INVESTIGATIONS UPON ABXORMAL ELECTROLYTIC DISSOCIATION BY A. N. SAKHANOV
1. Investigations upon Conduetivity Jnvestigations of the present day have shown to us that in many non-aqueous solutions the molecular conductivity of electrolytes decreases with increasing dilution. This phenomenon] which is known under the name of abnormal dissociation, obeys the following empirical laws : l I . The equation] connecting the molecular conductivity with dilution, is expressed in the following way: Xun = const., where X stands for the molecular conductivity, u for dilution and .tz is a constant. 2 . Abnormal dissociation is a property common to all the solvents with low dielectric constants. 3. The decrease of the molecular conductivity with increasing dilution takes place sooner for a given electrolyte, the smaller the dielectric constant of the solvent. 4. At a definite dilution the molecular conductivity reaches a minimum and then increases normally with further dilution. 5. A maximum of molecular conductivity is found in highly concentrated solutions. The purpose of our investigation is not only to treat the question of abnormal electrolytic dissociation in solvents with high dielectric constants] but also to explain the causes of that phenomenon. In order to test the first statement we have used silver nitrate as an electrolyte.2 The following substances have been taken as solvents : aniline, quinoline] pyridine, metachloraniline, acetonitrile] water, and also some mixtures of the above-mentioned substances have been used. The dielectric constants of the mixtures were estimated from For more details see my article in Zeit. Elektrochemie, 2 0 , 529 (1914). Messrs. Prscheborovsky and Rabinovitch.
* These investigations were carried out with
A: N . Sakhanov
170
Boutyl and Silberstein’s2equations, the average values being taken. Thus we have obtained the following thoroughly studied solvents :
1
I
Solvents
Aniline Mixture: 4 vol. of Quinoline Mixture: I vol. of Pyridine Meta-chloraniline Mixture: 2 vol. of , Acetonitrile Mixture: I vol. of Water
aniline
+
aniline(+ pyridine
I
vol. of pyridine
I
vol. of pyridine
+
acetonitrile
I
vol. of acetonitrile
+
I
vol. of water
Diel. const.
6.85 8 .o 8.9 9.7 12.56 13.35 19.7 36.1
59.7 81.7
Special care has been taken as to the purity of the substances. Aniline and quinoline were obtained from Kahlbaum (mark “Kahlbaum”). After being dried thoroughly by means of potassium hydroxide, they were distilled. Pyridine was also obtained from pyridine “Kahlbaum,” after being boiled for a long time with a-reflux condenser in the presence of potassium hydroxide. Acetonitrile was prepared by distillation of Kahlbaum’s product in the presence of phosphorus pentoxide. The conductivity of silver nitrate in aniline and quinoline had been determined b e f ~ r e . ~ My data for aniline agree closely with Pearce’s4 recent data but those for quinoline are somewhat different from Pearce’s. The character of the curve of the molecular conductivity is the same, however. On the contrary, Pearce’s conductivity data for silver nitrate in pyridine not only differ greatly from mine, but also show differences in the very character of the curve itself. Dutoit anti Friderichj have investigated dilute solutions of silver nitrate
3
Comptes rendus, 114, 1421 (1892). W e d . Ann., 56, 661 (1895). Jour. Russ. Chem. SOC., 42,683 (1910);Zeit. phys. Chem., 83,129 (1913). Jour. Phys. Chem., 19,1 2 (1915). Bull. SOC.chim. Paris, (3) 19,321 (1899).
Iwestigations u p o n Abnormal Electrolytic Dissociation I 7 I in acetonitrile and their results coincide with ours. The difference between our results for silver nitrate in pyridine and those of Pearce’s is due perhaps to the difficulty which arises in obtaining absolutely identical samples. Similar and even still greater deviations often take place in the case of boiling-point determinations, viscosity measurements, etc. As may be seen from the tables, we have also studied highly concentrated solutions. The viscosity of such solutions differs greatly from that of the pure solvent and of dilute solutions. I n order to compare the conductivities of concentrated solutions with each other, as well as with those of dilute solutions, it is necessary to correct for this and to express the results in terms of the viscosity of the pure solvent. A number of investigators1 have found that the ionvelocity is either proportional to the viscosity of the medium or the viscosity raised to a certain power. The latter must approximate unity. Therefore, the values of conductivities for solutions with higher concentrations are corrected after either of the following equations :
L
= X.
17
-or L
= X.
