as long as the p H < 7; Kl' is the first apparent dissociation constant of carbonic acid. If we make [Me HC03+] = x, [Me++] = y, and [ H C O . < - ] = ZJ,then c = x y and 2y x; = v. The constant for the formation of a complex according to (1) is then
+
+
For infinite dilutions this gives lim L
=
v
c=ov(v
(4)
- c)
Introducing K = K i ' [ C 0 2 ] s a t = 4.63 X lo-' X 0.0308 = 1.425 X a t 28.4' as a new constant,2 we obtain according to (2) v =
(5)
~f~02/(H+)fu
arid (4) becomes
Table I shows that 2c - K/(H+)and with it K , converge from negative values toward zero with decreasing c for both bicarbonates. Therefore, complex formation does not occur a t these ionic strengths. The pH of mixed bicarbonate solutions (Mg: Ca = 4 : 1, 3 : 2, 1: 1 and vice uersa) was measured at values of p ranging from 2.8 X to 1.75 X loF4. The finding that the p H is indepeiident of the Mg:Ca ratio for a given value of p indicates that there is no strong specific ionic interaction a t these concentrations. pH determinations were made on a series of solutions made up by dilution of A with increasing volumes of B in the p H meter. The plot of the p H values us. the volume ratios oacasioiitdly showed a maximum, ii finding which is inconsistent with the :Lbove conclusion. TABLE Ir 1 ~ E N S I T I E Sd 01" L)IFPERENT MIXEI) SOLUTIONS O F
(Hco,)~ AND ME, minole/l.
lG3.5 130.8 98.1 82.5 65.4 32.7 0.0
Vol. 62
NOTES
1136
-
Ca(HCOa)2I N WATER SATURATED (p 1 ATM.) Ca, mmole/l.
0.00 2.ox 4.06 5.08 6.09 8.12 10.15
Total hlg
+ Ca
163.5 132.8 102.2 87.6 71.5 40.8 10.15
WITH
hlg-
cop
dl9
1.0199 148 116 100 82 52 19
This anomaly proved t o be due to the adsorption of calcium and magnesium ions on the surface of the electrode; an equilibrium was not established bet,ween the more dilute solution and the adsorbed ions. The anomaly was never observed if the electrode was washed thoroughly between experiments, with highly diluted hydrochloric acid and distilled water, a precaution which was followed through our work. (2) Values of KI', which were taken from H. 6. Harned and B. B. Owen, "The Physical Chemistry of Electrolytic Solutions," Reinhold Publ. Corp., New York, N . Y., 1956, were interpolated. The dependence of the [COz)sn+,.on c is within the experimental error: see A. Seidell, "Solubilities," Vol. 1, D. Van Nostrand Co., New York, N. Y., 1940, pp. 221, 226.
Acknowledgments.-The p H meter used was designed by Mr. Van Aerschodt and built in our workshop.
THE CONDUCTANCE O F CONCENTRATED AQUEOUS SOLUTIONS OF POTASSIUM IODIDE AT 25' AND OF POTASSIUM AND SODIUM CHLORIDES AT 50' BY J. F. CHAMBER@ Chemistry Department, Universitv of Western Australia, Nedlands. Australza Received January BO, 1968
This paper continues the report of work carried out in t'he research program outlined in a previous paper.2 The aim, briefly, is to provide some precise experimeiit'al data of the conductance and viscosity of concentrated aqueous solutions of simple salts for study in the light of recent theoretical developments. 3-6 Experimental The apparatus and the experimental procedure used in the measurement of the present data are the same as described in ref. 2. The results presented in that paper, for iiqueous solutions of sodium and potassiuni chlorides, agreed within 0.02% with the results of Shedlovsky, et ~ 1 The . ~ purity of the stock of these two salts used in the present work was several times checked by a determination of the equivalent conductance of an aqueous solution a t 25". These values were always in excellent agreement with those in ref. 2. Sodium chloride was A.R. grade, purified by precipitation from Conductance water by hydrogen chloride gas. I t was then given a preliminary drying at 110' and hefore use was heated a t 400" for 48 hours. Potassium chloride was A.R. grade, purified