4176
J. Phys. Chem. 1988, 92, 4176-4180
Stabilities and Selectivities of Alkali-Metal Complexes of Tetraethylene Glycol Dimethyl Ether in Methanol and Acetonitrilet E. Buncel,* H. S. Shin, Ng. van Truong, Department of Chemistry, QueenS University, Kingston, Ontario, Canada K7L 3N6
R. A. B. Bannard, and J. G. Purdon Chemical Detection and Decontamination Section, Protective Sciences Division, Defence Research Establishment Ottawa, Shirley Bay, Ottawa, Ontario, Canada K1A 024 (Received: December 1.5, 1987)
Thermodynamic data (log K,,AHc, and TAS,)for alkali-metal-ion complex formation by tetraethylene glycol dimethyl ether (tetraglyme, TG) have been determined in methanol and acetonitrile as solvents at 25 O C by using a modified calorimetric technique. It has been shown that the stabilities and selectivitiesfor complexation by tetraglyme are lower than by macrocyclic polyethers. In methanol the selectivity order for tetraglyme is K+ > Rb+ > Cs+ > Na+ > Li+, but in acetonitrile the order is Na+ > Li+ > K+ > Rb' > Cs'. The stability constants, log Ks,of the TG-M+ complexes are higher in acetonitrile than in methanol. A linear relationship is found between log K,and Gutmann's donor number (DN) of the solvent. The sensitivity of stability constants to solvent variation is discussed in terms of free energies of transfer of the species involved. The plot of AG,(ML+) - AG,(L) vs AG,(M+) shows a much stronger dependenceon the nature of the cation for the TG-M+ complexes as compared to the corresponding [2.2.1] and [2.2.2] cryptand complexes, which indicates that solvent interaction with the cations of the TG-M+ complexes is still significant. Hydrogen-bonding interactions in methanol are identified as an important factor in decreasing the stability constants of the TG-M+ complexes relative to acetonitrile. The effect of anion on the stability constants of tetraglyme complexes is small except for KOPh in acetonitrile where the smaller K, value may be due to ion pairing. The stability of the TG-M+ complexes is enthalpy dominated, accompanied by an unfavorable entropy change in both methanol and acetonitrile, except for Li+ where the entropy term contributes to stability in both solvents. The macrocyclic effect is interpreted in terms of entropy changes.
The stability constants of metal-ion complexes have been Introduction measured by different methods,14 among which the calorimetric Due to their ability to form strong complexes with a variety technique showed some advantages, particularly for relatively weak of metal cations in organic solvents, the synthetic crown complexes where other methods fail or yield poor In and cryptand$ have found increasing use in various areas of chemistry, including transport through membranes, phase-transfer (1) (a) Pedersen, C. J. J . Am. Chem. SOC.1967, 89, 2495, 7017. (b) catalysis, and influencing ion pairing and reactivities of nucleoPedersen, C. J.; Frensdorf, H. F. Angew Chem., I n t . Ed. Engl. 1972, 11, 16. p h i l e ~as , ~well ~ ~ as in extraction studies6s7and cognate fields of (2) (a) Cram, D. J.; Cram, J. M. Science 1974, 183, 803; Arc. Chem. Res. chemistry.'-3 The open-chain polyethylene glycols (PEGs) and 1978, 11, 8. (b) Cram, D. J. In Applications of Biochemical Systems in Organic Chemistry; Jones, J. B., Sih, C. J., Perlman, D., Eds.; Wiley-Intheir ether derivatives have also been found to be capable of terscience: New York, 1976; Chapter 5. (c) Cram, D. J.; Trueblood, K. N. forming complexes with different metal cations in organic solvents, Top. Curr. Chem. 1981, 98, 43. in a similar manner to the macrocyclic polyethers.8-11 Like the (3) (a) Lehn, J. M. Srruct. Bonding 1973, 16, 1. (b) Lehn, J. M. Acc. synthetic crown ethers and cryptands, the PEGs and their ether Chem. Res. 1978, 11, 49; Pure Appl. Chem. 1978, 50, 871; 1978, 51, 979; 1980, 52, 244 1. derivatives have also been used as effective catalysts in phase(4) (a) Buncel, E.; Menon, B. C. J . Urg. Chem. 1979,44, 317. (b) Buncel, transfer catalysis,I2 while the lower molecular weight homologues E.; Menon, B. C.; Colpa, J. P. Can. J. Chem. 1979, 57, 999. can be used as effective solvents for many synthetic organic re(5) (a) Buncel, E.; Menon, B. C. J . A m . Chem. Soc. 1977,99,4457. (b) a c t i o n ~ .Thus ~ ~ the PEGs and their derivatives are of increasing Buncel, E.; Dunn, E. J.; Bannard, R. A. B.; Purdon, J. G. J . Chem. SOC., Chem. Commun. 