516
GBORGE WTONOFF
Indiscriminate use of the expressions “speed,” “velocity,” and “rate” without giving definition t o these terms 1eads.to confusion, of which the theory upheld by Luder and Glasdtone is a manifestation. REFEREKCES (1) ANTONOFF,G.: J. Chem. Education 21, 420 (1944). (2) ANTONOFF, G.: J. Chem. Education 22, 98 (1946). (3) GLASSTONE, S.: Tezt-book of Physical Chemistry. D . Van Noatrand Company, Inc., New York (1940). (4) GLASSTOHE, S.:J. Chem. Education 22, 201 (1945). (5) LUDER,W. F.: J. Chem. Educstion21, 559 (1944). (6) LUDER,W. F.: J . Chem. Education 23,201 (1945). (7) TAYLOR, H . S.: A Treatise on Phgsical Chemistry, 2nd edition, Vol. 11. D . T-an Sostrandcompsny, Inc.,NewYork (1931).
ISOTONIC SOLUTIONS : OSMOTIC AND ACTIVITY COEFFICIENTS OF LITHIUM AYD SODIUM PERCHLORATES AT 25°C.’ JAMES HOMER JONES Department of Chemistry, Indiana University, Bloominglon, Indiana
Received September 3,1946
The activity coefficients of lithium and sodium perchlorates up to 1 molal have been determined from freezing-point measurements by Scatchard and coworkers (4). No measurements a t 25°C. are available in the literature. The present investigation determines the osmotic and activity coefficients of the two salts a t 25’C. and over a much wider concentration range-about 0.2-6.5 molal for sodium perchlorate and 0.2-4.5 molal for lithium perchlorate. The method used is the familar isopiestic vapor-pressure measurement developed by Robinson and Sinclair and perfected by Robinson (3). The reference salt is sodium chloride, the activity and osmotic coefficients of which have been tabulated by Stokes and Levien (5). The apparatus, except for some modification, has been described previously (1). An all-brass desiccator about 8 in. in diameter and 6 in. deep replaced the one previously used, and a new rocking device was installed. This furnished a much superior heat reservoir and helped to moderate the effect of small fluctuations of temperature in the thermostat. EXPERIMENTAL
c.
P.
Purijicatim of materials anhydrous sodium perchlorate was recrystallized from isobutyl alcohol,
washed with anhydrous ether, and dried a t 100OC. The material was then crushed in an agate mortar, dried a t 250°C., and stored in a vacuum desiccator over anhydrone. The lithium perchlorate \vas made according t o the method 1 Presented before the Division of Physical and Inorganir Chemistry at the 110th meeting of the American Chrmical Soriety, Chicago, Illinois, September, 1946.
517
ISOTONIC SOLUTIONS
used by Scatchard and coworkers (4) for their freezing-point measurements. It was fused a t 300°C. to remove the last traces of moisture before each solution was prepared. c. P. sodium chloride was precipitated by hydrochloric acid, recrystalilsed twice from conductivity water, and dried a t 100OC. It was then crushed in an agate mortar and finally dried in a muffle furnace at 500-6OO0C.
Preparation of solutions All solutions were made to predetermined concentrations by weighing both dry salt and water. The solutions were added to the vapor-pressure cups from a weight buret, the weight of the sample being determined by difference. The TABLE 1 Concentrations of the isotonic solutions m/hTaC1
m/LiClOd
0.2267 0.2415 0.2459 0.3935 0.4283 0.5855 0.6343 0.8462 1.0429 1.3796 1.8530 1.9720 2.093 2.648 3.234 3.633 4.064 4.480 4.585 5.308 5.585
0.2178 0.2315 0.2361 0.3695 0.4000 0.5389 0.5785 0.7598 0.9208 1.1898 1 ,5572 1.6448 1.734 2.144 2.561 2.841 3.149 3.439 3.513 4.038 4.230
PATIO
1.0410 1.0431 1.0423 1.0650 1.0707 1,0864 1.0964 1.1138 1.1326 1.1595 1.18% 1.1996 1,2067 1.2351 1.2628 1.2788 1.2906 1.3027 1.3050 1.3151 1.3203
m/NaCl
m/NaCIOI
ut10
