T H E DISSOCIATION O F STRONG ELECTROLYTES. III* Complete Dissociation and Optical Properties BY MORRIS B. JACOBS AND CECIL V. KING
The supposed additivity of color or light absorption, of refraction and of dispersion by the ions in solutions of strong electrolytes has in recent years been cited in favor of complete dissociation, just as thirty years ago similar data were used to support the Arrhenius dissociation theory. The Arrhenius theory indicates that some properties of strong electrolytes should be found which can be assigned to the undissociated molecule; and the difficulty of specifically assigning any properties to such a portion of the electrolyte has recently been interpreted in favor of its complete absence. color and Absorption.-As far back as 1892 Ostwald' called attention to the fact that a solution of any electrolyte should show a t least three absorption spectra-one for each ion and one for the undissociated molecule; if complexes were formed, the absorption spectrum would be even more complicated. Ostwald showed that thirteen different metal permanganates, a t a dilution of one equivalent in 500 liters, had absorption bands in identical positions in the spectrum. Ten salts of fluorescein showed the edge of their absorption band in the same place, and eosin and anilin violet salts were similar in this respect. Ostwald took this, of course, to mean that these salts were highly dissociated in these dilute solutions, and that the metal ions and anions were highly independent; for the changing or substitution of atoms or radicals in non-conducting organic molecules is usually attended by pronounced color changes. Ostwald did not attempt to show evidence of absorption by the undissociated molecules of these salts. Also in 1892 Magnanini showed that two solutions of copper sulfate of the same concentration, one containing considerable sulfuric acid, had exactly the same light absorption. He interpreted this as evidence against ionization, but Ostwald replied that undissociated CuSOa might well have the same absorption, in solution, as copper ion. Magnanini then (1893) proceeded to show that the addition of KNOI to potassium violurate did not change the absorption; he was under the impression that violuric acid and the violurate ion (if it existed) were not colored and that undissociated potassium violurate was the only colored substance present. Donnan2 describes the above experiments and shows that Magnanini's work was marred by the impurity of his materials. Donnan himself shows that, assuming that the color is due to the violurate ion only, one can calculate the percentage ionization and the dissocia-
* Contribution from the Department of Chemistry, Washington Square College, New York University. Ostwald: 2. phyeik. Chem., 9, 579 (1892). Donnan: 2. physik. Chem., 19, 465 (1896).
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tion constant of violuric acid from the color a t various dilutions, and obtain results identical with those from conductivity measurements. This is pointed out by Halbanl who shows that p / p m = €/E, = a for this acid. ( p = equivalent conductivity, E = extinction coefficient). Bjerrum2 found that the color, in solution, of certain chromium salts was practically constant on dilution, Le., the molecular extinction coefficient did not change with the concentration; he interpreted this to indicate the absence of undissociated molecules. I n 1918 Bjerruma summarizes the evidence for complete dissociation and points to his earlier conclusion from the color of these chromium salts. He points out that a and i-I (from conductivity and from colligative properties) do not agree and remarks that if these are abandoned as a measure of the degree of dissociation the “anomaly of strong electrolytes” disappears. He emphasizes that his conclusions from the color of salts apply only when n o complexes are formed. Koyes4 on the basis of similar observations concluded that since the optical activity, color, etc. were independent of the concentration (when referred to equivalent quantities), these properties were additive with respect to the ions even a t concentrations where much of the salt was un-ionized. At that time, in other words, Noyes interpreted constancy of color or absorption to mean that ionzzation i s a n optically indifferent process, and that the color is the same whether the ions are “bound” or “free”, a view which has been held for many years by Hantzsch.j Hantzsch, like Bjerrum, found that many strong electrolytes show practically no change in molecular extinction coefficient on dilution. This was true for the Sr, Li, Ka, Cs, T1 walt,sof