Solubilities of Salts in Dioxane--Water Solvents

between dielectric constant and solubility of elec- trolytes, we have ... range of concentrations from 0 to 100yo dioxane ... and 0 X Short 7btd , 68,...
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Solubilities of Salts in Dioxane-Water Solvents' RY

T. W.



As a study in the problem of the relations between dielectric constant and solubility of electrolytes, we have made measurements of the solubilities of three insoluble compounds, barium iodate, silver acetate and silver sulfate in dioxane-water mixtures a t 25" over the complete range of concentrations from 0 to 100yodioxane in intervals of 10% by weight. The measurements extended over a dielectric constant range from 2.10 to 78.55 (the widest variation so far reported in solubility work) and the results have been treated in the light of the interionic attraction theory and of Born's theory of solvent interaction. I. Materials and Methods Barium ;Iodate.-A c. P. grade of barium iodate monohydrate was washed eight to ten times with distilled water, sedimented in tall cylinders to remove smaller sized particles, and dried a t 100" before use. Iodometric titration gave 96.5% Ba(IO& compared with 96.43% theoretical for the monohydrate. The several preparations gave a constant sohbility in water. Silver Acetate.-A c P product was used without further treatment. Volumetric analysis for silver gave 64.72% silver compared'with a theoretical 84.64%. Silver Sulfate.-A preparation from Merck was employed without further purification. On the basis of silver content, the salt was 99.9% Ag2S04. dard c. P grade chemicals Dioxane.-This solvent came from several sources The first step in the purification was usually either a rcfluxing with metallic sodium for several hours or treatment with sodium wire a t room temperature for several days This sodium treatment was then followed by distillation through a 24-inch (61-cm.) packed distilling column. In some cases the dioxane was first treated with alkaline permanganate, distilled, refluxed with lime and then redistilled before being treated with sodium: The density, refractive iudex and dielectric constant of the preparations checked closely with values to be found in the literature. The freezing points ran between 11.5 and 11.73", somewhat lower than those reported by several others$!' (about 11.8). Such d8erences have been neglected in the present work since the effect on the calculated dielectric constant of the mixed solvents would be very small The rapid deterioration of the material i p ascribed by He55 and he Baltimore nieetmg of t h e American Chemical ( 2 ) M Born Z Physsk. 1,4: (1920) (3) (a) C A Kraus and R &I Fuoss, THISJ o U R N A r , 55, 21 (1033), K A Tingee, A r d , 66, ,511 (1014) (b, G .kkedof and 0 X Short 7btd , 68, 1?41 10111) (4) K . Ire55 arid €1 lirahru 7'0r 1 I B 2 2 : 1 >?-, C A k'raus aiid

Frahm" to oxidation to give an aldehyde, peroxide and other subsequent products. Because of the unstable character of the liquid, it was prepared at frequent intervals during the course of the work. Solubility Determinations.-The dioxane-water mixtures were made up by weight for each experiment, in glass-stoppered Pyrex solubility bottles of 250 ml. capacity. The liquid was rotated a t 25 * 0.02" with excess solid for one or more days. Samples for analysis were withdrawn by suction through quantitative filters into calibrated 100-ml. pipets after allowing some time for the undissolved salt to settle There was no evidence of colloidal character in any of the solutions. The silver-containing samples were titrated, after appropriate dilution in the case of high dioxane ratios, against standard 0.01 N potassium thiocyanate sohtion using weight burets. Some difficulty was encountered in fixing the end-point of this titration in the presence of varying quantities of dioxane; the precipitation of the silver thiocyanate is apparently slow near the equivalence point, and the color is also somewhat abnormal. Very vigorous and prolonged agitation i s necessary to overcome the first difficulty, and comparison solutions were used as color standards to fix the final end-point in very high dioxane concentrations. 'The iodate solutions were first evaporated to dryness to expel the dioxane which seriously interferes with the observation of the starch-iodine end-point, and which accelerates the liberation of iodine from an acidified potassium iodide solution. The titration of the iodate was performed with a carefully regulated procedure, using frequently standardized thiosulfate solutions delivered from weight iiuicf 5, appropriate corrections being made for the blanks.

11. Solubilities in Presence of Dioxane The solubility figures appearing in Tables 11, I11 and I V are in every case the mean of one or more determinations from undersaturation and one or more from supersaturation. In Table I we have collected the data on the properties of the solvents which are useful in interpreting our solubility figures. The dielectric constants are taken from the report of Akerlijf and Short3b; the densities are interpolated from the measurements of Hovorka, Schaefer and D r e i s b a ~ h . ~The coefficients to be used in the Debye-Hiickel equation a t 25' are listed in the last two columns of the table. Column 5 of the table contains the differences between the reciprocal of the dielectric constants of the several media and the reciprocal of the dielectric constant of water. ( 5 ) F. Hovorka, R.A Schaefer and D Dreisbach, THISJ O U R h A L , 58, "304


