OSMOTICAND ACTIVITY COEFFICIENTS OF QUATERNARY AMMONIUM HALIDES
82 1
Structural Effects on the Osmotic and Activity Coefficients of the Quaternary Ammonium Halides in Aqueous Solutions at 25”’
by G. E. Boyd, A. Schwarz, and S. Lindenbaum Oak Ridge National Laboratory, Oak Ridge, Tennessee 8Y88l (Received October 18, 1966)
The quaternary ammonium halides derived from tetramethylammonium chloride and bromide by substituting either a P-hydroxyethyl or a benzyl group, or both, for methyl groups were employed in a study of the effect of cation structure on electrolyte behavior in aqueous solution. The osmotic and mean molal activity coefficients of the (P-hydroxyethyl) trimethylammonium (choline) salts were slightly smaller than those for the corresponding tetramethylammonium salts a t all concentrations. The coefficients for benzyltrimethylammonium chloride and bromide were appreciably smaller than those for Me4NC1 and MerNBr, respectively, while those for (P-hydroxyethy1)benzyldimethylammonium chloride and bromide were slightly smaller than those for the benzyltrimethylammonium salts at all concentrations. The “water-structure-breaking” action of the P-hydroxyethyl group was assumed to be the cause for the lower coefficients of the choline halides relative to the tetramethylammonium salts. The very small activity coefficients of the benzyland the ethanolbenzyl-substituted derivatives have provided a basis for understanding the relatively large selectivity coefficients for bromide over chloride ion shown by cross-linked, strong-base anion exchangers.
The quaternary ammonium halides are of interest in researches on the physical chemistry of aqueous electrolyte solutions because of the possibilities they present for studies of the effect of cation size, shape, and structure on thermodynamic and transport properties. In a previous research2 the influence of cation size in the homologous series of symmetrical tetraalkyIammonium ions on the solute activity coefficients was measured. An increase in size was accompanied by an increase in the activity coefficients of the chloride salts in dilute solutions, but with the bromides and iodides a decrease was observed. The behavior of the chlorides was explained in terms of the ability of large organic cations to enforce the water structure, and, the more carbon atoms in an ion, the larger this effect. With the bromides and especially the iodides, however, ion-pair formation occurred to an increasing extent with increasing cation size, and this was reflected by the opposite sequence in the activity coefficients. The effect of substituting other than naliphatic groups in the tetramethylammonium cation was therefore of interest. Accordingly, in this paper
the change produced in the activity coefficient by replacing a methyl group in tetramethylammonium chloride and bromide by a 6-hydroxyethyl group was measured as well as the change on substituting a benzyl for a methyl group. The consequence of substituting both P-hydroxyethyl and benzyl groups for methyl groups also was determined. Two of the compounds examined, trimethylbenzylammonium and (P-hydroxyethyl)dimethylbenzylammonium chloride and bromide, were of additional interest because these are “model compounds” for the strong-base anion exchangers, Dowex-1 and Dowex-2, respectively. Thus, additional insight was obtained as to the role of ionogenic group structure in determining anion-exchange selectivity.
Experimental Section Materials. (p-Hydrox y e t h yl) t rime t hyl ammonium chloride (choline chloride), obtained from Eastman (1) Research sponsored by the U. S. Atomic Energy Commission under contract with the Union Carbide Corp. (2) S.Lindenbaum and G. E. Boyd, J . P h y s . Chem., 6 8 , 911 (1964).
