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Cation-Ether Complexes in the Gas Phase: Bond Dissociation Energies and Equilibrium Structures of Li+(1,2-dimethoxyethane)x, x ) 1 and 2, and Li+(12-crown-4) Douglas Ray,*,† David Feller,*,† Michelle B. More,‡ Eric D. Glendening,†,§ and P. B. Armentrout*,‡ EnVironmental Molecular Sciences Laboratory, Pacific Northwest National Laboratory, Richland, Washington 99352, and Department of Chemistry, UniVersity of Utah, Salt Lake City, Utah 84112 ReceiVed: April 9, 1996; In Final Form: July 17, 1996X
Bond dissociation energies, equilibrium structures, and harmonic vibrational frequencies are reported for Li+(DXE), where DXE ) CH3O(CH2)2OCH3, Li+(DXE)2, and Li+(12-crown-4). The bond dissociation energies are determined experimentally by analysis of the thresholds for collision-induced dissociation of the cation-ether complexes by xenon (measured using guided ion beam mass spectrometry) and computationally by ab initio electronic structure calculations. For Li+(DXE)x, x ) 1 and 2, the primary and lowest energy dissociation channel observed experimentally is endothermic loss of one dimethoxyethane molecule. For Li+(12-crown-4), the primary dissociation channel is endothermic loss of the intact crown ether, although ligand fragmentation is also observed. The cross section thresholds are interpreted to yield 0 and 298 K bond energies after accounting for the effects of multiple ion-molecule collisions, internal energy of the complexes, and unimolecular decay rates. The calculated and experimentally-derived bond energies are in good agreement for Li+(DXE), are in reasonable agreement for Li+(12-crown-4), and differ by 32 ( 12 kJ/mol for Li+(DXE)2. On average, the experimental bond dissociation energies differ from theory by 9 ( 6 kJ/mol per metal-oxygen interaction. The equilibrium structures are determined primarily by strong electrostatic and polarization interactions between Li+ and the ligands. Charge transfer interactions are also important, as indicated by a natural energy decomposition analysis. Correlations between the bond dissociation energies and the equilibrium structures demonstrate that the orientation of the C-O-C subunits in the ethers relative to the metal cation is more important than the Li+‚‚‚O bond length in determining the stability of the complexes as predicted by Hay et al.1,2
I. Introduction Noncovalent interactions between ions and neutral molecules are of fundamental importance in molecular recognition phenomena occurring in complex chemical and biochemical systems.3 The study of a series of cation-ether complexes composed of different metal cations and a selection of ligands ranging from simple, monodentate ethers to cyclic polyethers provides an opportunity to examine the noncovalent interactions operative in “simple” ion-molecule complexes. Cation-ether complexes are also interesting from a practical point of view. Crown ethers have been proposed for use in new chemical separations technologies4 and in the development of advanced analytical methods.5 Computational models capable of reliably predicting ligand selectivity in a variety of condensed-phase environments would be valuable tools for the advancement of separations technologies. Such methods are currently under development; however, the development is hindered by a lack of suitable experimental data. One goal of the present work is to provide accurate experimental data to address this deficiency. The principal challenge to quantitative theoretical descriptions of ion-molecule interactions arises from the need to accurately reproduce both the dominant classical electrostatic and quantum mechanical contributions to the interaction. The relative importance of these two effects varies from system to system, but two recent studies6,7 of Li+-ether complexes have dem†
Pacific Northwest National Laboratory. University of Utah. § Current address: Department of Chemistry, Indiana State University, Terre Haute, IN 47809. X Abstract published in AdVance ACS Abstracts, September 1, 1996. ‡
