460
The Journal of Physical Chemistry, Vol. 83, No. 4, 1979
I. M. Kolthoff and M. K. Chantooni
infinitely with an increase of temperature. This suggests that phase separation into TBA-rich and water-rich phases occurs a t a higher temperature. Thus, a pseudo-lower critical solution temperature is expected to exist above the boiling temperature in TBA-water mixtures. In dioxane-water mixtures, the number ratio Of dioxane molecules in the local structure in solution is quite different
~
~~~~~~~~~~~~~~~~~~~~~
(8)
(9) (10) (11) (12)
~ f~~~~~ ~
The stability of the local structure in dioxane-water mixtures in the temperature range 18-49.5 "C also supports this conclusion.
References and Notes (1) K. Iwasaki and T. Fujiyama, J . Phys. Chem., 81, 1908 (1977). (2) D.N. Glew, H. D.Mak, and N. S. Rath, "Hydrogen-Bonded Solvent System", A. K. Covlngton and P. Jones, Ed., Taylor and Francis, London, 1968, p 195. (3) R. K. McMullan and G. A. Jeffrey, J. Chem. Phys., 42, 2725 (1965). (4) T. C. W. Mak and R. K. McMullen, J. Chem. Phys., 42, 2732 (1965). (5) K. Iwasakl, M. Tanaka, and T. Fullyam, Bunko Kenkyu(InJapanese), 25, 134 (1976). (6)K. Iwasaki, M. Tanaka, and T. Fujiyama, Bull. Chem. SOC.Jpn., 49, 2719 (1976). (7) The presentiy reported (R& values at 24 OC for TBA-water mlxtures are slightly larger than those reported in ref 1 over the entire
concentration range. Readers are invited to have more confidence in the present (R& values and, accordingly, in the present N((Ax)') values at 24 O C . J. Kenttamaa, E. Tommila, and M. Martti, Ann. Acad. Sci. Fenn., Ser. A2, 93 (1959). R. Arnaud, L. Avediiian, and J.P. Morel, J. Chim. Phys., 69, 45 (1972). V. S.Griffiths, J. Chem. SOC.,1326 (1952). R. D. Stallard and E. S. Amis, J. Am. Chem. SOC.,74, 1781 (1952). F. Hovorka, R. A. Schaefer, and D.Dreisbach, J . Am. Chem. soc.,
~
~
~
f
(13) The 56, 2264 similar(1936). temperature dependence of specific concentrations has been reported for the observed excess partial molal volume and for the observed excess partial molal heat capacity: C. De Vlsser, G. Perron, and J. E. Desnoyers, Can. J . Chem., 55, 856 (1977). (14) K. Iwasaki, Y. Katayanagi, and T. Fujiyama, Bull. Chem. SOC.Jpn., 49, 2988 (1976). (15) T. Kato and T. Fujlyama, J. Phys. Chem., 80, 2771 (1976). (16) T. Kato and T. Fujlyama, J. Phys. Chem., 81, 1560 (1977). (17) T. Kato and T. Fujiyama, Bull. Chem. SOC.Jpn., 51, 1328 (1978). (18) T. Fujiyama and M. Tanaka, Bull. Chem. SOC.Jpn., 51, 1655 (1978). (19) T. Kato, S. Hycdo, and T. Fujlyama, J. Phys. Chem., 62, 1010 (1978). (20) KO is defined as KD(nA,B/N")= (n,,/N")(n,/N*) where N" = nA,B -t nA, -k n ~ ~. A , B nAl, , and nBare the numbers of the local structures A,B and A, and the molecules B, respectively, in the mlxture of components A and B. (21) The dlssociatlon equlllbrium of the type, g(A,B) ABxl B, Is consldered here. (22) F. Franks, Ann. N . Y . Acad. Scl., 125, 277 (1965). (23) K. W. Morcom and R. W. Smith, Trans. Faraday Soc., 66, 1073 (1970).
