J. Chem. SOC., Faraday Trans. I, 1986, 82, 2547-2556 Thermodynamic Study of Organic Compounds in Octan- 1-01 Processes of Transfer from Gas and from Dilute Aqueous Solution? Paolo Berti, Sergio Cabani," Giovanni Conti and Vincenzo Mollica Dipartimento di Chimica e di Chimica Industriale, Universita degli Studi di Pisa, Via Risorgimento 35, 56100 Pisa, Italy Free energies and enthalpies of solvation of water and some hydrocarbons (hexane, cyclohexane), ethers (diethyl ether, tetrahydrofuran) and ketones (propanone, pentan-3-one, cyclopentanone) in octan- 1-01 have been deter- mined at 298.15 K from vapour-pressure measurements of dilute solutions and from limiting heats of solution. These solvation functions in octanol have been used together with the corresponding hydration functions in order to obtain water-octan- 1-01 partition coefficients and their dependence on temperature.A comparison is made with the practical partition coefficients relative to mutually saturated solvents. Many thermodynamic studies have been carried out these last two decades for the transfer process of organic molecules from ideal gas to aqueous so1ution.l The obtained standard thermodynamic functions of hydration, AXE (X = G, H , S , C,), were related to the molecular structure of the solutes through a scheme of group contributions. This procedure allowed some features arising from the water-functional group interactions to be classified according to the nature, number and position these groups occupy in the molecular structure of the so1utes.la Much attention has also been devoted to the transfer of organic and inorganic non-electrolytes from gas to dilute non-aqueous solution,2 but unfortunately a systematic study on the thermodynamics of solvation of organic molecules in non-aqueous solvents is hitherto lacking, although it presents many interesting aspects. In fact, a knowledge of the values of contributions of many organic functional groups to the thermodynamic functions of solvation in many different solvents could improve our possibilities of predicting and understanding of the role of the solvent on the following: chemical equilibria, kinetics of reaction, gas solubilities, partitioning of a solute between immiscible solvents etc.As part of a planned study on the thermodynamics of solvation of non-electrolytes in non-aqueous solvents, we report in this paper some results concerning the free energies, enthalpies and entropies of transfer of some small organic molecules from gas to dilute solution in octan-1-01.Free energies of solvation have been obtained from liquid-vapour equilibria and enthalpies of solvation have been obtained from calorimetric heats of solution and from known values of heats of vaporization. Finally, entropies of solvation have been calculated from the values of the above free energies and enthalpies. These functions have been used, together with the literature values of the corresponding thermodynamic properties of hydration,l a for calculating the thermodynamic functions of transfer of solutes from pure water to pure octanol. Many values of free energy for the transfer process from octanol-saturated water to water-saturated octanol are available from experimentally determined partition coef- ficients, Poct,w.3 These data are of pharmacological and environmental interest for their t Taken in part from a Ph.D.Thesis (P.B.), Universita di Pisa, 1985. 25472548 Study of Organic Compounds in Octan-1-01 correlation with drug activity4 and the solubility in water of liquid and solid non-ionic compo~nds.~ A few values of enthalpy of transfer have also been reported, obtained from studies of the dependence of log(Poct,w) on temperature,6 as well as from direct microcalorimetric measurements carried out by flow techniq~es.~-l~ Free energies, enthalpies and entropies of the transfer process, evaluated by considering separately the solution processes in water and in octanol, allow us to gain much more information on the causes driving the distribution.Moreover, the possibility of investigating both the distribution between the pure solvents and the distribution between reciprocally saturated phases is an important step in understanding the effects of the mutual water-octanol solubility on the partitioning of solutes. For this purpose some calorimetric measurements have also been carried out to study the differences between the enthalpies of solution in pure octanol and in octanol containing known quantities of water. Literature data have been used to examine this effect on the free energy of transfer. Experimental