onstrated that both contributions are essential. Treating relativelylarge chemical systems with the requisite level of theory is technically very challenging. In general, small basis set Hartree-Fock (HF) theory is incapable of accurately describing such interactions. Because theory is ultimately called upon to model a wide variety of molecular systems, a successful method must exhibit a certain degree of balance in its treatment of possibly competitive interactions. For example, the selectivity of macrocyclic ligands, such as 18-crown-6, toward alkali cations in aqueous solution hinges on a subtle balance between ion-ligand, ion-water, ligand-water, and water-water interactions. We have previously7,8 identified second-order perturbation theory coupled with polarized basis sets as a good entrylevel theoretical approach for cation-neutral systems. Accurate measurement of the intrinsic properties of cationether complexes is the principal challenge from the experimental perspective. Studies of ion-molecule complexes in the gas phase can provide insight into the intrinsic aspects of the interactions, because the isolated complexes are unperturbed by solvent-induced phenomena. Several groups have reported measurements on complexes between alkali metal cations and polyethers (crown ethers and acyclic polyethers) in the gas phase. Dearden and co-workers have measured rate constants for association reactions yielding cation-polyether and polyether-cation-polyether complexes,9-11 rate constants for cation transfer between various polyethers, and equilibrium constants for cation transfer between large crown ethers (e.g., 18-crown-6 and 21-crown-7).12,13 Brodbelt and co-workers have applied the kinetic method14 to adducts containing two different polyethers bound by a cation to determine the relative affinities
S0022-3654(96)01060-X This article not subject to U.S. Copyright. Published 1996 by the American Chemical Society
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of a variety of polyethers for a given alkali metal cation15 and to complexes containing one crown ether, one metal ion, and one metal halide to determine the relative affinities of the alkali cations for a given crown ether.16 Eyler and co-workers have reported semiquantitative estimates of the bond dissociation energies (BDEs) of K+(18-crown-6) and Cs+(18-crown-6) from collision-induced dissociation (CID) measurements with a Fourier transform ion cyclotron resonance mass spectrometer (FTICR).17 In the present study, CID measurements are made with a guided ion beam mass spectrometer, an instrument specifically designed for measurements of the kinetic energy dependence of collision-induced phenomena. In the work reported here, we present quantitative bond dissociation energies (BDE) and geometrical structures of gas phase complexes formed between Li+ and the ligands 1, 2-dimethoxyethane (DXE) and 12-crown-4 (12c4). The results, in combination with the results of our previous study6 of Li+ complexes with dimethyl ether (DME), facilitate discussion of the intrinsic factors operative in the complexation of cations by multidentate ligands. II. Theoretical and Experimental Methods A. Theoretical. Geometries and BDEs of the Li+(DXE)x and Li+(12c4) complexes were obtained from ab initio restricted Hartree-Fock (RHF) and second-order and fourth-order MøllerPlesset perturbation (MP2 and MP4(SDTQ)) theoretical methods using the standard 6-31+G* basis set18-20 on H, Li, and O and the smaller 6-31G* basis set21 for carbon. This basis set included a single d-type polarization function on Li, C, and O, as well as an extra diffuse shell of sp functions on Li and O. Previous calculations on cation-ether6,7 and cation-water complexes22 revealed that the diffuse functions on O reduce undesirable basis set superposition errors (BSSE), which result in an