Dec., 1939





media of different molecular composition is to use mole fractions as the more fundamental Wt. % 382.6 224 units. Scatchard6 has shown, for example, that diox- DenMole % Dielec. the reaction of thiosulfate ion with bromoacetate 78.55 0.000 0.505 0.3288 0 0 0.9970 ion in mixed solvents7is in better agreement with .GO6 .3490 69.71 .001614 10 1.0030 2.223 4.667 50.81 ,003714 .744 .374 20 1.0097 prediction when the volume concentrations are ,006533 ,943 ,404 8.063 51.91 30 1,0189 a mole fraction basis. Bell and corrected to 1.250 .444 43.00 .010525 40 1.0263 12.004 1.757 34.28 .01644 50 1.0311 16.987 .498 Gatty,8 however, have come to the conclusion on 25.86 .02594 2.681 .573 60 1.0345 23.485 theoretical grounds that volume concentrations 4.627 .693 17.70 .04377 70 1.0372 32.32 45.01 10.72 ,08055 10.04 .890 80 1.0349 should be used in calculations of interaction en15.12 1.020 8.16 .lo98 85 1.0320 53.68 ergies in solution. If this is correct, then all the .1654 26.54 5.61 1.230 90 1.0295 64.81 3.69 .2583 49.73 1.517 95 1.0280 79.52 data should appear in volume concentrations. 115.9 2.101 .4637 2.011 100 1.0277 100.0 We accordingly have listed the solubilities, in TABLE I1 Tables 11-IV, in terms of molarities and of mole SOLUBILITY OF SILVERACETATE IN DIOXANE-WATER fractions both, and have in fact carried out most MIXTURES AT 25" of the subsequent calculations for both sets of Solubility, Solubilit S X, mole units. The same tables also list the figures for wt. % moles/? fraction - log S X lo4 - log X 6 the square root of the ionic strength in the several dioxane X lo4 1.1763 12.027 2.9198 0.2581 0 666.3 solutions calculated on the basis of 100% dissocia9.819 3.0079 1.2980 .2244 503.5 10 tion throughout. 1,4391 7.714 3.1127 .1907 363.8 20 Since the barium iodate is a hydrated salt, its 5.693 3.2446 1.6105 .1550 245.2 30 activity and solubility must depend on the fu3.905 3.4084 .1232 151.7 1.8190 40 2.485 3.6048 .09253 gacity of water in the solutions. In every ex50 85.64 2.0673 1.409 3.8512 .06502 periment the solid phase a t saturation was the 42.27 2.3740 60 0.6818 4.1663 .04171 17.40 2.7595 70 monohydrate, as proved by actual analysis, even .2925 4.5339 .02472 6.11 3.2140 80 when the solvent was dioxane, the stability of the ,1552 4.8090 .01688 2.85 3.5452 85 solid in the last case being a result of the very 1.54 3.8125 .0949 5.0228 .0124 90 .0151 5.8220 .0046 0.21 4.676 95 low solubility of the salt. In no experiment with .0000 ,000 100 barium iodate as saturating phase, or with silver acetate or silver sulfate, was any double compound TABLE I11 with dioxane formed. The activity of barium SOLUBILITY OF SILVERSULFATE IN DIOXANE-WATER iodate monohydrate in solution being proporMIXTURES AT 25" S, X, mole tional to the activity of the water, we would Wt. % moles/l. fraction X lo' - l o g s X 10' - l o g X 6 dioxane find the solubility of the hydrate to be greater in 3.3121 0.2838 .1.5712 4.875 0 268.4 solutions of lower water activity, other factors 3.5194 .2156 1.8097 3.026 10 155.0 being the same. Accordingly all the solubility 3.7535 .1579 2.0802 1.764 20 83.14 figures have been multiplied by the ratio of the .lo98 2.3963 0.9326 4.0302 30 40.15 aqueous pressure from the solution to that from ,07157 2.7688 ,4387 4.3580 40 17.03 ,04245 pure water, using for this purpose the data of 3.2214 ,1747 4.7577 50 6.02 ,0232 3.7455 ,0600 5.2220 1.80 60 Hovorka, Schaefer and D r e i ~ b a c h .The ~ resulting 4.201 ,0247 5.6075 .0137 70 0.63 figure, which is that actually plotted, is the hy4.646 ,0110 5.9582 .0082 80 .23 pothetical solubility of barium iodate monohy.0043 ,00388 6.4110 90 .063 5.201 drate in pure water of dielectric less than 78.55. .00000 100 .0000 The measurements given in Tables 11-IV are The experimental figures on the solubilities of plotted in Figs. 1 and 2. The full curves of the three solutes investigated are presented in Fig. 1 show -log S (molarity) for the three 1 Tables 11, I11 and IV both in moles per liter and salts studied, as a function of D and of (-a1 - E ), in mole fractions, assuming normal molecular G.Scatchard, J . Chem. Phys., 7,657 (1939). weights in the liquid. We have included the . (6) (7) V. K La Mer and M. E. Kamner, THISJOURNAL, 67, 2669 data in terms of both units because there is ques- (1935). (8) R.P.Bell and 0. Gatty, Phrl. Mag., 19 [7],66 (1935). tion as to which should be used in equilibrium (9) F.Hovorka, R.A. Schaefer and D. Dreisbach, THISJOURNAL, studies. The general practice when dealing with 69,2753 (1937).

Vol. 61


S Mol;s/l. x 10'

yt. %


0 10 20 30 40 50 60 70

8.145 4.742 2.526 1.217 0.553 278 140




.003u . UP"

TABLE IV SOI,VBII,ITY OF Ra(IOa)gHzOIN DIOXANE-WATER MXTURESAT 25' X 5 .'.are_ - log s "20 y mole fraction x



- log.?

3 0891 8 3240 3 5976 3.9147 -1 2573 4 5360 4 8539 4.9586 .5. 13ox I; 3010

8.145 -1.687 2.463 1.167 0.520 .z55 184 .!)933 ,0589 .0035



p E70


-5 0891 3 3291 3 6086 3.9329 '4.2842 4.593-1: 4.90-53 5,0302 5.2302 6.4548



0.04943 .03772 .02753 .01911 ,01288 ,00913 .00648 '00574 ,00471 ,00122





4.8161 5.0386 5.2818 5.5669 5.8733 6.1307 6.3826 6.4371 6.5501 7.6651






15.273 9.147 5.226 2.711 1.3386 0.7401 ,4144 ,3855 .282 ,0216

15.273 9.256 5,361 2.827 1.424 0. 8067 4665 ,431 ,351 ,0308



PHtO - log x -


8.0 i










c 1








Fig. 1-Molar solubilities of A (silver acetate), I3 (silver sulfate) and C (barium iodate hydrate) in water-dioxane mixtures.

