Dec., 1953
ADSORPTION OF SILVER SALTSON SILVER
that electrical conduction can also readily occur across the amphiphilic electrolyte layers which, in this case, virtually form a quasi-cellular structure enclosing the water. The progressive change in conductivity with composition or temperature* in the SI-or Sz-phases indicates that the relative abundance of the more and less-conducting micellar forms changes gradually with changing conditions, Le., that the different forms co-exist in equilibrium and are mutually transformable.
895
The discontinuous form of the conductivity curve in the region of the existence of the liquid crystalline phase G indicates, when taken with the other special properties of this phase, that its micellar organination is fundamentally different from that of the S1-and &phases. The variation in structure between the latter two phases is however continuous in character. Acknowledgment.-The writers wish to thank the Shell Petroleum Co., Ltd., for permission to publish this paper.
THE ADSORPTION OF SILVER SALTS ON SILVER1 BY CECILV. KINGAND RUTHKIMMELMAN SCHOCHET Department of Chemistry, New Yorlc University, New Yo& Received March 10, 1066
The amount of adsorption of silver nitrate, perchlorate and sulfate on crystalline silver powder, from aqueous solutions, has been measured. The results at 25",when the metal surface has received no special treatment indicate less than monolayer adsorption. If the metal has been washed with dilute nitric acid, much larger amounts of siiver salt can be taken up. A few measurements have been made at 0". The results are compared with measurements found in the literature.
The amount of silver nitrate adsorbed on silver (and gold) was first determined by v. Euler and his co-w0rkers.~-5 The first measurements were made with very fine precipitated silver (particle diameter 0.6-0.8 p ) . The values found were consistent, but the authors were not satisfied with their estimation of surface area. Rudberg and v. Euler employed silver sheet, and found several times as much adsorption per unit of apparent area. Experiments by Proskurnin and Frumkin,' designed to find the concentration of silver ion which would make silver a "null electrode," allow calculation of the amount of adsorption of silver nitrate on sheet silver. The values cover an extreme range and are not very self-consistent, which is probably to be explained by the method of surface preparation employed. Experiments intended to be similar to those of v. Euler and Hedelius, using fine precipitated silver, have been reported by Tartar and Turinsky.' These authors concluded that any apparent adsorption was really due to reaction with traces of reducing agent left in the silver. However, in most cases their ratio of silver to solution volume was too small to detect adsorption, and their results are inconclusive. Work in this Laboratory on the rate of dissolution of silver in ferric sulfate*and perchlorateQsolutions has shown that silver salts, whether formed in (1) Based on a Ph.D. thesis submitted by Ruth Kimmelman Schochet to the Graduate School of New York University. Work done under U. 8. Atomio Energy Commission Contract No. AT (30-1)-816 with New York Univemity. (2) H. v. Euler and A. Hedeliua, Arkiv Kemi, Minsralogi Qeolopi, 7 , No. 31 (1920). (3) H. v, Euler and G.Zimmerlund, {bid., 8, No. 14 (1921). (4) E. G. Rudberg and H. V. Euler, 2. Physik, 18,275 (1923). (5) H. V. Euler, 2. Elektrochem., 28, 446 (1922). (6) M. Proskurnin and A. Frumkin, 2. phyaik. Chem., 166A, 29 (1931). (7) H. V. Tartar and 0. Turinaky. J . A m . Chem. Boc., 64,580 (1932). (8) H. Salrberg and C. V. King, J . Electrochem. Soc.. 97, 290 (1950). (9) C. V. King and F. 8. Lane, ibid., SS, 295 (1952).
