Spectra of Alkali Metal-Amine Solutions
323
On the Absor tion Spectra of Alkali Metal- Amine Solutions Y. Bar-Eli* and G. Gabor Department of Chemfstiy, re/-Aviv Univers\ty, Tel-Aviv, /Sfae/
(Received March 3, 7972)
'Ve have investigated very carefully the temperature dependence, solvent dependence, and decay rates of the spectra of alkali metals in solutions in methylamine and ethylamine. Results confirm earlier assumptions that the M band is associated with the negative metal ion. The decay rates of these solutions confirm the equilibrium between the metal anions and the solvated electrons. This equilibrium is shifted toward the metal negative ions on cooling. The existence of a new, hitherto unreported, absorption band at 440 nm in solutions of lithium is reported.
Introduction The absorption spectra of solutions of alkali metals in various amines cmd in ammonia have been studied extensively.1 The existence of two bands has been established: an M band which peaks between 660 and 1000 nm depending on the metal, and an S band which is metal independent and peaks at '1400 11111.3 The temperature dependence of the M band was investigated, and used as a clue in the identification of the chemical species associated with it. Also various other properties were reexamined, such as the solvent dependence, decay rates, etc., in order to shed more light on the behavior of these solutions. Experimental Section Solutions were prepared from alkali metals of highest available pur Ity Koch Light 99.9'70, which were distilled into the reaction vessel as described earlier.lf,2b The amines from the Matheson Co., Inc., were distilled from blue potassium solution onto the purified alkali metals. After being in contact with the metal for a few minutes, the solution was transferred through sintered quartz (in order to avoid the introduction of metal particles) into the spectroscopic c d l . As was indicated earlier,lg,2 the whole vessel, apart from the Fischer and Porter Co. Pyrex-Teflon needle valve, was made of quartz. Since the surface area of the vessel probably has great influence on the solutions, all vessels went through a careful and uniform cleaning procedure as follows: (1)washing in boiling solution of Alconox detergent, (2) washing the Alconox with distilled water, (3) boiling in distilled water, (4) washing twice in triple distilled water, ( 5 ) drying in an oven overnight g t 500", and (6) warming to 300" under vacuum for 24 hr In spite of the un,ihrmity of the cleaning procedure, there ivere variations 111 the stability of the solutions, although the general patterns of behavior were always reproducible. The cell was located in a dewar flask equipped with optical windows. The dewar was cooled by dry nitrogen as described earlier. The measurements were done in a Cary 14 spectrophotometer. The optical path in all experiments was 1 cm. The position of the maximum was found as follows: a few horizontal lines were drawn near the band maximum. The intersection of the midpoint line and the band was taken as h m a c . l C The spectra for these measurements were taken on an expanded wavelength scale in order to obtain greater accuracy. All the measurements of
temperature dependence were done by increasing the temperature slowly from -80", in order to avoid complications arising from the faster decay rates at high temperatures.,
Results and Discussion In Figure 1 typical spectra of potassium are shown in both solvents. It is seen that changing the solvent from ethylamine to methylamine causes a small red shift in the M band, and the S band is greatly increased. In the case of Li-methylamine and Cs-methylamine solutions, the S band was so large that the M band was not observed at the beginning. In lithium solutions it appeared later, as the decay rate of the S band is faster than that of the M band (see following paragraphs). On lowering the temperature, the M band is blue shifted. The extent of the blue shift depends on the alkali metal (Figure 2) and