*
+
Acid-Base Equilibria of Hydrochloric and Hydrobrornlc Acids in Isopropyl and fert-Butyl Alcohols I. M. Kolthoff" and M. K. Chantoonl, Jr. School of Chemistry, University of Minnesota, Minneapolis, Minnesota 55455 (Recelved June 8, 1978; Revised Manuscript Received October 17, 1978) Publication costs asslsted by the National Science Foundation
Acid-base equilibria of hydrochloric and hydrobromic acids have been studied at 25 "C in isopropyl and tert-butyl alcohols and to a lesser extent in n-hexyl alcohol. Values of pKd(HC1)are 3.1, 5.5, and 4.2, and of pKd(HBr) 2.0,5.0, and 4.1, respectively. Homoconjunction and protonation constants in t-BuOH of both acids, Kf(HX;), Kf((HX),X-), and Kf(H2X+)are 7 , 8 X lo1, and 8 for HC1 and 2.0 X lo1, 1.7 X lo2, and 1.6 X lo1 for HBr, Conductance and paH data provided conclusive evidence that with increasing concentration the dissociation 3HX F= H2X+ HX; becomes more predominant. The behavior of HC1 as a very weak base has been confirmed in acetonitrile as a solvent. From the above constants and from solubility products of several salts, transfer activity coefficients of various species with reference to methanol or acetonitrile as a solvent are reported. It is concluded that the above three alcohols are equally strong hydrogen bond acceptors as methanol and ethanol, but that isopropyl and tert-butyl alcohol are considerably weaker hydrogen bond donors than the normal alcohols. In this connection the methanolation (solvation) constant of the chloride ion in tert-butyl alcohol has been reported.
+
Introduction In the last 25 years there has been considerable interest in the dissociation of hydrogen halides in a homologous series of alcohols, particularly in primary alcohols. A review has been presented in 1960 by Janz and Dany1uk.l More recent contributions are referred to in the present paper. During the last few years we have made extensive studies of acid-base equilibria in isopropyl (i-PrOH) and tert-butyl (t-BuOH) alcohols with dielectric constants, D, a t 25 "C of 19.9 and 12.5, respectively. The present paper deals with a study of equilibria in solutions of hydrochloric and hydrobromic acids in these two solvents and, to a lesser extent, in n-hexyl alcohol (n-Hex) with D = 13.3, which is practically isodielectric 0022-3654/79/2083-0468$01.00/0
with t-BuOH. At the outset it may be pointed out that all experiments were carried out under conditions where no appreciable formation of alkyl halides occurred. From potentiometric paH and conductometric data obtained in solutions of the hydrogen halides and of mixtures of the hydrogen halides and the corresponding tetraethylammonium halides various constants have been calculated. Solubility products of some salts have been determined in order to obtain transfer activity coefficienta of the halide ions and of the undissociated hydrogen halides with reference to methanol (MeOH) and to the aprotic protophobic solvent acetonitrile (AN). These transfer coefficients are closely related to the extent of solvation of chloride and bromide ions in the higher alcohols as 0 1979 American Chemical Soclety
f
Acid-Base Equilibria
in i-PrOH and t-BuOH
The Journal of Physical Chemistry, Vol. 83, No. 4, 1979 469
compared to those in MeOH or AN. An indirect method was used to obtain transfer activity coefficients, Syt-BuoH, of chloride and bromide (and 3,5-dinitrobenzoate) ions. The solubility of the potassium salts of these ions in tBuOH is too small to yield a reliable value of the conductance of the saturated solutions and of the total concentration. The solubility products, KSP, could be obtained from saturated solutions of the salts in which K+ was complexed with the crown ether dibenzo-18-crown-6 (L).2 The complexed ion is denoted by LK+. From the conductance of a solution saturated with salt and crown ether the product a(LK+)a(X-) was found in t-BuOH. Methanolation constants of the chloride ion in t-BuOH have also been determined conductometrically from the effect of MeOH om the ionic solubility of tetramethylammonium chloride. These constants are a measure of the strength of hydrogen bonding of the anion with the solvent as compared to that by MeOH added as a solute. Equilibria of electrolytes in solvents of low dielectric constant are quite involved. The following equilibria in solutions oC a uni-univalent electrolyte MX must be considered: MX ,eM+ XKd(MX) = a(M+)a(X-)/[MX] (1)
