NON-COULOMBIC INTERACTIONS I N SOLUTIONS O F ELECTROLYTES' CHARLES A. KRAUS Y e l c a l j Research Laboratory, Brown University, Providence, Rhoda Island Receiced Ocfober 4, 10S8 INTRODUCTION
The present accepted theory of electrolytic solutions, as developed by Debye and Huckel, Onsager, and Bjerrum, accounts remarkably well for the properties of dilute electrolytic solutions on the basis of Coulombic interaction between the ions. The extension of this theory beyond the range of ordinary conditions has occupied the attention of physicists and physical chemists to such a n extent that sight has been lost of many other interactions which, in certain cases, play a significant r61e in determining the properties of solutions of electrolytes. I n this connection it is hardly necessary to recall the weaker acids and bases whose strength in aqueous solution is dependent upon the constitution of the substances in question. There are even salts which are weak electrolytes in aqueous solution and to account for whose properties it will doubtless be found necessary to take into consideration interactions other than the purely Coulombic. It is not the purpose of this paper to enter into a discussion of aqueous solutions of acids, bases, and salts, for there is already an extensive literature relating to this subject. It is in non-aqueous solutions that certain interactions of the non-Coulombic type may be studied with particular ease under conditions that permit of simple interpretation of the observed results. It is proposed to discuss several distinct classes of substances: first, acids and bases; second, weak salts; third, salts involving protonic interaction; and fourth, salts having ions containing dipoles. ACIDS AND BASES
In an aqueous solution the controlling factor that determines the strength of acids and bases as electrolytes is an active proton, due either to the dissolved substance or to the solvent molecules, which leads to interaction between the molecules of solute and of solvent. Water functions both as 1 Presented at the Symposiumon IntermolecularAction, held at Brown University, Providence, m o d e Island, December 27-29, 1938, under the auspices of the Division of Physical and Inorganic Chemistry of the American Chemical Society. 23 1
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acid and as base, as a result of which solutions of acids and bases in water differ from solutions of the same substances in most non-aqueous media. When a base, such as pyridine, is dissolved in a medium which has no active proton, there is no interaction and the solution shows no electrolytic properties. When, on the other hand, it is dissolved in a medium which has active protons, interaction occurs and the solution exhibits electrolytic properties. An acid may exhibit electrolytic properties whether the solvent in which it is dissolved has or has not an affinity for a proton of the acid. When the affinity of the solvent molecules for the proton is high, stable positive ions are formed and the solution of the acid exhibits the properties of a normal electrolyte, its strength being primarily determined by the size and configuration of the ions. This, for example, is the case for solutions of acids, in liquid ammonia, where the ammonium ion is very stable and there is no essential difference between the properties of ammonium salts and of corresponding salts of other ions such as the ions of the alkali metals. When an acid is dissolved in a medium which has no affinity for the proton, the properties of the resulting solution are largely determined by quantum forces acting between the proton and the negative constituent of the acid. With a few exceptions such solutions are extremely weak electrolytes, even when the dielectric constant of the solvent medium is high. Thus, hydrogen chloride dissolved in nitrobenzene, of dielectric constant 35, is a very weak electrolyte, while perchloric acid dissolved in the same solvent is a rather strong electrolyte. Obviously the energy necessary to separate the proton from the perchlorate ion, due to quantum forces, is much lower than the corresponding energy in the case of hydrogen chloride. As yet there are too few reliable data relating to the properties of solutions of acids in solvents having no affinity for the proton to enable us to obtain a clear picture of the underlying mechanism of their dissociation into ions. WEAK SALTS
Salts of strongly electropositive ions with negative ions behave like typical electrolytes in all solvents, regardless of whether the negative ion is strongly or weakly negative. The dissociation constant depends upon governing parameters, such as dielectric constant and temperature, and upon the energy necessary to separate the ions from their ion pairs under the action of Coulomb forces. Thus, in liquid ammonia sodium triphenylgermanide, or even sodium triphenylmethide, is a much stronger electrolyte than is sodium chloride, the reason for this being that the triphenylgermanide ion is much larger than the chloride ion and the energy necessary to separate the two ions due to Coulomb forces is correspondingly lower. I n the case of salts of weakly positive constituents with strongly negative constituents the situation is a different one. These electrolytes are weaker
