ELECTROCHEMISTRY OF PERMSELECTIVE COLLODIOK MEMBRANES.
11
299
T H E ELECTROCHEMISTRY OF PERMSELECTIVE COLLODION MEMBRANES. I1
EXPERIMEKTAL STUDIES Oh' THE COSCENTRATION POTENTIAL ACROSS VARIOUS TYPESO F PERMSELECTIVE COLLODION MEMBRbNES WITH SOLETIONS OF SEVERAL ELECTROLYTES' KARL SOLLNER
AWD
HARRY P. GREGOR?
Department of Physiology, Cniosrsity of Minnesota, Minneapolis 14, Minnesota Received August 8 , 1946
I In the first paper (17) of this series a study was made of the rates a t which final stable concentration potentials with various electrolytes are established across several types of permselective collodion membranes. The present paper is an experimental study of the final concentration potentials established by various electrolytes a t different concentrations across several types of permselective collodion membranes. All measurements reported in this paper represent final and stable potential values. The concentration potential across a membrane is a function of the nature, the concentration ratio, and the absolute concentrations of the electrolyte solutions separated by that membrane. In a systematic investigation it is therefore necessary to measure the concentration potentials arising with a variety of electrolytes a t several concentration levels. With the exception of the cases in which ideal ionic membrane selectivity prevails, the latter is strictly defined only with concentration ratios which differ infinitesimally from unity. In experime,ntal work the concentration ratio of the two solutions, therefore, should be as small as practicable; following an established procedure (6, 7, 8, 9, 12, 16) the concentration ratio 1:2 was used in t,he present investigation.
I1 The significance and the physical meaning of experimental concentrat,ion potentials measured across ion-selectire membranes can be visualized only by reference to the theoretical limits within which the properties of real membranes are confined. With membranes of high ionic selectivity by far the more important of these limits is the (calculated) potential, E,,,, which would arise if the membrane behaved under a given set of conditions as an ideal membrane for the reversible transfer of the critical ion. The critical ion in the case of electronegative membranes such as collodion membranes is the cation. The non-critical ion (the anion in the present case) is assumed to contribute nothing to the transport of electricity across the membrane; the concentration potential is defined only by the activities of the critical ion at the two sides of the membrane. Presented a t the Twentieth S a t i o n a l Colloid Symposium, nhich was held a t Lradison, Wisconsin, May 28-29, 1946. * Present address: Polytechnic Institute of Brooklyn, Brooklyn 2, S e w York.
300
KARL SOLLNER AND HARRY P. GREGOR
The other limit is the potential which would arise across a membrane which does not show any ion selectivity of its own; this lower limit therefore is the liquid junction potential E l arising from free, unhindered diffusion in the absence of any membrane. This criterion is of particular interest where a low degree of ionic selectivity of the membrane prevails. The two limiting values of the concentration potential can be calculatedwith considerable accuracy from available data. The calculations of the theoretically possible m a x i m u m values of the concentration potential, E,., are based on the general equation
E,,,
RT z+F
= - In
a:"'
-
atJ
and ay' its activities c:".y:' and c y ' . y f ' in the two solutions. In the case of uni-univalent electrolytes, the mean activity coefficients were used (a+ = a*). With the uni-bivalent electrolyte potassium sulfate the activity coefficient for the critical ion, the potassium ion, is calculated assuming that the activity coefficient of a given ion is determined only by the total ionic strength of the solution, s (s = &rc,z:). Then y K + in a solution of potassium sulfate is the same as y K + in a potassium chloride solution having the same ionic strength. The individual ion-activity coefficient of the potassium ion, y K + , in solutions of potassium chloride is equal to the mean activity coefficient yKC1.3 The majority of the numerical values necessary for these calculations were taken preferentially from Harned and Owen (3), MacInnes ( 5 ) , and the International Critical Tables (4),intermediate data being interpolated graphically. The maximum concentration potential can be calculated with considerable accuracy. The probable error in the case of potassium chloride and lithium chloride is undoubtedly considerably smaller than the accuracy of our experimental determinations. Although the calculated values in the case of potassium sulfate and hydrochloric acid may contain some small systematic error, it seems safe to assume that it is insignificant compared with the accuracy and reproducibility of the experimental data which are reported below. The potential arising from digusion in free solution, that is, the liquid-junction potential Q / C I , was calculated from the equation z+ being the valency of the critical ion and a:"
where El(ca,o) is the liquid-junction potential, t+ and t- the transference numbers of cation and anion within the membrane, a:" and a?' the single ion activities of the cations, and al" and a?) the single ion activities of the anions in the solutions (1) and (2). a For example, a 0.01 M potassium sulfate solution, the ionic strength of which is 0.03, has the same yr+ a s a potassium chloride solution of ionic strength 0.03, the concentration of the latter solution being 0.03 M . T h e yr+ of a 0.03 M potassium chloride solution is 0.848; therefore the yr+ in aO.01 M potassium sulfate solution is assumed to be 0 848.
