group, especially K. Katsuura, for many helpful discussions and suggestions during this work. Thanks
are also given to T. Ito for his assistance in the experimental work.,
Conductance Study of Model Hydrogen Bonding Solutes in Aqueous Solutions at 25"
by H. Olin Spiveyl and Fred M. Snell Department of Biophysics, State L'nicersity of N e w Y o r k , B u f f a l o , A'ew York
(Received J a n u a r y 15; 1564)
Conductance measurements a t 25' on dilute aqueous solutions of ethanoltrimethylammoniuin chloride, trimethylpropylamnioiiium bromide, and N-inethylbetaineamide chloride have been made. The limiting equivalent conductances of the cations with potential hydrogen bonding sites are compared with those of cations (alkyl analogs) of similar size and shape but lacking these sites. It is concluded that either the hydroxyl and amide groups on these cations are not associated with water molecules during transport or there exist mechanisms which compensate for their presence. Comparison of these data with mobility measurements on hydroxyl or amide containing nonelectrolytes in solution suggests that the affinity of solute hydrogen bonding groups for water molecules is affected by the neighboring atomic groups of the solute.
Hydrogen bond interactions between solute and solvent molecules may significantly affect both equilibrium and transport properties of these solutions. The purpose of this study is to assess, by means of conductance measurements, the extent of such interactions in the case of aqueous solutions of certain model solutes containing hydroxyl or amide groups.
Experimental Muteriuls. Ethanoltrimethylammonium chloride, (?\Ie3XEtOH)C1,and triniethylpropylamnionium bromide, (1leaNPr)Br,obtained from Eastiiian Co. and KH. X-methylbetaineamide chloride, (I\Ie3XCH2C0. CH3)C1, obtained from Starks A4ssociatesof Buffalo, N . Y. were recrystallized four to five times from isopropyl alcohol or a 50 vol. yo mixture of benzeiieisopropyl alcohol. Portions from the last two recrystallizations of each salt were stored in uucuo over P206until needed. Potassium chloride (A.R. grade) T h e Journal of Physical Chemistry
was recrystallized twice by hydrogen chloride precipi tation from water, then dried, and fused in an inert atmosphere. It was tested and found free of hydrovyl ion contamination.2 These crystals were used as the primary standard in titrations and conductanre measurements. Distilled water from the building supply was further distilled, once from dilute sulfuric acid and then from a potassium hydroxide-potassium permanganate solution in an all-glass still. This latter distillation provided water for the conductivity solutions as well as steam for cleaning glass vessels. (1) (a) This article is based on a thesis submitted by IT. Olin Spivey to Harvard Univervity in partial fulfillment of the requirements for the Ph.D. degree, 1963; (b) the work was begun at H a w a r d University and completed at the State University of ISew York at Buffalo; (c) Massachusetts Institute of Technology, Cambridge, Mass. 02139. ( 2 ) G. D. Pinching and R. G. Bates, J . R e s . A V d .Bur. S t d , 37, 311 (1946).
2127
CONDUCTANCE STUDY OF HYDROGEN BONDIXG SOLUTES