Fig. 2.--Solubilities in mole fractions of A (silver acetate), B (silver sulfate) and C (barium iodate hydrate) in water-dioxane mixtures.

values for which are given in Table I; Fig. 2 shows the corresponding variations of - log X. Values reported in the literature for the solubility of Ba(IO&-H2O at 25' are

As for silver sulfate, tbe only previous determinations at 25' are of rather old date:

Trautz and Anschutzlo Macdougall and DavieP Polessitskii13 Present investigation

5.7 x 10-4 M 7.895 x 10-4 M 8.10 x 10-4 ~LI 8.33 x 10-4 M P 135 x 10-4 M

The numerous determinations of the solubility of silver acetate were reviewed by MacDougall and Bartsch,l* who themselves found a value of 0.06634 MI as compared with the value here reported of 0.06663 M . (10) M. Trautz and A. Ansehdtz, 2.physik. Chem., 66,236 (1906)' (11) W. D. FIarkins and W J Winninghoff, THISJOURNAL, SS, 1827 (1911) (12) G.Macdougall and C.W. Davies, J. Chcm. S O L , 1416 (1935). (13) A Polessitskii, Compf. rend a d . x i . , U. R. S . S., 14, 193 (1985). (14)F N MacDnugall and C E. Barfsch J Phrs. Chmn., 40, 64'1 (1938)

M Druckerx6 (1901) Rotfimund16(1909) Hill and SimmonsH (1909) [email protected] (191J) Present investigation

0.0257 .0267 .02677 ,02676 02684

IIL Theory of Solubility Born2 has given the following equation for the free energy of charging spheres of radius Y;, in a uniform medium of dielectric constant D,in terms of the charge, &e, of the spheres, e being the charge of an electron

(15) K. Drucker. Z . onorg. Chem., 58,361 (1901). (16) V. Rothmund, 2. physsik. Chcnr., 89, 683 (1909). (17) A. E. Hill and J. P. Simmons, THIS JOURNAL. 31, 821 (1909) (18) W D. Ifarkins,&d., 33, 1807 (1911).

Dec., 1939


For a very dilute solution of an electrolyte considered as completely dissociated into spherical ions of radii r+ and r-, the energy of interaction of the ions with the medium per mole will be the sum of the individual energies E+

+ E-



Nv+ z+’@

D 2r+






+ 2+2



(v+X”)V+( v - X ” ) V = IiTh ( v + x ’ ) ~ + ( v ~ x ’ ) ~ -


( v + e * ~ +v-8-2) 2r the primes and double primes referring to the two different media and the z’s both being taken as v-), positive numbers. On dividing by RT (Y+ equation (4) becomes


The solubility of a slightly soluble electrolyte, however, is also dependent on the concentration of ions in the solution, including those from its own dissociation. The effect of the ionic atmosphere is to decrease the activity coefficients of the ions and hence to increase the solubility; and Born’s equation should be applied only to solubilities which have been corrected for this ion-ion interaction effect. The magnitude of the correction a t 25’ is given by the familiar Debye-Hiickel expression for the logarithm of the mean activity coefficient of the electrolyte


m x + a - f i


s, - 1 +B Af ai f i


Actually the effect of the Debye-Hiickel corrections is rather small, as the two terms, being about equal in magnitude, nearly cancel each other. It proves to be empirically true in many cases that the mean activity coefficient of an electrolyte in saturated solution is almost independent of the dielectric constant of the medium. While no especial mention has been made of this fact, several examples may be found in the literature, and data presented later in this report (Table VIII) illustrate approximately the same principle.14JD~20~21 This regularity and its possible application will be discussed in a later report. Equations similar to (8) (without the general valence factor) have been used by La Mer and Goldmanlg in connection with their study of thallous iodate in alcoholic solutions, and by Bronsted, Delbanco and Volqvartz.22 An equivalent valence factor appears in the equation for equilibrium constants in ionic processes, used by MacDougall and Larson,20while Dunning and Shutt23 use a similar equation for solubilities but with a different valence factor. The equation modified for the interionic attraction effects, as given by the Debye-Hiickel limiting law only, has been used by Flatt and Jordan.24 Bronsted, Delbanco and Volqvartz compared the solubilities of a number of electrolytes in water and in pure organic solvents. The unmodified Born equation was found to be in qualitative agreement with the actual “distribution (19) V. IC. La Mer and F. H. Goldman, THIS JOWENAL, 68, 478


(20) F.H. MacDougall and W. D. Larson. J . Phyr. Chem., 41,417

352 6 logf, =

log s o = log

Combining this equation with the Born equation, we find the solubilities in two different media related by the equation


independently of the concentrations as long as the solution is so dilute that interionic attractions can be neglected. In (2), r represents an averaged value of the two separate radii, and v+ and v- are the number of positive and negative ions, respectively, formed in the ionization of one molecule of the electrolyte. N is the Avogadro number. Combining equation (2) with the ideal relation between free energies, activities and mole fractions, and rearranging the dielectric term, the solubilities in two different media will be related as follows



If the solubility of the electrolyte in a solution of ionic strength p is S,,, then in a solution of zero ionic strength the solubility SOwill be given by an expression of the form


D 2r-

+ v-




1 +D’/rafi

where p represents the ionic strength and “a” the distance of closest approach of the ions (twice the average radius).

(21) P. Wdden, 2. physik. Chem., 66, 683 (1906),reported tetraethylammonium iodide to be about 48% dissociated in raturated solution in many solvents. (22) J. N. Brbnsted, A. Delbanco and K. Volqvarts, 2. physik. Chem., 6164, 128 (1932). (23) W. J. Dunning and W. J. Shutt, Trans. Faradoy Soc., 84,1192 (1938). (24) R. Flatt and A. Jordan,Helw. Chim. A c f q , 16, 37 (1937).

