W3 25¢ per day per item RETURNING LIBRARY MATERIALS: Place in book return to remove rge from circulation records MULTINUCLEAR MAGNETIC RESONANCE AND CALORIMETRIC STUDIES OF SODIUM CATION COMPLEXES WITH SOME CROWN ETHERS AND CRYPTANDS IN VARIOUS SOLVENTS BY Jy Dale Lin A DISSERTATION Submitted to Michigan State University in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry 1980 ABSTRACT MULTINUCLEAR MAGNETIC RESONANCE AND CALORIMETRIC STUDIES OF SODIUM CATION COMPLEXES WITH SOME CROWN ETHERS AND CRYPTANDS IN VARIOUS SOLVENTS By Jy Dale Lin Sodium (23Na), carbon (13C) and proton (1H) nuclear magnetic resonance methods were used to study sodium cation complexes with three crown ethers; 18-crown-6 (18C6), 15- crown-5 (15C5) and monobenzo-lS-crown—S (MB15C5), in aqueous and in several nonaqueous solvents. Concentration formation constants of these complexes were determined from the variation of the 23Na chemical shift as a function of ligand/Na+ mole ratio or from the variations of the 13C chemical shifts with Na+/ligand mole ratios. In general, the stability of the complex varies inversely with the solvating ability of solvent as expressed by the Gutmann donor number. For example, the formation constants for Na+-18C6 complex are log KE:4 in tetrahydrofuran (THF), acetone (MeZCO), propylene carbonate (PC) and nitromethane (MeNOZ), log K = 3.8 i 0.2 in acetonitrile (MeCN), log K = 1.A1 : 0.07 in dimethylsulfoxide (DMSO), log K = 2.31 i 0.05 Jy Dale Lin in dimethylformamide (DMF) and log K = 0.82 i 0.05 with NaI salt in water. The relative stabilities of the complexes are in the order Na+~18C6>*Na+-1505:>Na+-MB1505. The formation of a 2:1 (ligand:Na+ ion) complex is more favorable in a poorly solvating solvent such as MeNO2 if the ligand cavity size is smaller than the cation size. The 2:1 complex formation constants in MeNO2 solution were found to be K = 1.6 i 0.2 and K = 0.8 1 0.2 for sodium- 2 2 15C5 and sodium-MB15C5 complexes respectively. The apparent exchange rate of the sodium cation between the bulk solution and the Na+°18C6 complex in solvents of low dielectric constant such as THF and 1,3-dioxolane was found to be dependent on the anion. The exchange is slower with tetraphenylborate anion than with perchlorate or iodide anions. Complexation of the sodium cation by four cryptands, 0211, C221, C222 and C2228.in water and in several nonaqueous solvents was studied by the sodium-23 NMR technique. In most cases studied, the cation exchange between the bulk solutionand the complex is slow on the NMR time scale since two resonances of the 23Na nucleus were observed when the Na+ ion was present in excess. In all solvents used in this study, the three cryptands, C221, C222 and C222B, form very stable 1:1 complexes (log Kiih) with the sodium cation. The resonance linewidth of the Na+-0211 complex is so broad (about 150 Hz to 300 Hz) that no definite conclusion about Jy Dale Lin the stability of the complex can be made. For all the cryptates, the chemical shifts of the complexed Na+ ion are solvent-dependent indicating that the Na+ ion is not completely shielded by the cryptands from the external medium. The chemical shifts of the individual cryptate in the various solvents studied are in the same general area and the values are around +11 ppm, -h ppm, -9 ppm and -11 ppm for Na+-C211, Na+-C221, Na+-C222B and Na+o 0222 cryptates respectively. The thermodynamic parameters of sodium cation complexes with three crown ethers, 18C6, 1505 and MB15C5, were determined by 23Na NMR and calorimetric techniques in several nonaqueous solvents. For Na+-15C5 complex in DMF solution the values obtained from both methods agree within experimental error. The values are AHO = 44.7 1 0.8 kcal/ mole andASO = -7 j; 3 e.u. from the NMR method, and AH0 = -4.2 1 0.4 kcal/mole andASO = -5 i 1 e.u. from the heat of complexation determined calorimetrically and the stability constant determined by 23Na NMR. In most cases, the sodium-crown complexes are enthalpy stabilized but entropy destabilized. However, in few cases the complexes were found to be entirely entropy stabilized. To My Parents. ii ACKNOWLEDGMENTS The author wishes to thank Professor Alexander I. Popov for his constant guidance throughout this study. Professor James L. Dye is acknowledged for his helpful suggestions as second reader. Gratitude is extended to the Department of Chemistry. Michigan State University and the National Science Founda- tion for financial aid. Many thanks go to all the members of Dr. A. I. Popov's research group for their friendship and stimulation. Deep appreciation to my parents, brother and sisters for their love and encouragement during the course of this study. To my parents, I dedicate this thesis. iii Chapter LIST OF LIST OF CHAPTER 1. 2. 3. 4. CHAPTER 1. 2. 3. CHAPTER TABLE OF CONTENTS Page TABLES . . . . . . . . . . . . . . . . . . vi FIGURES . . . . . . . . . . . . . . . . . x I. HISTORICAL REVIEW INTRODUCTION . . . . . . . . . . . . . . . 1 COMPLEXATION 0F ALKALI METAL CATIONS WITH MACROMONOCYCLIC COMPOUNDS (CROWN ETHERS) ($_18 MEMBER RING) . . . . . . . . . . . . 4 COMPLEXATION 0F ALKALI METAL CATIONS WITH MACROBICYCLIC COMPOUNDS (CRYPTANDS); . . . 19 CONCLUSION . . . . . . . . . . . . . . . . 27 II. EXPERIMENTAL PART SALT AND LIGAND PURIFICATION . . . . . . . 36 SOLVENT PURIFICATION AND SAMPLE PREPARATION. . . . . . . . . . . . . . . 37 RECOVERY OF CRYPTANDS FROM CRYPTATES . . . #0 A. PREPARATION OF RESINS. . . . . . . . . 40 B. PROCEDURE . . . . . . . . . . . . . . A1 INSTRUMENTAL MEASUREMENTS AND DATA HANDLING . . . . . . . . . . . . . . . . . #2 A. NUCLEAR MAGNETIC RESONANCE . . . . . . 42 B. CALORIMETRY. . . . . . . . . . . . . . 44 C. DATA HANDLING. . . . . . . . . . . . . 49 III. NMR STUDIES OF SODIUM CATION COMPLEXES WITH SOME CROWN ETHERS IN VARIOUS SOLVENTS iv Chapter 1. INTRODUCTION. 2. RESULTS AND DISCUSSION. . . . 3. ANION EFFECTS ON THE COMPLEXATION 0F Na+ ION WITH 1806 IN TETRAHYDROFURAN AND 1,3-DIOXOLANE SOLUTIONS . . . u. SODIUM-23 NMR STUDY OF THE EXCHANGE KINETICS FOR SODIUM TETRAPHENYLBORATE COMPLEX WITH 1806 CROWN ETHER IN 1 ,3— DIOXOLANE SOLUTIONS . . . . . . . . CHAPTER IV. NMR STUDIES OF SODIUM CATION COMPLEXES WITH SOME 2—CRYPTANDS IN VARIOUS SOLVENTS 1. INTRODUCTION. . 2. RESULTS AND DISCUSSION. CHAPTER v. THERMODYNAMICS OF THE COMPLEXATION REACTION OF THE SODIUM CATION WITH SOME CROWN ETHERS IN SEVERAL NONAQUEOUS SOLVENTS 1. INTRODUCTION. . . . . . . . . . . . . . 2. RESULTS AND DISCUSSION. 3. SUMMARY APPENDICES 1. APPLICATION OF COMPUTER PROGRAM KINFIT To THE CALCULATION OF COMPLEX FORMATION CONSTANTS FROM NMR DATA . . . 2. SUBROUTINE EQN. LITERATURES CITED. . . . Page 50 51 88 98 103 103 136 137 155 156 158 159 Table 10 11 12 LIST OF TABLES Reported Thermodynamic quantities for The Complexation of Sodium Cation with 15- Crown-5, Monobenzo-15-crown-5 and 18- Crown-6 in Various Solvents at 25 OC. . . . . Reported Thermodynamic quantities for The Complexation of Sodium Cation with 2- Cryptands in Various Solvents at 25 OC. . . . Reported Kinetic quantities for The Complexa- tion of Sodium Cation with 15-Crown-5. Mono- benzo-15-crown-5 and 18-Crown—6 in Various Solvents at 25°C. . . . . . . . . . . . . . . Reported Kinetic quantities for The Complexation of Sodium Cation with 2- Cryptands in Various Solvents at 25 oC. . . . Key Solvent Properties and Correction for Magnetic Susceptibility on DA-60. . . . . . . The Physical Properties of the 23Na, 130 and Proton NUClei I I I I I I I I I I I I I I I I Sodium-23 NMR Chemical Shift-Mole Ratio data for Na+ Ion Complex with 1806 in Various Solvents at Ambient Temperatures. . . . . . . Sodium-23 NMR Chemical Shift-Mole Ratio Data for Na+ Ion Complex with 15C5 in Various Solvents at Ambient Temperatures. . . . . . . Sodium-23 NMR Chemical Shift-Mole Ratio Data for Na+ Ion Complex with MB15C5 in Various Solvents at Ambient Temperatures Formation Constants and Limiting Chemical Shifts of Na+-1806 Complex in Various Solvents at Ambient Temperatures. . . . . . Formation Co stants and Limiting Chemical Shifts of Na Ion Complexes with 15C5 in Various Solvents at Ambient Temperatures. . . Formation Constants and Limiting Chemical vi Page 28 30 32 33 39 52 57 60 63 70 71 Table Page Shifts of Na+ Ion Complexes with MB15C5 in Various Solvents at Ambient Temperatures . 72 13 Carbon-13 NMR Chemical Shift—Mole Ratio Data for Na+ Ion Complex with 1806 in Various Solvents at 31 1 1 oC . . . . . . . . . . . . 75 1h Carbon-13 NMR Chemical Shift- Mole Ratio Data for Na Ion Complex with 1505 in Various Solvents at 31 i 1 OC . . . . . . . . . . . . 77 15 Comparison of The Complexation Constants Obtained from Alkali Metal and Carbon-13 NW I I I I I I I I I I I I I I I I I I I I I 78 16 Proton Chemical Shifts of Na+-M31505 Complex in DMSO-d6 at Ambient Temperatures. . . . . . 81 17 Ionic Diameters of Alkali Cations and Ring Sizes of Some Crown Ethers. . . . . . . . . . 87 18 Comparison of The Formation Constants Currently Available for Na+ Ion Complexes with 1806, B1806 and DB1806 in Several SOlventSI I I I I I I I I I I I I I I I I I I 87 19 Sodium- -23 NMR Chemical Shift— Mole Ratio Data for Na Ion Complex with 1806 in THF Solutions with Various Anions at Ambient Temperatures. . . . . . . . . . . . . . . . . 89 20 Sodium-23 NMR Chemical Shifts for Na+ Ion Complex with 1806 in THF Solutions as A Function of Added n-BuuNI at 28 1 1 OC. . . . 91 21 Sodium- -23 NMR Chemical Shift- Mole Ratio Data for Na Ion Complex with 1806 in 1, 3- Dioxolane Solutions with Various Anions at 28 1 1 OC . . . . . . . . . . . . . . . . . . 93 + 22 Sodium-23 NMR Chemical Shifts at 1806/Na Mole Ratio of 0.5 in 1,3-Dioxolane Solutions at Various Temperatures . . . . . 96 23 The Coalescence Temperature and Approximate Exchange Rate of Na Ion Complex with 1806 in THF and 1,3-Dioxolane Solutions. . . . . . 97 vii '-able Page 23(b) The Observeda31a Chemical Shifts and Half- Height Linew ths of Solvated and 1806 Complexed Na Ion in 1,3-Dioxolane at Four Temperatures . . . . . . . . . . . . . . 100 23(c) The Observed 23Na Chemical Shifts of Solvated and 1806 Complexed Na+ Ion in 1,3-Dioxolane as A Function of Concen- tration at 28 1 1 OC. . . . . . . . . . . . 102 24 Sodium-23 NMR Chemical Shift—Mole Ratio Data for Na Ion Complex with 0211 in Various Solvents at Ambient Temperatures. . 104 25 Sodium— —23 NMR Chemical Shift- Mole Ratio Data for Na Ion Complex with 0221 in Various Solvents at Ambient Temperatures. . 109 26 Sodium- -23 NMR Chemical Shift- Mole Ratio Data for Na Ion Complex with 0222 in Various Solvents at Ambient Temperatures. . 114 27 Sodium- -23 NMR Chemical Shift- Mole Ratio Data for Na Ion Complex with 0222B in Various Solvents at Ambient Temperatures. . 117 28 Sodium—23 NMR Chemical Shifts of Solvated Na+ Ion and Na+-0221 Cryptate in Various Solvents at Ambient Temperatures. . 122 29 The 23Na Limiting Chemical Shifts of Na ~0221 Cryptate as A Function of Temperature in FY, MeZCO, THF and 2-Nitropropane Solutions. . . . . . . . . . 125 30 The 23Na Limiting Chemical Shifts of Na+° 0222 Cryptate as A Function of Temperature in DMF, MeCN, THF and 2-Nitropropane Solutions. . . . . . . . . . 128 31 The Limiting+ 23Na NMR Chemical Shifts of Na+ C211, Na 'C222 and Na+ C2223 Cryptates in Various Solvents at Ambient Temperatures. . . . . . . . . . . . . . . . 132 32 23Na NMR Chemical Shift-Mole Ratio Data for viii Table Page Na+ Ion Complex with 1505 in DMP. PI and PC Solutions at Various Temperatures. . . . . 143 33 23Na NMR Chemical Shift-Mole Ratio Data for Na+ Ion Complex with 1806 in MeCN and PC Solutions at Various Temperatures. . . . . 146 34 Formation Constants and L'miting Chemical Shifts of Na+-15C5 and Na ~1806 Complexes in Several Solvents at Various Temperatures. . . . . . . . . . . . . . . . . 147 35 Calorimetric Data for Some Sodium-Crown Complexes in DMF and DMSO Solutions at Ambient Temperatures. . . . . . . . . . . . . 150 36 Thermodynamic quantities for Some Sodium-Crown Complexes in Several Nonaqueous Solvents at Ambient Temperatures. . . . . . . . . . . . . . . . . 151 ix t“ F! U) *3 0 ”1 #11 I"! (J C3 :13 m U) Figure Page 1 Structural Formulas and Cavity Sizes of Macrocyclic Polyethers. . . . . . . . . 1 2 Structural Formulas and Cavity Sizes of Macrobicyclic Cryptands . . . . . . . . 3 The Three Conformations of 0222 Cryptand . 19 4 Calorimeter Calibration Response Curve . . 46 5 Calorimeter Response Curve for Fast Exothermic Reaction. . . . . . . . . . . . 46 6 Sodium-23 Chemical Shift Kg 1806/Na+ Mole Ratio in Various Solvents . . . . . . 53 7 Sodium—23 Chemical Shift is 15C5/Na+ Mole Ratio in Various Solvents . . . . . . 54 8 Sodium—23 Chemical Shifts Kg 15C5/Na+ Mole Ratio in Water, MeN02, and MeCN Solutions. . . . . . . . . . . . . . . . . 55 9 Sodium-23 Chemical Shift Kg ME15C5/Na+ Mole Ratio in Various Solvents . . . . . . 56 10 The Half-Height Linewidth Kg Lignad/Na+ Mole Ratio in MeNO2 Solutions. . . . . . . 67 11 Carbon-13 Chemical Shifts Kg Na+/1806 Mole Ratios in Various Solvents. . . . . . 74 12 Carbon-13 Chemical Shifts Kg Na+/1505 Figure Page Mole Ratios in Various Solvents. . . . . . 76 13 Proton (1H) Chemical Shifts Kg Na+/MB15C5 Mole Ratios in DMSO-d6 Solutions . . . . . 80 14 Sodium-23 Chemical Shift 1g Relative Mole Fraction of Solvated Sodium Ion for Na+/15C5 system in FY. DMSO and DMF Solutions. . . . . . . . . . . . . . . 83 15 Sodium-23 Chemical Shift Kg n-BuuNI/Na+ Mole Ratio in THF Solutions. . . . . . . . 90 16 Sodium-23 NMR Spectra at Various Temperatures for A Solution Contain- ing 0.05 M NaBPhu and 0.026 M 1806 in 1,3-Dioxolane . . . . . . . . . . . . . 94 17 Sodium-23 NMR Spectra at Various Temperatures for A Solution Contain- ing 0.05 M NaClOLL and 0.025 M 1806 in 1,3-Dioxolane . . . . . . . . . . . . . 95 18 Limiting Sodium-23 Chemical Shift at High, 0221/Na+, Mole Ratio 1g Temperature in Various Solvents. . . . . . 127 19 Limiting Sodium-23 Chemical Shift at High Mole Ratio, 0222/Na+, 1g Temperature in Various Solvents. . . . . . 130 20(a) Sodium-23 NMR Spectrum of 0.050 M xi a. 1L 211 CJ fily Figure 20(b) 21 22 23 24 25 NaI and 0.053 M 0222B in FY Solution. Sodium-23 NMR Spectrum of 0.050 M NaBPhu and 0.052 M 0222B in PY Solution. . . . . . . . . . . . . . . Sodium-23 Chemical Shift Kg 15C5/Na+ Mole Ratio in PY at Various Temperatures. . . . . . . . . . . . . Sodium-23 Chemical Shift Kg 15C5/Na+ Mole Ratio in DMF Solutions at Various Temperatures. . . . . . . . Sodium-23 Chemical Shift Kg 15C5/Na+ Mole Ratio in PC Solutions at Various Temperatures. . . . . . . Sodium-23 Chemical Shifts 1g 15C5/Na+ Mole Ratios in MeCN Solutions at Various Temperatures. . . . . . . . The Formation Constant 1g 103/T for Na+-1505 Complex in FY and DMF Solutions . . . . . . . . . . . . . . xii Page 133 134 139 140 141 142 148 CHAPTER I HISTORICAL REVIEW 1. Introduction Since Pedersen's first synthesis of macrocyclic polyethers (called crowns) which form stable complexes with the alkali and alkaline earth cations (l), the pro- perties of these ligands and related synthetic macrocyclic compounds and their complexes have been under active in— vestigation in many laboratories (2). Typical examples of some crown ethers are given in Figure 1. W0“, C” o a \-q\-‘;r:) (Kg_/p C) 15-Crown-5 o Monobenzo-15-Crown-5 fayity: 1.7-2.2 A (II) I _/ s. 0 0pm. E. @C. .10) Coy Cod 18-0rown-6 o Dibenzo-lB-Crown-6 cavity: 2.6-3.2 A (IV) (III) Figure 1. Structural formulas and cavity sizes of some macrocyclic polyethers. l :\ .C :- The cavity size and the properties of the crown ether can be modified by changing the number of methylene groups as well as the number and the nature of the hetero- atoms in the ring. Some of the polyethers have the useful property of solubilizing ionic compounds in organic sol- vents (l). The macrocyclic compounds not only form stable complexes with various ions but also can selectively bind certain cations in the presence of others in solution (3). Shortly after Pedersen's publications, Lehn and his co-workers introduced the diazapolyoxa-macrobicyclic ligands, termed cryptands (4), which form three dimen- sional inclusion complexes (cryptates) with various metal cations. If the size of the cation is equal to or some- what smaller than the size of the cryptand cavity, the complexed cation is essentially insulated from the medium. The cavity size and the properties of cryptands can also be varied by changing the number and type of the binding sites. The stabilities and selectivities of the cryptand complexes with suitable alkali and alkaline earth cations are several orders of magnitude larger than those of naturally occurring or synthetic monocyclic ligands (5). Some typical cryptands are shown in Figure 2. One of the remarkable properties of crown ethers and cryptands is that they can dissolve alkali metals in solvents in which the latter are normally insoluble or only slightly soluble (6, 7). Dye and co-workers first synthesizedéisalt containing an alkali metal anion, (Na+-C222)Na-, as gold colored crystals with a shiny metallic appearance by dissolving metallic sodium in ethylamine in the presence of 0222 (8). They studied the properties of various alkali metal anions in solutions and in solids by spectroscopic methods (9-11). The current status of research in this field has been recently re- viewed by Dye (12). o 0\\ it. COJ 0211 o cavity: 1.6 A (V) m C" N/\/o/\/O/\/N o O \__/ 0222 cavity: 2. 