0. D. BONNER
2512
Solute-Solvent Interactions in Aqueous Media by 0. D. Bonner Department of Chemistry, University of South Carolina, Columbia, South Carolina
29208
(Received January 3, 1968)
Spectroscopic measurements of the solvation of alkali halide salts in aqueous solutions are correlated with the colligative properties of these solutions as represented by their activity coefficients Evidence for ion pairing in some salts is based upon the concentration dependence of the absorption at 958 mp. The spectra of the aromatic sulfonates and certain nonelectrolytes indicate that these solutes exert a “structure-breaking” effect upon the solvent. Solutions of all of the nonelectrolytes in this category are found to exhibit positive deviations from Raoult’s law, in contrast to solvated solutes, such as dextrose, which have activity coefficients greater than unity. The small activity coefficients of the aromatic sulfonates also correlate with the spectral findings.
Preliminary studies of water and aqueous solutions in this laboratory,l in which the near-infrared absorbance of solutions is compared with that of pure water at the same temperature by a differential method, have yielded quantitative data on the solvation of both electrolytes and nonelectrolytes. The “solvation numbers” of the alkali metal chlorides were found to be unusual, in that the observed order of LiCl > CsCl > KC1 > NaCl does not correspond with the lyotropic series of activity coefficients or cation-exchange equilibria. It is significant to note that for proton magnetic resonance m e a ~ u r e m e n t s in ,~~ which ~ the magnitude of the effect of the ions on the water protons should be directly related to the chemical shift, the results are in complete qualitative agreement with those obtained from spectral measurements. As an explanation of the above unexpected findings, it was postulated that “hydration numbers” calculated from the colligative properties of the solution include both ion-solvent and ion-ion interactions, while those measured spectroscopically represent only the former. This article presents further experimental evidence to support this postulate.
efficient-concentration curves of aqueous solutions of the alkali halides was discussed by Frank.5 He pointed out, however, that many structural interpretations are available to explain each phenomenon and that more work needs to be done to clarify the picture. In the discussion of solutions, the solute-solvent interactions may be grouped together as “solvation” and the ionion interactions may be classified as “ion pairing.” In an estimation of the relative importance of these two types of interaction in aqueous solutions of simple electrolytes, one must take into account several peculiarities in the order of activity coefficients of the alkali metal halides, nitrates, and sulfonates (Table I). The sequence for salts of any anion is Li+ > Na+ > I R b + > Cs+. If tetraalkylammoniums alts are included, the RdN+ ion follows Cs+, as would be expected. If one notes a particular cation, however, it is observed that for the activity coefficients of lithium salts I- > Br- > C1- > NOa- > RSOI- > OH-. The sequence for the halide salts of Rb+, Cs+, and RdN+ is, on the other hand, the reverse of that of the lithium salts.
Experimental Section
Results T y p e s of Solvation. The activity coefficient data are
All spectra were recorded using a Cary Model 1411 spectrophotometer. The differential technique which was used has been described previously.’ All electrolytes except the sulfonates were reagent grade and were used without purification. The sulfonates were prepared and purified by methods which have been described in a previous p~blication.~Propylene carbonate and ethylene carbonate were triply distilled a t reduced pressure, while dimethylurea was twice recrystallized from benzene.
Discussion It has long been recognized that the structural properties of the solvent will need to be taken into account if the ultimate goal of electrolyte theory is reached. The “plausibility in the idea that water-structure influences play a part” in determining the activity coThe Journal of Physical Chemistry
consistent with the solvation datal if the latter are interpreted in terms of the two types of solvation which have been postulated and the effect that this would be expected to have upon ion pairing. First consideration may be given to the halides of the two extreme alkali metal cations Li+ and Cs+. The smaller lithium ion should polarize the water molecules strongly and should orient them so that the oxygen end of the dipole is nearest the ion. The cesium ion, on the other hand, (1) 0. D. Bonner and G. B. Woolsey, J . Phys. Chem., 72,899 (1968). (2) J. C. Hindman, J . Chem. Phys., 36, 1000 (1962). (3) M. S. Bergqvist and E. Forslind, Acta Chem. Scand., 16, 2069
(l9G2). (4) 0. D. Bonner and 0. C. Rogers, J . Phys. Chem., 64, 1409 (19GO). (5) H. 8. Frank, 2.Phys. Chem. (Leipiig), 228, 364 (1965).