Specific conductances between 1 and 2 x 10-7 ohm-’ cm.-’ were obtained for the water in all experiments. Apparatus. One quartz and one Vycor flask cell of Shedlovsky design and 1-1. size were used for the conductance measurements together with apparatus described elsewhere.a We feel that some improvements in performance were achieved in the construction of a solid state proportional temperature regulator and an inductance ratio arm bridge. Also, special Teflon pipets of 5- and 25-ml. capacity were constructed which offered advantages of quantitative transfer of small amounts of stock solution from solution weight bottles with negligible conductance contamination. The thermostat was maintained within f 0.001O of the same reference temperature for all experiments, as judged by a Beckrnann differential thermometer. ‘The absolute temperature of the bath was 24.97 f. 0.02’ asrneasured by anS.B.S. calibrated, mercury in glass thermometer. Procedure. A stock solution (0.05-0.10 M ) of the salt to be studied was made by weight and successive portions were Lransferred from the solution weight bottles to the water or solution in the conductance cell, by means o f the appropriate Teflon pipet. Conductance measurements were made on the resulting solutions covering the concentration range IO-c 2 X M . Sitrogen gas, which was cleaned and equilibrated with water by passage through a set of gas dispersion bottles, was passed through the cell prior to and during the measurements in order to maintain the conductance a t a stable minimum. Differential potentiometric titrations with silver nitrate solutions4 and duplicate density nieasureinents were made on portions of the stock solution to deterinine concentratioizs accurately in molarity units (nioles/l.). Equivalent conductance values a t each concentration were calculated after correcting for background conductance and polarization resistances and the resulting data were used to evaluate the unknown parameters in the Fuoss-Onsager equations5 A
=
A0
-
+
SdC EC log c 4- (J - FA0)C unassociated electrolytes
(1)
Results Three potassium chloride conductance experiments, each with eight concentration points, provided a test of our procedure along with cell constant calibratioim7 Cell constants with a standard deviation of 0.02% or better were obtained in each of the experiments. The first two experiments using the Vycor and quartz cells, respectively, were made just prior to the studies on the quaternary salts; the last one with the quartz, cell was performed following these studies. Average cell constants found in each experimeiit are 1.0271 Table I : Equivalent Conductances of Quaternary Salts in Water a t 25”
10ac
A
lO8c
10ac h (MesNCHzCO-
A
(MeaNEt0H)Cl
(MeaNPr)Br
R u n 1, 3 x R e x , S:.*
R u n 1, 4 X Rex, Q.
2,2797 4.1352 5.8827 7.5747 9.1443
110.43 108.98 107.94 107.09 106.41
Run 2, 3 X R e x , Q. 0.47715 112.65 1.3443 111.35 3,5268 109.39 5.4154 108.23 7.2402 107.29
2.7305 5.0237 7.0715 9.1693 11.866 15.564
NH,CHa)CI
110.19 108.54 107.36 106.33 105.21 103.81
R u n 2, 5 X Rex, 1%. 0.41915 113.03 1.2242 111.75 3.8324 109.32 6.0223 107.90 8.1522 106.80 10.919 105 57 14.173 104.32 I
R u n 3, 4 X Rex, V,
4.2693 5.0642 8.0669 10.640 14.002 19.017
108.93 108.41 106.90 105.87 104.61 103,17
R u n 1, 3 X Rex,
0.25937 3.5477 6,4844 9.1828 11.492 13.500
Q.
108.45 104.60 102.98 101.76 100.86 100,19
R u n 2, 4 X Rex, I., 0.46232 107,97 2.3276 105.66 4,3952 104.14 6.7616 102.85 101. S7 8.8613 10,389 101,26
Run 3, 5X Rex, Q.
0.81194 1.6625 4.5959 7.6727 11.1416
lln36 111.25 108.81 107.07 105.52
Abbreviations: V., Vycor cell; Q., quartz cell; R e x , times recrystallized.
( 3 ) See ref. l a , pp. 26-28, 51-54. (4) D. A. Maclnnes, “The Principles of Electrochemistry,” Dover Publications, New Tork. N. Y., 1961, p. 306. 15) It. M . Fuoss and F. Accascina, “Electrolytic Conductance,” Interscience Publishers, Inc., New York, N. Y.,1959, pp. 105, 234.
Least-squares analyses of both equations were obtained using a program of Kay6 transcribed for an I B I I 1620 computer.
(6) R . 1,. Kay, J . Am. Chem. Soc., 82, 2099 (1960); we are grateful to Dr. Kay for sending us his revised program in 1962. (7) J. E. Lind, J. J. Zwolenik, and R. M. Fuoss, ibid., 81, 1557 (1959).