ELECTROCHEMISTRY OF PERMSELECTIVE COLLODIOS MEMBRAKES.
Since l+
For uni-univalent electrolyte we assumed a, = a- = a*. equation 2 becomes
RT 5 y ) ELcenlc, = (2t+ - 1) - In -F a:)
11
+ t-
301 = 1,
(3)
In the case of the uni-bivalent electrolyte potassium sulfate, it is necessary t o use equation 2 and the single ion activities defined and calculated respectively as described farther above. The numerical values of t- ivere taken preferentially from the before-mentioned reference tables.4 The accuracy of the computed El(ealr)values is probably about as high as the accuracy and reproducibility of most of our corresponding measurements. With electrolytes other than potassium chloride, systematic errors of at least 0.10 millivolt may occur. I11 The degree of the ionic selectivity of the permselective collodion membranes is so great and the accuracy and reproducibility of the measured concentration potentials is so high that the full value of the assembled experimental data cannot be utilized unless proper corrections are made for the asymmetry of the two liquid-junction potentials potassium chloride-agar bridge electrolyte c1 and electrolyte c2 potassium chloride-agar bridge. The customary neglect of these corrections is not justified in the case of the permselective membra ne^.^ The best way to eliminate this difficulty ivould consist in the outright experimental elimination lvherever possible of the two doubtful liquid-junction potentials by the use of specific reversible electrodes, such as the Ag IAgC! or the Hgl Hg2S04 electrodes. (In cases where such specific electrodes are available their use seems indicated in further ivork.) For the present investigation we had recourse to the following method ivhich, although not completely free from objections, seems to be sufficiently satisfactory for the purpose on hand, no corrections being attempted in the case of potassium chloride. The electromotive forces El(,,,, ivere measured which arise in diffusion chains without membranes which were otherwise identical with those used for the measurement of the concentration potentials across membranes. These experiwere compcred with the mental values of the liquid-junction potential El(exp) corresponding calculated values of the liquid-junction pukntials Ei(ea~o), and the difference between these two values is taken as the correction.
1
!
, 4 Equations 2 and 3 assume t h a t the transfer numbers are independent of concentration. This assumption, a s can be seen from the last column of table 1, is not quite correct. Therefore in calculating the liquid-junction potential, values of 1, were used which correspond to the mean concentration of each concentration ratio. We also feel t h a t the introduction of this correction in many of the published d a t a on membrane potentials with electrolytes other t h a n potassium chloride would bring these d a t a more closely in line with one another and also into better agreement with theoretical conciderations.
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KARL SOLLNER AND HARRY P . QREGOR
If an experimental value was, e.g., lower than the calculated one, the difference had to be added to the measured concentration potential across a membrane in order to arrive a t the corrected, true value of the latter., All data referring t o the concentration potentials E across membranes which are given in table 1 are corrected in this manner.