8 A (VII) /"\ /\o O K/"\_/"‘ 0221 o cavity: 2.2 A (VI) 9 o o/\7 N/\/O/\/o N C. .D \_/ 0222B (VIII) 2" Figure 2. Structural formulas and cavity sizes of some macrobicyclic cryptands. Ligands containing three and four macrocycles have also been synthesized (l3, 14). The properties and applications of the complexation reactions of these cryptands with various metal cations, anions and molecules in solids and in solutions have been extensively studied and reviewed by Lehn et a1. (15-17). Several useful and important re- view articles (3, 18-22) have been published on the studies of macrocyclic ligands and their complexes. In this thesis, only studies on the alkali metal complexes of small crowns (i lB-membered ring) and 2-cryptands will be discussed. The literature up to early 1978 on complexes of large crown ethers can be found in the Ph.D. thesis of M. Shamsipur (23). 2. Complexation of alkali metal cations with macromono- cyclic compounds-crown ethers (i 18 member ring) The macrocyclic polyethers have been found to form primarily two-dimensional, one to one (1:1) polyether- metal ion complexes with a large variety of metal ions both in solutions and in the crystalline form (1, 3). How— ever, depending on the ratio of the polyether cavity to metal ion diameter, as well as on the solvents and the counterions, complexes with other stoichiometries such as 2:1, 3:2, and 1:2 (etherzmetal ion) complexes are also formed (24, 25). The crystal structures of a number of crown complexes have been determined (26-32). It was shown that for all 1:1 complexes except the lS-crown-G-NaSCN complex, the alkali metal cation is located approximately in the center of and sometimes slightly above a planar ‘ ring of the ether oxygen atoms. In a number of cases the metal ion is also bound to the solvent molecules and/or the anions, i.e. the polyether ring could only partially replace the solvation sphere of the cation. In the 18- crown-G-NaSCN complex, the ether oxygen ring of 18-crown- 6 is strongly distorted from its symmetrical conformation to accommodate the smaller cation, i.e. one of the oxygen atoms is drawn out of the mean plane of the other five to give a somewhat irregular pentagonal pyramidal coordination of the Na+ ion (28). The 2:1 complexes, e.g. potassium ion benzo-lS-crown-S (30) and sodium ion lZ-crown-4 com- plexes (31, 32), were shown to have a "sandwich" structure in which the metal ion lies between two ligand molecules and does not interact with either the solvent molecules or the anions. The solution structures of some crown ethers and their cation complexes in water, water-acetone, acetone and chloroform were investigated by Live and Chan (33) using proton (1H) and carbon (13C) nuclear magnetic resonance. The complexes of benzo-lB-crown-G and dibenzo- 18-crown-6 (DB1806) with Na+ and K+ ions were shown to have the same structure in the various solvents as do the DBlBCG complexes in the crystalline state. The "sandwich" complex with a Cs+ ion between two DBlSCG molecules was found to exist in acetone and chloroform solutions. In ’(1 ( ') DI The formation of complexes between the cyclic poly- ethers and alkali metal salts has been studied by a number of different techniques: solubility changes, extraction, potentiometry, calorimetric titration, polarographic re- duction, conductance measurement, UV, IR and NMR spectro- sc0py (2). Frensdorff (34) determined the complex stability constants of a variety of crown ethers with several cations in aqueous and methanolic solutions potentiometrically with cation selective electrodes. The stability constants were found to be three to four decades higher in methanol than in water, since methanol is a much weaker solvating medium and thus competes less with the polyether for the cations. Selectivity toward different cations varies with polyether ring size, the optimum ring size being such that the cation just fits into the hole. Comparison of the stabilities of the Na+, K+ and Cs+ ions each with the crown of optimum ring size in methanol reveals that the K+ complex is the most stable by at least one order of magnitude. This observation seems to reflect the competi- tion between complex formation and solvation. The largest cation has low stability constant because it has low charge density and thus does not attract the polyether very much, while the smallest cation is too strongly solvated for the polyether to compete successfully for the cation. It was noted that the substitution of nitrogen or sulfur atoms for oxygen in the polyether ring reduces the stabilities of potassium complexes but greatly strengthens those of the Ag+ ion complexes which involve covalent bonding. The K+ ion complex stability constants decrease in the order, 0 > NR > NH > S. Izatt et a1. (35-38) determined the formation constants and thermodynamic parameters (AH° and 85°) for the interaction of the two isomers of dicyclohexano-18- crown-6 with several uni- and divalent cations in water by using a calorimetric titration technique. The stability order for the alkali metal ions with either isomer is essentially the same as the permeability order for these same metal ions with the structurally related antibiotics, + + + . > L1 , valinomycin and monactin, i.e. K > Rb+ > Cs+, Na for the tranSport of these ions through natural and synthetic membranes. The behavior of these two isomers in terms of log K, AH°, AS° and the selectivity toward the alkali and alkaline earth cations for the complexa- tion reaction in aqueous solution is different due to the structural differences of the isomers. By using electrical conductance measurements, Evans et a1. (39) determined the formation constants of dibenzo- l8-crown-6 and dicyclohexano-lB-crown-6 complexes with sodium, potassium and cesium salts in methanol and acetoni- trile solutions. These authors found that the association constants of Na+-DBC in methanol and acetonitrile and K+o .DBC in acetonitrile are nearly independent of temperature fg .r~ Y? a» nu . Ts in the range 10-25°C. The fact that the sodium ion is more strongly complexed in acetonitrile than in methanol was explained by its weaker solvation in the former sol— vent. Similar results were also obtained by Koryta et a1. (40, 41) from polarographic studies of alkali metal cations complexed with crown ethers in methanol and acetronitrile solutions; the stability constants of the sodium complexes with some 18- and 24-membered crown ethers in acetonitrile are comparable or higher than the stability constants of the potassium complexes with the same crowns. This be- havior was attributed to the higher surface charge density of the sodium ion and, therefore, a stronger dipole-ion bond of the ion with the ligand. The difference in the solvation energies of Na+ and K+ ions in the weaker sol- vating medium of acetonitrile also becomes less signifi- cant. It was noted that the sodium ion forms more stable complexes with 18-membered crowns than with the 15-membered ones. The complexation constants of sodium, potassium and cesium salts with dicyclohexano-lB-crown-G were measured polarographically in methanol, ethanol and n- propanol by Agostiano et a1. (42). For each ion the equilibrium constant of the complex was found to increase in the order of methanol < ethanol < n-propanol. This was attributed to the decrease in the dipole moment of the solvent molecule from methanol to n-propanol by the authors. cg Stability constants for complex formation of dibenzo- 18-crown-6 with a series of metal ions in aqueous solutions were determined spectrophotometrically by Shchori et al. (43). They noted that the concentration formation constants of the complexes depend on the counter ion and on the ionic strength of the solution. The thermodynamic con- stants were obtained by extrapolating the experimental values at several ionic strengths to infinite dilution. A similar dependence of the concentration formation con- stants (Kc) on the ionic strength (I) of the solution was found by Smetana and Popov (44) for the 18-crown-6 sodium salt complex in anhydrous methanol solutions. It was also shown that the Kc values remained reasonably constant for I i 0.05 mol dm‘3. The effect of substituents in the aromatic ring on the complexation of sodium cation with dibenzo-lB-crown-G (DBC), cis-4, 4'-dinitrodibenzo-l8-crown-6 (NDBC) and cis-4, 4'-diaminodibenzo-l8-crown-6 (AmDBC) in dimethylformamide solution was investigated conductometrically by Shchori and Grodzinski (45). The order of the equilibrium constant was found to be as follows: KNDBC < KDEC ; KAmDBC The sodium ion complex of the dinitro derivative NDBC, which contains the electron-withdrawing -NO2 groups, is five times less stable than that of DBC. The AmDBC contains two (h .rh .C A: Cy :1 .3 C 10 strongly electron-donating -NH2 groups and thus is expected to form more stable complex with sodium ion than DBC. The nearly identical complex stabilities of sodium ion with DBC and with AmDBC was attributed to possible conforma- tional changes in the macrocyclic ring due to hydrogen bonding involving the two amino groups. Substituent effects on the stabilities of sodium and potassium cation complexes with a series of 4'-sub- stituted monobenzo-lS-crown-S and monobenzo-l8-crown-6 ethers were also determined conductometrically in acetone solution by Ungaro et a1. (46). The substituent effect for the complexes of Na+ ion with benzo-lS-crown-S is very pronounced, e.g. a nearly 25-fold difference in K between the 4'-amino- and 4'-nitrobenzo-15-crown-S, while the effect on the complexes of Na+/benzo-18-crown-6 is much smaller, and almost negligible for the electron-with- drawing substituents. The above results were explained by the possible conformational changes in the polyether ring of Na+/benzo-18-crown-6 complexes on varying the 4'- substituents. The effects are somewhat larger for K+/ benzo-lB-crown-6 complexes than for Na+/benzo-18-crown-6 complexes. Pannell et a1. (47) studied substituent effects on the selectivity of dibenzo-l8-crown-6 vis-a-vis sodium and potassium salts by extracting the salts in water with the crown ether in methylene chloride solution. The results 11 showed: (a) a nearly inverse relationship between the cumulative electron-withdrawing power of the substituents and the ability of the crown ether to extract the sodium and potassium salts due to the decreased basicity of the oxygen crown system, and, (b) crown ethers with cumula- tively large electron-withdrawing substituents exhibit a reversal of the "normal” ion extraction selectivity, K+ > Na+, of 18-crown-6 ethers. The reversal of the K+ - Rb+ selectivity of 18-crown-6 in methanol was also noted by Curtis et a1. (48) for a tetra-O-isopropylidene substituted 18-crown-6. Recently, Izatt et a1. (49, 50, 51) have extended their calorimetric titration studies of the interaction of several mono- and divalent cations with dicyclohexano- 18-crown-6 to lS-crown-S, 18-crown-6, thia-lB-crown-6 and benzo-lS-crown-S in aqueous, methanol and water-methanol solutions. The marked selectivities toward uni- and bi- valent cations shown by l8-crown-6 were not found with 15-crown-S which shows nearly identical log K values for all the alkali cations in water and shows very little selectivity in methanol solutions. The enthalpy change shows the same order as the change in log K for the complexation of alkali cations with each of these 18- membered crowns in water and with l8-crown-6 in methanol, while the entropy change shows a nearly opposite order of log K values. Thus AH° and AS° compensate each other with 12 AH° being the dominant quantity in determining the magnitude of log K. A change in the solvent composition from water to water-methanol mixtures results in different cation selectivity patterns for lS-crown-S and 18-crown-6. It was also shown that the substitution of one sulfur atom for oxygen in 18-crown-6, to give thia-18-crown-6, con- + and Ba2+ siderably reduces the stabilities of the Na+, K complexes in methanol solution due largely to the less favorable AH° values. With only a few exceptions, both the enthalphy and entropy changes of the above macrocyclic complexation re- actions are negative. These data clearly show that the origin of the macrocyclic effect might be different from that of the chelate effect. Cabbiness and Margerum (52) were the first to use the term "macrocyclic effect" to distinguish it fromifluechelate effect, since there is an additional enhancement in the stabilities of macrocyclic complexes beyond that of chelates. Since that time, the macrocyclic effect has been studied by many authors mostly on the complexation of transition metal ions with the macrocycles and their linear counterparts. Based on a detailed study of the thermodynamic properties of nickel (II)-tetramine complexes in water, Hinz and Margerum (53) concluded that the macrocyclic effect is almost entirely enthalpic in origin. The large negative enthalpy changes were attributed to the decreased l3 solvation of the macrocyclic ligand which has fewer hydrogen bonds with water molecules to be broken in complex forma- tion. The smaller loss in configurational entropy upon complexation of the macrocyclic ligand tends to be offset by less water being released from the solvated ligand. In solvents with weak or no hydrogen bonding, changes in con- figurational entropy of the ligand should be more important. In contrast, Kodama and Kimura (54) reported a polaroqraphic study on the thermodynamics of Cu (II)- tetramine complexes in water and stated that the macrocyclic effect is due solely to the entropy gain. Shchori and Grodzinski (45) studied the complexation of dibenzo-l8- crown-6 with a sodium salt at various temperatures by electrical conductance measurements and showed that the complexation reaction is both enthalpy and entrOpy driven in dimethoxyethane solution. 0n the basis of the results obtained from a microcalorimetric study of copper (II) and zinc (II)-tetramine complexes and the corresponding linear counterparts in aqueous solution, Anichini et al. (55) summarized that the macrocyclic effect is due to a favor- able entropy term and to a normally favorable enthalpy term. The magnitude of the latter is critically dependent on matching the size of the metal cation to that of the macrocyclic ligand. The thermodynamics of complexation reaction between the various cyclic polyethers and cations have been studied l4 extensively by Izatt et a1. (49, 50, 51). However, their data contain no trends in AH° and AS° among the complexes studied which would explain the macrocyclic effect. It is clear from the above discussion that although the chelate effect is definitely of entropic origin, no agreement has been reached as to whether the macrocyclic effect results from a more favorable enthalpy or more favorable entrOpy term in the complexation free energy. Mei et a1. (56, 57) studied the complexation of the cesium cation with 18-crown-6 (18C6), dibenzo-lB- crown-6 (DBC) and dicyclohexano-lB-crown-G (DCC) in various solvents by cesium-133 NMR technique. In all sol- vents studied, 18-crown-6 forms 1:1 and 2:1 (ligand to metal) complexes. In most solvents the formation of 2:1 complexes by the other two crown ethers was also in- dicated. The chemical shift of the 2:1 complex of Cs+ with 18-crown-6 is essentially independent of solvent, indicating that in the "sandwich" complex the cesium ion is effectively shielded from the solvent. In all solvents the one to one complexation constants (Kl) decrease in the order 18C6 1 DCC >> DBC while the two to one com- plexation constants (K2) follow the order DEC 3 1806 >> DCC. It was shown that there is no simple relationship between K1 values and either the dielectric constant or the 15 donicity* (58, 59) of the solvent, although the solvent plays an important role in the complexation process. The kinetic parameters for the release of Cs+ ion from some macrocyclic complexes in several solvents were also re- ported. The complexing ability of 1,10-diaza-18-crown-6 (DA18C6) with Tl+ and alkali metal cations in several non- aqueous solvents was recently studied in detail by Shamsipur (23). For the Cs+-DA18C6 complex in the sol- vents studied, there is no clear indication of the forma- tion of 2:1 (ligandzmetal) complex as found for other crown ethers by Mei et al. (56, 57). The Na+ and Cs+ ion complexes of DA18C6 were found to be less stable than those of 18C6 (57, 60). Similar results were reported by Frensdorff for the K+ ion complexes with these two crowns in methanol solution (34). The formation constants of Li+-DA18C6 in various solvents are unexpectedly large and greater than those of the Li+-18C6 complex (61) in the same solvents. This difference was attributed to possible covalent bonding between the Li+ ion and DAl8C6. With the exception of pyridine, the stabilities of the Li+, Na+ and *The Gutmann donor number is the positive enthalpy value ( in kcal/mole) of 1:1 complex formation between the given solvent molecule and antimony pentachloride in 1,2- dichloroethane solution: 1,2-DCE\ I S + SbCl5 S-SbCl5 -A donor number H S-SbClS Gutmann used the term "donicity" when referring to the dxonor ability of a solvent. 16 Cs+ ion complexes with DA18C6 increase roughly with de- creasing donicity of the solvent. Kinetic information about the complexation re- actions between crown ethers and alkali metal ions is very limited. The kinetics of complexation of sodium ion with dicyclohexano-l8-crown-6, dibenzo-lB-crown-6 (DBC) and its derivatives in several nonaqueous solvents were investigated by Shchori et a1. (62, 63) using sodium-23 NMR. They postulated that the dominant exchange mechanism involves the decomplexation step: + - , + - Na (X ), crown.