by precipitation from conductance water, then given a preliminary drying a t 110". Before use it was heated a t 400" for 48 hours. Potassium iodide was A.R. grade, recryst,alli&edtwice from conductance water, dehydrated over calcium chloride under vacuum for a week, then finally heated a t 150" in air for four days. The stock so prepared was used within a week or two. Regular check solutions of t>hesalt so prepared occasionally were found to have conductances which were up to 0.3% high. I n such cases the whole procedure was repeated untJil t,he conductances agreed with the data recorded in Table I. Conductance water was the laboratory distilled water redistilled through Pyrex and stored in polythene bottles. It had a specific conductance of about 1.2 X int. ohms-lcm.-' a t 25' and about 2.5 X 10-5 int. ohm-' cm.-I a t 50". A solvent correction was always made to the measured specific conductance of the solutions. Densities were h k e n from the International Critical Tables,? and vacuum corrections were used throughout. Molecular weights used in computing the molarities were: KCl, .74.553; NaCl, 58.454;KI,166.02. Temperatures were malntalned constant in oil bath thermostats a t 25 and 50' to within &0.002' by the use of calorimeter thermometers, whose true reading was found on a platinum resistance thermometer to within &0.01". The actual measurements a t 50" were made a t temperatures up to 0.09' low and later corrected to 50.00°. The Pyrex glass cells were as described in the first papera and had cell constants from 0.5 to 459 (1) Electrolytic Zinc Company Research Fellow, 1955. (2) J. F. Chambers, Jean M. Stokes and R . H . Stokes, THISJOURIFAD. 60, 985 (1056). (3) H. Falkenhagen, M. Leist and G. Kelbg, Ann. physik., [61 1 1 , 51 (1952); H. Falkenhagen and M. Leist, Naturwiss., 41, 570 (1954). (4) R. M. Fuoss and L. Onsager, Proc. Nutl. Acad. Sci., U. S., 41, 274 (1955). (5) B. F. Wishaw and R. H. Stokes, J . A m . Chem. Soc., 16, 2065 ,(1954). (6) T. Shedlovsky, A . S. Brown and D. A. McInnes, Trane. Electrocham. $oe., 66, 165 (1934). (7) "International Critical Tables," Vol. 111, BlcGraw-Hill Book Co., Inc., New York, N . Y.
,
cm.-l. The cells were calibrated a t 25" with the Jones and Bradshaw 1 and 0.1 D potassium chloride solutions.* For use at 50" the cells were regarded simply as long narrow glass tubes in order to derive the expression for the cell constants: a50 = m b [ l ( 1 - 25)(~g]where ag is the coefficient of linear expansion of Pyrex glass. This is undoubtedly a gross simplification for some of the cells used but is in agreement with previous workers: The resistance measurements were made with a Lee& and K'orthrup Jones conductivity bridge a t two frequencies, 1000 and 2000 c.P.s., and the resistances extrapolated to infinite frequency as recommended by Jones and Christian.'" The correction rarely exceeded 0.03% of the resistance.
-
TABLE I EQUIVALENT CONDUCTANCES OF KC1, NaC1 AT 50" KI AT 25" (A given as I n t . ohm+ cm.2 mole-') ---KI C, mole I . - '
0.010492 .012307 ,015422 ,019939 ,024356 .056218 ,068089 ,084311 ,13956 ,14665 ,17548 .22631 .36896 .49333 ,65169 ,72161 ,97751 1.2310 1.6595 2.0837 2.3417 2.7879 2.9347 3.6492 4.1946 4.6550 5.5041 5.6916
1137
NOTES
Sept., 1958
-KC1
25O--.. A
142.03 141.45 140.55 139.49 138.56 134.32 133.23 132.05 129.08 128.75 127.71 126.19 123.32 121.66 120 .08 119.51 117.74 116.27 114.06 111.86 110.45 107.87 106.94 101.95 97.62 93.34 85.31 83.40
C,
mole
-NaCl
50°---
l.-i
0.012171 .024630 .045221 ,062498 .078744 ,098784 .14151 .20104 ,29904 .35732 .41241 .55951 ,76679 .98528 1,4074 1.8187 2.2922 2.7382 2.8882 3.3107 3.6385 4.3'348
A
213.93 208.63 203.18 199.80 197.35 194.73 190.62 186.29 181.17 178.86 176.83 172.53 167.79 163.86 157.65 152.47 147.08 142.30 140.72 136.28 132.83 124.87
C.