acetyloxindon, for some copper salts and copper complexes in water, aqueous ammonia and other solvents, for permanganates, chloroplatinates, potassium diamminetetranitrocobalt, and for K and Ba salts of t,richloroacetic acid (in ultra-violet light). Further, he showed that trichloroacetic acid has exactly the same absorption in petroleum ether as in water, and that HBr is optically identical in water, alcohol and ether; also that many salts show the same constancy of molecular absorption in alcohol solution as in water. Recently, with more exact experimental work, he shows that the ultraviolet absorption by various halides is not strictly additive, and ascribes the discrepancies to minor chemical changes such as hydration. Remembering that substitution in un-ionized organic compounds has often a pronounced effect on the color, and that ionization should be less in some of these non-aqueous solvents than in wat,er, Hantzsch concludes that ionization is an optically indifferent process. Halban: 2. Elektrochemie, 34,489 (1928). Bjerrum: Det Kgl. Dansk. Vid. Sel. Skr., ( 7 ) 4, 26 (1906);2. anorg. Chem., 63, 140 (1909). 3 Bjerrum: 2. Elektrochemie, 24, 321 (1918). Noyes: Science, 20, j77 (1904). 5Hantzsch and co-workers: Ber., 41, 1216,4328 (1908);58,612 (192j); 59, 1096 (1926); 2. phvsik. Chem., 63,367 (1908);72,362 (1910); 84,321 (1913);86,624 (1913); 2. Elektrochemre, 29, 221, 434 (1923); 30, 194 (1924);31, I67 (1925). 2
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Hantzsch in this matter disagrees with Halban,’ who has also done much experimental work on absorption by strong and weak electrolytes. Halban at first agreed with Bjerrum’s view, and not Hantzsch’s; since Beer’s law is obeyed by strong electrolytes and not by weak ones, ionization cannot be an optically indifferent process. He even showed, as has been mentioned, that the percent ionization of a weak electrolyte might, be calculated from the extinction coefficients. Later, however, Halban came to the conclusion that Beer’s law is by no means generally valid even for strong electrolytes; he points out that the salts cited by these workers were mostly rather complex salts, while the relation is not obeyed by some of the simplest salts. For instance, while for the complex chromium salts E is practically constant with concentration while p is not, for the absorption in the ultraviolet by lithium nitrate e follows much the same type of curve as f i . Also, while according to Bjerrum’s idea addition of another salt should have no effect on the absorption by the first, actually enormous effects may be observed. Halban, while assuming practically cornplete dissociation in such cases, explains the invalidity of Beer’s law according to the electron-shell deformation theory of Fajam.* Halban and Ebert, in a careful piece of experimental work, show with exact measurements that Beer’s law is not strictly valid for several strong electrolytes and that the absorption of various salts is influenced great’lyby the addition of nomabsorbing salts. Halban and Fajans both admit t’he possibility of hydration or of equilibria which may include undissociated electrolyte or groups of ions of both signs in equivalent or non-equivalent amounts. Fajans3 admits the necessity of this since neither the molecular absorption or refraction is strictly additive for many strong electrolytes. Part of the discrepancy, he decides, is due to ion deformation; but this does not account for all the variations and the presence of molecules, ion groups or complex ions is indicated. In many cases color changes are apparently connected with changes in the degree of hydration of t’he salt’. The molecular absorption of cobalt chloride in water and various alcohols4 was found not to be constant with changing concentration, but to increase, decrease, or pass through a minimum as the dilution progresses. The gradual change in color in passing from a dilute to saturated aqueous solution and then to alcoholic solutions is perhaps best explained by a change in the degree of hydration of the salt. E. F. George5 found the absorption of light by certain electrolytes to change with concentration, to be presumably a function of their state of dissociation, and to be influenced by the addition of salts havinga commonion. Halban: Z. Elektrochemie, 29, 434 (1923);30, 601 (1924);34, 489 (1928);Halban and Ebert: Z. physik. Chem., 112, 321 (1924). * Fajans: Naturwissenschaften, 11, 165 (1923). Also other papers by Fajans and coworkers. 8 2. physik. Chem., A 137,361 (1928). ‘Hulbert, Hutchinson and Jones: J. Phys. Chem., 21, 150 (1917). 6 George: Disaertation, Ohio State (1920).
