370
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Vol. 62
Acknowledgments.-The support of this research by the National Science Foundation, Washington, D. C., is gratefully acknowledged. Distillations of the solvents were carried out by Miss Dolores Sicilia.
mentioned definition of the effective hydrodynamic volume is comparable with the value obtained from self-diffusion experiments. Hence, we re-affirm the statement3 that “it is impossible to relate the effective hydrodynamic volume of a dissolved protein molecule to its partial specific volume.” Tanford agrees that two hydrodynamic properON T H E INTERPRETATION OF ties have to be measured in order to deduce both HYDRODYNAMIC DATA FOR DILUTE the size and the shape of the effective hydrodynamic PROTEIN SOLUTIONS particle. However, he tries t o imply a greater senhitivity in the interpretation of hydrodynamic BYHAROLD A. SCHERAUA AND LEOMANDELKERN data than actually exists by stating that our 0 (and Department of Chemistyy, Cornell University, Ithaca, N e w Y o r k , and presumably also our 6) function represents a poor Polumer Structure Sectzon, Natzonal Bureau of Standards, Washington choice of measurements. As emphasized previ36,D. C. 0us1y,~Jwhereas a single property (such as viscosRereived September SO, 1967 ity) varies considerably with axial ratio for constant Several questions have recently been raised’ volume, one doesn’t know in advance what the volabout our method of interpretation of hydrody- ume is. Hence, one must use a function (e.g., our namic data on dilute protein solutions,2and about a or 6 function, or any other equivalent one) which recent application of this method to measurements pdepends on a pair of hydrodynamic measurements. on bovine serum a l b ~ m i n . We ~ believe that the A few calculations will show that any such funcargument developed by Tanford,’ in this connec- tion, which combines two hydrodynamic measuretion, is misleading and thus necessitates further ments, is very insensitive to changes in axial ratio. clarification. As stated in his equation 1, Tanford chooses to express the effective hydrodynamic volume of a dis- EXTRACTION OF INORGANIC SALTS BY solved protein molecule in terms of its partial 2-OCTANOL. 111. ZINC AND CADMIUM specific volume and a quantity 61,which he defines4 as “the number of grams of solvent incorporated CHLORIDES. AQUEOUS PHASE ACTIVITIES‘ in the hydrodynamic particle per gram of dry protein.” The arbitrariness of this assumption and its BY T. E. MOORE,NORMAN G. RHODEAND ROBERTE. disregard of physical reality have already been disWILLIAMS cussed in great detail both by us2 and by Sadron15 The Department of Chemistry, Oklahoma State University, Stallwater, Oklahoma whose treatment is essentially equivalent to ours. Received October 10, 1967 In addition to this arbitrary division of the effective hydrodynamic volume into two terms ( M / Preliminary experiments in these laboratories N ) G and 6I(M/N)vlo,Tanford’s procedure is also have shown that the extraction of both Zn(C104)z very misleading since one is thereby tempted to and Cd(C104)2from aqueous solutions (4 m) ocattach reality to 6’ as the mass of water actually curs readily. When solubilities of ZnC12and CdClz bound to one gram of protein. Tanford, in fact, is in 2-octanol are compared? however, a large difinconsistent on this point since he makes this lat- ference is found, ZnCL being over 1000 times as ter identity when he asserts’ that the validity of his soluble at 25”. This suggested that effective sepaequation 1 is confirmed by Wanq’s considerationsB ration might be achieved through the 2-octanol of the self-diffusion of water in dilute aqueous pro- extraction of aqueous mixtures of the chlorides. tein solutions. It is incorrect to obtain 61 from This is in general agreement with earlier observaself-diffusion since it is clearly stated by Wang and tions regarding the non-specificity of 2-octanol as quite apparent in his theoretical development that an extraction solvent for metal perchlorates conthe hydrodynamic behavior of the dissolved pro- trasted to the much more specific behavior of the tein molecule does not enter into his calculation. corresponding chlorides.2 Thus, it is not surprising that in Wang’s treatment To test this theory, six series of solutions were the appropriate and correct volume to be consid- extracted with the octanol at 25”. Figure 1 preered is that which describes the domain of the mole- sents the variation of the distribution coefficients, cule, with &, in this instance, being the specific kd, of ZnClz and CdClz in the different