NOTES
3356
Table I ---DV la -
1000eu
Itr
0.5 1.0 2.0 3.0 4.0 a
Dv X 106
0.969 0.941 0.891 0.848 0.812
1.022 1.044 1,074 1.111 1.140
From St.okes Table I.
0.983 0,971 0.954 0.952 0.955
0.981 0.960 0.916 0.870 0.822
= -C,[aqr
+ 2(1 - c~)q.z]Vp,
(10)
NOW,in the derivation of (1) and (2), it is assumed that Ju
(exptl)
1.344 1.305 1.234 1.163 1.107
1.340 1,308 1,244 1.188 1.144
DW*
1.340 1,303 1.238 1.192 1.160
With Dw*/qr
1.340 1.301 1.231 1.175 1.126
Using 41/42 = 0.7.
whence
J,
D,* x 105
X 10s (ca1cd)-
With
=
-cuquVpu
(3%)
and D,* is evaluated from D,* = q,RT
(54
By comparison of (10) and (3a), it f o l l o w s that a modified form of (1) and (2) is obtained by replacing qu with the term in square brackets in (10). Further, as D,* is now given by (7), it can be introduced by replacing Du* with D,*.f(a), where
y u given by Bower and Robinson,’ but the differences are not large enough to affect the calculation in Table I. The self-diffusion coefficient of water in urea solutions has not been measured. The value for water cmz/sec, and, as possible itself is about 2.5 X extremes, either this has been used or the value obtained by dividing by the relative viscosity of the solution, vr, using values given by Stokes. Values of concentrations below 0.5 mole/l. have not been included as D,* and DV are very close in this range. It can be seen that the experimental values of Dv lie on the whole between the two calculated values, up to the highest concentration of 4 moles/l. (7) 1‘. E. Bower and R . A. Robinson, J . Phys. Chem., 67, 1524 (1963).
Le., eq 2 modifies to
+
b In a, Dv = pw-(cuDu**f(a) cUDW*) (12) b In c, I n Albright and Mills’ paper, experimental values of
Du*are given, as well as expressions for Dv and Vu. From the latter, taking pw as 18.04 cc for the whole range of concentrations, cw can be derived, as cwVw = 1Stokes gives values of a,calculated on the assumption that activities of monomer and of dimer equal their respective mole fractions. With regard to the relative mobilities of monomer and dimer, a reasonable choice for qz/ql, following considerations advanced by Stokes, is 0.7, but f(a) is not very sensitive to this over the range of concentrations considered and gives almost the same vaIues for ratios between 0.6 and 0.8. The values for ( b In a,)/b In c,) in Table I have been calculated from a, using Stokes theory that monomer and dimer form ideal solutions. They are almost identical with the values calculated from In yu, as given by Albright and Mills, and tend to be a little lower than those calculated from the experimental values of
curu.
The Journal of Physical Chemistry
Effect of Inert Gas Pressure and Solubility on Fused Salt Conductance. 11.
Nitrogen with Sodium Nitrate by James L. Copeland and Steven Radak Department of Chemistry, Kansas State University, Manhattan, Kansas 66602 (Received M a y 2, 1966)
As a continuation o f the studies of Copeland and Zybko1*2on the manner in which the specific conductance of simple fused salts is affected by “inert” gas pressure and solubility, we wish to report our rather unusual and interesting results for Nz in high-pressure equilibrium with molten NaN08. (1) J. L. Copeland and W. C. Zybko, J . Am. Chem. Soc., 86, 4734 (1964).
(2) J. L. Copeland and W. C. Zybko, J. Phys. C h a . , 70, 181 (1966).
