Enthalpiesof Formation and Solvation of Some Organic Anions BY EDWARDM. ARNETT,DALEE. JOHNSTON,LEONARD AND E. SMALL AND DUMITRU OANCEA Department of Chemistry University of Pittsburgh Pittsburgh Pennsylvania 15260 U.S.A. Receiued 30th April 1975 Procedures are described for measuring AH, the enthalpy of deprotonation for organic Bronsted acids in DMSO. Evidence from a number of sources is presented to demonstrate that the deproton- ation process is clean and complete. Correlation with Bordwell's pKa data shows that AH; = AGb in DMSO for the compounds considered here. Good correlations are also found between A&(DMSO) and AHD(gas) for different classes of compounds. Solvation enthalpies are calculated by combining heats of deprotonation in the gas phase and in DMSO with heats of solution of the precursor acids.Comparison is made between resonance- delocalized carbanions enolate anions halide ions and alkoxide ions. The results are rationalized in terms of ionic size charge density and steric hindrance to solvation. The question of anion solvation is an essential one for elucidating acid-base interactions in solution. In addition such matters as salt effects on rates and equilibria the lyotropic series of protein denaturation salt rejection in reverse osmosis and solvent effects on nucleophilicity are of great topical significance in many areas of physical biological and applied chemi~try.~ Of particular concern to the organic chemist however are the many base-catalysed reactions-substitutions eliminations and condensations-which compose about 50 % of the " bread-and-butter " reactions of synthetic chemistry.These reactions are initiated through the conversion of an unreactive neutral organic molecule to a highly reactive anion through abstraction of a proton by a strong base. It is now recognized that solvation and ion-pairing of such organic ions can play a deciding role in their reactivity. This article presents a quantitative analysis of the effects of structural change on the solvation energies of a broad variety of organic anions from the gas phase to dimethyl sulphoxide (DMSO). We will use a simple Born cycle in conjunction with newly available experimental data. Although some of the data may require modification as new results are forthcoming the overall consistency of the results strongly enforces their validity and the (fairly obvious) conclusions which we will draw from them.It is widely understood that solvent effects on isodesmic proton-transfer equilibria A-+A'H+ A'-+AH (1) can be resolved into components attributable to the neutral acids AH and A'H and to their conjugate base anions. Some tentative conclusions can be made on the relative importance of specific solvation of neutrals and anions through studying their thermodynamic properties for transfer from one solvent to an~ther.~ However 20 E. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA the ultimate assignment for the role of solvation to each species can only be made by measuring reaction (I) in the gas phase in the absence of solvent and then separating the solvation energies of the neutrals and their ions through the following cycle.AHD A-+A’H -+ AH + A’-(gas) I1 I I AH,(A-) AH,(A’H) AHs(AH) I AHs(A‘-) 5.5. 1 1 AHD A-+A’H + AH + A’-(DMSO) We will be concerned here with enthalpy comparisons since it is much easier to measure or estimate heats of solution of neutral molecules [i.e. AH,(AH)] in DMSO than to measure their Henry’s Law constants as needed to obtain the corresponding free energy terms. Furthermore we have developed a simple procedure for corn-paring heats of deprotonation in DMSO which is applicable to a great number of organic acids including many compounds for which comparable free energy data (i.e.? pKa’s) cannot be easily measured.We shall see however that the enthalpy of deprotonation in DMSO and corresponding free energy are generally equivalent. To apply the above cycle to the calculation of anion solvation enthalpies the necessary data are obtained and treated in the following way (a)All acids are compared to cyclopentadiene as the standard acid. Its anion is the standard anion. Thus 6AHf’DMSo(AH)means “the partial molar heat of solution (i.e.? solvation enthalpy) of AH from the gas phase to DMSO minus the corresponding value for cyclopentadiene” . (b)Heats of deprotonation in the gas phase have been determined by McIver ’-’ using ion cyclotron resonance; by Bohme using flowing afterglow and by