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Proton-transfer equilibria in isodielectric acetonitrile–ethylene glycol mixtures at 298.15 K |
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Journal of the Chemical Society, Perkin Transactions 2,
Volume 1,
Issue 9,
1979,
Page 1208-1213
Kumardev Bose,
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摘要:
1208 ,J.CS. Perkiri I1 Proton-transfer Equilibria in lsodielectric Acetonitrile-Ethylene Glycol Mixtures at 298.15 K By Kumardev Bose and Kiron Kumar Kundu,' Physical Chemistry Laboratories, Jadavpur University, Calcutta- 700032, India The dissociation constants (K,) of tris(hydroxymethy1)methylammonium (TrisH4 ) and p-nitroanilinium (pNAH +) ions and p-nitrophenol (pNP) have been determined at 298.15 K both in ethylene glycol (EG) and its isodielectric mixtures with acetonitrile (ACN), an e.m.f. method being used for TrisH+ and a spectrophotometric method for the other two acids. These K, values have been used to derive the solvent effect on the free energies of dissociation of these acids. From independent solubility measurements on the neutral species : tris( hydroxymethyl) methyl- amine (Tris), p-nitroaniline (pNA), and pNP, the free energies of transfer of these species from EG to ACN-EG mixtures have been determined at 298.15 K.By a combination of the results of K, and solubility measurements and with a knowledge of the free energies of transfer of the hydrogen ion evaluated earlier, the free energies of transfer of TrisH+, pNAH+, and pNP- (p-nitrophenoxide ion) in these solvents have been computed. Plots of free energies of transfer for the individual species against solvent composition reflect a wide range of solvation behaviour which has been interpreted in terms of the tendency of protic solvents like EG to solvate specifically hydrogen- bonding solutes and the tendency of dipolar aprotic solvents like ACN to solvate specifically solutes undergoing primarily dispersion interactions, DESPITEextensive studies,l * the nature of tlic solveiit effect on the dissociation of weak acids and bases of various charge types is still iiiconipletely understood.For example, the observed minima in the dissociation constants of BH+ type acids in several mixed aqueous solvents 7u*8-12 has often incorrectly been attributed 5713 to the increased basicity of the mixed solvent resulting from a co-solvent-induced breakdown of water structure. It has become increasingly clear no~,'J~ however, that the apparent1y intriguing behaviour of dissociation constants results from the superimposition of the iiiediuni effects of the tliree species participating in tlie dissocia- tion equilibrium and that to understand the behaviour of the composite quantity entails an understanding of the behaviour of the individual species.The dissociation of a protogenic species HA(z+ l)' in a protic solvent SH involves the transfer of a proton to the solvent which acts as Broiisted base lequation (l)]: HA(*+ I)+-+ SH SH,+ + .Az' (1) where the charge z may be positive, negative, or zero. The thermodynamic dissociation constant K, is given by equation (2) where ai denotes the activity of the species i and for simplicity the charges have not been shown. The common problem encountered in the iriterpreta- tion of such acid-base equilibria in mixed solvents is the uncertainty of the ' electrostatic part' of the medium effecton the dissociation process, arising from the differ- ence in dielectric constants of the solvents, particularly in the dissociation of uncharged acids where charge separation occurs.For this reason, studies of acid-base equilibria in isodielectric solvents