IV The technique used for the measurement of the concentration potential e across the membranes was the.same as that described in the preceding paper: saturated calomel electrode [ potassium chloride saturated saturated potassium chloride-agar bridge electrolyte cz membrane electrolyte c1 saturated potassium chloride-agar bridge potassium chloride saturated saturated calomel electrode. In all instances the final stable concentration potentials e were determined with all the precautions discussed in the preceding paper. As in the case of the rate experiments (17), the membranes prior to the measurements had been aged by about 3 days' immersion in 0.1 Ai potassium chloride solutions. Occasional repetitions of previously made determinations which were interspersed between later experiments showed that the membranes did not undergo any detectable change in the course of the various consecutive experimental series. The electrolytes used were potassium chloride, lithium chloride, potassium sulfate, and hydrochloric acid. The investigated concentration ranged from 0.002 N/O.OOl N to 0.4 ,V/0.2 N. In addition to the three types of membranes Ox 8 - Hum 43, Ox 12 - Hum 43 and Ox 10 - Hum 43 - Alc 65 for which the time studies on the concentration potential have been reported (17), membrane Ox 14 - Hum 58 has also been used in the present work. All measurements reported below in table 1 for one type of membranes were performed successively with a single membrane, except in one case (Ox 14 Hum 58) where the original membrane broke accidentally and had to be replaced by a second specimen which, as shown by repeat experiments, had within the limits of experimental error the same properties as the original one. The reproducibility of the measurements was about =k 0.10 millivolt, except in the case of the highest concentrations of hydrochloric acid where the error may be as high as 0.30 millivolt. In this latter case the measurable potential changes appreciably with time when the potassium chloride-agar bridges are inserted in the solutions, so that here the error may be quite significant.6 T o determine the liquid-junction potential E l ( e x pthe ) membrane-free chains: saturated calomel electrode potassium chloride saturated saturated potassium chlorideagar bridge electrolyte cz 1 electrolyte c1 1 'saturated potassium
I
1
1
1
1
1
I
1
1
I
0 The instability of the concentration potential in this case is of course not due to a l a c k of definition of the situation a t the membrane, but to a lack of stability of the two liquidjunction potentials: saturated potassium chloride-agar bridge I electrolyte c1 and saturated potassium chloride-agar bridge electrolyte cl. T h e same difficulty, of course, arises also without a membrane is measured. when t.he liquid-junction potential
I
ELECTROCHEMISTRY O F PERMSELECTIVE COLLODION MEMBRANES.
I
I1
303
1
chloride-agar bridge potassium chloride saturated saturated calomel electrode were measured for all the electrolytes a t the same concentrations as were used with the membranes. A simple W-shaped tube with the glass stopcock in the middle proved satisfactory for these measurements. Otherwise the instruments used and the technique employed were the skme as were used for the determination of the concentration potentials across membranes. At medium and higher concentrations the measurements were reproducible within about 0.10 millivolt, except a t the highest concentration of hydrochloric acid. At the lowest concentrations the resistance of the system was too high for best performance of the potentiometer used; E l ( e x pwas ) difficult to measure xith the usual accuracy and the errors may he greater than 0.10 millivolt in some instances. In table 1 below are given the values of the liquid-junction potential, as measured in the manner outlined, together with the corresponding calculated values, Elccalc,,and their differences A, the correction which must be applied to the experimentally determined value of the concentration potential. Table 1 gives for four different electrolytes the following experimental and calculated data a t the temperature 25.00"C. f 0.05'. Column 1 presents the concentrations ca and c 1 of the electrolyte solutions used in equivalents per liter; column 2 gives the theoretically possible maximum value of the concentration potential E,,, calculated as outlined in the preceding section; columns 3, 4, 5 , and G present the concentration potentials t across the various membranes corrected as outlined above ; column 7a gives the calculated liquid-junction potential Eiccalr) nhich ~ o u l darise in free solution, computed as indicated before ; column 7b gives the liquid-junction potential as measured; column 7c gives A, the difference between 7a and 7b, which has been applied to the figures given in columns 3, 4>5 , and 6 . The data presented in table 1 are given in figure 1 in the form of graphs t o facilitate their visualization and evaluation. Following a previously established convention (6, 12, 16) the concentration c1 indicated in the graphs refers to the more dilute solution;? plotted in this man-
'
The conventional plot of concentration potentials against the lower concentration is arbitrary. This procedure involves a n appreciable error unless the concentration ratios used closely approach the 1:l ratio. The physical properties-of the membranes which are represented by the experimental d a t a correspond t o a concentration which lies somewhere between the two concentrations used. The use of the mean concentration for reference value, which was suggested by Michaelis (7, 8, 9), unfortunately ha.s not been accepted b y the later investigators. The mean concentration, although a better choice than the lower concentration, is, however, not necessarily the correct reference concentration; the latter can be assumed t o lie nearer t o the more dilute concentration (1). With a 2 : l concentration ratio of the two solutions the error which is introduced by the conventional assumptions is not too great for the purpose a t hand; the curves in figure 1 under no circumstances could require a shift of more t h a n & log 2 (this i s 0.15 unit) t o t h e left. For the comparison of the different membranes and various electrolytes the arbitrary selection of ci as the reference concentration is of course inconsequential.