Volume 68, Number 8
August, 1964
2128
H. OLIS SPIYEYAND FRED11. SNELL
Table I1 : Conductance and Density Results
---
An & #(An)'
8 (A)
114.55 1 0 . 0 1 114.84f0.01 109.80 1 0 . 0 1
( MesNEtOH)C1 (Me3NPr)Br (MeaNCH&O.NHaCH$)Cl
0.02 0,02
0.01
Density
7
Conductance---
Gb
(J
0.00 0.00 0.00
- FAo)' 96.9 57.2 93.1
-
vo
I
Po-
32
a
Y
5Q
1.6070
0 012
7.2649 4.0246
0.037
127 145 145
109 120 127
0,023
b u and G, standard and mean deviation. a ho, limiting equivalent conductance. J - PAo, term from eq. 1. m, moles/kg. of solution a t which densit,y, d = 0.99707 am, was measured. e (ml./mole), partial molal volume of the salt a t infinite dilution; VO- = 18 and 25 for C1- and Br-, respectively (K. H. Stern and E. S. Amis, Chem. Rev., 59, 1 (1959)), for calculations of thepartialmolal volumes of the cations,
vo
+
VO+-
~
5/
infinite dilution with their standard deviations, m(Ao), and the standard and mean deviations of all experimental points, .(A) and sh, respectively, froin the unassociated electrolyte eq. 1. In these calculations all data for a given salt were used in evaluation of the terms in eq. 1. Remaining columns in Table I1 give density data from duplicate measurements and partial calculated from molal volumes at infinite dilution, apparent molal volumes, 4, and use of the Redlich equation,84 = Po 1.86dc. Figure 1 shows cation limiting conductance values, lo+, derived from these and similar measurements of others plotted as a function of n, the number of atoms other than hydrogen in the groups attached to the quaternary nitrogen. Table 111 summarizes various radii for cations of interest, calculated as indicated footnotes of Table 111.
4
4 40 51
35
vo,
+
25
-
Table 111: Ionic Radii,
Figure 1. Ionic conductance as inverse function of size. Ion
n
ha+
Ion
n
Xo+
4 4 . 9 [MegNCHaCO.NH.CH3]* 8 3 3 , 4 " [Mea?jEtl 5 41 , Z b [Et4Nl+ 8 32.7 [ MeaNEtOH] 6 3 8 . 2 " [Pr&] + 12 23.4 [Me,Nl'r] + 6 36.7" [BUN] + 16 1 9 . 5 [Me&(EtOH)z] 8 33 6* [Am4Nl 20 1 7 . 5 Ref. 9. All other Xo+ values from ref. 13, p. a This study. 463. The broken curve represents the corrected Stokes law (ref. 13 and Discussion) ; the line representing the uncorrected Stokes law, X = 91.7/~(.k,),was estimated assuming 0.5 A. change for each four - C H 2 groups removed from Am4N to which r = 5.3 A. was assigned (ref. 13). Chloride and bromide limiting conductances were taken as 76.4 and 78.1, respectively, for calculating A,,+ values (ref. 13, p. 463), whirh are probably all accurate to within 1 0 . 1 X unit.
[Me+X']+
4
+
+
+
I
+
+
em.-', Vycor cell; 0.62887 and 0.62872 em.-', quartz cell, listed in order of determination, Equivalent conductances, A, and molarities, c, calculated from the experimental measurements on the quaternary salt solutions are given in Table I. Table I1 lists values of equivalent conductances a t The Journal of Physical Chemistry
A. Tea
[Me,NEtOH] + [MesNPr] [R/IeJTCH&O.NH.CH~] +
+
[EtaNl+
[Me2N(EtOH)2] +
2.40 2.50
2.74 2.80 2.73
rhb
rVc
TCd
3.72 3.77 3.90 3.92 3.90
3.5 3.6 3,7
2.9-3.9 2.9-3.9 2,9-4 7 4.0 2.2-4 . O I
a T , = O.82O1zi/XO~o, Stokes law radius. * T h = corrected T., = 1 0 8 [ ( 3 / 4 1 r N ~ ) ~ ~ + + ]r'c/ ~ esti. Stokes radius, ref. 13. mated from Courtauld molecular models as minor-major semiaxes, assuming ellipsoid of revolution.
Discussion Experimental Accuracy. Our conductance values for potassium chloride solutions are within 0.02oJ, standard deviation of the equation suggested by Lind, Zwolenik, and Fuoss7as best representing selected data from the literature. The close agreement between &-values obtained using the quartz or Vycor (8) 0 . Redlich, J . P h y s . Chem., 67,496 (1963).