series. The solvation of the actual protein molecule. However, distrlbution coefficient is here defined as the ratio this latter conclusion is limited to problems involv- of the molal concentration in the non-aqueous phase ing the self-diffusion of water and has no applica- to the molal concentration in the aqueous phase. tion to the present matter concerned with the in- The equilibrium mixtures of octanol and water are terpretation of the hydrodynamic data of protein considered as solvents in each phase. solutions. A similar error is committed by Tanford It is evident from the figure that separation and Buzzell,4 who assume that the value of 61 ob- factors, s, of the order of 50-60 (s = kd(ZnClz)/ tained from intrinsic viscosity data and the afore- kd(CdClz)) are obtained with the ZnClz-CdCln (1) C. Tanford, THISJOURNAL, 61, 1023 (1957). mixtures investigated. These values, however, (2) H.A. Scheraga and L. Mandelkern, J . Am. Chem. Soc., 76, 179 are only about 50y0 of the values calculated from (1953). (3) G. I. Loeb and H. A. Scheraga, THISJOURNAL, 60, 1633 (1956). (4) C. Tanford and J. G. Buzsell, THISJOURNAL, 60, 225 (1956). (5) C. Sadron, Prou. in Biophys.. 3, 237 (1953). (6) J. H. Wan& J . Am. Chem. Soc., ‘76,4755 (1954).
(1) Supported under Contract AT(11-1)-71 No. 1 with the U. 8. Atomic Energy Commission. (2) T.E. Moore, Roy J. Laran and Paul C. Yates, THISJOURNAL, 19, 90 (1955).
I
. I
1 I
a*. ?
t
c
*.
March, 1958
.
the ratio of the distribution coefficients in the binary solutions, the experimental results showing a definite mutual extraction-promoting effect of ZnClz on the extraction of CdClz and of CdClz on the extraction of ZnClz. Whereas this mutual promoting effect (curves C and C‘) may be regarded as normal, the existence of maxima and minima in the distribution coefficient curves (for the mixtures with CaCI2, curves D and D’) has not been previously observed. In order to understand better the interactions which are responsible for the concentration dependence of the distribution coefficientsin these systems, the activity of CdC12 in 1 m CdC12 mixtures with ZnCla and the activity of ZnClz in 0.5 m ZnClz mixtures with CdClz were measured. Figure 2 shows ratios of the stoichiometric CdC12 activity coefficient in the mixtures to the activity coefficient of a 1 m CdCL solution. In the same figure there are shown also the non-aqueous-phase acvitity coefficients calculated from the aqueousphase activities and the distribution coefficients.a These are expressed relative to the octanol-phase activity coefficient of CdCL calculated for zero molal ZnC12. Unfortunately hydrolysis and the, precipitation of a solid phase prevented measurements at concentrations above 1.5 m ZnCL; it is evident, however, that over the concentration range up to 1.5 m ZnClz the observed increases in the CdClz distribution coefficient can be largely attributed to the increases in the aqueous-phase activity coefficient of CdC12,arising probably from the effects of the ionic hydration of ZnClz.4 From the constancy of the octanol-phase activity coefficients it is concluded that, as in the case of its aqueous solutions, CdCL is only slightly dissociated in octanol. The unusual shape of the distribution coefficient curve for ZnClz in mixtures with CaClz (Fig. 1) is better understood by reference to Fig. 3 where the aqueous-phase activity of ZnCl2 is plotted as a function of the CaClz concentration. Qualitatively, the initial rise in the activity curve can be interpreted as showing the effect of the common chloride ion and of ionic hydration. At higher concentrations of CaC12,however, ionic association reactions leading to the formation of ions such as ZnClf and ZnCL+ would cause the activity to reach a maximum and then fall to smaller values as the reactions become more complete. The literature value for the dissociation constant for the 1:1 chloro complex shows that it is but a weak complex ( K = 0.65), and since the 1:1 complex is usually the most stable, higher complexes would be of importance only at high CaClz/ZnC12ratios in concentrated solutions. The minimum at about 4 m CaC12 and the subsequent rise in the activity values then follow from the removal of additional solvent by CaC12hydration at still higher concentrations of CaC12. The similarity of curve D of Fig. 1 and the curve for the aqueous-phase activity (3) T.E. Moore, R. W. Goodrioh, E. A. Gootman, B. S. Slezak and P. C. Yates. ibid., 60, 564 (1956). (4) R. H. Stokes and R. A. Robinson. J. Am. Chem. Soc., 7 0 , 1870 (1948). ( 5 ) R. A. Robinson and R. 0. Farrelly, Tmo J O U R N A L , 51, 704 (1947).