1984, 162. interest in many areas, because of their useful properties as ef(6) (a) Buncel, E.; Shin, H. S.;Bannard, R. A. B.; Purdon, J. G. Can. J . fective complexing agents as well as their relative inexpensiveness Chem. 1984,42,926. (b) Buncel, E.; Shin, H. S.; Bannard, R. A. B.; Purdon, and nontoxicity. J. G.; Cox, B. G. Talanra 1984, 31, 585. (c) Cox, B. G.; Buncel, E.; Shin, While the metal ion complexes of the synthetic crown ethers H. S.;Bannard, R. A. B.; Purdon, J. G. Can. J. Chem. 1986, 64, 920. (7) Takeda, Y. Top. Curr. Chem. 1984, 121, 1. and cryptands have been extensively studied in recent years with respect to structure, stability constants and rate p r o c e ~ s e s , ' - ~ ~ ' ~ (8) Tummler, B.; Maass, G.; VBgtle, F.; Sieger, H.; Heimann, U.; Weber, E. J. Am. Chem. SOC.1979, 101, 2588. much less is known about the metal-ion complexes of the PEGS (9) (a) Toke, L.; Szabo,G. T. Acra. Chim. Acud. Sci. Hung. 1972,93,421. and their derivatives. The alkali-metal complexes of the PEGS (b) Toke, L.; Szabo, G. T.; Aranyosi, K. Ibid. 1979, 100, 257. (c) Ono, K.; Konami, K.; Murakami, K. J . Phys. Chem. 1979,83, 2665. and their derivatives have been studied mostly in alcoholic sol(10) Chaput, G.; Jeminet, G.; Jillard, J. Can. J . Chem. 1975, 53, 2240. vents8-I0 and no systematic study has dealt with solvent effects (11) (a) Weber, E.;VOgtle, F. Angew. Chem.;Int. Ed. Engl. 1978, 18,753. on the stabilities of metal-ion complexes of these ligands. The (b) Smid, J. Agnew. Chem., Int. Ed. Engl. 1972, 11, 112. stabilities of metal-ion complexes with different naturally occumng (12) (a) Gokel, G. W.; Goli, D. M.; Schultz, R. A. J . Urg. Chem. 1983, 48, 2837. (b) Neumann, R.; Sasson, Y. Ibid. 1984, 49, 3448. and synthetic polymers have been found to be very sensitive to (13) (a) Santaniello, E. In Crown Ethers and Phase Transfer Catalysis variation of in Polymer Science; Mathias, L. J., Carraher, Jr., C. E. Eds.; Plenum: New As part of our studies6 using complexing agents to transport York, 1984; p 397. (b) Santaniello, E.; Fiecchi, A,; Manzocchi, A,; Ferraions from water into nonaqueous organic solvents, we present a boschi, P. J . Urg. Chem. 1983, 48, 3074. (14) (a) Izatt, R. M.; Bradshaw, J. S.; Nielsen, S. A,; Lamb, J. D.; systematic study of the stability constants of alkali-metal-ion Christensen, J. J.; Sen, D. Chem. Rev. 1985, 85, 271 and references cited complexes of tetraethylene glycol dimethyl ether [CH30(CH2therein. (b) Burgermeister, W.; Winkler-Oswatitsch. R. Top. Curr. Chem. CH2-0)&H3, tetraglyme, TG] in methanol and acetonitrile as 1977, 69, 91. solvents. The effect of counteranion on the stabilities was also (15) Kolthoff, I. M.; Chantooni, Jr., M. K. Anal. Chem. 1980,52, 1039. (16) Cox, B. G.; Garcia-Rosas, J.; Schneider, H. J . Am. Chem. SOC.1981, determined. 'Part 6 in the series on Metal Ion Complexation by Macrocycles; part 5 , ref 6c.
0022-3654/88/2092-4176$01.50/0
103, 1054, 1384. (17) (a) Cox, B. G.; van Truong, Ng.; Schneider, H. J . Am. Chem. SOC. 1984, 106, 1273. (b) Cox, B. G.; van Truong, Ng.; Rzeszotarska, J.; Schneider, H. J. Chem. SOC.Faraday Trans. I 1984, 80, 3275.
0 1988 American Chemical Society
Alkali-Metal Complexes of Tetraglyme this work, a modified calorimetric technique was employed to determine directly the stability constants (log K,) and enthalpies (AHc)of alkali-metal-ion complexation by tetraglyme, as described in the Experimental Section.