Kebarle using high pressure mass spectroscopy. We shall use Kebarle’s recent results l3 for most of our carbon acids.These will be compared with published data for alcohols and alkoxide ions. (c)Heats of deprotonation in DMSO using the potassium lyate salt (K+DMSYL-) are readily determined by solution calorimetry using a method described fully in previous publications.14* (d) The most dubious property in the cycle for many of the compounds reported here is the heat of solution from the gas phase to DMSO i.e. 6AH:’DMSo (AH)* In principle this is readily calculated by subtracting the heat of vaporization (or sublimation) from the heat of solution of pure liquid or solid AH in DMSO. In practice we find that the necessary heats of vaporization or sublimation have either not been measured or (even worse) have been reported but are obviously wrong.’ 6* l7 In order to compensate partially for the lack of vaporization enthalpies we have introduced a substitute approach (see Experimental and Results).We note that the heats of vaporization per se are not needed for our analysis or even the actual heats of solution of the various compounds from the gas phase to DMSO. What we do need are the rehire heats of solution from the gas phase to DMSO. There is a considerable body of evidence l8 to show that in general such differences are rather small and additive unless hydrogen-bonding to solvent occurs. Furthermore the relative strength of hydrogen-bonding interactions in 6AHj’DMSo(A) should be equal or proportional to the “ relative ” heat of transfer from an “ inert” solvent (such as CCl,) to DMSO.We may therefore propose that relative heats FORMATION AND SOLVATION ENTHALPIES of transfer from CC14 to DMSO should be roughly proportional to relative heats of solution from the gas phase to DMSO. As seen below this correlation holds to about k 0.5 kcal/mol for five model compounds whose heats of vaporization are reliably known. We doubt if it produces an error greater than 2 kcaI/mol in any compound in this study. Finally (e) In order to calculate relative heats of solution for anions from the gas phase to DMSO we combine terms as follows l4 6AH!'DMSo (A-) = SAHD(DMS0) -6AHD(gaS) -k 6AHf'DMSo (AH). EXPERIMENTAL MATERIALS All compounds were commercially available but were purified until homogeneous to gas liquid chromatography and the physical properties such as refractive index or melting point agreed with reliable literature values.DMSO was carefully purified as before15 by vacuum distillation from n-butyl lithium. Water concentration was maintained below 50 ppm (Karl Fischer titration) by storage over 4-A Molecular Sieves under argon in a black-taped bottle. K+DMSYL- was prepared as before15 using potassium hydride the working concentration of 0.1 M of lyate salt being established by quenching an aliquot in water and titrating with standard HCl. CALORIMETRY The 250 ml adiabatic solution calorimeter and its use for determining AHs and A& have been described fully.15 All measurements were made at 25°C and involved seven to twelve replica injections of 50-1OOmg.of AH into the calorimeter liquid -DMSO K+DMSYL- solution or CC14. Most enthalpy values are precise within a standard deviation of & 0.25 kcal/mol. Reproducibility of the system was checked frequently by determining A& for fluorene the value -18.2k0.4 kcal/mol being a well established standard in this laboratory. RECOVERY EXPERIMENTS As part of the evidence (infra vide) that the acids in this study were undergoing clean reversible deprotonation a standard quenching and recovery routine was developed. A 2 pl sample of the acid under study at 0.005 M in DMSO was injected into the water carrier stream of a DuPont 830liquid chromatograph at 1500 p.s.i. pressure. After passage through an octadecylsilane column the elution time and peak area of the acid were determined using a u.-v.detector at 254 nm. Following this a sample of the K+DMSYL- solution containing deprotonated acid in DMSO was injected. The anion of the acid and DMSYL- anion were neutralized instantly and in all cases except cyclobutanone cycloheptatriene and acetonitrile the elution volume and peak area of the quenched acid corresponded to that for the original DMSO solution of the acid. RESULTS Evidence that deprotonation of all compounds considered here was clean complete instantaneous and reversible in 0.1 M K+DMSYL/DMSO rests on the following facts. (1) Exothermic heats of reaction were displayed on the calorimeter strip chart recorder immediately after sample injection following which the recorder trace continued parallel to the base line.Heats of reaction were not concentration-dependent. Furthermore addition of dicyclohexyl-18-crown-6 polyether which we have shown l5 to dissociate potassium ion pairs through cation complexing had no effect on AH, for acetone fluorene phenylacetonitrile cyclopentadiene or acetyl-acetone B. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA (2) Proton magnetic resonance spectra of DMSO-DMSYL-d6 solutions of deprotonated acids corresponded in every way with DMSO-d6 solutions of the acid precursors except for the presence of the acidic proton. (3) Recovery experiments using liquid chromatography described above showed no decomposition for the compounds considered here. In view of the sensitivity of many of these compounds to base-catalysed reactions other than deprotonation this evidence is crucial.(4) A fourth piece of evidence (with many more ramifications) arises through comparison of our AHD data with the pK,’s for the same compounds determined in Bordwell’s laboratory. Although Bordwell’s measurements involve an indicator titration at roughly one-hundredth the base concentrations used by us the correlation of our results with his for 43 compounds has a correlation coefficient of 0.988 and a slope of 1.32f0.03 as shown in fig. 1. This slope corresponds within experimental error to 2.303 RTat 25” with the extraordinary implication that the relative standard free energies for deprotonation are exactly equal to their heats of deprotonation. This close agreement strongly supports the validity of both sets of measurements.pK vs AHEMSO -20--15--10-8 10 12 14 16 I8 20 22 24 26 28 30 32 PK FIG.1.-Correlation of pKa’s of Bronsted acids in DMSO (kindness of Prof. Bordwell) with heats of deprotonation in DMSO AH,. DISCUSSION The heat of deprotonation AH,(DMSO) measured in K+DMSYL-in DMSO is a physical property of singular value for comparing the Bronsted acidities of organic acids. Its particular virtues are (a)it is the only method currently available which can be applied directly without extrapolation or interpolation to the entire range of proton donors from strong oxygen acids such as carboxylic or phenolic ones whose pKa’s lie midway in the pH range to relatively weak carbon acids such as acetonitrile or diphenylmethane with pKa’s close to 34-the value for DMSO itself; (b)the FORMATION AND SOLVATION ENTHALPIES DMSO/DMSYL- system has been investigated extensively so that our results are directly relatable to pKa studies by Bordwell and Ritchie,lg to a host of kinetic and ion-pairing studies 2o and to an enormous variety of recent synthetic work; (c) it is relatively convenient to determine heats of solution of neutral organic acids in DMSO in order to calculate the solvation enthalpies of their conjugate base anions as we shall do here.In this section we shall comment briefly on the relationship between AHD(DMS0) values and two other criteria of acidity for the same compounds-their pKa’s in DMSO and their acidities in the gas phase.We shall then consider the factors which influence solvation energies of the anions by comparing them with the heats of anion formation in the gas phase. CORRELATION OF AHD(DMS0) WITH pKa (DMSO) Fig. 1 portrays the surprisingly good correlation between AHD and pKa for a rich variety of organic acids in DMSO. As noted the correlation is significant not only because of its generally clean linearity over a wide range of energy for many structural types but also because its slope implies equality between AH; and AG;. One consequence here is that AS& the standard entropy for deprotonation for these manifold species is nearly constant and probably equal to that for deprotonating DMSO since the deprotonation reaction is AH+K+DMSYL- K+ +A- +DMSO. In an earlier publication l5 we noted a rather good correlation between the first AHD values we determined and a few pKa’s in DMSO which were available at that time.It is both gratifying and surprising to us to find further experimentation not only supporting the original conclusions but strengthening them. At this point the only series of compounds which shows a systematic displacement from the line is the nitroalkanes. Cyclobutanone is also far from the line but as we have shown by the recovery experiment it undergoes a fast exothermic secondary reaction. CORRELATION OF AHD(DMS0) WITH