; J t i a> I C -

T, W. r m S



ratio of the ions” or the ratio of the solubility in water to that in the organic solvent; but quantitative tests were not made, and intermediate strengths of the mixtures were not used. The data of Flatt and Jordan were extremely limited (the solubility of the perchlorates of potassium, rubidium and cesium in 0, 50 and 175% alcohol, by volume). With only three points for the solubility curve of a given electrolyte, it is no surprise that the theory was found to be in “splendid agreement” with observation. What we are attempting here is a test of the combined Born and Debye-Hdckel equations on a whole series of media of varying dielectric constant, using different types of electrolytes, and with a sufficient number of dependable determinations to furnish a drastic test of the mathematical predictions.

in the presence of varying concentrations of potassium nitrate, barium nitrate, magnesium chloride and potassium bromate; and that of silver acetate in lO7, dioxane in the presence of potassium nitrate, sodium nitrate and potassium chlorate, and in 20% dioxane in the presence of potassium chlorate. The results are presented in Tables V, VI and 1’11; in these tables, m * is the mean ionic Concentration of the saturating salt, i. e., P+ -k y-j~/i
IV. Determination of “a” Values

Vol. 61


S 14.5

- log mi: 2 8884

8 513 Y 383 11 74 16 43

KNQ 2 8692 2 8270 2 7296 2 5831

x l y X IO4


0 002


a calcd., A.

0,04943 0.06748


In attempting to test the validity of equation o1 .1132 1.64 (S), which we may call the “corrected” Born 05 .2314 1.82 ,4527 2 2.01 equation, the question immediately presents itself as to the correct values of a to be used in the Ba(ru’C& several terms of the equation. 0 000.5 6 9.55 2 8786 0.0599 3.89 001 b 206 2 8676 2.07 ,0697 Ionic diameters for the three solutes examined in .0025 4 963 2 8433 ,0948 1.19 this investigation may be calculated from data on 01 3 426 2 7712 ,1762 1.49 the solvent effects of strong electrolytes. Some 02.5 .2754 1.76 2 854 2 6947 of the data and computations are to be found in M gClz the literature. The values for silver acetate in 0 001023 8 ,596 2 8650 3.53 0.07515 water, alcohol-water and acetone-water solvents, 005069 3.53 9.591 2 8174 ,1345 in the presence of the various electrolytes shown in :3 5:: ,2820 132532 11 99 3 7203 parentheses, given in Table IX below, are taken 3.3) 1014 16 11 2 5922 ,5558 from the publications of MacDougall and [email protected] workers. 14,30,25 That given for silver sulfate in 0 000997 8 3 the presence of potassium nitrate was calculated 002996 8 5 U0504<5 9 0 from some measurements of Harkins on the effect of potassium nitrate on the solubility of silver The figures in Tables VI and VI1 are sumsulfate.18 The value for barium iodate in the marized in Table VIII, which shows S, the obpresence of potassium nitrate was calculated served solubility, SO,the solubility calculated at both from our own measurements and from the zero ionic strength, and f, the mean activity data given by MacDougall and Davies12: that for coefficient ( = So/S). In agreement with the barium iodate in the presence of potassium chlo- regularity pointed out above, the mean activity ride is taken from the measurements and calcula- coefficient of silver acetate in its pure saturated tions of Naidich.zO solution seems to be roughly independent of the To supplement to some extent the data already dielectric constant of the solvent. found in the literature, we have made further The silver acetate solubilities in presence of similar determinations, studying the solubility salts, in various dioxane-water mixtures, are of barium iodate monohydrate in water at 2.5’ shown graphically in Fig. 3 . The curve for silver 1251 F Fr MacDoulrall nnd J Rehner, TXISJOURNAL, 56. 3fi8 acetate in water solvents is taken from the measurements of MacDougall and R e h n e ~ - . ~The ~ ~



Dec., 1939


Soly., M 0.05032

- log m


0.009978 .04995 .lo02 .2001 .4992 ,9985

0.05099 .05377 .05619 .06476 .06959

1.2980 KNOa 1.2925 1.2695 1.2503 1.2270 1.1887 1.1575

0.005084 .009924 .04979 .09073 .09996 .1999 ,4927 .9979 1.999 2.999 3.172 4.002

0.05081 .05116 .OS13 .05585 .05649 ,05985 ,06583 .07209 .07781 .08034 .08067 .OS076

0.004987 .01007 .04995 .05017 ,09987 .2018 .3045 .4002,

0.05077 ,05127 .05429 .05429 ,05695 ,06093 ,06453 ,06637




a calcd., A.

0.2244 0.2469 .3221 .3954 .5093 .7510 1.033

(5.47) 4.64 4.44 4.50 4.44 4.41

NaNOa 1.2941 1.2910 1.2666 1.2530 1.2481 1.2229 1.1816 1.1429 1.1090 1.0951 1.0933 1.0928

0.2392 .2472 .3224 .3829 .3955 .5096 ,7474 1.034 1.441 1.762 1.804 2.021

4.01 3.93 4.21 4.01 4.07 4.04 3.90 3.93 4.01 4.04 4.21

KClOa 1.2944 1.2901 1.2653 1.2653 1.2445 1.2152 1.1903 1.1780

0.2361 .2477 .3228 .3232 .3960 .5126 .6075 .6831

3.21 3.24 3.50 3.53 3.58 3.30 3.21 3.41


Soly., M

0 0.009331 ,05005 ,09988 ,1506 ,2000 .2996

0.03613 ,03703 .04023 ,04275 ,04487 ,04660 .04913

- logm*


1.4421 1.4316 1.3955 1.3691 1.3480 1.3313 1.3087

0.1902 .2153 .3005 ,3777 ,4422 .4966 .5905


(4.31) 3.32 3.67 3.56 3.51 3.53

TABLE VI11 MEANACTIVITY COEFFICIENT OF SILVER ACETATE IN PURE SATURATED SOLUTION Solvent so D S f WaterI4 78.55 0.0531 0.0663 0.800 10% dioxane 69.71 .0396 ,0503 .786 20% dioxane 60.81 .0279 .0361 .771

limiting slopes for the different dioxane solvents involved, namely, 0.505 for O%, 0.606 for 10% and 0.744 for 20% dioxane are shown as dashed lines in the figure for comparison with the experimental curves. In the case of the barium iodate


solubilities, in which the measurements extend down to a much lower value of the ionic strength, the observed limiting slope, as has been shown also by Bronsted and La Mer2' and by Macdougall and Davies,12is in actual agreement with the theoretical value of 1.01.