0.2003 0.2415 0.2458 0.4241 0.6186 0.8557 1.0328 1.1208 1.3415 1.0486 1.5493 1.7295 2.213 2.467 2.825 2.845 3.234 3.633 4.064 4.480 4.585 5.308 5.585
0.2013 0.2433 0.2473 0.4283 0.6284 0.8730 1.0574 1.1522 1.3800 1.4562 1.6018 1.8016 2.336 2.634 3.037 3.057 3.508 3.990 4.513 5.032 5.159 6.075 6.436
0.9951 0.9926 0.9940 0.9902 0.9845 0.9802 0.9767 0.9728 0.9721 0.9673 0.9672 0.9598 0.9474 0.9400 0.9302 0.9305 0.9219 0.9105 0.9005 0.8914 0.8890 0.8738 0.8678
-~
precision in weight was a t least 0.5 mg. The final weight of the solution after equilibrium had been established m s obtained by subtracting the known weight of the cup from the final weight of cup plus sample. The cups were provided with covers to reduce loss of water after they had been removed from the reaction vessel. Experiment has shown that the loss of weight from removal until time of weighing was usually of the order of 1 mg. or less. The final concentration of each solution was computed from the known initial concentration, the weight of the sample, and the loss or gain of water. DATA OBTAINED
The concentrations of the isotonic solutions are collected in table 1. In most of the case8 a t least two duplicate cups were used, and so the recorded concentra-
518
JAMES HOMER JONES
tions are the average of the two. These agreed with each other for the most part within 0.05 per cent. TRISATYENT OF D.kT.4
The osmotic coefficients of lithium and sodium perchlorates were calculated from the isopiestic ratios and known values for sodium chloridr, using the equation
Thr osmotic coefficients for sodium chloride were taken from thc' tabulat~ionI)? Stokes and Levien. Thr activity coefficients bvcre computed from the osmotic coefficients, using tho m~t~htrd outlined by Harned and O w n (1 1. Thc equation w e t 1 was
The activity coefficientsfor sodium perchloratr up to 1 molal itere also computed b y the original rquation of Robinson, whrrr the iubscript z refer%t o sodium
perchlorate and r to sodium chloride. The answers by both equations \\-ere identical. Very good agreement should be expected, since the value of thr integral in equation 2 up to esperimental Concentrations is very small ant1 could be missed by several pcr c r n t ivithout atrecting the answer appreciably. This is certainly not the case ivith salts such as lithium perchlorate, ivherr thr integral has a quite large value. r , l o evaluate t>heintegral in equation 1, it is necessary to use sonic extrapolation 1-4.
equation based on theory, since the value of the function ill7 ' 1- IF so srisceptible to experiment,al errors in dilute solut,ions. Such an equation may be derived The from the Debye-Huckel theory in the form of 1 - = 0.3888~,rn~/~. evaluation of urndepends on t,he finding of a suitable PL parameter (distance of closest approach of t,he ions) from the experimental data at the lower concentrations. For sodium perchlorate, the five lower concentrations gave a value of L = 4.4 A. ( k g s t r o m units). For lithium perchlorate, however, a very high value of 7-8 A. was indicated. This value is too high, indicating that the solutions a.t the lower concentrat,ions do not approach the theoretical value closely enough to be used. Since lithium perchlorate has activity coefficients close t,othose of lithium iodide, for which a value of &of 5.5 A . has been suggested, this value was also chosen as 3, for lithium perchlorate. .I linear term in con-
+
5 19
ISOTONIC SOLUTIONS
mntration was then added to the hitingequation to make it fit the experimental data a t the lower concentrations. The extrapolation was then made with the equation: 1 - 6 = 0.3888u,,,mljz - 0.148m
The u,,, values for various values of m were taken by interpolation from the table given by Harned and Owen (1). TABLE 2 Acfiuity and osmolic coe.