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233
the less electropositive the positive constituent. The weakly positive ions, in general, have a valence greater than unity, and this change in valence complicates a study of the relationship between the strength of the electrolyte and the electropositiveness of thc positive constituent. This difficulty may be overcome by satisfying all but one of the valciices of the positive constituent by organic groups, leaving only one valence that may have electropolar properties. Thus we have such substances as methylmercury chloride, trimethyltin chloride, triphenylmethyl chloride, and the like These substances, as a rule, are soluble in water, a$ well as in other solvents, and form electrolytic solutions, although they are always more or less hydrolyzed in aqueous solutiolis and sometimes the degree of hydrolysis is too high to permit of their study. All thesc substances may conveniently be studied in solution in non-aqueous solvent.;. Triniethyltin chloride is only a moderately strong electrolyte in ethyl alcohol yolution, having a Other substances of this type in dissociation constant of 0.35 X alcohol, or other suitable solvents, are distinctly weaker electrolytes than are normal salts, indicating that quantum forces are involved. When substances of the type of trimethyltin chloride or triphenylmethyl chloride are dissolved in a solvent medium, the appearance of electrolytic properties depends upon interaction hetwecn the positivc constituent of the compound and thc solvent molecules. As stated, trimethyltin chloride is a fairly good electrolyte when dissolved in ethyl alcohol (dielectric constant 26), but it is a very weak electrolyte when dissolved in nitrobenzene (dielectric constant 35). So, also, triphenylmethyl chloride is a good electrolyte in liquid ammonia, of dielectric constant 22, but a yery poor electrolyte in nitrobenzene. The interaction which occurs is illustrated by the following equation :
The reaction is closely analogous to the corresponding reaction between triphenylboron and ammonia, which takes place as follows :
The only difference between the two is that in one case the complex formed is positively charged and in the other case it is neutral. A compound such as triphenylmethyl chloride, if dissolved in an inactive solvent of high dielectric constant, may exhibit electrolytic properties if the energy necessary to separate the two constituents is not too high. This energy is dependent upon the quantum forces involved. When this energy is small, as it is in the case of certain groups, electrolytic properties result. Thus, while triphenylmethyl chloride is a very poor electrolyte in nitrobenzene solution, triphenylmethyl Perchlorate is a rather good electrolyte. In this
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respect the triphenylmethyl compounds, as well as other similar compounds, resemble the corresponding acids which were mentioned above. SALTS INVOLVING PROTON INTERACTION
Any negative ion tends to interact with an active proton attached to a positive ion. The extent to which such interaction may influence the properties of the resulting solution will depend upon the nature of the solvent molecules themselves. When a partially substituted ammonium salt is dissolved in a solvent medium the molecules of which have no a&ity for the proton, interaction takes place between the negative ion and the proton of the positive ion. In other words, there is formed what is commonly known as a hydrogen bond, the energy of rupture of which may be much higher than the energy necessary to overcome the Coulombic forces acting between the ions. When, on the other hand, the same substance is dissolved in a solvent which has a marked a f b i t y for the proton, such as ammonia, the electrolyte behaves normally and in solution will be as strong as any normal electrolyte having ions of the same size. In table'l are given dissociation constants for a number of ammonium salts dissolved in nitrobenzene, ethylene chloride, and pyridine, respectively, together with the constants for several quaternary salts in the same solvent. Considering, first, substituted ammonium picrates in nitrobenzene, it will be noted that the dissociation constants for monobutyl- and dibutylammonium picrates have practically the same value, 1.5 X lo+. The dkociation constant of ammoniumpicrate in the same solvent has not been precisely determined, but it cannot M e r from this value by more than a few per cent, On substituting the third hydrogen of the ammonium ion by a butyl group, the dissociation constant is raised to 1.9 X lo4, indicating that the energy due to Coulombic interaction is appreciable although small. That the Coulombic forces affect