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Huttig and Keller‘ found that exact measurements showed the molecular absorption in the ultraviolet of all the halides of lithium to change with concentration. Houstoun and his coworkers2 studied the absorption of many salts in aqueous and non-aqueous solutions; mostly salts which show extensive hydrate and complex ion formation, as ferric, nickel, cobalt, and copper halides. With these salts no constancy of extinction coefficient with concentration is found; no additivity of absorption for anion and cation. Houstoun reaches the conclusion that solvation plays an important role in the color of ions; and that in dilute solutions where additivity is to be expected, the spectrophotometric methods in use at that time were inadequate to prove constancy of molecular absorption. Very often the factor of complex ion formation and its effect on color have been neglected. Bjerrum repeatedly emphasizes that absorption may be additive for the ions and one can expect this to mean complete dissociation only when no complexes are jormed. Considerable work has also been done to show a parallelism between color change and complex formation in certain cases.3 From the above can be seen the difficulties involved in using color or light absorption in support of the theory of complete dissociation. The possibility of complex ion formation even in dilute solution limits the application of this theory ~onsiderably.~For color and absorption data to be entirely satisfactory evidence of complete dissociation, Beer’s law should hold unqualifiedly for strong electrolytes, and color and absorption should be strictly additive for anion and cation; but these conditions are met only with good approximation in some cases and not a t all in others. Hence one can adopt any of several views : ( I ) The older view of Xoyes: that undissociated molecules of strong elect,rolytes are potentially ionized and dissociation of these ions is an optically indifferent process; The view of Hantzsch: that ionization itself is an optically indifferent (2) process, and change in color is connected, rather, with a change in chemical constitution; ( 3 ) The view of Bjerrum: that ionization is not an optically indifferent process and that validity of Beer’s law indicates complete dissociation (minor divergences being explained by Bjerrum’s ion-association theory or other factors) ; (4) The view of Halban and Fajans: that ionization is not an optically indifferent process, that even large departures from additivity or from Beer’s ‘Huttig and Keller: Z. Elektrochemie, 31, 390 (1925). Physik. Z., 14,424 (1913); Proc. Roy. SOC.Edinburgh, 33, 35, 44, 137, 147, 156 (1913). For example-Denham: Z. physik. Chem., 65, 641 (1909); Mecke and Ley: 111, 385 1924); Ley and Heidbrink: Z. anorg. allgem. Chem., 173, 287 (1928); Getman: J. Phys. hem., 26, 377 (1922). See also McBain and Rysselberge: J. Am. Chem. SOC.,50, 3009 (1928); 52, 2336 (1930). See, however, Freed and Kasper: J. Am. Chem. SOC., 52, 2632 (1930).
t:
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law can be explained by the ion-shell deformation theory, and that some small percent of the electrolyte may be in the form of groups of ions or even neutral molecules; or ( 5 ) The view that we wish to emphasize here, namely that the percent dissociation can no more be calculated with certainty from color or absorption data than from CY ( = p / p L , ) or G I . Refractive Index.-The above statements apply almost verbatim to the supposed additivity of refraction by the ions. The smallness of change in molecular refractivity with concentration has often been cited in support of complete dissociation. Thus Bjerrum’ commented on the work of Fajans. Fajans answered that if such small changes are real, they must be significant. U’hen the Arrhenius theory was first presented, the approximate additivity of refractive index was used as a support for the theory; now it is used as evidence of complete dissociation, and the discrepancies are explained by ion association, deformation, hydration, complex formation, etc. It will be seen that almost every experimenter has his own explanation for the small or great divergences from true additivity. Additivity of refractive index has been cited as evidence of complete dissociation by Lewis2 in his famous address “The Use and Abuse of the Ionic Theory”; by Lewis and Randall3; by Lange,4 who says the existence of appreciable numbers of molecules in solutions of st,rong electrolytes has not yet been provided by optical (or other) means; by Harnedj who says “The additivity of the refractive indices, rotatory powers . . . is well known”; etc. With refractivity we have the same difficulty