Volume 7 0 , .\‘umber
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G. E. BOYD,A. SCHWARZ, AND S. LINDENBAUM
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Organic Chemicals Co., Rochester, N. Y., was recrystallized from ethanol and dried. Trimethylbenzylammonium chloride was purchased from Matheson Coleman and Bell, East Rutherford, N. J., as a 60% aqueous solution. Trimethylbenzylammonium bromide and choline bromide were prepared from the chloride by treatment with silver hydroxide and neutralization of the filtrate with HBr. Dimethyl-/3hydroxyethylbenzylammonium chloride and bromide were prepared by reacting 2-dimethylaminoethanol with benzyl chloride and bromide. The various salts were recrystallized from appropriate organic solvents, and their purity was determined by gravimetric analysis for their halide content by precipitation with AgN03. Osmotic Coeficient Measurements. Osmotic coefficients of binary aqueous solutions of the various quaternary ammonium chlorides and bromides were measured with two different, complementary techniques. The gravimetric isopiestic vapor pressure comparison method3 was used on solutions more concentrated than ca. 0.2 m following procedures already described.2 However, with dilute solutions equilibrium was attained extremely slowly with the gravimetric method. Accordingly, measurements on choline chloride (ChCI) and bromide (ChBr) in the concentration range 0.01 to ca. 0.5 m were conducted with a thermoelectric ~ s n i o m e t e r . ~I n this method a steady-state temperature difference between a reference solution (KCl) and the test solution and pure water was determined. Thermistors were employed as temperature sensors, and their electric resistance was measured with a Wheatstone bridge arrangement. The osmolality ( W Z ~ of) ~the test solution was computed from the osmolality (m4), of the reference solution and the measured resistance ratio (AR,/AR,) with the relation
(m4L
=
(AEJARJ(m4)r
(1)
A plot of the concentration dependence of the osmotic coefficients for dilute aqueous solutions of choline chloride and bromide is shown in Figure 1, where it may be seen that satisfactory precision as well as a reasonable approach of the data to the Debye “limiting law” behavior was obtained. Moreover, these measurements appeared to lie on the same smooth curves as those determined by the gravimetric method. The measured osmotic coefficients, eq 1, and the isopiestic solution concentrations are given in Tables I and 11. Osmotic coefficients were calculated from the isopiestic solution concentrations with the relation
vXmx4,= Vrmr4r (2) where v is the number of ions (v = 2), m, is the molality The Journal
0.f
Physical Chemistry
Table I Osmotic coefficients for dilute solutions from thermoelectric osmometer
a.
mchci
9ChC1
mChBr
0.0126
0.9606
0.0122 0.0253 0.0849 0,1291 0.2788 0.4034 0,7447
...
...
0,0801
0.9083
...
I
0.2955 0.4587 0.7602
.
.
0.8675 0.8547 0.8370 b.
“aC1
1.1044 1.4525 2.0784 2.5324 2.9706 3.3918 4.0083 4.2903 4.7994 5.1191
...
9ChBr
0.9569 0.9367 0,9047 0,8901 0.8516 0,8309 0.8015
Molalities of isopiestic solutions mchci
mtiaci
mChBr
1.2250 1.6228 2.3552 2.8569 3.3472 3.8423 4.5740 4.9133 5.5017 5,9076
0.9937 1.1044 1.4525 2.0784 2,5324 2.9706 3.3845 4.0083 4.2903 4.7994 5.1191
1,1930 1.3407 1.8028 2.6674 3.2960 3.9191 4.5172 5.4649 5.8986 6.6624 7,1514
...
of the quaternary ammonium salt, and 4, is its molal osmotic coefficient; mr and 4r are the molality and osmotic coefficient of the reference electrolyte (NaCl), respectively. The necessary values of 4r were computed from a least-squares “fit” of the concentration dependence of the osmotic coefficients of the reference NaCl solution^.^ Mean molal activity coefficients were calculated from the osmotic coefficients with the Gibbs-Duhem equation JO
The integration required in eq 3 was performed numerically taking into account the fact that6 lim (1 -
m+O
+)/G = a / 3 = 0.3903
(4)
Computed osmotic and activity coefficients for interpolated concentrations are given in Table 111. (3) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” 2nd ed, Academic Press Inc., New York, N. T.,1959, p 177 ff. (4) A Model 301 vapor pressure osmometer manufactured by Mechrolab, Inc., Mountain View, Calif., was employed. ( 5 ) M. H. Lietzke and R. W. Stoughton, J . Phys. Chem., 6 6 , 508 (1962). (6) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” Butterworth and Co. Ltd, London, 1955, p 232 ff.
823
OSMOTIC AND ACTIVITY COEFFICIENTS OF QUATERNARY AMMONIUM HALIDES
Table 11: Molalities of Isopiestic Solutions MezOEtBzNCl
MezOEtBrNBr
0.2026 0.2726 0.2877
0.1944
0.2059
... ...
...
0.6056 0,6054
0.7157 0.7154
...
...
0,6208 0.6233 0.7900
0.7379 0.7409 0.9906
...
...
...
...
...
1.715
0.8462 0.9105 1.348
1.070 1,175 1.835
1.018 1,039 1,100
1.370 1.401
1.851 1.906
... ...
...
...
... ...