170
where I, denotes the corrected molecular conductivity, X the directly measured molecular conductivity, 17 and v0 are viscosity coefficients of the solution and the pure solvent, respectively. We have corrected the values of the molecular conductivities after the simplest of the equations, the reason for this being that in the second equation the exponents are approximating unity and because the minimum of corrected molecular conductivity is obtained a t the comparatively low value of the ratio q / v 0 where both equations give similar I,-values. The values for the corrected molecular conductivity from the two equations differ greatly from each other only in the case of the highly concentrated solutions, where the value of the 1 Kohlrausch: Proc. Roy. SOC.,71, 338 (1903); Walden: Zeit. phys. Chem., 55, 207 (1906);Johnston: Jour. Am. Chem. SOC., 31, I O I O (1909);Washburn:
Ibid., 33, 1469 (1911).
A. N .
172
Sakhanov
ratio q f t q o is too great. As the problem of the present investigation is t o find minima of molecular conductivity, therefore, the real values of the corrected molecular conductivity for highly concentrated solutions are of little importance for us. The following Table I contains the results of our measurements (in recipr. ohm) at 2 5 ' * 0.05'.
TABLE I Conductivity and Viscosity of Solutions
Dilution in lit. V
186.6 112.9 52.4 32.6 26.38 18.35 10.41 6.76 3.24 I .560 0.911 0.570 0.543
.
Mol. conduct.
x
0.37 0.33 0.32 (Minim.) 0.34 0.36 0.39 0.64 0.85 1.54 I .96 (Maxim.) 1.57 0.76 0.62
Ratio'
Corr. mol. cond.
9 :90
L
.oo .oo
I I
I .02
.03 .04 I .06 I
I
I . IO I . 19 I .46
253.4 61.3 47.9 27.3 15.6 7.52 3.58 1.712 0.774 1
x 4.39 3.77 3.72 (Minim.) 3.81 4.~. I3 4.89 5.53 (Maxim.) 4.96 2.47
I .OI
2.25 4.57
2.33
-
-
-
-
IT. Silver nitrate in mixture: 4 vol. of aniline V
0.37 0.33 0.33 (Minim.) 0.35 0.37 0.41 0.70
9
:?lo
I
.oo
I .02 I .02 I
.04
I .07 I . 16 I .40
2.05 5.3
+
I
vol. of pyridine L 4.39 3.85 3.79 (Minim.) 3.96 4.42 5.67 7.35 10.2
13.0
For absolute data of viscosity see Jour. Russ. Chem. SOC.,47, 849 (1915).
Imestigations upon Abnormal Electrolytic Dissociation
I 73
111. Silver nitrate in quinoline. w
x
92.65 35.73 13.j 1 7.17 4.78. 3.48 2 .404
3.05 2.73 2.61 2.62 (bending) 2.56 2.37 2.08
IV. Silver nitrate in mixture:
I
.41 .64 2.08 I I
vol. of aniline
x 220.4 62. I O 35.85 16.09 8.37 3.83 1.736 0.804 0.431
3.08 2.81 (Minim.) 2.90 3.17 3.61 3.89 4.33
I .OF I .03 1.11 I .24
14.86 11.76 11.06 10.79 I O . 92 (bending) 10.94 9.26 5.15 1.365
I
+
?l : ?lo
I .oo I .OI I .03 I .06 I . 13
1.34 I .92 4.38 20.8
I
vol. of pyridine 4
14.9 11.9 11.4 I I , 4 (Minim.) 12.3 14.7 17.8 22.6 28.4
V. Silver nitrate in pyridine
x
V
175.7 89.26 57.10 23 . j 6 15.46 6.37 3.98 2.04 I ,250 0.945 0.572
41.6 36.6 33.8 30.0 28.7 26. 7 25.2 22.3 18.2 15.6 8.9
?l
:?lo
I .oo I .oo I .oo I .02 I .os I . 14
.22 1.44 I .92 2.48 4.96 I
1
41.6 36.6 33.8 30.6 30. I (Minim.) 30.4 30.7 32.1 34.9 38.7 44.1
A. N . Sakhavlov
I74 '
VI. Silver nitrate in rneta-chloraniline
x I
1
L
9:90
I
176.8 75.26 27.54 11.80 4.42 I .660
.007
1.57
I
I , I2
I .02
0.805 0.694 0.683 0.605
I
I .58 I . 14
.05
0.84 0.78 (Minim.) 0.95 1.59
I . I2
I .38 2.63
VII. Silver nitrate in mixture:
2 vol. of pyridine nitrile
+
I
vol. of aceto-
V ~
263.8 18.02 13.30 5.55 2.27 I .866
0.964 0 * 473
~
I .oo I .02 I .03 I . IO I .29
85. I 61.2 57.6 48.3 36.7 33.2 24.0 10.3
1.43 2.19 5.16
85.1 62.5 59.4 53.1 47.3 (Minim.) 47.5 50.2 53 . O
VIII. Silver nitrate in acetonitrile
x
V
61.8 31.9 8.70 3.73 I .630 I ,076 0.593 0,359 0.294 0 . I94
113.0 93.6 61.25 44.20 31.80 26.00 19.30 13.70 I1
.so
6.56
9 : 90 I .o I .o I . IO I . 17 I .40 I .60
2.3 4.6 6.2
18.0'
L
113.0 93.6 67.4 51.7 44.5 41 . 6 (Minim.) 44.4 63.0 71 .o 118.0