should be specially informative, since here only the ' chemical ' nature of solvation of the reacting species is reflected in the observed medium effect. In previous papers,15 we have studied the solvent effects on the free energies of solvation of a number of electrolytes and ions in ethylene glycol (EG) and in its isodioelectric mixtures with acetonitrile (ACN), and Iiave derived useful information regardiiig ion-solvation in these media. On the basis of tliese studies we are now able to exanline tlie effect of this unique series of solvents on thermodynaiiiic processes such as proton-transfer equilibria in terms of the related solute-solvent inter-actions.Here, we report the measurement of the dis- sociation constants of three weak acids, representing two charge types, in EG and its isodielectric mixtures with ACN. These, taken with associated solubilit>- measure- iiients on the neutral species in these solvents and the free energies of transfer of 13 ' evaluated earlier,15(f provided the free energies of transfer of tlie other species participating in the dissociation equilibria. An insight regarding the transfer behaviour of tlie individual species concerned in the solvents should facilitate an under-standing of the overall medium effect on the proton- transfer equilibria.Of the three acids chosen, two are of the cationic type, z~iz,,the protonated fornis of tris(liydroxymethy1)- niethylaiiiine (Tris) arid p-nitroaniline (pNA), while the third, $-nitrophenol (pNP) , is neutral. An important distinction between pNA (or pNAH ') and Tris (or TrisH+) is that while the first may undergo strong dispersion interactions with the solvent, the second should be solvated primarily through hydrogen-bonding. In the present solvent system, where one solvent coin- ponent (EG) is protic whereas the other (ACN) is dipolar aprotic, specific solvation of these solutes by one com-ponent or the other should give rise to some interesting trends in their respective inediuin effects.The dissociation constants were determined from e.1n.f. measurements for TrisH + , and from spectrophotornetric measurements for pNAH' and pNP, while solubility measurements were carried out on the neutral compounds Tris, pNA, and pNP. The mixed solvents contained 20, 40, and 60 percent by weight of ACN, and all measure- ments were performed at 298.15 * 0.05 K. EXPERIMENTAL The purification of the solvents EG and ACN has been described earlier.15n Tris (Fischer primary standard grade 1979 was used after recrystallization from 75% methanol. Tris-hyrlrocliloride (‘Tris-HC1) was prepared by the slow addition of c‘o~ic.HCl (Analalt, E.T).H.) to a saturated aqueous solution of Tris; tlie precipitate was filtered off and re-crystallized three times froni water.pNA (Rietlel) was twice recrystallized from 90% ethanol. pNP (Rietlel) was recrystallized from benzene until perfectly white. -411 compounds were dried in vacuo. Since pNP slowly de- composes on exposure to light and air, the product used was always freshly crystallized and dried. Dissocintio.~zConstants.-The electronietric method used for the determination of the dissociation constants of TrisH’ was essentially similar to that used for the same acid in water 14c and in 5001; E.1n.f. measurements were performed on the cell where nz, and ~n,are the molali- Pt, H, (g, 1 atm)/?’ris-HCl (ml),Tris (m2), solvent/AgC1-.4g (A) ties of Tris-HCl and Tris respectively.The ratio m, : ma was kept constant for a given solvent, while the ionic strength was varied. E.1ii.f. measurements on these solutions were performed in a manner essentially similar to that earlier described.15“ Equilibrium was reached in 7-8 Ii on the average. No poisoning of the hydrogen electrodes was observed. The spectrophotometric method employed for the deter- mination of the dissociation constants of pNXH + resembled that used by Bates and his co-workers.