NOTES
371
-1
-2 e 4 M
I4 -3
-4
-5 I
I I I I 2 4 6 8 Aqueous phase molality. Fig. 1.-Distribution coefficients: A, ZnClz 1 m CdClz; B,ZnCl2; C, 0.5 m ZnClz CdC12; D, 0.5 rn ZnClz CaClz; A’, CdClz 0.5 m ZnCl2; B’, CdCL; C’, 1 m CdCh ZnClz; D’, 1 m CdClz CaC12. 0
+ +
+
+ +
+
0.5 1.0 1.5 ZnClz molality. Fig. 2.-Activity coefficients of CdCll in mixtures with ZnC1,: 0,aqueous-phase coefficients relative to 1m aqueous CdClz; 0 , calculated octanol hase coefficients relative to that of an octanol solution of &Clz in equilibrium with 1m aqueous CdCL
0
of ZnCL (Fig. 3) is obvious and suggests a close correlation between the aqueous-phase activity and the distribution coefficient. It appears likely, therefore, that the anomalous behavior of the distribution coefficient in this system arises principally from aqueous-phase interactions of the kind described. A similar explanation should hold for the effect of CaClz on the extraction of CdC12. The results of these experiments further emphasize the difficulties in developing a really compre hensive theory for the extraction of metal halides in concentrated solutions. DiamondG has attempted this but has nbt quantitatively dealt with the effects of solvation in the aqueous-phase. (6)
R. h l . Diamond, i b z d . , 61, 69 (1957).
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Vol. 62
THE SURFACE STRUCTURE OF SODIUM CHLORIDE BYL. G. HARRISON’ AND J. A. MORRISON Diuiszon o f Pure Chemistry, National Research Laboratories, Ottawa, Canada Receiued September 26,1967
Studies of isotopic exchange between gaseous chlorine and small particles of sodium chloride already have been reported.2~3 Exchange of the surface layer of the solid obeyed first-order kinetics and proceeded rapidly at room temperature (velocity constant -0.012 mh-1). The energetics of the reaction suggested that the surface of the 0 2 4 6 8 particles differed markedly from a normal crystal CaC12 molality. Fig. 3.-Activity of 0.5m ZnClp in mixtures with CaClz. plane; however the possibility that their surface structure was strongly affected by adsorption of Experimental atmospheric gases during preparation and storMaterials.-All chemicals were C.P. or reagent grades. age had to be considered. Accordingly, a similar The 2-octanol was the best grade furnished by the Matheson investigation has been begun with films of sodium Coleman and Bell Co. Extraction Procedures.-The equilibrations were made chloride prepared in vacuo and immediately exfrom approximately equal weights of the aqueous solutions posed to carefully purified chlorine. Preliminary and octanol sealed in glass-stoppered flasks and mechani- results indicate that exchange of the surface layer cally agitated overnight at 25.0 f 0.1”. The phases were still occurs very readily, but that the first-order separated and analyzed. Analytical Procedures .-The analytical procedures were process is preceded by an extremely rapid exchange for the most part standard. Octanol phases were back- (within the first minute or so) amounting t o about extracted with water and zinc and cadmium determined haIf the totaI reaction. This type of behavior had polarographically either alone or together. Chloride was been observed in some of the earliest experiments on determined in the back-extracted samples volumetrically particles prepared and handled in dry nitrogen and with AgNOa. The aqueous phases were similarly analyzed, calcium dry air, and was attributed to impurities in the being found by difference. chlorine gas.2 However, re-examination of these Activity Measurements.-The activities of both CdC12 and ZnClz were determined by electromotive force measure- results in the light of the current experiments sugments. The cell was a conventional H-type cell with the gests an alternative explanation.
