J. T. NELSON,R. E. CUTHRELL, AND J. J. LAGOWSKI
1492
Liquid Ammonia Solutions. 111. The Nature of Solutions of the Alkali and Alkaline Earth Iodides
by J. T. Nelson, R. E. Cuthrell, and J. J. Lagowski Department of Chemistry, University of Terns, Austin, Texas (Received October 25, 1965)
The spectra of liquid ammonia solutions of the alkali and alkaline earth iodides, except Ca12, exhibit an absorption band at about 2500 A, the position and intensity of which is temperature dependent; the position of the band is also dependent upon the concentration of added inert salt. Analysis of the data indicates that the band arises from a chargetransfer-to-solvent transition. The absence of this band in solutions of CaIz suggests that a molecular complex is present in this system.
Introduction The position of the first ultraviolet absorption band for the iodide ion is solvent and temperature dependent; this band has been described as arising by a charge-transfer-to-solvent (ctts) mechanism.' We present here the results of a detailed study of the ultraviolet absorption band reported2 for iodide ions in liquid ammonia which support this model. Experimental Section The spectra of liquid ammonia solutions containing iodide ions were determined using the equipment3 and rinsing technique* previously described. The temperature of the solution in the optical cell was determined with a calibrated thermistor probe sealed into a glass well. Commercially available reagent grade lithium, sodium, potassium, rubidium, cesium, calcium, strontium, and barium iodides were dried at 200" in vacuo and stored in a glove box containing helium which was equilibrated with sodium-potassium alloy. Lithium and potassium perchlorates were dried a t 200 and 300",respectively, and stored in a glove box.
LiI, NaI, KI, CsI, Sr12, and BaIz obey Beer's law in m at -78" the concentration range 1-10 X (Table 11), indicating that the iodide ion and aggregates containing this species cannot be distinguished spectrophotometrically under these condition^.^ However, the addition of a 100-fold excess of inert salt (Kc104 or LiC104) lowers the wavelength of the iodide absorption band by 6-8 A at a given temperature (Table I). The temperature dependence of the band maximum is characteristic of ctts spectra, the transition corresponding to excitation of an electron to the first layer of solvent molecules. Platzman and Franck6 proposed a model for the excited state of the solvated iodide ion in which the electron moves in an orbital centered on iodine and defined largely by the field of the polarized solvent molecules around the ion. Smith and Symonsla-c suggested that the excited electron is trapped in a solvent cavity. I n this interpretation the excited (1) (a) M. Smith and M. C. R. Symons, Discussions Faraday Soc., 24, 206 (1957); (b) Trans. Faraday SOC.,54, 338 (1958); (c) ibid., 54, 346 (1958); (d) G.Stein and A. Treinen, ibid., 55, 1086 (1959). (2) W. Jolly, University of California Radiation Laboratory Report 1952,p 2008. (3) E. C. Fohn, R. E. Cuthrell, and J. J. Lagowski, Inorg. Chem., 4, 1002 (1965). (4) (a) D. F. Burow and J. J. Lagowski, Advances in Chemistry Series, No. 50, American Chemical Society, Washinpton, D. C., 1965,p 125; (b) R. E. Cuthrell, Ph.D. Dissertation, The University of Texas, 1964. (5) R.E.Cuthrell, E. C. Fohn, and J. J. Lagowski, Inorg. Chem., 5, 111 (1966). (6) R. Platzman and J. Franck, Z. Physik, 138,411 (1954). ~
Results and Discussion Liquid ammonia solutions of the alkali metal iodides, as well as of Sr12 and Ba12, exhibit an absorption band a t about 2500 A, the position of which is temperature dependent (Table I). The general characteristics of the band, however, do not vary with the nature of the cation. Within experimental error, solutions of T h e J O U T ?of ~Physical Chemistry
LIQUIDAMMONIA SOLUTIONS OF ALKALIAND ALRALINEEARTH IODIDES
~~~~~~
~
~
Table I : Values of Xmax a t - 50" and the Temperature Dependence of the Iodide Ion Band in Solutions of Alkali and Alkaline Earth Iodides
Aa
dX,,. , dt A deg-1
dEm*x, dt oal deg-1
2500 2493 2500 2500 2499 2491 2493 2499 2491 2485 2501
1.28 f 0.05 1.42 f 0.07 1.36 4 0.04 1.34 4 0.06 1.24 4 0.05 1.31 4 0.03 1.33 4 0 . 1 0 1.39 4 0.02 1.32 4 0.03 0.74 f 0.03 1.06 zk 0.08
-59.0 -72 6 -62.8 -61.5 -56.0 -66.8 -67.8 -64.0 -67.4 -32.0 -30.3
Amsx,
Compd
LiI LiI NaI KI RbI RbI RbI CSI CsI SrIz BaIl
+ LiC1O:
a
+ KC102 + LiC104b + LiCIOP
Using a 1Wfold excess of inert salt.