*& For pNAH’, a series of solutions having varying ionic strengths were pre- pared by mixing weighed amounts of standard solutions of pNA and HC1 (G. R.Merck) with the appropriate solvent. From preliminary iiieasurenients it was found that the wavelength of maximum absorbance for pNA in these solvents is 380 nm, $0all measurements were performed at this wavelength.The optical absorbances were measured in a 1.00 cni cell using a Beckinann DLJ (G 2400) spectro-photometer. The limiting ahsorbsnces oi the untlis-sociated fornis of pX-A were determined by measurement on pNA solutions containing small amounts of concentrated solutions of HCl and sodium lyate respectively. The absorbance of the undissociated foriii, i.e. pXAH+, was found to be negligibly small at 380 nm and was therefore taken to be zero. For pNP, the solutions for measurement contained pNP, Tris, and Tris-HCl dissolved in the appropriate solvent. The ionic strength was varied by varying the concentration of Tris-HC1, while the buffer ratio WZT~~~-HC’~: WIl*,is was kept at a fixed value.The spectrophotometric measurements were essentially similar to those for pNA, the wavelength chosen being 405 nm at which the dissociated foriii of pNT’ showed the iiiasiinuni absorbance. SoZubilities.--For the determination of the satura tecl solubilities of Tris, pNA, and pNP, a small quantity 01 each solute was adclecl to the appropriate solvent taken in a Jena bottle, shaken and allowed to equilibriate in a thermo-stat at 298.15 f0.05 I<. Aliquois of each solution were withdrawn at 48 h intervals and estimated after appro- priate dilution with a large voluine of water so that the medium was essentially aqueous. For pNA and pNP, the estimation was performed spectrophotoinetricaIly at 380 and 820nm respectively, the corresponding molar absorp- tivity values being 1.28 x 101 and 9.72 x 103respectively.Tris was estimated by titration with standard aqueous HCI using Methyl Orange as indicator. The saturation equilibrium was generally reached in 7-10 days, when the solutioii concentrations showecl no further change. RESULTS The e.ni.f. values (I?)for the cell (A) in different solvents, duly corrected to 1 atm partial pressure of hydr0gen,l5~ are listed in Table 1 with tlie corresponding values of mI,m2, TABLE1 Neccssary quantities for the evaluation of PK’B=+ for Tris in ACN-EG niistures at 298.15 K %l mol kg-.’ %lniol kg-1 din01 dm-3 El\.’ ~~K’BH+ 20% ACN 0.004 90 0.004 96 0.005 11 0.7713 10.49 0.009 64 0.009 64 0.009 93 0.7600 10.53 0.0144 0.0144 0.0148 0.7529 10.53 0.0193 0.0193 0.0199 0.7488 10.64 0.0241 0.0141 0.0248 0.7450 10.53 0.0340 0.0340 0.0350 0.7393 10.50 0.0388 0.0388 0.0400 0.7380 10.51 4070 ACN 0.005 27 0.006 21 0.00504 0.7550 10.62 0.0103 0.0102 0.009 85 0.7437 10.60 0.0208 0.0204 0.0197 0.7313 10.63 0.0258 0.0256 0.0247 0.7281 10.64 0.0309 0.0300 0.0295 0.7264 10.65 0.036.’ 0.0358 0.0346 0.7230 10.63 0.0417 0.0412 0.0399 0.7202 10.61 6076 ACN 0.005 25 0.005 20 0.004 67 0.7369 10.85 0.0166 0.0166 0.0148 0.7150 10.83 0.0222 0.0222 0.0198 0.7142 10.89 0.0279 0.0280 0.0248 0.7101 10.88 0.0330 0.0337 0.0299 0.7090 10.90 0.0393 0.0394 0.0350 0.7058 10.88 0.0449 0.0450 0.0899 0.7007 10.82 and til where p is the ionic strength given by p = m,.d,, ds being the solvent density. An extrapolation function pI<’1513+ defined t>ythe equation was constructed. Values of j?I\I’BE[+ :-(E -E,$)/h 111 10 + log m, -1 log (II$*/?M.2) -2AoBods-)p~ (I 1 u,Bods--~p?)