the energy appreciably is shown, moreover, by the fact that the dissociation constant of trimethylammonium picrate is 1.46 x 104, which is about 25 per cent lower than the dissociation constant of tributylammonium picrate. The dissociation constant of tetrabutylammonium picrate is greater than 0.2; the substitution of the last hydrogen of the ammonium group, therefore, raises the constant approximately one thousand times. The dissociation constant of butylammonium perchlorate is 25.2 X lo+, as against 1.50 X 10-4 for the corresponding picrate. Here we see the influence of the markedly weaker a i h i t y of the perchlorate ion for the proton as compared with that of the picrate ion. I n ethylene chloride the influence of the proton closely parallels that found in benzene. Tetramethylammonium picrate has a dissociation constant of 0.326 x 10-4 in ethylene chloride, while pyridonium picrate
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has a dissociation constant of approximately 4 X lo-*. Again, the dissociation constant of a quarternary salt is approximately one thousand times greater than that of the partially substituted ammonium salt. It is interesting to note that pyridonium perchlorate has a dissociation constant of 50 X (approx.), which is in accord with the results found for perchlorates and picrates in nitrobenzene. All the ammonium type of salts when dissolved in pyridine have a nearly norma.1 dissociation constant. Thus, pyridonium nitrate has'a dissociation constant of 0.51 X lo+, while phenylpyridonium picrate has a dissociation constant of 11.5 X lo4. The nitrates are generally somewhat weaker than the picrates because of the smaller size of the nitrate ion, but TABLE 1 Dissociation constants of electrolytes i n different solvents* BALT
LIOLVENT
Ammonium picrate. . . . . . . . . . . . . . . . . . . . Butylammonium picrate. . . . . . . . . . . . . . Dibutylammonium picrate. . . . . . . . . . . . Tributylammonium picrate. . . . . . . . . . . . Trimethylammonium picrate. . . . . . . . . . Tetrabutylammonium picrate. . . . . . . . . Butylammonium perchlorate. . . . . . . . . . Pyridonium picrate, . . . . . . . . . . . . . . . . . . Pyridonium perchlorate, . . . . . . . . . . . . . . Tetramethylammonium picrate.. . . . . . . ! Pyridonium nitrate. . . . . . . . . . . . . . . . . . . Piperidonium n i t r a t e . . . . . . . . . . . . . . . . . Phenylpyridonium picrate, . . . . . . . . . . . . Ammonium picrate, . . . . . . . . . . . . . . . . . . .
Nit,robenzene Nitrobenzene Nitrobenzene Nitrobenzene Nitrobenzene Nitrobenzene Nitrobenzene Ethylene chloride Ethylene chloride Ethylene chloride Pyridine Pyridine Pyridine Pyridine
K
1.5 X 10-4 (approx.) 1.50 x lo-' 1.51 X lo-' 1.90 x 10-4 1.46 X 10-4
>0.2 15.2 x 10-4 4 X 10-1 (approx.) 50 X 10-8 (approx.) 0.3213x 10-4 0.509 x 10-4 0.184 X lo-' 11.5 x 10-4 2.87 X lo-'
* The values presented in tables 1 and 2 are due to Messrs. J. B. Ramsey, E. G. Taylor, D. J. Mead, L. M. Tucker, and C. R. Witschonke. Details will be published in due course.
the value obtained for pyridonium nitrate may indicate a slight interaction of the proton with the nitrate ion. Piperidonium nitrate, quite surprisingly, has a dissociation constant of 0.184 X lo4, indicating an interaction due to the proton. Apparently a proton interaction is influenced somewhat by the saturation or unsaturation of the groups attached to nitrogen. The dissociation constant of ammonium picrate is 2.87 X lo4, indicating that ammonium picrate is a normal type of electrolyte when dissolved in pyridine. Taking into account the difference in the behavior of ammonium salts in pyridine and in nitrobenzene or ethylene chloride, it seems clear that the low value of the dissociation constants in the latter solvents is due to
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interaction between the negative ion and the proton of the positive ion. The absence of such interaction in the case of pyridine solutions is doubtless due to the fact that, because of the relatively high affinity of the pyridine molecule for the proton, the interaction of the proton of the positive ion takes place with a pyridine molecule rather t h m with a negative ion. ,
SALTS OF POSITIVE IONS CONTAINING DIPOLES
Heretofore we have had very little information as to the possible influence of substituents other than alkyl or aryl groups in positive ions, All salts are so highly dissociated in water that aqueous solutions do not lend themselves to a study of the relation between constitution and dissociation constant. In solvents of lower dielectric constant, the dissociation constants of salts may be readily determined and the influence of various substituents may be investigated. Choline, hydroxyethyltrimethylammonium hydroxide, is a base in which one hydrogen of the ethyl group of ethyltrimethylammonium hydroxide has been substituted by a hydroxyl group. It gives rise to a series of salts which are kndwn as choline salts. Other atoms or groups of atoms may be similarly