Th!e Journal ofphysical Chemistry, Vol. 92, No. 14, 1988 4177 TABLE I: Stability Constants (log K , ) of Alkali-Metal Complexes with Tetraglyme (TG) and Other Ionophores in Methanol and Acetonitrile at 25 OC 1% Ks K+ methanol TG" 1.11 1.68 (1.28)* (1.72)b 15-C-5' 1.23 3.30 3.35 (3.24)d (3.43)d Enniatine Be 1.28 2.41 2.92 acetonitrile TG" 2.17 2.40 2.02 15-C-Sr 3.60 5.28 2.98 solvent
Experimental Section
Materials. Tetraglyme was obtained from Aldrich and distilled at least three times from sodium metal under vacuum. Methanol was purified by distillation from magnesium turnings and iodine. Acetonitrile was dried over molecular sieves overnight and distilled from calcium hydride under nitrogen atmosphere. The following reagents were used as purchased without further purification except for drying: lithium, sodium, and potassium perchlorate (Fisher), potassium thiocyanate (BDH, AnalaR), potassium iodide (Alfa, ultrapure grade), and rubidium and cesium perchlorate (Alfa). Potassium tetraphenylborate was prepared from sodium tetraphenylborate (Aldrich) and potassium chloride (Fisher), in aqueous solution. The white precipitate of the salt was recrystallized from acetonewater mixture and washed twice with water. Potassium picrate (KPic) was prepared from potassium hydroxide (Fisher) and picric acid (Aldrich) in ethanol. Potassium p nitrophenoxide (PNPhOK) was prepared from potassium hydroxide (Fisher) and p-nitrophenol (Aldrich) in methanol. The salts were recrystallized from the respective alcohols. Potassium phenoxide (PhOK) was prepared by mixing potassium ethoxide and phenol in ethanol. The white phenoxide salt, PhOK, was obtained on evaporation of the ethanol under a flow of nitrogen. All salts prepared in this work were dried under vacuum at room temperature and subsequently at 60-100 OC. Procedure and Calculations. The most widely used calorimetric method in determination of stability constants of cation-ligand complexes is the isoperbol titration method developed by Izatt and co-workers.18J0 However, Abraham2' has recently highlighted the use of a batch calorimetric technique to measure stability constants of metal-ion complexes. This method is based on one reported earlier by Drago.22 In our experiments the batch calorimetric technique was further modified as follows. The ampule containing 100 pL of tetraglyme (or 400 pL aliquot of ligand solution) was broken in the reaction vessel containing 50 mL of a solution of the salt in a given solvent. The total heat released (QT) was measured via a chart recorder and includes the heat of the complexation reaction (Qc),together with the heat of dilution of tetraglyme in the given solvent and the heat of breakage of the ampule The value of Q L can be determined independently by breaking the ampule containing the same volume of tetraglyme into 50 mL of the given solvent. Thus the heat of complexation, Qc, can be calculated from Qc = Q T - QL. The process was repeated by adding successive increments of tetraglyme (100 pL) via the ampules into the same batch of solution of given metal salt in a given solvent. The heat of complexation, Qcl, between the added tetraglyme and the remaining uncomplexed cation in the reaction vessel at the ith successive increment was then given by Qci= QTi - QL. The heat of dilution and breakage was determined several times with equal of the ampule successive portions of ligand (100 pL of tetraglyme), which were added in turn into the same batch of a given solvent contained in the reaction vessel. The values of Q L agreed within experimental error and the average value was calculated. According to the principle of the Born-Haber cycle, the sum of the Qcivalues at a stage that n increments of ligand have been
(e,).
(eL)
(18) (a) Christensen, J. J.; Ruchman, J.; Eatough, D. J.; Izatt, R. M. Thermochim. Acta 1972, 3, 203. (b) Eatough, D. J.; Christensen, J. J.; Izatt, R. M. Thermochim. Acra 1972, 3, 219, 233. (19) Lamb, J. D.; Izatt, R. M; Swain, C. S.; Christensen, J. J. J . A m . Chem. SOC.1980, 102, 415. (20) Davidson, R. B.; Izatt, R. M.; Christensen, J. J.; Schultz, R.; Dishong, D. M.; Gokel, G. M. J . Org. Chem. 1984, 49, 5080. (21) Abraham, M. H. In Thermochemistry and its Applications to Chemical and Biochemical Sysrems; Ribeiro da Silva, M.A.V., Ed.; Reidel: Dardrecht, The Netherlands, 1984; p 275. (22) (a) Rose, N. J.; Drago, R. S. J . Am. Chem. SOC.1959,81,6138. (b) Bolles, T. F.; Drago, R. S. J . Am. Chem. SOC.1965, 87, 5015.
ligand
Li+ 0.89
Na+
Rb+
Cs+
1.52
1.45 (1.45)b 2.62
2.74 1.92
2.34 1.53
'This work (log K , f 0.05), using alkali-metal perchlorate salts. *Reference 10. 'Reference 20, 14a. dReference 26. 'Reference 14b. 'Reference 28.