+
MX + M+ + M2X+ @(M2X+) = a(M2X+)/a(M+)[MX]
(2)
MX + X- F? MX2Kf(MX2-)= a(MX,-) /a(X-) [MX]
(3)
3MX + M2X+ + MX2Kd(3MX) = u(MX~-)U(M,X+)/[MX]~ In this paper MX is regarded as being monomeric. In solutions of MX, in which only 1:l complex formation occurs, the electroneutrality condition is [M+] + [MZX'] = [MXZ-] + [X-] (4) Formation of higher complexes, i.e., (MX),M+ and (MX;12X-,is also considered. Using the symbol y for the Debye-Huckel activity coefficient (IUPAC notation) we find in a solution of MX
Parentheses refer to Kd,brackets refer to concentration. The total ion ccncentration, C[M+] = [M'] + [M2X+],to a first approximation, is taken proportional to the conductance yC[M+] = (Kd(MX)[MX]{l+ Kf(MX,-)[MX]} X (1+ KYMZX+)[MX]})1/2(6) When either M2X+ or MX2- is present, the equivalent conductance, A, would essentially be independent of concentration of MX.3 The slope of the paM vs. C(MX) plot is unity when formation of MX2- is extensive3 (Kf(MX,-)[MX:] >> 1and Kf(M2Xt)[MX]0.003 M in t-BuOH and n-Hex, >0.02 M in i-PrOH) which indicate bilateral triple ion formation. Assuming the formation constaints of M2X+ and MX2- equal and Ao(M-MX2)= ho(M2X.X),iplotb of h y ~ ~ ' ~ / F vs. ( z )c9 were constructed (not presented) and found to be linear. From the ratio of the slope to intercept the following values of Kf(M2X+) were found tetramethylammonium chloride 3.6 in i-PrOH, 1.6 X lo1 in t-BuOH, and 9 in n-Hex; while the values in t-BuOll of tetrabutylammonium bromide and iodide are 6.6 X lo1 and 4., X lo', respectively. For these salts the intercept yields Kd(MX),the values of which are in Table 11. In a previous paper5 ion mobilities of tetrabutylammonium, tetraethylammonium, proton, and picrate have been reported in i-PrOH and t-BuOH, and of the first two in n-Hex. In i-PrOH Evans et al.1° report Xo equal to 13.10,10.55, and 11.45 for tetramethylammonium, chloride, and bromide ions, respectively. As ho values of tetraalkylarnmonium halides in t-BuOH and n-Hex could not be determined with reasonable accuracy from the Fuoss and K,raus plots, owing to triplet ion formation, ionic mobilities in these alcohols were calculated from the Walden product in MeOH, yielding the following values: Me4N+7.9, C1- 6.0, lBr- 6.5, and I- 7.2. It is reasonable to regard Xo(LK+)= Xo(Bu4N+),as the ionic radii are similar (5 A). Also, Kd(Bu4NPi)= Kd(LKPi) in t-BuOH. In i-PrOH Xo(K+) was obtained from the Walden product in MeOH and found equal to 13.2. Viscosities at 25 "C of MeOH, i-PrOH, t-BuOH, and n-Hex are 0.545, 2.018, 4.77, and 4.59 cP, respectively. Acid -Base Equilibria in Solutions of H X . Conductance Section. Conductance data of freshly prepared solutions of HCI and HBr in i-PrOH, t-BuOH, and n-Hex are presented as plots of A vs. c1I2(phoreograms) in Figure 2. In n-propyl, isopropyl, and n-butyl alcohols steep phoreogrms of HC1, HBr, and HI have been reported,ll with leveling off at higher concentrations (C > 0.06 M). We find a similar phoreograni of HBr in i-PrOH (Figure 2 ) . The value ofKd(HBr)in i-PrOH (Table 111) has been estimated from the linear portion of the Fuoss and Kraus plot (C < 0.01 M). In the present study a well-defined minimum at -0.04 M is found in the phoreograms of HC1 and HBr in t-BuOH and n-Hex in Figure 2. Values of pKd(HX), Kf(HXJ, KfC(HX),X-), and Kf(H2X+)are obtained for HCl and HBr in t-BuOH by combining conductivity and paH data (vide infra). In n-Hex, values of these constants for HCL and HBr in Table I11 have been approximated from plots of dye'~'/F(z) vs. c(HX) (C < 0.02 M), neglecting formation of (HX)2X-and assuming Kf(HX,-) =
0
04
02
06
010
0
020
E
E
Figure 2. Plots of A vs. cl'* of hydrogen halides in alcohols. (A) 1, HBr in n-Hex; 2, HCI in n-Hex; 3, HBr in t-BuOH; 4, HCI in t-BuOH. (B) HBr in i-PrOH.
I
05
,
,
,
,
I
15
IO
-log
,
,
,
,
J
20
,
,
,
,
I
,
25
CHX
Figure 3. Plots of paH vs. log C(HX) of hydrogen halides in t-BuOH: (A) HCI; (8) HBr.
