Effect of Ethanol 011 Hydrclphobic Interactions
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Effect of Ethanol on Hydrophobic Interactions. Conductometric Study of Ion-Pair Formation by Double-Long-chain Electrolytes
. Olakenfull" and D. E. Fenwick Division of Food Research, CSIRO, North Ryde, New South Wales 21 13, Australia (Received December 27, 1973) lJublicarion costs assisted by the CommonwealthScientific and hdustrial Research Organization
The equivalent conductance (A) of double-long-chain electrolytes such as decyltrimethylammonium decanoate was measured in a series of ethanol-water mixtures in which the mole fraction of ethanol was 0.05, 0.10, 0.15, and 0.20. With increasing electrolyte concentration, A was found to decrease by more than was predicted by the Onsager equation. This discrepancy was treated as resulting from ion-pair formation, and ion-pair association constants ( K ) were calculated. The electrostatic and hydrophobic contributions to the free energy of ion-pair formation (AGlon pair = -RT In K ) were separated by examining the effect of hydrocarbon chain length on AGion pa,p The change of the electrostatic contribution with ethanol concentration, iis a result of accompanying change in the dielectric constant, agreed with Bjerrum's equation. The hydrophobic contribution (AGHI) initially decreased (hydrophobic interactions became stronger) with increasing concentration of ethanol, with the minimum value observed when the mole fraction of ethanol was 0.10. A linear relationship was obtained when A G H was ~ plotted against the reciprocal of the isothermal compressibility (a measure of the free energy required to make a cavity in the solvent to accommodate a hydrophobic rdute). Contrary to popular belief, AGHI does not appear to be related to the surface tension of the solvent.
Introduction The importance of hydrophobic interactions in maintaining the v lability of living organisms is widely recognized.1-8 Studies of effects of solutes on hydrophobic interactions are important since they further our theoretical understanding o Ithese interactions and also provide a means of distinguishing hydrophobic from other forces (such as electrostatic and hydrogen bonding) involved in stabilizing the native conformation of proteins and nucleic acids. While studying the effects of hydrophobic interactions on the kinetic3 of reactions of some long-chain alkylamines with long-chain carboxylate esters of p-nitrophenol, we noticed that low concentrations of ethanol (mole fraction 50.1) apparently zacreased the effect of hydrophobic interaction betwee n the hydrocarbon chains, compared with the effect in pure ~ a t c r In . ~this paper we report a more detailed investigation of this phenomenon. This investigation is based on measurements of the equivalent conductance of solutions of hydrophobic ions. Some preliminary experiments10 on purely aqueous solutions of decyltriniethylammonium undecanoate, decanoate, nonanoate, etc. had suggested that hydrophobic interaction between the hydrocarbon chains might lead to the formation of ion pairs (at concentrations below the critical micelle concentration), but the effect was too small to obtain reliable ion. pair association constants. Addition of ethanol can be expected to favor formation of ion pairs because it lowers the dielectric constant of the so1vent.l' Our investigation shows that when the mole fraction of ethanol is 20.05, ion-pair formation is sufficiently increased to obtain reliable association constants and consequently to obtain an estimate of the fret. energy of hydrophobic interaction be tween the h ydimarbon chains. Experimental Section 1. Materials. Decyltrimethylammonium bromide (Eastman) was recrystallized from 50% (v/v) ethanol-ether. Car-
boxylic acids were redistilled before use. Water was distilled from glass and passed through a mixed-bed ion-exchange column [Bio-Rad AG 501-X8 (D)]. The conductiviohm-l cm-l. ty was always below 1 X Decyltrimethylammonium hydroxide was prepared by shaking a solution of the bromide (0.15 M ) with a twofold excess of freshly precipitated silver hydroxide, for 2 hr. The filtrate was titrated potentiometrically (Radiometer PHM26/ABU IC),and carboxylate solutions were prepared by neutralizing (6 < p H < 7) aliquots with the appropriate amount of neat carboxylic acid. An appropriate weighed quantity of ethanol was added a t this point. 2. Methods. Conductance measurements were made with a Wayne-Kerr universal bridge (B224). The cell was fitted with platinized platinum electrodes and was of a type designed for use with solutions which have a tendency t o froth (see Figure 13 of ref 12). The cell constant (15.72) was checked before and after each set of experiments by using standard potassium chloride.11 The carboxylate solutions were diluted by weight with the appropriate mixture of ethanol and water. All solutions were therrnostated at 25' (&0.05O). We could detect no change in conductance with time, once the test solution had reached the thermostat temperature. When appropriate, the conductivity was corrected for the conductivity of the solvent. An Olivetti Programma 101 programable calculator was used for the repetitive calculations.