in explaining the experimental data as with optical rotation and light absorption. Actually the molecular refraction is usually not additive for the solid electrolyte and the solvent; and sometimes the calculated molar refraction of the solute remains nearly constant with change of concentration, sometimes it does not. If it does remain nearly constant] we can interpret this fact in two ways: ( I ) the electrolyte is completely dissociated at all concentrations; or ( 2 ) the electrolyte has very nearly the same refractive power whether ionized or not. If the refraction does not remain constant, we can interpret this fact in two ways: ( I ) the electrolyte is completely dissociated, a.nd the change in refraction is caused by ion association, ion deformation, or some other factor; or ( 2 ) the ions and the molecules have widely different refractive powers. Walden6measured the molar refractivity of tetramethylammonium iodide, tetrapropylammonium iodide, and phenylethyldimethylammonium iodide in water and in a series of other solvents of high and low dielectric constants. He found only I or z percent change over a wide concentration range in most ’Trans. Faraday SOC., 23, 376 (1927). Chem., 70, Z I Z (1909). 3Lewis and Randall: J. Am. Chem. Soc., 43, I I I Z (1921); “Thermodynamics”, 318, 319 (1923). Lange: Physik. Z., 29, 760 (1928). Taylor’s “TreatLse of Physical Chemistry”, 2, j90 (1925). Walden: Z. physik. Chem., 59, 385 (1907).
* G. Tu’. Lewis: Z. physik.
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of the solvents; also a rather small difference in solvents of high and low dielectric constants. It is hard to see how these results can be interpreted in favor of complete dissociation in aqueous solution; if the salts are really unionized in the non-aqueous solvents of low dielectric constant, then the molecular form has nearly the same refractive power as the ionized form. H . C. Jones and his coworkers’ carried out many experimental studies on the refractive index of salt solutions. They found that, in general, the refractivity of the electrolyte was proportional to its concentration. In many cases the refractivity of the pure electrolyte could be predicted from the study of its solutions; they concluded that when this could not be done, the electrolyte formed hydrates in solution. Rimbach and Wintgen* used the refractive index to determine the formation of complexes in solution. They found the refraction to be additive for pairs of salts which form no complexes, as NaCl-KC1, AgN03-LiX03etc. Likewise the refraction was additive for the components of mixtures which deposit double salts, as alums, which led to the conclusion that these double salts are completely dissociated in dilute solution. Mixtures of salts which are known to form complexes, as CdCl*-KCN, AgX03-KCN, gave entirely different results, the refraction being not a t all additive. Thus, they say, measurements of refractive index can be used to give qualitative knowledge of complex formation; but the change in refractivTindex is too small for quantitative work (usually not over I % even for concentrated solutions). SchwersS shows that the ratio of change of density to change of refraction with change in concentration, for sulfuric acid, increases with the dilution. From a study of this coefficient he concludes that in many cases-HC1, HBr, HI, HnS04-the refractive index change, for the same specific volume change, is less in dilute than in concentrated solution; or in other words, that ionization (assuming the Arrhenius theory) lowers the refractive index (at constant density). Schwers points out that LeBlanc came to the opposite conclusion, i.e., that the ions have a greater refractive power than the molecules. Other authors reached the conclusion that the ions and molecules have the same refractive power; in other words, that ionization is an optically indifferent process. Ch6veneau4 made careful determinations of the refractive indices of solutions of KC1, KH&O, and hIg(sOs),. He found practical constancy of specific refraction of the electrolyte over a considerable concentration range, but in very dilute solutions there was a pronounced increase or decrease. In 1910 he indicated the possibility that this was due to lack of adequate temperature control; but in 1921 the experiments were repeated with all precautions for keeping the temperature constant and the previous results were completely verified. Chkveneau interpreted these results to indicate that Many papers in Am. Chem. J.; 2. Elektrochemie; Z. physik. Chem.; Carnegie Institution Publications; etc. * Z. physik. Chem., 74, 233 (1910). Z. physik. Chem., 75, 621 ( 1 9 1 1 ) . Compt. rendu, 150, 866 (1910); 172, 1408 ( 1 9 2 1 ) .