Materials All examined solutes were commercial products of the best grade available.The ketones (propanone, pentan-3-one, cyclopentanone) were dried over anhydrous calcium sulphate. The alkanes (hexane, cyclohexane) and the ethers (ethyl ether, tetrahydrofuran) were dried over calcium hydride. All compounds were then fractionally distilled in an inert atmosphere. The purity of the fractions collected at constant boiling point was checked by g.1.c. and was in all cases better than 99.8%. The water used in the experiments was first deionized and then distilled from alkaline KMnO,. The solvent, octan-1-01, was a Fluka product of purissimum grade ( > 99.5 % ). It was refluxed over metallic sodium and then distilled under nitrogen at reduced pressure. The samples used showed a g.1.c. purity better than 99.9% and a water content < 0.01 % , determined by Karl Fisher titration.Vapour Pressure Measurements The total vapour pressure, pt, of the solutions was measured at 298.15 K with a precision of 2.5 Pa by a static apparatus described in aprevious paper.ll Particular care was taken in deaerating the solutions by repeated freezings and meltings under vacuum. The partial pressure of solutes, P,, was then calculated using the relationship : P, = pt - (1 - X,) Po,,, where X , and Pact, are the solute mole fraction and the vapour pressure of pure octan- 1-01, respectively. Po,, was determined in separate experiments using octanol samples treated as described above. The measurements were made in the temperature range 293-303 K. At 298.15 K we found Poet = 14.5 Pa and its derivative with temperature (dPoct/dT)298 = 1.43 Pa K-l.These data are in good agreement with literature data.I2 The composition of the liquid phase in equilibrium with vapour was determined, according to the various solutes, by U.V. absorbtion or by density measurements of the mixtures. The density values were obtained using an Anton Paar digital vibrating density meter (DMA 602) with a precision better than 3 x lop3 kg m-3. Heats of Solution The heats of solution in octanol for the compounds considered were measured by an isoperibol calorimeter built for this purpose. The calorimetric cell was a glass bottle of ca. 100 cm3 volume, with a very reduced vapour space (< 1 cm3). It was contained in a copper cylindrical vessel within a sophisticated air circulation thermostatic bath, capable of maintaining the temperature within f 5 x K even for long periods of time.Successive additions of solutes were made through an external digital microburette having a total volume of 1 cm3 and a resolution of lop4 cm3. Before entering the cell,P. Berti et al. 2549 the liquids flowed through a heat exchanger constituted by a loop of PTFE tube wound around a massive brass cylinder situated beside the cell. A miniaturized relief valve at the end of the dispensing tube prevented diffusion of the solute in the solvent during intervals between successive runs. The temperature changes in the cell caused by addition of solutes were read by means of a thermistor (5000 R at 293 K) connected to a Wheatstone bridge. Signals from an unbalanced bridge, via an AD converter, were transferred to a AIM 65 Rockwell microcomputer for the collection and the elaboration of calorimetric data.The heat capacity of the system was obtained for each run by repeated electric calibrations. The stirring inside the calorimeter was maintained constant during the experiments. It was possible, however, to change the stirring speed in a wide range, using a stepping motor, in order to obtain rapid mixing with low noise. Repeated experiments allowed us to evaluate a A T resolution within 5 x lop5 K, corresponding to 1 x J on the heats, owing to the heat capacity of our system. The accuracy of the apparatus was tested by neutralization experiments using aqueous HC1 and NaOH solutions. For each compound at least 10 measurements were carried out in the concentration range 0.002-0.04 mole fraction.No appreciable dependence on concen- tration was observed for all the compounds examined and the uncertainty on the limiting heats of solution did not exceed 0.10 kJ mol-l. Results In table 1 are reported the thermodynamic quantities related to the isothermal transfer at 298.15 K of a mole of solute from the ideal-gas state to infinitely dilute solution. The free energy change is associated with the process: S (ideal gas, C, = 1 mol dm-3) + S (ideal solution in octan-1-01, Coct = 1 mol dm-3) where S is a generic solute, Cg and Coct are the molar concentrations of S in the two states. The standard free energy of solvation has been