128
ROBERT R. DEWALD AND JAMESL. DYE
Solubilities and Conductances of Alkali Metals in Ethylenediamine
by Robert R. Dewald’ and James L. Dye k‘edzic Chemical hboratory, Michigan State University, Euat Lanuing, Michigan
(Receiaed August 12, 1068)
The condurtances of cesium, rubidium, potassium, and sodium in ethylenediamine have heen measured a t room temperature as a function of concentration. The equivalent conductance for solutions of cesium, potassium, and rubidium showed the same general concentration dependence while that for sodium solutions exhibited a marked difference in behavior. The apparent limiting values of the equivalent conductances of cesium, potassium, and rubidium solutions were evaluated using the Shedlovsky conductanre function, and values of 204, 139, and 117 kohlrausch units, rcspectively, were obtained. The solubilities of the alkali metals in cthylenediamirw were also determined.
Introduction The absorption spectra of solutions of lithium, sodium, potassium, and cesium in liquid ammonia measured up to 0.03 M a t the absorption maximum and up to 0.2 M a t other wave lengths have recently been reported by Gold and Jolly.2 Their results showed that the alkali metal spectra are practically identical and follow Beer’s law from 350 to 1800 mp over the concentration range investigated. Conductivities of dilute solutions of sodium and potassium in liquid ammonia are nearly the same,3 and the small differences which do occur may be attributed to the different conductivities of the solvated metal ions. Metal solutions in amine solvents exhibit absorption spectra which are markedly different from those of metal-ammonia solutions. 4-6 Spectroscopic data for solutions of metals in amines and in mixed solvents show the presence of one to three very intense absorption bands depending upon the alkali metal. One of these is in the visible region (-600--700 mp), one in the nearinfrared region (-800-1000 mp), and the other in the infrared region (-1200-1500 mp), while metal-ammoriia solutions exhibit only an infrared absorption band with a maximum near 1500 mp, Conductivities of the different alkali metals in a given amine solvent cannot be readily compared because of the limited conductance data available. Gibson and I’hipps’ reported the conductance of potassium and cesium in methylamine, but their data have been questtioned* because of the probable presence of impurities which tended to dwompose their solutions. The Journal o j Physical Chemistry
Evers, Young, and Pansons made a careful study of the conductance of lithium in methylamine a t high concentrations. Later, Berns, Evers, and Frankg successfully reproduced their experimental data in more dilute solutions using a conductance function that had been previously derived by Evers and FrankLoand used to fit the data of Kraus3 for solutions of sodium in ammonia. This conductance function uses the same mass action equations as Reeker, et al.,lI and a modified form of the Shedlovsky conductance function. Since the different metal solutions in a given amine solvent have different spectra, it was thought that they might also show differences in conductance. ThereWork performed in partial fulfillment of the rcquirements for the Ph.D. degree, Michigan State University, 1963. Paper presentcd in part before the Division of Physical Chemistry at the 145th National 1Tecting of the American Chemical Society, New York, N. Y . , Sept,einber, 1963. ill. Gold and W. L. Jolly, Inorg. Chcm., 1 , 818 (1962). C . A. Kraus, J . Am. Chcm. Soc., 43, 749 (1921). G. W. Fowles, W. It. McGregor, and M. C. It. Symons, J. Chem. Soc., 3329 (1957). T. P. Dtls, “Advances in Chemical Physics,” Vol. IV, Interscience Publishers. Inc.. New York, N. Y., 1962. R. R. Dewald and J. 1,. Dye, J . Phya. Chem.. 6 8 , 121 (1964). G. E. Gibson and T. E. Phipps, J. Am. Chem. Soc., 48, 312 (1926). E. C. Evers, A. E. Young, and A. J. Panson. ibid., 79, 5118 (1957). D. S. Berns, E. C. Everu, and P. W. Frank, Jr., ibid., 82, 310 (1960). (10) E. C. Evers and P. W. Frank, Jr., J. Chem. Phys., 30, 61 (1959). (11) E. Beeker, It. H. Lindquist, and B. J. Alder, i b i d . , 2 5 , 971 (1956).