Kf(H2X+).Because of an uncertainty of about 3% in the conductivity of the hydrogen halides extrapolated to zero time and uncertainty in the values of paH, values of pK(HX) in Table I11 are good to within 0.1 unit and those of Kf(HX2-)and Kf(HX),X- in t-BuOH are within 15%. Bezrnanl, presented phoreograms of HC1 in isoamyl alcohol (D = 14.7) exhibiting a minimum, which upon addition of benzene in increasing amounts becomes more pronounced and moves to lower concentration as the dielectric constant of the mixtures decreases. Potentiometric Section. Potentiometric paH measurements with the glass electrode were made within 0.5 h after preparation of 3 X to 0.04 M HC1 and IlBr solutions in t-BuOH in absence of salt. Data are in Table a, column 7.13 Plots of paH vs. -log C(HX) in Figure 3 are almost linear for both acids throughout the concentration range, the slopes being -O& and -0.71, for HC1 and HBr, respectively. Conductivity and paH data taken together suggest formation of the cationic complex H2X+ and anionic complexes HX2- and (HX)&. In solutions of HC1 and HBr more dilute than 0.03 M the agreement between C[H+] = [H+] + [H2X+]from conductance (eq 6) and [H+] from paH data (eq 5 ) is as good as can be expected. Concentrations of H2X+found from eq 4 are listed in the third from the last column in Table a. Values of [HX], equal to C(HX) - C[H+] are in the next to the last column in Table a, while values of Kf(H2X+)are in the last column. Values of the latter are essentially constant from about 0.04 to 0.4 M HX in t-BuOH. No higher cationic complex is formed. The evaluation of anionic complexation constants involves the product of C[H+] determined conductome-
472
The Journal of Physical Chemistry, Vol. 83,
I. M. Kolthoff and M. K. Chantooni
No. 4, 1979
TABLE I11 : Dissociation and Homoconjugation Constants of Hydrogen Halides in Alcohols and Acetonitrile solvent
pKd( HX)
pKd( 2HX)a
K~(Hx,-)
K~(H,x+)c
Kf((HX),X-)c
HCl MeOH EtOH i-PrOH l-BuOHC n-HexC
AN
l.lb
2.0b 1.9e 3.1e 5.5 (4.2) 10.4
-0 7 (2.8 x 101) 4.3 x 103c 2 x 104f
4.7 (2.8) 6.F 6.V
ll.OB
7.,
x 10'
8
( 2 * 8x 101)
5 * 4 x 10'
HBr MeOH EtOH i-PrOH t-BuOHC n-HexC AN
( 0.8)d
1.7d 2.0c 5.0 (4.1) 5.F
3.7 (2.6) 2.6c 3.1g
?.Of
2.0 x 10' ( 3 x 101) 1.4x 103c 8.2 x 103f
1 . 7 X 10'
1.6 X 10' (3 x 10')
pKd( 2HX) = pKd(HX) - log Kf(HX,-). Average value from several authors: A. Ogston, Trans. Faraday SOC.,32, 1679 (1936); T. Shedlousky and R. Kay, J. Phys. Chem., 6 0 , 1 5 1 (1956); I. Bezman and F. Verhoek, J. A m . Chem. SOC., 67,1330 (1945); A. El Aggan, D. Bradley, and W. Wardlaw, J. Chem. SOC.,2092 (1958). This work. G. Charlot and B. Tremillon, "Chemical Reactions in Solvents and Melts", Pergamon Press, New York, 1969, p 278. e Reference 16. Value of pKd(HPi) = 11.0 used to calculate Kd(HX) and Kf(HX,-) of HCl and HBr in acetonitrile from reaction of HPi with Et,NX (ref 3). g pKd( 2HX) estimated from reaction of HX with nitroaniline indicators (see footnote f ) . a
PO H
13 12
IO
8
9
7
8
6
7
5
-I 0
-0 5
0
05
Pa H
10
log c o / c s
Flgure 5. Plots of paH vs. log [C(HX)/C(Et,NX)] of HX in AN. Curve M, C(Et,NCi) = 2.6 X 10-3-2.6X A: HCI (A)C(HCI) = 2.8 X lo-' M; (A)C(HCI) = 4.4 X 10-3-l.lX IO-' M, C(Et,NCi) = 2.3 X M. Curve B: HBr (0)C(HBr) = 3.5 X M, C(Et,NBr) = 2.6 X 103-2.4 X lo-' M; (0)C(HBr) = 9.5 X 103-4.0 X lo-*M, C(Et,NBr) M. Curve A, left-hand ordinate, curve B, right-hand = 3.2 X ordinate. Figure 4. Plots of y z [ H + ] a ( H + ) / [ H X ] vs. C(HX) for HBr and HCi in t-BUOH.