1,581
2.306
1.274 1.348 1.363
1.773 1,882
...
2.547 2.736 2.776
1.915 2.039 2.066
2.825 3.047 3.111
1.623 1.668 1.838
...
..
2.385
3.537
...
...
2.514 2.619 2.898
3,812 3.976 4.391
2.358 2.530 2.641
3.442 3.685 3.850
5.130 5.474 5,727
3,845 4.139 4,344
5.765 6,216 6.472
2.746 3,102 3.128
...
...
4,536
6.687
4,416 5.058 5.159
6.589 7.407 7,605
4.105 4.213 5,002
...
6.780
8.914 10.457
...
9.862
6.069 6.995
...
... ...
9.268 9.314
12.662
,..
waci
MesBzNCI
0.1799 0,2350 0.2470
0.1933 0.2547 0.2679
0.5043 0.5063 0.6215 0.6542 0.6974 0.9453
5.760 5.762 6.144
... .
I
.
... 8.451
MesBzNBr
.
I
.
. I .
.
.
I
... ... 13.42 13.67
...
The osmotic and activity coefficients for choline chloride in Table I11 do not agree with previously reported values7 derived from measurements made with the gravimetric isopiestic solution comparison technique. The published activity coefficients are larger than those for tetramethylammonium chloride at all concentrations, whereas our measurements show that choline chloride and bromide have smaller coefficients than the corresponding tetramethylammonium salts. Several attempts were made to determine the cause for the discrepancy between these two series of measurements on choline chloride. The isopiestic solution concentrations given in ref 7 were employed in a complete digital computer recalculation of the osmotic and activity coefficients according to eq 2 4 . The recalculated values differed from those published by ap-
'
I
0.i
I
I
I
I
I
I
0.2 0.3 0.4 0.5 0.6 0.7
m
I
I
I
I
0.8 0.9 t.0 1.1
I
t
1.2
Figure 1. Concentration dependence of the osmotic coefficients for aqueous solutions of choline chloride and bromide a t 25" (measurements with thermoelectric osmometer).
proximately 1.5%, so that computational errors do not appear to be a cause for the discrepancy. A direct comparison of approximately 1 m solutions of NaCl, KC1, Me4NCl, and ChCl was made with the gravimetric isopiestic apparatus in a special study with the results given in Table IV. The osmotic coefficients for KCl and Me4NC1 were in good agreemerit with published values, and it may be seen that the coefficient for choline chloride was distinctly less than that for Me4NCl. The fact that the substitution of a phydroxyethyl for a methyl group in trimethylbenzyl chloride (and bromide) also lowered the osmotic and activity coefficients relative to the values for the latter salts also lends support to our belief in the correctness of the 4 and y values for choline chloride given in Table
111.
Discussion The introduction of a /3-hydroxyethyl &roup either in tetramethylammonium chloride or bramide or in trimethylbenzylammonium chloride or bromide produced a lowering in the mean molal activity coefficient a t all molalities (Figure 2). The substitutioh of a CHzOH group increased the size of the quaternary ammonium cation as measured by the apparent mala1 volume a t infinite dilution, t&O, from 89.1 ml for tetra(7) R. Fleming, J. Chem. Soc., 3110 (1961).
Volume 70,Number S March 1966
G. E. BOYD,A. SCHWARZ, AND S. LINDENBAUM
824
~
Table 111: Osmotic and Activity Coefficients at 25' for Interpolated Molalities ,__-
m
0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 1.2 1.4 1.6 1.8 2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5 6.0 7.0 8.0 9.0 10.0 11.o 12.0 13.0
Y-ChBr-
ChCl--
Q
Y
Q
Y
9
Y
Q
0.715 0.642 0.592 0.557 0.530 0.508 0.490 0.474 0.459 0.447 0.425 0.409 0.396 0.385 0.375 0.356 0.343 0.335 0.328 0.324 0.320 0.317 0.316 0.316
0.900 0.860 0.830
...
0.900 0.872 0.849 0.834 0.823 0.814 0.807 0.800 0.794 0.788 0.779 0.774 0.772 0.770 0.770 0.771 0.775 0.783 0.793 0.802 0.811 0.820 0.831 0.856
...
...
...