129
SOLURILITIES A N D CONDUCTANCES OF ALKALI ILIETALS IN ETHYLENEDIAMINE
fore, wc measured the conductances of alkali mctalcthyleriedianiirie solutions as a function of concentration. Windwcr and Sundheirn12 h a w reported the electrical cwnductivity of a solution of potassium in ~~thylrncdiamint~ as a function of optical absorbance. With grc>atclraccuracy, we have measured the conductances of sodium, potassium, cesium, and ruhidium solutions up to saturation. In addition, the solubilities of the alkali mctals in ethylenediamine were dctermineti at room tempcrattire.
Experimental Mntcrials. Purification of materials has been described in the previous paper.6 Apparatus. The assembly used in this work was specifically designed for the study of the conductance of these rather unstablc systems. It was possible to correct the raw data for decomposition so that the major difficulty encountered in thcse systems could be overcome. l’hr apparatus is shown in 1:ig. 1 A solution flask (500-ml.) with a side arm for metal distillation was attached to the solvent reservoir flask of the ethylcnediaminc purification train, the decomposition vcsscl of the gas analysis assembly, and the highvacuum line. To high votuum
,-
Waste vessel
-Solution
storage vessel (C)
Figure 1. Assembly for conductance and solubility memurement8: (a), metal solution make-up flmk; ( b ) , gas analysis assembly; ( c ) , conductivity cell apparatus.
12 oonductancc cell with small platinum ball electrodes was constructed of soft glass. The cell was scaled via a graded seal to a Pyrex storage vcssel and via a hall joint t,o a waste vessel (Fig. 1). The entire apparatus was so eonstructcd that solution could be tipped into the conductivity cell and then into the waste vessel. This allowed rinsing the cell with solution if desired. The cell constant \E” determined using a series of‘ standard KCl solutions. Conductivity was measured with a Siemens Wheatstone bridge a t 100,000 C.P.S. The conductances were det)ermined a t room temperature (‘23-25’) without temperature control since the main source of error was considered to result from decomposit]ion of the solutions. The gas analysis assembly consisted of a decomposition vessel, two liquid nitrogen traps, a leveling hulb with mercury, and a gas buret. Operation. A glass tube or capillary of purified metal was broken and introduced into the side arm of the solution make-up vessel. The complete assembly, including the gas analysis apparatus, was then evacuated and flamed until thc pressure stabilized a t less than torr. After sealing off the side arm under vacuum, the metal was vacuum distillcd into the bottom of the 500-ml. make-up flask. Sext, stopcocks 1, 2, and 3 (Fig. 1) were closed and the purified ethylenediamine was introduced into the make-up flask via t,he stopcock on the solvent reservoir flask. The blue solution which formcd immediately was then degassed by the high vacuum system through a liquid air trap and stirred constantly with a magnet sealed in glass for about 10 min. after all metal had apparently dissolved. In preparing saturated solutions of the metals, a large excess of metal was distilled into the makc-up flask. The solution was left in contact with the excess metal for ‘2 hr. with continuous pumping and stirring before transfer. Using this technique, consistent solubilities and conductance valucs were obtained. Ihrlier runs had shown that saturation is only slowly attained. After the blue solution had been degassed, previously purified nitrogen or argon gas, which was stored in 2-1. flasks on the vacuum line, was introduced via stopcock 1. The gas pressure forced the blue solution through the frit such that the solution filled the delivcry tube up to stopcock 2. The first fraction of the solution was then t,rnnsferred into the waste flask (A) of the conductivity apparatus via stopcocks 2 and 3. Then, the ncxt fraction of the solution was transferred into the decomposition vessel. I4nally, the remainder of the _______
(12)