trically and a(HS) determined potentiometrically. Combining eq 5 and 6, and taking formation of (HX)zXinto account, yields
+
yC[H+]a(Hf) = Kd(HX)[HX](l Kf(HXz-)X [HX] + Kf((HX)&) [HX12}( 7 )
+
The term (1 Kf(H,X+)[HX]] cancels. Plots of YE[H*]a(H+)/[HX]vs. C(HX) according to eq 7 are presented in Figure 4 for HC1 and HBr in t-BuOH taking C(HX) = [HX]. The intercept is equal to Kd(HX),while Kf(HX,) and @((HX),X-) are found by curve fitting. The stability of (FIX)& accounts to a large extent for the sharp minimum in the phoreograms in Figure 2. For a comparison of paH vs. log (Ca/Cs) plots in i-PrOH and t-BuOH with similar curves in a protophobic dipolar aprotic solvent, such plots of freshly prepared solutions of HC1 and HBr in acetonitrile are presented in Figure 5. Calculation of Kd(HX) and Kf(HXz-)from paH data in AN when the salt MX is incompletely dissociated has been I4 described previously.8 Dissociation constants 2.9 X and 0.2115for tetraethylammonium chloride and bromide,
respectively, in AN were used. Table I11 summarizes values of pKd(HX), Kf(HXz-),and Kf((HX),X-)) of HC1 and HBr in the alcohols and AN. In addition, de Lisi and GoffredP report pKd(HC1)as 2.5, 3.08,and 3.25 in n-PrOH, n-BuOH, and i-BuOH, respectively. In AN the agreement in pKd(2HX) from potentiometric and previous spectrophotometric nitroaniline indicator data3 is good, but values of Kf(HX,-) differ considerably. Potentiometric values of Kf(HX,-) in the present study are more reliable than those reported previously3 from the reaction of picric acid with tetraalkylammonium halides in poorly buffered solutions. In DMF values of pK(HX) are 3.417 and 1.818for HC1 and HBr, respectively. Corresponding values in MezSO are 2.019 and 1.0.20 Protonation of Hydrogen Chloride in Acetonitrile. Qualitatively, the basic character of HC1 in AN was easily demonstrated by observing the effect of HC1 on the color (log of two Hammett indicators, 2-nitro-4,5-dichloroaniline Kf(IH+) = 3.lZ1)and 2-nitro-4-chloroaniline (log Kf(IH+) = 4.2,I) in freshly prepared solutions which were 5 X loe3 to 2 X lo-, M in trifluoromethanesulfonic acid and 0.1-0.6 M in HC1. A definite increase in intensity of the yellow color of the alkaline form of the indicators was observed
The Journal of Physical Chemistry, Vol. 83, No. 4, 1979 473
Acid-Base Equilibria in 1-PrOH and t-BuOH TABLE IV: Effect of Methanol on Solubility of Me,NCl in t-BuOH C(MeOH), M
0
0.243
0.515
1.0
specific conductance total solubility
2.80 X l o w 6 1.81 X 12.5 2.24 x 10-4 0.60
5.05 X lo-' 1.95 X lo-' 12.0 4.20 x 10-4 0.48 37.3 1.3 x 10-3 1.2,
8.07 X 2.26 X lo-' 11.3 7.13 x 10-4 0.40 27.2 4.1 x 10-3 0.8,
2.01 x 10-5 3.82 X 9.5 2.12 x 10-3 0.23 4l.1 1.8, X 1.0,
-.
A
[MeJ+It Y2
( V - l)/C*(MeOH.Cl-) [ Me4NC1-2MeOH] ITf(Me4NC1.2MeOH)
-
TABLE: V: Transfer Activity Coefficients between Alcohols
S log MrS(BPh,-) log M r S ( TAB') log My (K+) log My (H') log MrS(Pi-1 log MrS( c1-) log log log log
My (Br-)
MrS(HPi) MrS(HCl) MrS(HBr)
EtOH
i-PrOH
t-BuOH
n-Hex
An
DMFb
Me,SOb -2.4
_ .
t 0.5c 1.oc
-0.1c 0.9c 1.3c
1.Od 0.6c 0.3f 0.of
1.5,'
2.5,a
1.7'
1.7b
-2.6
-0.8" 3.1" 5.4e 5.68 5.08 1.1" 0,1e,f 0.OfJ
0.4" 2.3' 3.8e
-0.4b 6.2b 0.5b 5.2b
-3.5 -4.4 (-0.7) 5.9
-3.9 -5.2
3.6b
4.1
2.5
-0.8f -1.3f
-1.5f -2.9f
2.2e
-1.4' 1-P 2.8e 2.6e 0.4 ' -0.d O.lf
0.9" (0.9)eIf
4.6
-0.5' 2.lf 4.88
a Reference 5. I. M. Kolthoff and M. K. Chantooni, Jr. J. Phys. Chem., 7 6 , 2024 (1972). 0. Popovych, A. Gibofsky, and D.H. Berne, Anal. Chem., 44,811 (1972). Derived from pKSP values of AgBPh, and A Br of R. Alexander, A. J. Parker, J. H. Sharp, and W. E. Waghorne, J. A m . Chem. SOC.,94,1148 (1972 , and using log rEtoH(BPh,-) = 0.5. e From pKsP(KPi)QKsP(KX) or#KSP(Me,NPi), pKsp(Me,NC1), and log M7S(Pi-). From the relation Sa'pKd(HX) = log M y S (H*) t log yS(X-) - log rS(HX), knowing the first three quantities in M and in S. 8 From ionic solubilities of KX or KDNB in t-BuOH, also saturated with dibenzo-18-crown-6and KSP of KX and KDNB in AN (see text).