AHD(gaS) Bordwell and coworkers 21 have recently reported a close correlation between the pKa’s of a number of carbon acids in DMSO and corresponding gas phase acidities reported by McMahon and Kebarle.22 In view of the correlation of our data with Bordwell’s a linear relation of AH,(DMSO) to gas phase acidities is required.Such a plot is shown in fig. 2 in which we include many new data graciously provided by Prof. Kebarle. Also included are some previously reported l4 gas phase data from McIver’s laboratory. It is important to recognize here that heats of deprotonation in the gas phase are usually equal within experimental error to standard free energies of deprotonation for isodesmic processes like those considered here. Kebarle’s values were obtained by high pressure mass spectrometry at 600 K while McIver’s acidities were determined by ion cyclotron resonance at 298 K. Taft 23 has shown the generally good agree- ment between data collected by the two groups.However some discrepancies do exist the most glaring being 3.7 kcal/mol for p-chlorophenol. In order to minimize artifactual differences between methods and observers all of the AHD(gas) values used here are taken arbitrarily from Kebarle’s work except for the aliphatic alcohols. The overall trend towards correlation of our results with the gas phase acidities is obvious from fig. 2. Furthermore we find as did Bordwell a partial separation of hydrocarbon acids from oxygen acids such as alcohols ketones phenols and carboxylic acids. The separation suggests (but does not require) an added exothermic E. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA term for the ionization of the latter compounds in solution.Special factors also seem to be affecting the ionization of the alcohols thus rendering them much more sensitive to structural change than the other compounds. 310 t -AHD(DMSO)/kcal mol-' FIG.2.-Plot of enthalpies of deprotonation in DMSO [AHD(DMSO)] against enthalpy of deproton-ation in the gas phase [AH~(gas)]. All values are in kcal/mol and gas phase data were provided by Prof. Kebarle and McIver. Correlation coefficient of points on top line = 0.990 and on bottom line = 0.998. Obviously the differences between acidity orders in solution compared to those in the gas phase can be caused by solvent interactions with the neutral precursors or the anions or both. It is important to resolve these factors since such knowledge can suggest the proper point of attack for manipulating proton transfer energies through variation of molecular structure or solvent.This we shall do in the next section. SOLVATION ENTHALPIES OF ORGANIC ANIONS Heats of solution of a number of typical organic anions from the gas phase to DMSO were calculated using the approach described in the introductory section of this paper. The necessary data for the calculation have been published recently 14* 24 for many of the anions treated here. However we have added several more for the present discussion. The solvation energy of an ion is that part of its energy in solution which derives from its net interactions with the solvent. It is informative to compare this external interaction of the ion with the energy required to form it in the absence of solvent- AH,(gas).In fig. 3 the two properties are plotted. Although the overall pattern is FORMATION AND SOLVATION ENTHALPIES a random scatter we have been so bold as to separate it into several lines which include ions with common structural features. The interpretation of the factors which lead to these groupings seems so reasonable that we do not hesitate to repeat them here. 8 -s-23 --E 0;- -50 -40 -30 -20 -10 0 10 20 30 4 I deprotonation energy SAH:/kcal mol-l FIG.3.-Correlations of solvation enthalpies of anions from the gas phase to DMSO [8AH:*DSMo (A31 with deprotonation enthalpies of corresponding acid precursors in the gas phase SA&(gas). All values referred to cyclopentadiene and cyclopentadienyl anion as standards.All ions except alkoxide ions fall on three lines which are mutually parallel as shown. On each of these lines the ions which enjoy the highest degree of charge delocalization because of resonance or large size have the least exothermic values of 8AH:-DMSo(A-). The series of ions whose solvation energies are the least are the hydrocarbanions at the top of the graph. Included with these are p-nitrophenolate and malononitrile anion whose charges must be very widely dispersed. The next line includes