on the solvent (Figure 3). The results are summarized in Table I. Following Stein and Treining we plot the values of the slopes of Figure 3, namely, d;/dT us. the peak absorption at 20". A straight line for this plot is a test for a CTTS absorption. Indeed a straight line is obtained (Figure 4) in agreement with the results of others"7 that the species associated with the M band is the metal negative ion. The CTTS mechanism for the absorption involves the solvation shell of the negative ion. The thermal vibrations of this shell decrease on lowering the temperature, and will cause narrowing of the band, together with absorbance increase, in order to preserve the oscillator strength8 (1) (a) H. Blades and W. Hodgins, Can. J. Chem., 33, 411 (1955); (b) L. R. Dalton, J. D. Rynbrandt, E. M. Hansen, and J. L. Dye, J. Chem. Phys., 44, 3969 (1966); (c) M. Gold and W . L. Jolly, lnorg. Chem., 1, 818 (1962); (d) R. C. Douthit and J. L. Dye, J. Amer. Chem. Soc., 82, 4472 (1960); (e) D. F. Burrow and J. J. Lagowski, Advan. Chem. Ser., No. 50, 125 (.1965); ( f ) M Ottolenghi, K. BarEli, A. Linschitz, and T. R . Tuttle, Jr., d. Chem. Phys., 40, 3729 (1964); (9) G. Gabor and K. Bar-Eli, J. Phys. Chem., 75, 286 (1971). . (2) (a) I . Hurley, T. R. Tuttle, Jr., and S. Golden, J. Chem. Phys., 48, 2818 (1968); (b) Pure Appl. Chem., Suppi., 449, 503 (1969) (3) One has to be careful when working in quartz vessels to eliminate as far as possible the leaching of sodium from the walls.* (4) M. Ottolenghi, K. Bar-Eli. and H. Linschitz, J. Chem. Phys., 43, 206 (1965). (5) G. Stein and A . Treinin, Trans. Faraday Soc., 55, 1086, 1091 (1959). (6) (a) S. Matalon, S. Golden, and M. Ottolenghi, J. Phys. Chem., 73, 3098 (1969); (b) D. Huppert and K . H. Bar-Eli, ibid., 74, 3285 (1970). (7) M. G. DeBacker and J. L. Dye, J . Phys. Chem., 75,3092 (1971). (8) (a) G. E. Gibson and N. S. Bayliss, Phys. Rev., 44, 188 (1933); (b) G. E. Gibson, 0. K. Rice, and N. S. Bayliss, ibid., 44, 193 (1933): (c) C. N. R. Rao, "Ultraviolet and Visible Spectroscopy," Butterworths, London, 1961, p 142.
The Journal of Physical Chemistry, Vol. 77,
No. 3, 7973
K Bar-Eli and G . Gabor
32 TABLE I: Spectral Data on Solutions of Alkali Metals in Methylamine-Ethylamine Mixtures Mole fraction of MeNH2
Solute Band Absorbance Ratio M I S Width at half-
intensitye -dG/dT,
cm-' deg-'
0.98
1
K M
K
Rb S
0.54
0.365 Rb
K
M S M S ' M S 1.9 1.9 1.11 0.7 $20" 1.9 1.6 1.85 0.58 1.8 0.43 -..80° 2.9 1.1 1.7 0.38 3.05 1.2 2.56 0.51 2.7 0.45 20" 1.o 1.58 1.19 3.2 4.2 4.48 --ma 2.64 2.5 5.0 6.0 4-20' 4100 3450 -- 80" 2700 3000 10.8 11.8 9.8 10.8 k0.3 f0.2 f0.5 k1.2 M
S
+
a.0
0.368a NaC K I V
Rb Cs M M 1.13 0.32 3.67 0.56 1.34d 0.32 0.53 1.02d 1.63 0.19 3.25 0.1 0.74 0.46 0.81 0.98 3.52 6.55 8.6 3.2!ib 3450 3600 4100 3600 2750 2550 3200, 2630 8.9 8.2 9.4 13.9 * ? . I 1 0 . 6 f0.5 A l " 1 K
M
S
K S
M
M
I
*
a Solution saturated with KCI. At low temperatures split spectrum is obtained (ref lg). Measured in lithium Solutions. Absorbance results are doubtful because of fas! decay rate. e Error is approximately + l o 0 cm-'
1 2 . 4 ~
I O r
,e--.
/' /
m c b
11.6
k -60 -40
10'8-100-80
-20
0
20
u 40 60 80
roc
A
nm
Figure 1. Spectra of potassium in (- - - --)
)-(
ethylamine and
methylamine.
This is indeed borne out by experiment, as can be seen from Table I. In solutions where the S band is significant, the ratio M/S increases with the ethylamine mole fraction and with temperature decrease, as seen from Table I. One expects the S band to be stabilized by methylamine, since in pure ethylamine this band is hardly seen. The temperature influence is twofold: (a) the increase, described in Table I, of the M band, due to the freezing of the solvation shell, and (b) the decrease of the S band, due to the shift of equilibrium in (1) to the left on lowering the temperature. This deduction i s given further support by adding potassium chloride; the ratio M/S is further increased as is expected from the above equilibrium (see Table I).