CONDUCTAKCE QTLDY
OF
2129
HYDROGEX BONDING SOLUTES
cell for the quaternary salt experiments further cor-. roborates the accuracy of the cell constants used Duplicate potentiometric titrations with silver nitrate solutions were made on each stock solution of quaternary salt except for N-niethylbetaiiieaniide chloride 3 X R e x , where one of the titrations was lost. Silver nitrate solutions were standardized by duplicate titrations with potassium chloride solutions prior to each stock solution determination. At least two endl point determinations were made for each titration. Standard deviations of concentrations calculated from, all end points averaged 0.01% and never exceeded 0.02570 for the stock solutions. For each standardization, a 0.01% agreement between weight determinedl concentrations and titration determined concentrations of silver nitrate was obtained, giving additional evidence of the purity of potassium chloride which served as the standard for the conductance data. The above considerations would lead us to judge that the errors in A- and &-values of these measurements are not in excess of +0.05 A unit. To our knowledge, the only cation in these experiments which has been similarly studied is [Me&EtOH] +. Varimbj and Fuossg report a io+ = 38.4; FleminglO reports io+ = 42.0, to be compared with the value of 38.2 herein reported. Conductance Equation. As indicated in Table 11, the data for each salt studied can be represented bv the unassociated electrolyte eq. 1 within a standard deviation of 0.02% and negligible mean deviation, i.e., the data are randomly distributed about this equation. Analysis of the data by the associated electrolyte equation provided association constants less than unity which are negligible in dilute solutions and of the order of experimental precision. However, more recent theoretical and experimental considerations” have indicated that association constants less than 401 may be in considerable error and that in general, little physical significance may be attached to the distance of closest approach parameter, a, derived from J . We, therefore, report only the lumped coefficient of the linear term in the Fuoss-Onsager equation, ( J - FRo)c. These uncertainties in the highel. order terms have little effect on the intercept, A,, however. Molecular Interpretations. Each hydroxyl group has the potentiality of two hydrogen bonds, and the amide groups the possibility of three (not to mention the possibility of secondary hydration). On the other hand, Courtauld molecular models and the derived radii, yo, listed in Table I11 indicate that, unless there are special compensating mechanisms, a single water molecule hydrogen bonded to the hydroxyl or amide
groups of the model cations would lower their mobiljtiesI2 below that of the alkyl analogs of comparable size. This lowering would be a t least one X-unit and probably two or more, judging from the effect of an added -CH2 on Xo+ seen in Fig. 1 in the region n =r 6-8. Instead, we have found quite similar, but slightly higher mobilities for the model cations (n = 6 and 8) than for their alkyl analogs. This is most simply interpreted as evidence that the -OH and -CO.KH*CH, groups of these solutes are not hydrogen bonded to water. We consider the data next in terms of classical hydrodynamic theories. The dashed curve in Fig. 1 is drawn through the conductance values of the symmetrical tetraalkylammonium ions with n = 4, 8, 12, 16, and 20 and represents the corrected Stokes law suggested by Robinson and Stokes.la This curve demonstrates the experimental fact that as the radii of solute particles decreases (below 5 A. in water) relative to the solvent, the mobility of the particles increases progressively faster t@n predicted by the linear Stokes law, h = 91.7/r(A.). Thus, corrected Stokes radii, rh, calculated for the ions of interest here, (Table III)la differ still less in percentage than ’ the mobilities, due to the nonlinear relation between them. For example, the 4,1% mobility difference between [Y!e3NEtOH]+ and [Me3NPr]+requires only o 1.301, difference in l^h or 0.05 8. Furthermore, on the assumption that there is negligible solvent interaction with the symmetrical alkylammonium ions (or interaction similar to that of other ions), deviations from this curve may be attributed to molecular asymmetry and/or specific solvent interactions. Indeed, the low mobility of the [Me3NPr]+ion compared to the symmetrical cations seen in Fig, 1 is not unexpected, considering its asymmetry and the effects of rotational Brownian motion. Thus, the small differences that are found in the cation mobilities seem explicable in terms of existing hydrodynamic theories without invoking any specific interaction with solvent molecule^.^^
(9) J. Varimbi and R. M . Fuoss, J . Phys. Chem., 64, 1335 (1960). (10) R . Fleming, J. Chem. Soc., 4914 (1960). (11) D. J. Karl and J. Dye, J . Phys. Chem., 6 6 , 477 (1962); R. M. Fuoss and L. Onsager, ibid.,66, 1722 (1962); R . L. Kay and J. Dye, Proc. N d . Acad. Sci. U . S., 49, 5 (1963); H. 0. Spivey, ref. l a , p. 7ff. (12) Mobility, u = 6.466 X 1OeX/ I z l (cm. set.-' dyne-’), Bee, e.g., ref. 13, pp. 42, 43. (13) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” 2nd Ed., Butterworth and Go., Ltd., London, 1959, pp. 123-125. (14) Radii derived from the partial molal volumes, T ” , are consistent with this view a s well as reasonable in comparison with the other radii, though molecular interpretations from these small differences are certainly questionable.