0.805 0.785 0.768 0.753 0.740 0.729 0.720 0.704 0.693 0.686 0.681 0.677 0.675 0.680 0.692 0.708 0.727 0.750 0.775 0.800 0.852 0.913
0.733 0.641 0.587 0.546 0.513 0.486 0.462 0.441 0.423 0.407 0.378 0.356 0.338 0.324 0.312 0.291 0.276 0.265 0,259 0,255 0.252 0.251 0,252 0.261 0.272
...
...
... ... ...
0.882 0.826 0 783 0.742 0.708 0.679 0.654 0.633 0.615 0.599 0,570 0,547 0.528 0.512 0.498 0.475 0.465 0.459 0.458 0.458 0.462 0.468 0,477 0,498 0.519 0.541 0.561 0.579 0.602
0.709 0.596 0.527 0.476 0.435 0.402 0.375 0.350 0.329 0.310 0.278 0.251 0.230 0.215 0.202 0.175 0.157 0.144 0.134 0.126 0.119 0.114 0.109 0.103 0.099 0.095 0.093 0.091 0.090
0.887 0.854 0.823 0.797 0.777 0.755 0.738 0.723 0.710 0.697 0.677 0.661 0.648 0.638 0.631 0.619 0.613 0.614 0.618 0.626 0.636 0.646 0.658 0.689 0.726 0.766
0.730 0.639 0.585 0.538 0.500 0.469 0.445 0.423 0.405 0,386 0,358 0.333 0.314 0.299 0.286 0.260 0.240 0.226 0.217 0,209 0.203 0.198 0.194 0.188 0.186 0.184
... ...
...
...
...
...
0.876 0.812 0.766 0.725 0.694 0.664 0.638 0.615 0.595 0.576 0.544 0.518 0.498 0.481 0.466 0.440 0.422 0.411 0.406 0.404 0.405 0.409 0.416 0,432 0,452 0.475 0.496 0.513 0,526 0.532
0.914 0.886 0.870 0.859 0.853 0.848 0.846 0.844 0.843 0,844 0.847 0.850 0.854 0.858 0.863 0.881 0.908 0.932 0.953 0,975 0.999 1.025 1,046
0.737 0.664 0.621 0.590 0.567 0.549 0.535 0.523 0,513 0.505 0.492 0.482 0.474 0.468 0.464 0.458 0.462 0.467 0.474 0.482 0.493 0.506 0.519
... ...
...
...
...
...
... ... ...
... ...
-MenOEtBzBr--
Y
Y
...
-MerBzNCI--
Q
Q
... ... ...
... ...
... ...
.
.
I
...
...
... ...
... ...
...
Y
0.701 0.596 0.522 0.464 0.418 0.380 0.349 0.325 0.304 0.286 0.256 0.234 0.215 0.199 0.185 0.158 0.140 0.127 0.116 0.109 0.102 0.097 0,092 0.086 0.081 0.077 0.074 0.072 0.070 0.068
proach of its osmotic coefficient to the "limiting law" slope shown in Figure 1. Even with choline iodide, conductivity measurements'O on dilute solutions Salt m # have indicated virtually no ionic association. A possible explanation for the activity coefficient lowering 0.9401 ... NaCl is that the P-hydroxyethyl group in the quaternary am0.9794 0.896 KC1 MeaNCl 1.0147 0,865 monium ion interacts with water and disrupts its 1.0372 0,846 ChCl structure slightly, thereby raising its free energy and, hence, lowering the free energy of the cation in accordance with the Gibbs-Duhem equation. methylammonium* to 106.9 ml mole-1 for ~ h o l i n e . ~ Independent evidence bearing on this hypothesis This change of 17.8 ml mole-' may be compared with appears to be somewhat contradictory. Conductivity the increases of 15.1 ml mole-' in going from (CHa)qN+ measurementslo on the ethanolammonium iodides in to (C2H&N+. An augmentation in size alone would water have been interpreted as indicating the absence be expected to cause a small increase in the activity of interaction between the cation and the hydrogencoefficient of the choline salt similar to that found bond structure of water: the cationic conductances previouslyz in researches with tetra-n-alkylammonium were determined by the number of heavy atoms in the chlorides. The observed lowering, therefore, must be side chains of the quaternary ammonium ions, inderelated to the introduction of a large dipole moment into the cation. The lowered activity coefficients in (8) W. Y . Wen and S. Saito, J . P h y s . Chem., 68, 2639 (1964). choline chloride cannot be attributed readily to the (9) R.Fleming, J . Chem. Soc., 4914 (1960). formation of ion pairs in view of the nature of the ap(10) J. Varimbi and R. M. Fuoss, J . P h y s . Chem., 64, 1335 (1960). Table IV : Concentrations of Isopiestic Solutions and Computed Osmotic Coefficients at 25"
The Journal of Physical Chemistry
OSMOTICAND ACTIVITY COEFFICIENTS OF QUATERNARY AMMONIUM HALIDES
CONCENTRATION DEPENDENCE OF MEAN MOLAL ACTIVITY COEFFICIENTS OF AQUEOUS. SOLUTIONS OF QUATERNARY AMMONIUM CHLORIDES AT 25%
0.8
I
I
1.0
2.0
I
I
3.0 4.0 MOLALITY
I 5.0
I 6.0
I 7.1
Figure 2. Concentration dependence of mean molal activity coefficients of several quaternary ammonium chlorides in aqueous solution a t 25": curve 1 : tetramethylammonium chloride (data from ref 2); curve 2: (p-hydroxyethy1)trimethylammonium(choline) chloride; curve 3 : trimethylbenzylammonium chloride; curve 4: (p-hydroxyethy1)dimethylbenzylammoniumchloride.