ienls
NsCIOd
LiClOI
m Q
0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 .o 1.2 1.4 1.6 1.8 2.0 2.4 2.7 3.0 3.4 3.7 4.0 4.5 5.0 5.5 6.0 6.5
0.9190 0.9142 0.9115 0.9102 0,9096 0.9100 0.9110 0.9122 0.9135 0.9166 0.9202 0.9244 0.9290 0.9338 0.9445 0.9528 0.9618 0.9742 0.9836 0.9925 1 .o090 1.0250
1.0412 1.0570 1.0745
7
0.728 0.701 0.681 0.667 0.656 0.647 0.640 0.634 0.629 0,621 0.616 0.612 0.610 0.608 0.608 0.610 0.612 0.617 0.622 0.627 0.637 0.649 0.662 0.675 0.691
0.9575 0.9710 0.9855 0.9995 1.0135 1.0285 1.O430 1.0575 1.0720 1.1030 1,1350 1.1680 1.2025 1.2375 1.3175 1.3680 1.4210 1.4935 1.5450 1.5985 1.6515
0,792 0.792 0.79'3 0.808 0.821 0.836 0.851 0.869 0.888
0.930 0.978 1.031 1.090 1.156 1.316 1.440 1.585 1.803 1.984 2.170 2.415
The activity and osmotic coefficients of sodium and lithium perchlorates are tabulated in table 2. DISCUSSION O F RESULTS
The general trend of the variation of activity coefficients with concentration can best be illustrated graphically. Figure 1 shows such a graph. In behavior these two salts are quite different. The activity coefficients of lithium perchlorate are distinguished by their high values. They pass through a minimum at about 0 2 molal, or just at the edge of the experimental region,
520
JAMES HOMER JONES
and increase rapidly with concentration. Among the uni-univalent electrolytes measured, these activity coefficientsare exceeded only by those of lithium iodide and hydriodic acid. The value of the activity coefficient in a 0.2 molal solution a t its freezing point is 0.805. The value of 0.792 a t 25% and 0.2 molal is in the right direction and of the correct order of magnitude. The activity coefficientsof sodium perchlorate go through a long flat minimum between 2 and 3 molal. The value of the activity coefficient in a 0.2 molal solution a t its freezing point is 0.720. The value of 0.728 at 0.2 molal and 25°C.
1
2
3
4
5
6
Molality FIG.1 is also in the right direction, and the difference is of the order of magnitude expected. BUMMARY
The concentrations of isotonic solutions of sodium perchlorate-sodium chloride and lithium perchlorate-sodium chloride were determined from approximately 0.2-5.5 molal sodium chloride. From the data obtained and the known values for the osmotic and activity coefficients of the reference salt, the corresponding values for sodium and lithium perchlorates were computed. Check computations on sodium perchlorate were made by two independent methods.
THE SYSTEM
Fe203-Cr203
521
REFERENCES (1) HARXED,H. S., A N D OWES, B. B.: The Physical Chemistry of Electrolytic Solutions. Reinhold Publishing Corporation, New York (1943). (2) JONES, J. H.: J. Am. Chem. SOC.66, 1353 (1943). (3) ROBIKEON, R. A., AND SINCLAIR, D. .4.: J. Am. Chem. SOC.66, 1830 (1934). G., PRENTISS, S.S.,A N D JONES,P . T . : J. Am. Chem. SOC.60, 803 (1934). (4) SCATCHARD, R . H . , A K D LEVIES,B. J.: J. Am. Chem. S O C .68, 333 (1946). (5) STOKES,
X-RAY DIFFRACTION STUDIES Ii\; THE SYSTEM Fe203-Cr203L W. 0. MILLIGAK
AXD
L. MERTEZi
Department nf Che?nzstrU, The Race Institute, Houston, Texas Recezced October 8 , 1946
An unusual mutual protective action has been observed in mixed gels of cupric and ferric oxides (4)and of nickel and aluminum oxides ( 5 ) . In both oxide pairs, the constituents mutually protected each other against crystallization, even a t high temperatures. Thus ferric oxide prevented or retarded the crystallization of cupric oxide, and nickel oxide prevented marked crystallization of aluminum oxide in samples heated below 1000°C. for a period of 2 hr. In this paper these results have been extended to include thesystem Fe203-Cr203,in which the components are closely similar in crystal structure and in lattice constants. EXPERIMENTAL
Preparation of samples Mixed gels of hydrous ferric and chromic oxides were prepared by the addition of an equivalent amount of ammonium hydroxide to mixtures of solutions of ferric nitrate (0.5 X with respect to Fe203)and chromic nitrate (0.5 M with respect to Crz03),using a rapid mixing device described elsewhere (IO). The amounts of the ferric and chromic nitrate solutions were chosen so that there was obtained a series of eleven samples containing 0, 10, 20, 30,40, 50, 60, 70,80, 90, and 100 mole per cent of ferric oxide. The dual gels were washed in a centrifuge until the supernatant liquid no longer gave a test for nitrate ions. After the moist gels were dried in air a t room temperature, separate portions of each of the mmples were heated for 2-hr. periods a t various temperatures. In a similar manner, a second series of eleven gels was prepared, using sodium hydroxide as the precipitant. In order to attempt to ascertain any possible effect of adsorbed sodium hydroxide or silica from the alkali solution employed, separate experiments were 1 Presented before the Division of Colloid Chemistry a t the 109th Meeting of the American Chemical Society, which was held in Atlantic City, New Jersey, April 8-12,1946.