2
in the presence of HC1, indicating that log Kf(H2C1+)lies between 1 and 2. Acid-base equilibria in solutions of CF3S03Hin AN have been found4to be unusually involved and the system becomes still more complex in the presence of HC1. The main objective of the above experiment was to confirm that HCW definitely has basic character. Methanoltrtion of' Chloride Ion in t-BuOH. The dimethanolation conslants of chloride ion and of the ion pair Me4N"C1- were estimated in t-BuOH from the ionic and total solubilities, [Me4N+Itand St, respectively, of tetramethylammonium chloride in presence of MeOH. Tetramethylammonium ion was regarded not to be methamolated, but triple ion formation was taken into account, the formation constants of (Me4N)2C1+and Me4NC12-assumed equal. Under these conditions, the total solubility is St = [Me4NC1]t [Me4NC1.2MeOH] t [Me4N+] t [(Me4N)2C1+],while the ionic solubility is [ Me4N1' = [Me4NS]t [ (Me4N)2C1+]= [C1-1 t [Me4NC12-] t [2MeOH.C1-]. Hence y2[Me4N+I2= KsP(Me4NC1)(1 Kf((Me4N)zC1+)[Me4NC1])(1t fl((Me4N)2Clf)[Me4NC1] + Kf(2MeOH.C1-) [MeOHI2). Specific conductivities of saturated solutions of tetramethylammonium chloride in presence of MeOH and corresponding ionic solubilities are presented in Table IV. From these data and values of pKap(Me4NCl) = 7.73, [Me4NG1],+,, = 1.8 X M (Table I), and Kf((Me4N)2C1t) = 1.6 X lo1, resulting values of Kf(2MeOHC1-) and Kf(Me4NC1*2MeOH)are 3.5 X lo1 and 1.0, respectively.
+
Discussion Protonation of HX. In the gas phase both HC1 and HBr can act as bases; their proton affinities are 135 and 140
kcal/mol, as determined by Polley and Munson.22 One of the very interesting results in this study is that both HC1 and HBr exhibit very weakly basic properties in t-BuOH, n-Hex, and AN which must be attributed to the weakly ionic character of both acids in these solvents. In fact, HCl in acetonitrile appears to be a stronger base than acetic acid (Kf(CH&OOH2+) = 1 X 101)4and is only slightly stronger in AN than in t-BuOH even though both the
proton and HC1 are solvated to a much lesser extent in AN, as is evident from the transfer activity coefficients in Table V. From these relations it may be concluded that H2C1+ is much stronger solvated by hydrogen bonding to t-BuOH or n-Hex than to AN. It is of interest to mention that Christie23isolated a white solid product (likely C1H2+SbFc) upon protonation at -45 "C of HC1 in the super acid solvent HF--SbC15. This product melted below room temperature with decomposition. Transfer Activity Coefficients of H X and HX;. In a subsequent paper we conclude that all the alcohols from MeOH to the hexanols are equally strong hydrogen bond acceptors with reference to carboxylic acids. Apparently, this conclusion also holds for HC1 and HBr as the values of log MyS(HX)in Table V are close to zero, S being alcohols. When interpreting transfer activity coefficients of HX from MeOH to basic dipolar aprotic solvents (IIS), such as dimethyl sulfoxide, formation of n relatively stable adduct HSaHX must be considered. As an example, the crystalline adduct HCLDMF was isolated by S10an.~~ The solvation of HX by hydrogen bonding increases with basicity of HS as is shown by values of log MyS(HX:)in Table V, On the other hand, HC1 is poorly solvated in AN (Table V) and is easily volatilized even from dilute, freshly prepared solutions. To a large extent, the fact that the homoconjugation constants of HC1 and HBr are smaller by at least 2 orders of magnitude in t-BuOH than in AN can be attributed in part to the stronger hydrogen bonding of HX to the alcohol than to AN, and also to larger solvation of the chloride ion by t-BuOH than by AN. Corrected for the Born effect log t-BuoHyAN(C1-) N- 3 (vide infra). Transfer Activity Coefficients of X-. In recent publication^^^-^' convincing evidence has been presented attesting to the validity of the S1ySa(PhP+)= SlrS2(BPh4-) assumption, originated by Grunwald.qs Because of the extremely Iow ionic solubility of Ph4AsBPh4in the higher alcohols we have used the closely related assumption of Popovych,2gS1rSz(TABt) = S1yS2(BPh4-)(TAB+ = triisoamylbutylammonium). In a previous paper6 values off log
474
I. M. Kolthoff and M. K. Chantooni
The Journal of Physical Chemistry, Vol. 83, No. 4, 1979