‘_____ Na (X ) + crown and showed that the rate of decomplexation is very strongly affected by substituents in the aromatic rings of DBC. It was found that the rate constants of complexa- tion increase with an increase in the stability of the complex, while the rate constants of decomplexation vary in the Opposite direction. It was also shown that the effect of solvent on the equilibrium constants for complexa- tion is very pronounced, while the activation energy of the decomplexation of DBC and of its derivatives is in- dependent of solvent (to within i 1 kcal). Because of this constant activation energy and the coincidence of this value with the energy of activation of a conformational change of the macrocyclic ring, these authors speculated that the energy barrier for decomplexation may be l7 determined by the energy required for a conformation re- arrangement of the complex. Similarly, the kinetics of complexation of potassium and rubidium ions with dibenzo- 18-crown-6 (DBC) in methanol were studied by Shporer and Luz (64) using potassium-39 and rubidium-87 NMR techniques respectively. The rate constant and activation energy of the decomplexationof K+-DBC are similar to those of Na+-DBC complex (62, 63). But the dissociation rate con- stant of the Rb+-DBC complex is much larger possibly be- cause of the weaker binding between the larger rubidium ion and the DEC molecule. Kinetics studies of the conformational equilibrium of lB-crown-6, and its complexation with Li+, Na+, K+, Rb+, + Cs , Tl+ + 4 I . . . + + + + + . S and Its complexation with Na , K , Rb , T1 and Ag ions , Ag+, NH and Ca2+ ions as well as of lS-crown- in aqueous solution have been carried out using ultrasonic absorption method by Eyring et a1. (65-68). According to these authors, the basic reaction scheme for the complexa- tion which best fits their experimental data is the two- step process proposed by Chock (69) k CR1 12 ‘ CR2 (1) ‘ k 21 + k23 + M + CR ‘ MCR (ii) 2 X k 2 32 where CR1 and CR2 denote the two different conformations of the polyether. The conformation CRl is unreactive, M+ 18 is a metal ion and MCRZ+ is the complex ion. The mechanism involves a fast ligand conformational change followed by a stepwise substitution of the coordinated solvent molecules by the ligand. The equilibrium constant for the rapid conformational rearrangement of the ligand, equation (i), was determined - - -2 - - to be K21 - k21/k12 - (2 i 2) x 10 for 18 crown 6 and was estimated to be K21 : 0.1 for lS-crown-S. This means that most of the free ligands are in the CR form. The 2 variations 4). The chemical shift of the lithium cation complexed by C211 is essentially independent of the solvent and of the counterion used, indicating that the lithium ion is completely insulated from the external medium, i.e. it forms an inclusive complex. Hourdakis and Popov (79) studied the complexation of Li+, Na+ and Cs+ ions with cryptand C222-dilactam (X) (which is the product of the penultimate step in the synthesis of cryptand C222) in various solvents by lithium- 7, sodium-23, cesium-133 NMR and far infrared measurements. The complexing ability of the dilactam is similar to, but weaker than that of the crytand 0222. The formation con- stants of 0222-dilactam complexes with Li+, Na+ and Cs+ ions in several nonaqueous solvents were reported. Och—Km ‘N”\VO/\\/0/\VNO O/\V/’W\b_/ (X) In order to have a better understanding of the thermodynamics of cryptate formation, Kauffmann et a1. (80) obtained the enthaplies (ARE) and entropies (ASS) of 23 ‘complexation of alkali and alkaline earth cryptates from calorimetric measurements of AH; and the previously deter- mined stability constants. It was found that both enthalpy and entropy changes play an important role in the stability and selectivity of the complex, but the cryptate effect and the selectivity peaks observed in the stability con- stants of the cryptates are of enthalpic origin. The cavity-radius/cation-radius effect used as an empirical criterion for discussing the selectivity of cryptand com- plexation (73) was shown to incorporate both enthalpic and entropic effects and is not just a measure of the steric fit. They also reported that the marked increase in stability on transfer from water to methanol/water mixture is entirely due to an increase in enthalpy of complexation in the mixed solvent, which was then explained by them as an increased electrostatic interaction of the cation with the ligand in the medium of lower dielectric constant and the smaller interaction of the cation with the solvent. Abraham et al. (81), however, report that for both cryptates, Na+-C222 and K+-C222, the enthalpically favored complexation in methanol solution is entirely due to the effect of solvent on the free cryptand C222. These authors obtained an extraordinarily large enthalpy of transfer, AH; a +13.9 kcal-mol-1, from water to methanol for the ligand. This means that the free cryptand is much more strongly solvated in aqueous solution than in methanol. 24 On the other hand, the solvent effect on M+ and M+-C222 in terms of enthalpy both disfavor complex formation in methanol, i.e. the M+ ions are more solvated in methanol while the M+-C222 ions are more solvated in water. Mei et a1. (82, 83) investigated the complexation reaction of cesium cation with 2-cryptands in various sol- vents by cesium-133 NMR. It was found that both the dielectric constants and the solvating abilities of sol- vents influence the complexation reactions. The relative stabilities of the complexes were shown to be in the order of Cs+-0221 _>_ Cs+-C222 > Cs+-C2228 >> Cs+-0211. The Cs+-C222 complex was found to be enthalpy stabilized, but entropy destabilized in all the nonaqueous solvents studied. It was noted that there is a temperature- and solvent-dependent equilibrium between "exclusive" and "inclusive" conformations of the Cs+-C222 complex, with the inclusive complex (in which the Cs+ ion is located in the center of the cavity) favored at lower temperatures. In the exclusive complex, the cation is not completely within the cavity and so it is only partially insulated from the external medium. Thermodynamic parameters (AHi and A83) for the formation of the exclusive complex and for its conversion to the inclusive complex (ARE and A33) in acetone, propylene carbonate and N, N-dimethylformamide solutions were reported. The possibility of the existence of both exclusive and inclusive Cs+-C222 complex in water 25 was shown recently by Desrosiers and Morel (84) from the determination of the standard volumes of complexation of the alkali chlorides with cryptand C222. The kinetics of cryptate formation was first studied by Lehn et al. (85) by temperature dependent pro- ton NMR in D20 solutions. The authors assumed that the exchange mechanism proceeds by a dissociation-complexation process rather than a bimolecular process and the exchange becomes slower as the stability of the cryptate increases, as was found for the metal cation-crown ether complexes (63). Cahen et a1. (86) investigated the kinetics of the complexation reaction of the lithium cation with cryptand C211 in water and in several nonaqueous solvents by lithium-7 NMR. The energy of activation for the re- lease of Li+ from the Li+-C211 complex was found to in- crease with increasing donicity of the solvent. By con- trast, Shchori etflal.(63) found that the activation energy for release of Na+ from the complexes of DBO and its derivatives is independent of the solvents used. However, two of the three solvents used, methanol and dimethylformamide, have the same donicity while that of the third solvent, dimethoxyethane, is not known. Cahen et a1. pointed out that the transition state for cryptate formation appears to be on the reagent side, i.e. it involves substantial solvation of the cation. Activation energies (Ea), dis- sociation rate constants, and values of ARE, ASE, and AG: for the release of Li+ from the cryptate were reported. 26 Sodium-23 NMR kinetics studies on C222-cryptate in four solvents were performed by Ceraso et a1. (87, 88). It was shown that the exchange proceeds through a dis- sociation-association process in ethylenediamine solution. The activation energies were found to depend upon the sol- vents used. However, the correlation between the activa- tion energy and the donicity of the solvent reported by Cahen et al. for Li+-cryptate decomplexation (86) was not observed. The rate constants and values of AG:, AH: and as: for the decomplexation reaction of Na+-C222 cryptate were also reported. Cox and Schneider (89) measured the dissociation rates of metal cryptates in aqueous solution using a con- ductance method. It was found that in some cases, e.g. the Li+-Cle and Na+-C221 cryptates, the dissociation of the metal-cryptand complex is acid catalyzed, i.e. the observed rate constants increase as the acid concentrations increase. A mechanism was suggested for the dissociation reaction in which there is a preceding conformational change from in-in (Mn+) to in-out (Mn+) or out-out (Mn+) complex that leaves one of the lone pairs on nitrogen free to be trapped by H+, prior to the dissociation of the metal ion from the cryptate. Cox et a1. (90) have also studied the kinetics and thermodynamics of some alkali metal complexes with three cryptands, 0211, 0221 and 0222, in methanol solutions by electrical conductance measurements and 27 potentiometric titrations. It was shown that the pro- nounced selectivity of the ligand is reflected entirely in the dissociation rates, with the formation rates which are about 108 larger than the dissociation rates increasing monotonically with increasing cation size. These results strongly suggest that the transition state for the forma- tion reaction lies very close to the reactants as was found in the lithium-7 NMR kinetics studies (86). In the transition stage there is no specific interaction between the cryptands and the cations that strongly differentiates between the various cations. The specific size-dependent interactions between the metal ions and the cryptands must then occur subsequent to the formation of the transition state. The quantitative data currently available on sodium salt complexes of lS-crown-S, benzo-lS-crown-S, lB-crown-6, and 2-cryptands are presented in Tables 1-4. 4. Conclusion From the previous discussions, the parameters which influence the stabilities and selectivities of macrocyclic complexes may be summarized as follows: (a) the relative sizes of cation and ligand cavity, (b) the type and number of donor atoms in the ring, (c) substitution on the macrocyclic ring, (d) type and charge of cation, (e) sol- vent properties. Most of the investigations in solution were done in water, methanol and methanol-water mixtures, while studies in other nonaqueous solutions are quite sparse. 28 am Add mo.o H ~m.o as Hoo e.mu ca.o H mm.~- oa.o H om.o em you m.o v nouns outsouolma on and c.eau mo.o H mm.mu ~o.o H c~.~ moo: mom on and s.mu so.o H ~m.mn oH.o H aa.a mom: was om Hoo ~.mu mo.o H ms.m- so.o H ec.a mom: was om goo m.m- HH.o H mc.~: NH.o H AH.H moo: mos om moo c.~: No.o H se.an mo.o H Ne.o row: row om Hmo o 2 cv.o noun: He mace nmm.e zoo: we coo oo.o H mm.m occumoa muszowonmamz Hm ado om.ou mo.o H ma.e- Ho.o H me.m row: as nod mo.o H no.0 as Hoo m.~- so.o H om.~- oa.o H os.o hours muczononma .wmm ooosuoz Ammo Hmw\aooc AHoE\mmoxv HWImmm .HmmmHmm .mmmmmm . . 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H Canoe 30 om om om mm mm om mm mm mm om mm mm mm a me .wwm some mmu< Ho> Hoao uom pom Ho> Ho>o Hmo a you Hon 2 Hod Hon Hoe uom ate a mom uom Dom ate a you moored: m H ~.o+ N + ml Ammo HOE\HmoV omc .N.o + mm.m| «.0 H «.m- AHOE\HmoxV o=< comm um muco>~0m msoHum> :H mpcmumxuolm coHumo EQHOOm mo coHumonmEoo may new MOHHHHBMSU oHEmc>OOEumnu pouuommm w n x mos Headrooz umumz $002 wmm Honored: noun: :Omz mmm Honorees Hmumz ucw>aom NNNU HNNU Haao commas .N CHQMB 31 .mwcmp one opHmpso MH x on once esp Homnma mosooon nomno one can .mve onvH so.“ To H op OMHooLQ ohm Aow H mm .wom Eoewv mosam> omen as com as Don sh Ho> Hoao om a me Soc a mod N H m.~: ~.o H c.oas .mom moored: Ammo Hos\amoe Aaoe\aooxv one one xuumaoucmmufl MH omud auuoEmuHo> UHHo>o MH Ho> Home anuoEHuono MH Hmo xuuoEoHucmuom MH you s.e zoo: wmm o.q nouns mmmmo e.m omzo H~.e moo: mmm we mom Deo>aom panama Aconcaocoo. m canoe 32 co .mno .woas sea x v.m so .mnc .AOH: hoe x m.e .mom moored: land. or coHumu0mnm oHcommHuH: oH x m.m z Hmumz oHanHm>m camp 02 OH x v.~ z umumz . m :2 oco>Hom ma .mn< .ODHDm m|c30HOImH mnczouolmamz mICBOHOImH pcmqu oomm um muco>H0m meoHum> :H mlc3ouolmH cam mlczouolmauoNcmnocoe .m nczououma :DH3 :oHumo ESHOOM mo coHumonQEoo one “Om MOHHHHCMSG oHuocHx cmuuomom .m wanna 33 mm mm mm om mm om Hm mm om Hm .wom mZZ mZZ mZZ pcoo mzz o o op cu z coHuoonoH mean ousumnomfiou MH .m.n.e .oocmuozczoo MH ccoo cceemaoocoaseum .wmm H no >uchuHoocs cm o>ms mosam> omone .mm + oxa.eco.mma Ols~.ovmo.m olmo.ocea.e mace x e.~ osm.~ sea x ~.H ole.~ve.eva nmoa x s.a owned x mm.~ ooa x 1m H as m H as sea x o.m m.ea noon x H.m oom.~ «OH x AA H ac Noe x 1H.o H v.HV m m 1a- a-sc In. t or ex comm um muco>H0m msoHuo> :H mpcmum>HOIm nuHB coHumo E=HOOM mo coHuoondeoo on» “cm moHqucmso oHuocHx couuomom 0 «on .cmusmowowcmuuoa u one Q .coHuoH>o© Unopcoumo ©m Hocmnuoz nouns Hocosuoz woumz soonest: noun: uco>Hom Q «o huchuHoocs cm o>mn mozao> omosem NNNU ammo HHNU tcmqu .v oanoe 34 Especially for sodium cation complexes with lS-crown-S, monobenzo-lS-crown-S, l8-crown-6 and 2-cryptands, very few data exist on the thermodynamics and kinetics of the com- plexation reaction in nonaqueous solutions as can be seen from Tables 1-4. The use of macrocyclic compounds in organic synthesis, electrochemistry and analytical separations has been demonstrated. Sam and Simmons (95) showed that the KMnO4 complex of dicyclohexano-l8-crown-6 oxidized olefins, alcohols, aldehydes and alkylbenzenes under mild conditions in benzene. In the absence of the crown ether, KMnO ‘has no detectable solubility in benzene and no re- 4 action occurs with organic substrates. Valinomycin has been used as a membrane component in a potassium-selective liquid-membrane electrode (96). The above electrode could make measurements of potassium ion activities in the range of l M to at least 10.6 M in unbuffered aqueous solu- tions with a selectivity of potassium over sodium cations of more than 4,000. It has been shown that crown ethers and cryptands may serve as model compounds for investigat- ing and understandingtflmaphenomenon of ion transport through cellular membranes (2, 22, 35, 36, 97). Numerous other actual and possible applications of the macrocyclic com- pounds in synthetic organic, electro-chemistry, separations, biology and drugs have been discussed in references 2, 3, 18, 19, 21, and 22. 35 Many uses of macrocyclic compounds depend on their marked ion specificities. Consequently, a knowledge of log K values for the interaction between these compounds and ions becomes very important in predicting their be- havior and determining their uses. Therefore, this dis- sertation reports an investigation of the thermodynamic properties of the complexation reaction between sodium cation and several crown ethers and cryptands in aqueous and various nonaqueous solutions by proton, carbon-l3, and sodium-23 NMR as well as calorimetric techniques. CHAPTER II EXPERIMENTAL PART 1. Salt and Ligand Purification Sodium salts used in this study were reagent grade quality and not further purified before use except for drying. Sodium tetraphenylborate (J.T. Baker) and sodium thiocyanate (Mallinkrodt) were dried under vacuum at 50°C for 3 days. Sodium perchlorate, iodide (Matheson, Coleman and Bell) and chloride (J.T. Baker) were dried at 110°C for 3 days. Tetrabutylammonium iodide (Aldrich) was re- crystallized first from water and then from a 9:1 ethyl- acetate : 95% ethanol mixture. The purified tetra- butylammonium iodide (TBAI) was dried under vacuum at 50°C for 3 days. The macrocyclic polyether 18-crown-6 (18C6, Aldrich) was twice recrystallized from acetonitrile, dried under vacuum at i 40°C for three hours and then at room temperature for three days. The dried 1806 melted at 37-38°C (lit. mp 36.5-38.0°C (98), 39-40°C (1)). Macrocyclic 15-crown-5 (15C5, Aldrich) is a liquid and was purified by vacuum distillation and dried under vacuum over Drierite at room temperature for three days. The dried 15C5 in CCl solvent has a large singlet proton NMR 4 peak at 6 3.53 ppm (lit. 6 3.58 ppm (49)). Dibenzo-18C6 (D818C6, Parish) was recrystallized twice from benzene and dried under vaccum over Drierite at room temperature for 36 37 three days; mp 165-166°C (lit. mp 164°C (1)). Monobenzo- 15C5 (MBlSCS) was synthesized according to Pedersen's method (1) by M. Shamsipur in this laboratory. It was re- crystallized from n-heptane twice and dried under vacuum over Drierite at room temperature for three days; mp 79-80°C (lit. mp 79-79.5°C (1)). Cryptand C222 (E.M. Laboratories, Inc.) was re- crystallized twice from Hexane and dried under vacuum over Drierite at room temperature for three days; mp 69°C (lit. mp 68-69°C (99)). 