the reaction or added separately, have a pronounced inhibiting effect. Equations based on surface coverage following isotherms of the Langmuir type were used to express the dissolution rates. A few adsorption experiments with silver perchlorate on reduced silvergwere found to follow approximately a Langmuir isotherm, and the adsorption per unit apparent area, extrapolated to infinite concentration (l/c = 0), was about twice that estimated from the values of Rudberg and v. Euler for silver nitrate. Because of the great discrepancies in published values, and to find more about the influence of the anion, it was thought desirable to make quantitative measurements of the adsorption of silver nitrate, perchlorate and sulfate on the metal. Experimental The Silver.-In order to adsorb appreciable amounts of the salt, it is necessary to have a large silver surface area in contact with the solutions, and for this reason precipitated silver was employed. Most of the measurements reported here were made with one lot of Mallinckrodt silver. The analysis indicated 0.005% other metals, a similar per cent. of chloride, and 0.05% sulfate. Preliminary measurements were done with the lot of silver used by King and Lang,D which was prepared as described by Walden, Hammett and Edmonds.10 Both samples were similar in appearance under the microscope, consisting originally of quite uniform rectangular particles whose dimensions could be measured easily. After some use t,he particles were in general less uniform, the edges and corners were less sharp, and the estimated area had increased as much as 15%. About 80 particles were examined to calculate the area. Another lot of Mallinckrodt silver was purchased, but not used in the experiments. The particles were smaller and less uniform in size. The estimated area was 1780 cm.g/g.; low temperature nitrogen adsorpt,ion" gave an area of 2380 cm.a/g. If the true areas of all the samples are considered to have this ratio to the microscopic area, the adsorption values given below should be multiplied by 0.75. (10) G. H. Wslden, L. P. Hammett and S. M. Edmonds, J . A m . Chem. Soc., 66, 350 (1934). (11) Courteay of Sylvania Electric Produota, h a .
CECILV. KINGAND RUTHKIMMELMAN SCHOCHET
896
The Measurements.-Bottles containing silver and solution were rotated in a thermostat at 25 f 0.02’ except for a few experiments done in an ice-bath. It is essential to have a silver surface-solution volume ratio which will result in a reasonable concentration change. King and Lang used 10 g. of silver, estimated area 500 cm.2/g., with 25 ml. of silver perchlorate solution, and. found 4 1 0 % loss from the solution. When the same silver was used with silver sulfate solution in the present work, the concentration change was found to be 2% or less; with silver nitrate the change was even smaller. The discrepancy was later traced to treatment of the silver with dilute nitric acid before use by King and Lang. This apparently “activates” the surface in some way, and the effect is not permanent. Without such activating treatment, it was found necessary to increase the amount of silver to 55 g. with 25 ml. of solution. This was the largest amount which would move about freely when the bottles were rotated slowly. This resulted in a loss of &12% silver sulfa& from solution, somewhat less perchlorate and nitrate. All of the silver powder was used in preliminary runs, so that it had been in contact with silver salt solution for several hours before use in final experiments; this should be effective in removing less noble metals from the surface, and possibly in reducing local cell action. Before each experiment the silver was washed very thoroughly with water, rinsed with alcohol and fat-free ether, and dried with air on a glass filter. Since surface contamination with oxide or sulfide was feared, a few samples were rinsed with 5’% potassium cyanide before washing with water and drying. There was no pronounced effect on the adsorption measurements. The time necessary to ensure equilibrium was explored thoroughly. Two hours was sufficient for the nitrate and sulfate; perchlorate was taken up very slowly for a much longer time, and the experiments reported were continued for 72 hours. Analysis of the Solutions .-Analytical grade potassium bromide was used as the primary standard, and titrations were made with Rhodamine 6 G as adsorption indicator. Successive samples of approximately 10 ml., dehvered carefully from a pipet, could be titrated with a precision approaching one part per thousand. Samples of original and final solutions were titrated in succession with the same bromide solution to obtain maximum accuracy in measuring the concentration changes. From two to six samples were run a t each concentration, and the results averaged. The average deviation in amount of adsorption was ordinarily 1 to 6%, in only two expcriinents above 10%.
Results The amount of adsorption of the three salts is represented in Fig. 1 in terms of the reciprocal of 0.6 2
equivalents per cmB2apparent areal2 vs. the reciprocal of equivalent concentration; this is a convenient way of extrapolating to “infinite concentration” (l/c = 0). Since silver sulfate is not very soluble, the concentration range for this salt is limited. The other salts were used to the highest practical concentrations to assist in the extrapolation. To find whether dilute nitric acid treatment was responsible for the much larger sorption values found for silver perchlorate by King and Lang, silver samples were washed with 4.5% nitric acid (7 1. commercial acid per 100 ml.) as was done by ‘tp,em. Brief rinsing on a suction filter had little effect, and the silver treatment was repeated as follows: 30-g. samples were washed with six to eight 50-ml. portions of acid, allowing 2 or 3 minutes each time before applying suction to remove the acid. The silver was then washed and dried as usual. Measurements are given in Table I, and are to be compared with values ranging from 3.5 to 18 X equiv./cm.2 plotted in Fig. 1. TABLE I NITRICACIDTREATED SILVERAT 25 O 10.0 g. silver, area 675 cm.*/g., bottles rotated 72 hr.