Investigations u p o n Abnormal Electrolytic Dissociation I 75 IX. Silver nitrate in mixture:
46.93 13.25 7.70 3,511 1.644 0.883 0.547 0.370 0.203 0.1149
98.2 89.7 84.7 76.9 65.6 53 1 42.3 32.7 20.5 12.36
0.7128 0,3587 0.2061 0.1182
71.27 56.88 43.84 20.52
'
I
vol. of acetonitrile
I
I
.o .o
I .02
.07 1.23 I .48 I .87 2.45 3.93 7 21 I
'
1.11
.27 .63 2.74 I I
+
I
vol. of water
98.2 89.7 85 .o 82.4 80. 7 78.5 (Minim.) 78.9 80.2 80.6 89.6
79.11 72.24 71.46 (Minim.) 80.88
The results obtained show how important for the right characterization of electrolytic dissociation are the corrections for the change of viscosity. It is only in solvents with dielectric constants not over 8.0 that the curve of molecular conductivity for silver nitrate reveals rather sharply the phenomenon of abnormal dissociation with the formation of a minimum and maximum (Solvents I and 11). I n solvents with dielectric constants between 8.5-10.0 the curve of molecular conductivity (A) forms only bendings (Solvents I1 and IV). And finally in solvents with dielectric constants over I O the molecular conductivity increases continually with the dilution. If, however, the conductivity values be corrected for the changed viscosity, the curve of corrected molecular conductivity for silver nitrate shows distinctly marked minima in all solvents. T h u s in solvents with dielectric constants over I O the continual increase o j molecular conductivity (A) with dilution depends u p o n the great increase o j viscosity o j the solutions with increasing concentration. The position of the minima of corrected molecular con-
A . N . Sakhanov
176
ductivity depends strictly upon the dielectric constant of the solvent; that is, with the increase of dielectric constant m i n i m a are gradually displaced to the region of more and wore concentrated solutions. For this reason in the case of aniline the minimum occurs a t a comparatively high dilution. I n solvents with still lower dielectric constants, such as amylamine, or benzylamine, minima of molecular conductance for silver nitrate were not noticed, the reason being, of course, because they lie in higher dilutions, where the conductivity of the solutions is too slight. However, in the case of such solvents as water, minima can be detected only with difficulty for just the opposite reason, because they lie in the region of highly concentrated solutions. Waldenl showed that there exists a certain relation between the position of the minimum of molecular conductivity (dilution = V,) and the dielectric constant (D) of the solvent : -
DV34V, = const.
Nevertheless he proved that this equation can be applied only to solvents with small dielectric constants, since the values of the molecular conductivity were not corrected for the changed viscosity. The following Table I1 contains the results obtained concerning the position of minima of corrected molecular conductivity for silver nitrate. As m a y be seen f r o m the table, W a l d e n ' s equation can be applied perfectly to solvents with dielectric constants up to 35. It is rather important to note that the constancy of the product D 3 d V , occurs only when the corrected molecular conductivities ( L ) are used. For -solvents with dielectric constants over 35 the product (D "V,) increases. The reason for this phenomenon may be that there are real deviations from Walden's equation in the case of solvents with rather high dielectric constants. However, we must not forget that the correction of conductivities made by means of the above-mentioned equation was only ap1
Bull. Akad. Sci. Petersb., 1913,1083.
Investigations upon Abnormal Electrolytic Dissociation I 77
TABLEI1 3iel. const.
