NOTES
1168
is fixed and that of the perchlorate varied (Fig. 3), a slope of 0.5 for a tenfold change in concentration is observed. In both cases, a linear relationship is present as the Concentration of one of the moieties is varied while the other is maintained constant. These results may be rationalized by assuming that sodium perchlorate undergoes greater dissociation than sodium acetate (a greater degree of "ion-pair" separation), hence the sodium ion concentration of the solution may be attributed mainly to the sodium perchlorate. By expressing the dissociation constants for the two salts as K N ~ C I=O(Naf)(C104-)/(NaCl04) ~
(1)
K ' N ~ =A (Na+)(Ac-)/(NaAc) ~
(2)
and employing the simplifying assumption6 above that the concentration of sodium ion is determined essentially through the dissociation of sodium perchlorate, we obtain the relationship
+ 3)
(2
(Na +) E { K( NaClO4)] '/n (Ac-) S K'(NaAc)/K'/z(NaC104)'/2
(3)
(4)
Equation 4 shows a first-order dependency on the sodium acetate concentration and an inverse square root dependency on the sodium perchlorate concentration. Since experimental data are in close agreement with those predicted by equation 4, it may be concluded that sodium perchlorate is a ('stronger electrolyte" than sodium acetate in acetic acid, hence repressing the dissociation of sodium acetate, resulting in a decrease in the relative basicity of the solution.
VOl. 58
It is evident from these behaviors that the common practice of estimating pK values of acids and bases in water by determining the pH of half neutralized solutions cannot be used in certain nonaqueous systems. The apparent basicity of a base! as in the present instance, is dependent in part on the strength of the acid employed in its neutralization. Stronger the acid, weaker the base would appear since the resulting salt formed would be more highly dissociated, leading to a greater suppression of the dissociation of the parent base. Experimental Analytical grade reagents were used throughout. Sodium acetate solution, 0.05 molar, was standardized against perchloric acid with quinaldine red as indicator. Sodium perchlorate, 0.025 molar, was prepared in situ by mixing equal volumes of 0.05 molar acetic acid solutions of perchloric acid (previously rendered anhydrous with acetic anhydride) and sodium acetate. All other solutions were prepared from these stock solutions. The potentiometric measurements were carried out using a Beckman Model G pH meter, e uipped with a glass electrode and an external calomel eqectrode joined with the solution by means of an agar-gel bridge (4% agar and 5% potassium chloride). All potentiometric readlngs were determined relative t o that of a 0.100 molar solution of sodium salicylate in acetic acid taken just prior to the determination of the unknown solution. Duplicate determinations on independently prepared solutions agreed, in most part, within one or two millivolts.
DENSITIES OF MOLTEN SODIUM AND RUBIDIUM HYDROXIDES BY DONALD BOQART N A C A Lewis Flight Propuleion Lab., Cleveland, Ohio Received J u l y 86, 1964
The densities of molten sodium and rubidium hydroxides have been determined up t o temperatures of 920" using a method based on Archimedes' principle of buoyancy. The present data for sodium hydroxide in the temperature range 690 to 920" are shown in Fig. 1.
0.02000
2 0.01000 8 0 J
a8$ 0.00500 a
s
3
I
m 0.00250 ' -
I.. .____
60 80 100 120 Millivolts. Fig. 3.-Potentiometric determination on sodium acetatesodium perchlorate system; concn. of sodium acetate in M : ( l ) ,0.00125; (2), 0.00250; (3) 0.00500; (4),0.01000.
40
(5) This assumption may be shown valid b y employing the alternative case, that the sodium ion concentration from the acetate is greater than that from the perchlorate, whereby experimentally determined data do not conform to the derived equation.
The data for rubidium hydroxide (measurements taken on two successive days) are shown in Fig. 2. (1) K. Arndt and G. Ploetz, Oak Ridge Nat. Lab. Index No. Y-F35-5, Mar. 21, 1952 (translated from Z . physik. Chem., 121, 439 (1926)). (2) M. Nishibayashi. Wright Air Development Center Teohnical Report 53-308, Nov. 1953.
e
3a3
1169
NOTES
Dec., 1954
400
500 800 TEMPERATURE, ' G .
Fig. 2.-Density
700
800
SkI
of molten RbOH.