Fig. 3.Solubility of silver acetate in dioxanewater mixtures in the presence of added salts: A, 0% dioxane, with NaNOa; B, 10% dioxane, with (1) KNOa, (2) NaNOa, (3) KCIOa; C, 20% dioxane, with KClOa.

The various Debye-Hiickel values of the a parameter for silver acetate, silver sulfate and barium iodate are summarized in Table IX. Chloupek and Danes2* have reported another Debye-Hiickel value of a for silver sulfate, determined, however, from studies on the e. m. f.'s of cells. Their values of 1and 3 A. for the diameters of the silver and sulfate ions, respectively, may be compared with the figure for this electrolyte given in the table. VALWSOF a Solvent Water

Alcohol (10-30%) Acetone (10-30%) Dioxane (10%) Dioxane (20%)



AgCzHgOn AgzSO4 Ba(1Oa)rHnO 4.35 (KNOa) 2 . 4 3 (KNOa) 1 . 8 0 (KNOd 3.69 (NaNOa) 3 . 2 0 (KCl) 3.52 (MgCla) 1.70 (Ba(N0s)r) 3.99 (KNOa) 3 . 7 2 (NaNOa) 3.99 (KNOs) 3.87 (NaNOa) 4 . 4 8 (KNOs) 4 . 0 3 (NaNOa) 3.37 (KCIOa) 3.52 (KCIOI)

It is observed that the values are not and cannot be characteristic of the saturating salt itself (27) J. N. Briinsted and V. K . La Mer, THISJOURNAL, 46, 555 (1924). (28) J. B.Chloupek and V. 2. DaneS, Coll. Czech. Chcm. Comm., 4, 165 (1932); B.C.A., 572 (1932).

F. W. nA!v1sAND J. E.



but are averages including the added solvent ions too. It is obvious that we cannot arbitrarily take an average of these parameters in applying Born's equation although we expect the a which should apply to be of the same order of magnitude. We have, therefore, thought it best to calculate a values by applying the "corrected" Born equation (Eq. 8) to solubilities experimentally observed in the several solvents. There may be some question whether the same a should be used in the Debye-Hiickel terms as in the Born term, but inasmuch as the latter term far outweighs the others in numerical significance, appreciable variations in a have little effect on the Debye-Huckel VALVES OF a FOR




I N ' b S T I N G EQUArION 8 .4gzso1 Ba(iOal? HtO



3 '

3 05


3 30



3 2



3 57

4 59


1 2s



2 34

Vul. 61

corrections, and it seems reasonable to use a single figure for the ion diameters. We have calculated the a values for the several solutions, combining the solubility in water with the solubility in the several mixed media as indicated in equation 8. The resulting a's are listed in Tables XI, XI1 and XIII. In each case the values increase with increasing proportion of dioxane. As a result of the mathematical transformations involved in the change of units the a's calculated from molarities are smaller than those based on mole fractions. But whichever units are used, the a values are of about the same magnitude and of the same relative consistcncy over the entire range; furthermore, they are comparable with a parameters appearing in Table VIII, a t least in the water-rich solutions. For the purpose of comparison the various a values for the three separate salts which might possibly be used in testing the Born equation (Eq. 8) are collected in Table X. The (L'S so determined and their variations as D changes, test and disprove the validity of the fundamental assumption made in the simple application to solubilities of Born's equation corrected for interionic attraction, namely, that there is a value for n (an effective ionic diameter) which


Limiting slope of Fig. i (molarities) Equation 8, &lo% dioxane (using molarities) Limiting slope of Fig. 2 (mole fractions) Equation 8, O-lO% dimane (using mole fractions) Equation 8, 0-307, dioxane (using mole fractions) Average of Debye-Huckel values, Table VI11



10 20 30 40 50 60

70 80 90




Bared on molarities Using a = 3.05 S calcd. S calcd./S obsd. x 101 x 100

3.05 3.2G 349 4 3.49 206 9 3.80 98 6 4.29 32 2 5.01 5 30 0.37 0.186 8.98 2 1 x 10-4 14.1 z 5 x 10-14


Using X calcd.





4.36 459 4 91 5 48 6 26 7 99 11 2 17 3

7.48 51s 3.03 1.34 0 355 .0293 1 9 x 10-4 2 1 x 10-9


10 20 30 40 50 60 70 80 90

a calcd ,




= 4.14

calcd /X obsd x 100

96.9 90.9 77.5 53.9 35.2 4.30 0.066 2 . 2 x 10-6

Using a = 4.60 X calcd /X obsd x 100

X calcd.