added must be equal to the Q, value resulting from a single addition of an equivalent total quantity of ligand, since the equilibrium condition should be the same for a given system regardless of the number of steps taken for it to be reached. Hence the method yields a set of data which is equivalent to that obtained by the batch calorimetric technique, in which separate experiments are repeated n times with different fresh batches of salt solution. The data were then treated by the Rose-Drago method,22 using a basic program provided by Dr. M. H. Abraham (University of Surrey, U.K.) which was modified to allow for an increase of total volume resulting from addition of ligand. This program gives the values of log K, and AHc through an iterative method. Therefore this technique could also be considered as a discontinuous calorimetric isoperbol titration. For a given complexation reaction, several sets of data are performed and average values of log K, and AHc were calculated. The ligand concentrations used were in the range of 7 X to 4.5 X M, and metal-ion concentrations in the range of 2 X to 1.2 X M . No correction was made for activity coefficients since the concentrations are rather low and the terms due to the ionic species appear in both the numerator and denominator in the expression for K,. The calorimetric measurements were performed using a Tronac Model 1250 or a Tronac Model 450 calorimeter. For acetonitrile, the solutions were prepared and transferred under nitrogen atmosphere. The reliability of log K , and AH, values obtained from this calorimetric method was checked by comparing with the results determined by other methods; generally good agreement was found. For complexation of NaCl by 18-C-6 in methanol, the values obtained by this technique (log K, = 4.64 and AHc = -7.99 k ~ a l / m o l )are ~ ~in good agreement with results given by Abraham (log K, = 4.49, AH, = -7.66 kcal/mol)21 using the batch calorimetric method. They are also comparable with those obtained by calorimetric isoperbol titration.20 The values for the calcium complex of 18-C-6 in methanol (log K , = 3.81, AH, = -3.08 k ~ a l / m o l )obtained ~~ by our method are in excellent agreement with the results of Izatt and co-workers (log K, = -3.86, AH, = -2.75 kcal/mol) obtained by isoperbol titration.20 For some alkali-metal complexes of tetraglyme the results obtained by the present method and the batch calorimetric method are also very similar to those reported in the literature.'O It is hence concluded that the calorimetric method used has yielded reliable results on thermodynamic parameters of metal-ion complexation by various ligands. Results and Discussion 1. Alkali-Metal Complexes in Methanol. Several investiga-
t i o n ~ *have $ ~ ~reported that the linear, open-chain, tetraethylene glycol dimethyl ether, tetraglyme (TG), and its derivatives form 1: 1 cation:iigand complexes with alkali-metal ions in solution. ~
~~~~~~~~~
(23) Buncel, E.; Shin, H . S., unpublished results. (24) (a) Shinohara, M.; Smid, J.; Szaarc, M. J . Am. Chem. SOC.1968, 90, 2175. (b) Grandjean, J.; Laszlo, P.; Vogtle, F.; Sieger, H. Angew. Chem., Int. Ed. Engl. 1978, 17, 856.
4178
The Journal of Physical Chemistry, Vol. 92, No. 14, 1988
5t
24 I-
maFigure 1. Stability constants (log &) of alkali-metal complexes: (0) M+-tetraglyme complexes in methanol, (0)M+-tetraglyme complexes in acetonitrile, (0)Mt-15-C-5 complexes in methanol, (m) M+-l5-C-5 complexes in acetonitrile.
Crystallographic studies" have also shown that, in the solid state, these ligands form 1:l complexes even with the large rubidium cation, Rb', in which the open-chain ligand wraps around the cation in a quasi-cyclic structure. Such a crown ether-like conformation may well be expected also in solution for tetraglyme and its derivative complexes.24a The thermodynamic parameters (log K,, AHc) obtained here for alkali-metal complexes of tetraglyme thus refer to 1:1 species. In Table I are presented the results on the stability constants of 1:1 complexes of tetraglyme with alkali-metal ions together with those of macrocyclic polyethers. The results determined in the present work are in good agreement with those obtained for tetraglyme complexes by potentiometric and conductometric techniques.I0 The results also show that tetraglyme forms less stable complexes than those formed by cyclic synthetic and naturally occurring ionophore^^^^^^ and cryptands.I6 For example, the binding constants for tetraglyme complexes of alkali-metal ions in methanol are lower than those for the cyclic antibiotic Enniatine B and 15-crown-5, by a factor of 2-20. The higher stabilities of cyclic polyethers toward alkali-metal cations, over tetraglyme, result in part from the so-called macrocyclic effect25but also from other factors which will be discussed subsequently. In Figure 1 are shown the stability constants (log K,) of different alkali-metal complexes with tetraglyme and 15-C-520(both ligands have the same number of oxygen binding sites) in methanol and acetonitrile. In methanol, tetraglyme forms the strongest complex with K+ among the alkali-metal ions. Thus, the selectivity order for TG-M+ complexes in MeOH is K+ > Rb' > Cs' > Na' > Li+. The greater stability of K+ over Na' was also observed for macrocyclic crown ethers,26 even for the smaller 12-C-4 and 15-C-5, whose cavity sizes correspond closer to Li' and to Na+ than