940
NOTES
T'ol. G6
TABLEI HYDROLYTIC DEQRADATION D A T A - N ~ ~ P , O ~ ~
Ref.
Chemical state
R. N. Bell
Temv.. ..
Aqueous soh. (1 w/o) (10 w/o) (1 w/o) Aqueous soln.
0. T. Quimby
(1 w/o) (10 w/o) A. G. Buyers
Aqueous soh. (14 w/o)
0. T. Quimby
Solid Na6P8Ol06H20
A. C.
Solid NasP30106H20
Zettlemoyer, et a1.b
OK.
IOa/Temp.
Decimal fraction NaaPrOlo remaining at 30 min.
373
2.68
373 343 298
2.68 2.92 3.36
355 394 385 381 423 393 378 368
2.82 2.54 2.60 2.62 2.36 2.54 2.65 2.72
393 388
2.54 2.58
.70 .85
383
2.61
.88
Pa01/P04 for the temp. indicated
0.96
1.13
.89 99 .96a
*
1.12 1.70
..
..
.99 .37 .74 .77 *
Estimated av. mole ratio
18
.58 .88 .96
-1
1.78 1.30 1.24 1.34 1.63 1.63 1.35
Rate constant X 101 (2nd order) dec. fraot.-', min. -1
0.13 .40 '03 .00002" .03 5.66 1.17 1.00 15.18 2.4 * 47 .13 1.43 .60
1.1
.47 --j
373
2.68
to
.90
.37 1.73
368 358
2.72 2.79
.93 .97
.27 .10
-_I
Calculated from 146-day analysis. remaining data by these authors.
Data for 378" were not tabulated as these values were not consistent with the
0
I
lo-'
.-
AQUEOUS SOLUTION 13 3 SOLID No5 p3Olo6H20
A 3 SOLID Nag f?0,06H,0
IO4
z
U k-
I n
0 0 W
k-
U (z
1 6 ~
w
5
0
z"
IO5
0 0 W
hydrolytic process is the breaking of the P-OP bond, a fact not established, but plausible. It then is interesting to note that apparently the energies of activation are not different for first or second-order interpretations. First-order mechanisms, eq. 1or 4, require that initially one P-OP bond be broken by hydrolysis. The second-order expressions, eq. 2 or 3, suggest that two P-OP bonds are broken and one P-OP bond is formed, representing a net energy change for breaking one bond, ie., 20-25 kca1.l mole.6 Admittedly, the mechanism for the hydrolytic degradation of sodium triphosphate to orthophosphate has not been established by this note. However, it is believed that the examined hydrolysis data and their interpretation render dubious the simple hydrolytic mechanism depicted by eq. 1. Furthermore, evaluation of available research has indicated that the initial step in this hydrolysls can be described in terms of a second-order reaction. ( 8 ) J. R. Van W a ~ e rEncyclopedia , Chem. Technol., 10,403 (1953).
IO+
T H E ACTIVITY COEFFICIENTS OF ,4MMONIUM PERCHLORATE I N WATER AT 4 1 5 O 2.3
25
2.7
2.9
3.1
3.3
1 3 T X IO
Fig. 1.-Arrhenius relation for second-order reaction based upon the disappearance of sodium triphosphate hexahydrate.
BY 0. E. ESVAL AND S. Y. TYREE, JR. Department of Chemistry, University of North Carolina. Chapel Hall, h'orth Carolina Receiued October I d , 1981
The extensive measurements of activity coefficients for 1-1 electrolytes indicate that the activity
May, 1962
NOTES
941
TABLE I coefficients of animonium perchlorate should be similar to those of ammonium nitrate.' Harned ISOPIESTIC SOLUTIONSOB AMMONIUM PERCHLORATE AND and O ~ e have n ~ made a rough classification of the POTASSIUM CHLORIDE AT 25" thermodynamic properties of 1-1 electrolytes based m"4cIO4 mECl mNHICIO4 mECl on their ion size and structure. Basically it can 0.104 0.101 1.495 1.106 be assumed that any deviation from the Debye.205 .197 1.544 1.346 Huckel equation can be attributed to the hydration ,283 .270 1.602 1.394 energy of the cation or anion. This is directly .450 .420 1.670 1.445 related to the charge, size, and symmetry of the I.730 1.488 .603 ,558 ion. In addition it can be shown that an ion with .759 .703 1.811 1.550 a low hydration energy will tend to have a low .918 .837 1.go5 1.622 value for a given activity coefficient as compared to 1.091 .983 2.006 1.692 an ion which is more highly hydrated. Therefore, 1.241 1.106 2.100 1,768 the perchlorate ion should have a slightly lower 1,455 1,268 hydration energy than the nitrate ion, since the perchlorate ion is more symmetrical and larger calculated from the experimental data by the while having the same charge. Thus it may be relation' concluded that the activity Coefficients of ammod = mnR$R/m nium perchlorate should be slightly lower than where qi is the osmotic coefficient of ammonium those of ammonium nitrate. perchlorate a t a concentration m, and q i ~is the The osmotic coefficients of ammonium per- osmotic coefficient of potassium chloride, the chlorate were obtained by the isopiestic method of reference salt, a t a concentration mR. The osmotic Sinclair and R0binson,*~4from which the activity coefficients of potassium chloride were those as coefficients of ammonium perchlorate were calcu- listed in the new edition of Lewis and RandalL8 lated over a range of 0.1 to 2.1 m. These values are 0.0020 greater than the values tabulated by Robinson and stoke^.