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LOYSJ. NUNEZ AND M. C. DAY
Vol. 65
THERMODYNAMIC STUDIES OF HYDROBROMIC ACID I N ANHYDROUS ETHANOL1 BY LOYSJ. NUNEZ AND M. C. DAY Coates Chemical Laboradones, Louisiana State University, Baton Rouge, La. Received August 4* 1060
Electromotive force measurements were made a t 25” on the cell without liquid junction Pt, Hz( 1 atm.) IHBr(m) ethanol] AgBr, Ag. The standard electrode potential of the silver, silver bromide electrode in anhydrous ethanol was determined to be -0.1939 v . on the molar concentration scale and -0.1816 v. on the molal concentration scale. Mean molal activity coefficients are tabulated and the &valuefor hydrobromic acid in ethanol is determined. A new method for the preparation of silver-silver bromide electrodes is reported.
Introduction Although there have been several determinations of the standard electrode potential of the silversilver chloride electrode in anhydrous ethanol2 and the standard potentials of both the silver-silver chloride and silver-silver bromide electrodes in anhydrous methan01,~~~ there have been no comparable studies of the silver-silver bromide electrode in ethanol. From both a practical and a theoretical point of view the potential of the silversilver bromide electrode in ethanol should be of interest. The greater solubility of the bromide salts of the alkali metals in ethanol compared to the corresponding chloride salts would indicate the possibility of a broader range of concentration studies. And in terms of simple solution concepts, a knowledge of the standard potential might lead to some general correlations. For this study the cell without liquid junction Pt, HE(1 atm.) IHBr(m), ethanol1 AgBr, Ag
was used to determine the standard potential of the silver-silver bromide electrode and to evaluate the activity coefficients of hydrobromic acid. Experimental Chemicals.-!rhe anhydrous ethanol was prepared using a modification of the Bjerrum Method.& In order to remove trace uantities of benzene, 2% water by volume was added to U1.I. absolute ethanol, U. S. P. grade, and the mixture was fractionally distilled. When the benzene had been removed as the benzene-water-ethanol azeotrope, magnesium alcoholate prepared from dry ethanol and magnesium was added in the ratio of 1 g. of magnesium alcoholate to 6.5 ml. of benzene-free ethanol. This solution was refluxed until free of water and was then immediately fractionally distilled. The middle fraction was taken as the pure anhydrous ethanol. All joints in the distillation apparatus were glass to glass with no stopcock grease being used. Both distillations were monitored by means of the Beckman Model D.K. Spectrophotometer. Using this method of detection, the maximum weight per cent. of benzene present should be less than 0.008 and that of water should be less than 0.02. The anhydrous HBr was obtained from the Matheson Co., Inc., and all HBr solutions were prepared immediately before their use to minimize reaction between the ethanol and theHBr. All ethanol and ethanol solutions were kept in a nitrogen dry box with PzOsused as a desiccant. Electrodes.-The silver-silver bromide electrodes were (1) Taken from ,t portion of the thesis submitted to the Louisiana State University by Loys J. Nunez in partial fulfillment of the requirements for the degree of Doctor of Philosophy. (2) H. Taniguchi and G. J. Janz, J. Phys. Chem., 61, 688 (1957). P (3) R. A. Robinson and R. H. Stokes, ”Electrolyte Solutions,” Butterworths Scientific Publications, London, 1955, p. 457. (4) E. W. Kanning and A. W. Campbell, J. A m . Chem. SOC.,64,517 (1942). ( 5 ) H.Lund and J. Bjerruru Ber., 64B,210 (1931).
prepared by sealing a silver wire of 1 mm. diameter in one end of a section of glass tubing having a slightly larger diameter than the silver wire. Approximately one cent;meter of silver wire was allowed to protrude from the glass tube a t both ends, one end serving as a means of electrical contact to the external circuit and the other end serving as the surface of the electrode. One end was then heated in a flame until molten, and a spherical droplet was allowed to form. After cooling, a paste of silver oxide was applied, and this was heated until a spongy surface of metallic silver had covered the spherical surface. The electrode was then anodized in a dilute KBr solution as usual until a layer of AgBr had covered the surface of the electrode. All silversilver bromide electrodes used in this investigation were prepared in this manner. Within the limitations of our experimental mebhods, they gave results identical to those prepared by the usual thermal electrolytic method.6 However, they showed greater stability in the ethanol solvent. The hydrogen electrode was a commercially available one of the Hildebrand type. The electrode was replatinized every two or three runs or whenever the system showed signs of instability. All electrodes were stored in anhydrous ethanol. Procedure.-The cell was constructed so as to permit the use of two silver-silver bromide electrodes and one hydrogen electrode. For all measurements, the cell was placed in a water-bath thermostated a t 25.25 Z!Z 0.05’. The criterion for equilibrium was a stable reading for a period of one hour. This usually required several hours after the initiation of hydrogen bubbling. Before introduction into the ccll, the electrodes were soaked for approximately one hour in an ethanol solution of the same HBr molality as that used for the given run. Prepurified hydrogen, sold by the Matheson Co., Inc., was passed first through a hydrogen catalytic ,purifier. It then went through a silica gel column and into a bubbling tower filled with a solution of identical composition as that present in the cell. Measurements were made a t random choices of the concentration. After completion of a run, aliquots of the ethanol-HBr solution were withdrawn from the cell and titrated with a standardized solution of NaOH using phenolphthalein as an indicator. All concentrations should be accurate to within 0.02%.