to K+. A similar selectivity order is found for 15-C-5, except that the positions of Na+ and Cs' are reversed, Le., K+ > Na+ > Cs+ >> Li+.14a320 It is likely that the higher selectivity of Kf over Na' for tetraglyme and the crown ethers in methanol is due to the stronger solvation of Na' compared to K+ in this solvent,27 since the essentially two-dimensional structure of these ligands should still allow significant interaction of solvent with the complexed cation. However, this would hold to a much lesser degree for the cryptands, with their three-dimensional structure, which shield more effectively the cations from solvent interactions.I6 The selectivity of tetraglyme toward alkali-metal cations in MeOH is also lower than that of 15-C-5. For example, the stability constants of TG-M+ complexes vary only within 1 order of magnitude, compared with more than 2 orders of magnitude (25) (a) Cabbiness, K.; Margerum, D. W. J . Am. Chem. SOC.1969, 91, 6540. (b) Frensdorff, H. K.J . Am. Chem. SOC.1971, 93, 600. (26) Gokel, G . W.; Goli, M. M.; Minganti, C.; Eschegoyen, L. J . Am. Chem. SOC.1983, 105, 6786. (27) Cox, B . G . Annu. Rep. Prog. Chem. Sect. A 1973, 70, 249.
Buncel et al. displayed by 15-C-5. Despite the fact that both ligands form stronger complexes with K+, the ratio of binding constants of 15-C-5 and tetraglyme, K,( 15-C-5)/KS(TG), to a given cation is highest for Na+ and lowest for Li+; the values of K,( 15-C-5)/ K,(TG) are 2, 130, 47, and 15 for Li', Na+, K+, and Cs', respectively. This may originate from contribution of the macrocyclic effect for 15-C-5, in which the Na+ cation is optimally fitted into the cavity of 15-C-5, thereby increasing the stability of the Naf complex. However, the macrocyclic effect may be modulated by larger or smaller cations.3a A larger cation like Cs' will be situated appreciably outside the ligand cavity and therefore exposed to solvent interaction, while the Li' ion is too small to interact effectively with all of the donor sites of 15-'2-5. This leads to a strained conformation of the complex, thereby decreasing the stability. On the other hand, tetraglyme with its flexible open chain can easily wrap around the small Li+ cation. Consequently, the selectivity between 15-c-5 and tetraglyme should be lower for Cs+ and Li' than for Na'. 2. The Stabilities of Alkali-Metal-Tetraglyme Complexes in Acetonitrile. The results on the stability constants of alkalimetal-tetraglyme complexes in acetonitrile are included in Table I and illustrated in Figure 1. It is seen that, in contrast to the situation found for methanol, in acetonitrile as solvent tetraglyme forms the most stable complex with Na+ among the alkali-metal ions. Thus the selectivity order for tetraglyme in acetonitrile is Na+ > Li' > K+ > Rb' > Cs'. In methanol, the selectivity order was found to be K+ > Rb' > Cs+ > Na+ > Li+. Though only few results are available for 15-C-5 in acetonitrile, the same order of selectivity was obtained, namely Na+ > Li' > K+.28 The selectivity for TG-M+ complexes in acetonitrile is lower than for 15-C-5, as had also been observed in methanol. The stability constants of TG-M+ complexes vary within 1 order of magnitude, compared with about 3 orders of magnitude displayed by 15-C-5 in acetonitrile. The selectivity of 15-C-5 compared to tetraglyme, K,( 15-CS)/K,(TG), in acetonitrile, is also highest for Na'; Le., the values of K,(lS-C-S)/K,(TG) are 27, 759, and 9 for Li', Na', and K+ respectively. 3. Solvent Effects on Stabilities of Complexes. The results in Table I clearly show that the stability constants of alkalimetal-tetraglyme complexes are higher in acetonitrile than in methanol. A larger increase is observed on going from methanol to acetonitrile for the small cations; the stability constant is increased by a factor of -20 for Li' and Na+ but only by -2 for K', Rb', and Cs+. Consequently, the selectivity of tetraglyme toward alkali-metal cations is shifted from the K+ complex being the most stable in methanol to the Naf complex in acetonitrile. This is reflected in the fact that the solvation energies of Li' and Na+ cations are much less in acetonitrile than in methan01.~' The free energies of transfer from methanol to acetonitrile are reported as 26 and 5 kJ/mol for Li+ and Na', respectively, compared to -2, -3.3, and -4.6 kJ/mol for K+, Rb+, and Cs+, respecti~ely.~' Several s t u d i e ~ ' ~ indicate J ~ 3 ~ ~ that crown ethers and cryptands are solvated more strongly in methanol than in the polar aprotic solvent acetonitrile, presumably because of hydrogen-bond formation involving the protic solvent and the donor sites of the ligand. This should also hold for tetraglyme. Thus the slight increase in the stability constants of the K', Rb', and Cs+ complexes in acetonitrile compared to methanol suggests that the larger solvation energies of these cations in acetonitrile may be compensated by the less effective solvation of tetraglyme and its cation complexes in this solvent. The change of stability constant with solvent variation is related to the effect of solvent on the free energies of the species involved in complexation. This may be expressed quantitatively by eq 1 AG,,(ML+) - AG,,(L) = -2.303RT log [K,(S2)/Ks(S1)]+ AG,,(M+) (1) ~~~~~
~~~
(28) Hopkins, Jr., H. P.;Norman, A. B. J . Phys. Chem. 1980, 84, 309. (29) Nakamura, T.; Yumoto, Y.; Izutsu, K. Bull. Chem. SOC.Jpn. 1982, 55, 1850.