-l/ln10 -2 In(1 4-0.002 M,d<1p.) = PKBH+-tf(p) (3) Em” (the standard potential of the Ag-AgC1 electrode)] Ms (the average molecular weight of the solvent), d,, and the characteristic solvent parameters A,Bo and B, were taken from a previous and uo = 0; KB=+is the dissocia- tion constant of TrisHf. For pNAHkl a similar extrapolation function may be clefinecl by equation (4) where msH,t is tlie niolality of HC1 in pIi’l<HI = -log Tj/slf2 t -log [2t/(uI,-~)]= PKBH++fW (4) the solution and tlie ionic strength is given by u = msH,+ . d,; 1.1 = D/m and TL~,-= Db/?nb? where pn is the total niolality of pNA in any solution and vnb that in the solution containing completely dissociated pNA ; D and Db are the respective optical absorbances of these solutions.The values of ?%SH1t, mind(i.e. 1% or mb)]v, and D are given in Table 2. For pNP, an analogous extrapolation function ph”ah may be defined IIIc’1I.Z = PICR,H+ + log (~~zU,/fl~B,H+)-1% I(u -ua)/(ua-41 + 2 A,B,d,-~p~(1+ a,Bods-*~*)-l/ln10 4-i’k In (1 $-0.002 M,dL1 p) == PKHA4-f (p) (5) TABLE 2 Kecessary quantities for the evaluation of pKIHH+ for pNA in ACN-EG mixtures at 298.15 K 105mind mol kg-1 D PK'BH+ Pure EG 0.0100 0.0111 7.27 0.591 2.032 0.0196 0.0218 7.36 0.382 2.060 0.0399 0.0443 14.57 0.397 2.113 0.0569 0.0632 15.47 0.305 2.122 0.0693 0.0769 15.90 0.267 2.133 0.0778 0.0864 15.49 0.225 2.134 Completely basic 7.06 1.190 20% ACN 0.002 68 0.002 76 7.71 1.100 1.977 0.004 98 0.005 14 7.69 0.910 2.012 0.0101 0.0104 7.59 0.662 2.018 0.0201 0.0207 7.66 0.440 2.023 0.0296 0.0305 7.58 0.304 2.069 0.0498 0.0513 7.54 0.188 2.0965 Completely basic 7.54 1.350 40% ACN 0.005 25 0.005 02 7.68 0.748 2.080 0.009 94 0.009 50 7.66 0.529 2.117 0.0198 0.0189 7.64 0.323 2.143 0.0301 0.0288 7.78 0.244 2.130 0.0402 0.0384 7.62 0.193 2.119 0.0490 0.0469 7.54 0.144 2.174 Completely basic 7.68 I.220 60% ACN 0.004 93 0.004 39 6.15 0.546 2.110 0.009 94 0.008 84 6.17 0.378 2.139 0.0198 0.0176 6.13 0.227 2.169 0.0303 0.0270 6.15 0.163 2.170 0.0403 0.0359 6.15 0.125 2.184 0.0506 0.0450 6.15 0.102 2.185 Completely basic 6.13 0.890 as where u and ub have similar significances as in equation (4), and u, = D,/nz,, nz, is the total molality of pNP in the solution containing the completely undissociated form, and D, is the absorbance of this solution. KH~and KBlII+ are the dissociation constants of pNP and TrisH+; wzgl and P.vB,H+ are the molalities of Tris and TrisH+ in any solution.Table 3 lists the values of p, rnillc1, D,and as obtained by taking a, = 0. Dissociation constants (pK,) of the three acids deter- mined from the respective pK' us.f~.plots extrapolated to f~.= 0 are given in Table 4 which also shows the pK, value for TrisH+ in EG which was not determined in this study but was obtained from literature.l6a For comparison, the pK value for pNAH+ in EG obtained in an earlier study 7a is also shown and is seen to agree with the present value within 0.02 unit. The uncertainty in these pK values is about j0.02 unit. The saturated solubilities (mi)in niol kg-l at 298.15 I< for Tris, pNA, and pNP are presented in Table 5. These are correct to within &2%. The standard free energies of dissociation of each acid in a mixed solvent relative to pure EG, on the mole fraction scale, were calculated at 298.15 K by equation (6) where AGto(A -B)sys= (In 10) BT [p(sKa) -P ((&a)] -RT In (Ms/MG) (6) and sK, are the dissociation constants in pure EG and the mixed solvent respectively, and MG and Ms are the molecular weight of EG and the average molecular weight of the mixed solvent respectively.15a The standard free energies of transfer on the mole fraction