calculated by extrapolating at Xs = 0 the function : At least five different values of AGb,, have been determined for X s ranging from 0.01 to 0.15 for each solute examined.Fig. 1 and 2 show, in this concentration range, a linear relationship between AGict and Xs, which allowed us to obtain extrapolated AG$t values with an uncertainty of ca. 0.17 kJ mo1-l. The limiting value, AGZCt, corresponds to the transfer of a mole of solute from an ideal gaseous state at P = 1 atmt to a hypothetical ideal solution state in octan-1-01 at X s = 1. This quantity was then converted to the standard free energy of solvation, AG:ct, relative to process (1) by: (1) AG;,, = RT In (Ps/Xs). (2) AGEct = AGzct - RT In (RT) - RT In (Poct/M,,,). (3) A value of poet = 821.72 kg mV3 (determined by us) for the density of pure octanol at 298.15 K was used, in accordance with the value of 821.77 kg m-3 reported in ref.(1 3). In eqn (3) Moct represents the molecular weight of octan-1-01. The enthalpies of solvation, A g c t , have been calculated at 298.15 K using the relations hip : (4) where AKol, is the limiting heat of solution in octanol and AWvap is the enthalpy of vaporization. These latter values have been taken, except for water,14a from ref. (14b). A K c t = A%ol* - m a p t 1 atm = 101 325 Pa.2550 Study of Organic Compounds in Octan-1-01 Table 1. Thermodynamic standard functions of solvation, AXEct ( X = G, If, S ) and heats of solution, AKoln in octan- 1-01 at 298.15 K" ~ ___ Solute AGct W C t - TAS0,ct A f C O h water hexane cyclohexane ethyl ether tetrahydrofuran propanone pentan-3-one cyclopentanone - 16.48k0.17 - 13.89f0.08 - 15.56 f 0.04 - 12.35k0.17 - 16.18f0.21 -12.93f0.13 -17.99kO.13 - 20.88 f 0.21 -40.57 k0.08 - 30.79 f 0.08 - 31.55 f 0.06 - 24.86 0.36 -28.32 +O.15 -22.37k0.16 - 32.86 k 0.21 - 34.34 0.24 24.09 k 0.25 16.90 f 0.16 15.99 f 0 . 10 12.51 k0.53 12.14f0.36 9.44 f. 0.29 14.87 k 0.34 13.46 0.45 3.44 f 0.07 0.76 f 0.04 1.50 & 0.02 2.29 k 0.1 1 3.68 0.06 8.47 & 0.12 5.72 f 0.08 8.38 k 0.06 a All data are in kJ mol-l. AGEct refers to the transfer of solutes from 1 mol dm-3 ideal gas state to 1 mol dmP3 hypothetical ideal solution state in octan-1-01. AWoct is the calorimetric enthalpy change for the transfer from ideal gas to infinitely dilute solution. T A g c t is obtained from - T A g c t = AGZct - A K c t . The standard entropies for process (1) have been finally calculated from the free energies and enthalpies of solvation using the well known relationship : A x k t = (Wet -AG&t)/T* Our value of the free energy of solvation of hexane is in good agreement with the value of AG:,, = - 13.56 kJ mol-1 reported by Abraham.2b An acceptable agreement also exists between our AG:ct value for water and those calculated by using the solubilities of water in the octanol phase equilibrated with aqueous saline solutions of known water activity, reported by A ~ e l b l a t .~ In effect, this author considered solutions in octan- 1-01 of higher water content than ours. However, both our data and his may be satisfactorily described by the same quadratic equation (fig. 2), which gives an intercept (AG$, = -4.01 kJ mol-l) very near to the value obtained from our data alone (AG,*,, = -4.02 kJ mol-I).An appreciable difference is instead found between the value AGZct = - 15.56 kJ mol-1 obtained by us for cyclohexane and the value calculated from data reported by Abraham AG:,, = - 13.26 kJ mo1-1.2a In this latter comparison the results are affected by the fact that Abraham considered the free energy of partitioning referred to a system whose phases are water-saturated octanol and octanol-saturated water. The effect of the mutual water-octanol solubility on the partition coefficients is showed by fig. 3, where the AGf-,,,, data, obtained from direct measurements of partition coefficients, are compared with the values of standard ideal free energy of transfer AG;,,,, calculated from the free energies of solvation in the pure phases: It may be seen that differences of ca.1-2 kJ mol-l are common and may be positive or negative according to the nature of the solute. No comparable data are reported in literature for the enthalpy of solvation in octanol for the compounds examined here. As far as the dependence of the enthalpy of solvation on the water content of the octanol solvent is concerned, some experiments were conducted for propanone. In fig. 4 is reported the trend of the limiting heat of solution of propanone when octanol (with increasing amounts of water) is used as a solvent. The difference between the heat of solution in octanol and in water-saturated octanol is 2.8 kJ mol-l. Noticable differences are also exhibited by the phenol derivatives studied by Beezer et aZ.,7 whereas Riebesehl did not find any appreciable difference in the enthalpy of solution of n-alkanols in pure and water-saturated octan01.