SOLUTE-SOLVENT INTERACTIONS IN AQUEOUS MEDIA
251 3
Table I : Activity Coefficients of Aqueous Solutions a t 0.5 m Concentrations“ OH
Li Na K Rb cs MeaN EtaN
0.583 0.688 0.712 0.782
F
0.632 0.670
c1
Br
I
NOa
Tolb
M eac
0.739 0.681 0 649 0.634 0.606 0.597 0.600
0.753 0.697 0.657 0.632 0.603 0.588 0.505
0,824 0,723 0.676 0.629 0.599
0.726 0.617 0.545 0.534 0.528
0.659 0,627 0.605
0.632 0.598
0.406
R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” Academic Press Inc., New York, N. Y., 1959; 0. D. a Data are from: Bonner and 0. C. Rogers, J. Phys. Chem., 64, 1499 (1960); S. Lindenbaum and G. E. Boyd, ibid., 68, 911 (1964). Toluenesulfonate. Mesitylenesulfonate.
’
because of its more shielded nuclear charge, is not so effective in polarizing the solvent but appears to enhance the structure of water around it in the manner of the tetraalkylammonium ion or the inert gases. The anions should not be nearly so effective as the lithium ion in orienting the water dipole, but the fluoride and hydroxide ions should be the most effective. One might expect that ion pairing would be almost nonexistent for the lithium salts, except perhaps for the cooperative type shown in Figure 1 which involves a solvent molecule. This model is consistent with the large activity coefficients of the lithium salts and with their order. Cesium salts should be much more susceptible t o ion pairing, since the cation is not protected by a sheath of polarized solvent molecules. The chloride ion, because of its greater solvation, is more effective than bromide or iodide ion in resisting ion pairing, and the order of activity coefficients for the cesium salts results. The other alkali halides which are listed show behavior intermediates between lithium chloride and cesium iodide. The hydration numbers which were reported in the first article’ were those obtained by extrapolation of the data to infinitely dilute solutions. It was subsequently realized that, if the above postulated type of ion pairing occurred, it should be capable of experimental verification. The absorption at 958 mp, which is used as a measure of the solvation of the solute, should change less rapidly with increased concentration of LiCl solutions, since the concentration of “free water” is decreasing and the lithium ion solvation in concentrated solutions is consequently decreased. Concentrated solutions of cesium salts should, however, exhibit increased solvation, since the water structure is enhanced around an increasing number of ion pairs which are larger than the individual ions. This behavior is verified in Figure 2, where a plot is given of I I C as a function of the concentration. Dextrose is included as an example of a highly solvated nonelectrolyte for which ion pairing is not possible. The slopes of the molal absorbance curves are found to have opposite arithmetic signs for the two different types of solvation.
H
W
H
Figure 1. A solvated ion-pair “model.”
0
1.0
2.0
3.0
4.0
5.0
6.0
mlality
Figure 2. Molal absorbance of solutions a t 958 mp: (A) LaClS, (B) dextrose] (C) LiCl, (D) CsI, (E) KNOI.