J.C.S. Perkin I1 scale at 298.15 K of the neutral species Tris, pNA, and pNP were calculated by equation (7).Here the ratio of the AGbo (i) = RiT In (aImj/smj) -RT In (MsIMa) (7) activity coefficients of i in EG and in the mixed solvent has been assumed to be unity; the same assumption has been used by Rates and his co-workers 8c and Kundu and his co- workers. TABLE3 Necessary quantities for the evaluation of pWHLk for pNP in ACN-EG mixtures at 298.15 K p/mol ~lrn-~ Cornpletely acidic 0.005 22 0.009 89 0.0196 0.0295 0.0393 0.0489 Completely basic Completely acidic 0.004 82 0.0101 0.0196 0.0296 0.0396 0.0492 Completely acidic Completely acidic 0.004 87 0.009 73 0.0195 0.0295 0.0396 0.0493 Completely basic Completely acidic 0.005 03 0.009 7G 0.0198 0.0297 0.0398 0.0499 Completely basic 105mind/ niol kg-l D PK'HA Pure EG iu/??2T1iaHi = 0.11 7 8.54 0.930 8.32 0.798 10.51 8.45 0.780 10.45 8.50 0.751 10.45 8.49 0.739 10.47 8.59 0.722 10.50 8.54 0.704 30.52 8.49 0.168 20% ACN mTris/nzTrisH+0.999= 7.00 0.725 6.99 0.446 10.76 6.98 0.415 10.75 6.97 0.387 10.77 6.98 0.372 10.81 7.04 0.372 10.86 7.08 0.370 10.90 6.9G 0.116 40% ACN ~T~~~/NZT~~~H+= 1.603 7.59 0.350 7.62 0.254 11.27 7.69 0.241 11.27 7.64 0.234 11.31 7.65 0.232 11.38 7.72 0.226 11.40 7.67 0.227 11.47 7.63 0.064 60% ACN 1'IZ~r,~/lM~r~i~~.+= 1.497 20.54 1.850 20.49 1.475 11.75 30.55 1.480 11.73 20.45 1.310 11.71 20.61 1.275 11.74 20.64 1.200 11.72 20.79 1.175 11.75 20.35 0.256 TABLE4 pK,, Values for various acids on the molal scale in ACN-EG mistures at 298.15 I( 0 Wt."/o ,4CN 10.475 a TrisH+ 2.04pNAH+ 10.47 PNP 20 10.52 2.00(2.06 I) 10.7 1 40 10.62 2.09 11.22 60 10.86 2.12 11.72 a Ref.16. Ref. 7a. TABLE5 Solubilities wzi /mol kg-l of some neutral species in ACN-EG mixtures at 298.15 K Wt. % ACN Tris PNA PNP 0 0.665 0.843 2.87 20 0.497 0.770 4.45 40 0.281 1.36 6.01 GO 0.132 1.92 6.49 1979 Finally, using the AGtO (A-B)sys and AGtO (i) values just determined and the AGtO (Hi) values reported in a previous paper,lSd the values of the standard free energy of transfer at 298.15 K on the mole fraction scale of the conjugate acids of Tris and pNP and of the conjugate base of pNP were computed using relations (8) and (9). The values of AGto Acto (BHt-B)sp 1 AGO (B) + AGln (HI) -AGto (BH') (8) AGt' (HA-A-),,, = AGt' (A-) + AGLO (H') -AGtO (HA) (9) (A-BjSys, i.e. either Acto (BH+-B),,, (for B = pNA, Tris) or Acto (HA-A-)sys (for HA = pNP), for various ACN + EG mixtures are listed in Table 6, together with those of AGf,O (R),AGto (A-), AGto (BH+), and AGp (HA).The probable uncertainty in the values of AGto (A-B)s4.s,AGto (B), and AGt," (HA) is &O.lG kJ mol-l. TABLE6 Standard free energies of transfer (kJmol-l) on the mole fraction scale in ACN-EG mixtures at 298.15 K B = Tris Wt. % r h ACN AGto (BH+-B),,,, AGt *(B) AG~(BH~) 20 0.50 0.96 0.8 40 1.29 2.60 2.0 GO 2.87 4.68 3.5 B = PNA Wt. 7; B = pNA ACN LGto (BH+-B),,, AGto (B) (BH+)AGO 20 0.02 -1.76 -1.5 40 0.69 -2.95 -2.9 60 1.14 -3.60 --3.0 HA = PNP hwt. yo r ACN AGto (HA-A-),,, AGto (HA) AGp (A? 20 1.61 -0.84 0.4 40 4.74 -1.37 2.7 60 7.81 -1.35 4.8 DISCUSSION In Figure 1, AG,o (A-B)sys for pNA, Tris, and pNP is plotted against the mole percent of ACN in ACN-EG mixtures.The corresponding plots for pNA and Tris in the H20-EG system are also shown, being drawn from the results obtained earlier 7a In the absence of AGto (A-B)sys values for pNP in H20 + EG mixtures, AGto (A-B)sys values for another uncharged acid : acetic acid (HAc) 166 have also been plotted. The variation of AGto (A-B),,, with solvent composi- tion for any acid-base system is observed to be sharply contrasting in the two solvent systems one of