~
S. W i n d w r arid B. 11. Sundhrim. ,I. Phip. Chem., 66, 1254 (1962).
Volume 68,Number I
Januaru, 1964
130
ROBERT
--
blue solution was transferred into the storage vessel of the conductivity cell apparatus. The time of transfer was recorded. The total time required to make the complete solution txatisfer was iiornially less than 15 see. Conductivily Readin.gs. Immediately after solution transfer, thc conductivity cell was disconnected and the electrodes of the cell were rinsed with blue solution, and then the cell was filled. The resistance was recorded as a function of time for about I hr. After this time, the solution was transferred into the waste vessel and a new sample was tipped into the cell for further measurements. l'he conductivity of the second sample was always lower than the first because of decomposition of the solution in the storage vessel of the conductivity cell apparatus. I t has been reported4,I2that platinum acts catalytically on the alkali metalethylenediamine solu.tions, accelerating the decomposition. This was also observed in this work, and the rate of decrease of the specific conductance with time was usually two to four times greater in the conductance cell than in the storage vessel. We also observed that during vigorous shaking, the conductance of the solution increased noticeably. This suggests that the solution in ttie vicinity of the electrodes differs in nature from the solution in the bulk because of decomposition. Before and during each conductivity reading, therefore, the cell was shaken vigorously until the highest steady conductivity reading was obtained. The log of t,he specific conductivity is a linear function of time, consistent with the decomposition behavior observed in the spectral cell which also showed essentially first-order decay. The specific conductance as a function of time could be extrapolated back to the time of transfer into the conductance cell. The several values at transfer could then be extrapolated to the time of initial transfer into the storage vessel which was also the time of transfer into the analysis vessel. Usually, this final correction amounted to less than 2% of the initial conductivity reading and rlormally the conductivity decreased between 5 and 20% during the first hour. Concentration Determinations. The volume of the solution transferred into the calibrated decomposition vessel was noted. Solid ammonium bromide was then tipped, via a side arm, into the blue solution. The evolved 13, gas was then pumped through two liquid nitrogen traps using a mercury leveling bulb, and collected in a gas buret in which the volume and pressure were measured. During the pumping, the liquid in the decomposition vessel was frozen slowly to ensure removal of all the hydrogen from the decomposed solution. This freezing and pumping process The Journal of Physe'eal Chemistry
R. DEWALI) AND JAMESL.
DYE
was repeated until the reading became constant. Usually after the first freezing, no more hydrogen could be pumped off by repeated freezing, melting, and pumping. Finally, the amount of metal could be determined from the stoichiometry of the reaction NH4Hr
+ R/I
--c
0.5H2
+ NH, + MBr
To ascertain that this was indeed the correct stoichiometry, a weighed sample of potassium was partially distilled into the analysis vessel and the amount remaining in the side arm was determined by titration. The amount of hydrogen collected agreed with that calculated to within 1%. This method of analysis has the advantage that metal present as decomposition products can be distinguished from the metal present in the zero oxidation state. A number of decompositions were made using water instead of NH413r and gave the same results. Mass spectrometric analysis of the gas obtained from the decomposition showed only hydrogen to be present. However, when the solution was allowed to decompose without adding KH4Br, requiring several days, mass spectrometry showed the presence of significant amounts of low molecular weight decomposition products of ethylenediamine in addition to the hydrogen. Lithium Conductivity. Since lithium cannot be melted in glass, a bent side arm was attached to the solution make-up flask by a standard taper joint so that it could be rotated into an upright position to tip the metal into the previously evacuated and flamed flask. The lithium metal solution was stirred and pumped on for 20 min. before transfer. Only the conductance of the saturated solution was studied because of the instability of the lithium solutions. Also, these solutioris were extremely uristable when in contact with platinum.
Results Solubility.
The solubilities of the alkali metals in ethylenediamine are given in Table I. It was noted that excess sodium, potassium, and rubidium had no noticeable effect on the stability of the solutions, while excem cesium metal appeared to catalyze the decomposition reaction. Conductance. The values of A, K , and C are given in Table 11. The equivalent conductance, A, is expressed in kohlrausch units, the specific conductance, K , in ohms-' em.-', and the concentration, C, in moles per liter. The values of A for sodium, potassium, rubidium, and cesium are plotted us. C'" in Fig. 2. 1,og-log plots of the specific conductance us. concentration (illustrated by Fig. 3) for these same metals give straight lines, and potassium, rubidium, and cesium have the same
SOLUBILITIES AND CONDUCTANCES OF
ALKALIJ f E T A W
Table I1 : Conductanc:e-Concentration Data for Solutions of Alkali Metals in Ethylenediamine
Table I : Solubilities of the Alkali Metals in Ethylenediamine, a t Room Temperature
Metal and run
Average
Run no.