304 0
KARL SOLLNER AND HARRY P. GREGOR
Loq c, -I
c20
-2
-3 0 Lo9 c, -I t2 0
m v.
C
i
mv.
I5
-2
t t ? ~ I
t 15-
t 10-
-0
..........................................
-5
-5
'IT
............................................
.-g +" It #
-10
mv.
+ I5 w04
--............
O ................ 0+ ILoq c,
t5+
-I
-2
Fiq.1 Coneentrotion potmtials 0
-
Lo9 E , -I
-2
-3
c2:c,=2:I aemss various pennrolaetiva collodion membranes
-3
ELECTROCHEMISTRY OF PERMSELECTIVE COLLODIOX MEMBRANES.
305
11
TABLE 1 Concenlmtion potentials E ( c 2 : c L= 2 : 1) o j seceral electrolytes across rarious pernzselectice collodion membranes
0x12-
ox 10-
Kum 43
Hum 58
Hum 13
DIFFUSION POTENTIAL El IN FREE SOLUTION
bl EY B R A Y E
MEMBRANE Ox 14-
UEYBBANE
- Alc 65
A . Potassium chloride cqui:.;Iiier
~
1
0.002/0.001 0.004/0.002 1 0.01/0.005 1 0.02/0.01 0.04/0.02 0.1/0.05 1 0.2/0.1 0.4/0.2 '
I ~
mv.
17.5 17.3 17.1 16.9 16.6 16.3 16.1 16.0
mz.
n D.
nv.
mo.
17.5 17.3 17.1 16.9 16.6 16.0 14.8 13.4
17.5 17.3 17.0 17.0 16.6 16.1 15.2 13.9
17.5 17.3 17.1 16.9 16.5 15.6 14.5 13.0
17.5 17.3 17.1 16.9 16.4 14.8 13.2 10.4
1
'
1
I
mu.
1
-0.3 -0.3i -0.31 -0.38 -0.31 -0.3,
I -0.31 1 ~
'
nv.
mo.
nil nil nil nil nil nil nil nil
-0.31
B. Lithium chloride
0 002/0 001 0 004/0 002 0 01/0 005 0 02/0 01 0 04/0 02 0 1/0 05 0 2/0 1 0 4,O 2
'
17 5 17 3 17 1 16 9 16 7 16 5 17 0 169
17.4 17.1 16.9 16.7 16.4 156 14.7 13.4
17 5 17.1 16.7 16.7 16.4 15.4 14.5 12.6
17.5 17.2 16.9 16.7 16.3 15.0 13.9 12.3
17.4 17.2 16.6 16.4 15.8 14.2 12.0 8.7
I
-5.8 -5.81 -5.81 -5.81 -5.9' -6.0; -6.3, -6.51
-5.7 -5.7 -5.6 -5.6 -5.5 -5.4 -5.1 -4.7
-0.1 -0.1 -0.2 -0.2 -0.4 -0.6 -1.2
-1.8
C. Potassium sulfate 0 00210 001 17 4 0 004/0 002 17 1 0 01/0 005 16 9 0 02/0 01 1 16 7 0 04/0 02 1 16 4 0 1/0 05 1 16 1 0 2/0 1 15 9 0 4/0 2 15 8
1
17.4 17 .1 16.9 16.7 16.2 16.2 15.9 155
17.4 li.2 16.8 16.7 16.4 16.1 15.9 15,5
li.4 17.1 16.8 16.7 16.4 16.2 15.7 152
17.5 17.2 16.9 16.6 16.4 16.0 15.5 14.7
4.2 4.4 4.6 4.8 5.0 5.4 5.7 6.1
4.2 4.3 1.5 4.6 4.8 4 9 5.0 5.1
0.0 +0.1 +0.1
+0.2 +0.2 +0.5 +0.7 f1.0
D . Hydrochloric acid 0.004/0.002 ' 17.3 0.01/0.005 17.1 0.02/0.01 1 17.0 0.04/0.02 16.8 0.1/0.05 16.8 0.2/0.1 16.8 0.4/0.2 17.6 ~
I
17.2 16.8 16.7 16.1 15.1 15.0(?) 16.0(?)
, ~
'
lZ.l 16.8 16.7 16.2 15.1 l5.0(?) 15.9(?)