which is a non-isodielectric protic + protic combination while the other is an isodielectric protic + dipolar aprotic one. For the dissociation of HAc which involves charge separation, a large portion of the increasing ease of ionization in H20 + EG mixtures must come from the increasing dielectric constant of the medium, whereas the isoelectric dissociations of TrisH and pNAH k, being little affected by dielectric constant changes, become readier owing to the increased solvation of the H+ ion by increasing amounts of added water; l7 the flat portion in the curve for Tris and the upcurving portion in the curve for pNA may be attributed to the increase in Acto (H+) at water rich compositions.17 On the other hand, in the isodielectric ACN-EG system, we find a decreasing ease of ionization for all three acids, pNAH+, TrisH+, and pNP, as the proportion of ACN in the solvent increases.Because of their closely similar structures, the dispersion interactions with the dipolar aprotic solvent ACN are likely to be similar for the two components of each of the conjugate acid-base pairs pNA-pNAH' and pNP-pNP-, pNP-being the 9-nitrophenoxide ion. Consequently, the increasing resist- ance to ionization must be explained by the increasing desolvation of the hydrogen ion15d and, for pNP, also by the increasing desolvation of the pNP- anion in ACN-EG mixtures.Similarly, the hydrogen-bonding 1°.OtI I I I 0 30 60 90 100 Mol '10ACN or H,O FIGURE 1 Variation of AGto (A-B)w (A = pNP, pNA, Tris, or HAc) at 298.15 K with solvent composition in ACN-EG 0and H,O-EG A mixtures tendencies of Tris and TrisH+ are expected to be almost identical in a given solvent and here too the increasing difficulty of ionization must arise from the increasing desolvation of the hydrogen ion.To confirm these contentions, it will be necessary to examine the behaviour of the individual species concerned. In Figure 2, the fIee energies of transfer for the neutral species Tris, pNA, and pNP in ACN-EG mixtures are shown plotted against the mole percent of ACN. For comparison, the corresponding plots for the H,O-EG 7~ system and the isodielectric methanol (MeOH) + propy-lene glycol (PG) Tf) system are also shown (excepting those for pNP for which no data are available). In interpreting the plots shown, it must be noted that while Tris, by virtue of its three OH and one NH, groups, interacts with the solvent mainly through *i 1 1;;T 5.0 30 600 Mol %ACN,MeOH or H20 loo 2FIGURE Variation of AGtO (B)(B = pNA or Tris) at 298.15 in ACN-EG 0,H,O-EG A,and MeOH-PG 0,as well as of AG~,, (p~p)at 298.15 K in ACN-EG mixtures with solvent composition hydrogen-bonding, pNA and pNP, by virtue of their aromatic nuclei, exhibit primarily dispersion inter-actions with the solvent.Thus, increasing values of AGto (Tris), e.g. for the ACN-EG and MeOH-PG systems, show that the hydrogen-bonding prowess of the mixed solvent in each case progressively decreases as the proportion of ACN or MeOH increases. The relative diminution of this prowess in ACN-EG mixtures far exceeds that in MeOH-PG mixtures, probably because MeOH possesses a single centre of hydrogen- bonding, EG and PG have two each, and ACN fares poorly either as a hydrogen-bond donor or ac~eptor.