It may also be noted that the two types of solvation postulated above offer a reasonable explanation of the behavior of the alkali halides at elevated temperatures.e Osmotic coefficients of LiCl are lower a t 121” than a t 25” because of a decrease in solvation (water polarization). The coefficients of CsCl increase with temperature since ion pairing, which is enhanced by the water structure a t the lower temperature, decreases. Structure-Breaking Solutes. Positive solvation,” as evidenced by a decrease in the fraction of monomeric water in dilute solutions, was found for all of the salts except KN08 in the initial report. One other class of compounds for which activity coefficient data have been previously reported4 appeared to be a good candidate for further investigation. Activity coefficients of the aromatic sulfonates are in the usual sequence, Le., Li+ > Na+ > K+, etc., and those of the p-toluenesul(6) B. A. Soldano and 6. S . Patterson, J . Chem. SOC.,937 (1962). Volume ‘78, Number 7‘
July 1968
0. D. BONNER
2514 fonates are of the same order of magnitude as the nitrates. With progressive substitution of the aromatic ring by methyl groups, the activity coefficients become quite small. This would indicate a type of “negative solvation” if one believes that the solvation of KNOa is nearly zero. This phenomenon was first observed by Sugden7 from experiments in which acetic acid was distributed between salt solutions and amyl alcohol. The explanation which was proposed was a postulated depolymerizing effect on the water structure. This phenomenon is again capable of experimental verification. Pure water at 25” has been found to have an absorption maximum at 977 mp. Elevation of the temperature causes a shift of the maximum toward shorter wavelengths and it asymptotically approaches 958 mp, the characteristic band for liquid monomeric water. Solutions of salts which polarize the water molecule or solutes which cause an increase in the hydrogen bonding in water by enhancement of the structure cause a shift toward the limiting value of 986 mp. The data presented in Table I1 reveal that the band -
Table 11: Absorbance Maxima of Water and Certain Saturated Aqueous Solutions Position of the 2 n 4us
band,
Solution
mfi
Magnesium chloride Water Sodium mesitylenesulfonate Propylene carbonate
986 977 974 973
shift for aromatic sulfonates is in the direction opposite to that of the alkali metal halides. A more striking contrast of the behavior of the aromatic sulfonates and the alkali halides is shown in Figure 3. Solutions of sodium mesitylenesulfonate and potassium chloride are compared with water a t 25” in the double-beam instrument with water in the reference beam. The water us. aromatic sulfonate has a broad maximum a t 990 mp, indicating that the pure water is hydrogen bonded to a greater extent than the water in the solution. The hump at about 960 mp is due to the solvation of the sodium ion, with this band atop the shoulder of the broad band of hydrogen-bonded water. The water us. KC1 spectrum, on the other hand, has only the expected sharp peak at 958 mp. The spectroscopic evidence thus indicates that the lower activity coefficients of the aromatic sulfonates are related to their “structurebreaking” effect on the solvent and are not primarily due to ion pairing. While it is not possible to propose a definite “model” for this solute-solvent interaction, one may note that the sulfonate group is conjugated with the aromatic ring and that this permits the unit negative charge to be spread over a large area. The The Journal of Physical Chemistry
9w
920
I
I
I
910
960
980
1000
Yare Lcopth (4’)
Figure 3. Relative absorbance of solutions: (A) propylene carbonate, (B) sodium mesitylenesulfonate, (C) potassium chloride.
positive end of the water dipoles may be attracted sufficiently to “depolymerize” the water without substantial polarization of the molecules. The charge distribution on the large aromatic anions is quite different from that of the tetraalkylammonium cations which enhance the water structure in that the unit positive charge resides on the nitrogen atom of the RdN+ ion. Nonelectrolytes. Solvation numbers of certain nonelectrolytes were also reported in the previous article. Sucrose and dextrose were found to be highly solvated and it was noted that their solutions exhibit negative deviations from Raouit’s law. An equation relating the solvation of these solutes and their osmotic coefficients has been proposed by Stokes and Robinson.8 No explanation other than“ association” has been given, however, for nonelectrolytes, such as sym-dimethylurea,g in which the aqueous solutions exhibit positive deviations from Raoult’s law. It may be noted in Figure 3 that a nonelectrolyte, propylene carbonate, has a depolymerizing effect on the water structure. The spectra for solutions of ethylene carbonate and sym-dimethylurea are also quite similar and are, therefore, not included. It would be anticipated that osmotic and activity coefficients of aqueous solutions of ethylene carbonate would be substantially below unity, and preliminary results substantiate this expectation. Spectroscopic evidence of structure breaking by nonelectrolytes seems, in all cases, to be supported by colligative property measurements which indicate a positive deviation of the solutions from Raoult’s law.