MyS(Pi-) were reported, which were derived from the solubility products of TABBPh4 and TABPi in MeOH, i-PrOH, t-BuOH, and n-Hex. Potassium chloride, bromide, and picrate are sufficiently soluble in i-PrOH to allow the estimation of log Myi-P"oH(X-) from the solubility products of these salts in MeOH and i-PrOH. As these potassium salts are extremely slightly soluble in t-BuOH and n-Hex, values of log Myt-BuoHiHex(X-) were estimated indirectly in two ways: (a) from solubility products in Table TI of tetramethylammonium chloride, bromide, and picrate in i-PrOH, t-BuOH, and n-Hex, and (b) from the ionic solubility of potassium chloride, bromide, and 3 5 dinitrobenzoate (DNB-) in t-BuOH solutions also saturated with dibenzo-18-crown-6(L). The solubility product, a(LK+)a(X-), which is equal to [LK+I2y2= KsP(KX)[L],,t,Kf(lAK+),was determined conductometrically. The formation constant of LK+ is denoted by Kf(LK+). In a subsequent paper the value of log ANyt-BuoH(DNB-), equal to 0.3, will be reported. Use of this value was made to r-) calculate values of log My"BUoH(C1-) and log Myt-BUoH(B from those of log MyAN(X-)(Table V), the values in tBuOH of a(LK+)a(X-) in Table 11, and the relation log AN t-BUOH(DNB-) - log AN yt-BuOH (X-) = log [Kap(KX)KfY (J-JK') [Lls~t.lt-BuOH+ PKap(KX)AN 1% [KSp(KDNB)Kf(LK")[L],,&.BuOH - pK"P(KDNB),. The quantities [L],, and Kf(LK+) in t-BuOH cancel in the above relation. Agreement in log Myt-BUoH(C1-) obtained from solubilities of tetramethylammonium and LK+ salts is good (Table V). The solid salts which are in equilibrium with the saturated solution remained uncomplexed. An attempt was made to calculate log Myt-BUoH(C1-) by using KBPh4 or KPi instead of KDNB in the above scheme involving LK+ salts, but the complexed LK+ salts appeared to be sparingly soluble. Corrected for the Born effect, values of log MyS(Cl-)in Table V become 0.6, 1.6, 2.1, and 0.8 and log MyS(Br-),0.4, 1.5, and 1.8, S being EtOH, i-PrOH, t-BuOH, and n-Hex, respectively. Ionic radii of 1.81 and 1.95 & . were used30for C1- and Rr- respectively. In order to convert free energies of transfer from the molar scale to those per mole of hydrogen bond donor a further correction involving the logarithm of the ratio of the molarities of the pure solvents has been applied. Resulting corrected values of log MyS(Cl-)are 0.4, 1.3, 1.7, and 0.3, and log MyS(Br-)= 0.2, 1.2, and 14 (expressed in log K units), S again being EtOH, i-PrOH, t-BuOH, and n-Hex. The neutral contribution to the transfer coefficients of halide ions may be approximated from the solubilities of the noble gas analogue^.^^ For the neutral contribution to log MyS(X-)of 61- and Br- the following solubilities of argon and krypton have been reported by Chiang,321.09 x 2.72 x M in MeOH, 1.1 x M (A) in AN, and 0.42 X M (A) in Me2S0. Hence, log "yAN(A) 0, and log MyMezSo(A)= 0.4. Assuming the solubility of Kr to be the same in AN as in acetone33 (4.37 X M), log MyAN(ICrj= -0.3. It appears that the neutral conis very small, as already tribution to log MyAN,Me2So(X-) concluded by Parker.33 If the same relation is true between M and the higher alcohols, the corrected values of log M-yS(X-)indicate much stronger hydrogen bonding of the primary alcohols to X- than that of the secondary and
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tertiary alcohols. Extensive steric crowding of the OH group in the latter two classes of alcohols iphibits such hydrogen bonding, in spite of the higher monomer content as compared to that in the primary alcohol^.^ The strength of hydrogen bonding of primary alcohols to X- is practically the same. The values of log MyS(HX)in Table V suggest that the free energy of solvation of HC1 and of HBr are about the same in the alcohols studied. Qualitatively, it appears that going from t-BuOH to n-Hex the hydrogen bond involving ROH and the proton in HX is weakened, while that involving ROH and X in HX is strengthened. Acknowledgment. We thank the National Science Foundation for Grant CHE-7522642 in support of this work. Supplementary Material Available: Table a containing conductivity and paH data for HC1 and HBr in t-BuOH (2 pages). Ordering information is available on any current masthead page.