~.l**l9 On the other hand, H,O is a better hydrogen-bonding solvent than EG, if only marginally so, as reflected in the slowly decreasing values of AGto (Tris) with the increasing addition of H20to EG. The behaviour of AGto (pNA) in these solvent systems is seen to be strikingly different. The strength of dispersion interactions undergone with a solute increases in the sequence H20 < EG < ACN as shown by the solvation behaviour of the Ph,As+ or Ph,lB-This is confirmed by the sharp increase and decrease of AGto (pNA) in the H20-EG and the ACN-EG systems respecti- vely. MeOH and PG are probably equally effective (or ineffective) in undergoing dispersion interactions and virtually no change in AGto (pNA) with solvent composi- tion is noticed in MeOH-PG mixtures.A further point worth noting is that although pNA and pNP have identical centres of dispersion interactions (namely, quinonoid nuclei), the relative increase of solvation of pNA exceeds that of pNP in ACN-EG mixtures, thus implying that P~~ is less solvated in pure EG than pNP, even if both solutes are identically J.C.S. Perkin I1 solvated through dispersion interactions in a given solvent mixture. Figure 3 shows a comparison of AGt: (BH+) (B = Tris, pNA) us. solvent composition plots in ACN-EG and H,O-EG mixtures. For the latter system, equation (8) was used to calculate AGto (BH!) values from liter-ature values of AGto (BH+-B),,s,7a AGho (B),7aand AGto (H+).17 Here, contrasting behaviour is exhibited by AGto (pNAH+) in the two systems, whereas the AGto (TrisH+) curves are very similar.What is significant is that in spite of the increasing dielectric constant, both these cationic species are desolvated on passing from EG to H,O-EG. For pNAH+, this trend may be attri- buted to the marked loss of dispersion interaction stability not adequately compensated for by increasing electrostatic or hydrogen-bonding effects, but for TrisH + this trend is apparently inexplicable unless we assume some sort of specific solvation with EG. In ACN-EG mixtures the decreasing AGto (pNAH+) indicates once more that dispersion interactions for this cation are overwhelmingly stronger than hydrogen-bonding so that its solvation by ACN rather than by EG is favoured.The diametrically apposite nature of the AGto (TrisH+) curve shows that the reverse is true for TrisH+ and a loss of hydrogen-bonded stability in passing from EG to ACN-EG is indicated. Figure 4 presents a collection of the various free energies of transfer arrived at in this study by direct experiment or by indirect derivation. In the ultimate analysis, it now becomes clear that the features present in the AGtO (A-B)sys curves arise from the super-imposition of the features of the individual curves for AGt, (A), AGto (B), and AGto (H+). As conjectured earlier, the relative changes of solvation of the cationic acids BH+ from EG to ACN-EG mixtures are not much -< +2 30 60 90 100 MOl O/o ACN Ot H20 FIGURE3 Variation of AGto (BH+) at 298.15 11; (B = pN4 or Tris) with solvent composition in ACN-EG 0,and H20-LG A different from those of their conjugate bases B, so that AGto (BH+-B),, is almost wholly dictated by AG,o (H+).Now since the values of AGto (H') were obtaineclI5d using the tetraphenylarsonium tetraphenylboride assumption,20u such close correlation betu.een AGto (BH+-B),,, and AGto (H') is strongly indicative of the 1213 positions of the ACN-EG system before being increas- ingly desolvated at higher ACN proportions.15" This difference between these two ions of a similar nature arises most probably from the more extensive charge clelocalization achieved by the Pi-ion compared to the pNP-- ion, which enables the former to mask successfully effectiveness of this assumption in estimating ionic its anionic character, at least up to moderate concen-AGt: values, as suggested by Popovych 20b who found a trations of ACN, in favour of its tendency to undergo increased dispersion interactions with the solvent.The authors express their thanks to the C.S.I.R., New Delhi, for providing a senior research fellowship to one of Tris them (I<.