A N D OF
50"--1
mole
0.009626 ,019176 .030925 .G48570 ,068474 ,11379 .1455ti ,18299 ,27966 .37615 ,50448 .65880 ,80532 1,0007 1,2063 1.4931 1.6744 2.0900 2.2154 2.36-13 2.5153 2,9739 3.2748 3.9907 4.5035 5.0776
A
184.89 180.35 176.61 172.61 169.29 163.75 160.84 157.98 152.27 147.92 133.26 138.64 134.87 130.13 126.29 121.07 118.02 111.58 109.78 107.U 105.51 '39.41 !)5.62 87.07 81.26 55.11
Results The experimental concentrations a i d equivalent, conductances are listed in Table I. The conceiitrations are in moles per liter a t the temperature of measurement and the equivalent conductances are in international ohm-' cm.2 mole-1 units. The 25' results are as actually recorded but the 50' results have been adjusted from the actual experimental temperature to the 50.00' values utilizing separately determined temperature coefficients of conductance. At 50°, (lOO/A) (dA/dt) varied from 1.46% a t C = 0 to 1.14% a t C = 4 M for potassium chloride and from 1.55% a t C = 0 to 1.48% at C = 4 M for sodium chloride. The magnitude of the adjustment to the equivalent conductance did not exceed 0.03% for sodium chloride or 0.1% for potas(8) G. Jones and B. C. Bradshaw, J. A m . Cham. Soc., 55, 1780
(1933).
(9) B. B. Owen and I". > I . Sweet,on, ibid., 68, 2811 (1911). (10) 0.Jones and S. M. Christian, ibid., 61. 272 (1935).
sium chloride. By mealis of large scale graphs of the deviation fullctioll = A + ~ d c where , A is chosen to give as flat a curve as possible, the results were plotted and the data a t round concentrations were illterpolated and are presented ill Table 11. At 25' the limit,s Of possible error are f0.030Jo. A t 50°, the limits of possible error are f0.07~~. TABLE I1 EQUIVALENT CONDUCTANCES AT ROUNDCONCENTRATIONS C, mole 1. - l
A K,I
AKC1 50'
ANaC1 500
0.05 .10 .15 .20 .25 .30 .40 .50 .60 .70 .80 .90 1.0 1.2 1.4 1.6 1.8 2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5
134.97 131.11 128.69 126.89 125.57 124.51 122.68 121.58 120,53 119.68 118.92 118.21 117.61 116.43 115.37 114.35 113.34 112.30 109.55 106.47 103.04 99.17 94.84 !IO. 12 85.34
202.14 194.59 189.93 186.35 183.51 181.13 177.28 174.14 171.50 169.18 167.14 165.31 163.63 160.57 157. 76 155.14 152,69 150.35 144.82 139.55 134.31 128.80
172.35 165.26 160.48 156.66 153.86 151.27 146.98 143.41 140.29 137.52 135.00 132.65 130.46 126,42 122.71 119.26 116.00 112.94 105.72 99.08 92.88 86.96 81.30 75.92
25
.
.... ...
Comparison with Previous Results.-With the exceptio11 of the potassium iodide data, the method of checking these results with previous work was t,o inFert the values for A into a theoretical equation giving Ao in terms of A and a parameter d and to vary d to give the best constancy for Ao from data a t concentrations up to 0.03 AT. This is the method used by Robiiison and Stokes,ll but instead of their equation, one was used in,which the term for the relaxiition effect was that derived by Fakenhagen and which also corresponds to the first-order term in the expression due to Pitts.12 Potassium Iodide at 25".-The A values agree to within fO.OlOj, with the data of Jervis, ct aZ.,la and with the data of Longsworth14 and these data are incorporated in the deviation function graphs. Potassium Chloride at 5O0.-The use of d = 2.58 A. gives Ao = 228.95 from the data up t o 0.03 M . Hariied and Owen'5 give 228.92 f 0.04 and the re(11) R. A. Robinson and R. H. Stokes, ibid., 76, 1991 (1954). (12) E . Pitts, Proc. Rov. Soc. (London),A a l l , 43 (1953). (13) R. E . Jervis, D. R. Muir, J. P. Butler and A. R. Gordon, J. A m . Chem. Soc., 7 5 , 2855 (1953). (14) L. G. Longsworth, reported in D. A. RIacInnes, "Principles of Electrochemistry," Reinhold Publ. Corp., New York, N. Y . , 1939, p. 339. (15) H. S. Harned and B. B. Owen, "The Physical Chemistry of 1Clectrolytic Solutions," 2nd ed., Reinhold I'ubl. Corp., New York, N. T., 1950, p. 589.