17.1
1
17.0 16.3 16.0 16.1 15.3 1 14.2 14.8 11.6(?) 14.1(?) 15.4(?) 15.2(?)
if3:;
1
11.3 11.2 11.1 11.1 11.0 11.1 11.1 11,s
' 1
1
I ~
11.3 11.2 11.1 10.6 10.2 9.8 9.0(?) 8.0(?)
0.0 0.0 0.0 +0.5
f0.8 +1.3 f2.2(?) +3.8(?)
306
KARL SOLLNER AND HARRY P. GREGOR
ner the data become easily comparable to many analogous data published in the literature. In many instances the experimental concentration potentials obtained with different membranes and the same electrolyte coincide. In order to make the graphic representation of these points feasible, the expedient has been chosen in figure 1 of plotting some of the points not in their proper position but of arranging them outside the curves and indicating their proper position by arrows.
v The first impression one gains when looking a t the data of table 1 and figure 1 is that of their high degree of regularity and consistency. In the subsequent publication an attempt will be made to apply to these data two kinds of objective quantitative considerations, one of these already used in principle by Michaelis ( 7 , 8 , 9 ,lo), the other one developed to utilize fully for the elucidation of membrane action the experimental data given here. In the present paper we shall evaluate these data merely in the conventional qualitative and semiquantitative manner. The criterion for this is the difference between the calculated theoretically possible maximum values E,,, and the corresponding experimental concentration potentials E . The greater the difference between two corresponding such values, the lower is the “ionic selectivity” of a membrane under the particular conditions. When considering the data presented in table 1 and figure 1 the emphasis may be placed on the one hand on the comparison between the different membranes; on the other hand, one may focus one’s attention on the differences between the various electrolytes. With regard to the former problem it is evident that with the lowest concentrations used all membranes behave as ideal-at least within the limits of experimental error. The differences between the various membranes become apparent in the differences of the concentrations a t which a significant deviation of the experimental from the calculated concentration potential can be observed. The disparity between the membranes of the various types is most pronounced with the highest concentrations investigated, except in the case of hydrochloric acid, which shows a peculiar behavior (see below). Any comparison therefore has to be based preferentially on the results obtained with the medium high and the highest concentrations. Although the concentration potentials across membranes Ox 8 - Hum 43 and Ox 12 - Hum 43 coincide with both potassiym sulfate and hydrochloric acid within the limits of experimental error,* the ionic selectivity of the membranes follows the sequence: membrane Ox 12 - Hum 43, Ox 8 - Hum 43, Ox 14 8 This observation probably is explained by the fact t h a t on the one hand membrane Ox 8 - H u m 43 is less highly oxidized t h a n membrane Ox 12 - H u m 43, and on t h e other hand i t is of lesser porosity, these two antagonistic factors seem t o compensate each other within the limits of experimental error.
ELECTROCHEMISTRY OF PERMSELECTIVE COLLODIOX MEMBRANES.
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307
Hum 58, Ox 10 - Hum 43 - Alc 65, the latter membrane having the lowest degree of ionic selectivity. This sequence of ionic selectivities of the different membranes is identical with that found previously (2) with 0.1 N notassiiim chloride/O.Ol A; potassium chloride. One may turn now to the other part of the conclusions which can be drawn from the data presented in table 1 and figure 1-namely, a comparison of the effects which are observed with the solutions of various electrolytes. For a detailed comparison of the various electrolytes one must keep in mind the limits of the accuracy of the calculated as well as of the experimental data. Some discretion must be used not to transgress the inherent limits of the usefulness of these values in any attempt to establish minor differences between various electrolytes. Differences between calculated and experimental concentration potentials of =t0.20 millivolt can hafdly be taken as significant; discrepancies of this magnitude may easily be due to a combination of limited experimental accuracy and systematic error in the computation and correction methods used. For this reason any conclusions which are drawn with respect to the differences of behavior of different electrolytes should be based primarily on the trend of the curves over wide concentration ranges and on the data obtained with higher (> 0.05 3') concenwations. I t is immediately evident from table 1 and figure 1 that the agreement betwen the calculated and the experimental concentration potentials is best with potassium sulfate. In other words, the ionic selectivity of the membranes is highest in this case, far surpassing that observed with the other electrolytes. The ionic