Metal
Lithium
1 2 3 la lb 2 la lb 2 3
Sodium Potassium
Rubidium
1
Cesium
2 I
Solubility, mole/l.
0 0 0 2
283 264 314 37 x 2 35 x 2 45 x 1 02 x 1 03 x 1 05 X I 07 X 1 29 X 1 33 x 5 43 x
solubility," rnole/l.
uo.
0 287 f 0 015 10-3 10-3 10-8
o
2 39
10-2 10-2
04
x
10-3
1 04 f 0 02 X
1 31 f 0 02 X
IO-* 10-2 10-2
5 4
131
IN ETNYLEICEDIAMINE
x
10-2
The limits of precision given are average deviations from the mean.
slope but are displaced from one another, while sodium gives a straight, line with a different slope. It should be emphasized that each value given in Table I1 and shown in Fig. 2 represents an independent solution make-up.
Specific conductance
Concentration, moles/l
cs-1 cs-2 CS-5 CS-6 cs-3 cs-7 cs-4 Na-2 Na-lb Na-4 Na-3 K-3 K-2 K-lb K-10 K-6 K-8 K-4 K-9 K-5 K-7 K-11 K-12 Rb-1 Rb-7 Rb-2 Rb-6 Rb-3 Rb-5 Rb-4 Li-3 Li-4
54 48 20 7 3
x
X 10'
3 9
10s
1420 1350 742 342 193 113 78.6 62.5 60.0 27.0 6.36 30 1 297 299 201 139 135 103 99.0 92.6 90.8 29.7 12.1 321 199 102 98.6 70.0 35.0 30.6 15300 14000
8 65
55
1 68 0 960 2 45 2 35 1 04 0 242 10 7 10 5 10 3 6 02 3 68 3 66 2 43 2 37 2 10 2 04 0 416 0 135 13 3 6 94 2 85 2 78 1 78 0 696 0 526 314 287"
Equivalent conductance
26 27 35 44 54 67 81 25 25 25 26 28 28 29 33 37 36 42 42 44 44 71 89 24 28 35 35 38 50 58 48
2 7 6 8 4 3 8 5 5 9 3 1 2 0 4 8 9 5 7 1 5 3 7 2 7 8 5
5 3 3
49
The average solubility value. Littiiuni results are only semiquantitative.
The specific conductivity for the concentration range investigated may be expressed by the function log K
= U
log
c+b
Spectral data showed that solutions tend to decay by a first-order rate law, so that lOgC
=
-kt
+d
If the decomposition in the conductance cell is also first order, we have log
20
1 0
1 4
1
1
1
1
8 C'/l
x
12 102.
1
1 16
1
I 20
Figure 2. Equivalent conductance us. C ' / 2for sodium, potassium, rubidium, and cesium in ethylenediamine.
I 24
K
= -(k/U)t
+ (d - b )
so that the decay of conductance is consistent with first-order decomposition. Lithium Conductivity. Semilog plots of the specific conductance us. time for lithium solutions display Volume 68,n'umher 1 January, 1.964
132
ROBERT R. DEWALD A N D $JAMES 1,. DYE
0
A
I No loo
A
ductivity is considered to be due to metal ioris and ethylcncdiamide ions which result from thc decomposition reaction. Conductivity data for tho decomposed solutions are given in Table 111. These data arc of interest because they emphasize t he low equivalent conductariccs of the decomposition products, consistent with the low conductance found for sodium cthanolarnidel3 in ethylenediamine. I t is evident from t>hese data that a small accumulation of decomposition product in a metal -cthylencdiaminc solution uould lead to no significant error i n the conductance attributed to the metal in the zero oxidation state.
Cs in Ethylenediamine Rb in Elhyiermdiamine in Ethyienedlamine
K in Ethylenediamine Nain Ammcnia (31
0 Cd in Methylamlns
(71
X K in Methylamine (71 0 L i in Methylamine
(91
10
s X
F 1
Table 111: Conductance-Concentrution I h t a for Decbomposed Solutions of Sodium, Potassium, Rubidium, and Cesium in Ethylenediamine Concentration. IIIO~CR/~.
, 0.1
I
I I I I
I
,
1
1
1
10
1
cx
x
108
Cesium
34 48 20 1 2 2 1 0 10
3 9 8 68 45 35 04 242
10 2 2 2 0 0 13 12 6 2 0
3 82 42 10 412 138 3 9 92 78 497
108.