Solvent
Aniline Mixture: 4 vol. of aniline of pyridine
+
Quinoline
+
I
Dilution
I 3.\ic D
6.85
60 Lit
29.5
8.0 8.9
50 Lit 45 Lit
29.5 31.7
vol.
Mixture: I vol. of aniline I vol. of pyridine 35 Lit 9.7 31.7 12.56 Pyridine 31 .o I j Lit Mixture: 2 vol. of pyridine I vol. of acetonitrile 28.4 3 Lit 19.7 1 2 Lit Metachloraniline 30.6 13.35 I Lit Acetonitrile 35.8 35.8 Mixture: I vol. of acetonitrile I 0 . 7 Lit vol. of water 59.7 53.0 Water 81.7 0 . 2 Lit 47.4 proximative, especially for highly concentrated solutions, where minima are found for such solvents as water. The conductivity values corrected by means of the equation are too high for concentrated solutions and the real minima also must lie in the region of more concentrated solutions. Therefore, the values of V, for the last two solvents, IX and X, are smaller than those given in the table and the application of the Walden equation to the last two solvents is not excluded. A perfectly regular displacement of minima of corrected molecular conductivity with the increase of the dielectric constant of the solvent forces us to draw a conclusion that a t some concentration or other for each of the above-tested solvents there is found the phenomenon of abnormal dissociation, upon which depends both the formation of minima and the further increase of molecular conductivity with the concentration. This conclusion holds true for the solvents with dielectric constants up to 36. Nevertheless for the last two solvents, IX and X, extrapolation is necessary as in these cases the abnormal dissociation can be found only in the region of highly concentrated solutions. Bowden’sl investigation of this ques-
+
+
Jour. Chem. SOC., QQ, 194 (1911);McBain and Taylor: Zeit. phys. Chern., 76, 179 ( 1 9 1 1 ) ;Bunbury and Martin: Jour. Chem. SOC.,105, 417 (1914).
178
A . N . Sakhanov
tion as well as investigations of others have not only removed all doubts concerning this question, but have also proved that some electrolytes in aqueous solutions form minima of molecular conductivity a t comparatively low concentrations and without any corrections for changed viscosity. Thus we draw the following very important conclusions: The corrected molecular conductivity ( L ) of silver nitrate j o r m s m i n i m a in all the solvents studied with dielectric constants between 6-82 and the position of these m i n i m a i s determined by the dielectric constaiat. T h e abnormal dissociation decreases with the increase of the dielectric constaiat of solvent. Therefore, oidy jor solvents with low dielectric constants i s the abnormal dissociation a very characteristic property. T h e phenomenon of electrolytic dissociation i s determined by the action of two factors. Ion formation proceeds by means of normal as well as abnormal dissociation. The corrected molecular conductivity, having passed through the minimum, increases continually with the concentration. This is shown clearly by measurements in which we have reached, at least in some cases, exceedingly high concentration of salt. This phenomenon in its turn leads to the following very important conclusion : the curve o j the corrected molecular conductivities intersects the axis of the ordinates (the dilution being = 0 ) at a certain positive value of the ordinate. I n the Solvents I and I1 the directly measured molecular conductivity (A) forms distinctly marked maxima. I n the Solvents I11 and IV bendings are formed, i. e., faintly marked and brought together maxima and minima. In Solvents V-X, with still higher dielectric constants there are already neither maxima nor bendings. Thus we arrive a t the conclusion that m a x i m a o j molecular conductivity (A) are formed in solvents with dielectric constants up to I O . All these conclusions are drawn as the result of the investigation of one electrolyte, silver nitrate, but unquestionably they can be applied as well to other “normal” electrolytes of the silver nitrate type : salts of ammonium, tetraethylam-
Investigations u p o n Abnormal Electrolytic Dissociatiovl
I 79
monium, etc. At the same time they hold true only for so-called “normal” electrolytes. 2. Investigations upon the Transport-Numbers
Notwithstanding the fact that the most important laws, governing the abnormal electrolytic dissociation, are fixed experimentally, up to the present day we have not any universally accepted hypothesis or theory concerning the nature of this phenomenon. I n 1905-1906, Steele, McIntosh and Archibald suggested a hypothesis, the essential point of which was that they assumed that in some solvents only the associated molecules of salt are capable of electrolytic dissociation. I n I 9 I 2-1 9 I 3, I developed this hypothesis further. The essential point of this hypothesis that associated molecules are more capable of electrolytic dissociation than simpler ones, I have explained from the point of view of Abegg and Bodlander’s theory, according to which the electro-affinity of complex ions is greater than that of simple ones. Anyhow, the principal point of the hypothesis of conducting current complexes is in the firmly fixed phenomenon of polymerization of electrolytes in solvents with small dielectric constants. Likewise the increase of polymerization is always accompanied by the increase of molecular conductivity with concentration : determinations of molecular weights in solvents with low dielectric constants show that in these solvents the salts are associated and that the association increases with the concentration. Therefore, we see the close relation between the two phenomena of abnormal dissociation and of polymerization of the salt. If an electrolyte is polymerized, the result of its electrolytic dissociation is, as is known, the complex ion. Therefore, the complex ion formation is looked upon as being natural for solvents with low dielectric constants. Kraus and Bray found that the transport numbers for different salts in liquid ammonia contradict the hypothesis of current conducting complexes. As the transport-numbers for most salts in this solvent change but little with dilution,