SORPTIONOF AgClO,
ON
C orig.,
C final, N
Equiv. cm.-* X 1010
0.0226 ,0496
0.02075 ,0458 ,0561 .2495 ,437
68 141 167 330 360
N
.0606
,2585 ,447
Rudberg and v. Euler reported no significant difference in the adsorption of silver nitrate on gold between 0 and 5 0 ° , but did not carry out similar experiments with silver. A few experiments with silver nitrate a t 0” are given in Table 11. The bottles were immersed in an ice- and water-bath and were shaken frequently during the day. TABLE I1 ADSORPTIONOF AgN03 ON SILVERAT -0’ 25 ml. solution, 55.0 g. silver, area 6!5 cm.2/g., two samples each concentration Hours in ice-bath
C orig.,
C final, N
8 25 25
0.01885 .0 1922 .09825
0.01855 ,01730 ,08715
El 0 . 5 X N
0.4
Vol. 57
N
Equiv. cm.-z X 1 0 1 0
2.3 14.1 82
It is well knomii that silver ion in solution undergoes exchange with silver metal, to such an extent that penetration t o the depth of several atomic layers, or its equivalent in surface electrolysis or recrystallization must be assumed.13 A few adsorption experiments were carried out with silver nitrate containing a small amount of radioactive Ag1l0 in order to measure the adsorption and exchange simultaneously. Samples mere titrated as
0
d
.C
$0.3 al
u
3 0.2
e
.-a, 2 0.1
(12) R. B. Dean, THIS JOURNAL, 65, 611 (1951), has proposed that mole om.-) be regarded as en adsorption unit and called the
10-10 I
Fig. 1.-The
I
I
1
I
I
20 40 60 80 100 120 Reciprocal of concentration. adsorption of silver salts on silver at 25“.
Gibbs. (13) B. V. Rollin, J . Am. Chem. Soc., 6B, 86 (1940); C. C.Coffin and I. I. Tingley, J . Chem. Phya., 17, 502 (1949); H. Gerischer and W. Vielstich, 2. Elektvochem.. 66, 380 (1952); M.Ants and M. Cottin, Compl. rend., 234, 1686 (1952).
ADSORPTION OF SILVER SALTSON SILVER
Dec., 1953
usual, and 2-ml. samples of initial and final solutions were placed in small Petri dishes and counted with an end window GM tube. The initial count of each solution was about 800 per minute above background. Two and three duplicates at each concentration gave the usual agreement in amount adsorbed (-5%) and similar agreement in amount of exchange. The removal of silver nitrate and of radioactivity from the solutions is shown in Table 111.
