636
(6) (7)
(8) (9) (10) (11) (12)
C. W. CARR, W. F. JOHKSOK .4KD I. M. KOLTHOFF nomena, and their Dependence on Colloid Stability; Paper So. 27 of the Symposium on the Stability of Colloidal Dispersions, which was held under the auspices of the Division of Colloid Chemistry a t the 110th RIeeting of the American Chemical Society, Chicago, Illinois, September, 1916. LANGMUIR, I.: J. Chem. Phys. 6,873 (1938). STAMBERGER, P. : Flocculation of Lyophobic Dispersions by Slow Mechanical Stirring; Paper S o . 33 of the Symposium on the Stability of Colloidal Dispersions, which was held under the auspices of the Division of Colloid Chemistry a t the 110th 3Ieeting of the American Chemical Society, Chicago, Illinois, September, 1946. VERWEY, E. J. W . : Chem. Xeekblad 39, 563 (1912). VERWEY, E . J. W.: Contribution to a symposium held by the Xederlandsche Chemische Vereniging on July 3-4, 1944. VERWEY, E. J. W . : Philips Research Reports 1.33 (1945). VERWEY, E. J. W., A K D OVERBEEK, J. TH.G . : Trans. Faraday Soc., in press. VERWEY,E. J. W., ABD OVERBEEK, J. TH. G., with the collaboration of K . van Xes: Theory of the Stability of Lj/ophobic Colloids, in press. Elsevier Publishing Company, Amsterdam, Holland.
T H E USE OF MEMBRASE ELECTRODES I S THE STUDY OF SOAP SOLUTIONS' C. W. CARR,
W.F. JOHSSOS, . ~ N DI. 11,KOLTHOFF
School of Chemistry, Institute of Technology, C-niversity of Minnesota, IMinneapolis, Minnesota Received November 14, 1946 I. INTRODUCTION
In studies of the properties of soap solutions several different physicochemical methods have been applied. These methods include such diverse measurements as that of the electrical conductance, the degree of solubilization of an oil or dye, the freezing-point loivering, the viscosity, and the surface tension of the solutions. The present paper is concerned with the use of recently developed membrane electrodes (1, 2, 4, 8) for the determination of the cation and anion activity in soap solutions. The collodion membranes are negatively charged in electrolyte solutions and behave as electrodes for cations in much the same may as the glass electrode behaves toward hydrogen ions. Similarly, the protamine-collodion membranes, being positively charged, behave as anion electrodes. These membranes differ from the glass electrode in that they are not specific for any one ion. For that reason they cannot be used for determining directly the activity of one ion in the presence of another of the same sign. However, they have been applied with success to a number of solutions containing a single electrolyte (3). Since the membranes are not always perfectly selective, especially a t concen]This investigation was carried out under the sponsorship of the Office of Rubber Reserve, Reconstruction Finance Corporation, in connection with the synthetic rubber program of the United States Government.
1IEMBR.LUE ELECTRODES I S STCDT O F SOAP SOLCTIOXS
637
trations above 0.05 S,the simplest method to determine ionic activities is by means of a potentiometric titration. The solution of unknoirn concentration is placed on one side of the membrane and a known volume of xater on the other side. -1strong solution of k n o m concentration of an electrolyte u-hich has the ion TL-hose activity is to be determined in common with the electrolyte studied is added to the ivater from a buret. I t is not necessary that the ions of opposite charge in the unknoxn and linonn be the same. .ifter each addition of electrolyte the membrane potential is measured, and the addition of electrolyte is continued until the potential changes in sign. The point of zero potentisl indicates that the activity of the ion being measured is the same on both sides of the membrane. At the point of zero potential the activity of the ion under consideration is equal to that of the same ion in the outside solution. From the volume and concentration of strong electrolyte added to a given volume of water the concentration of this electrolyte a t the zero-point potential is calculated, and from this concentration the activity of the ion under investigation is estimated. EXPERIMESTAL PROCEDURE
The membranes used in this work were prepared as described by Sollner and coworkers (1, 2, 4, 8). They are bag-shaped, being formed over 25 x 100 mm. test tubes, and hold about 30 ml. of solution. A known volume of water is added to a beaker of such size that the beaker is about one-half to tn-o-thirds filled. In most of the experiments a 250-ml. beaker filled with 150 ml. of xater was used. Sext, the membrane is filled uith the solution to be investigated. About tvo-thirds of its length is immersed in the heaker containing the xater, and the membrane is then clamped in position. For a clamp a wooden test tube holder held t o a ring stand with a right-angle clamp is convenient. The temperature is measured and titration is started. The electrolyte solution used as titrant should be a t least ten times as concentrated as the solution being titrated, to avoid the addition of excessively large volumes needed to reach the end point. After each addition the E.>r.F. of the system is determined: outside inside Hg/Hg&lz(s) KCl(satd.) solution I membrane solution1 KCl(satd.) Hg,ClZ(s)1 Hg Ct cz
I
1
In our measurements a Leeds 6 Sorthrup Type K potentiometer was used. The saturated calomel electrodes are connected n ith the t u o solutions by means of specially constructed agar bridges saturated n i t h potassium chloride. These bridges are made with 3-mm. glass tubing, and the tips which make contact with the solutions are about 1 mm. in diameter. The small contact area is necessary to minimize the diffusion of potassium chloride into either solution. T o minimize this diffusion still further, the bridges are removed from the two solutions after each measurement of the potential and placed in saturated potassium chloride solution. Just before use they are wiped dry and placed in the solutions being measured.