10.04 102.2 7.99 102.4 5.69 99.9 90.6 3.54 1.71 69.0 0 522 37.1 ,0494 7.24 5 . 9 x 10-4 0.197 1 . 9 x 10-8 2 . 0 x 10-6



Based on moltrritier Using a 3 41 S calcd. S calcd./S obsd x 104 x 100

3.41 3.68 74.36 4.00 27 29 4.42 6.639 5.00 0.858 5.96 0353 e 19 1 02 x lo-' 12 6 8 9 x 20 9 5 4 <


4 14 96. (1 84 3 65 0 37 6 12 5 1.07 o 0035 2 ii x 1 0 - 9


Based on mole fractions


89.4 67.9 38 9 14 3 1 96 0 016 3 . 9 X lo-? 8 6 X

Based on mole fractions


calcd ,


3 93 4.24 4.59 5 07 5 72 6.78 9.36 14.6 24 3

X calcd. Using a X* calcd./X 3.93 obsd



1.601 0.671 .196 .0326




72.0 44.8 18.7 .00202 3 38 1.23 X loTs 0 050 3.6 X 3.3 X 13X 3.3 X

X calcd. Using aX=calcd./X 4.60 obsd.



3.297 1.948 0 940 .329 ,0703 .00588 6 3 X 1.0 X IO-* 12 X



109.0 110.4 100 8 75.1 40.3 9.81 0.262 9 . 3 X 10-6 3.2 X 10-lS


Dec., 1939



rHE SOLUBILITY OF Based on molarities Usinga 3.30 S calcd./


yt. %




10 20 30 40 50 60 70 80 90

3.30 3.51 3.80 4.32 5.34 6.95 11.1 17.9 23.6

S obsd. X

S calcd.




2.274 0.862 .219 .0291 .00115 2 . 8 X lo-’ 1.1 x 10-12 4 . 5 x 10-24


Using a = 3.57

e a P’HtO calcd.,

X calcd.

x 10’



92.3 73.9 42.2 11.4 0.93 .0030 1 . 9 x 10-8 1 . 3 x 10-19

3.57 3.93 4.28 4.89 6.11 8.07 13.3 21.9 27.8

F ;:&

242 aoDo


1.6 X 1.6 X 5.1 X


BO6 0 1 +Boa&

X obsd. X


X calcd.






4.684 1.908 0.537 ,0832

should remain constant from one medium to another and should be related to the solubility and dielectric by the equations given above. At this point it may be appropriate to note that deviations from theoretical curves similar to those noted in this communication have been observed and reported by Tomlinson in a study of reaction velocities.” We could proceed, on the other hand, on the more practical basis that the effective a changes as D changes. Then, assuming in the original equation that there is a characteristic value of a (= a ~ for the electrolyte in water (this could be taken as the value approached by the composite a when the percentage of dioxane decreases), an appropriate a ( = al) representing the behavior in each successive solvent can be calculated using the modified Born equation log S’ = log so+z+z- (1

Using a = 4.30 X calcd./

X calcd./

242) a&

Since the a as already calculated (Tables XIXIII) is in effect a composite of the diameters for each pair of media, although it is not an arithmetic average, the a values determined in the new way would be even larger than those already calculated. The principle, however, of an effective diameter varying with the medium would offer no possibility of testing the Born equation for the prediction of solubility, inasmuch as we have for the present no independent determination of the change of the diameter, a, for any given electrolyte with respect to change of dielectric. The a would in other words be an adjustable parameter only. We have consequently omitted (29) H.M.Tomlinson, “Influence of Dielectric Constant upon the Kinetics of the Bromoacetate-ThiosulfateReaction,” Thesis, Columbia, 1939; cf. also V. K. La Mer, J . Franklin Insf., 426,709 (1939).

5.75 2.735 0.955 .201 .0168 1 . 6 x 10-4 1 . 2 x 10-8 3.2 X

89.6 70.4 40.1 11.2 1.01 0.0044

5.8 X 2.4 X

X obsd. X


$‘Ha 100

109.4 109.8 100.9 71.3 27.2 4.06 0.044 4 . 2 X lo-’ 1 . 5 x 10-14

such additional figures in the preparation of our tables.


Solubilities Calculated from Born’s Equa-

tion While the variations of the a’s cited in the last three tables themselves indicate the failure of Eq. 8 to account for our experimentally determined solubilities, a clearer picture of the extent of this failure may be drawn from a comparison of predicted and determined solubilities. The ) calculations have been made in terms of mole fractions using first the a values for the 0 and 10% dioxane solutions and again with the a for the 0 and 30% solutions. The calculations have been repeated on a volume-molar basis using the a which fits the 0 and 10% observed solubilities. The calculations, as before, involve successive approximations; first the solubility is computed without correcting for the interionic attraction, then with this preliminary value the interionic attraction terms are added to the equation and a new solubility is computed. Two approximations are all that are needed except for silver sulfate. The calculated solubilities using the a values just noted together with the ratio of the calculated to the experimental figures are given in the three tables. The calculated and observed curves for log S are plotted together in the two sets of concentration units in Figs. 1 and 2. As can be seen from these figures, the calculated values for log S for both molarity and mole fraction units lie in every case on a very nearly straight line with a slight curvature in the first two or three points due to interionic effects where the ionic strength is appreciable, and with a nearly constant slope, given by slope