The Journal of Physical Chemistry, Vol. 92, No. 14, 1988 4179
Alkali-Metal Complexes of Tetraglyme
-
TABLE 11: Free Energies of Transfer of Alkali-Metal Cations [AC,(Mt), kJ/mol] for Methanol Acetonitrile, and Differences in Free Energies of Transfer between the Complexed and Free Ligands AC,(L), kJ/mol] for Methanol Acetonitrile, at 25 [AC,(MLt)
-
-
O P
cation Lit Na' K+
Rb+
cs+
AG,,(Mt)' 25.5 5.4 -2.1 -3.3 -4.6
TGb 18.2 -2.0 -4.0 -5.6 -5.1
15-C-SC (2.2.1)d 12.0 -4.7 0.6
-2.7 -6.0 -7.6 -6.3 -9.3
3t
(2.2.2)d -6.0 -5.0 -7.1 -8.0 -7.6
"Reference 27, based on the assumption that AG,,(BPh4-) = AG,,(Ph4Ast). bThis work (Table I). cReferences 14a, 20, 22, and 28. dReference 16.
't I
0
1 20
10
30
DN
Figure 3. Plot of stability constant (log K,)of Nat-tetraglyme complex in various solvents versus the donor number of the solvent. TABLE III: Effect of Counteranion on the Stability Constants of Potassium Complexes of Tetraglyme in Methanol and Acetonitrile at 25 OC loa K. solvent KBPh, KI KSCN KC104 KPic KPNPhO KOPh 1.68 1.60 1.64 1.61 methanol 1.81 1.64 1.67 1.97 1.74 2.02 2.00 aceto2.08 2.06 2.03
nitrile -104
,
?
2
-
10
20
30
AGtrIM') (KJ/mol)
Figure 2. Plot of AG,,(MLt) - AG,,(L) versus AG,,(M+) for transfer from methanol to acetonitrile at 25 OC for various alkali-metal complexes by (0)tetraglyme, (A) (2.2.1)cryptand, and (0)(2.2.2)cryptand.
for solvents SIand S2,where K, is the stability constant of complex ML+ and AG,, refers to the free energy of transfer from SI to S2. The value of AC,,(ML+) - AG,,(L) can be determined via eq 1 from the stability constants of the ligands and the available values of AG,,(M+) for the two solvents studied here.27 The variation of AG,,(ML+) - AG,(L) for different cations may give information on solvation of the ML+ complexes as function of the cation. In Table I1 are presented the results on the values of AG,,(M+) and AG,,(ML+) - AG,,(L) for tetraglyme, 15-C-5, (2.2.1), and (2.2.2) cryptands,I6 for transfer from methanol to acetonitrile. Figure 2 shows the plot of AG,,(ML+) - AG,(L) against AG,(M+) for tetraglyme and the cryptands. Although the points are somewhat scattered, it is apparent that the values of AG,,(ML+) - AG,,(L) for tetraglyme are strongly dependent on AG,,(M+), while the values for the cryptands are nearly constant for the various alkali-metal cations. The slopes of the plots are 0.65 and 0.1 for tetraglyme and the cryptands, respectively. It is difficult to determine the value of the slope for 15-C-5, since the results for this ligand are few and show appreciable scatter. However, the results from Table I1 show that the values of AG,,(ML+) AG,,(L) for 15-C-5 are also strongly dependent upon AG,,(M+). Moreover, Kolthoff and Chantooni have recently reported a slope of 0.37 in the corresponding plot for dibenzo-18-C-6 on transfer from methanol to dipolar apr'otic solvents such as acetonitrile, propylene carbonate, N,N-dimethylformamide, and dimethyl s~lfoxide.'~ Since the value of AC,,(L) is constant for transfer of a given ligand between two solvents and does not affect the slope of the above plot, it follows that the slope will reflect the dependence of AG,,(ML+j on the metal ion contained in the ML+ complex. It thus appears that there is still a significant interaction between the solvent and the complexed cation (slope > 0). This interaction is strong for tetraglyme complexes but decreases for crown ethers and even more so for cryptands (slope = 0) where the solvent interaction with the complexed cation may be regarded as negligible.