Results Figure 1 shows equivalent conductance ( A ) plotted against the square root of the molar concentration, in 0.1 mole fraction ethanol-water, for decyltrimethylammonium dodecanoate, undecanoate, decanoate, nonanoate, octanoate, heptanoate, and acetate. The long-chain carboxylates form micelles within the concentration range of these experiments (indicated by a sharp decrease in A), but as the points observed above the critical micelle concentration (cmc) are not relevant here, we have simplified the figure The Journal of Physical Chemistry. Vol. 78, No. 17. 1974
D. 6.Oakenfull and D. E. Fenwick
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TABLE I: Decyltrimethylammonium Carboxylates in 0.10 Mole Fraction Ethanol-Water at 25 Carboxylate ion
Acetate Weptanoate Octanoste Nonnnoate Decanoate Undecanoate Dodecanoate
doRCO2Ka
hob
hoc
64.27 57.59 57.42 56.50 55.88 55.42 54.89
29.72 29.55 28.63 28.01 27.55 27.02
36 .40d 29.83 29.71 28.66 28.03 27.53 27 -05
K, M-1
Cmc, M
>o .02 >o .02
5.8 i 1.2 10.6 f 1 . 2 38.6 =t1 . 2 91.6 f 1 . 7 190 & 2 192 f. 3
0.014 0,0080 0.0056 0.0046 0.0036
Limitsing equivalent conductance (cm2 ohm-1 mol - 1 ) of the potassium salt. Limiting equivalent conductance of the decyltrimethylammonium salt, calculated from eq : ? ~ Limiting equivalent conductance obtained by using Davies' procedure. Obtained, directly, from the Onsager equation.
TABLE II: Calculation of'K for D e c y l t r i m e t h y l a m m o n i u m Decanoate in 0.10 Mole Fraction Ethanol-Water at 25 O
0,00397 0.00199 0.00156 0.00120 0.00079
20.70 22.97 23.81 24.49 25.25
0.804 0.872 0.898 0,918 0.938
0.854 0.887 0.897 0.908 0.923
89.5 94.8 90.1 88.6 95.0
a Observed equivalent conductance (cmz ohm-' mol--'). Fraction of dissociated ions, calculated from eq 1. Mean activity coefficient calculated from the Debye-Huckel equation, Mean value ia 91.6 with standard error f1.2.
TABLE 111: Comparison of Observed and Calculated Values of A for Decyltrimethylammonium Decanoate in 0.10 Mole Fraction Ethanol-Water at 25" Calcda A
0
0.05
0.1 0
vc
C, M
0.1 5
0.00397 0.00199 0.00156 0.00120 0.00079
Figure 1. Equivalent (conductance (A) of decyltrimethylammonium carboxylate solgtions iin 0.1 mole fraction ethanol-water at 25' plotted against & where Cis the molar concentration: (1) acetate, (2)
heptanoate, (3) octanoate, (4) nonanoate, (5) decanoate, (6)undecanoate, ~(7)dodecanoate. Broken lines show the theoretical (Onsager) slope. Vertical lines represent the cmc. by terminating the plots at the cmc. (Approximate values of the cmc are given in Table I.) The points for the acetate conform with the Onsager equationll (within the limits of experimental error) but those for the long-chain carboxylates lie consistently below the theoretical lines, indicated in Figure 1 by broken lines. The discrepancy is large enough (