Vol. 61

where a has the value used in the given series of independent method of estimating these ionic calculations. All the predicted values are too diameters in solution, so that our conclusions resmall except a t a few points where the 3UojG a is garding them must remain speculative. used for the computations since herc the straight Scatchard30has shown that the e. m. f. of cells line of the calculated values cuts through the curve made up of hydrogen chloride in aqueous alcoholic ttf the observed solubilities at the point for :?Oyo solutions may be accounted for by taking a as dioxane. equal to 1 A. up to 50% alcohol and decreasing While there is general agreement as to order a to about 2 absolute alcohol. On the of magnitude of the solubility down to 40-507, other hand, the values for sodium chloride and dioxane computed on the basis either of mole potassium chloride in aqueous alcohol are not fractions or of molarities, even this degree of con- much different from the a's for these salts in cordance disappears for solutions containing more water.27 The observation also of rather constant values than SO($$ dioxane. This disagreement is a little worse for the two higher valence type electrolytes of a over a wide temperature range for cadmium than for silver acetate. The use of a larger a im- sulfate and zinc sulfate in water3i and the very proves the agreement in the high-dioxane mix- small variation of a for hydrogen chloride in tures but makes the figures for the low-dioxane several dioxane-water solvent mixtures32 may solvents higher than the experimental, as already be noted as further evidence for the belief that ;or a given salt the values of a should be charpointed uut. acteristic only of the electrolyte and independent VI. Discussion of the medium. It would be pointless then to The disagreement between the observed and try to bring the Born equation into agreement the calculated solubilities is seen to be substan- with our experiments by artificial adjustments of tially the same for either the volume or the mole the a s. 'The assumption of continuous character of the fraction units of concentration. The question solvent medium is apt to be a serious source of then of the proper concentration units to use error in mixed solvents. The question is related would seem to involve factors of minor importance of the effective diameter of solvated ions. to that as compared to the larger discrepancies noted beIt is to be expected that in a mixed solvent the tween observed and theoretical solubilities. more polarizable substance, water in the present In seeking some explanation for the wide deviacase, will be preferentially oriented around the tions between prediction and experiment in these ions, with more soluble electrolytes such as solubihties, we observe that three assumptions sodium nitrate, this phenomenon may lead to the invohed in Eq. 8 may have to be modified in formation of two liquid layers. While this factor such a way as to bring about better agreement. has often been mentioned (for example, by ScatchThese assumptions are (1) a constant diameter for the several electrolytes independent of the ard30in 1925), few calculations of its consequences medium, ( 2 ) a homogeneous structureless char- have been made. Using the Debye expression for acter for the dielectric medium and (3) the absence the distribution of molecules in a dielectric mixture around a univalent in conjunction of molecular or ionic associations. with Born's equation, Dunning and Shutt2*report While it may seem reasonable to suppose that a calculation of the effective radius of silver chlovariable solvation will lead to a changing effective ride when dissolved in a one molar aqueous soludiameter in different media, it is likely that ions tion of urea, this radius agreeing more closely with in solutions containing even small percentages of crystaIlographic values than the radius deterwater will be fully hydrated or nearly so, so that mined from the Born equation alone. The calcufor practically the whole range of dioxane-water mixtures, a should be nearly constant for a given lation, however, is very tedious, and while the electrolyte. Moreover, if the extent of solvation correction introduced will be in the right direcis actually smaller in dioxane-rich solutions, tion of increasing the solubilities calculated for a (30) G Scatchard, T m s JOURNAL,47,2098 (1925). smaller diameters should be expected, whereas, as (31) V. K La Mer and a'.G Parks, r b d , 53, 2040 (1931); I A. already seen, the observed solubilities require Cowperthwdite and V. K L a Mer, tbrd , 63,4333 (1931). 13 S I I ~ r n e d J, 0 hlorrison, F Walker, J. G. Donelson and much larger values to bring about agreement C 132) Calmon, h t d , 61, IC) ilc(39). with the Born equation, 'There seems to he no (3) P U r l q e Z P h ~ s z k C h w . , Cohen Festband, 56 (1927).

Dec., 1939



given ion size, there is still doubt whether a constant a can be determined for the whole series of solvents, especially when ion association becomes important. The effect of selective dipole concentration around the ions would not enter into solutions of salts in pure organic solvents of low dielectric; but in most of these liquids the solubility of electrolytes having slight solubility in water will be below the limit of quantitative estimation, and no tests can be applied. One may suggest that the small but detectable dissolving in 90% dioxane of such a difficultly soluble substance as barium iodate, while its solubility in a single solvent of the same dielectric capacity would probably be zero,34is evidence that the preferential adsorption of water around ions has a tremendous effect in determining solubility in solvents containing water. Another effect particularly important in solvents of low dielectric constant is ion clustering in the Bjerrum sense.3s Such clustering will result in an increased solubility over that calculated assuming complete dissociation. A formula for the formation of binary pairs for a uni-univalent electrolyte has been given by Fuoss and Kraus as a function of the diameter of the ions and the dielectric of the medium,aain the approximate and simplified form 1 = moa44nN


(k D+T)

where K is the dissociation constant and the other letters have their usual significance. Application of this equation gives approximately 3.1 X and 1.1 X 10-l2 as the values of K for silver acetate in 80 and 90% dioxane, respectively, where our value of a, namely, 3.05, is taken to deline the corresponding value of K . Such results would indicate negligible quantities of silver acetate in molecular form when the ion concentrations are of the magnitudes given in column 3 of Table XI. Smaller a values in the Fuoss and Kraus equation will give, it is true, a K sufficiently small to account for observed solubilities. If a is taken as one-half of 3.05 8.,roughly correct solubilities are predicted. The magnitude of K calculated according to this equation is extremely sensitive to small changes in a. Uncertain, therefore, as a itself is, any value of K derived in (34) Thallous iodate has negligible solubility in alcohol, of much higher dielectric than 90% dioxane; cf. La Mer and Goldman, ref. 19. (36) Cf.evidence from absorption curves reported by H. v. Halban and B. Szigeti, Helv. Chim. Acto, 90,746 (1937). JOURNAL, 65,1010 (1938). (36) R.M.Fuoss and C.A. Kraus, THIS


the Fuoss and Kraus manner must be considerably in doubt. It is unlikely, furthermore, that the mean ionic diameter for silver acetate can be much smaller than 3.05 d.,for it usually has been found that the parameters fitted into the ordinary Debye-Hiickel equation are smaller than expected in the sense that sometimes negative values are found, while a more extended treatment such as that of Gronwall, La Mer and Sandveda7 leads always to positive and hence physically plausible values of a. It is possible that the a's we have used, which are in general agreement with Debye-Hiickel values, are too small also. Several authors have declared that in media below about 60 in dielectric capacity, most of the divalent ions will be associated and that below a dielectric of about 30 univalent ions will be largely associated into neutral molecules. If this be the case in our experiments, the solutions should behave like non-electrolytes where the dioxane percentages exceed 50 approximately, and the solubilities should be nearly proportional to the mole fractions of water in the mixtures. This is not observed, as seen from the data in Tables 11-IV, so that we may infer that the solutions of the un-ionized salts here studied are not ideal. A complete prediction of solubilities in the light of modem theories would involve an estimation of the solubility by the corrected Born equation followed by a calculation of solvent segregation around the ions, a redetermination then of the ionic solubility and a determination of the Bjerrum ion association using the Kraus and Fuoss method of calculation. Such a series of calculations has not yet been made, but while laborious it would have some interest as a test of the combined theories.