€3.) and partly also the N.C.E.R.T., New Delhi, for ;financial assistance. [S/1605 Received, 7th September, 19781 REFERENCES R. P. Bell, ' The Proton in Chemistry,' Cornell University Press, Ithaca, New York, 1959.E. J. King, 'Acid-Base Equilibria,' Pergamon, London, 1965. E. J. King, in ' Physical Chemistry of Organic Solvent Systems,' ed. A. K. Covington arid T. Dickinson, Plenum Press, London, 1973, p. 331. K. G. Bates, ' Determination of pH, Theory and Practice,' Wiley-Interscience, New York, 2nd edn., 1973. R. G. Bates, in ' Hydrogen-Bonded Solvent Systems,' ed. A. K. Covington and P. Jones, Taylor and Francis, London, 1968, p. 4.9. (a) C. H. Rochester and H. Rossall, Trans. Faraduy SOC., 1969, 65, 993, 1004; (b) G. €1. Parsons, C. H. Rochester, and ('. E. C. Wood, J. Chem. SOC.(B),1971, 533; (c) G. H. Parsons and C. H. Rochester, J.C.S. Faraday I, 1975, 1058, 1069. K. I<. Kundu, A. L. De, and M. N. Das, (a)J.C.S. Pevkin 11, 1972, 2063; (b) J.C.S.Dalton, 1972, 356. (a) E. E. Sager, R. A. Robinson, and R. G. Bates, J. Res. Nut. Bur. Stand., Sect. A, 1964, 68, 305; (b) M. Woodhead, &I. Paabo, R. A. Robinson, and R. G. Bates, zbid., 1965, 69, 263; (c) P. Schlindler, R. A. Robinson, and R. G. Bates, ibid., 1968, 72, 141; (d) M. Paabo, R. G. Bates, and R. A. Robinson, J. Phys. Chem., 1966, YO, 247. C. L. DeLigny, Rec. Trav. claim., 1960, 79, 731. lo K. G. Bates and G. Schwarzenbach, Helv. Chim. Actu, 1955, 38, 699. l1 M. Merle, G. Douheret, and M. L. Dondon, Bull. SOC.china. Fvance, 1966, 159. l2 A. K. Boniand H. A. Strobel, J. Phys. Chem., 1966'70,3771. l3 E. A. Braude and E. S. Stern, J. Chem. SOC.,1948, 1976. l4 (a) R. G. Bates, J, S. Falcone, jun., and A.Y. W. Ho,Analyt. Chein., 1974, 46, 2004; (b) Y. Bokra and R. G. Bates, ibid., 1975, 47, 1110; (c) R. G. Bates and H. R. Hetzer, J. Phys. Chem., 1961, 65, 667. l5 K. Bose and I<. K. Kundu, (a)J.C.S. Faraday I, 1977, 73, 284; (b)J. Solution Chem., 1979, 8, 175; (c) Indian J. Chew., in the press; (d) Canad. J. Chem., in the press. l6 (a)K. K. Kundu, P. K. Chattopadhyay, and M. N. Das, J. Chem. SOC.(A),1970, 2034; (b) S. K. Banerjee, Ph.D. Thesis, Jadavpur University, Calcutta-700032, India. (a)A. K. Das, Ph.D. Thesis, Jadavpur University, Calcutta- 700032, India: (b)A. K. Das and K. K. Kundu, Indian J. Chem., in the press. 1' I I 25 50 Mo~ ACN FIGURE Variation of AGp (A-BBYa),AGt0 (A), AGp (B), and4 AG,o(H+) at 298.15 K with mol yo ACN in ACN-EG mix-turcs similar correlation for several primary ammonium acids in aqueous mixtures of methanol and ethanol. On the other hand, we find widely different curves for AGt() (pNP) and AGto (pNP-). While increasing dispersion interactions with ACN increasingly stabilize the neutral pNP molecule in ACN-EG mixtures, the pNP- ion, in spite of its aromatic nucleus, is strongly desolvated as the ACN content of the solvent increases. This emphasizes the strong anion-desolvating character- istics of a dipolar aprotic solvent like ACN 3918919921 which outweighs the effect of dispersion interactions in this case. It is pertinent to note that the picrate ion (Pi-) was found to undereo increased solvation at intermediate com-la R. G. Bates, in ' Solute-Solvent Interactions,' ed. J. F. Coetzee and C. D. Ritchie, Marcel Dekker, New York, 1969, p. 45. 19 E. Price, in ' The Chemistry of Nonaqueous Solvents,' ed. J. J. Lagowski, Academic Press, London, 1966, p. 67. 20 0. Popovych, (a) Critical Rev. Anulyt. Chem., 1970, 1,73; (b)Analyt. Chem., 1974, 46, 2009. 21 A. 1. Parker, Chem. Rev., 1969, 69, 1.
ISSN:1472-779X
DOI:10.1039/P29790001208
出版商:RSC
年代:1979
数据来源: RSC
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