Measured a t -50".
Table 11: Molal Extinction Coefficient" a t 2475 A for Iodide Ion in Liquid Ammonia a t - 78" Compd
6
x
1.86 i 0.06 1.81 zk 0.04 1.70 f 0.06 1.73 ==! 0.07
KI CSI 8rI2 BaIz
a Determined from the slope of the Beer's law plots. lated per equivalent of iodide ion.
* Calcu-
electron is constrained to move within a solvent cavity containing an iodine atom, the energy of the transition being approximated by a simple electron-in-a-squarewell model from the energy cycle shown in Scheme I . l a ~ b I n this cycle 1-("3), represents the iodide Scheme I E1
I-("&
J.
(I
--f
("3)Z
+
kp
Emax
+ e-)("&
("3)z
y. \ E,
I("3)z
+ I(g) + e-@ /
k7
+ e-(d
ion in a cavity composed of x ammonia molecules, (NH3), the cavity consisting of similarly oriented solvent molecules, I(",), an iodine atom in the cavity, e-)(NH3), an iodine atom plus an electron and (I occupying the same cavity. The energy required to
+
remove the ion from the cavity leaving the latter intact is represented by El, the ionization potential of the iodide ion by I P , the energy gained by replacing the iodine atom in the cavity by E3, the energy required to put the electron into the lowest orbital defined by the first solvation shell by E4, and the energy of the transition by E,,,. It follows that Emax =
IP
+ E1 + E3 + Ed
(1)
The energy of an electron in a well of radius ro is given' by
E = hz/8mro2 (2) where the other symbols have their usual meanings. By previous arguments1c E4 = -E1
+ h2/8mro2
(3)
since the relaxation time of the solvent cavity is very much less than the time required for the transition. Assuming El to be negligibly smallla-c and substituting eq 3 into eq 1 gives
E,,,
10-b
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=
IP
+ h2/8mro2
(4)
The absorption band for ammonia-iodide systems occurs at 2475 A at -78". Using 72.3 kcal/mole for the ionization potential of the iodide ion, the radius of the solvent cavity is calculated to be 4.48 A, which extrapolates to 4.80 A at 25". The latter value should be compared with the value of 4.05 A reported for the water-iodide system at 25". IC The observed decrease in E,,, with temperature (Table I) suggests that, on the basis of the model proposed by Smith and Symons, the radius of the solvent cavity decreases with temperature. I t might be expected that the radius of the cavity would become essentially that of the iodide ion at 0°K and that the value of E,,, should approach that for the alkali metal iodides at 0°K. The extrapolated value of E,,, at 0°K is 130 kcal/mole, which is in good agreement with the values of 134.5 1 kcal/mole calculated for a variety of and 135.3 kcal/mole observed for crystalline K I at 20°K.* The extrapolated value of E,,, corresponds (eq 4) to a cavity radius of 3.89 A. Stein and Treinens point out that the value of the cavity radius for the iodide ion in water calculated from the energy cycle shown in Scheme I is about 1 A greater than any value expected. A modified Platz-
*
(7) K. Pitzer, "Quanhm Chemistry," Prentice-Hall, Inc., Englewood Cliffs, N. J., 1960,p 512. (8) H.Fesefeldt, Z.Phyaik, 64, 623 (1930). (9) G.Stein and A. Treinen, Trans. Faradav SOC., 55, 1086 (1959).