NOTES
3357
Experimental Section The general apparatus and method are identical with those reported in ref 2. The cell constant was 384.2 cm-l, as calculated by assigning an interpolated literature valup of 1.230 ohm-l cm-l to the specific conductance of NaN02 at 369” and atmospheric pressure.2~3 Reagent grade NaN03 was obtained from the Baker and Adamson Co., as in the prior workU2 Nitrogen was at Ieast 99.98% pure (Bone Dry grade), and was procured from the National Cylinder Gas Co. A constant salt temperature of 369” was employed to allow comparison with the He and Ar studies in the same system at the same temperaturee2 The initial Nz pressure at this salt temperature was 379 atm. After equilibration and the first electrical resistance measurement, increments of Nz pressure were released at intervals, the system was allowed to equilibrate at each new pressure (with frequent agitation), and equilibrium resistance measurements were made. The “zero” pressure measurement was made with the system being agitated and pumped by a mechanical vacuum pump. The entire experiment was performed over a period of 9 days. The solubilities of N2 in fused NaN03 were taken from the work of Copeland and Seibles.4
Results Empirically, the joint effect of increasing NZpressure and solubility was to depress linearly the specific conductance, K , of the WaT\’08at 369” over the range of pressures employed. The data are summarized in Table I. The empirical least-squares relationship resulted as KNz -
(1.231 f 0.001) - (7.59
f
0.18) X
P N ohm-’ ~ cm-1
(1)
where P N is ~ the saturating pressure of N2 in atmospheres. The errors are the least-squares probable errors.
Table I : Summary of Specific Conductance of NaNOa at. 369” under NZPressure Saturatingpressure, P N ~atm , Specificconductance, K N ~ , ohm-’ em-’
379
309
236
162
69
0
1.201 1.208 1.213 1.218 1.226 1.230
The Henry’s law constant for NaNOSat 369” is4
NZsolubility in fused
KN*= (19.7 f 0.4) X
mole of Nz ~ m atm-1 - ~
(2) where the error is the probable error for a single solubility-pressure point. Relationships 1 and 2 are to be compared to those for Ar and He in the same system, as summarized below.216 KAr
= (1.228
0.002) - (1.94 f 0.06)
x
lov4Par ohm-’ cm-l
(3)
K A =~ (17.2 f 1.7) X lo-’ mole of Ar cmp3 atm-I
(4)
K H ~ =
(1.230 f 0.001) - (9.19 f 0.25) X
loW6 Pne ohm-’ cm-l
(5)
K H =~ (22.7 f 0.7) X lo-’ mole of He omd3 atm-l
(6)
The striking features to be noted are that whereas the solubility of Nzlies between those of Ar and He (cf. eq 2, 4, and S), the effect of Nz in diminishing K is less than those of the other two gases (cf. slopes of eq 1,3, and 5). Discussion The dilution mechanism, hypothesized earlier to account for the different extents of the depression of K by the solubilities of different nonreactive gases,2may still play a dominant role in the system under study here, but may not be the entire story. With an assumption of ideal dilution effects alone, as a consequence of the effect on K we would have to conclude that E2 dilutes the melt to a lesser extent than does either Ar or He, even though the Nz solubility is intermediate between those of the other two gases. This implies, of course, a smaller partial molal volume for Nz than for Ar or He in this system. This is not unrealistic in view of the negative enthalpy of solution of -2.73 f 0.09 kcal mole-’ for Nz in ?SaNOs observed by Copeland and S e i b l e ~ . ~ On the other hand, the large negative standard entropy of solution of -16.6 f 0.1 eu4 may imply some subtle ion-molecule interactions which contribute to the effect on K in an undetermined manner. Detailed determinations of the densities of such gas-salt solutions, for calculations of equivalent conductance, should cast more light on the question of the presence or absence of such influences of solute molecules on the ions. Acknowledgments. The authors gratefully acknowledge support of this work by the National Science (3) A. Klemm, “Molten Salt Chemistry,” M. Blander, Ed., Interscience Publishers, New York, N . Y.,1964,p 566. (4)J. L. Copeland and L. Seibles, J . Phys. C h a . , 70, 1811 (1966). (5) J. L.Copeland and W. C. Zybko, ibid., 69, 3631 (1965).
Volume 70,Number 10 October 1966
3358
Foundation, Grant GP-4274. This work will constitute a part of the thesis of S. R. which will be submitted to the Graduate School of Kansas State University in partial fulfillment of the requirements for the degree of Master of Science.