A. N . Sakhavzou .
I80
.
therefore, the above-mentioned authors conclude that complex ions are not formed in liquid ammonia. Bringing this conclusion to a general form, the authors consider the complex ion formation as being an exceptional phenomenon. Anyhow, the absence of change or, more properly speaking, the slight change of transport numbers for many electrolytes in liquid ammonia must be noticed. The fact is that Franklin and Kraus’s ebullioscopic measurements in liquid ammonia show that electrolytes in this solvent are polymerized. Since the complex ion formation comes as a result of the electrolytic dissociation of associated molecules, therefore, the presence of complex ions in liquid ammonia is more than probable. If, notwithstanding this, the transport-numbers do not give any hint as to the presence of complex ions, there is, therefore, a contradiction between the ebullioscopic data and those of the transport-numbers. The second example of similar contradiction we find in Serkov’s researches. These investigations, carried out after the method of transport-numbers, show that the complexity of lithium chloride in acetone is greater than that of lithium iodide, in spite of the fact that according to numerous cryoscopic and ebullioscopic measurements iodide salts reveal a greater tendency to form complexes than bromide and especially chloride salts. Thus we perceive here also a difference between the results of the two methods. These contradictions as well as the chance to come to the question of the cause of abnormal dissociation were the reasons for the following investigation. For this purpose we have measured2 the transport-number for silver nitrate in aniline and two mixtures of aniline with pyridine (I1 and IV). On the other hand, Schlundt3has measured the transportnumber for the same salt in pyridine and acetonitrile. If to these data be joined Hittorf’s data for silver nitrate in aqueous solutions, we shall obtain representation of the change of 1 2
a
Zeit. phys. Chem., 73, 500 (1910). These investigations were carried out by Mr. Grinbaum. Jour. Phys. Chem., 6 , 159 (1902).
Investigatiows upow Abnorwaal Electrolytic Dissociation I 8 I transport-numbers for silver nitrate, depending both upon the concentration of solution as well as upon the dielectric constant of the solvent. The apparatus, in which the electrolysis took place, was constructed like that of Nernst and Loeb. The anode consisted of a bent silver rod and the cathode of a silver plate. Electrolysis lasted for 3-5 hours. The current strength was from 4 to 8 milliamperes. After the electrolysis was finished the liquid from the apparatus was poured off into weighed bottles through the tap, which was in the lower part of the long tube, where was the anode. Therefore, the first bottle contained the anode layer. Afterwards in the same manner were taken away the two small middle layers. The difference between the whole weight of the liquid in the apparatus and the sum of the anode layer and two middle layers represented the weight of the cathode layer. The content of the coulometer, as well as the solutions of the layers, were titrated according to Volhard’s method with nearly N / 4 0 solution of ammonium rhodanate. The titre of that solution was determined before each experiment. When we have titrated the cathode and anode layers, we have taken only a part of it in order not to add too great a quantity of ammonium rhodanate solution. By adding a sufficient amount of dilute nitric acid we have dissolved the aniline and pyridine of the solutions in the state of nitrate salts. After that the solution thus obtained was diluted with water and was titrated according to Volhard in the presence of ferrous alum. Different experiments proved that such titration is possible in the presence of aniline nitrate, in case the titration is made sufficiently rapidly. On standing, the solution gradually becomes green because of the reaction between the ferrous salt and aniline nitrate. This very slight greenish or bluish tint, which appears in the diluted solutions a few minutes after the addition of alum, interferes with the titration. The intensity of the color increases with the time. Therefore, we were obliged to make two titrations for each
I 82
A. N . Sakhanov