Standard Electrode Potential.-Assuming the existence of extensive ionic association, we chose t o use the extrapolation procedure7based on the equation (Eo’- E a ) = E + 2k log CUY, C = EO (1) Here (Y is the degree of dissociation of HBr in ethanol and ya is the molar activity coefficient for an ionic concentration of aC. The relation between a! and ye is given by
where K is the ionization constant for HBr in ethanol. The determination of Eo involves, then, (6) 11. Taniguchi and G. J. Jana, J. Electrochem. SOC.,104, 123 (1957). (7) H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolytio Solutions,’’ 2nd Ed., Reinhold Publ. Corp., New York, N. Y.,1950.
Jan., 1961
THERMODYNAMIC STUDIES OF HYDROGEN BROMIDE IN ANHYDROUS ETHANOL
165
the evaluation of a and ua at various concentrations and the extrapolation of a plot of (EO' - Eu) 120 A. vs. HBr concentration to infinite dilution. I n the region where this approach is valid, a straight line of zero slope should be obtained that intersects the axis at EO. I n order to obtain a and ya, a first-order approximation of yu is obtained by calculating y~ from the Gronwall, La Mer and Sandved extended terms of the Debye-Huckel theory using a particular value of the ion size parameter 8. Once a yu has been approximated, an a may be determined from eq. 2. This a is nom used to recalculate y~ -40 by again determining y* from the extended terms -eo! 24 20 16 12 28 expression a t a new concentration of a C. This - l o g m. method of successive approximations is continued Fig. 1.-Corrected e.m.f. of the cell as a function of -log until a final a and YU are obtained. mHBn The observed potentials for the cell, corrected to a hydrogen pressure of 1 atm. are presented in Table I and are summarized in Fig. 1, where they are plotted as a function of the log of the HBr molality.
i
t
TABLE I CORRECTED ELECTROMOTIVE FORCE DATAOF THE CELL Pt, Hz(l atm.) IHBr(m)Ethanoll AgBr, Ag Molal
Molar
conen.
concn.
E, v.
0.00291 .00325 .0103 .01510 .02300 .0290 .0364 .0440 .0530 .0631 .0682 .0881
0.00229 .00255 ,00807 .0119 .OB04 ,0228 .0286 ,0346 ,0416 .0496 .0536 .0692
0.1363 .1330 ,0824 .0630 .0454 .0351 .0211 .0137 .0097 ,0010 ,0011 .0134
1901
4
Fig. 2.-Extrapolation
8
16 l2 moles liter-! x103.
20
-
of (Eo' Ea)in terms of molar concentration.
-
The empirical equation for the curve within the concentration range studied may be expressed as E = -0.1021 log m - 0.1221 (3) where m represents molal concentration. The determination of the molalities was based upon the following expression for the density in g./ml. of the solutions a s a function of HBr normality D = 0.07993N + 0.7855 (4) Extrapolations based on 12 values of 4.0,5.0,6.5 and 8.0A. were made. These are given in Fig. 2, 0 2 4 6 8 IO 12 14 16 and the pertinent calculations for the extrapolation -& x1 4 based on an ion size of 5.0 A. are given in Table I1 Fig. 3.-Evaluation of the dissociation constant of HBr in using selected points based on eq. 3. Of these four ethanol from the data of Goldschmidt and Dahll. choices, a value of 5.0 8. for the ion size parameter gave the only curve approaching a zero slope, and TABLE I1 it approached this slope only a t concentrations less DETERMINATION OF (E"- E a ) ON THE MOLARCONCENthan 0.003 M . Using this extrapolation, the Eo TRATION SCALEFOR a = 5.0 A. value for the Ag, AgBr electrode has been taken C Eoato. - (E'' - Ea) (moiefl.) (V.) Ya (-7.) as -0.1939 v. on the molar concentration scale and -0.1816 v. on the molal concentration scale. 0.0009 0.1783 0.970 0.818 0.1939 The reliability may be estimated to be 10.5 mv. .0016 .1528 ,950 ,772 .1939 Dissociation Constant of Hydrogen Bromide..0025 .1330 .940 .730 .1942 In order to use the extrapolation procedure chosen .0036 .1168 .694 .1953 ,920 for this work, it is necessary to know the thermo.0049 .lo31 .910 .661 .1963 dynamic dissociation constant of HBr in ethanol. .0064 .0913 .895 .633 .1974 For this purpose, we have used the conductivity .010Q .0715 .870 .584 .1999 01
166
L. C. CRAIGAND W. KONIGSBERG;
Vol. 65
data of Goldschmidt and Dahlls along with the TABLE I11 method of proposed by boss and DETERMINATIONOF THE IONIZATION CONSTANT OF HBr IN K r a ~ s . This ~ method should be valid as long as ETHANOL ionic interactions greater than pairwise are negC CA2/ ligible. The: relation between the conductance wZ) m 2 ’(?’ F(Z) and the diseociation coilstant can be expressed as o.oo625 o.1624 0.82073 0.911 o.359 O.o13O1 0.l723 (5)