i138
NOTES
Vol. 62
calculations by Robinson and Stokesi1of the data due to Benson and GordonlBand t o Owen and Zeldesl'lead to 228.83 f 0.05. Sodium Chloride at 5o0.-The use of a = 2.55 A. gives A,' = 197.81 from the data up to 0.02 M . Harned and OwenI6 give 197.77 =k 0.05. Acknowledgments.-The author wishes to express his thanks to the Electrolytic Zinc Company for receipt of their Fellowship in Electrochemistry and to Professor R. H. Stokes for much helpful advice and assistance.
sisting of only two hydrocarbons that separates into two liquid phases. The present system is the first to come to the authors' attention. Multiple fluid phases have been obtained3 upon addition of methane and propane to a natural crude petroleum. The two liquid phases so obtained were certainly multicomponent and probably did not consist only of hydrocarbons, since it is unlikely that all of the sulfur-, nitrogen- and oxygen-containing compounds in the oil were precipitated into the asphaltic solid phase. The isobutaiie used in the present study had a (16) G. C. Benson and A. R. Gordon, J. Chem. Phus., 18,473 (1945). minimum purity of 99% (Phillips Petroleum, re(17) B. B. Owen and H. Zeldes, i b i d . , 18, 1083 (1950). search grade). The phenanthrene was Eastman Kodak Co. white label product; its melting point PHENANTHRENE-ISOBUTANE : BINARY was observed to be 99.5'. The materials were HYDROCARBON SYSTEM HAVING TWO used without further purification. The major item of equipment used was a pressure LIQUID PHASES1 cell with internal window4 that permitted visual observation on the volumetric and phase behavior BYJ. G. ROOFAND N. W. CRAWFORD, JR. at pressures up to 500 atmospheres and temperaShell Development Company, Houston, Texas tures as high as 150". The phenanthrene was Recezved March IS, 1968 placed in an auxiliary vessel, melted and degassed During the course of an investigation on the before being transferred through a heated line into phase behavior of hydrocarbon systems, we have the evacuated windowed cell. The quantity of made observations on the phenanthrene-isobutane phenanthrene so introduced was determined by obsystem, Mixtures of these two hydrocarbons ex- servation of the volume occupied by the liquid in hibit a behavior unusual t o binary hydrocarbon the cell. The density of liquid phenanthrene was systems in that two liquid phases are found to co- determined by pycnometer to be 1.066 g./cm.a at exist over a considerable range of pressure and tem- 101'. Known amounts of isobutane were then perature. I n 1950 it was stated2that the literature introduced by volumetric metering from a vessel a t seemed to contain no description of a system con- known temperature and pressure. Density of the isobutane is available from the l i t e r a t ~ r e . ~ 2500 Results of the observations on phase behavior in this system at 101' are given in Fig. 1. Experimental points are shown by circles. Phase boundaries so established are shown by solid lines. S, L , and G designate solid, liquid, and gas, respectively. The less dense liquid is designated by LI, I500 the more dense by Lz. It was not practical to expand samples sufficiently to establish the dew-point 1000 locus, which is shown only schematically by the dotted line. No attempt was made t o determine compositions of coexistent phases; hence, it is not known whether the solid phase is pure phenanthrene or a solid solution of isobutane in phenanthrene. If solid solutions do exist, an X area should be shown to the left of the S L area. Within the L1 Lz area there are two liquids whose bubble point occurs a t the pressure shown by the horizontal line ab. The liquid-liquid miscibility gap is of Considerable size a t temperatures slightly above the normal melting point of phenanthrene. A single observation a t a higher temperature (120') indicated that the gap decreases rapidly with temperature and probably would disappear below 150". Perhaps the lack of complete miscibility between these two hydrocarbons should not be surprising when one considers the differences in shape and
+
Fig. 1.-Phase
behavior in the phenanthrene-isobutane system at 101'.
(1) Publication No. 158, Exploration and Production Research Division, Shell Development Company. (2) A. W.Francis, Chapter 7 in "Physical Chemistry of t h e Hydrocarbons," edited by A. Farkas, Academic Press, New York, N. Y.. 1950.
+
(3) D. J. Vink, A. M. Ames. R. A. David and D. L. K a t & Oil and J., 89, No. 28. 34 (1940). (4) F. W. Wells and J. G. Roof, Rev. Sei. Instr., 2 6 , 403 (1955). (5) (a) B. H. Sage and W. N. Lacey, Ind. Ens. Chem., 80, 673 (1938); (b) "Thermodynamic Properties of the Lighter Hydrocarbons and Nitrogen," Amerioan Petroleurn Institute, New York, N. Y.. 1950 Gas
I