0211, 0221 and 02228 (E.M. Laboratories, Inc.) are liquids and were used without further purifica- tion except drying under vacuum over Drierite at room temperature for three days. Potassium hexafluorophosphate (Pflaltz and Bauer) was recrystallized from water and dried for several days at 110°C under vacuum. 2. Solvent Purification and Sample Preparation Tetrahydrofuran (THF) was refluxed over metallic potassium and benzophenone for 24 hours and then fractionally distilled using a 30 cm Vigreux column. Acetone (MeZCO) was refluxed over Drierite for 24 hours and then frac- tionally distilled. 2-Nitropropane was refluxed over Drierite for 24 hours and then fractionally distilled under reduced pressure. Acetonitrile (MeCN) and 1,3-dioxolane were refluxed over calcium hydride for 24 hours and then fractionally distilled. Tetramethylguanidine (TMG) was 38 refluxed for 24 hours under reduced pressure and then fractionally distilled. Nitromethane (MeNOZ), pyridine (PY), propylene carbonate (PC), dimethylsulfoxide (DMSO) and N,N-dimethylformamide (DMF) were refluxed over calcium hydride for 24 hours under reduced pressure and then fractionally distilled. Before use all the distilled sol- vents were further dried overnight over freshly activated 4A or 3A molecular sieves.. The water content of all sol- vents, except acetone, was measured with Karl Fischer automatic titrator (Photovolt Aquatest II) and was found to be always less than 100 ppm. The water content of acetone was less than 100 ppm as measured by gas chromato- graphy analysis. Deuterated DMSO-d6 (Stohler Isotope Chemicals) was used as received. Useful solvent properties are listed in Table 5. The molecular sieves used were washed with dis- tilled water and dried at 100°C overnight. The dried molecular sieves were further heated at 500°C under nitroqen for 24 hours. All solutions were prepared in a dry box under nitrogen atmosphere. Ligands were quickly weighed out in the desired amount into a 2 ml volumetric flask and then transferred to the dry box for subsequent manipulation. Each sample solution was prepared by mixing appropriate amounts of stock salt solution, ligand and solvent. 39 .AONHV .Mom* #23. .Aooav mzz ozmm so ooeoaoona ** .Asm .mmv .eoma --- --- aaaa.sa.< as.a --- ocoaoxoaa-m.a Nsm.o- cam.o nu- nu- o.HH ocapacmsmaseooEmwuoe ~44.o- mom.o us- me.m m.m~ ocoaowdonuazum mmo.o- Ham.o s.m ae.m a.mm macsuoeonuaz oam.o- emm.o H.4a oa.m m.em oeawoacooooa omH.on «mo.o H.ma m.m o.mo oomconnoo occasaona mem.o- ooe.o o.aa mm.~ s.o~ onerous e-.o- mam.c o.o~ mc.a c.e spasmonoxronooe ame.o- mam.o «as.m~ oe.a e.mm Hococooz mom.ou mum.o c.o~ cm.m e.om moneoEHomassooEAp-z.z He~.o- moc.o m.mm a.m s.oe moaxouasmflsnuoEHQ ooo.o oms.o «clo.mm. o.mH em.H m.me nous: cN~.o- «no.6 H.mm m~.~ v.~a ocfloaasa Aammv omu « censusw oHOQHQ oHuuooHoHo condo co quHHnHumoomsm oHuocmmz MOM coHuoouuoo can ooHuquoum uco>Hom hox .m oanoa 40 3. Recovery of Cryptands from Cryptates A. Preparation of Resins Cation-exchange resin Dowex 50Wx8 (100-200 mesh) in the H+ form was purified by soaking in warm (60°C) 4 M hydrochloric acid with stirring for three hours, followed by several washes with distilled water. The product was further purified by carrying through the same procedure once again. The resin was then washed with distilled water by suction using a dispersion tube until the pH of the filtrate was about 4. Anion-exchange resin Dowex 1x2 (50-100 or 100-200 mesh) in the Cl- form was purified by stirring in a beaker with 4 Mhydrochloric acid for three hours, and washing several times with distilled water. The same procedure was likewise repeated once again. The resin was then washed with distilled water until the pH of the filtrate was about 7. The purified resin was stirred in a beaker with 50% v/v water and acetonitrile mixture before trans- ferred to the column. Anion-exchange resin Dowex 1x2 (SO-100, or 100- 200 mesh) in the Cl- form was converted into the OH- from by stirring the chloride form of the resin in a beaker with 3 M NaOH for three hours, followed by several washes with distilled water. The same procedure was likewise re- peated once again. Then the resin was washed with deionized water until the pH of the filtrate was 7. 41 B. Procedure Shih et al. (101) developed a recovery process for cryptands from used solution mixtures of the complexes. The procedure was followed in this work apart from minor modifications which are given below. (i) To remove a large quantity of solvent with a regular vacuum pump is quite inconvenient because the trap of the pump is very quickly plugged with frozen solvent. Instead 2 torr (101), solutions of being dried in vacuo at ~ 10- of metal cryptates in various solvents were dried with a rotavapor to remove the bulk solvent mixtures and then the last traces of solvents were removed in vacuum at ~ 10-2 torr. (ii) Shih et al. (101) pointed out that if the original solution mixtures contained a variety of anions, it was useful to convert the salts to the chloride form by passing the solution through an anion exchange column in the Cl- form. However, they did not give details of the pro- cedure which is described as follows: After the metal ion cryptates were converted to the diprotonated cryptand salt by addition of excess HCl, the solid (~ 0.5 g) was dissolved in a 50% v/v aqueous acetonitrile mixture (~ 5-10 ml). The solution was then passed through the anion exchange column (1.5 X 25 cm) in the Cl- form. The chloride salt of the metal cation and the diprotonated cryptand were eluted with about 40 m1 of the same aqueous -42 acetonitrile mixture at a flow rate of about one ml per minute. 4. Instrumental Measurements and Data Handling A. Nuclear Magnetic Resonance Most of the sodium-23 NMR measurements were obtained by using a Varian DA-60 spectrometer operating at a field of 1.409 Tesla and a frequency of 15.871 MHz in the pulsed Fourier transform mode. The spectrometer is equipped with a wide-band probe capable of multinuclear operation (102) and with a home-built lock probe (103) which uses the DA-60 console to lock the magnetic field on an external proton resonance. A Fabri 1080 computer was employed to carry out the time averaging of spectra and the Fourier trans- formation of the data. Some sodium-23 NMR data were ob- tained with a Bruker HEX-90 spectrometer Operating at a field of 2.114 Tesla and a frequency of 23.81 MHz in the pulsed Fourier transform mode. A Nicolet 1080 computer was employed to carry out the time averaging of spectra and the Fourier transformation Of the data. A 3.0 M aqueous sodium chloride solution was used as an external reference and the reported sodium-23 chemical shifts are referred to this 3.0 M aqueous NaCl solution. Ten mm o.d. precision NMR tubes were used. The reported data were also corrected for the differences in bulk diamagnetic susceptibilities between sample and reference solvents according to the following equation 43 (104) for a cylindrical sample in a spectrometer with a perpendicular magnetic field. . _ - ref. _ sample Ocorr. _ Gobs. 2H/3 (xv Xv ) (2'1) where xief' and Xjample are the volume susceptibility (105) of the reference and sample solvents respectively and dobs. and acorr. are the Observed and corrected chemical shifts respectively. It was assumed that the con- tribution of the added salt to the susceptibility of the solution was negligible as shown by Templeman and Van Geet (106). The magnitude of the correction for various solvents is given in Table 5. The paramagnetic shift from the reference (downfield) is designated as a positive value. Carbon-13 NMR measurements were performed on a Varian OFT-20 spectrometer operating at a field of 1.868 Tesla and a frequency of 20.0 MHz in the pulsed Fourier transform mode. The sample solution was in an 8 mm o.d. NMR tube which was coaxially centered in the 10 mm o.d. NMR tube containing the mixture of reference and lock sol- vents (50% v/v acetone:D20). The methyl carbon NMR peak of acetone was used as the external reference. All carbon- 13 chemical shifts were corrected for the differences in bulk diamagnetic susceptibilities of solvents according to equation (2.1) and referred to the internal TMS resonance in acetone. The paramagnetic shifts (downfield) are designated as positive. 44 Proton (1H) NMR measurements were Obtained by using a Bruker WH-180 superconducting spectrometer operat- ing at a field of 4.228 Tesla and a frequency of 180.05 MHz in the pulsed Fourier transform mode. A Bruker ENC-80 computer was used to perform the Fourier trans- formation. Deuterated solvents were used to lock the magnetic field. All the proton chemical shifts reported are referenced to the internal TMS resonance. B. Calorimetry Enthalpies of Na+ ion-macrocyclic complexation re- actions were determined with a Guild Model 401 isoperibol solution calorimeter (107, 108) under nitrogen atmosphere. The calorimeter cell, calorimeter insert (including thermister, calibration heater, stirrer and stainless steel cooling coil) and stirrer motor were contained in a glove bag (12R, Model X-27-27) which was continuously purged with a flow of nitrogen during the entire experi- mental process. A Beaker containing some fresh P205 was placed inside the glove bag to absorb any traces of moisture. A Sargent-Welch Model SRG millivolt strip chart recorder was used to record the temperature changes. The voltage applied to the calibration heater was measured with a Keithley Model 169 digital multimeter and an 8.5:1 voltage divider fabricated with 1% precision resistors. Potentials could be read in a range of i 2.000 V with i 0.5 mV accuracy. 45 About 50 ml of the sodium salt solution was allowed to equilibrate for ~ 1.5 hours in a calorimeter cell which is the smaller one of the two silvered glass dewars fabricated in the Michigan State University glass shop (61). When the temperature of the solution inside the calorimeter cell was higher than the ambient temperature, the horizontal baseline was kept as Close as possible to the ambient tem- perature by adjusting the flow rate of cooling nitrogen gas which came from an external cooling coil immersed in an ice bath and went into an internal cooling coil immersed in the solution. If the temperature of the solution inside the calorimeter cell was lower than the ambient temperature. the baseline was obtained by heating the solution to slightly above the ambient temperature with the calibration heater and then adjusting the flow rate of cooling nitrogen gas to compensate for the heat of stirring. The system was then calibrated by electrically adding a definite quantity of energy into the calorimeter cell and measuring the recorder response. Because it re-' quires a finite time period to generate the calibration energy electrically, and there is a change in the rate of heat loss as the solution temperature changes due to the imperfect insulation, the final solution temperature is not reached instantaneously. A typical recorder response curve for calibration is shown in Figure 4. The recorder deflection corresponding to the energy of 46 energy of calibration ._; -__- —-———————..._—__ ---S/....- -4- - V Figure 4. Calorimeter calibration response curve. .1F- -- ‘N‘ heat of reaction ._4 Figure 5. Calorimeter response curve for fast exothermic reaction. 47 calibration was Obtained by using a temperature extrapola- tion and time averaging procedure as shown in Figure 4. The energy (Q, calories) generated was calculated by using equation (2.2) (108) Q = ———————' (2.2) where R(ohms), the resistance of the calibration heater, V(volts), the voltage across the calibration heater and t(sec), the time of heating, were all known and the factor 4.184 converts joules to calories. The recorder deflec- tion was then calibrated in units of calories per division on the chart paper. After this first calibration run, the system was cooled down to the initial temperaturetnrincreasing the flow rate of the cooling gas and re-equilibrated. The entire calibration procedure was repeated. Generally, the precision of the calibration procedure was found to be better than i 1%. Then, the heat of dilution of the salt solution (as the ligand solution was added) was determined by adding a known volume (1 ml or less) of pure solvent to the calorimeter cell with a syringe and measuring the recorder response. The system was brought back to the initial temperature as previously described, and the same volume of a ligand solution was added to the calorimeter cell with a syringe. After the complexation reaction was complete, the system was brought back to the initial 48 temperature, re-equilibrated, re-calibrated twice. A typical response curve for an exothermic reaction is shown in Figure 5. For a rapid reaction, the linear portion of the temperature decay was extrapolated back to the initial time of reaction to determine the heat released. For a slower reaction, the response curve resembled a calibra- tion curve and was analyzed similarly. A separate experiment was performed to determine the heat of dilution of the ligand solution by adding the same volume and concentration of the ligand solution as in the previous complexation reaction run into about 50 m1 pure solvent in the calorimeter cell. The heat of reaction was calculated by comparing the recorder deflec- tion with the post-reaction calibration data and was corrected for any heat of dilution of the salt and ligand solutions. Since the system is calibrated directly in calories per division on the recorder chart paper, solution heat capacity measurements may be omitted. For reactions which go virtually to completion, the enthalpy of complexation is obtained by dividing the heat of reaction by the number of moles of complex formed. For incomplete reactions, the complex formation constants must be known in order to calculate the number of moles Of complex formed or else a large enough excess of one of the reaction components must be used to drive the reaction completely to the right. 49 C. Data Handling The stability constants of complexes were cal- culated by fitting the NMR chemical shift-mole ratio (ligand to metal ion) data to appropriate equations using a weighted nonlinear least squares program KINFIT (109) on a CDC-6500 computer. Details on the use of this program are given in the appendices. CHAPTER III NMR STUDIES OF SODIUM CATION COMPLEXES WITH SOME CROWN ETHERS IN VARIOUS SOLVENTS 1. Introduction The importance Of nuclear magnetic resonance spectroscopy, especially proton and carbon-13 NMR, as applied to the elucidation of the structures of organic molecules and electrolyte solutions is well established. In recent years, an increasing number of investigators are utilizing magnetic resonance of other nuclei, particularly the alkali metal cations and the halide ions, to study the behavior of ions in aqueous and various nonaqueous solu- tions (110-114). Alkali metal and halide NMR studies allow a more direct observation of ionic interactions. Moreover, the sensitivity (to the immediate chemical environments) of alkali metal NMR is considerably larger 13C NMR which in turn is larger than that of than that of PMR. Hence, weak ion-ion, ion-solvent and ion-ligand interactions can be observed more easily by alkali metal NMR methods with the aid of computerized Fourier transform instrumentation. By monitoring the chemical shifts and linewidths of the resonances and the relaxation times of various nuclei, one may Obtain information about ionic solvation, association and complexation. Thus, 7Li, 23Na, 39K and 13305 NMR have been shown to be very sensitive and powerful probes of the immediate chemical environments of the 50 51 individual cations (2, 115-1I8). The physical properties Of the 23Na, 13C, and proton nuclei are shown in Table 6 (119). As indicated previously in the historical review, with the exception of methanol, very few studies on sodium cation complexes with crown ethers have been done in non- aqueous solvents. This chapter reports studies by sodium- 23, carbon-13 and hydrogen-1 NMR techniques of Na+ ion complexes with three crown ethers; 15C5, MBlSCS and 18C6, in several solvents. 2. Results and Discussion The variation of the sodium-23 chemical shift as a function Of ligand/Na+ mole ratio in various solvents at ambient temperatures is shown in Tables 7-9 and Figures 6-9. In all cases, except 18C6 and sodium tetraphenylborate in THF solutions, (which will be discussed in detail in part 3), only one population-averaged resonance was ob- served indicating that the exchange of the Na+ ion between the bulk solution and the complex is fast on the NMR time scale. From Figures 6-9, it is obvious that the solvent plays an important role in the complexation reaction. Generally, in solvents of poor solvating ability the chemical shifts begin to level off at mole ratio of about one indicating the formation of a stable 1:1 complex. How- ever, in stronger solvating solvents the chemical shifts change gradually and do not reach the limiting values as 52 Table 6. The Physical Properties of the 23Na, 13C and Proton Nuclei 23 13 1 Na C H Resonance frequency in MHz for a 1.409 Tesla field 15.87 15.08 60.00 a 2.114 Tesla field 23.81 22.63 90.00 Natural abundance, % 100 1.108 99.985 Relative sensitivity (vs. 1H) for equal 9.25xlo‘ number of nuclei at constant field 2 2 1.59x10’ 1.00 Magnetic moment u, in multiples of the 2.2161 0.702199 2.79268 nuclear magneton (eh/4WMC) Spin I, in multiples of 3/2 1/2 1/2 h/ZW Electric quadrupole 0.14-0.15 - - moment Q, in multiples of e x 10.24 cm2 Figure 6. 