~ The influence of water content on the thermodynamic functions of solvation of organicP.Berti et al. 2551 Fig. 1. AGb,, [eqn (2)] us. the solute mole fraction in octan-1-01, X,. Solute: (a) cyclohexane, (b) hexane, ( c ) cyclopentanone, (d) pentan-3-one, (e) THF, (f) ethyl ether. compounds in octanol will be considered elsewhere. In this paper we will limit ourselves to a first comparison between the thermodynamics of solvation in octan-1-01 and in water and to deduce their role in the transfer process of solutes between these two pure liquids. Discussion In fig.5 are reported the standard thermodynamic functions of solvation in octan-1-01, together with the values of the thermodynamic functions of hydration, AX:, and the standard thermodynamic functions of transfer of solutes from an ideal 1 mol dmP3 solution in water to a hypothetical ideal 1 mol dm-3 solution in octanol. Values of AX; (j = h, oct, w -+ oct) for methane, obtained from studies of solubility at various temperatures, are also reported.17 A first glance analysis of this figure allows us to recognize that: (a) Except for hydrocarbons in water and methane in octanol, the solution is more stable than the gas phase. (b) The stabilization of the solution is due, for both solvents, to an enthalpic effect, which prevails on the unfavourable entropic term.(c) The order of the compounds on2552 Study of Organic Compounds in Octan-1-01 \ \ I I I Fig. 2. AGA,, us. A', for (a) propanone and (b) water: 0, this work; ., calculated from ref. ( 5 4 . the scale of increasing enthalpies of solvation in octanol differs from the order found on the scale of the corresponding entropies. The same behaviour is observed when the solvent is water. This means that the rule of enthalpy-entropy compensation is not strictly observed.lB ( d ) The order of the compounds in the scales of the functions of solvation in octanol is different from that found in the corresponding scales of hydration. Obviously, this is a consequence of the different interactions the polar and the non-polar groups have with the two solvents considered. ( e ) As far as the hypothetical partitioning between pure octanol and pure water is concerned, the octanol phase is favoured in respect to the aqueous state, except for propanone.The reason lies in the positive entropic effect, which prevails on the positive value of the enthalpy of transfer. Riebesehl found an endothermic enthalpy of transfer from water to octanol even for the n-alkanols studied by flow-microcalorimetry.9 This behaviour is not exhibited by phenol derivatives, which show negative enthalpies, as shown from measurements of heats of solution in water and in octanol carried out by Beezer et al.' and Haberfield et aZ.,lo or from the dependence of experimental partition coefficients on temperature.6 c $ d * l9 From fig. 6 we find that the entropies of solvation in octanol and in water are linearlyP.Berti et al. 2553 -30 - 20 -70 0 A G: - OCt /kJ mol-* Fig. 3. Free energy of transfer from octanol-saturated water to water-saturated octanol, AGF +,,., us. free energy of transfer from pure water to pure octanol, AG;,,,,. Points 1-8 are as follows: 1, methane; 2, ethane; 3, hexane; 4, heptane; 5, octane; 6, cyclohexane; 7, pentan-3-one; 8, tetrafluoromethane. AGf,,,, values were calculated from partition coefficient data [AGf,o,t = --RTln (Poctlw)] taken from ref. (2b), (15) and (16). AG~,,,, values were obtained using eqn (5) from AGZct and AGE data taken from ref. (1 a ) and from this work and ref. (2b), respectively. c -. ---aw- l I I 1 x w 0 .I 0.2 0.3 Fig. 4. Limiting values of the enthalpy of solution, A g o , , , of propanone in an octan-1-01-water mixed solvent us.the water mole fraction, X,. related to the intrinsic volume of solute molecules, estimated as the van der Waals volume, Vw.20v 21 Contrary to what happens in water, the free energies and enthalpies of solvation in octanol are mainly related to the size of solutes. Water is an evident exception, in fact it exhibits a large affinity for octanol, in spite of its small van der Waals volume and its substantially hydrophobic environment. To this purpose, it may be noted that the solubility of water in 2,2,4-trimethylpentane is Xw = 8.76 x 10-4,22 whereas in octan-1-012554 40 Study of Organic Compounds in Octan-1-01 - AH" 20 - I .