Volume 68, Number 8
Auguet, 1964
2130
Alternative explanations should be considered which might permit hydrogen bond interactions of water to the cations without causing a decrease of their mobility. Thus, e.g., Me&+Pr ions might be surrounded by solvent with a larger local viscosity than the bulk value due to the structure promoting effect of the nonpolar groups as suggested by Frank and Evans.'6 Indeed, viscosity'6 and heat capacity" measurements on the symmetrical quaternary ions do indicate an unusually large structure promoting influence of these cations on aqueous solutions. In contrast, such a structured region about the more polar [?tle,NEtOH]+ ion might be prevented or less favored by solvent interactions with the hydroxyl group. The ideas of Samoilov18whereby the frequency of exchange between water molecules adjacent to a solute is increased may lead to the same result. Also, the attraction between the charge on the nitrogen and the polar functional groups may cause a sufficient contraction in radius to mask the effects of bound water. It seems unlikely, however, that these special mechanisms could compensate for the effects of bound water molecules (particularly for two or more) in the case of each model solute. Finally, n7e compare these ideas and results with those from related studies of others. Viscosity measurements on the system dimethylacetamide and waterlg and dimerization constant measurements of N-methylacetamide in water, in comparison with measurements in dioxane, carbon tetrachloride, or benzene, *O have been interpreted as evidence for water hydrogen bonded to these amides. It is likely, however, that the high concentrations of amides in water encountered in these experiments (7-12 M ) could prevent the normal hydrogen bond structure of water and hence shift the equilibrium in favor of solute-solvent interactions on this basis. On the other hand, Longsworth21 calculates hydration numbers a t infinite dilution of 3.0,2.5,and 1.9 for acetamide, methylacetamide, and dimethylacetamide, respectively; for hydroxy
The Journal of Physical Chemistry
H. OLIKSPIVEYAND FRED M. SNELL
compounds, hydration numbers of approximately one molecule per hydroxyl group are found. If these hydration numbers are accepted and attributed to hydrogen bonding, comparison with our conductance data, where little evidence of hydrogen bonding appears, suggests that such interactions may be affected by neighboring atomic groups of the solute. Such differences in behavior between the cations of this study and the nonelectrolyte solutes in other studies mentioned might be due to the charge on the nitrogen atom interfering with the neighboring hydrogen bonding sites or they may be due to alternative mechanisms as referred to above. In any case, they indicate that a variety of behavior may be observed with such solutes in water and suggest that considerable caution should be used in deductions and generalizations based on these models.
Acknowledgments. We wish to express our thanks to Dr. Robert A. Spaiigler for advice and assistance on electronic problems, Mr. Edward Gill for assistance with some of the measurements, and Mr. Basilio Scofidio for work on the computer program and calculations. We are also grateful t o Dr. T. Shedlovsky for timely help in regard to the potentiometric titrations and construction of the conductance cells. Economic support for this study was provided in part by U. S. Public Health Service Research Grant RG6730 and Training Grant 571GM718. (15) €3. S Frank and M.W. Evans, J . Chem. Phvs., 13, 507 (1945). (16) E. Huckel and H. Schaaf, 2. physik. Chem. (Frankfurt), 21, 326 (1959). (17) H. S. Frank and W. Y . Wen, Discussions Faraday SOC.,24, 133 (1957).
(18) 0 . Ya. Ramoilov, ibid.,24, 141 (1867). (19) R. C. Petersen, J . Phys. Chem., 6 4 , 184 (1960). (20) 7. M. Klots and J. S. Franaen, J. Am. Chem. Soc., 84, 3461 (1962); M. navies and D. K. Thomas, ,I. Phvs. Chem., 6 0 , 767 ( 1 9 5 6 ) ; G. E. Ribberd and A. E. Alexander, ibid., 66, 1854 (1962). (21) L. G Longsworth, ibid., 67, 689 (1963),and personal communioation.