Figure 3. Comparison of the conductanre curves for metalamine systems (log-log plot of K V us. ~ C ) ; K = specific conductance, q0 = viscosity of the solvent.
three distinct decay rates before the solution is totally bleached. Initially, the decay rate was similar to, but faster than, the rate observed for the other metals and was, of course, accelerated by the catalytic effect of the platinum. After 10-20 min., the specific conductance dropped by about two orders of magnitude within 1 or 2 min. Finally, the lower specific conductance decreased more slowly until it reached the value of the bleached solution. When lithium solutions decompose, precipitation occurs and the lithium spcctral data show that precipitate formation results in rapid disappearance of the infrared absorption. The behavior observed during lithium decomposition in the conductance cell is explained as follows. The solutions decompose normally until precipitation occurs. When the precipitate forms ncar the platinum electrodes, it then speeds the decomposition and rcsults in further precipitation. The infrared species then is rapidly destroyed by the precipitate. This rapid removal of the infrared species caiises a sudden drop in the conductivity, followed by a slower decrease as the visible species decomposes. This is in keeping with the idea that the infrared species is the main contributor to the conductivity of the solution. Conductivity of Decomposed Solutions. The resistances of the solutions approached constant values as the blue color disappeared. The remaining conThe Journal of Phuaical Chemistry
Metal
1
Sodium
Potassium
Rubidium
7 10 R
Specific conductivity
x
106
!)ti, 0
87.4 72.5 11.4 G.40
5.88 3.26 1.06 31.7 27.3 30.3 12.7 12.3 10.1 3.36 2.12 38.7 37.4 22.8 12.7 4.89
Equivalent conductivity
1.76 1.78 3.31 6.80 2.61 2.50 3.13 4.38 2.96 2.60 2.94 4.67 5.06 4.82 8.16 15.4 2.91 2.90 3.29 4.57 9.85
Discussion Comparison o j Conductance Curoes. The general form of the conductance curves for potassium, rubidium, and cesium in ethylenediamine is similar to that for alkali metals in mcthvlaminc but difTers somewhat from that of sodium in liquid ammonia. A method of comparing the different, svstems, which tends to correct for differences in viscosity, is to plot the log of the product of the specific conductance and viscosity 1)s. the log of the concentration as shown in Fig. 3. This ~
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W. B. Schaap, R. E. Bayer. J. R. Sicfker. J. Y . Kim, I-’. W. Brewster, and F. C. Schmidt, Record Chem. Progr., 22, 197 (1961).
133
SOLURILlTIES A N D CONDUCTANCES OF ALKALI JIETALS IN ETHYLENEDIABII~E
figure shows that cesium-amine systems tend to be better conductors than systcnis of the other alkali metals in amincs. I t should be rioted that the data of Herns9 show a deviation from tlie usual straight-line log -log plot a t about 0.02 M and then continut at higher concentrations with a straight line having the same slopc as t)efore. I t is tempting to attribute this ariornalous brhavior to a systematic error introduccd in measurements in the high concentration region. H o ~ e v e r thc , rcsults of a number of runs are consistent. The conductance brhavior of sodium in cthylencdiamiric (h'ig. 2 arid 3 ) differs markedly from that of the other alkali metals in amines. Thc equivalent conductance of sodium solutions is much lower than for the other alkali metals in cthylenedianiine and is almost independent of concentration over the tenfold concentration range investigated. This indicates that the number of current carriers is proportional to the concentration and that the mobility is nearly independent of concentratioii. Such behavior would not be expected for conventional dectrolytcs in this medium since extensive ion-pairing would cause the equivalent conductance to be strongly concentration-dependent. The apparent limiting values of the equivalent coriductanccs, A,,, were evaluated by using the method of Shedlovsky. l 4 ,Ilso, the apparent pairing constants, K1, for solutions of potassium, rubidium, and cesium in ethylenediamine wrre evaluated. Constants used for the Shedlovsky aiialysis of thc data were ?lo = 1.54 X l o b 2 poise arid D = 12.9, where 90 and D are, respcctivrly, thc viscosity and dielectric constant of tlie solvcnt To compute activity coefficients, the extended Dcbyc- Hiickel expression logs*
_=
- 7-~. 6 4 d E 1 1- 4.8GdC, --
--2
was used. Thiso corresponds to a closest approach distance of 6.0 -1. The results for potassium, rubidium, and cesium are plotted in Fig. 4. The intercept and slope, from which apparent values of A. and K , were ot)tained, werr dctcrmined by the method of least squarcs for each set of points in Fig. 4. In order to have enough points in the dilute region for cesium, it was ncccssary to estimate K by extrapolation of the log--log plot of specific conductance 2's. coriccritration (Fig. 3 ) . This is felt to be valid in vitw of the data for potassium arid rubidium a t low concentrations. Evers and Frank'" have cvaluatcd the limiting equivalcnt conductance arid the pairing constant for the sodium-ammonia syst em and 13crns9 has done so for the lithium -methylamine system. These values, as well as the values obtained i n this work, are given in
0 Cs--Slope = 0.167
------ 0.0490 2
8
6
4
CA/%S(Z) X 10'.