16r
14
Figure 3. M band maximum vs. temperature in solution of potassium in 0 ,methylamine; 0 , methylamine-ethylamine mixture = l : l ; and A , ethylamine.
I-
M+M++2e-
t
oc
Figure 2. M band vs. temperature in solutions of alkali metals in ethylamine: 0, Na (see footnote c, Table I ) ; 0 , K; A , Rb; A ,
cs.
The Journal of Physical Chemistry, Vol. 77, No. 3, 7973
(1)
The decay of both M and S bands, follows first-order law;9 the M band decays twice as fast as the S band. This result conforms with the fact that equilibrium 1 is f a ~ t . ~Taking + ~ b logarithms and the derivative with respect to time of equilibrium 1,one obtains (9)J L. Dye, Accounts Chem. Res., 1,306 (1968)
Spectra of Alkali Metal--AmineSolutions
325
,-u
r
q ( + 2 0 ° )cm-lx 10-3 '03
10.6
105
'104
2o c
10.7
10.8
T
/
C
i 121 x
u
6
0 25
11.0
11.2
1'1
113
11.1
11.5
%
V (+20')cmlx 10-3 Figure 4. "Blue chift" of M band vs. M band maximum at room temperature. I:
I C
I
I
OE >.
t L1
z L!J OF i
a
u !-
0
01
0 ;
OC
IO
W A V E L E N G T H , A. n m
Figure 5 . Visitile part of spectrum of solution of lithium in ethylamine. Notice the fast decay of 690-nm band (5 min between each plot). This spec;tnm is recorded after the complete decay of the S band.
. R = (d In [M-]/dt)/(d In [e]/dt) = 2 + ( d l n [M ]/dt)/(dln [e]/dt) (2) Since le1 and l p V P - - l decrease with time, and it is reasonable to assume &at'[M+] does not, then the ratio of the
0.30
I 0.35 I R ~ ( A- - ' )
0.40
0.45
Figure 6. Difference between M band maximum (room temperature) and electron affinity vs. reciprocal of anion radii of the various alkali metals (the 440-nm band is assumed to be associated with lithium and the 690-nm band with sodium). rates is 0 5 ' R 5 2. The upper value is obtained in the case of potassium and rubidium in methylamine; the lower ratios are obtained for lithium in ethylamine. In the latter case one obtains an intense S band which decays very fast, leaving a band at 690 nm which decays more slowly; also, in a freshly prepared solution, we obtained a hitherto unreported band at 440 nm which decays slower than the 690-nm band (Figure 5 ) . After its decay this 440-nm band does not reappear even when the solution contacts the metal again, while the S and the 690-nm bands do reappear. Other reports on lithium solutions in ethylamine show ,~ that in only an S band.l.2 However, Hurley, et ~ k l . report very concentrated lithium solutions, a 690-nm band exists and starts to decay only after there is no more S band; they do not report the existence of a 440-nm band. Similar results were obtained in ethylenediamine by Dewald and Dye.10 It is possible to associate the 690-nm band with the Liion, i.e., an M band, and the 440-nm band with an unidentified unstable organic species. Another possible explanation is to associate the 690-nm band with Nawhich is leached by lithium from the vessel walls (quartz contains ~ 0 . 1 %of sodium), and the 440-nm band with Ei-, so that both bands are M bands of the respective metals. Support for the latter conjecture can be obtained by the plot of the difference between the absorption energy and the electron affinity11 us. the reciprocal of the anion radii of the various alkali metals according to Stein and Treinin.5 The assumption that the 440- and 690-nm bands are associated with lithium and sodium, respectively, will put all the alkali metals on one straight line (Figure 6). In spite of the feasibility of each of the above explanations, it is impossible to definitely associate the bands with any particular species, until a stable solution of lithium is prepared and a method to dissolve sodium is found. (IO) N. R. Dewald and J. L. Dye, J. Phys. Chem.. 6 8 , 121 (1964).
(11) (a) A . W. Weiss, Phys. Rev., 166, 7 0 11968); (b) B. Yaakobi, Phys L e t t , 23,655 (1966).
The Journal of Physical Chemistry, Vol. 77,
No. 3, 7973