overestimation of BDEs. The diffuse sp functions on C, which are present in the standard 6-31+G* set, were not included in order to keep the size of the calculations manageable. Selected tests with the full 6-31+G* basis set indicated only marginal changes on calculated geometrical features, BDEs, and BSSE. For example, bond lengths in the Li+(dimethyl ether) complex differed by (0.001 Å and BDEs changed by DXE > DME). C. The Macrocyclic Effect. The data presented here may provide some insight into the origins of the macrocyclic effect.63 The macrocyclic effect refers to the greater thermodynamic stability of a complex with a cyclic polydentate ligand in comparison to the complex formed by the corresponding acyclic ligand (e.g., Li+(12-crown-4) and Li+(1,2-bis(2-methoxyethoxy)ethane) (CH3O(CH2)2O(CH2)2O(CH2)2OCH3 or triglyme)). The macrocyclic effect is manifested in solution by higher binding constants and greater selectivity for a given metal ion by cyclic Vs acyclic ligands. The entropy of association, which is smaller for cyclic ligands because their donor atoms are prearranged to accommodate a guest species, and the enthalpy of desolvation, which is greater for acyclic ligands because of their conformational flexibility, are invoked to explain the macrocyclic effect. The enthalpy of association is not considered to make a significant contribution to the macrocyclic effect as the difference between the enthalpies of association of metal ions to cyclic and acyclic ligands is presumed to be negligible.64,65 Dearden et al.11 have shown that the efficiencies of 1:1 metal ion-ligand association reactions in the gas phase are greater for cyclic ligands (i.e., crown ethers) than for the corresponding acyclic glymes at 350 K. This manifestation of the macrocyclic effect was presumed to be entropically driven (because solvation is not operative in the gas phase). The results reported here, which demonstrate that preorganized multidentate ligands have smaller enthalpies of association than corresponding unconstrained ligands, provide support for this interpretation. A more quantitative discussion of the gas phase macrocyclic effect in Li+(12c4) requires determination of the binding energy of Li+ to triglyme in the gas phase. V. Conclusions Values for the BDEs of Li+[1,2-dimethoxyethane]x, x ) 1 and 2, and Li+(12-crown-4) are obtained by a combination of experimental and theoretical methods. Values for the BDEs obtained from ab initio calculations employing RHF and MP2 methods using 6-31+G* hybrid basis sets are in relatively-good agreement with the experimentally-derived values. Theory finds that the equilibrium structures of these complexes are determined primarily by strong electrostatic, polarization, and charge transfer interactions. Correlations between the BDEs and the equilibrium structures demonstrate that the orientation of the C-O-C subunits relative to the metal cation is more important than the Li+‚‚‚O bond length in determining the stability of the complexes. We conclude that electronic structure calculations of this type are of utility to a quantitative description of the structure and energetics of cation-ether complexes. We also conclude that measurement of collision-induced dissociation cross sections as a function of translational energy can be used to yield accurate thermochemistry of cation-ether complexes when effects associated with multiple collisions, the internal energy of the complexes, and the lifetime for dissociation are carefully controlled and analyzed. We anticipate that the methods used here will be of utility in the development of accurate models of the noncovalent interactions between ions and neutral molecules operative in molecular recognition and related phenomena in condensed phases. Calculations and measurements of similar