Figure 4. Shedlovsky analysis plots for cesium, potassium, arid rubidium in ethylenediamine.
Table I\'. The value of A. for lithium in methylamine reported by Berm may be too high. When the five highest concentration points are eliminated (which also fall in the anomalous region of the plot of log K vs. log C using t,he data of Rems), one obtains a value for A. of about 180 rather t,han 228. Table I\' shows that the Walden product, novo,
Table I V : Comparison of Limiting Equivalent Conductances Solvent
Ethylenediamine" hlethylarriine' Ammoniad a
Metal
AQ
Cesium Potassium Ituhidium Sodium Lithium Sodiuni
204 139 117 27' 228.3 1022
Ki
Aorlo
3.14 2.14 1.80
1 . 4 4 X 10-4 1 . 5 4 X 10-4 1 . 6 9 X 10-4 ...
...
2.07 2.61
5.5 7 23
x
x
10-6 10-8
Referred to in this work. * Estimated from a plot of A vs. See ref. 9. See ref. 10.
C''l.
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T.Shedlovsky, J . Franklin IT&.,
225, 739 (1938)
Volume 68,.Vumber I
January, I064
134
for cesium in ethylenediamine is somewhat larger than for the sodium-ammonia system while the Walden products for the other metals in amines are considerably lower than this. Since the spectra of dilute cesium-ethylenediamine solutions exhibit only an infrared absorption band and the conductance behavior appears to be similar to that of the sodium-ammonia system, we conclude that the conductance processes for these two systems are essentially the same. The spectra of dilute solutions of potassium and rubidium in ethylenediamine show absorption bands other than the infrared absorption band. If the contribution to the total conductivity of these species is small and essentially a constant fraction over the concentration range investigated, one concludes that the major contributor to the conductivity of these solutions is the infrared absorbing species. This requires that the general form of the conductance curves be the same for rubidium, potassium, and cesium in ethylenediamine. This is shown to be the case in Fig. 2 and 3. The ionpair constants given in Table IV are essentially the same for these three metals, and this further indicates the same conduction process. The dissociation con-
The Journal of Physical Chemistry
ROBERTR. DEWALD AND JAMESL. DYE
stant for sodium in liquid ammonia is somewhat larger than for these metals in ethylenediamine. This would be expected, however, on the basis of the higher dielectric constant of ammonia. The large differences in the apparent limiting equivalent conductance values are difficult to understand and seem to indicate that the other species besides the infrared absorbing species which are present in the metal-ethylenediamine systems either do not dissociate into solvated electrons and solvated metal ions at the dilutions studied, or else that the cation has a strong influence upon the mobility of the solvated electron. Acknowledgments. The experimental work described in this paper was done at the Max Planck Institut fur physikalische Chemie, Gottingen, West Germany. The authors wish to thank Dr. Manfred Eigen for his kind invitation to work in this laboratory as well as the many personnel of the Institut who made their talents available to us. Part of this work was done under a Science Faculty Fellowship from The National Science Foundation (J. L. D.), and it was also supported in part, by the U. S. Atomic Energy Commission (Contract AT(11-1)-958).