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ionization has no appreciable influence on the specific refractive power in solutions of concentration over gram per liter; the changes at higher dilution he attempts to explain on the basis of a change in vibration frequency of the constituents of the electrolyte, and compares the dissolved substance in dilute solution with an attenuated gas. It can be seen that the earlier work on refraction by electrolytes in solution by no means gave the conclusive proof of additivit,y that the exponents of complete dissociation would have one believe. The difficulty of exact enough measurements, the uncertainty of calculation of the molar refractive power of the dissolved electrolyte, the lack of knowledge concerning formation of hydrates and complex ions all contributed to the uncertainty regarding the interpretation of refractive index data. The most accurate experimental work done prior to 1920 could scarcely distinguish between complete or 95% ionization, even if other complicating factors were absent. This was the status of the problem until the work of Debye and Huckel stimulated further research. Fajans and his coworkers' have done much experimental work on refractive index with the express intention of learning the condition of strong electrolytes in aqueous solution. They point out that very few of the older measurements are sufficiently accurate to test rigorously additivity of refractive power of the ions; and show the difficulty of calculating the true refractivity of the solute in those cases where the refractivity of solvent and solute are not additive. The molar refractivity of sulfuric acid, Fajans says, is constant up to a concentration of 3 0 2 ; this is, however, by no means true for HC1 and many salts. For many salts the refraction extrapolated to infinite dilution is additive for the ions, but this is not true at finite dilution. The equivalent refraction of most salts shows a linear relationship with the concentration. Fajans concludes from refractometric and other optical data and from vapor pressures, that in solutions of strong electrolytes ion deformation plays an important role but that there are undoubtedly some neutral molecules or a t least pairs of ions and complexes perhaps of the type H2Cl+. He points out emphatically that we haye every gradation from the strongest to the weakest electrolytes. Strong electrolytes differ from weak electrolytes only in degree; the essential modification of the classical theory lies in taking account of the inter-ionic forces in any quantitative treatment of the subject. The refractive index of the halides of lithium was measured by Hiittig and Keller? it was found that the refractivity of the salts and the water was not additive, and the value for the salts varies with the dilution. Fajans remarked that the refraction of the solution was in all cases greater than the sum of that of the salt and the water, indicating greater ion deformation, the closer the salt atoms. Huttig ascribed the effect to hydration, whether in conjunction with ion deformation or not. Z. Physik, 23, I (1924); Ber., 59, 249 (1926); Trans. Faraday Soc., 23, 357, 375, 408 (192jj; 2. physik. Chem., 130, 724 (1927); 4 137, 361; B 1, 42; (1928); 2. Elektrochemie, 34, I (1928). 2 Z. Elektrochemie, 31, 390 (1925).
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Conclusions.-We can only conclude from the above that the “additivity of optical properties” is a myth; t’ruein some accidental cases, nearly enough true in others so that the discrepancies were not noticeable in the older measurements or could be overlooked. Closer examination shows that from the very nature of the properties under consideration, since the experimental variations are so small, the older, less exact work could be interpreted in favor of the Arrhenius theory as well as the complete dissociation theory. In either case the most exact recent work shows that there are discrepancies to be explained by some other means than the degree of ionization. The observations considered here, in connection with the other experimental evidence, undoubtedly indicate that the conductivity ratio gives values too small for the degree of ionization; but there is mounting evidence of the exist,ence of some small percent of undissociated molecules, or ion-pairs or groups, in even the best examples of highly dissociated electrolytes. Too many workers forget that the inter-ionic attraction theory does not require complete ionization. And it must be remembered that a t the high dilutions where there seems to be fair agreement with the Debye-Huckel equations, . even the Xrrhenius theory postulates nearly complete dissociation. S e w York, N. Y.