References and Notes (1) G. Janz and S. Danyluk, Cbem. Rev., 60, 209 (1960). (2) C. J. Pedersen, J . Am. Cbem. Soc., 89, 7017 (1967). (3) I. M. Koklwff, S. Buckenstein, and M. K. Chantooni, Jr., J. Am. Cbem. Soc., 83, 3927 (1961). (4) I.M. Kolthoff and M. K. Chantooni, Jr., J . Am. Cbem. Soc., 95, 8539 (1973). (5) M. K. Chantooni, Jr., and I. M. Kolthoff, J . Pbys. Cbem., 82, 994 (1978). (6) C. Gracias, Ph.D. Thesis, University of Minnesota, 1961. (7) M. K. Chantooni, Jr., and I. M. Kolthsff, J. Am. Cbem. Soc., 89, 1582 (1967). (8) I. M. Kolthoff and M. K. Chantooni, Jr., J . Am. Cbem. Soc., 87, 4428 (1965). (9) R. M. Fuoss and C. A. Kraus, J. Am. Cbem. SOC.,55, 2387 (1933). (10) M. Matesich, J. Nadas, and D. Evans, J . Pbys. Cbem., 74, 4568 (1970). (11) H. Goldschmidt, Z.Pbys. Cbem., 124, 23 (1926); H. Goldschmidt and E. Matheison, ibid., A121 (1926); H. Goldschmidt and L. Thomas, ibid., 126, 124 (1927). (12) I. Bezman and H. Watson, J . Cbem. SOC.,2101 (1927). (13) See paragraph at end of text regarding supplementary material. (14) C. W. Davies, "Ion Association", Butterworths, Washington, D.C., 1962, p 96. (15) G. Forcier and J. Olver, flectrocbim. Acta, 14, 135 (1969). (16) R. de Lisi and M. Goffredi, flectrocbim. Acfa, 16, 2181 (1971). (17) B. Clare, D. Cooke, E. KO,Y. Mac, and A. J. Parker, J. Am. Cbem. Soc., 88, 1911 (1966). (18) P. Sears, R. Wolford, and L. Dawson, J . Electrocbem. Soc., 103, 633 (1956). (19) R. L. Benoit and C. Buisson, Electrochim. Acta, 18, 105 (1973). (20) J. Bolzan and A. Arvia, Nectrocbem. Acta, 16, 531 (1971). (21) I.M. Kolthoff and M. K. Chantooni, Jr., J . Am. Cbem. Soc., 95, 4768 (1973). (22) C. W.Polley and B. Munson, Int. J . Mass Spectrum Ion Pbys.,26, 49 (1978). (23) K. 0. Christie, Inorg. Cbem., 14, 2230 (1975). (24) W. J. Sloan, private communication. (25) A. J. Parker, Electrochim. Acta, 21, 671 (1976). (26) J. J. Kim, J. Pbys. Cbem., 82, 191 (1978). (27) S. Villermaux and J. J. Deipuech, J. Cbem. Soc., Cbem. Commun., 478 (1975). (28) E. Grunwald, G. Baughman, and C. Kohnstam, J . Am. Cbem. Soc., 82, 5801 (1960). (29) 0. Popovych and A. Dill, Anal. Cbem., 41, 456 (1969). (30) L. Pauling, "Nature of the Chemical Bond", Cornell university Press, Ithaca, N.Y., 1940, pp 346, 350. (31) M. A. Alfenaar and C. L. de Ligny, Red. Trav. Cbim., Pays-Bas, 86, 929 (1967). (32) F. Chiang, Ph.D. Thesis, University of Minnesota, 1967, p 203. (33) A. J. Parker in "Non Aqueous Electrochemistry", IUPAC, Paris, 1970, Butterworths, London.