Summary The solubilities of silver acetate, silver sulfate and barium iodate have been determined a t 25' in a series of dioxane-water solutions extending from 0 to 100% dioxane. The dielectric constants cover the range 2.1 to 78.55. The solvent effect of potassium nitrate, barium nitrate, magnesium chloride and potassium bromate solutions on barium iodate in water has been examined experimentally. The solvent effects of potassium nitrate, sodium nitrate and potas(37) T.H.Gronwall, V. K . La Mer and K. Sandved, Physik. Z . , 99, 358 (1928).



sium chlorate on silver acetate in 10% dioxane solutions, and of potassium chlorate on silver acetate in 20% dioxane, have also been investigated. The corresponding Debye-Huckel a parameters have been calculated. The Born equation even when corrected for interionic effects fails to account for the course of the solubilities as the percentage of dioxane increases, predicting in each case lower solubilities than those observed. The differences increase in media of higher dioxane content, pre-


Vol. 61

diction and observation disagreeing even as to order of magnitude in solvents containing more than 50% dioxane by weight. The disageernent is not diminished appreciably by substitution of mole fractions for volume concentrations, showing that this question is of minor importance as a factor in the application of the Born equation. Modifications to be superimposed on the Born equation to bring it into line with the observations are discussed. NEWSORK,N. T.




Kinetic Studies on Iodine Derivatives. I, The Thermal Decomposition of Acetyl Iodide1 BY J. LESLLEJONES The thermal2 and photochemical3 decomposi- it is possible to study the thennal decomposition tion of acetaldehyde into methane and carbon of the product, acetyl iodide, in the same temperamonoxide has been investigated by numerous ture range as the alkyl iodides. Such a study authors. Since the deduction of a possible free casts light upon the mechanism of the thermal radical mechanism for the thermal and photo- decomposition of acetaldehyde, supporting the chemical decomposition of acetaldehyde by Rice free radical mechanism. I n addition information and Herzfeld,4 a great deal of controversy has is gained concerning the secondary reactions of arisen concerning its validity. The opposing the free radicals involved. theory2c assumes that the activation process, Experimental Section thermal or photochemical, is followed by a reThc rate of decomposition of the acetyl iodide was folarrangement of the acetaldehyde directly into the final products, methane and carbon mon- lowed by the pressure change method, employing essentially the apparatus of Jones and Ogg.a The quartz spiral oxide. manometer prevented the iodine evolved in the reaction On the basis of previous investigations6it seems from reacting with the mercury in the absolute maprobable that the thermal decomposition of alkyl nometer. A hollow barrel stopcock was employed t o close iodides occurs by the breaking of the carbon- the reaction vessel, employing a plasticizer as lubricant. iodine bond (temperature range of 250-350') as The compound, di-o-xenzyl-monophenyl phosphate,a has B vapor pressure of 5 mm. at 285" and a viscosity of 170 the primary process. This reaction is followed by centipoises a t 60". The liquid was warmed in a high the secondary reactions of the free radicals. By vacuum system for several hours in order to remove all the substitution of an iodine atom for the hydro- solvents. All observations indicated that it was ungen atom of the aldehyde group in acetaldehyde affected by the iodine generated in the system, as its structure suggested. It was necessary to heat the stopcock

(1) A preliminary report of this paper WPS presented at the S y n closing the reaction vessel to about 110 in order t o prevent pcsium on the Kinetics of Homogeneous Gas Reactions, University of condensation of iodine. Wisconsin, June 20-22,1939. (2) (a) Fletcher and Txis JOURNAL, 68, 2136 (193t~), The acetyl iodide was prepared by the treatment of (b) Letort, J . chim. phy (1937); (c) Staveley and HinshelEastman Kodak Co. acetyl chloride with an excess of dry wood, J . Chcm. SOC.,1 (d) Brenschede and Schumacher, hydrogen iodide gas at 0". The brown crude product Ber., ?OB,452 (1937), ( e ) Rollefson and Fauil, THISJOURNAL, 69, (90% yield) was protected from the air and vacuum dis625 (1937). (3) (a) See P. A. Leighton, J. Phys. C h m . , 42, 749 (1937), for a tilled in a n all-glass fractionating system two times, large complete review t o that date; (b) Blacet and Volman, THISJOURNAL, amounts of the first and third fractions being discarded. 60,1234 (1938); fc) Gorin, J . Ckem. Phys.,?, 256 (1939); (d) RollefA third fractionation was carried out in high vacuum, the son and Grahsme, paper presented a t the "Symposium on the product being distilled directly into the reservoir of the Kinetics of Homogeneous Gas Reactions," University of Wisconsin June 20-22, 1939, system employed for the study of the reaction rates. In (4) Rice and Herzfeld, TEESJ O U R N 66, ~ , 2 8 4 (1934). ( 8 ) Furnished through the courtevy of the Dow Chemical Corn(5) See Jones, Ibrd., 60, 1877 (1938), for complete references to earlier research pany, Midland, Mich.