Polyhedron, 1988,7, 1227. M. S. Greenberg and A. I. Popov, Spectrochim. Acra, 1975, 31A, 697. J. P. Birk, P. B. Chock and J. Halpern, J. Am. Chem. SOL, 1968, 90, 6959. W. Likussar and D. F. Boltz, Anal. Chem., 1971, 43, 1262. J. M. Lehn, Sfrnct. Bonding (Berlin), 1973, 16, 1. H. B. Burgi, S. Diuric, M. Dobler and J. D. Dunitz. Helv. Ch&Acta,w197i, 56, 1771. R. L. Martin. A. H. White and A. C. Willis. J. Chem. Sot. Dalton i%ans., 1977, 1336. V. Gutmann, Coordination Chemistry in Nonaqueous Solvents,Springer-Verlag. Vienna, 1968. E. Schmidt, A. Hourdakis and A. I. Popov, Inorg. Chim. Acta, 1981, 52, 91. M. Shainsiour. G. Rounaahi J. Soln. - and A. 1. Ponov. _ Chem., 1980, 9, 701. M. B. Gholivand and M. Shamsipur, Inorg. Chim. Acta, 1986, 121, 53. M. B. Gholivand, S. Kashanian and M. Shamsipur, Polyhedron, 1987, 6, 535.
9. S. Kashanian,
15. 16. 17. 18. 19. 20.