Only a few results have been reported hitherto on stability constants of tetraglyme-cation complexes in different solvents. Smid and his cc-workers1Ibhave reported that the binding constant (log K,) for sodium fluorenyl-tetraglyme complex in T H F (dielectric constant, e = 7.4) is 2.23, while for the lithium fluorenyl complex with tetraglyme in dioxane (e = 2.2), log K , = 2.38. These values can be compared with our results in methanol ( e = 32.7) and in acetonitrile (AN) (e = 37.5). The stability constants for the TG+-Na complex thus decrease in the order A N > T H F > MeOH. It is hence apparent that the dielectric constant of the solvent does not play an important role in determining the stabilities of the cation complexes. However, the stability constant for a given complex seems to be related to the electron-donor character of the solvent toward the cation, defined as Gutmann's empirical donor number (DN).30 Figure 3 shows the fair relationship between the stability constant of the TG-Na+ complex and the donor number of the solvent.30 This is not unexpected, if one considers the ion-solvent interaction in the framework of electron pair donor-electron pair acceptor interactions. Thus the solvent capable of strong electron pair donation (as reflected by the donor number) can strongly solvate the cations and consequently reduce the stabilities of the cation-ligand complexes. A similar relationship between the DN value of the solvent and the stability constant has recently been observed for crown ether31 and cryptand complexes.16 4. Anion Effect. In Table I11 are shown the results on the effect of the counteranion on stability constants of the TG-K+ complexes in methanol and in acetonitrile. It is seen that the effect of the anion is almost negligible in methanol, and also in acetonitrile, with the exception of phenoxide anion in the latter solvent for which K, is appreciably lower. The relative lack of a significant dependence of K, on the anion is consistent with results for modified 18-C-6 complexes.32 The salts may exist in both free ion and ion-paired states, depending on the nature of the solvent and the counteranion. As discussed above, the cation complexed by tetraglyme is still exposed to interaction by the solvent and the counteranion. In poor solvating media like acetonitrile, ion-pair formation between the uncomplexed or tetraglyme-complexed cation and the counteranion may occur as well. The lower K, value for KOPh in acetonitrile (30) (a) Gutmann, V. Coordination Chemistry in Nonaqueous Solvents; Springer Verlag: Vienna, 1968. (b) Popov, A. I. In Solute-Solvent Interactions; Coetzee, J. F., Ritchie, C. D., Eds.; Marcel Dekker: New York, 1976; Vol. 2,p 271. (31)Shamsipur, M.; Popov, A. I. Inorg. Chim.Acta 1980, 43, 243.
4180
J . Phys. Chem. 1988, 92, 4180-4184
groups of the ligand in a cyclic-like conformation of the Li' complex, as well as a high solvation energy of the Li' cation. As mentioned above, tetraglyme may transform its conformation from a flexible linear chain in the uncomplexed state to a fairly rigid quasi-cyclic conformations.*lain the complexed state. This would lead to a significant loss in entropy upon complexation. However, this loss may also be, partly, compensated by an increase in entropy resulting from solvent release on desolvation of the cation on complexation. This increase will be largest for the smallest cation. Thus the negative entropy of cation complexation in both solvents may primarily be due to a conformational change could hence result from significant ion pairing in this system. of tetraglyme, except with the small Li' ion, in which the entropic However, in the polar protic solvent methanol, capable of gain on desolvation of the cation should effectively compensate solvating both cations and anions, such a difference in association the loss in entropy from a conformational change of tetraglyme, between a given cation and various anions would be d i m i n i ~ h e d , ~ ~ . ~since ~ the value of TAS, is positive in both solvents. which is in accord with the present results. The TAS, values for a given complex is found to be more 5 . Thermodynamics of Complex Formation. In Table IV are negative in methanol than in acetonitrile. This suggests that the listed the thermodynamic parameters for alkali-metal complexation metal ion and tetraglyme may not be completely desolvated in by tetraglyme in methanol and acetonitrile at 25 OC. The T U , methanol, and hydrogen bonding between methanol molecules and values are calculated from experimental values of AG, and AH, the cation complexes may still exist. On the other hand, the origin by means of the relationship