Pulse Radiolysis of Anhydrous Amines by Larry R. Dalton, James L. Dye, Department of Chemistry, Michigan State University, East Lansing, Michigan
E. R4. Fielden, and Edwin J . Hart Chemistry Division, Argonne Katwnal Laboratory, Argonne, Illinois (Received May 2, f966)
In their original studies, Hart and Boag' demonstrated that the hydrated electron absorption band is altered considerably in 12.2 M aqueous ammonia and in 12.5 Ai' methylamine solutions. In each case the transient absorption peak formed during pulse radiolysis lies beyond 9000 A, the limit of sensitivity of their photographic plates. From the time of these first spectrophotographic observations of the hydrated electron,' it has been assumed that solvated electrons produced by the radiolysis of water and other solvents are analogous to those formed when alkali metals are dissolved in liquid ammonia and amines. Comparison of the spectra of metal solutions with those of the transients formed during pulse radiolysis of the same solvent should form a good test of this assumption. Compton and co-workers2 have shown the spectrum obtained by pulse radiolysis of anhydrous ammonia to be similar to that of metal-ammonia solutions. On the other hand, Anbar and Hart3 reported an absorption maximum a t 9200 A for the transient produced by pulse radiolysis of anhydrous ethylenediamine. They attributed this band to the solvated electron in disagreement with the assignment of Dewald and Dye4 based upon the spectra of alkali metals in et,hylenediamine. The latter investigators found no peak NnnmOn to the various alkali metals in this region and assigned an absorption maximum a t 12,800 A to the solvated electron and loosely bound aggregates of the electron with cations. This disagreement, and the observation of an anomaly in the hydrated electron absorption spectrum in the regon of solvent (water) absorption bands5 (7500 A), prompt,ed US to reexamine the spectra of the transients produced by puke radiolysis of anhydrous The J O U Tof~Physical Chemistry
NOTES
amines. I n the case of the hydrated electron in water, the anomaly had been shown to be an artifact essentially produced by the effect of scattered radiation on the particular photomultiplier used, complicated by effects due to different illumination intensities. Although an apparent decrease in the radiation-induced absorbance occurs near the ethylenediamine solvent absorption band (10,500 A), by minimizing stray radiation and operating the photomultiplier under constant illumination conditions in a range of proven linearity, the previous results3 were shown to be incorrect. For each of the amines tested, the absorbance continued to increase up to the cutoff wavelength of the photomultiplier used (11,200 A). The band shape is similar to that attributed to the solvated electron in metal solutions in eth~lenediamine,~ although the comparison suffers from our inability to search far enough into the infrared to locate the absorption maxima in the radiolysis studies. Experimental Section Anhydrous ethylenediamine (Matheson Coleman and Bell) was first purified by fractional freezing and . ~ was distillation as described by Feldman, et ~ 1 This followed by two distillations in vucuo into a quartz bulb attached to the quartz irradiation cell. The first of these final distillations was from a sodium mirror and the second was from an intermediate vessel containing no metal. I n this way, the carry-over of ions by the spray can be avoided. A second sample of ethylenediamine was prepared using a potassium mirror in place of the mirror of sodium. A similar purification procedure was used to prepare the sample of 1,3-propanediamine ("Itheson Coleman and Bell). Ethylamine (Eastman Organic Chemicals) and n-propylamine (Matheson Coleman and Bell) were purified by two distillations in a nitrogen stream followed by two distillations in vacuo as described above. The absorption spectrum of the solvated electron in these amine solvents was determined by examining the decay of absorbance of the transients produced by a 0.4-psec pulse of 15-34ev electrons. The pulse intensity was estimated by determining the absorbance (at 7000 (1) E. J. Hart and J. W. Boag, J. AWL Chem. Sac., 84, 4090 (1962). (2) D. hf. J. Compte;: J. F. Bryant, R. A. Cesena, and B. L. Gehman in "Pulse Radiolysis, M. Ebert, J. P. Keene, A. J. Swallow, and J. H. Baxendale, Ed., Academic Press, London and New York, N. Y.. 1965, D 43. (3) hl. Anba; and E. J. Hart, J. Phys. C h a . , 69, 1244 (1965). (4) R. R. Dewald and J. L. Dye, ibid., 68, 121 (1964). (5) E. M. Fielden and E. J. Hart, unpublished results. (6) L. H. Feldman, R. R. Dewald, and J. L. Dye, Advances in Chemistry Series, No. 50, American Chemical Society, Washington, D. c.,1965, 163.