Cation-Ether Complexes in the Gas Phase accuracy as those presented here on complexes of other alkali cations with acyclic and other cyclic ethers are currently underway. Measurements on hydrated cation-crown ether complexes are also in progress. Acknowledgment. Funding for this work was provided by the Division of Chemical Sciences, Office of Basic Energy Sciences, U.S. Department of Energy (D.R., D.F., M.B.M., E.D.G.) and by the National Science Foundation under Grant CHE-9530412 (P.B.A.). E.D.G. acknowledges the support of Associated Western Universities, Inc. under Grant DE-FG0689ER-75522 with the U.S. Department of Energy. Pacific Northwest National Laboratory is operated for the U.S. Department of Energy by Battelle under Contract DE-AC06-76RLO 1830. Supporting Information Available: Calculated bond dissociation energies and equilibrium structures of Li+[1,2dimethoxyethane]x, x ) 1 and 2, and Li+[12-crown-4] (9 pages). Ordering information is given on any current masthead page. References and Notes (1) Hay, B. P.; Rustad, J. R. J. Am. Chem. Soc. 1994, 116, 6316. (2) Hay, B. P.; Rustad, J. R.; Hostetler, C. J. J. Am. Chem. Soc. 1993, 115, 11158. (3) See, for example: Dougherty, D. A. Science 1996, 271, 163. (4) See, for example: Horwitz, E. P.; Dietz, M. L.; Fisher, D. E. SolVent Extr. Ion Exch. 1991, 9, 1. Moyer, B. A; Delmau, L. H.; Case, G. N.; Bajo, S.; Baes, C. F. Sep. Sci. Technol. 1995, 30, 1047. (5) See, for example: Grate, J. W.; Strebin, R.; Janata, J.; Egorov, O.; Ruzicka, J. Anal. Chem. 1996, 68, 333. (6) More, M. B.; Glendening, E. D.; Ray, D.; Feller, D.; Armentrout, P. B. J. Phys. Chem. 1996, 100, 1605. (7) Glendening, E. D.; Feller, D.; Thompson, M. A. J. Am. Chem. Soc. 1994, 116, 10657. (8) Feller, D.; Apra, E.; Nichols, J. A.; Bernholdt, D. E. J. Chem. Phys., to be published. (9) Zhang, H.; Chu, I.-H.; Leming, S.; Dearden, D. V. J. Am. Chem. Soc. 1991, 113, 7415. (10) Zhang, H.; Dearden, D. V. J. Am. Chem. Soc. 1992, 114, 2755. (11) Dearden, D. V.; Zhang, H.; Chu, I.-C.; Chen, Q. Pure Appl. Chem. 1993, 65, 423. (12) Chu, I.-H.; Zhang, H.; Dearden, D. V. J. Am. Chem. Soc. 1993, 115, 5736. (13) Chu, I.-H.; Dearden, D. V. J. Am. Chem. Soc. 1995, 117, 8197. (14) (a) Cooks, R. G.; Kruger, T. L. J. Am. Chem. Soc. 1977, 99, 1279. (b) McLuckey, S. A. ; Cameron, D.; Cooks, R. G. J. Am. Chem. Soc. 1981, 103, 1313. (c) McLuckey, S. A.; Schoen, A. E.; Cooks, R. G. J. Am. Chem. Soc. 1982, 104, 848. (15) Liou, C.-C.; Brodbelt; J. J. Am. Soc. Mass Spectrom. 1992, 3, 543. (16) Maleknia, S.; Brodbelt, J. J. Am. Chem. Soc. 1992, 114, 4295. (17) Katritzky, A. R.; Malhotra, N.; Ramanathan, R.; Kemerait, R. C.; Zimmerman, J. A.; Eyler, J. R. Rapid Commun. Mass Spectrom. 1992, 6, 25. (18) Dill, J. D.; Pople, J. A. J. Chem. Phys. 1975, 62, 2921. (19) Clark, T.; Chandrasekhar, J.; Spitznagel, G. W.; Schleyer, P. v. R. J. Comput. Chem. 1983, 4, 294. (20) Hehre, W. J.; Ditchfield, R.; Pople, J. A. J. Chem. Phys. 1972, 56, 2257. (21) Hariharan, P. C.; Pople, J. A. Theor. Chim. Acta 1973, 28, 213. (22) Glendening, E. D.; Feller, D. J. Phys. Chem. 1995, 99, 3060. (23) Boys, S. F.; Bernardi, F. Mol. Phys. 1970, 19, 553. (24) Frisch, M. J.; Trucks, G. W.; Head-Gordon, M.; Gill, P. M. W.; Wong, M. W.; Foresman, J. B.; Johnson, B. G.; Schlegel, H. B.; Robb, M. A.; Replogle, E. S.; Gomperts, R.; Andres, J. L.; Raghavachari, K.; Binkley, J. S.; Gonzalez, C.; Martin, R. L.; Fox, D.J.; Defrees, D. J.; Baker, J.; Stewart, J. J. P.; Pople, J. A. Gaussian-92/DFT, ReVision C; Gaussian, Inc.: Pittsburg, PA, 1992. (25) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Gill, P. M. W.; Johnson, B. G.; Robb, M. A.; Cheeseman, J. R.; Keith, T. A.; Petersson, G. A.; Montgomery, J. A.; Raghavachari, K.; Al-Laham, M. A.; Zakrzewski, V. G.; Ortiz, J. V.; Foresman, J. B.; Cioslowski, J.; Stefanov, B. B.; Nanayakkara, A.; Challacombe, M.; Martin, R. L.; Fox, D. J.; Binkley, J. S.; Defrees, D. J.; Baker, J.; Stewart, J. J. P.; Head-Gordon, M.; Gonzalez, C.; Pople, J. A. Gaussian-94, ReVision C; Gaussian, Inc.: Pittsburg, PA, 1995.
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