AG, = -RT log K, = AH,- TU,. of the more negative values of AH, in methanol is not clear at The free energy of complexation, AG,, results from contributions present. Recently Takeda and his c o - ~ o r k e r shave ~ ~ also found of both enthalpy and entropy terms, generally including the binding that the values of AHc and TAS, for alkali-metal complexes of energy between complexone and cation, the energy of conforthe flexible dibenzo-18-C-6 are more negative in methanol than mational change of the ligand as a consequence of complex forin acetonitrile and propylene carbonate. mation, and the energies of desolvating the ligand and the cation. The AH,values for alkali-metal complexes of tetraglyme are The main energy contribution to stability varies from one comcomparable: to those of 15-C-5,14a,20,29 but the TAS, values for the plexone to another depending on the solvent, the cation, and the tetraglyme complexes are much lower. Thus the high selectivity flexibility of the complexone, as well as on the nature of binding between 1542-5 and tetraglyme toward a given cation, Ks(15-Csites on the complexone, etc. Nevertheless, the experimental data 5)/Ks(TG) >> 1, results mainly from a loss in entropy in preomay indicate whether the complexation is enthalpic or entropic rientation of the open-chain tetraglyme on complexation. This in origin. result is in accord with previous findings on the origin of the The results in Table IV show that the stability of the TG-M+ macrocyclic e f f e ~ t . ~ ~ , ~ ~ complexes is enthalpy dominated, accompanied by an unfavorable Acknowledgment. This research was supported by Supply and decrease of entropy (TAS, < 0), except for the Li' complex in Services Canada, Department of National Defence. We thank both solvents, for which the stability constant has contributions Dr. B. G. Cox for helpful discussions and Dr. M. H. Abraham from both enthalpic and entropic terms ( A H , < 0, TAS, > 0 ) . and Dr. H. C. Ling for provision of computer programs for In both solvents, the values of AH,are rather similar for Na', calculation of stability constants. K+, Rb', and Cs' cations, and more negative than that for the Li' cation. The less favorable complexation enthalpy of Li' Registry No. AN, 75-05-8; MeOH, 67-56-1. suggests that tetraglyme does not bind optimally with this small cation, probably because of a steric effect between the two terminal TABLE I V Thermodynamic Data for Alkali-Metal (as Perchlorate Salts) Complex Formation by Tetraglyme in Methanol and Acetonitrile at 25 O C (in kJ/mol) solvent Li+ Na' K+ Rb' Cs+ methanol AG, -5.1 -6.3 -9.6 -8.7 -8.3 -30.5 -25.0 -32.4 -27.3 AH, -2.8 -21.8 -16.7 -26.1 -17.7 TAS, +2.3 acetonitrile AG, -12.4 -13.7 -11.5 -10.9 -8.7 -20.3 -19.7 -21.5 -21.1 AHc -5.7 -9.4 -11.0 -7.8 -9.6 TAS, +6.7
(32) Tusek-Boric, L. J.; Danesi, P. R. J. Inorg. Nucl. Chem. 1979, 41, 833. (33) Jackman, L. M.; Lange, B. C. Tetrahedron 1977. 33, 2737.
(34) (a) Kodama, M.; Kimura, E. J. Chem. Soc., Dalton Trans. 1976, 116. (b) Kodama, M.; Kinura, E. Bull. Chem. SOC.Jpn. 1976, 49, 2465. (35) Takeda, T.; Kudo, Y.;Fujiwara, S. Bull. Chem. Soc. Jpn. 1985, 58, 1315.
Reactions of Singlet and Triplet Methylene with a C-H Bond of Ethylene. An ab Initio Study Miquel Moreno, Jose M. Lluch, Antonio Oliva, and Juan BertrPn* Departament de Quimica, Universitat Autbnoma de Barcelona, 08193 Bellaterra, Barcelona, Spain (Received: August 13, 1987)
The insertion reaction of singlet methylene into a C-H bond of ethylene and the hydrogen abstraction from ethylene by triplet methylene have been theoretically studied by using the 3-21G and 6-31G* basis sets and introducing electron correlation and thermodynamic corrections. The obtained results have permitted us to analyze the mechanistic differences between both processes. The competition between them and the well-known addition reactions to olefinic double bonds is also discussed.
Introduction Methylene reactions have become the topic of an increasing number of experimental and theoretical studies in recent years. 1*2 It has been found that the product composition of many systems (1) Baird, M. S. Annu. Rep. Prog. Chem., Sect. B 1984, 81, 79. (2) Skell, P.S. Tetrahedron 1985, 41, 1427.
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involving methylene is profoundly influenced by the presence of hydrogen atoms3f4 It is now firmly established that triplet methylene abstracts hydrogen atoms from hydrocarbons while the analogous reactions with singlet methylene yield very fast insertions (3) Buehler, C. A. J. Chem. Educ. 1972, 49, 239. (4) Isaacs, N. S. Reactiue Intermediates; Wiley: London, 1974; p 375.
0 1988 American Chemical Society