where A and A0 are the equivalent conductance and equivalent conductance a t infinite dilution, respectively, and F ( 2 ) is a function defined and tabulated by Fuoss. If P ( Z ) / A is plotted against y a 2 C h / F ( Z ) , a straight line of slope 1/KAo2and intercept l / A o should be obtained. The results of these calculations on the data of Goldschmidt and Dahll are given in Table 111,and the corresponding plot is seen i n Fig. 3. These lead to a value for the dissociation constant of 0.0187 mole/l. for HBr in ethanol a t 25’. This is a quite reasonable result in the light of the accepted value for the dissociation constant for HC1 in ethanol1° of 0.0113 mole/. Activity Coefficients of Hydrobromic Acid.(81 H. Goldschmidt and P. Dahll, 2. p h w i k . Chem., Leipaig, 114, 1 (1925). (9) R. M.Fuass, J . Am. Chem. Soc., 67, 488 (1935); R. M. Fuoss and C. A. Kraus, ibid., 66, 476 (1933). (10) I. I. Bezman and F. H. Verhoek, ibid., 67, 1330 (1945).
.00312 ,00156 .000781 .000391
.1194 .OS70 .0629 .0453
,87214 ,90873 ,93495 ,95362
,923 .947 .959 .976
,483 ,593 .690 .767
,01284 .01250 ,01235 .01213
,1174 .0737 ,0436 .0247
The molal activity coefficients were determined from the equation
in which y * is defined such as to approach unity at infinite dilution in ethanol. Values of y A are listed in Table IV, for convenient choices of the concentration. TABLE IV MEAN MOLALACTIVITYCOEFFICIENTOF HYDROBROMIC ACIDIN ETHANOL Molal canon.
0,005 .Ol .02 .03
Molal
Y t
0.649 ,590
,537 ,507
concn. 0.04 .05 .07 .lo
T*
0.488 ,474 .452
.431
DIALYSIS STUDIES. 111. MODIFICATION OF PORE SIZE AND SHAPE I N CELLOPHANE MEMBRANES BYL. C. CRAIGAND WM.KONIGSBERG Rockefeller Institute Laboratories, New YO&,N . Y . Received Auwst 6 , 1960
A search has been made for various ways of modifying the porosities of cellophane dialysis membranes. Mechanical stretching, acetylation and zinc chloride treatment have been found effective. Optimum porosities for studying solutes of molecular weights varying from the size of dipeptides to proteins of molecular weight over 100,000can be made a t will.
Introduction Previous :membrane diffusion studies from this L a b ~ r a t o r y l -have ~ shown considerable promise in the use of this approach as a tool for separation and characterization of solutes with respect to their shape and size. Many factors, such as porosity of the membrane, temperature, pH, solvent, ionic strength, eta., were found to influence more or less the rate of diffusion of a particular solute through a membrane but perhaps even more interesting, the relative rate of diffusion of one solute as compared to another. The various factors showed considerable interdependence but the role of each factor could be revealed largely by comparative studies with solutes of known size and shape under standardized conditions. It soon became evident in our studies4 that the (1) L. C. Craig and T. P. King, J . A n . Chem. SOC.,17, 6620 (1955). (2) L. C. Craig, T. P. King and A. Stracher, ibid., 79,3729 (1957). (3) L. C . Craig, Wm. Konigsberg, A. Stracher and T. P. King, in Symposium on E’rotein Structure, p. 104,A. Neuberger, Ed., IUPAC Symposium July 1957,John Wiley and Sons, Inc., New York. (4) This investigation was supported in part by a rasearch grant, A-2493 B.B.C., from $he National Institute of Arthritis and Metabolic Diieaiei of the National Inrtitutea of Health, Publio Health Service.
greatest selectivity in distinguishing one solute from another or of the effect of change of one of the factors, such as temperature, was achieved when the solute of interest was barely able to diffuse through the membrane. It, therefore, became highly desirable to learn how to change the porosity of the membrane without altering appreciably its properties in any other way. Mechanical stretching seems to offer one possibility. When a wet cellophane casing is slowly stretched to near the bursting point by hydrostatic pressure, it will only partially return to its original size on release of the pressure. Such an enlarged casing will have a thinner wall but more important will permit those solutes which would barely diffuse through before treatment now to pass through at a much accelerated rate. On the other hand, when a wet inflated casing with no hydrostatic pressure inside is stretched longitudinally to near the breaking point and then released, the tubing will return again only partly to its original size. It will have a permanently decreased diameter and those solutes which diffused through the unstretched membrane a t a reasonable