897
Reciprocal of concentration. 100 200 300 400 0.2 I I
I
I
I
I
500 I
600 1
I
I
I
I
I
I
TABLE I11 ADSORPTION A N D EXCHANGE WITH AgNOs CONTAINING Agllo ON SILVERAT 25 25 ml. solution, 55.0 g. silver, area 675 cm.2/g., bottles rotated 20-25 hours C orig.,
AgNOn removed,
Activity removed,
0.01876 ,04855 .09875
2.03 1.55 0.81
63 60 77
N
%
% '
Exchange Adsorpadsorption, equiv. tion, equiv. om.-' X 10" cm.-* X 1010
+
2.6 5.1 5.4
79 196 510
Discussion A Langmuir adsorption isotherm can be put in the form
I
240 160 200 80 120 Reciprocal of concentration. Fig. 2.-The adsorption of silver nitrate and perchlorate on silver: upper curve, v. Euler and Hedelius, precipitated silver; lower curves, A, King and Lang; B, Rudberg and v. Euler, silver sheet. 40
values for silver sulfate are in the same range as those of v. Euler and Hedelius for the nitrate, and !=a 1 both extrapolate to similar values at l / c = 0. x C + G It seems more probable that a factor such as surWhere z -= aniount adsorbed per unit area, a is a face cleanness is involved, with other adsorbed maconstant, c = concentration and xo is the saturation terials being displaced most easily by the sulfate, adsorption. In Fig. 1, values for silver sulfate and but by the other salts at sufficiently high concenfor more concentrated nitrate solutions are in rea- trations. While the measurements discussed above seem to sonable agreement with such isotherms, while values for the perchlorate approach those for the ni- indicate that a simple adsorption process occurs trate a t higher concentrations. It is significant normally, other results show that more complicated that the three curves extrapolate to about the same sorption can take place under suitable conditions. value, near 0.05, at l / c = 0, indicating saturation Curves A and B of Fig. 2 show the measurements equivalents per cm.2. of King and Lang on nitric acid treated precipitated adsorption near 20 X A single atomic layer in the (100) plane of the silver, and of Rudberg and v. Euler on sheet silver. silver crystal contains 19.9 X 10-lo gram atoms These curves extrapolate to 87 and 200 X 10-lo per while the reticular density of other princi- equiv./cm.2, and the smallest value measured is pal planes is somewhat smaller. Since the true about 40 X If the silver used by King and area is almost certainly not less than the meas- Lang had not failed to give similar sorption in laured area, we can say that the adsorption at 25" ter use, we should ascribe the large values to enon silver not treated with nitric acid is always less hanced surface area. The silver used by Rudberg than monolayer in nature. One is inclined to be- and v. Euler was treated with steam and alcohol lieve that, with precipitated silver crystals which vapor, rinsed with water and dried a t 100". Mehave been in contact with silver salt solution for chanical polishing or nitric acid etching (presumseveral hours, the real area is not much different ably followed by the above treatment) did not affrom the apparent area, and that the limiting fect the results by more than the experimental error amount of adsorption is a monolayer. (about 25%). The measurements of v. Euler and Hedelius with Cooling to 0" certainly does not affect the sursilver nitrate on precipitated silver have been face area, but can have a pronounced effect on the plotted in the same manner in the upper part of sorption (Table 11). The surface treatments menFig. 2. The straight line has been drawn arbitrar- tioned above probably do not increase the real area ily to an intersection of 0.05 X lolo, but could to 5 or 10 times the apparent area but, instead, fahardly be drawn much differently. While the au- vor multilayer sorption. Consistent measurements thors felt uncertain of the area within f50%, it then demand consistent surface treatment, and seems more than coincidence that the extrapolated each experimenter finds different values. value agrees with ours. However, adsorption at According to the measurements of Proskurnin the low concentrations employed by these authors, and Frumkin,@silver acquires a positive charge in while representing less than a monolayer, is much solutions containing its ion in concentrations above greater than was found in the present experiments 10" N . The charge density on the mota1 results in with silver nitrate. To make the two sets of data an electrical double layer capacity in the neighoverlap in the dilute range, one of the area estima- borhood of 20 microfarads per Since the potions would have t o be in error by a factor of.13 tential difference between metal and solution is times. This is rather unlikely, especially since our only a few tenths of a volt, this would require less
LAWRENCE M. KUSHNERAND WILLARDD. HUBBARD
898
adsorption of silver ion alone than any of the values give above; substitution in the relation between capacity, charge and potential 20
x
96500x 10-6 = 0.1
gives x = 0.2 X 10-loequiv./cm.2 per 0.1 volt. Consequently adsorption is specific, involving attachment of both anions and cations to the surface, and not merely a process of charging the double layer. We may assume that silver ions are adsorbed first, but become a part of the metal lattice; the positive charge carried does not remain localized. To obtain the equivalent of monolayer adsorption, it would not be necessary that silver ions deposit over the whole surface but, rather, that anions become intimately attached in a monolayer. If the anions should be bound approximately as rigidly as in the salt lattice, further sorption might reasonably be expected to occur. The process would not continue indefinitely, because the effect of extra charge on the metal and the lattice forces would be balanced by thermal agitation and dissolution forces acting on the outer layer.