solution and in the second titration alum was added a t the very end of the titration. It must be stated that, after some practice, the first titration, when alum is added a t the beginning, was accomplished so quickly that it coincided entirely with the second. The difference between the two titrations rarely reached 0.05 cc. The middle layers show that the concentration of silver nitrate does not change during the electrolysis in all the experiments included in Table 111. Only in two cases have we found that the concentration of the middle layers had changed, evidently in consequence of their being mixed either with the anode or with the cathode solutions. For this reason these two experiments were omitted as being evidently faulty. The whole quantity of silver nitrate in all the layers, being compared with the quantity before electrolysis took place, was found to be unchanged within the limit of * I .5 mg. As we have determined both the increase of silver at the anode and the decrease of it a t the cathode, therefore, in every experiment we have obtained two values for the transportnumber: at the anode and at the cathode. The maximum difference between these two values is, as may be seen, 3 percent. The result is quite satisfactory, if we take into consideration that even in aqueous solutions the difference is the same. The character of the silver deposited on the cathode depends upon the concentration of the solution. In highly concentrated solutions the silver is deposited in the form of a very dense layer. In more dilute solutions the silver is deposited in the form of dendrites or of a black powder sornetimes very friable and easily broken off from the cathode. The solutions of silver nitrate at the anode and cathode did not change color during the electrolysis. Only in the very dilute solutions did the cathode solution grow quite yellow during the electrolysis. Evidently the pulverulent silver accelerates catalytically the oxidation of the aniline. When the yellow cathode solutions were treated in dilute nitric acid,
Investigations u p o n Abnormal Electrolytic Dissociation 183 tinted solutions were obtained which were titrated with some difficulty. Therefore] in the case of dilute solutions the transport-number was determined only with the anode solution. The experiments were carried out a t the room temperature, say, a t 18". The results of the experiments are given in Table 111. TABI,E I11
1 I
Time of electrolysis
Current strength
i 4 4 4 4
hrs. 30 min. hrs. I O min. hrs. hrs. 30 min.
5 hrs. 5 hrs. 5 hrs. 3 hrs.
1 I
Dilution in l&rQ liters
I Ca 8 8 8 6
I I
Transport-number I
a t anode
1 I .40 4.0 10.0
25.2
.60
8 8
2.10
7 6
10.35 19.6
I
a t cathode
I
0'342 0.353 0.372 0.403
0.326 0,336 0.352 0.395
I
1
0,335 0.363 0,374 -
1 0.329 0.337 -
A . N . Sakhanov
184
port-numbers which we obtained are expressed as average of the measurements for the anode and cathode solutions.
TABLE IV The transport-nurhbers of silver for silver nitrate (multiplied by Solvent dilution:
1.0
1.5
2
4
5
IO
100)
20
---50.0 - 48.3 47.2 - - 38.3 - - 42.2 - 44.8 --_.
Water Acetonitrile Pyridine Mixture: I vol. of pyridine I vol. of aniline Mixture: I vol. of pyridine 4 vol. of aniline Aniline
+
+
32.6 - 34.2 - - 39.0 -
33.8 - 35.6 - 37.3 -
-
35.239.: - 32.733.6 37.3 36.6 - 34.4 - -
It is seen from Table IV that a decrease of transportnumber with increasing concentration takes place in acetonitrile, pyridine and two mixtures. In water and aniline we observe the reverse phenomenon. All the solvents enumerated in the table (with the exception of aniline for investigated concentration of the salt) demonstrate well the general law : the transport-numbers j o r the give?z electrolyte in dilute solutions do not depeizd upon the nuture OJ the solvent. In fact, the more dilute the solutions, the more closely the transport-numbers of silver approach the value of 0.47. The considerable and regular change of the transportnumbers with dilution was explained by Hittorf and there is hardly anything to be said against this ordinary interpretation. Carrara,I Abegg and Neustadt,2 Schlundt3 and others likewise admit that the transport-numbers change because of the presence of complex ions-be it cation or anion according to the sign of these changes. Elektrochemie nichtwasseriger Losungen, 25 (1908). Zeit. phys. Chem., 69, 486 (1910). 3-Jour. Phys. Chem., 6 , 159 (1902). 1 2