membrane selectivity ivith potassium chloride, while being considerably less than with potassium sulfate, is decidedly greater than with lithium chloride. The behavior of hydrochloric acid deviates from that of the neutral salts potassium chloride and lithium chloride, in that the membrane selectivity is not quite as high a t medium concentrations as with the neutral salts; a t the highest concentrations, on the contrary, the selectivity of the membrane in the case of hydrochloric acid is higher than tvith potassium chloride and lithium chloride. With the membrane of the highest ionic selectivity reported on in this paper (membrane O s 12 - Hum 43), ideal ionic selectivity within the limits of the error of experiments and computations is observed with potassium sulfate a t all concentrations up to 0.2/0.1 N; n ith potassium chloride, lithium chloride, and hydrochloric acid this limit is approximately 0.04/0.02, 0.02/0.01, and 0.004/ 0.002, respectively. With the membranes of lesser ionic selectivity these figures are correspondingly lon-er, but even with the alcohol-swelled membrane Ox 10 - Hum 13 - Alc 65 ideal ionic selectivity exists with potassium sulfate up to a concentration of 0.1/0.05 'V. These results, generally speaking, are in agreement with the concept put forn-ard by Michaelis ( 7 ) that electrical and steric pore blocking account for the characteristic ionic selectivity of membranes of porous character, According to the theory T! hich has been stated of late more precisely ( G , 11, 12, 13, 16, IS), the electrochemical structure of an electronegative membrane is due t o the presence of an essentially constant number of anionic groups perma-
308
KARL SOLLKER AKD HARRY P. GREGOR
nently built into the malls of the pores of the membranes; these groups, carboxyl groups in the case of collodion membranes (G, 11, 12, 14, 15, l G ) , are compensated for electrically by the same species of cations as that present in the adjacent solutions. Since the “degree of dissociation” of these surface compounds does not vary appreciably with the different alkali ions, the electrical situation a t the pore walls is very similar in the various neutral electrolyte solutions.9 In acid solution the surface carboxyl groups are combined n-ith hydrogen ions as counter ions; the resulting free carboxylic acids cannot be expected to act like strong electrolytes. The electrochemical structure of the membranes in acid solution, a t all concentrations except perhaps the very lowest, therefore, must be expected to differ considerably from that existing with fair uniformity in all neutral solutions where the acidic (anionic) surface groups with their counter ions, K+ or Li+, act as strong electrolytes (16). Let us consider now in the light of the above-outlined theory the results obtained with the various electrolytes, focussing our attention first on the neutral salts. In the case of potassium sulfate, the electrical and the steric blocking reinforce one another. The sulfate ion is prevented from penetrating the membrane both on account of its size and on account of the electric repulsion which arises between the fixed negative groups a t the pore walls and its two negative charges. On account of its double charge alone the sulfate ion would be screened out much more effectively than ions irhich carry single charges only, such as the chloride ion. With potassium chloride the steric factor does not play an important rBle, since potassium and chloride ions have very nearly the same size; only the electric factor of pore blocking is operative. In the case of lithium chloride the steric factor of pore blocking on account of the greater size of the hydrated lithium ion results in a lower degree of ionic selectivity of the membrane as compared with potassium chloride ; certain pathways across the membrane which are accessible on a purely steric basis to potassium and chloride ions are inaccessible to the lithium ion. The behavior of the membranes with hydrochloric acid cannot be interpreted in the same straightfonvard manner as that of the neutral salts. On the basis of steric considerations only, one would expect that the hydrogen ion would go particularly easily across the membrane; in other words, one would expect the situation to be somewhat similar to that observed with potassium sulfate, where the cation is much smaller than the anion. This, however, is obviously not the case; the different electrochemical structure (and its dependence upon concentration) which prevails in acid solution must be considered. The lower degree of dissociation of the electrochemically active surface compounds within acid solution results in an electrochemical structure of the membrane under considera9 Some pores of a membrane for geometrical reasons may not be available for the one or the other species of ions of a n electrolyte; nevertheless, those critical acidic (anionic) wall groups which are operative are compensated for electrically by the cation of the electrolyte under consideration.