53 NaCIO4 in H20 / /Nalln H20 NoCI in H20 I l l l 1.0 2.0 3.0 4.0 5.0 6.0 7.0 C l8C6 C No 1' Sodium-23 chemical shift Kg 1806/Na+ mole ratio in various solvents. The solutions, except H20, were 0.05 M in NaBPhA. 54 ‘8 I -6._ DMF A g '4 — ' THF CL 9: 4 '2 0.0 I |.0 2.0 3.0 4.0 5.0 6.0 Cl5€5 2 Figure 7. Sodium-23 chemical shift ngSCS/Na+ mole ratio in various solvents. The solutions were 0.05 M in NaBPh4 except in H20 where the salt was NaI. 55 I 1 l 1 l l I 0.0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0 10,0 11.0 C15C5 N¢+ Fig 8. Sodium-23 chemical shifts vs. 1505/Na+ mole ratios in various solvents. The solutions were 0.05 M in NaBPh except in H20 where the salt was NaCl. 4 56 o 1 1 1 1 1 1 / 0.8 (.6 2.4 3.2 4.0 4.8 2L C815c5 CNa” Figure 9. Sodium-23 chemical shift vs. 131505/Na+ mole ratio in various solvents. The—Solutions were 0.05 M in NaBPh4. 57 Table 7. Sodium-23 NMR Chemical Shift-Mole Ratio Data for + . . . . Na ion Complex w1th 18C6 in Various Solvents at Ambient Temperaturesa Water Water Water [18C6]b 5 [lBCGA 6 [18C6] *5 5 [NaCl] ppm [NaI] ppm [NaClO4] ppm 0 -0.46 0 -0.45 0 -0.65 0.12 -0.54 0.17 -0.88 0.30 -l.32 0.32 -1.00 0.35 -1.23 0.53 -1.73 0.42 —1.23 0.52 -1.58 0.58 -l.69 0.56 -1.46 0.58 -1.73 0.86 -2.4 0.72 -l.77 0.72 -2.04 1.06 -2.6 0.83 -2.0 1.23 -2.9 ' 1.23 -2.9 0.97 -2.1 1.60 -3.3 1.83 -3.8 1.12 -2.5 1.65 -3.4 2.01 -4.1 1.32 -2.7 2.13 -4.0 2.39 —4.5 1.63 -3.3 2.55 -4.4 2.87 -5.1 1.95 -3.7 2.84 -4.8 3.30 -5.4 2.13 -3.8 3 03 -4 9 3 60 -5 6 2.56 -4.3 4.20 -5 8 4.13 -5.9 3.17 -4.9 6.06 -6.8 4.53 -6.3 3.77 -5 4 5.03 -6.4 4.24 -5.7 7.00 -7.5 58 2.12 Table 7 (continued) MeCN MeZCO -—E::‘ifl ——::.:s:1- _;;:s:; 0 -7.56 0 -8.08 0 -9.7 0.23 -9.48 0.24 -10.30 0.36 -12.4 0.45 -ll.12 0.51 -12.53 0.52 -13.4 0.55 -11.91 0.73 -l4.4l 0.74 -l4.8 0.83 -l3.80 1.04 -l6.3 0.84 -15.3 1.00 -14.94 1.46 -l6.4 1.02 -16.0 1.22 -15.22 2.00 -l6.4 2.02 -16.2 1.46 -15.22 3.58 -l6.4 3.12 -16.1 1.78 -15.32 2.34 -15.35 3.25 -lS.30 MeNO c DMSOC PyridineC 0 -14.68 0 -0.5 0 +0.75 0.20 -15.11 0.51 -4.0 0.50 ~ -0.3 0.52 -15.7 1.00 -6 0.99 ~ -13 0.74 -l6.l 1.06 -l6.8 1.26 -16.7 1.51 -16.7 Table 7 (continued) 59 DMF TMG [18C6] [18C6] [Na+]e PPm [Na+]c ppm 0 -5.0 0 -10.2 0.98 ”‘13 0.60 ~ 4 l 07 ~ 9 2.12 ~12 aThe temperature was either 27 : 1°C or 28 : 1°C. b The solutions were 0.05 M’in sodium salt. CThe solutions were 0.05 M in NaBph4. dInsoluble. 8The solutions were 0.10 M in NaClO4. 60 Table 8. Sodium-23 NMR Chemical Shift-Mole Ratio Data for Na+ ion Complex with 15C5 in Various Solvents at Ambient Temperatures Water Water DMF [15C5] 6 [15C5] [15C5] O [NaCl]a 99“ [NaI]a ppm [Na+] PP“ 0 -0.46 0 -0.44 0 -5.06 0.30 -0.78 0.14 -0.54 0.14 -5.31 0.46 -0.84 0.19 -0.58 0.27 -5.42 0.67 -0.97 0.32 -0.64 0.38 -5.52 0.79 -1.02 0.65 -0.81 0.52 -5.77 1.04 -1.10 0.77 -0.85 0.76 -5.96 1.30 -l.21 1.08 -l.04 0.94 -6.08 1.47 -1.30 1.35 -1.08 1.13 -6.15 2.13 -l.S6 1.62 -1.19 1.18 -6.23 2.93 -1.77 2.00 -1.35 1.25 -6.31 3.73 -2.05 2.49 -1.54 1.47 -6.38 4.84 -2.28 3.23 -1.77 1.67 -6.38 5.60 -2.48 3.88 -2.00 2.16 -6.50 6.63 -2.74 4.02 -2.04 2.59 -6.58 7.30 -2.79 5.31 -2.38 2.83 -6.56 10.44 -3.21 5.58 -2.35 3.07 -6.62 7.14 -2.69 3.19 -6.60 3.36 -6.65 3.53 -6.64 61 Table 8 (continued) DMSO Pyridine MeNO2 [15C5]~ 5 [15C5] 5 [1505] [Na+]E ppm [Na+]b ppm [Na+]b ppm 0 -0.5 0 +0.75 0 -14.55 0.23 -1.2 0.18 +0.07 0.14 -13.4 0.36 -1.5 0.31 -0.31 0.24 -12.5 0.52 -1.8 0.34 -0.41 0.51 -10.2 0.74 -2.1 0.57 -l.27 0.73 - 8.1 1.01 -2.4 0.80 -1.85 0.85 - 6.8 1.17 -2.5 0.97 -2.49 0.97 - 5.8 1.44 -2.9 1.03 -2.72 1.04 - 5.3 1.52 -3.0 1.24 -2.81 1.34 - 5.5 2.24 -3.5 1.34 -2.93 2.26 - 5.8 2.35 -3.5 1.45 -2.97 3.22 - 6.1 4.14 -4.3 1.57 -3.04 3.94 - 6.3 1.87 -3.03 4.52 - 6.5 2.27 -3.08 5.65 - 6.8 2.31 -3.16 6.45 - 7.1 2.72 -3.18 8.32 - 7.4 3.16 -3.20 9.90 - 7.6 3.56 -3.20 3.75 -3.20 62 Table 8 (continued) MeCN THF PC [15C5] [15C5] [1505] [Na+]b ppm [Na+]b ppm [Na+]b Ppm 0 -7.56 0 -7.72 0 -9.7 0.24 -6.89 0.15 -7.18 0.11 -9.1 0.41 -6.43 0.31 -6.65 0.37 -8.0 0.64 -5.86 0.58 -5.8 0.52 -7.7 0.87 -5.33 0.75 -5.2 0.70 -6.9 1.04 -4.89 0.94 -4.7 0.90 -6.3 1.99 -4.93 1.05 -4.4 1.19 -5.9 2.11 -5.54 2.10 -4.4 1.48 -5.9 4.20 -5.77 2.93 -4.4 2.01 -5.8 6.27 -6.00 2.70 -6.0 7.56 -6.19 3.12 -5.9 9.95 -6.46 aThe solutions were 0.05 M b The solutions were 0.05 M in in sodium salt. NaBph 63 Table 9. Sodium-23 NMR Chemical Shift-Mole Ratio Data for + . . . . Na lon Complex with MB15C5 in Various Solvents at Ambient Temperatures Pyridine DMSO MeNO2 [MBlSCS] 0 [MBlSCS] 0 [MBlSCS] 5 [Na+]a ppm [Na+]a ppm [Na+]a ppm 0 +0.75 0 -0.5 0 -14.55 0.19 +0.25 0.13 -0.7 0.21 -12.5 0.30 -0.08 0.33 -0.9 0.30 -ll.5 0.52 -0.7 0.55 -l.l 0.76 -l.3 0.71 -l.2 0.42 -10.6 0.89 -l.5 0.92 -1 2 0.52 - 9.2 1.03 -l.8 1.00 -1.3 0.60 - 8.5 1.22 -2.1 1.13 -l.3 0.78 - 6.8 1.49 -2 2 1.34 -1.5 0.92 - 4.9 1.81 -2.2 1.45 -1.6 1.02 - 4.2 2.08 -2.2 1.63 -l.6 1.40 - 4.5 2.16 -2.3 1.72 -l.7 1.58 - 4.6 2.43 -2 3 2.02 -l.8 1.83 - 4.7 3.10 -2.2 2.22 -2 0 1.95 - 4.8 4.15 -2.4 2.49 -2 0 2.20 - 5.1 2.77 - 5.3 64 Table 9 (continued) MeNO2 MeCN THF [MBlSCS] [MBlSCS] [MBlSCS] [Na+]a ppm [Na+]a ppm [Na+]a ppm 4.10 -6.3 0 -7.56 0 -7.76 4.82 -6.5 0.11 -7.14 0.19 -7.00 6.01 -7.2 0.34 -6.30 0.43 -6.0 8.33 -8.2 0.56 -5.61 0.55 -5.5 0.68 -5.10 0.61 -5.3 0.91 -4.5 0.79 -4.6 1.04 -4.1 1.01 -3.8 1.40 -4.0 1.29 -3.8 2.16 -4.1 1.56 -3.8 1.70 -3.8 2.11 -3.7 0MPa PCa 0 -5.05 0 -9.7 0.14 -5.15 0.09 -9.3 0.36 -5.23 0.30 -8.4 0.48 -5.31 0.53 -7.4 0.77 -5.46 0.74 -6.7 1.03 -5.54 0.95 -5.6 65 Table 9 (continued) DMF PC [MBlSCS] 0 [MBlSCS] 0 [Na+] ppm [Na+] ppm 1.20 -5.6 1.09 -5.3 1.40 -5.6 2.08 -5.2 1.48 -5.6 3.33b 2.00 -5.7 2.61 -5.8 3.16 -5.8 aThe solutions were 0.05 M in NaBph4. b Insoluble. 66 ligand is added to the Na+ ion solution. The addition of 15C5 crown ether to a nitromethane solution of sodium tetraphenylborate results in a paramagnetic shift with a sharp break at mole ratio of one followed by a diamagnetic shift which is changed gradually as the concentration of the ligand is increased. This behavior clearly indicates the successive formation of a 1:1 and 2:1 (ligand : cation) complex. Both 1:1 and 2:1 complexes are also formed with 15C5 crown ether in acetonitirle and MBlSCS in nitro- methane solutions. Similar behavior was observed in the complexation of 12C4 with Li+ ion in nitromethane and PC solutions (61) as well as 1806 with Cs+ ion in several nonaqueous solvents (57). In all the above cases where two complexes are formed the cavity size of the ligand is smaller than the cation size. The successive formation of the 1:1 and 2:1 com- plexes is also shown in the variation of the half-height linewidth which is plotted versus the ligand/cation mole ratio in Figure 10 for the complexation of 15C5 and MBlSCS with the Na+ ion in nitromethane solutions. The line- width increases rapidly to mole ratio of one and then de- creases as more and more ligand is added to the Na+ ion solution. This variation strongly indicates that the electronic environment of the Na+ ion in the 2:1 "sandwich" complex is much more symmetrical than that in the 1:1 complex. 6? 260M- 240;- 220. - 200x- 160. P k 140. ‘Vyrriz 120, “XL GD. 1 lj l l 1 I l 00 30 40 50 60 20 80 90 0...»... Cad? Fig 10. The half-height linewidth vs. ligand/Na+ mole ratio in MeN02 solutions. The solutions were 0.05 M in NaBFh4. 68 The formation constants of the 1:1 Na+- crown complexes were obtained by analyzing the chemical shift vs. mole ratio data with a weighted nonlinear least-squares curvefitting program, KINFIT (109), according to the equilibrium (3.l) M+ + Lgm‘“ _ (3.1) + + . . where M , L and ML represent the metal ion, the ligand and the complex respectively. The concentration equilibrium constant (K) is given by the equation (3.2) C + K = —ML—— (3.2) C - C M+ L where C +, C + and CL denote the equilibrium molar ML M concentration of the complex, the metal ion and the ligand respectively. If only the formation of 1:1 complex is important and the exchange between the free and the com- plexed cation is fast on the NMR time scale, a population- averaged chemical shift is observed. éobs = XMéM + XMLéML (3.3) where the éobs’ 6M and 6ML are the observed, solvated and complexed cation chemical shifts, XM and XML are the relative mole fractions of solvated and complexed cation respectively. From equations 3.2-3.3 and the mass balance equations it can be easily shown that, t t 2 t2 2 t2 2 t t Ix = T, - - - . oobs [(xcM KcL l) + (K CL + K CM 2K chM + o - o 1/ zxct + 2Kct + l)2]( M ML) + 5 (3.4) L M t ML 2KcM In equation (3.4), the total concentration of the cation and the ligand (CE and c: respectively) are known, 6M is determined by measuring the cation chemical shift in the absence of the ligand. The equation, therefore, has two unknown quantities, K and 5 which are obtained ML’ by an iteration method using the program, KINFIT, starting with reasonable estimates of K and 5 The results of ML' the above analysis, the formation constants (K) and limiting chemical shifts (SML) of the Na+-crown complexes in various solvents, are shown in Tables lO-12. When both 1:1 and 2:1 complexes are formed, an upper limit for the 2:1 complexation constant (K2) was calculated by assuming the formation constant of the 1:1 complex (K1) is so large that in solutions with ligand/Na+ ion mole ratios greater than one,the concentration of the free Na+ ion is negligible, that is, only the data above mole ratio of one were used to evaluate K2. For very stable Na+-crown complexes, only the lower limit of K = 104 could be obtained by our technique. In the cases of the Na+ ion complexed with 18C6 in FY, DMF and DMSO solutions, the 23 Na resonance peak is so broad (about 220 to 400 Hz) that precise measurements of the chemical shifts are impossible with either the SCH so .poms ooqfisnoou may wo uHEHH woman on» ccoamm o v H.o H m.mHu m.o + m.m nomoz zoo: ~.o H e.oan e A ozomoz N02% «6 H H67 v A vzmmmz om N.o H «.mal v A vnmmmz OUNmz ~.o H o.oau o A ocmmoz one mo.o H om.oh bmo.o H Hm.~ enomoz tzo Mm mo.o H mo.oe oeo.o H He.~ vnomoz cmzo mo.o H oo.oo o- voomoz mononuso m.o H ~.mH- mo.o H em.o Hooz m.o H ~.oH- mo.o H mm.o Hmz o.o H m.HH- mo.o H on.o oodooz om: .amm oxowmsoo mo omega mwlmmm uHom .mmmwwmm Hmofieono mewquflA ouaumuodan unofinEd um muco>aom m:0aum> :H ondEoo oUmH.+mz mo muuwcm HmowEonu mcwufiswq cam mucmumcou cowumauom .OH manta 71 .sucflszHH womun >uo> onu mo omsmoon com: mosqwccoou man an cmCHEHmump on poccmo uHo .coumHDOHmo on \O O +| NI—INHHN OOOOOO +| +I +4 -+I +l +I N O +l ou HHmEm oou ma Acofl Esflpomuncmmflav xoamfioo Hum mnu mo H.HHI Emmgqxodmsoo mo uwwsm Hmoflsmnu mcflufleflq ousumuomsoa acmflnec um muco>aom msofium> ca moma suwz moondEou :ofi +mz mo muuficm HmoHEono mcwufifiwq can mucmumcoo cofiumfiuom N.o + m.H @OH mo.o mo.o mo.o mo.o mo.o +l +4 -+l +I +| w .memEoo Hum Gnu mo Hmowsmso mcfluweflq A aux hm.a Hm.H mo.~ v¢.c av.o x mos v ooomoz ozomoz «oomoz enomoz eoomoz voomoz «Emmmz aUmz Hmz uaom Q Nos obeo OZGE 20m: Um hmfi bio OmZQ mcfipfluam om: uco>aom .HH manna .quH3mcHH nmoun >Ho> 0:» mo mmsmoon 0mm: mmschnomu map xn pmcHEuwuop on uoccmo pH 3 .xoaosoo Hum one to bongo Hoerooo moeerMao om H mm- ~.o + m.o u sea A 7o H To: 2 I fix m.o + ~.m- ~.o H m.m.. m.o H o.b- H.o H m.o H «.m- ~.o H m.o H v.~- H.o H smatdxoamioo Mo HMHnm w HmoHEmno mCHuHEHq x ousuwuomEoB ucoHnsd um muco>aom mooHum> :H momamz :qu moonmEoo :oH +mz mo mHMHnm HmoHEozo mCHuHEHq can mucmumcoo coHumEuom mum me meg L eoomoz enomoz e53oz ¢£mmmz enomoz ebomoz «nomoz udmm 02oz zoo: Um mme bio Omzo mcHGHuxm acm>aom .NH mans? 73 DA-6O or the Bruker-9O spectrometer. Therefore, carbon-13 NMR was used to study the complexation reaction of 18C6 with Na+ ion in these three solvents. The variation of the carbon-l3 chemical shift as a function of Naf/18C6 mole ratio at 31 i 1°C is shown in Table 13 and Figure 11. The data were also analyzed with the proqram, KINFIT, and the results are shown in Table 10. The formation constant of the Na+-18C6 complex in pyridine is beyond the upper limit of our technique. It has been shown that the carbon- 13 NMR determination with a 100 MHz NMR spectrometer is accurate for the equilibrium constants within the range 100 l K i 0.01 (120). It has also been shown that the complex formation constants, within the range 200 1 K > 5, obtained from the carbon-13 NMR (using an 80 MHz NMR spectrometer) agree very well with those from the alkali metal NMR for the complexation of potassium ion with 12C4 in several nonaqueous solvents (118). The same agreement was also found for the complexation of the Na+ ion with 15C5 in DMF and DMSO solutions as shown in Tables 14-15 and Figure 12. The upper limit of the complexation con- 13 stant which can be obtained by C NMR is lower than that 23 obtained by Na NMR because of the smaller range of the 13C NMR chemical shift. Thus the experimental error in the determination of the chemical shift is relatively larger with 13C NMR. 7L» 18C6 0.0 04 0.8 1.2 1.6 2.0 24 28 6&6 l I' I r r l 1 698- O 7t2 714 7t5 Fig 11. Carbon-13 chemical shifts vs. Na+/18C6 mole ratios in various solvents. The solutions were 0.05 E in 18-Crown-6 except in DMF where the concentration was 0.06 g. 75 Table 13. Carbon-l3 NMR Chemical Shift-Mole Ratio Data for Na+ Ion Complex with 18C6 in Various Solvents at 31 :_1°C PY DMSO DMF [NaBph4] c [NaBPh4] c [NaBph4] 6C [18C61a ppm [18C61a ppm [lace]E PP“ 0 71.02 0 71.42 0 71.26 0.22 70.81 0.10 71.35 0.14 71.16 0.43 70.59 0.21 71.31 0.27 71.04 0.64 70.38 0.31 71.24 0.42 70.93 0.84 70.15 0.50 71.19 0.51 70.85 0.93 70.05 0.70 71.08 0.65 70.77 1.01 69.98 0.90 71.03 0.77 70.69 1.10 69.99 1.10 70.97 0.92 70.60 1.14 69.98 1.30 70.93 1.00 70.58 1.29 69.98 1.51 70.86 1.24 70.50 1.52 69.99 1.72 70.82 1.37 70.47 1.73 70.00 1.93 70.79 1.52 70.44 1.90 70.00 2.14 70.76 1.74 70.44 2.35 70.75 1.93 70.42 2.12 70.42 2.31 70.40 aThe solutions were 0.05M in 18C6. b The solutions were 0.065M in 18C6. cValues are precise to :0.05 ppm. 76 C:Nd+ (:15C5 w °:* 28 1:2 #6 27° 2.4 £3 688 - 6&0 _ W—é—‘l—é— py 6&2 - 694 - 6°“ ' .ah—DMF sea — /*/#’ A,ppm /* 700 - é *X¥* 702- ¥, 704-- f/ ’4r¥/+4r4”'[)lv155‘:’ 705 70.3 ;/////V+/+ 712‘7/l/g/ /) 1 715 L ml- Figure 12. Carbon-13 chemical shifts vs. Na*/15c5 mole ratios in various solvents. T C1scs= 0-05 Mo 77 Table 14. Carbon-13 NMR Chemical Shift-Mole Ratio Data for Na+ Ion Complex with 15C5 in Various Solvents at 31 i 1°C PY DMSO DMF [NaBph4] 6b [NaBph4] 6b [NaBph4] 6b [lscsla ppm [lscsla ppm [lscsla ppm 0 71.18 0 71.50 0 71.33 0.20 70.74 0.20 71.35 0.20 71.05 0.50 70.03 0.39 71.21 0.44 70.72 0.60 69.76 0.59 71.10 0.64 70.47 0.79 69.29 0.78 70.95 0.84 70.28 0.90 69.09 1.04 70.81 0.94 70.18 0.99 68.94 1.24 70.73 1.04 70.14 1.02 68.92 1.43 70.62 1.14 70.06 1.09 68.90 1.63 70.56 1.23 70.02 1.20 68.90 1.67 70.53 1.43 69.92 1.39 68.90 1.82 70.50 1.63 69.82 1.79 68.90 2.06 70.41 1.83 69.77 2.17 68.90 2.21 70.39 2.02 69.72 2.43 68.90 2.41 70.34 2.27 69.70 2-42 69.66 aThe solutions were 0.05g in 15C5. b Values are precise to i 0.05 ppm. 78 .com: oqucnoou onu mo uHEHH woman on» ccoxomn .Ananv .mom scum ouooo o- ao.o H os.m coma boom ago neo.o H b~.~ N.o H ~.~ coma oboe zoo: oso.o H sm.a ~.o H m.a coma oboe oomoz oa.o H n.o so.o H Hm.o coma boom Oman n- mo.o H mo.~ moms comoz ocaoansm so.o H mm.a mo.o H no.H moms enamoz use so.o H ~H.H oo.o H Hm.~ moms ocomoz omzo mmmlmwum omzz omnx mzz mmuoz .mmmmw oaom mmmwmmm we mos ex mos ex mos mzz malconumu can mzz Hmuoz Hawxa< Eoum oochuno mucmumcou coHumonmEOU wo :omHusmEoo one .mH manna 79 Since 15C5 and MBlSCS crown ethers form both 1:1 and 2:1 complexes with Na+ ion in nitromethane solution, the possibility of the formation of both complexes in other solvents could not be ruled out although our 23Na NMR measurements did not seem to show the existence of 2:1 complexes. The possibility of 2:1 complex formation was examined by using proton NMR of MBlSCS and its complex with Na+ ion. The variation of the proton chemical shift as a function of Na+ ion/MB15C5 mole ratio in d6-DMSO solutions is shown in Table 16 and Figure 13. The results seem to show that there is a very small amount of 2:1 complex in the solution as indicated by the fairly slight curvature at mole ratio of 0.5 (Figure 13) for proton number 2 as well as numbers 3 and 4, which cannot be separated by the limited resolution of the instrument used. However, the quantity of 2:1 complex is so small that it is negligible compared with the 1:1 complex. In fact, as shown in Tables 11 and 12, the Stability constants of the 2:1 complexes are quite small even in nitromethane solutions. It is expected that the formation of the 2:1 complex is more favorable in nitro- methane solution because of its poor solvating ability. By contrast, in solvents of higher donicity, the 2:1 complex may be quite unstable due to the solvation of the 1:1 complex so that the second ligand molecule cannot readily attach to the 1:1 complex. 80 Cw CB15c5 do as 1.0 15 2,0 2.5 I l I I 1 A ,H2 H3 650. é, & / H4 652, Figure 13. Proton(H1) chemical shifts vs. Na+/ B1505 mole ratios in DMSO-d6 solutions. T CB15C5= 0.05 E. Table 16. Proton Chemical Shifts of Na+/MB15C5 Complex in DMSO-d6 at Ambient Temperaturea [NaBph4] H H3 and H4 [MB15C5]b 5Hz 8ppm 6H2 ppm 6H2 5ppm 0.00 726.6 4.036 677.2 3.762 651.1 3.617 0.50 728.5 4.046 677.5 3.763 651.4 3.613 1.00 729.7 4.053 677.2 3.762 651.1 3.617 2.00 731.4 4.063 676.3 3.756 650.4 3.612 aProtons are designated by the number of the carbon to which they are attached (see Figure 13). b The solutions were 0.05M in MBlSCS. 