delocalized oxyanions of all sorts enolates carboxylates and phenolates. We believe it is significant that the carbanions and enolate anions of most localized charge fall at the intersections with the lines generated by the alkoxide ions whose charge is mostly localized on oxygen.We have rationalized the steep slope of the alkoxide ion line in terms of steric hindrance to solvation of the localized charge a factor which differentiates alkoxides from the other ions in the plot. In the interests of objective scepticism we also warn the reader that this series of ions was studied in the gas phase under different conditions from the others. Also it is possible but unlikely that in solution they are not completely ionized due to their notorious proclivity for homoconjugate ion formation RO--HOR.25 We doubt if these factors are a problem here. Finally we note that the halides which are relatively small and sterically unpro- tected have relatively high solvation energies and generate a line parallel to the two above it.Solvation energies of ions and molecules are generally related to their volumes-in E. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA the former case charge density and in the latter the energy of cavity formation are the chief variables. One might suppose in the present case that in keeping with electro- static principles the solvation energies of the organic ions considered here would be related to their reciprocal radii. Since the ions are generally non-spherical and their partial molar volumes unknown we have crudely estimated r-l for the ion through using the molar volume of the neutral acid precursor and then taking the reciprocal of its cube root. Fig. 4 shows that even with the many rough approximations in this approach a clear trend towards linearity is found.This gives added support to the general consistency of the methods and data used in deriving the solvation energies. 15 t ._ c -13- 0 EtOH rcl x m WE -17-2l- 0 MeOH cl .- crj u >- -25- % -29 - H2O -331 I I 1 0.I 0.2 0.3 0.4 FIG.4.-Correlation of anion solvation enthalpies with estimated reciprocal ionic radius. We are glad to acknowledge the helpful contributions of Prof. Bordwell Kebarle and McIver to this research. Supported by NSF grant G.P. 6550-X. On leave from the Department of Physical Chemistry University of Bucharest Romania. See E. M. Arnett N. J. Hornung and R. J. Minasz Proceedings of C.N.R.S. Conference on Water in Biological Systems Roscoff France June 1975 for recent references relating to these topics.B. G. Cox in Ann. Rep. Progr. Chem. A 1973 70. ’R. T. McIver JI. J. A. Scott and J. M. Riveros J. Amer. Chem. SOC.,1973 95 2706. R. T. McIver Jr. and J. H. Silvers J. Amer. Chem. SOC., 1973 95 8462. ’R. T. McIver Jr. and J. S. Miller J. Amer. Chem. SOC.,1974 96 4323. * D. K. Bohme E. Lee-Ruff and L. B. Young J. Amer. Chem. SOC.,1972 94 5153. D. K. Bohme E. Lee-Ruff and L. B. Young J. Amer. Chem. Soc. 1971,93,4608. FORMATION AND SOLVATION ENTHALPIES lo T.B. McMahon and P. Kebarle J. Amer. Chem. SOC.,1974,96 5939. P. Kebarle presented at Structure-Energy Symposium San Juan R.R. ; 1974. l2 P. Kebarle in Modern Aspects of Electrochemistry ed. B. E. Conway and J.O’M. Bockris (Plenum Press N.Y. 1974) 9 1. l3 P. Kebarle personal communication. l4 E. M. Arnett L. E. Small R. T. McIver Jr. and J. S. Miller J. Amev. Chem. SOC.,1974 96 5638. Is E. M. Arnett and T. C. Moriarity J. Arner. Clzem. SOC.,1971 93,4908; 1973 95 1492. l6 J. S. Chickos J. Chem. Ed. 1975 52 134. l7 G. W. Thomson in Technique of Organic Chemistry Vol. 1 Part 1 ed. A. Weissberger (Wiley Interscience New York 3rd Edn. 1965 ) p. 401. R. Fuchs and R. F. Rodewald J. Amer. Chem. SOC.,1973 95 5897. Also personal com- munication with Prof. Fuchs. l9 F. G. Bordwell Faraday Symp. Chem. SOC. 1975 10 100. Provides references to the relevant literature. 2o D. J. Cram Fundamentals of Carbanion Chemistry (Academic Press New York N.Y. 1965) ; M. Szwarc Ions and Ion Pairs in Organic Reactions Vol.2 (John Wiley and Sons New York N.Y. 1974). 21 F. G. Bordwell J. E. Bartmess W. S.Matthews,G. E. Drucker and Z. Margolin,J.Amer. Chem. SOC.,1975 97 3226. T. B. McMahon and P. Kebarle J. Amer. Chem. Soc. 1974,96 5940. 23 R W. Taft in Proton Transfer Reactions ed. E. Caldin and V. Gold (Chapman and Hall London 1975). 24 E. M. Arnett Dale E. Johnston and Leonard E. Small J. Amer. Chem. SOC., 1975 97. 25 J. H. Exner and E. C. Steiner J. Amer. Chem. SOC.,1974 96 1782.