638
C. W. CARR,
W. F. JOHSSON A S D
I. M. KOLTHOFF
As more and more electrolyte is added to the outside solution, the potential decreases until it reaches zero, changes sign, and begins to increase. The end point, which is the point of zero potential, is determined graphically. The membrane potential is plotted on the linear axis of semilogarithmic paper and the concentration of the outside solution is plotted on the logarithmic axis. The slope of the lines obtained with the negative membranes is close to 55 mv. a t 2 5 T . for a tenfold change in concentration 11-hen the membranes are perfectly cation-selective. Since this is not aln ays true, especially a t concentrations above 0.05 31, the slopes will tend to be slightly less than 55 mv. With the positive membranes the largest slope that has been obtained is 53 mv. For the titrations with either type of membrane it is not necessary that the maximum slope be obtained. I t is only necessary that it be constant in the concentration range near the end point. 31aterials used The soap solutions used mere in all cases prepared by the neutralization of the fatty acid with an equivalent amount of standard aqueous base. The capric, lauric, and myristic acids nere obtained from the Eastman Kodak Company. The oleic acid was a sample of 9 i per cent purity supplied to us by Dr. W. C. Ault of the Department of Agriculture, Eastern Regional Research Laboratory. The sodium rosinate \vas made from a sample of dehydrogenated rosin acid, obtained from the Hercules Poivder Company, which was mostly a mixture of dihydroabietic acid, tetrahydroabietic acid, and dehydroabietic acid. EXPERIMESTAL RESULTS
The first titration carried out was with a solution of sodium laurate of known Concentration. Freshly prepared 0.01 J1 sodium laurate was placed inside a negative membrane, and a known volume of vater Tyas placed outside, using 0.1 31 sodium laurate as the titrating solution. The results are plotted graphically in figure 1. The concentration of sodium laurate in the outside solution a t the end point is found to be 0.0101 -11,in good agreement with the known concentration (0.0100 X)of the inside solution. To determine whether or not the counter ion was of any influence in these dilute solutions, another titration was carried out using 0.02 AM laurate in the inside and 0.2 M sodium chloride instead of laurate as the titrating solution. The results are also shown in figure 1. In this titration the concentration of sodium chloride a t the end point vas found to be 0.0201 ;M. Thus the sodiumion activity in 0.0200 31 sodium laurate is the same, at least nithin 1 or 2 per cent, as in 0.0200 31 sodium chloride. .lfter these preliminary esperiments the cation activity of several soap solutions \vas determined. The solution in question was titrated with standard sodium or potassium chloride solution. h sample calculation of the cation activity coefficient is as follows: In the titration of 0.1 31 sodium laurate the concentration of sodium chloride a t the end point was 0.0512 X . The sodiumion activity in this solution is 0.0130. The activity coefficient of the sodium ion in 0.1 31 sodium laurate then is 0.043/0.1 = 0.43.
MEJIBRASE ELECTRODES IS STCDY OF SOAP SOLUTIONS
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7
-
z
J
2z
I l l 1 1
I
-
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639
-
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I
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FIG 1. Potentiometric titrations of sodium laurate solutions with the use of membrane electrodes. Curve I, 0 01 .If sodium laurate titrated with 0.1 .If sodium laurate; curve 11, 0.02 .If sodium laurate titrated with 0 2 '11 sodium chloride.
The curves in figure 2 shox that the cation activity in soap solutions below the critical point is the same as in solutions of strong mi-univalent electrolytes. When the critical point is reached, holyever, the cation activity begins t o drop below that of a strong uni-univalent electrolyte. This drop in activity indicates
640
C. W. CARR, W. F. JOHNSON AKD I. &KOLTHOFF I.
t$at as the soap concentration increases beyond the critical concentration alarge fraction of the sodium ions becomes inactivated with the formation of lamellar micelles and does not contribute to the sodium-ion activity measured in the soap solution. A rough estimate of the activity of the sodium ions in the dissolved soap micelles can be made as follows. The critical concentration of sodium laurate is approximately 0.02 JT. The activity coefficient of sodium ions in this solution corresponds to an activity of sodium ions of 0.018. In the concentration range between 0.02 M and 0.04 AI sodium laurate, the activity of the sodium ions increases from 0.018 to 0.027. Assuming that the un-micellized concentration of sodium laurate remains unaltered above the critical concentration, xve find that an increase of the concentration of the micellized soap from 0 (0.02 Jf sodium laurate) to 0.02 (0.04 ; I Isodium laurate) gives rise to an increase in the sodiumion activity of 0.009. This would correspond to an “average activity coefficient” of the sodium ions of the micellized soap of 0.45. Making similar calculations TABLE 1 The critical concentration in several soap solutions as determined by two different methods SOAP
! ~
Sodium caprate... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Sodium laurate.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Potassium myristate.. . . . . . . . . . . . . . . . . . . . . . . . . . . Sodium oleate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Sodium rosinate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
PEISUREYEICT OF xa- XCTLVlIT
I
::A: 0.003