82 The possibility of the formation of two complexes in solutions was also checked by the following method. Equation (3.3) (see page 68) can be rearranged to give equation (3.5) because XV + XML = l. sobs = XM6M + XMLSML (3.3) éobs = (5M ‘ 6ML)XM + 5ML (3'5) The relative mole fraction of solvated Na+ ion, X was MI calculated from the formation constant obtained using equation (3.4) (see page 69) and then Gobs was plotted versus X Some typical plots are shown in Figure 14. M' The data points fall on a straight line with the pre- dicted slope (5M - 6 L) and intercept (6ML), indicating M that the data fit very well the 1:1 complex model. If a 2:1 complex is also formed, its concentration is negligible as compared with the concentrations of other species in the solution. From the data in Tables 10-12, it seems that in the complexation reactions the donor ability of the sol- vent plays a more important role than the dielectric con- stant. In general, the stability of the complex increases as the Gutmann donor number (see page 15) of the solvent decreases. It is obvious that in strongly solvating sol- vents the competition between the solvent molecules and the ligand for the coordination sites of the cation should decrease the stability of the complex. However, in 83 -7.0r- -5.0)- -4.o- \ - 3.0 -\+++ 49¢ m AM) I -100- \ PY *+ oon'a‘o‘x'o‘o'fifio' CM 03.} Figure 14. Observed chemical shift vs relative mole fraction of solvated sodium ion for Na+/15C5 system in FY, DMSO and DMF solutions. 84 pyridine solution which has the highest donicity among the solvents used, the complex formation constants are quite large. Similar results have been shown in the complexa- tion of the Li+ ion with 15C5 (61), the Na+ and Cs+ ions with 1,10-dithia-18C6 (23) and the Cs+ ion with (2)- cryptands (82) in various nonaqueous solvents. A possible reason for the behavior in pyridine is that this solvent, being a nitrogen donor or a "soft" base, does not solvate strongly a "hard" acid such as an alkali ion (121). There is another exception in the complexation of Na+ ion with 18C6 that the stability constant of this complex in MeCN solution is smaller than in PC, THF and acetone which have higher donor numbers than MeCN. Similar behavior has been found in the complexation of Li+ ion with 18C6 in pc and acetonitrile solutions (61). The above exception in MeCN solution could be due to the strong interaction between the ligand and solvent molecules which decreases the stability of the complex. It has been shown that 18C6 forms a stable complex with acetonitrile (98). In the calculation of the complex stability con- stants possible cation-anion interactions were not taken into account since sodium tetraphenylborate was used in all the nonaqueous solvents studied. It was shown from conductance studies that sodium tetraphenylborate is com- pletely dissociated in acetonitrile (122) and sodium perchlorate is essentially completely dissociated in DMF 85 and DMSO solutions (115, 123). It has also been shown that the behavior of sodium tetraphenylborate is quite similar to that of sodium perchlorate in these solvents (115). Therefore, it is reasonable to assume that at the concentrations studied the competition from the ion pairing is negligible in these three nonaqueous solvents and in nitromethane solution which has a comparable di- electric constant (35.9) to that of DMF (36.7). In sol- vents of low dielectric constant such as pyridine ion pairing should be more important. Unfortunately, there are no data available in literature about the ion pair formation constants of sodium perchlorate or sodium tetraphenylborate in pyridine. Therefore,the formation constants given in Tables 10-12 for the 1:1 complexes in pyridine are lower limits which represent the relative complexing abilities of the three ligands in this solvent. It is difficult to understand the small but noticeable variation in the formation constants of Na+-18C6 complex 4! in Table 10. However, it may be due to the formation of in water with three anions, Cl-, I- and C10 as shown ion pairs to different extent with various anions. In the solvents studied the stability constants de- crease in the order Na+-18C6 > Na+-15C5 > Na+-BlSC5. It has been shown that one of the major factors affecting the stabilities of macrocyclic complexes is the relative sizes 18C6, it seems that the cavity size factor is much more important, i.e. the D818C6 fits Na+ ion much better than 18C6. Therefore, the greater stability of Na+-18C6 than Na+-15C5 complex in the solvents studied may be due to the larger number of oxygen atoms in the polyether ring. As expected, (57, 82) 15C5 crown ether forms more stable sodium complexes than MBlSCS in which the decreased basicity of the oxygen atoms, flexibility of the ring and cavity size disfavor the complex formation. 87 Table 17. Ionic Diameters of Alkali Cations and Ring Sizes of Some Crown Ethers . o a Crown 9 b Cation Ionic Diameter (A) Ether Ring Size (A) Li+ 1.72 12c4 1.2 - 1.5 Na+ 2.24 15C5 1.7 - 2.2 K+ 2.88 18C6 2.6 - 3.2 Rb+ 3.16 Cs+ 3.68 aReference 125. bReference 24. Table 18. Comparison of the Formation Constants Currently Available for the Na+ Ion Complexes with 18C6, 818C6 and DBlBC6 in Several Solvents 2915 Solvent 18C6 B18C6 DBlBCG Reference H20 0.8 50 1.16 43 DMF 2.31 60 2.69 62 MeOH 4.32 34 4.36,4.16,4.5 34,39,2 MeCN 3.8 60 4.90 41 5.04 39 88 3. Anion effects on the complexation of Na+ ion with 18C6 in THF and 1,3-dioxolane solutions It was previously indicated in part 2 that the ex- change rate of the sodium cation between the bulk solution and the complex in the system NaBPh4/18C6 in THF solutions is slow on the NMR time scale since two resonances of 23Na nucleus were observed when free Na+ ion was present in excess. The complexation reaction of Na+ ion with 18C6 in THF was studied with three salts, NaBPh4, NaClO4 and NaI, and the results are presented in Table 19. The data show that anions play an extremely important role in the complexation process. With sodium iodide or sodium perchlorate only one population-averaged 23Na NMR signal was observed under the same experimental conditions. As shown in Table 20 and Figure 15, when a small amount of tetrabutylammonium iodide was added to a solution, which was 0.05 M in NaBPh and 0.025 M in 18C6, the two sodium 4 resonance peaks collapsed to give one broad peak. As the concentration of tetrabutylammonium iodide was increased, the linewidth of the peak decreased and the frequency of the peak moved downfield toward the value of 18C6/NaI in THF at mole ratio of 0.5. In order to understand the behavior of the Na+/18C6 system in THF solutions, an attempt was made to study the complexation reaction of dibenzo-18C6 with three sodium salts, NaBPh4, NaClO4 and NaI,in THF. Unfortunately, in Table 19. 89 Sodium-23 NMR Chemical Shift-Mole Ratio Data for Na+ ion Complex with 18C6 in THF Solutions with Various Anions at Ambient Temperatures NaBph4 NaClO4 NaI [18C6] 6 [18C6] [18C6] [Naba Ppm [Na+]a ppm [Na+]a ppm 0 -7.72 o - 8.3 o 6.50 0.20 -7.9 0.51 -12.8 0.51 -2.1 c 1.02b 1.01 -11.4 0.59 -7.8 2.19 -12.6 -l7.l 2.98 ~12.s 1.04 -16.9 1.21 -17.1 3.00 -l7.l aThe solutions were 0.05 M in sodium salt. b CThe chemical shift could not be measured precisely because The complex is not soluble. the peak was very broad and noisy. 90 _1 l l l l l I l I J l 0.00102030405050708091011 C3040" Cu...» Figure 15. Sodium-23 chemical shift vs. Bu Nl/Na+ mole ratio in THF solutions. The solfitions were 0.05 M in NaBPh4 and 0.025 M in 18-Crown-6. 91 Table 20. Sodium-23 NMR Chemical Shifts for Na+ ion Complex with 18C6 in THF Solutions as a Function of Added Bu NI at 28 i 1°C 4 [Bu4NI] _________ 6 AV (Hz) [Na+]a ppm 1/2 0.10 9 218 0.37 6.9 182 0.52 6.0 161 0.78 5.4 155 0.97 4.9 160 1.02 4.8 150 2.04b - - aThe solutions were 0.050 M in NaBPh4 and 0.025 M in 18C6. bInsoluble. 92 all three cases either the free crown ether or the complex is not soluble in the solvent at total salt concentration of 0.05 M. Then the complexation reaction of the Na+ ion with 18C6 was studied in 1,3-dioxolane solutions. The results are presented in Table 21. The same phenomena as those in THF were observed, i.e. at 18C6/Na+ mole ratio of 0.5 two 23Na NMR signals were observed with the tetraphenylborate anion, while only one signal was observed with the perchlorate anion. Two typical spectra for the cation exchange between the two sites as a function of temperature are shown in Figures 16-17 and the corresponding NMR data are presented in Table 22. In Figure 17 at -14 and -21°C there is another peak coming out at higher field, corresponding to the complexed sodium ion, and the frequency of the peak at lower field is equal to that of the solvated Na+ ion and the linewidth of this peak is smaller than those at -6 and 3 -21 0C, the peak at higher field is smeared out by 00. As the temperature of the solution is lower than viscosity broadening and. therefore, two distinct peaks are not observed in this case. The coalescence temperature and approximate exchange rate at coalescence calculated from equation (3.6) for the Na+-18C6 complex in THF and 1,3-dioxolane solutions are shown in Table 23. (\)f - v ) (3.6) 93 Table 21. Sodium-23 NMR Chemical Shift-Mole Ratio Data for Na+ ion Complex with 18C6 in 1,3-dioxolane Solutions with Various Anions at 28 : 1°C NaBPh4 NaClO4 NaIb [18C6] .obs. [18C6] obs. + a Oppm + a 5 pm (Na 1 [Na 1 P 0 -9.90 0 —9.40 0.57 ~9.8 0.54 ~13.6 -l6.8 1.04 -16.9 1.11 -17.0 2.16 -l7.0 2.01 -17.0 aThe solutions were 0.05 M in sodium salt. b 23Na NMR measurements. NaI is not soluble enough in 1.3—dioxolane to perform 94 64 42 28° 20 ppm I Figure 16. Sodium—23 NMR spectra at various temperatures for a solution containing 0.05 M NaBPh and 0.026 M 18-Crown-6 in 1,3-Dioxolane. 95 28 -21 ‘-30 -41 1.1 30 ppm L rfi —' Figure 17. Sodium-23 NMR spectra at various temperatures for a solution containing 0.05 M NaClO4 and 0.025 M 18-Crown—6 in 1,3-Dioxolane, 96 Table 22. Sodium-23 NMR Chemical Shifts at [18C6J/[Na+] Mole Ratio of 0.5 in 1,3-dioxolane Solutions at Various Temperaturesa NaBph4 NaClo4 Temp(°C) GObS AV (Hz) Temp(°C) sobs AV (Hz) PPm 1/2 ppm 1/2 28 - 9.7 54 28 -13.1 45 -l7.l 93 3 -12.8 114 42 -10.7 -b -6 -10.5 123 -l6.8 -b -14 - 9.7 92 50 -11.8 -b -21 - 9.6 82 -16.3 -b ~30 - 9.8 84 57 —15.5 118 -41 - 9.5 89 64 -l4.8 74 -59 -9.4 133 aThe solutions were 0.05 M in sodium salts. bThe two peaks were so close to each other that the AV could not be measured. 1/2 97 Table 23. The Coalescence Temperature and Approximate Exchange Rate of Na+ ion Complex with 18C6 in THF and 1,3-dioxolane Solutions [18C6] Coalescence b Exchange Rate Solvent Salt [Na+]a Temperature (°C) at Coalescence -1 (s ) THF NaBph4 0.59 41 3.3 x lo2 NaClo4 0.54 -20 -c NaI 0.50 -30 6.7 x 102 1,3- 2 dioxo- NaBPh4 0.52 54 2.5 x 10 lane 2 NaClO4 0.50 -5 2.7 x 10 aThe solutions were 0.05 M in salt. bValues are precise to : 3°C. cIt can not be calculated because the resonance frequency of the complexed Na+ ion is unknown due to the limited solubility of the complex. 98 where of and Vc are the resonance frequencies of free and complexed Na+ ion respectively at total salt concentra- tion of 0.05 M and k is the exchange rate. It was shown previously that the ion pair dissocia- tion constant of NaBph4 is 8.8 x 10"5 at 25°C and that NaI forms a much more stable ion pair than NaBPh in THF 4 (125, 115). It has also been shown that the dissociation constant of the D818C6 complexed NaBPh4 ion pair i.e. 5 Na+-DBC-BPh ' is 6.0 x 10' at 20 °C in dimethoxyethane 4 which has dielectric constant (7.2) as low as that of THF (7.6) (45). Therefore, the observed variation in the cation exchange rate may be due to ion pair or even triple ion formation of the solvated salts and the ion pair formation of the complexed salts to different extent with different anion. 4. Sodium-23 NMR study of thelexchange kinetics for sodium tetraphenylborate complex_with 18C6 crown ether in 1,3-dioxolane solutions It was previously shown in part 3 that the apparent exchange rate of the sodium cation between the bulk solution and the complex in the system NaBPhu/18C6 in THF and 1.3-dioxolane solutions is slow on the NMR time scale at ambient temperatures. Therefore, an attempt was made to study the kinetics of the complexation reaction between the 23Na Na+ ion and the 1806 crown ether in 1,3-dioxolane by NMR technique from changes of the resonance lineshape as a 99 function of temperature (116, 117). In order to determine the linewidths of the two sites (free and complexed Na+ ion) at various temperatures in the absence of exchange, separate lineshape analyses were made of the salt and the completely complexed Na+ ion at various temperatures. The results (chemical shifts and linewidths) are listed in Table 23(b). It is immediately seen that the linewidth of the free Na+ ion increases as the solution temperature increases and the chemical shifts of the free Na+ ion and the complex change with temperature. As far as the viscosity is conCerned. the linewidth of a resonance in the absence of exchange should decrease as the solution temperature increases because the viscosity of the solution decreases. Therefore, the variation of the linewidth of the free Na+ ion with temperature seems to indicate that there is an exchange between the free Na+ ion and the ion pair in the salt solution. Similarly, it may have an exchange between the complexed Na+ ion and the ion pair of the complexed salt in the solution containing stoichiometric amount of the salt and the ligand. It is not surprising that there is substantial ion pairing between the free (and the complexed) Na+ ion and the tetraphenylborate anion in this solvent of low dielec- tric constant (see part 3). Therefore, the complexation Table 23(b). 100 The Observed 23Na Chemical Shifts and Half- Height Linewidths of Solvated and 1806 Complexed Na+ ion in 1.3-Dioxolane at Four Temperatures Na+a Na+ Na+-18C6b Na+-1806 Temp (QC) 8ppm AV1/2 Hz 8ppm Avl/Z Hz 28.3 1 0.1 -9.8C 31.3d -16.9 i 0.2 88.9 i 0.5 37.6 -10.7 35.6 -16-5 75.5 46.4 -ll.2 39,0 -l6.2 60.8 54.9 -11.9 “3.8 ~15.7 49.8 a. The solution was 0.051 M in NaBPh“. b. The c. The d. The solution was 0.051 M in NaBPh“ and 0.055 M in 1806. values are precise to + values are precise to + 0.1. 0.5. 101 reaction in solutions containing free Na+ ion in excess in- volves exchanges of more than two sites (free and complexed Na+ ion). As seen in Table 23(c), however, the chemical shifts of the free and of the complexed Na+ ion are independent of concentrations. Both salts are preferentially ion-paired in the measurable (by 23Na NMR) concentration ranges. Since the Chemical shifts and the linewidths of the free Na+ ion. the complex and the ion pairs of the free and the complexed salt cannot be obtained, the kinetics of the complexation reaction in 1,3-dioxolane (and likewise in THF) solution cannot be studied by 23Na NMR method. Thus, the observed variation in the apparent cation exchange rate with the anion (part 3, Table 23) in these two solvents may attribute to the different exchange rates between the various species in the solutions. 102 Table 23(c). The Observed 23Na Chemical Shifts of Solvated and 18C6 Complexed Na+ ion in 1,3—Diox01ane as A Function of Concentration at 28 i 1 OC + + Na Na -18C6 Conc.a(M) sppm Sppm 0.010 -10.0 i 0.1 -17.0 3:. 0.2 0.030 -9.9 -17.0 0.050 -9.9 -l7.0 0.10 -16.9 a, The salt was NaBPh“- 1. Introduction It has been shcwn previously that the stabilities and selectivities of the cryptand complexes with suitable alkali cations are several orders of magnitude larger than those of crown ethers (5). Cahen et al. (78) found that the chemical shift of the lithium cation complexed by cryptand C211, which has a cavity radius nearly equal to that of the unsolvated lithium ion, is essentially in- dependent of the solvent and the counterion used. Mei et al. (83) showed that the limiting Chemical shift of the Cs+-0222 complex is dependent on both the solvent and the temperature but converges to a solvent-independent value at low temperature. Since the cavity size of cryptand C221 is nearly equal to that of the unsolvated sodium cation and the radius ratio of Na+/0211 (=1.4) is close to that of Cs+/0222 (=1.3), it would be interesting to study the complexation reaction of Na+ ion with the cryptands. This chapter reports the studies of sodium cation complexes with four cryptands; C211, 0221, 0222 and 02228 in several solvents using sodium-23 NMR technique, 2. Results and Discussion The variations of the sodium—23 Chemical shifts as a function of cryptand/Na+ mole ratio in various solvents at ambient temperatures are shown in Tables 24-27. In most of the cases studied, the cation exchange between the bulk solution and the complex is slow on the NMR scale since 103 CHAPTER IV NMR STUDIES OF SODIUM CATION COMPLEXES WITH SOME 2-CRYPTANDS IN VARIOUS SOLVENTS 104 'Table 24. Sodium-23 NMR Chemical Shift-Mole Ratio Data + . . . . for Na ion Complex with C211 in Various Solvents at Ambient Temperature + [C211] Solvent Salt [Na ](M) [Na+] EEEE. AVl/2(Hz) ‘Water NaI 0.0502 0 -0.45 7 0.56 -0.23 31 12.6 262 1.11 11.8 329 2.17 12.3 264 2.95 12.4 272 NaCl 0.0501 0 -0.46 8 2.02 12.7 268 Pyridine NaBPh4 0.100 0 0.84 18 0.65 -a - 0.99 10.7 251 2.00 10.8 208 3.35 10.9 297 NaClO4 0.0508 0 -0.7 39 0.43 -l.0 39 ~11.5 -b 1.03 10.9 217 2.18 10.8 219 NaI 0.0500 0 7.3 40 o 61 7.5 77C 1.00 11.5 230 1.79 11.5 242 105 Table 24 (continued) + [C211] Solvent Salt [Na ](M) [Na+] 522$ AVl/2(Hz) DMSO NaBPh4 0.0501 0 - 0.5 51 0.88 - 0.5 95 11.8 -b 1.09 10.0 294 2.26 9.7 289 DMF NaBPh4 0.0502 0 - 5.04 28 2.30 10.1 194 NaClO4 0.0501 0 - 5.13 27 0.56 - 4.8 46 -d ~195 1.15 10.1 177 2.21 10.0 187 N31 0.0502 0 - 4.95 28 0.59 - 4.6 56 10.2 192 1.15 10.0 174 2.48 10.0 171 THF NaBPh4 0.0502 0 - 7.72 25 0.62 - 7.4 60 10.5 ~198 1.40 10.7 252 2.00 10.6 264 Table 24 (continued) Solvent Salt NaClO4 NaI MeZCO NaBPh4 NaI 106 (Nah (M) 0.0500 0.0251 0.100 0.0503 0.100 0.0501 0.0249 [C211] [Na+] 0 -8.29 ~8.06 6.55 9.4 -7.94 -7.74 ~11.5 10.5 10.2 -8.08 10.5 -4.99 10.8 10.9 AVl/2(Hz) 16 27 198 203 232 203 22 162C 17 20 ~176 200 193 15 161 151 164 19 142 142 107 Table 24 (continued) + [C211] Solvent Salt [Na ](M) [Na+] 322$ AVl/2(Hz) MeCN NaBPh4 0.0502 0 -7.56 10 0.55 -7.56 20 ~12.1 ~156 1.00 11.2 182 2.38 11.2 183 NaClO4 0.0506 0 -7.8 8 0.52 -7.5 12 11.5 145 2.54 11.2 168 NaI 0.0536 0 -6.51 8 0.62 -6.53 13 11.5 172 1.02 11.4 164 2.06 11.5 168 MeNO NaBPh4 0.0502 0 -14.55 10 0.51 -13.90 37 -d - 1.06 ~11 ~471 PC NaBPh4 0.100 0 -9.7 62 1.06 -d - MeOH NaBPh4 0.0502 0 -3.71 17 108 Table 24 (continued) aThe two peaks were not well-separated. bThe two peaks were so close to each other that AVl/2 could not be measured precisely. cAn unsymmetrical peak. dThe peak was so noisy that the chemical shift could not be measured precisely. eInsoluble. 