pendently of whether these were carbon or oxygen. However, a more recent conductance study'l of model hydrogen-bonding solutes in aqueous solutions at 25", which included (P-hydroxyethyl)trimethylammonium chloride, has concluded that alternative explanations should be considered which might permit interactions of water with such dipolar cations without causing a decrease of their mobility. Calorimetric measurenients12 on aqueous solutions of the related compound, tetra(@-hydroxyethyl)ammonium bromide, have shown that the introduction of OH groups into an n-alkyl quaternary ammonium ion greatly diminished the heat capacity effect observed earlier with BurNBr, suggesting that a significant interaction with water must have occurred. Quite recently, activity coefficient and apparent molal volume measurements on tetra(@-hydroxyethyl) ammonium fluoride and bromide solutions
825
have appearedI3which also indicate that the OH groups interact with water and make the behavior of the cation much less abnormal. The substitution of a benzyl for a methyl group in tetramethylammonium, or in choline chloride or bromide, produced a large decrease in the osmotic and activity coefficients, especially at high concentrations (Figure 2 ) . A significant increase in the size of the quaternary ammonium ion was produced by this substitution, so that an increase in the activity coefficients at least for the chloride might have been expected. The occurrence of extensive ion-pair formation in the case of the bromide salt, at least, was suggested by the fact that its osmotic coefficients fell below the "limiting law" values (Figure 1) at the lowest concentrations. The relatively small activity coefficients for the benzyl-substituted tetramethylammonium chloride and bromide help to explain the fact that Dowex-1-type strong-base anion exchangers show greatly increased selectivity coefficients for bromide over chloride ion when compared with strong-base exchangers prepared by reacting polyvinyl chloride with ethylenediamine. In the latter exchanger the ionogenic groups are trimethylethylammonium salts, while in Dowex-1 the groups are trimethylbensyl. With the former exchanger the selectivity coefficient for the uptake of bromide ion by the chloride form is only slightly greater than unity,I4 while for Dowex-1 a value of 1.9 was found15 even with very lightly cross-linked preparations. This increase follows from the general principle that the stronger the ion binding by a polyelectrolyte, the greater the ionic selectivity. The selectivity coefficient is proportional approximately to the square of the ratio of the mean molal activity coefficient of the chloride to the bromide salt. Thus, for example, from previously published data2 and Table 111, we estimate values of 1.3 and 1.8 at 1 m for the selectivity coefficients of anion exchangers where the cation is tetramethylammonium and trimethylbenzylammonium, respectively. (11) H. 0. Spivey and F. M.Snell, J . Phys. Chem., 68, 2126 (1964). (12) (a) H. S.Frank and W. Y . Wen, Abstract 30R, 135th National Meeting of the American Chemical Society, Boston, Mass., April 5-10, 1959; (b) H. S. Frank, private communication, Sept 1965. (13) W. Y. Wen and S. Saito, J . Phys. Chem., 69, 3569 (1965). (14) W. Slough, Trans. Faraday Soc., 5 5 , 1036 (1959). (15) G. E. Boyd, S. Lindenbaum, and G. E. Myers, J . Phys. Chem., 65, 577 (1961).
Volume 70, Number 3 March 1966