On the first day, the rubidium hydroxide was found to melt gradually over the temperature range 200 to 250". Density measurements from 250 to 500" exhibited the reversal shown which may be attributed to the change in density encountered by driving off water from the partly hydrated hydroxide. For higher temperatures the density variation was essentially linear. During the second set of measurements, rubidium hydroxide melted above 350". Density measurements in the range 400 to 650" being in agreement with the linear portion of the curve obtained from the first data, indicates no further changes in composition. The density-temperature relationship may be expressed as
360;
10
20
30
40
50
60
COWNG TIME, MIN,
Fig. 3.-Cooling
curve for RbOH.
was washed out, a considerable quantity of black crystalline material remained which had been formed during the cooling period of about 10 hours; this would not have interfered with the density measurements which were completed during a heating time of about two hours. Loss of weight due t o corrosion of the nickel plummet was negligible. Rb0H.-White vapors were again observed leaving the crucible a t temperatures above 500". The nickel wire supporting the plummet failed a t a temperature of about 900". The melt was further heated t o about 1000" a t which temperature the thermocouples failed, No evidence of boiling wa.s p = 3.11 - 0.000782' observed and the melt was light yellow and similar The melting point of the dehydrated RbOH was more accurately determined by the cooling curve in appearance to NaOH a t these temperatures. shown in Fig. 3. The transition occurred a t 383"; Small black crystals similar to those mentioned this value disagrees with a melting point of 300" in above were found when the crucible was washed the l i t e r a t ~ r e . Since ~ in the present investigation out. Qualitative analysis of these crystals indicated a lower melting point was obtained with undried the presence of some nickel carbonate. material, it may have been possible that the hydroxide used in the earlier experiments3 was not completely dehydrated. THE CATALYTIC ACTIVITY OF BISULFATE Experimental ION I N THE HYDROLYSIS O F ETHYL Both hydroxides were obtained from commercial sources. ACETATE1
Analysis indicated the following compositions, given in per cent. by weight: NaOH 96.2% RbOH 90.4% RbzCOa 2.1% NazCO, 1.0% HzO 7.5% HzO 2.8% Each hydroxide was contained in a cylindrical nickel crucible and was gradually heated in a tube furnace controlled through a rheostat. The hydroxides were exposed to air during the tests. Cylindrical nickel plummets were suspended in the molten hydroxides by nickel wire. The loss of weight of the plummets due t o the buoyancy of the hydroxides was measured with a precision analytical balance. The volume of the plummets (approximately 5 cc.) was corrected for thermal expansion; this correction is about 5% a t 1000". Temperatures were measured by means of two chromelalumel thermocouples spot-welded to the outer wall of the crucible and connected to a direct-reading potentiometer. Temperatures were also measured by immersing a thermocouple directly into the hydroxide; all temperatures agreed within 10'.
Discussion NaOH.-White vapors were observed above the crucible during the tests. When the nickel crucible (3) G. Von Hevesy, Oak Ridge Nat. Lab. Index No. Y-F35-6, Apr. 1 , 1952 (translated from R . phueik. Chem., 7 8 , 667 (1910)).
BY DENNISW. BARNUM AND GEORGE GORIN Contributed from the Department of Chemistry, University of Oregon, Eugene, Oregon Received August 9,1964
A very large amount of research has been done upon the acid-catalyzed hydrolysis of esters, but there still exists a difference of opinion on whether the reaction is susceptible to acidic species generally, or only to hydrogen i0n.~13 The very existence of such uncertainty suggests that the catalytic activity of acidic species other than hydrogen ion, if appreciable, cannot be large. It is therefore surprising to find that the catalytic constant of bisulfate ion has been estimated as 1.5 X mole-' min.-', about one-fourth as (1) Presented a t the Northwest Regional Meeting, American Chemical Society, Moscow, Idaho, June 13, 1953. (2) I t is understood that the hydrogen ions are solvated; in this and to paper, the term "hydrogen ion" refers to the species HsO', more or less hydrated forms. (3) R. P. Bell, "Acid-Base Catalysis," Oxford University Press, London, 1941, p. 48-81, especially p. 80-81; K. J. Laidler, "Chemical Kinetics," McGraw-Hill Book Bo., Inc., New York. N. Y., 1950, p. 282-300.