109 Table 25. Sodium-23 NMR Chemical Shift-Mole Ratio Data for Na+ ion Complex with C221 in Various Solvents at Ambient Temperature + [C221] Solvent Salt [Na ](M) [Na+] 022$ AV1/2(Hz) Water NaI 0.0502 0 -0.45 8 0.52 -0.69 16 -a - 1.08 -4.8 102 2.18 -4.9 106 3.31 -4.8 108 NaCl 0.0506 0 -0.46 9 0.99 -4.8 99 2.59 -4.9 107 NaClO4 0.201 0.14 Pyridine NaBPh4 0.0501 0 0.76 15 1.07 -5.0 67 1.71 -5.0 69 NaClO4 0.0509 0 -0.8 43 1.32 -5.0 67 0.100 0 -0.7 53 1.13 -5.1 69 3.26 -5.1 76 NaI 0.0510 0 7.3 42 0.61 7.4 53 -6.1 75 3.01 -5.2 73 Table 25 Solvent Pyridine DMF Me CO (continued) Salt NaSCN NaI NaBPh4 NaBr NaF NaBPh 0.0504 0.103 0.100 0.0501 0.0506 0.0501 0.0500 0.0500 110 [Na+](M) [C221] [Na-1+1 0 0.46 3.19 -5.1 -4.89 -5.2 -5.3 -5.2 AVl/2( Hz) 61 85 66 70 31 41 75 76 77 27 67 44 45 111 Table 25 (continued) + [C221] Solvent Salt [Na ](M) [Na+] £222 AVl/2(Hz) NaClO4 0.0501 0 -7.96 14 1.47 -4.0 43 NaSCN 0.0507 0 -6.1 42 1.10 -4.2 46 2.32 -4.2 43 THE NaBPh4 0.0500 0 -7.85 19 1.05 -4.2 64 1.52 -4.2 67 NaSCN 0.0508 0 -3.2 60 1.38 -4.6 74 2.03 -4.8 72 NaClO4 0.0501 1.11 NaI 0.0509 0.56 PC NaBPh4 0.100 0 -9.6 68 0.52 -d 1.05 -4.3 190 2.11 —4.4 198 0.0501 0 -9.7 66 1.01 -4.6 180 2.14 -4.5 181 NaI 0.0503 0 -9.4 64 1.24 -4.6 170 2.02 -4.6 171 3.22 -4.6 171 Table 25 (continued) Solvent Salt MeCN NaBPh4 NaClO4 NaI Z-Nitro- prOpane NaBPh4 MeNO2 NaBPh4 112 [Na+] (M) 0.100 0.0500 0.0500 0.0507 0.0501 0.0501 [C221] [Na+] 0 1.05 1.87 -13.25 -12.9 AVl/2(Hz) 10 13 47 52 48 43 44. 41 43 46 18 48 79 83 78 80 71 74 113 Table 25 (continued) + [C221] Solvent Salt [Na ](M) [Na+] 0 pm AV1/2(Hz) DMSO NaI 0.0506 0 -0.7 35 3.05 -4.7 133 aThe peak was so noisy and weak that the chemical shift could not be measured precisely. bInsoluble. CThe two peaks were so close to each other that AVl could not be measured precisely. /2 dThe two peaks were not well-separated. 114 Table 26. Sodium-23 NMR Chemical Shift-Mole‘Ratio Data + . . . . for Na ion Complex with C222 in Various Solvents at Ambient Temperature + [C222] Solvent Salt [Na ](M) [Na+] 3229 AV1/2(Hz) Pyridine NaBPh4 0.100 0 0.84 20 0.50 0.84 25 ~12.4 42 0.74 0.84 20 -12.4 35 1.00 -12.5 40 2.00 -12.5 45 3.06 -12.4 42 NaI 0.0506 0 7.3 42 0.52 8.3 43 -12.2 34 1.36 -12.4 37 DMF NaBph4 0.100 0 -S.10 29 0.23 -5.4 38 -12.0 -a 0.50 -5.5 56 -ll.6 75 0.70 -6.0 58 -ll.7 -a 0.82 -1l.5 56b 1.00 -11.5 38 1.92 -ll.5 43 2.90 -11.5 43 115 Table 26 (continued) [C222] 1921].. Solvent Salt [Na+](M) PC NaBPh4 0.100 MeCN NaBPh4 0.0501 THF NaBPh4 0.0503 NaClO4 0.0502 NaI 0.0500 2-Nitro- propane NaBPh4 0.0502 0.83 1.02 2.51 0.50 1.39 1.00 l.62 0.51C 0.51C 0.60 fees -9.8 -10.4 -10.9 -11.0 -10.9 -11.4 -1l.4 -7.76 -9.2 -11.18 -1l.22 -7.71 -7.80 -1l.93 -11.93 -12.10 -13.5 -12.5 -11.9 -ll.8 AV 1/2(I 73 76 77 87 90 90 114 12 50 23 23 20 22 33 29 31 21 33 37 37 116 Table 26 (continued) Solvent Salt [Na+](M) 192231 6 AV (Hz) __ _______ [Na ] ppm 1/2 MeNOZ NaBph4 0.0500 0 -l4.50 11 0.25 -13.62 17 0.41 -13.00 22 0.56 -12.54 24 0.80 -11.58 28 0.96 -ll.06 32 1.01 -10.82 32 1.17 -10.90 31 1.53 -10.90 31 2.21 -10.81 33 aThe two peaks were so close to each other that AVl/2 could not be measured precisely. bAn unsymmetric peak. CInsoluble. Table 27. for Na+ Solvent Salt Pyridine NaBPh4 NaI NaSCN DMSO NaBPh4 117 (Nah (M) 0.0502 0.0504 0.0504 0.0501 [C222B] [Nab 0.55 1.05 1.49 2.07 0.47 1.06 1.56 1.46 0 0.69 -10.1 Sodium-23 NMR Chemical Shift-Mole Ratio Data ion Complex with C2228 in Various Solvents at Ambient Temperature AVl/2(Hz) 18 67 69 69 69 40 42 62 67, 62 63 64 61 93 66 64 66 46 58 118 Table 27 (continued) [C2228] Salt Solvent DMS O DMF MeCN MeNO 2-Nitro- propane NaBPh4 NaBPh NaBPh NaBPh NaBPh4 [NJ] (M) 0.0501 0.0500 0.0501 0.0500 0.0501 [Na+] 1.05 1.06 2.05 1.03 1.36 1.89 62 62 16 31 32 42 32 47 47 54 60 64 119 Table 27 (continued) + [C2228] Solvent Salt [Na ](M) [Na+] £223 AVl/2(Hz) MeZCO NaBph4 0.0519 0.43 -8.4 24d 1.02C NaClo4 0.100 0 -8.4 10 0.12 -8.5 12 0.33 -8.6 16 0.52C THF NaBPh4 0.0502 0.59C NaI 0.0500 0.52C aThe peak was so weak and noisy that the chemical shift could not be measured precisely. bThe two peaks were so close to each other that the AV I 1/2 could not be measured preCisely. cInsoluble. dAn unsymmetrical peak. 120 two resonances of the 23Na nucleus were observed when the Na+ ion was present in excess. The results shown in Table 24 indicate that in all solvents used in this study, the three cryptands, C221, C222 and 0222B form very stable 1:1 complexes (log K214) with the sodium cation, since the limiting Chemical shifts of the complexes are already reached at 1:1 ligand/Na+ mole ratio. The resonance linewi- dth of the Na+-C211 complex is so broad that no definite conclusion about the stability of the complex can be made. The formation constants of the above cryptates cannot be obtained accurately by our techniques (see previous chapter) which depends on the variation of population averaged resonance with the ligand/metal ion mole ratio. As the data in Table 25 indicate that the addition of cryptand 0221 to a DMF solution of NaBPh“ results in a very small change in the 23Na chemical shift, although the half-height linewidth of the resonance broadens. This observation suggests that either a very unstable comp— lex is formed or the electronic environment of the Na+ ion in the complex is very similar to that of the solvated Na+ ion in the DMF solution. In order to find out the reason for the above observation, the system Na+/0221 in DMF solution was studied with two other salts, NaCl and NaBr. The results (Table 25) Show that stable Na+-C221 cryptate is indeed formed in the solution as indicated by the variation of 121 chemical shift (in the case of NaBr salt), the large broadening of the linewidth (in all three cases) upon addition of cryptand C221, and the same chemical shift of the complex in all three cases. The data shown in Table 24 seem to indicate that the cryptand 0211 forms an exclusive complex with Na+ ion since the limiting Chemical shifts of the complexed sodium ion are both solvent- and anion-dependent. The broad half- height linewidth of the complexed sodium resonance peak indicates that in the Na+-C211 cryptate, the electronic environment of the Na+ ion is quite unsymmetrical. The possibility of equilibrium between the inclusive and exclusive Na+-C211 cryptates, such as in the case of 03+- 0222 cryptates (83), was studied by the 23Na NMR at low temperatures. Unfortunately, the chemical shift of the complexed Na+ ion cannot be measured precisely by the inst- rument used because of the very broad linewidth of the resonance at low temperatures. Therefore, no conclusion about the equilibrium can be made. Although x-ray studies of the Na+-0221 cryptate Show that in the solid state the sodium ion is located in the center of the cryptand cavity (71), the limiting chem- ical shift of the Na+°0221 cryptate is solvent-dependent (Table 28). It has been pointed out previously (78) that the chemical shift of the Li+°C211 cryptate is independent of the solvent and the counter-ion used and the limiting 122 Table 28. 23Na NMR Chemical Shifts of Solvated Na+ Ion and Na+-C221 Cryptate in Various Solvents at Ambient Temperature + + Solvent Salt [Na+](M) dNam 633m C221 Water NaCl 0.051 -0.46 -4.8 NaI 0.050 -0.45 -4.8 Pyridine NaBPh4 0.050 0.76 -5.0 NaClO4 0.100 -0.7 -5.1 NaClO4 0.051 -0.8 -5.0 NaI 0.051 7.3 -5.2 NaSCN 0.050 2.6 -5.1 DMF NaBPh4 0.100 -5.1 -5.2 NaBPh4 0.050 -5.1 —5.3 NaI 0.051 -4.89 -5.2 NaBr 0.050 -3.7 -5.3 THF NaBPh4 0.050 -7.85 -4.2 NaSCN 0.051 -3.2 -4.7 MeCN NaBPh4 0.100 -7.53 -4.0 NaClO4 0.050 -7.86 -4.1 NaI 0.051 -6.68 -4.0 MeNO2 NaBPh4 0.050 -l4.58 -3.7 MeZCO NaBPh4 0.050 -8.1 -4.1 NaClO4 0.050 -7.96 -4.0 NaSCN 0.050 -6.1 -4.2 123 Table 28 (continued) + + +. Solvent Salt [Na ](M) 3223. 522m C221 PC NaBPh4 0.100 -9.6 -4.4 NaBPh4 0.050 -9.7 -4.5 NaI 0.050 -9.4 —4.6 DMSO NaI 0.051 -0.7 -4.7 2-Nitro- propane NaBPh 0.050 -l3.25 -3.8 4 124 chemical Shift of the inclusive CS+'C222 cryptate is solvent independent at low temperature. Thus, the 23Na chemical Shifts of Na+°0221 cryptate at high C221/Na+ mole ratio (>1.4) in pyridine, MeZCO, THF and 2-nitropr0pane solutions were mearsured at various temperatures and the results are shown in Table 29 and Figure 18. When the temperature of the solution decreases, the chemical shift of the Na+- C221 cryptate moves upfield instead of converging to a solvent independent value as in the case of Cs+-0222 cryptates. This behavior suggests that the sodium ion is not completely shielded by 0221 from the external medium at all temperatures studied. For comparison, the variation of the sodium-23 chemical shift at high 0222/Na+ mole ratio (>1.3) as a function of temperature was determined in THF, DMF, MeCN and 2-nitropropane solutions. The results are presented in Table 30 and Figure 19. The changing be- havior of the chemical shift of the Na+.0222 cryptate is quite different from that of the Na+-C221, i.e. when the solution temperature decreases, the limiting chemical shift moves slightly upfield and then downfield in THF, DMF and 2-nitropr0pane solutions and it moves downfield gradually in MeCN solution. The reasons for the changing behavior of the limiting chemical shifts of the Na+-0221 and Na+'0222 cryptates as a function of temperature are not clear. However, one possible reason is a conformational change of the complex at various temperatures. Table 29. Solvent Pyridine THF MeZCO 23 The Cryptate as a Function of Temperature in Pyridine,Acetone, THF and 2-nitropropane Solu- tions. [Na+ 1.55 125 Na Limiting Chemical Shifts of Na+-C221 [C22 1] ]a Temp. (°C) b 80 62 28 10 -4 -22 -28 51 28 -3 -18 -30 -48 41 28 12 -4 -16 -22 -30 -48 AV1/2(Hz) 40 54 69 96 128 198 231 50 67 107 138 183 305 39 45 58 72 93 99 118 190 126 Table 29 (continued) Solvent %§§%%5 Temp. (°C)b dppm AV1/2(HZ) 2-nitro- propane 1.54 47 -3.4 66 28 -3.8 86 11 -4.2 108 -12 -4.8 167 -22 -5.4 218 -34 -6.3 300 aThe solutions were 0.050 M_in NaBPh4. bValues are accurate to : 1°C. 127 . PY 9 THF I Me2CO ' 2-NITROPROPANE \+ \ \\\ + / )— p— )- LLHLL 1 l I l 1, 1 l l 1 l l --50. -40. -30. -20. -10. 0.0 10. 20. 30. 40. 50. 60. 70. 80. TEMPERATURE.°C Figure 18. Limiting 0221/Na+ various solvents. 0.05 M in NaBPh sodium-23 chemical shift at high mole ratio fig. temperature in The solutions were 4. Table 30. Solvent DMF THF MeCN The 23 [C222] [Nah ‘3 1.49 128 Temp. (°C) b 28 12 Solutions 6 PPm -11.4 -1l.8 -1l.6 -11.4 -11.1 -10.9 -10.5 -10.1 -12.1 -12.6 -12.6 -12.4 -12.2 -11.7 -11.4 -ll.l -1l.0 -10.2 -1l.2 -11.0 -10.9 -10.8 -10.5 Na Limiting Chemical Shifts of Na+-C222 Cryptate as a Function of Temperature in DMF, THF, MeCN and 2- NitroprOpane AVl/2(Hz) 42 51 68 78 88 95 124 195 34 45 60 62 79 106 108 147 165 213 23 29 26 32 36 129 Table 30 (continued) [C222] Solvent [Na+] Temp. (°C) MeCN 1.39 -20 2-Nitro- propane 1.56 28 aThe solutions were 0.050 M in NaBPh bValues are accurate to : 1°C. 4. 0 ppm -10.1 -10.1 - 9.6 -11.8 -12.2 -12.3 -12.0 -1l.7 -10.9 AVl/2(Hz) 42 46 56 37 47 52 69 91 174 130 . THF T44. 0 2-Nitr0propane m I DMF App «~13. e M CN .1 e {/74 ...12 ¢\ L. 1 1 TL 11 l l 31 ad; I 1. 1 -80. -70. -60. ~50. -40. -30. -20. -10. O. 10. 20. 30. 0 TEMPERATURE. C Fig. 19. Limiting Sodi -23 chemical shift at high mole ratio, C222/Na , vs. temperature in various solvents. The solutions were 0.05 M in NaBPh4. 131 As expected (78), the limiting chemical shifts of Na+-0222 and Na+.C2228 cryptates are solvent-dependent (Table 31) because the cavity sizes of these two cryptands are larger than the Size of sodium cation. For the com- plexation reaction of Na+ ion with cryptand 02228 in pyridine and DMSO solutions, two 23Na resonance peaks (free and complexed Na+ ion) were observed at ligand/ Na+ mole ratio of one (Table 27 and Figure 20). For the system NaI/02228 in pyridine solution, the separation of these two peaks is large enough to allow a semiquantitative determination of the areas under the peaks. The ratio of the areas was estimated to be 10:1 (complexzfree). Hence, the formation constant of the Na+°0222B complex in pyridine was estimated to be log K g 5 by integrating the peak areas, taking the ion pair formation constant of the NaI salt in pyridine as Kip: 2200 (126) and assuming that the extent of ion pair formation of the complex is negligible. When the chemical shifts of these four cryptates (Na+.C211, Na+-0221, Na+-0222 and Na+-0222B) in various solvents are compared, the following features are obtained. The similarity of the large paramagnetic chemical shifts (around +11 ppm) of the Na+-0211 cryptate in the various solvents studied indicates that the Na+ ion is fored very much into the 0211 cavity and thus the short range repul- sive interaction leads to the large downfield shift. The chemical Shifts (around -9 ppm) of the Na+-0222B cryptate are slightly downfield from those (around -11 ppm) 23 Table 31. The Limiting Na NMR Chemical Shifts of Na+°C21l, Na+-C222 and Na+-C2228 Cryptates in Various Solvents at Ambient Temperature Solvent Salta Na+-C211 Na+°C222 5Na+-C2223 _______ ppm H20 NaI 12.3 NaCl 12.7 pr NaBph4 10.8 -12.5b -9.9 NaClO4 10.8 NaI 11.5 -12.4 -9.8 NaSCN -10.0 DMSO NaBPh4 9.8 -8.4 DMF NaBph4 10.1 -ll.5b -8.8 NaClO4 10.1 NaI 10.0 THF NaBPh4 10.7 -12.1 NaClO4 8.1 NaClO4C 8.1 MeZCO NaBPh4 10. 5 NaIc 10.8 MeCN NaBPh4 11.2 -ll.2 -8.9 NaClO4 11.2 NaI 11.4 MeNO2 NaBPh4 ~11 -10.86 -8.7 PC NaBPh4 -ll.4 2-Nitro- propane NaBPh4 -ll.8 -9.2 aThe solutions were 0.05 M in sodium salt unless other- wise noted. bThe solutions were 0.10 M in NaBPh4. CThe solutions were 0.025 M in NaCLO 4. 133 wwmcnm mo. Amy moQHCSImu 23m momowdsa om 0.0mo B zmH moo 0.0mw B cmmmw H: ownpoezm mowcdeos. 134 wpmcno mo. Adv modwcalmu 23w momoanzs ow 0.0mo B zmmw: H: demawsm mowzewos. a has 0.0mm R ommmm 135 of the Na+°0222. It shows clearly that the orbital overlap between the Na+ ion and the lone electron pairs of 02228 is greater than that in the Na+-0222 cryptate. This suggests that either the attractive interaction between the Na+ ion and the lone electron pairs of the ligand is stronger in Na+-02228 complex or the slightly smaller cavity size of 02228, because of the attachment of a benzo group (74), causes a short range repulsive interaction between the Na+ ion and the ligand. As far as the decreased basicity of the oxygen atoms in 02228 is concerned, Na+.02228 cryptate is expected to be less stable than Na+-0222 cryptate. However, Na+-02228 was found to be slightly more stable than Na+-0222 in 95% MeOH solution and they are about equally stable in water (Table 2). Therefore, the greater stability of Na+- 02228 than Na*-0222 may be due to the better fit of the Na+ ion into the 02228 cavity. Thus, the attractive interaction between the Na+ ion and the lone electron pairs of the ligand is stronger in Na+°02228 than that in Na+-0222, and that causes the slightly downfield shift of the Na+o 02228 resonance. CHAPTER V THERMODYNAMICS OF THE COMPLEXATION REACTION OF THE SODIUM CATION WITH SOME CROWN ETHERS IN SEVERAL NONAQUEOUS SOLVENTS 1. Introduction The free energy change (4:00) of a complexation reaction can be separated into two components: the enthalpy change,AHO, and the entropy change,AS°, i.e.I-‘eGo =AHO - T4350. The enthalpy change is determined by the bonding interaction between the metal ion and the ligand as well as between the solute Species and the solvent molecules. The entropy change depends on the overall change in the order of the system. The thermodynamic parameters of a complexa- tion reaction can be evaluated based on a determination of the formation constant as a function of temperature. The formation constants are related to the relevant thermodyna- mic parameters by the following relationships: AGO=-RT an (5.1) AGO =AHO - TASO (5.2) ln K = -AH° + 48° (5.3) RT' R Thus, ifAHo is independent of temperature, a plot of 1n K y§ 1/T (van't Hoff plot (127)) should give a straight line with a Slope of -AH°/R and an intercept ofASo/R. Another method for evaluating the thermodynamic parameters is to obtain the enthalpy value calorimetrically 136 137 at a given temperature and then, knowing-AGO at the same temperature,ASo can be easily calculated. This chapter reports thermodynamic studies of the complexation reactions between the Na+ ion and three crown ethers: 18-Crown-6, 15-crown-5, and M815-Crown-5'in several nonaqueous solvents by 23Na NMR and calorimetry. 