Vol. 57
One can understand why low temperatures might favor such multilayer sorption, but we have no explanation for the effect of nitric acid too dilute to etch the surface, or possibly the alcohol vapor treatment of Rudberg and v. Euler. Exchange experiments show that the surface atoms of a metal, to a depth of several atomic layers at least, have a great deal of mobility. The experiments of Table I11 show that 60 to 7701, of the silver ions originally in solution become part of the metal in some 20 hours, while an equal amount of the original metal dissolved. Exchange in depth may be aided by local cell action, but there is probably the equivalent of an abnormally large self diffusion coefficient in the surface layers, as postulated by Gerischer and Vielstich.la The mobility should be an aid in establishing adsorption equilibrium and especially in the displacement of less firmly adsorbed substances. The tremendously rapid exchange must serve to repair minor irregularities of the surface when the metal is immersed in a solution of its own ions. The true area would tend to become equal to the measured area unless the metal is actually porous or has major faults such as abrasion marks, corroded areas, etc.
ON THE DETERMINATION OF CRITICAL MICELLE CONCEKTRATIONS BY A BUBBLE PRESSURE METHOD BY LAWRENCE M. KUSHNER AND WILLARD D. HUBBARD Surface Chemistry Section, Division of Chemistry, National Bureau of Standards, Washington 26, D. C. Received March IS, 1966
Bubble pressure measurements have been made on aqueous solutions of pure and commercial grade sodium dodecyl sulfate. The curves obtained with the commercial material have considerable structure depending on the bubbling rate. The curves for the pure samples are what one might expect on the basis of the surface tension of the solutions. The interpretation of bubble pressure and conductance data for solutions of impure surface active materials is discussed.
Introduction In a recent publication,' Brown, et al., have described a bubble pressure method for the determination of the critical micelle concentration of surface active agents in aqueous solution. The method involves measurement of the air. pressure necessary to maintain a stream of bubbles from a small diameter tube immersed in the detergent solution. Plots of bubble pressure versus concentration of detergent, at different rates of bubbling, show distinct and reproducible irregularities at particular concentrations. The electrical conductances of the same solutions show sudden, small changes at some of these concentrations. Brown, et al., interpret these effects as indicating more than one critical micelle concentration and also feel that the data present a strong argument for the existence of more than one micellar species in these dilute solutions. They do however point out that all of the surface active compounds used were commercial grade, and suggest that some of the effects noted may be due to the presence of impurities. Since their apparatus is easily reproduced, it was decided to make bubble pressure (1) A. 8. Brown, R. U. Robinson, E. H. Sirois, H. G. Thibault, W. MoNeill and A. Toflas, TEXIS JOURNAL, 66, 701 (1962).
measurements of this type on a sample of commercial sodium dodecyl sulfate and a sample of pure sodium dodecyl sulfate which had been synthesized in this Laboratory for an earlier researchS2
Experimental Materials.-The commercial materials was used as received. The ure sodium dodecyl sulfate was synthesized, as described by j h e d l o ~ s k yfrom , ~ a vacuum-distilled sample of n-dodecyl alcohol. The chlorosulfonic acid for the synthesis was distilled immediately before use. All other reagents and subsequent purification conformed with American Chemical Society specifications. The final step in the purification procedure consisted of extracting the detergent crystals with diethyl ether for about 8 hours in a Soxhlet extractor. Apparatus and Procedure.-The apparatus design and procedure were the same as described by Brown, et al.1 A 20-gage hypodermic needle was used for all the measurements reported here.
Results and Discussion Typical plots of bubble pressure versus concentration of detergent for both pure and commercial grade sodium dodecyl sulfate are shown in Fig. 1 . (2) L. M. Kushner, B. C. Duncan and J. I. Hoffman, J . Research
Natl. Bur. Standards. 49,No. 2, 85 (1952),RP 2346. (3) Obtained from the Fisher Scientifio Co., Silver Spring, Md., under the label, Sodium Lauryl Sulfate, U.S.P. (4) L. Shedlovsky, Ann. N. Y. Acad. Sci., 46, 427 (1946).
t