^ The physical Experimental Materials.-T wo preparations of ammonium perchlorate constants for the Debye-Huckel limiting equation were used, one prepared by the method of Hartley6 and one also were obtained from the same source.l0 obtained from J. T. Baker and Company. Both preparations were recrystallized three times with demineralized water (specific conductance 10-7 em.-' ohms-') or until a negative chloride test was obtained as indicated by silver nitrate. The salt was dried in a vacuum desiccator over PZOS instead of an oven at 110" in order t o prevent possible slow decomposition of the perchlorate into the chloride over long periods of time. Both prepsrations were analyzed for chloride by reduction of the perchlorate to chloride by the Parr bomb and gravimetric determination of the resulting chloride. Three samples were run in this manner and all agreed to aithin a few tenths of 1 %. Anal. Calcd. for "&IO4: C1, 30.17. Found: C1, 30.37. Reagent grade KC'1 was used without any further purihcation. The salt wm dried in an oven at 110'. The apparatus for the isopiestic equilibrations and the experimental technitque involved have been previously described.6
Results and Discussion The isopiestic solutions of potassium chloride and ammonium perchlorate are listed in Table 1. Deviations greater than 3 parts per thousand between the ammonium perchlorate solutions or the potassium chloride solutions were not used. This happened in only two instances out of 21 experimental points; in both cases, it appears that splattering of the solution from the cups had occurred. The osmotic coefficients and the mean molal activity coefficients of ammonium perchlorate are listed in Table 11. The osmotic coefficients were
TABLE I1 OSMOTICAND MEAN MOLALACTIVITYCOEFFICIENTS OF AMMONIUM PERCHLORATE AT 25" m
c
0.1 .2 .3 .4 ,5 .6 .7 .8 .9
0.901 .879 .861 .848 .838 .834 ,825 .820 ,816
Y*
m
0.730
1.0 1.2 1.4 1.6 1.8 2.0 2.1
.662
.617 .583 .556 .540 .521 .507 .493
c .811 0.801 .791 .783 .775 .767 .756
Y b
.482 0.460 .442 .426 ,411 .398 .394
The activity coefficients were determined by" where a plot of (1 - +)/m''Z z's. ml/s is sufficient t o evaluate the integral of the above equation and consequently log yq. However, since the range of concentration of the experimental data was 0.1 to 2.0 m, an extrapolation function was needed which would extend the plot of (1 - qi)/m'/a vs. ml/z to zero concentration. The extrapolation function used is the one based on the Debye-Huckel equation and described by Harned and Owen.12 In addition, the equation used for extrapolation was found to give a very close fit over the entire range of the experimental results. By using a value of 0.606 for Am', the deviation from the experimental
(1) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," 2nd Ed., Butterworthn Scientific Publ.. London, England, 1959,p. 495. (2) H. S. Harned and 13. B. Owen, "The Physical Chemistry of Electrolytic Solutions," 2nd Ed., Reinhold Publ. Corp., New York, N. Y., (7) Reference 2. p. 288. 1950,p. 394. (8) G. N. Lewis and M. Randall, Revised by K. 8. Pitzer and L. (3) D. A. Sinclair, J . r h u s . Chem., 37,495 (1933). Brewer, "Thermodynamica," 2nd Ed., McGraw-Hill Book Co., New (4) R. A. Robinson a n 3 D. A. Sinclair, J. Am. Chem. Soc., 56, 1830 York, N. Y., 1961. p. 643. (1934). (9) Reference 1, p. 484. (5) H.Hartley, Proc. Roy. SOC.(London), A132, 429 (1930). (10) Reference 8,p. 640. (6) C. S. Patterson, 8. Y. Tyree, and K. Knox, J. Am. Chem. SOL, (11) Reference 2, p. 292. 77, 2195 (1955). (12) Reference 2, pp. 289-294.