(7) J. N. Sugden, J . Chem. SOC.,174 (1926). (8) R. H. Stokes and R. A. Robinson, J . Phys. Chem., 70, 2126 (1966). (9) 0.D.Bonner and W. H. Breazeale, J . Chem. Eng. Data, 10,326 (1965).
NMRSTUDY OF PROTON EXCHANGE
2515
Conclusions Spectroscopic evidence presented above indicates two types of positive solvation by electrolytes. (1) Small or highly charged ions, such as lithium and fluoride, are very effective in orienting the water molecules in their vicinity and tend to ion pair only with other ions having a high charge density. (2) Larger ions, such as cesium and tetraalkylammonium, tend to enhance the water structure around them; ie., the water is hydrogen bonded to a greater
extent. These ions tend to form ion pairs, especially with other large ions having a lower charge density. A type of negative solvation is exhibited by aromatic sulfonate anions which may be associated with a structure-breaking effect on the solvent. Nonelectrolytes are found to be both positively and negatively solvated, and, in all observed instances, the spectroscopic solvation evidence may be directly correlated with the direction of the deviation of their aqueous solutions from Raoult's law.
Nuclear Magnetic Resonance Study of Proton Exchange Involving Methyl-Substituted Pyridinium Salts in Methanol by Michael Cocivera Bell Telephone Laboratories, Inc., Murray Hill, New Jersey
07974
(Received January 6,1968)
The nuclear magnetic resonance line-broadening technique has been used to study proton-exchange reactions involving methyl-substituted pyridinium salts and methanol. For a reversible acid-dissociation reaction, ka BH+ MeOH F* B MeOHz+,changing the position of the methyl group from the para position to the k ortho position does not alter the value of either k , or k-,. In addition, the value of k-. does not depend upon the degree of methyl substitution, and this value is comparable to the value expected for diffusion-controlled kz reactions. The reaction BHf MeOH B -+ B MeOH BH+ was found to involve only one methanol molecule. For the monomethyl compounds, the value of kz does not depend upon the position of the methyl group. For the dimethyl compounds, the value of kz depends upon the position of the methyl groups and indicates that the rate for this reaction is retarded when both methyl groups are ortho to the nitrogen of the pyridine ring.
+
--&
+
+
+
+
+
Introduction Nuclear magnetic resonance studies of proton exchange involving substituted ammonium salts and such solvents as water and methanol have revealed considerable information concerning the mechanisms by which the exchange can occur.' However, little information has been obtained concerning steric hindrance in these reactions. The only available evidence that steric effects can be significant in these reactions comes from a study of proton exchange involving triethylammonium ion in water.2 I n that study, Ralph and Grunwald concluded that the ethyl groups were sterically hindering one of the exchange reactions. I n the present article, the results of a study of the exchange involving methyl-substituted pyridinium salts in methanol are presented. The salts which were studied were the hydrochlorides of 2-picoline (I), 4-picoline (11), 2,4-lutidine (111),and 2,6-lutidine (IV).
f;"3
I
31
y 3
m
llTL
These salts were chosen because of the possibility that the methyl group could sterically hinder one or more of the exchange reactions when it is ortho to the nitrogen of the pyridine ring. The advantage of using these salts is that the pK.4 of the salt does not change significantly when the methyl group is moved from the ortho to the para position on the pyridine ring. Consequently, the interpretation of the results is not (1) For a brief review, see E. Grunwald and M. Cooivera, Discussions Faraday Sac., 39, 105 (1965). (2) E. K. Ralph, 111, and E . Grunwald, J . Amer. Chem. Soc., 89, 2963 (1967).
Volume 7f2, Number 7 July 1968