2. Results and Discussion It was Shown in Chapter III that at ambient tempera- tures most of the stability constants of the complexes in poorly solvating solvents, such as MeCN, MeNOZ. THF, MeZCO and P0, are beyond the upper limit of the 23Na NMR method. Thus, if the complexation reaction is exothermic (which is true for most Na+-crown complexes), the stability constants at higher temperatures should be smaller and may be determined by the NMR techniques. However, with the exception of solutions in PC, the range of temperatures in which the stability constants can be determined is limited by the low boiling points of the solvents. On the other hand, in better solvating solvents such as DMSO, DMF and H20, the techniques are limited by the broadening of the 23Na resonance at low temperatures and/or the small range of chemical shifts between the free Na+ ion and at high ligand/Na+ mole ratio. The broad linewidth makes the chemical shift measurements less accurate, and the small range of chemical Shifts makes the error in the chemical 138 shift measurements relatively larger. Considering all of the above factors, the few systems, which seem to be more 23Na favorable for such studies, were investigated by the NMR method. The data are presented in Figures 21-24 and Tables 32-33. The complexation Constants were calculated by the KINFIT program as described in chapter III, and the results are listed in Table 34. It can be seen that for the systems Na+/1806 in P0 and in MeCN,and Na+/1505 in PC, the variations of the formation constants with temperatures are so small that the 23Na NMR method may be ineffective. For the system Na+/1505 in DMF and in FY solutions, the plots of 1n Kuyg 1/T are shown in Figure 25. The Slopes and intercepts of the straight lines were calculated by KINFIT and the correSponding thermodynamic quantities are presented in Table 36. Indeed, for the systems Na+/ 1806 in MeCN and Na+/1505 in PC, which have small changes in formation constants with temperatures and for which only three data points are available in the calculations, the standard deviations of the values are very large. These results, together with the data obtained from the calorimetric method, will be discussed later. It is readily seen from the above results that the thermodynamics of the sodium-crown complexation reaction can be studied by the 23Na NMR technique only in a few (- 45° -4.o — - 3.0 — 28.5" - 52° 5 x 70° gE-ao — w - - l .o — 0.0 l I l I I 1 I I I 1 LC 2.0 3.0 4.0 50 C Iscs CNa+ Figure 21. Sodium-23 chemical shift 15- Iscs/Na+ mole ratio in pyridine solutions at various temperatures. The solutions were 0.05 M in NaBPh4. 140 D M F 27n5° _.—-—-l—I""""-'"_ °——-——"‘0 ° °"'28xf ....——-—-"'."".— 67 5° 1 I I I I I I l‘ I I 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5 01505 CNa+ Figure 22. Sodium-23 chemical shift yg. Iscs/Na+ mole ratio in DMF solutions at various temperatures. The solutions were 0.05 M'in NaBPh4. 141 -1LO -100 p c -90 t.\ \ s- _. o. 8'0 \ CL ‘1 _ ’ZO- " 6.0 "' 1 l 28.20 4—4— 63.00 _V— ‘ 950° - l l l I l l I I 5.000 1.0 2.0 3.0 4.0 C1505 CNa+ Figure 23. Sodium-23 chemical shift lg. 1505/Na+ mole ratio in P0 solutions at various tempera- tures. The solutions were 0.05 M in NaBPhA. 142 +418 °c a 28.0 °c Figure 24, Sodium-23 chemical shifts vs. 1505/Na+ mole ratios in MeCN solutions at various temperatures. The solutions were 0.05 M in NaBPhA. 143 Table 32. 23Na NMR Chemical Shift-Mole Ratio Data for Na+ ion Complex with 15C5 in DMF,Pyridine and PC Solutions at Various Temperaturesa DMF 67.5°C 48.0°C l.5°C [15C5] [15C5] [1505] [Na+] Sppm [Na+] dppm [Na+] 6ppm 0 -5.08 0 -5.15 0 -6.1 0.12 -5.19 0.17 -5.31 0.27 -6.7 0.26 -5.31 0.47 -5.54 0.50 -7.0 0.57 -5.43 0.64 -5.65 0.74 -7.4 0.74 -5.54 0.83 -5.77 0.87 -5.58 1.07 -5.92 1.04 -7.7 1.16 -5.62 1.30 -6.00 1.27 -7.9 1.45 -5.69 1.53 -6.08 1.50 -8.1 1.62 -5.77 2.05 -6.23 1.66 -8.0 1.93 -5.77 2.67 -6.31 1.88 -8.2 2.26 -5.89 3.58 -6.34 2.32 -8.3 2.60 -5.96 4.06 -6.34 2.59 -8.2 3.24 -5.93 4.68 -6.38 3.30 -8.3 3.63 -6.00 5.08 -6.46 4.03 -8.3 5.35 -8.3 144 Table 32 (continued) Pyridine 112%21 70.0°C 52.0°C 4.5°C [Na ] 0 +0-38 0.57 +0.46 0.15 -0.08 0.15 -0.24 0.33 -0.39 -0.43 -1.1 0.50 -0.90 -0.81 -1.8 0.64 -l.16 -l.39 -2.5 0.82 -l.62 -l.81 -3.5 0.99 -l.93 -2.24 -3.8 1.10 -2.00 -2.35 -4.0 1.26 -2.08 -2.43 -4.2 1.46 -2.08 -2.46 -4.2 1.65 - -2.52 -4.4 1.83 -2.16 -2.58 -4.5 2.02 - -2.62 -4.5 2.48 -2.24 -2.58 -4.5 3.08 -2.39 -2.74 -4.6 3.63 -2.47 -2.85 -4.6 4.79 -2.70 -3.08 -4.8 145 temperatures are given in_Table 8. Table 32 (continued) PC -Ll§%§l- 141.0°C 95.0°C 63.0°C [Na ] 0 -9.36 -10.36 -9.8 0.11 -8.74 ~ 9.95 -9.5 0.37 -7.59 - 8.72 -8.4 0.52 -7.05 - 8.10 -7.7 0.70 -6.18 - 7.33 -7.1 0.90 -5.23 - 6.45 -6.4 1.19 -4.51 - 5.82 -5.8 1.48 -4.36 - 5.72 -5.7 2.01 -4.36 - 5.59 -5.7 2.70 -4.43 - 5.57 -5.7 3.12 -4.36 - 5.59 -5.7 3.36 -4.36 - 5.64 -5.8 aThe solutions were 0.05 M in NaBPh and the data at ambient 146 Table 33. 23Na NMR Chemical Shift-Mole Ratio Data for Na+ ion Complex with 18C6 in MeCN and PC Solutions at Various Temperatures Me CN [18C6] Temperature [Na+] 69.0°c 45.5°C -11.8°C 0 - 7.86 - 7.60 - 7.85 0.23 - 9.71 - 9.58 - 9.86 0.45 —ll.01 -10.96 -1l.4 0.55 -11.83 -11.76 - 0.83 -l3.63 -13.70 -l4.4 1.00 -14.76 -14.96 -15.7 1.22 -14.91 -15.17 -16.0 1.52 - -16.0 1.78 -15.19 - 2.34 ~15.l7 - 3.25 -15.09 -16.0 PC 47.0°C 0 -10.0 0.16 -11.2 0.50 -13.0 0.71 -l4.2 1.05 -15.8 1.53 -15.8 2.12 -15.9 3.04 -15.8 aThe solutions were 0.05 M in NaBPh temperatures are given in_Table 7. and the data at ambient 147 Table 34. Formation Constants and Limiting Chemical Shifts of Na+-15C5 and Na+-1806 Complexes in Several Solvents at Various Temperatures 4. Crown Solvent Temp(°C) Log Kf 6Na crown ppm 15C5 Pyridine 70.0 i 0.5 2.4 i 1 -2.54 52.0 2.5 :.l -2.88 28.5 2.68 i 0.08 -3.25 4.5 2.71 i 0.08 -4.7 DMF 67.5 1.5 i 0.1 -6.16 48.0 1.64 i 0.05 -6.56 27.5 1.97 :_0.05 -6.85 1.5. 2.17 :_0.08 -8.4 PC 141.0 3.2 i 0.2 -4.33 95.0 3.0 i 0.1 -5.55 63.0 3.3 i 0.2 -5.7 a 28.2 ‘ -5.8 18C6 MeCN 45.5 4.2 i 0.6 -15.16 28.0 3.8 i 0.2 -15.3 -11.8 3.9 i 0.4 -l6.0 PC 47.0 > 4.0 -15.8 28.0 > 4.0 -16.1 aIt cannot be determined by the techniques used because of the broad linewidth. 148 30 FY 2. M DMF 2. _ log Kf 1.5__ 1.0), 0.5L °°3Hrat——532 3.5 716—556 _1__03 T Figure 25. The formation constant vs. 103/T for Sodium-1505 complex in pyridine and DMF solutions. 149 systems. Therefore, we determined the enthalpy of the complexation reaction for some Na+-crown complexes by a calorimetric method and the data are presented in Table 35. It is very difficult to get the horizontal baseline, espectially in viscous solutions such as DMSO and P0, possibly because of the inefficient stirring due to the poor design of the stirrer and the fluctuation of ambient temperatures (the calorimeter cell was 393 immersed in a constant temperature bath). An attempt was made to determine the enthalpy of Na+-1806 complex in MeCN solution by the calorimetric method. Unfortunately, after the necessaryr42 hours of stirring for equilibration, the thermister lead was above the solution level due to the evaporation of the solvent. At least 45 ml of solution is needed to cover the thermister lead and the maximum volume of the calorimeter cell is about 50 ml. In the actual ex- periment, less than 1 m1 of a concentrated solution of the ligand (n20.4 M) in a given solvent was added to about 50 ml of 0.05 M salt solution. Thus, with the design of the calorimeter cell and insert which was used in this study, we cannot measure the heat of reaction in solvents of low boiling points. Since in most of our cases the complex stability constants are small, the complexation reactions were in- complete under the experimental conditions used and thus 150 Table 35. Calorimetric Data for Some Sodium-Crown Complexes in DMF and DMSO Solutions at Ambient Temperaturesa Ligand Solventb cL (g x 103) Heat of . (ca1.)° Complexation MB1505 DMF 7.66 -0.9 DMSO 6.11 -0.7 1505 mm 5.90 -1.0 DMSO 5.86 -0.9 1806 DMF 7.73 -1.6 DMSO 10.34 -1-3 a. The temperaturewes either 24‘: 1 0C or 25 1,1 00. b. The solutions were 0.05 M in NaBPhu. c. The values are precise to 1 0.1 calories. 151 Table 36. Thermodynamic Quantities for Some Sodium-Crown Complexes in Several Nonaqueous Solvents at Ambient Temperatures 23Na NMR Method Ligand Solvent ACT298 OK AHO ASO (kcal/mole) (kcal/mole) (cal/mole deg) 181505 DMF -2.2 1 0.1 DMSO -1.5 1 0.2 1505 DMF -2.69 1 0.07 -4.7 1 0.8 -7 1 3 DMSO -1.8 1 0.1 PY -3.7 1 0.1 -2.1 1 0.7 +5 + 2 P0 -443 -1 1 3 +12 i 9 1806 DMF -3.17 1 0.06 M80 -1.9 1 0.1 MeCN -5.2 1 0.2 o 1 5 +18 1 l7 Calorimetry, 1481505 DMF -3.7 1 0.4 -5 1 1 DMSO -5.5 1 0.8 -13.: 3 15C5 DMF -4.2 1 0.4 -51 1 DMSO -5.9 1 0.6 -14 1 2 1806 DMF -4.4 1 0.3 -431 1 DMSO -4.2 1 0.3 -8 1 1 a. Estimated from the values ofeHO andisSO. 152 the quantity of the complex formed was calculated from the stability constant obtained by 23Na NMR. The enthalpy change for the complexation reaction was obtained by divi- ding the heat of reaction by the number of moles of complex formed. The results are shown in Table 36, together with the calculated entropy changes. The values (Table 36) for the complexation reaction of the Na+ ion with MBI5C5 in DMSO solution are less precise than others, because the formation constant is very small and thus less complex is formed and less energy is released upon complexation. It is immediately seen that for the Na+-15C5 complex in DMF solution the values obtained from both methods agree within experimental error. For the Na+-1505 complex in pyridine solution only approximate values are obtained because of the possible ion pair formation in this solvent of low dielectric constant. In the solvents studied, most of the Na+'crown complexes are enthalpy stabilized but entrOpy destabilized. The Na+-15C5 complex in PC and Na+'1806 complex in MeCN are entirely entropy stabilized. However, no Na+-crown complex is entrOpy stabilized but enthalpy destabilized. Similar results have also been shown previously for the complexa- tion reactions of alkali metal cations with crown ethers in various solvents (23, 45, #9, 50, 51): with only a few exceptions, both the enthalpy and entropy changes are negative. 153 In DMF and DMSO solutions, all complexes are enthalpy stabilized but entropy destabilized. The enthalpy change follows the same trend as the complex stability in DMF solution while this trend is not followed in DMSO solution. The negative entrOpy changes for the complexes in these two solvents may be attributed to the increased ligand rigidity upon complex formation because the free crown ether is more flexible than the complex. The possible contributions to overall entropy changes include the different solvation entrOpies of the metal ion, the ligand and the complex as well as the changes in the ligand configurational entropy, in the translational entropy and in the total number of particles. Thus, other factors may also contribute to the above entropy changes. The more negative entropy changes for the complexes in DMSO than in DMF solutions may be due to the rearrange- ment of DMSO structure upon complexation. It is known that DMSO is a structured solvent due to dipolar interactions through the S - 0 bond (128), and the sodium ion (a small inorganic cation) acts as a structure breaker because the solvation of the cation disturbs the organization of the bulk solvent, but the complexed ion (a large organic cation) acts as a structure maker. The complexation reaction transforms a small inorganic cation into a large organic cation and the effect is a loss in entropy. 154 As the data shown in Table 36, the higher stability of the Na+-1806 complex in DMF than in DMSO is due to both the more favorable enthalpy and entropy changes in DMF solution. The more stable Na+-15C5 and Na+oMB15CS complexes in DMF than in DMSO are both due to the more favorable entropy change in DMF, while the enthalpy change disfavors the comp- lex formation in DMF. In DMSO solution, the lower stability of the Na+-1505 than Na+-1806 complex is because of the unfavorable entrOpy change for Na+-1505, while the enthalpy change actually favors the Na+o15C5 complex. 0n the other hand, the less stable Na+oMB15C5 than Na+-15C5 complex in DMSO is due to the less favorable enthalpy change of NaTMBlscs while the entropy change slightly favors the Na+1MB15C5 complex. In acetonitrile solution, the considerable interaction between the solvent molecules and the 1806 molecules (98) may account for the large entrOpy gain in the complexation reaction of Na+ ion with 1806. 155 3. Summary It was shown in chapter III that in general, the stabi- lity of the Na+ocrown complex varies inversely with the solvating ability of the solvent as expressed by the Gut- mann donor number, and the relative stabilities of the com- plexes are in the order Na+-18C6 > Na+'15C5 > Na+oMB1505. The above differences in the stabilities of the complexes were explained by the differences in the solvating ability of the solvent, the flexibility of the macrocyclic ring, the steric fit between the cation and the ligand cavity, and the number and basicity of the donor atoms in the macrocyclic ring. The thermodynamic quantities shown in this chapter indicate that the contributions of enthalpy and entropy changes to the complex stability are quite different in different system. Thus, in the various solvents the more stable Na+°18C6 than Na+-15C5 complex is not simply due to the larger number of donor atoms in the polyether ring as explained in chapter III. Similarly, the higher stabi- lities of the complexes in DMF than in DMSO is not simply because of the lower solvating ability of DMF. For the thermodynamics of the Na+-crown complexes, this chapter reports only some preliminary work from calorimetric method. More work is necessary for a better understanding of the thermodynamic behavior of the Na+ ion complexes with macrocyclic ligands in nonaqueous solvents. APPENDICES 156 APPENDICES 1. Application of computer program KINFIT to the calcula- tion of complex formation constants from NMR data. The KINFIT computer program was used to fit the sodium-23 and carbon-l3 NMR chemical shift vs. mole ratio data to equation (3.4) which was inserted by the user into the SUBROUTINE EON. _ t _ t _ 2 t2 2 t2 _ 2 t t éobs - [(KCM KCL l) + (K CL + K CM 2K CLCM + t t 1 5M ' GML L M t ML ZKC M Equation (3.4) has two unknown quantities, 6ML and K, designated as U(l) and U(2) respectively in the FORTRAN code. The two input variables are the analytical concen- tration of the ligand (CE, 3) and the observed chemical shift (6 ppm) which are denoted as XX(1) and XX(2) obs' respectively in the FORTRAN code. Starting with a reason- ably estimated value of K and 5ML' the program fit the calculated chemical shifts (the right hand side of equation (3.4)) to the observed ones by iteration method. The first control card contains the number of data points (columns l-S (Format 15)). the maximum number of iterations allowed (columns ll-lS (F 15)), the number of 157 constants (columns 36-40 (F 15)) and the convergence tolerance (0.0001 work well) in columns 41-50 (F 10.6). The second control card contains any title the user desires. The third control card contains the values of CONST(1) t M' in columns ll-ZO (F 10.6) and other constants can be (C E) in columns l-lO (F 10.6) and CONST(2) (6M, ppm) listed in columns 21-30, 31-40, etc. The fourth and final control card contains the initial estimates of the unknowns U(l) = oML and U(2) = K, in columns 1-10 and 11-20 (F 10.6) respectively. The fifth through N cards are the data cards which contain XX(l) = c: in columns l-lO (F 10.6), the variance on XX(l) in columns 11-20, XX(2) = the chemical shift at XX(l) in columns 21-30 (F 10.6) and the variance on XX(2) in columns 31-40 (F 10.6) followed by the same parameters for the next data point. Each card may contain two data points. The SUBROUTINE EQN and a sample data listing are given below. 158 "QN "MTV" H ‘ ad ’T A. “v BRVU ---‘H S I]... . .l...~ {V ..—.._.. '0. Or.» ) \l V O .J p. 3C O r z. u Tvl ‘- Ole cl 0.! l\)i\ bl... V72" 0... 5 71: d. 0 Cal. 'JYI IIVI) l .3 9.9.1. VD. ,1)? IF\ l-Cn‘ t! C. [on VA Jag?! LrJ'L 3:0- C II' 110.; 096‘. oDJ )3. 1‘4 0 a.)) SF?! 300 \l I 0‘ ‘J’h 1 «Dvcfl; (s..( l‘ J]...- ‘p D U .V 9 — (Y, 0 ?\;VI E 1.9!.) .1 p QC. 0‘11 nL 1a . ’VF‘ \I.& . vr. M. o ,0 OrnL 1‘!" \Ill 9 I O IfuClp a.) nsnvl I 0!. 0V. (\- u Or. .‘ll"... I O X'rc \I‘vo.rn.llh (In. U I... gentle-5 0 O 90 O . l‘ 1‘ \I \l.. 2:7. 1r:.9\l ‘b ‘1. a. V C 0719...}. l\ 1: . LV..O. . ...-I‘ 'I ‘0 91‘ ‘))«U0 3 Vin» YI O Q ‘avAv Pk't .. a“. 1.37 TOR. 01 0 ass. 0.53. V. Cl lo, 0 c al.. .p. O I 0" v04 ,0 a... r 0'. .llr...o;...!b f. .1) ‘l‘ .1: 2.. o On_72.(0 v. o) .1 OS 2 o-.-J V0.1. O..V.1 )015 I. 0.. V...J —I\.l OY O .x (‘1‘ 1 .ca -voJ .J .‘IVQ V 017)‘- ls.‘ 9| 0‘. O JO.¢,>OO o‘l’vl) [Y1 {.11. H 0" 016,? Al, ’Y‘IS‘I ,3 A, 2. 9 9|». 4...... .3 ‘J.I\.o‘ ,0 . n... o f.) ..3..21t ‘.I)'g-¥l).l(. VI} 0 in.‘.(.1( 0 ) !«.41/.(...\1(O \l . ‘09]. I1 Or ('1 ‘. ‘5 ‘F..‘l\!(-a, ‘6 ....c. //‘-.r.°. O “..Ia..'..’ O ..v.|.. . Yo-(..'l.49 O OO))\I.\I\J1~ o . 0' ‘10,]..a u. .c. ,)J.\I)l\01uol\ .- . . 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