ELECTROCHEMISTRY O F PERYSELECTIVE COLLODIOS >lEMBR.%NES. 11
309
tion n-hich is analogous to that prevailing in neutral solution with membranes having a smaller surface concentration of potentially dissociable anionic groups ; for the case of the collodion membrane this means that the behavior of a highly oxidized membrane in acid solution is similar t o that shown in neutral solution by less highly oxidized membranes. The above-outlined explanation of membrane behavior in acid solution cannot, as it seems, be applied t o the highest acid concentrations for which data are available if the latter are taken a t their face value. At the highest acid concentrations used the membrane selectivity is slightly greater than,at lower concentrations. An explanation of this latter observation is still missing. I t may possibly be due to a systematic error in the measurement or the computation of the liquid-junction potential (or a combination of these two factors) Tvhich mas used as a correction for the measured empirical values of the concentration potential. However, there is no proof that this situation really prevails. I t is equally possible that some unknown physical factor causes the apparent anomaly. Only further experiments will permit a decision regarding this question.’” The concentration range of useful highest ionic selectivity of these membranes can be expanded to an appreciable extent beyond that indicatecl in table 1 and figure 1 by keeping the concentration at the one side of the membrane fairly IOK, e.g., ten or one hundred times lower than that of the more concentrated solution. With this precaution the membrane selectivity as determined by potential measurements can be fully satisfactory JTith electrolyte concentrations considerably higher than is indicated by the above presented experimental data with the 2 : 1 concentration ratio. This fact is ivell known from measurements x i t h the 10:1 concentration ratio of the characteristic concentration potential 0.1 S potassium chloride/0.01 S potassium chloride and analogous measurements with other electrolytes. The higher degree of membrane selectivity observed under these conditions can be explained in the folloiving manner. The permselective collodion membranes are relatively thick, about 40 p. When one side of the membrane is in contact with a sufficiently dilute solution, an adjacent layer of the membrane retains a very high degree of ionic selectivity. This layer of the membrane is not greatly influenced by the presence of a relatively concentrated solution on the other side of the membrane, although the layer of the membrane which is in immediate contact n-ith this solution may show a someivhat impaired degree of ionic selectivity. As was pointed out above, all the potential measurements and calculated I o I n this connection i t is interesting t o note t h a t Michaelis, Ellsworth, and Weech (8) with membranes much denser than the permselective membranes have found a much higher membrane selectivity with hydrochloric acid than with the alkali halides, the investigated concentration range being from 0.02 S/0.01 S to 0.16 S l 0 . 0 8 A’. A final explanation of this discrepancy between the results of the above-mentioned authors and our own d a t a is still lacking. I t probably has t o be looked for in the much smaller size of the majority of the pores of the membranes used by Michaelis and collaborators. 4 s will be shown in a later paper, a similar discrepancy erists betxeen the two sets of d a t a in the case of the membrane resistance.
‘
310
KARL SOLLNER AND HARRY P. GREGOR
potential values dealt with in this paper have a definite limitation in significance. Each individual calculated or experimental potential value has a probable error of 0.10 millivolt, in a few instances even more. Although this situation is somewhat ameliorated by the fact that the consistent trend of curves has a greater significance than the individual numerical values, small deviations of the membranes from ideal selectivity, small differences between different membranes, and minor differences between different electrolytes can better be established by methods other than direct potential measurements. As such methods one may mention the combination of “leak” (2) and resistance (2) measurements, studies on the transfer of ions in an electric field across the membranes (7, IO), studies of both the rates of cation and anion exchange across the membranes, csrried out preferentially with tracer elements11 and others. SUMMARY
1. Concentration potentials of potassium chloride, lithium chloride, hydrochloric acid, and potassium sulfate solutions across several types of permselective collodion membranes were measured a t several concentration levels between 0.001 Ar and 0.4N,the concentration ratio being 2 : l . 2. The experimentally obtained concentration potentials are compared with the calculated values of the concentration potentials which would arise with membranes of ideal ionic selectivity, in other words, with the potential which would originate if the membranes would act as ideal reversible electrodes for the critical (cat)ion. 3. The agreement between the calculated and the experimental concentration potentials in all instances is closest with the more dilute solutions. With the membrane of the highest ionic selectivity reported on in this paper (membrane Ox 12 - Hum 43), ideal ionic selectivity within the limits of the error of the experiments and computations f 0.2 millivolt is observed with potassium sulfate a t all concentrations up to 0.2/0.1 N; with potassium chloride, lithium chloride, and hydrochloric acid this limit is approximately 0.04/0.02, 0.02/0.01, and 0.004/0.002, respectively. With the membranes of lesser ionic selectivity these figures are correspondingly lower. 4. The results are in agreement with the fixed-charge theory as applied to membranes of porous character. Also, the existing differences in behavior between neutral electrolytes on the one side and hydrochloric acid on the $her side can be explained essentially on the basis of this theory, by reference to the changed electrochemical structure of the membranes in acid solution owing to a lower degree of dissociation of the electrically active surface compounds. REFERESCES
(1) GREGOR,H. P.: Ph. D . Thesis, University of Minnesota, 1945. (2) GREGOR, H . P., A N D SOLLIVER, K.: J. Phys. Chem. 60, 53 (1946). (3) H ~ R K EH. D ,S . , A N D OWEN,B.: The Physzcal Chemistry of Electrolyte Sohtions. Reinhold Publishing Corporation, New York (1943). 1 1 The reasons for this l a t t e r restriction will become apparent in a forthcoming paper on the bi-ionic potential.