substituted. Thus the halogens (chlorine, bromine, and iodine) may be substituted in place of the hydroxyl group in the case of the normal choline salts. Many similar substitutions may be made. In table 2 are given dissociation constants for a number of salts of the choline type, dissolved in nitrobenzene, ethylene chloride, and pyridine. Considering, first, solutions of various choline salts in nitrobenzene, i t will be observed that there is a marked decrease in the dissociation constant due to the hydroxyl group. Thus, ethyltrimethylammonium picrate has a dissociation constant of 575 X while hydroxyethyltrimethylammonium picrate has a dissociation constant of 67 X lo4. The corresponding methyl derivative is markedly stronger, having a dissociation constant of 199 x 10-4. Methoxytrimethylammonium picrate and methoxymethyltrimethylammonium picrate have dissociation constants of 269 X lo+ and 262 X lo4, respectively, indicating a slight, and *nearly identical, effec7 due to the introduction of the polar groups C&o- and CH,OGH2-. In ethylene chloride the relationships are similar to those just described in nitrobenzene. The dissociation constant of ethyltrimethylammonium picrate is 0.46 X lo4, while that of choline picrate is 0.066 X lo-'. In this solvent the hydroxymethyltrimethylammonium picrate, having a constant of 0.088 X lo4, is again stronger than the corresponding ethyl derivative. There seems to be a rather close parallelism between the dissociation constants in nitrobenzene and in ethylene chloride. It is interesting, also, to note the dissociation constants of some of the
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halogen-substituted cholines, such as bromoethylcholine picrate, with a dissociation constant of 0.132 X IO"', and bromomethylcholine picrate, whose constant is 0.078 X Other cholines have been measured with results similar to those given above. When dissolved in pyridine the cholines seem to have normal strength. For example, ethyltrimethylammonium picrate has a dissociation constant of 8.2 X lo4, while hydroxyethyltrimethylammonium picrate has a dissociation constant of 9.5 X The hydroxymethyl compound, having a dissociation constant of 5.5 X IO+, is distinctly weaker than the corresponding ethyl derivative. Both bromoethyltrimethylammonium picrate and bromomethyltrimethylammonium picrate are distinctly weaker than choline picrate. TABLE 2 Dissociation constants of some choline salts LIALT
1b.n.ens Nitro-
Tetrabutylammonium picrate.. ...................... >0.2* Hydroxyethyltrimethylammonium picrate . . . . . . . . . . . . 67.0 Hydroxymethyltrimethylammonium picrate. . . . . . . . . . 199 .O Phenylhydroxydimethylammonium picrate. . . . . . . . . . . 0.177 Hydroxytrimethylarnmonium picrate.. . . . . . . . . . . . . . . . 0.173 Methoxytrimethylammonium picrate.. . . . . . . . . . . . . . . .269.0 Methoxyrnethyltrimethylammonium picrate.. ........ 262 .O Bromoethyl trimethylammonium picrate. . . . . . . . . . . . . . Bromomethyltrimethylammonium picrate. ........... Ethyltrimethylammonium picrate. . . . . . . . . . . . . . . . . . . . 575.0
R X 104 Ethylene chloride
Pyridine
2.18 0.0659 0.0879
12.2 9.49 5.55 12.27
0.254 0.152 0.078
0.460
6.85 4.79 8.21
At first sight, one would be inclined to ascribe the lower dissociation constant of the choline type of salts to the presence of dipoles in the positive ion. These might well affect the energy of separation of the ion pairs to an appreciable extent. On the other hand, it remains to account for the fact that in the case of pyridine the cholines have normal strength. This would rather seem to indicate that a n interaction which occurs in nitrobenzene and in ethylene chloride does not occur in pyridine. I n this respect the choline salts resemble the partially substituted ammonium salts in their behavior, alt>houghto a much lesser degree. Very striking are the values found for phenylhydroxydimethylammoniurn picrate and hydroxytrimethylammonium picrate in nitrobenzene, the constants having values of 0.177 X and 0.173 X 10-4, respectively. The dissociation constants of these salts are about one-tenth that
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of the corresponding partially substituted ammonium salts. In pyridine phenylhydroxydimethylammonium picrate has a dissociation constant of 12.27;if anything, it is a rather strong salt in that solvent. This indicates that low values of the dissociation constants of these electrolytes in nitrobenzene are probably due to interaction of the negative ions with the proton of the hydroxyl group attached to nitrogen. When this group is attached to a carbon atom which, in turn, is attached to nitrogen, the dissociation constant is much higher, as in the case of hydroxymethyltrimethylammonium picrate, for which K = 199 X 104. It seems, therefore, that the two last named compounds should be classed with the ammonium type of compounds, in that their interaction appears to be due to the presence of an active proton located in the positive ion.