-. E c, -. A Y 0 - 2 c -20 -40 - A G O TAS" 7.2.1 4';!6 ."~' 3 1 I Fig. 5.Standard thermodynamic properties of solvation, AX; ( X = G , H , S ) , in octan-1-01, ( j = oct), in water, ( j = h), and of transfer from water to octanol, AXO,,,,, (= AXZCt-AX;). 1, ethyl ether; 2, THF ; 3, hexane; 4, cyclohexane; 5, propanone ; 6, pentan-3-one; 7, cyclopentanone; 8, methane. AXE,, data are taken from this work, except for methane, ref. (1 7); AX; data are from ref. (1 a). Hatched areas show AX:ct, shaded areas show AXO,,,,, and unfilled areas show AX;. X, == 0.275.5d The high solubility of water in octanol may be justified by the formation of a tetrahedral complex in which a molecule of water is hydrogen-bonded to four hydroxy groups of octanol m01ecules.~~ The formation of this species is accompanied by a large heat of solvation, which is similar in magnitude to the enthalpy of vaporization of pure water.This enthalpic effect prevails on the unfavourable entropy of solvation, which is, however, noticeably more negative than the entropy of solvation of the organic compounds examined here. Although less evident, we can however recognize the different behaviour of hydro- carbons in octan-1-01 from that of the other organic compounds of similar size but containing heteroatoms, the former showing less stability. Cyclisation increases the stability of the compounds in octanol by ca. 3-4 kJ mol-1 and is due to an enthalpic effect, being the entropy of solvation of the open-chain and their analogous cyclic compounds practically the same. A more detailed analysis in terms of the contributions that the polar and non-polar groups bring to the thermodynamic functions of solvation in octan-1-01, is at the moment premature because of the few experimental data available.This work has been supported by Minister0 della Pubblica Istruzione (Roma).P. Berti et al. 2555 vw Fig. 6. Thermodynamic properties of solvation in octan-1-01, AXct, and in water, A%, ( X = G, H , S ) us. the intrinsic molar volume, V,. AX,",, data are taken from this work and ref. (2b), AX: data from ref. ( l u ) . The V, values were estimated according to ref. (20) and (21). ., water; 0, alkanes (methane, ethane, cyclohexane, hexane, heptane, octane); 0, ketones (propanone, cyclopentanone, pentan-3-one); 0, ethers (THF, ethyl ether); (A), octan-1-01. (a) AHgct, (b) AH:, (4 TAgct, (4 TAG, (4 AGiCt, (f) AGE.References 1 See e.g.: (a) S. Cabani, P. Gianni, V. Mollica and L. Lepori, J . Solution Chem., 1981, 10, 563; (b) E. Wilhelm, R. Battino and R. J. Wilcock, Chem Rev., 1977, 77, 219; (c) M. H. Abraham, J . Chem. Soc., Faraday Trans. I , 1984, 80, 153. 2 See e.g.: (a) M. H. Abraham, J . Am. Chem. Soc., 1979, 101, 5477; (b) M. H. Abraham, J . Am. Chem. Soc., 1982, 104, 2085; (c) C. L. De Ligny, N. G. van der Veen and J. C. van Houwelingen, Znd. Eng. Chem. Fundam., 1976,15, 336; ( d ) C . V. Krishnan and H. L. Friedman, in Solute-Solvent Interactions, ed. J. F. Coetzee and C . D. Ritchie (Marcel Dekker, New York, 1969), vol 11, chap. 9, pp. 1-103; (e) E. Wilhelm and R. Battino, Chem. Rev., 1973, 73, 1 ; (f) H. L. Clever and R. Battino, in Solution and Solubilities, Techniques of Chemistry, ed.M. R. Dack (John Wiley, New York, 1975), vol. 8, part 1, chap. 7. 3 See e.g.: (a) T. Fujita, J. Iwasa and C. Hansch, J . Am. Chem. Soc., 1964, 86, 5175; (b) C. Hansch and S. M. Anderson, J , Org. Chem., 1967,32,2583; (c) A. Leo, C. Hansch and D. Elkins, Chem. Reu., 1971, 71, 525; ( d ) A. Leo, J . Chem. Soc., Perkin Trans. 2, 1983, 825; (e) I. Moriguchi, Y. Kanada and K. Komatsu, Chem. Pharm. Bull., 1976, 24, 1799.2556 Study of Organic Compounds in Octan-1-01 4 See e.g.: C. Hansch, Acc. Chem. Res., 1969, 2, 232; (b) R. F. Rekker, The Hydrophobic Fragmental Constant (Elsevier, Amsterdam, 1977); (c) N. H. Anderson, M. James and S. S. Davis, Chem. Znd., 1981, 677; ( d ) C. Hansch and A. Leo, Substituent Constants for Correlation Analysis in Chemistry and Biology (John Wiley, New York, 1979); (e) R.N. Smith, C. Hansch and M. M. Ames, J. Pharm. Sci., 1975, 64, 599. 5 See e.g. : (a) C. Hansch, J. E. Quinlan and G. L. Lawrence, J . Org. Chem., 1968,33,347; (b) Y. B. Tewari, M. M. Miller, S. P. Wasik and D. E. Martire, J. Chem. Eng. Data, 1982, 27, 451; (c) D. Mackay, A. Bobra and W. Y. Shiu, Chemosphere, 1980, 9, 701 ; ( d ) A. Apelblat, Ber. Bunsenges. Phys. Chem., 1983,87,2; (e) M. M. Miller, S. Ghodbane, S. P. Wasik, Y. B. Tewari and D. E. Martire, J. Chem. Eng. Data, 1984,29, 184; (f) Y. B. Tewari, D. E. Martire, S. P. 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