942
NOTES
gregate t o form micelles. Below the c.m.c. it generally has been assumed that, in the absence of hydrolysis, such electrolytes consist of simple ions though the suggestion has been made from time to time that aggregation occurs below the c.m.c. We have made accurate measurements of the electrical conductivity of aqueous solutions of sodium dodecyl sulfate at 25’ a t concentrations below the c.m.c. domi to 4 X M in order to investigate the extent of aggregation of anions to form dimers or small pre-micelles which, alone, presumably would increase the conductivity, or of association to ion pairs which would decrease the conductiyity. The results have been analyzed in the light of the theory of electrolytic conductance given by FLIOSS and Onsager , 2
0.0
\
-0.1
+ii 8C
Vol. 66
-0.2
-0.3
-0.4
0.5
1.0
1.5
4112.
Figure 1.
log y& was found to be =kO.OOl over a range of 0.1 to 1.2 m and .tO.OlO up, to 2.0 m. This value of A,’ corresponds to 1.84 A. for (2, the mean distance of closest approach. Use of the further extended Debye-Huckel equation with the added term of B, did not give a better fit after the proper adjustment of Am’ and B. i l s is expected for ammoniuni salts, the value of B is much lower than the values for the alkali halides, which range between 3.5 and ~ to indicate 6.2 8. The small value of B T V O L I ~seem that other important effects are present in addition to the very low hydration energies of the cation and anion. The logs of the mean molal actiyity coefficients of ammonium perchlorate along with some other 1-1 electrolyte~l~ are plotted in Fig. 1. It can be seen that the plot is slightly lower than that of ammonium nitrate, as m s expected. The L. L. line represents the Debye-Huckel limiting slope. (13) Reference 1, pp. 479-480.
CONDUCTIT’ITY OF SODIUM DODECYL SULFATE SOLUTIONS BELOW THE CRITICAL MICELLE CONCEKTRBTION BY G. D. PARKITT AND 8.L.
SMIT€ll
Department of Chemistrg, linzversity o f ATottingham,Kottzngham, Endand Recezved October 16, 1.961
Colloidal electrolytes such as sodium dodecyl sulfate shorn a more or less abrupt discontinuity in physical properties over a relatively short concentration range termed the critical micelle concentration (c.m.c.). Above this concentration it is well established that the amphipathic ions ag(1) A t the time of this vork a t the College of Technology, Northampton, England.
Experimental The sodium dodecyl sulfate was a pure sample kindly supplied by Thomas Hedley & Co., Ltd., having a purity of 99.9% as determined by partition end-point titration.8 It was purified further by a liquid/liquid extraction technique.* The water used was obtained from an ion exchange column and, equilibrated with air, had a conductivity 1.1 X 10-aohm-l cm.-l. Conductivities were measured on a conventional 1000 cycles/sec. bridge incorporating a Wagner earth using resistance boxes of 0.05% grade. The cell, similar in design t o that of Flockhart and Graham,6 was of 400 ml. capacity and required 50 ml. to cover the electrodes. Dilution additions were by weight and concentrations calculated from measured densities. Equilibrium was reached about 30 min. after dilution. Resistances were taken after about 1 hr. and remained constant for a t least 7 days.
Results and Discussion The results are shown in Fig. 1 by a plot of A us. c X lo3 which is of the form expected for a 1:l strong electrolyte. Both the internal consistency and the agreement between the three overlapping series of measurements were -0.05%. Fuoss and Onsager2 give for an unassociated 1:1 electrolyte (neglecting a small viscosity term) A = Ao - SC‘/Z-j- EC log c JC (1) where X is the Onsager limiting slope, E is a function of .io and the solvent properties, and J is a function of do,the solvent properties, and the ion size parameter “a” For an associated electrolyte this becomes
+-
A =
A0
- SC‘/Zy ’ / z f ECy IOg cy
+
JC+y
- KaCyf’h
(2)
Treating the sodium dodccyl sulfate as a simple unassociated 1:l electrolyte do was found from the data by a Shedlovsky extrapolation6 so that X and E could be calculated and eq. 1 applied. A plot of -1 SC’/~- Ec log c us. c (Fig. 2 ) is straight up t o concentrations very close to the c.m.c., after which an abrupt change is obvious. The slope, J , of the portion below the c.m.c. is $95, which corresponds to a value for “u” of 5.0 A. This is rather less than the aT-erage value of 5.5 8.found’ for this electrolyte in dioxane-water mixtures a t lower dielectric constants with appreciable values of the
+
(2) (a) R. M.Fuoss and L. Onsager, J . Phvs. Chem., 61,868 (1957); (b) R. M. Fuoss and L. Onsager, sbzd., 62, 1339 (1958). (3) T.Barr, J. Oliver, and W. V. Stubbings, J . 8 o c . Chem. Ind.. 67, 45 (1948). (4) S.P. Harrold, J. CoElozd Sez., 15, 280 (1980). (6) B. D. Flockhart and E. Graham, %bad.,4,367 (1949). (6) T.Shedlowky, J . Am. Chem. Soc., 64, 1405 (1932). (7) G.D.Parfitt and A. L. Smith, awaiting publication.