Investigations u p o n Abnormal Electyolytic Dissociation I 85 The considerable and regular decrease of the transportnumber of silver with increasing concentration in acetonitrile, pyridine, and two mixtures is due to the fact that the negative electricity is being conducted not only by the ions N03’, but also by the complex ions evidently of the composition -Ag(IK03)2’. The concentration of the complex ions increases with concentration of the solution according to the mass law. It may be seen from Table IV that the normal value of the transport-number of silver (for silver nitrate), 0.47, is obtained a t different dilutions, which depend on the dielectric constant of the solvent. This result is a consequence of the theory of complex ions, since the dielectric constant determines the dilution, a t which the electrolyte is already to a great extent changed into normal molecules. This explanation of the change of the transport-number is in perfect harmony with the measurements of the molecular weights of silver nitrate in pyridine and aniline. Walden and Centnerschwer, and Walden2show that silver nitrate is polymerized in these solvents and that the degree of polymerization increases with the concentration. If silver nitrate is polymerized in both pyridine and aniline it is doubtlessly polymerized in the mixtures of these two solvents. Quite unexpectedly the transport-number of silver (for silver nitrate) in aniline (as far as the tested concentrations are concerned) changes exactly as in water; that is, increases with the concentration. Moreover the values of the transport-numbers (in tested concentrations) are greater in aniline than in the mixtures. Since silver nitrate is polymerized in aniline, the increase of the transport-number of silver with concentration must be explained by the formation of a complex cation -Ag2N03’. From the formation of complex anions in the mixtures of pyridine with aniline we may also decide that such anions are present also in aniline. Drawing a converse conclusion (from the formation of complex cations in &it. phys. Chem., 55, 321 (1906). Bull. Acad. Sci. Petersb., 1913, 1083.
186
A . N . Sakhanov
aniline) we see that the complex cations Ag2NO3' are also formed in the mixtures of pyridine with aniline. Thus as a general case electrolytic dissociatian of associated molecules gives com9lex anions as well as cations. For example, the electrolytic dissociation of silver nitrate in any of the solvents leads as a general case to the following processes: AgN03 (AgN03)z (AgN03)z
Ag'
+ NO3'
e Ag' + &(NO&'
AgzN03' 4- Nos'
These phenomena may be complicated on the one hand by solvation and on the other hand by the formation of still more highly polymerized molecules than double, as is the case with solutions in chloroform. Probaljly the simultaneous formation of complex anions and cations is a frequent phenomenon in non-aqueous solutions. I n that case this phenomenon throws light not only on the measurements described in this paper, but also on those contradictions between the data of transport-numbers and the determinations of molecular weights, which we have been speaking of a t the beginning of the chapter. Evidently when the formation of only one complex ion (anion or cation) takes place, the transport-numbers will change abruptly with dilution. Simultaneous formation of both complex ions makes this abrupt change of transport-numbers less marked. If both complex ions are formed in the same quantity, the transport-numbers will be independent of the dilution. Thus in this case the phenomenon will follow exactly as in the case when the formation of complex ions does not occur. In such conditions the method of transport-numbers does not enable us to detect the existence of complex ions. Therefore, when deciding the question of complex ions, the method of transport-numbers must by all means be combined with determinations of molecular weights by cryoscopic or ebullioscopic method. The slight change of transport-number with dilution in liquid ammonia in connection with determinations of molecular weights can be explained by formation of complex anions and
Investigatiolzs upon Abnormal Electrolytic Dissociation 187 cations: Likewise the greater change of transport-number of lithium for lithium chloride in acetone, than for lithium iodide can be explained by the fact that lithium iodide, being more polymerized, dissociates so that both ions are in the complex form. Thus our investigations carried out with the transportnumber are in complete harmony with the hypothesis under consideration. 3. Investigations upon the Electromotive Force of Concen-
tration Cells Because the question of abnormal dissociation is considered to be of great importance, electrometric measurements of concentration cells were made, for example:
+
Ag
1
1
0.1 N
0.001
AT
I
-
Ag
We are giving in the present paper the data only for solutions in pyridine. All the measurements were carried out a t a temperature of 25.0 =t 0.05'. The method of measuring is that of Poggendorff and the results are given in the following Table V. TABLEV. Electromotive force of concentration cells in pyridine Solution of AgNOa (conc.1) 2 N I .5 I .o N 1.0
n:
N I.ON
0.2
"1
0.1
N
0 . I 11: 0.01
AT
0.001 N
Solution of AgNOs (conc.?)