Volume '70,Number 6 M a y 1966
1494
J. T. NELSON, R. E. CUTHRELL, AND J. J. LAGOWSKI
man and Francks model incorporating the conceptla-c of a variable-radius parameter was suggested based on the energy cycle shown in Scheme 11; all energy terms Scheme I1
5
I-(sol)z JEW
sol),
x-(g)
+
Table 111: Iodide Ion Cavity Radius (SOlL
kp
-E x(g> + e-(g> + (sol>z
Emax
=
E1
+ + IP E2
are relative to a nonpolarizing electron in the medium at infinite distance from the ions. The energy, El, is the sum of the potential energy of the ion due to the persistent polarization of the organized medium and the electronic polarization caused by the ion in solution; Ezis the energy involved in introducing the radical into the organized medium (taken to be equal, but of opposite sign, to the heat of solvation of the radical) and introducing the electron to its excited level.6 Equation 5 is the basic relationship of the Stein and Treinen theory
where L, is the heat of solvation of the iodine radical, Do,and D, are the optical and static dielectric constants, respectively, Re is the mean distance (5.8 A) of the electron from the center of the cavity for the 2s state (chosen by Platzman and Franck for the halides), and the remaining terms have their usual significance. Assuming that the difference in the heat of solvation by ammonia and by water for the iodide ion (Le., ea. -4 kcal/molelO) is the same as for an iodine atom, and using -4.5 kcal/mole for the heat of hydration of an iodine atom," L, can be estimated as -8.5 kcal/ mole. Equation 5 can be used to calculate the cavity radius for the iodide ion in liquid ammonia using the data for the static4&and optical dielectric constants12 that appear in the literature. The results of this calculation for several temperatures appear in Table 111. The cavity radius calculated from eq 5 is markedly smaller than that calculated on the basis of Smith and Symons' model (eq 4). The interpretation of ctts spectra of anions in aqueous solutions according to eq 5 has yielded anion radii that correspond to the apparent ionic radii obtained from partial ionic volumes.9 Of the cavity radius from eq 5 to absolute zero gives a value of 2.43 A; the Crystallographic radius of the iodide ion is 2.16 A.13 There The J o u r 4 of PhysieaZ Chemistry
are no independent data available on apparent ionic volumes for species dissolved in liquid ammonia.
Temp,
Emas*
O C
A
25
-78 -273
2599 2475 2204
-Cavity Eq 4
4.80 4.48 3.89
radius, AEq 5
3.04 2.84 2.43
The absorbance a t the band maximum decreases with increasing temperature, as might be expected on the basis of the decrease in the density of the solution with temperature. However, the extinction coefficient a t the band maximum decreases with increasing temperature even after a correction is made for the change in density. Thus, for example, tmsxfor K I varies linearly from a value of 1.86 X lo4 at -72.2' to 1.58 X lo4 at -38.0'. A variation of this nature has been interpretedlc in terms of an increase in the breathing vibrations of the cavity with temperature, which affects the relative shape of the configuration-coordinate curve for the ground and excited states. The presence of excess cations causes a slight decrease in the value of E,,, for the iodide ion (Table I), a phenomenon which has been observed in other solvents.1b On the basis of the model suggested, it would appear that under these conditions cations can influence the potential field about, and the effective radius of, the cavity. An obvious suggestion is that cations become incorporated in the solvent layer a t higher over-all cation concentrations. However, the temperature dependence of the iodide band in solutions containing excess cations is the same as that for solutions containing only the alkali and alkaline earth iodides. The absorption bands attributed to an iodide-containing, contact ion pair occur at about 2900 A ; the band position is virtually temperature independent. The band associated with iodide-containing, solvent-shared ion pairs is significantly less sensitive to temperature than is the position of the band attributed to solvent-separated ion pairs.14 Thus, it is likely that the slight shift in the position of the iodide (io) w. J ~ U Y ,them. R ~ v .so, , 351 (1952). (11) H.L. Friedman, J . Chem. P h w , 21, 319 (1953). (12) D. F. Burow, Ph.D. Dissertation, The University of Texas, 1966. pauling, '$TheNature of the Chemical Bond,?,3rd ed, Cor(13) ne11 University Press, Ithaca, N. y., 1960, P 514. (14) G.stein and A. Treinen, T T ~Faradw . SOC.. 56, 1393 (1960).
LIQUID AMMONIA SOLUTIONS OF ALKALI AND ALKALINE EARTHIODIDES
band in solutions containing excess cations arises from slight, nonspecific perturbations by the cation on the solvent cage of the iodide ion." The spectral characteristics of solutions of BaI2 and SrIs are different from those of the alkali iodides, the extinction coefficients of solutions of the former compounds being about 6% less than those of the latter compounds. The marked change in the temperature dependence of Ba12 and Sr12solutions suggests that either the nature of the species or the cavity (or both) has been altered. It is possible that the band observed for Ba12 and SrIz solutions is due, in part, to solvent-shared ion pairs, because the band attributed to these species is far less sensitive to temperature
1495
changes than is the band associated with free solvated iodide ion. l6 The spectra of liquid ammonia solutions of CaIa are featureless in the ultraviolet region, indicating that the concentration of solvated iodide ions is negligible in this system. This obsemation leads to the suggestion that the predominating species present are solvated molecules, probably of the type CaI2 -4NH8, where the calcium atom is octahedrally coordinated.
Acknowledgment. We gratefully acknowledge the support of the National Science Foundation in the form of a grant (NSF G-15734). (16) T.R. Grifiiths and M. C. R. Symons, Mol. Phye., 3, 90 (1960).
Volume 70, Number 6 Mall 1066