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(4) International Critical Tables, Vol. VI. McGraw-Hill Book Company, Inc., Kew York (1929). (5) M A C I K N E SD,. A . : The Principles of Electrochemistry. Reinhold Publishing Corporation, Xew Tork (1943). (6) M E Y E RK , . H., ASD SIEYERS,J.-F.:Helv. Chim. Acta. 19, 649, 665 (1936). RIEYER, K. H . : Trans. Faraday Soc. 33, 1073 (1937). (7) RfICHIELIS, L . : Bull. K a t l . Research Council, No. 69, 119 (1929); Kolloid-Z. 62, 2 (1933). (8) MICHAELIS, L . , ELLSWORTH, R. McL., ASD WEECH,A . A , : J. Gen. Physiol. 10, 671 (1927). (9) MICHAELIS, L . , ASD WEECH,A . A . : J. Gen. Physiol. 11, 147 (1927). (10) l I I C H A E L I S , L., WEECH, A . 8 . ,A S D YAlL4TOR1, A . : J. Gen. Physiol. 10,685 (1927). (11) SOLLSER, K . : J. Phys. Chem. 49, 47 (1915). IC: J . Phys. Chem. 49, 171 (1915). (12) SOLLKER, (13) SOLLNER, K . : J . Phys. Chem. 49, 265 (1945). (14) SOLLKER, K., ABRAXS,I . ,A N D CARR,C. W.: J. Gen. Physiol. 24,467 (1941). (15) S O L L S E R , K . , .&BRaMS, I., AKD C.4RR, J . Gen. Physiol. 26, 7 (1941). (16) SOLLNER, K., A X D CARR,C. W . : J. Gen. Physiol. 28, 1 (1944). (17) SOLLNER, K., AND GREGOR, H. P . : J. Phys. Chem. 60,470 (1946). (18) TEORELL, T . : Proc. SOC.Exptl. Biol. N e d . 33, 282 (1935) ; Proc. Natl. Acad. Sci. 1.; S. 21, 152 (1935).
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T H E COLLOID CHENISTRY OF T H E CLAY MIKER.1L ATTAPULGITE’ * C. E. MARSHALL
AND
0. G . CALDWELL
Department of Soils, Missouri Agricultural Experiment Station, Columbia, Missouri Received August 8 , 1046
httapulgite, the most recently characterized of the clay minerals, is of especial interest t o colloid chemists. Its colloid-chemical properties have been little investigated, but sufficient is known of its structure t o act as a stimulus to further research. It affords a beautiful example of the close relationship between atomic structure and colloidal properties in the group of the silicates. In 1935 de Lapparent (G) decided that a certain fuller’s earth from Attapulgue, Georgia, and a similar clay from Mormoiron, France, were quite distinct from the clays of the montmorillonite group in spite of some similarity in chemical composit,ion. Both were hydrous aluminum magnesium silicates, but the x-ray data and thermal dehydration curves indicated that the new clays were more closely related t o the hydrous magnesium silicate sepiolite (or meerschaum) than t o montmorillonite. Further work by de Lapparent (7, 8) and by Longchambon (9, 10) strongly suggested that attapulgite and sepiolite are members Presented at the Twentieth National Colloid Symposium, which was held a t Madison, Wisconsin, May 28-29, 1946. 1 Contribution from the Department of Soils of the Missouri Agricultural Experiment Station, Journal Series No. 1007.