Electromotive force
N
0.054 Volt 0.042 0.035
0.2
0.15 N
A:
0.1 0.01
N
0.001 AT 0.01 0.01 N 0.001 N
N
0,001
N N
0.0001
0.070
0.114 0.050
0.038 0.080 0.041 0,053
As may be seen from the table, in very dilute solutions in pyridine the electromotive force of concentration cells
A . N . Sakhanov
I88
approaches the normal value. For example, the electromotive force of the cell Ag I 0.001 N AgN03 I 0.0001 N AgN03 Ag in pyridine is equal to 0.053 volt, whereas that for the corresponding cell in aqueous solution is 0.055 volt. This fact proves that Nernst’s theory can be applied to solutions in pyridine. On this basis we calculate the ion concentration and the degree of dissociation of silver nitrate in pyridine after Nernst’s simplest equation, neglecting the diffusion potentials : C1 1 E = RT -log - = 0.059 log c F Cz cz I n these calculations we consider the dissociation of silver nitrate in 0.0001 N solution as being complete, which, of course, is not quite correct. The results of these estimates are given in Table VI. TABLEVI I
AgNOa Conc. of solution 0.0001 0.001 0.01 0. I 0.2 I .o
1.5 2 .o
Conc. of Ag-ion 0.0001
0.00079 0.0039 0.018 0.028 0.066 0.123 0.231
Percentage dissociation IO0
79 39 18 I4 6.6 8.2 11.6
From the electrometric measurements as is seen from the table we can also draw the conclusion that the degree of dissociation of silver nitrate in pyridine passes through a minimum in accordance with the results of measurements of conductivity. Since only corrected molecular conductivities (I,) give minimum values, therefore, the expedience of the suggested corrections for the change of viscosity is proved. There is a striking difference between the position of the minimum according to the data of conductivity and that as shown b y the electromotive forces. I n the first case we ob-
c
Investigatiom upon Abvlormal Electrolytic Dissociation I 89 tain a minimum a t a dilution of 15 liters and in the second a t a dilution of I liter. Such difference from the point of view of the theory of current conducting complexes could be foreseen, since the formation of complex cations of silver must inevitably produce this difference. This question will be particularly treated in our next investigation. LITERATURE CONCERNING THE ABNORMAL DISSOCIATION Steele, Macintosh and Archibald: Zeit. phys. Chem., 55, 129 (1906). Archibald: Jour. Am. Chem. SOC.,29, 665, 1416 (1907); 34, 584 (1912). Franklin andKraus: Ibid., 27,216 (1905); 29,1395 (1907). Franklin: Jour. Phys. Chem., 15,675 (1911). Lewis and Wheeler: Zeit. phys. Chem., 56, 190 (1906). Foote and Martin: Am. Chem. Jour., 41,451 (1909). Sakhanov: Jour. Russ. Chem. Soc., 42, 683 (1910); 42, 1363 (1910); 43, 526 (1911);44,324, 1794 (1912);Zeit. phys. Chem., 80, 13 (1912); 83, 129 (1913); Zeit. Elektrochemie, 20, 529 (1914). Sakhanov and Prscheborovsky: Zeit. Elektrochemie, 20, 39 (1914); Jour. Russ. Chem. SOC.,47, 879 (1915). Sakhanov and Rabinowitch, Jour. Russ. Chem. SOC.,47, 859 (1915). Sakhanov and Grinbaum: Ibid., 1769 (191j); second paper is printed. Plotnikov: Ber. deutsch. chem. Ges., 39, 1794 (1906); 42, 1154 (1909). Walden: Bull. Imper. Akad. Sci. Petersb., 1913,922, 987, 1100; 1915, 789. Hopfgartner: Wien. Akad. Ber., 120, IIC, I (1911); 122, IIC, 603 (1913). Fitzgerald: Jour. Phys. Chem., 16, 621 (1912). Pearce. Ibid., 19, 14 (1915). Anderson: Ibid., 19 (1915). Fischler: Zeit. Elektrochemie, 19, 126 (1913). Beckmann: Zeit. anorg. Chem., 77, 280 (1912). Beckmann and Waentig: I b i d , 67, 36 (1910). Kraus and Bray: Jour. Am. Chem. SOC., 35, 1315 (1912). Laboratory of Inorganic and Physical Chemistry
Imperial University Odessa, Russia