metal amalgam electrode in one arm and the Ag-AgC1 electrode in the other. The latter was prepared by the thermal reduction of AgSO according to the recommendations of Taniguchi and Janz.7 Interagreement among the electrodes was of the order of 0.05 mv. The zinc amalgam electrode was an amalgam pool (ca. 2.5%). The e.m.f. values of the cell a t equilibrium were very constant, varying by not more than 0.03 niv. over a period of several hours. The cadmium amalgam electrode (ca. 5.5%) also was a pool electrode and again the e.m.f. values at equilibrium were quite constant, the average deviation being only about 0.03 mv The eIectrolytic preparation of the amalgams followed the directions of Richards.* In order to determine the standard potentials of the cells, the electromotive forces of standard solutions of ZnClz and of CdC12were measured. For ZnCl, the eoncentrations were 0.4320 and 1.031 m , and for CdCI? they were 0.0500 and 0.2325 m. Zinc chloride activity values were taken from those listed by Robinson and StokesQ; CdClz activity values were those given by Harned and Fitzgerald.’O The Eo values a t the two reference concentrations differed by 0.06 and 0.2 mv. for ZnClz and CdC12, respectively. A potential of 0.2224 v. was used for the Ag-AgC1 electrode. All solutions were prepared by weight from stock solutions whose concentrations had been determined electrolytically. Anhydrous CaClz was added to the ZnCl solutions to obtain the mixtures of CaC12 and ZnClz. Oxygen was removed from the solutions by a stream of oxygen-free nitrogen before the measurements were made.
.
(7) H.Taniguchi and G. J. Janz, “The Preparation and Reproducibility of the Thermal-Electrolytic Type Silversilver Chloride Electrodes,” ARDC Contract No. A F 18(600)-333,Technical Note No. 3, 1955. (8) T. W . Richards, “Electrochemical Investigation of Liquid Amalgams of Thalium, Tin, Zinc, Cadmium, Lead, Copper and Lithium,” Carnegie Institution Publication, Washington, D. C., 1909. (9) R. A. Robinson and R. H. Stokes, Trans. Faraday Sac., 36, 740 (1940). (10) H. S. Harned and M. E. Fitzgerald, J . A m . Chem. Sac., 68, 2624 (1936).
Experimental Chlorine gas was circulated by convection in a closed system consisting of the jacket of a Geiger counter and a 15 mm. diameter Pyrex tube upon which the sodium chloride films were deposited. The counter was connected to a counting rate meter and graphic recorder. The sodium chloride (0.1 to 0.4 g.) containing radioactive aaCl was evaporated into the tube from a platinum crucible which was held on a quartz carriage and was heated by an inductlon furnace. The system, with the crucible in situ, was previously baked out at 400” and evacuated to -lo-* mm.; the pressure was read with a Bayard-Alpert ionization gage. The evaporation took about an hour and the pressure never exceeded mm. The crucible was then withdrawn by a magnetic device and the system sealed off (maximum pressure -10-6 mm.). The chlorine gas, purified by two distillations (at 760 and -1 mm.) and stored over a film of sodium chloride, was admitted through a break seal to give a pressure between 10 and 40 cm. The exchange reaction was followed for a t least 24 hours, and the vessel was thereafter connected to an adsorption apparatus. The surface areas of the films, estimated by adsorption of nitrogen at 77’K., were -10 m.2/g.; the surface present in an experiment was 1-4 m.9. The maximum amount of gas available to contaminate the film was estimated from the pressures during evaporation and sealing and the pumping speeds in the system.
Results and Discussion Figure 1 is a typical record of the radioactivity of the gas phase, C (the background has been subtracted), showing an initial very rapid reaction up t o C1 and a subsequent process which obeys the first-order law (1) National Research Laboratories Postdoctorate Research Fellow. (2) L. G. Harrison, J. A. Morrison and G. S. Rose, Second International Congress of Surface Activity, London, April, 1957. (3) L. G. Harrison, J. A. Morrison and G. 8. Rose, THISJOURNAL, 61, 1314 (1957).
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