M.ANBARA N D EDWINJ. HART
1244
The Effect of Solvent and of Solutes on the Absorption Spectrum of Solvated Electrons'"
by M. Anbarlb and Edwin J. Hart Chemistry Division, Argonne National Laboratory, Argonne, Illinois
(Received October 17, 196.4)
The properties of electron-pulse-generated solvated electrons in ethylenediamine and in some concentrated aqueous electrolytes are reported. In ethylenediamine, a single intense optical absorption band with a peak a t 9200 A. is found and is assigned to the solvated electron. The maximum of the hydrated electron (eaq-) absorption band shifts from 7200 A. for pure water to shorter wave lengths in concentrated solutions of MgC12, KF, NaOH, KOH, NaC104,and LiCl. These results are discussed and compared with the absorption spectrum of solvated electrons in other solvents and with the charge-transferto-solvent spectrum of the iodide ion in the same media. In a 12.4 M solution of KF, the rate constants of eaq- reaction with benzoate ion, acetone, nitrous oxide, and nitrate ion are lower than the corresponding rate constants in water.
The detailed studies on the absorption spectrum of dilute solutions of alkali metals in ammonia, which revealed the agsorption spectrum of the solvated electron in ammonia,2 were followed by the discovery and characterization of the absorption spectrum of the hydrated electron in aqueous solution^.^-^ Recently the absorption spectra of solvated electrons in ethanol6 and in methanol' were also determined. A recent study on alkali metal solutions in ethylenediamines has assigned a certain absorption peak to the solvated electron in this medium. We have checked on this assignment by studying the absorption spectrum of solvated electrons formed by pulse radiolysis in ethylenediamine. It has been found that the absorption spectrum of the hydrated electron is modified in concentrated solutions of a m m ~ n i a . ~It? ~was of interest to investigate the effect of other solutes on the absorption spectrum of eaq-. In contrast to the case of ammonia where a shift to longer wave lengths was observed, various electrolytes tested affected the absorption spectrum of eaq- by a shift to shorter wave lengths. I t is the purpose of this paper to describe these findings and to discuss them in relation to solvent effect on charge-transfer-to-solvent (c.t.t.s.) spectra.
Experimental Anhydrous ethylenediamine (Fluka, Puriss.) was further purified by refluxing over potassium metal and The Journal of Physical Chemistry
redistilling under partial vacuum under a helium atmosphere. The pure air-free solutions were handled in syringes and transferred into the optical cell, as described e l s e ~ h e r e . ~The half-life of the solvated electron in this solvent was about 2 psec. Other chemicals used were K F (Baker and Adamson), NaC104prepared from NaOH (Baker Analyzed) and 70% HCIO, (G. F. Smith), KOH (Mallinckrodt), LiCl (Baker Analyzed), and MgC12(Mallinckrodt). The absorption spectrum of the solvated electron in ethylenediamine and in concentrated aqueous solutions was measured by determining the absorption of the transient formed on pulsed radiolysis a t different wave lengths. Full details of the experimental setup have been previously described.lo, l 1 (1) (a) Based on work performed under the auspices of the U. S. Atomic Energy Commission; (b) on sabbatical leave from the Weizmann Institute of Science, Rehovoth, Israel. (2) W. L. Jolly, Progr. Inorg. Chem., 1, 235 (1959). (3) J. W. Boag and E. J. Hart, Nature, 197, 45 (1963). (4) E. J. Hart and J. W. Boag, J. A m . Chem. Soe., 84, 4090 (1962). (5) J. P.Keene, Nature, 197,47(1963): Radicrtion Res., 27, 1 (1964). (6) I. A. Taub, D . A. Harter, M. C. Sauer, and L. M . Dorfman, J . Chem. Phys., 41, 979 (1964). (7) G. E. Adams, J. H. Baxendale, and J. W. Boag, Proc. Roy. SOC. (London), A277, 549 (1964). (8) R . R. Dewald and J. L. Dye. J . Phys. Chem., 68, 121 (1964). (9) S. Gordon, E. J. Hart, M. S. Matheson, J. Rabanl, and J. K. Thomas, Discussions Faraday SOC.,36, 193 (1963)
ABSORPTIONSPECTRUM OF SOLVATED ELECTRONS
The rate of reactions of eaq- with different organic and inorganic. reagents was determined by following tbe rate of decay of the absorption spectrum a t the 7200 A. maximum. Details of the experimental procedure anb calculation of the specific rate constants have been dec.ribed.g,12.13 The absorption spectrum of Sa1 (Baker and Adamson) in ethylenediamine was measured using a Cary recording ultraviolet spectrophoto5eter. As the solvent did absorb strongly below 2500 A. the spectra of 2 X lo-*, 5 X and 5 X M NaI solutions in ethylenediamine were measured down to 2500 A. using the pure solvent as reference. From these results the onset (e = 5 ) of the absorption band was determined. Further, the spectra of aqueous NaI solutions were measured a t the same concentrations and the parallel shift in the absorption curve was determined a t half the height of the peak value. From these measurements it is concluded that the iodide band in ethylenediamine is shifted 170 A. toward longer wave lengths, compared with the iodide band in aqueous solution. In this treatment we assume that the extinction coefficient and the shape of the absorption curve of the iodide ion in ethylenediamine and water are not significantly different. This assumption seems justified in view of results with iodide ion in other media.l4!l5
Results The absorption spectrum of the electron-pulsegenerated solvated electron in pure ethylenediamine (e,d,-) shows a single symmetrical band with a maximum a t 9200 d. falling to half its height a t 7000 and 11,200 A. (see Figure 1). Since these experiments were carried out a t eeda- concentrations of the order of Ab, we attribute this band to the isolated solvated electron. No other absorption bands could be detected in the range 5000-12,000 21. This implies that the band a t about 12,000 A. reported by Dewald and Dyes in rather concentrated solutions of alkaIi metals in ethylenediamine is due, not to eedn-, but probably to units such as ion pairs and more complex ions as suggested by them. Attempts to detect any of the other bands observed by Dewald and Dye, by adding potassium ions M to the pulse-radiolyzed solution, up to 2 X failed to produce additional bands. As the G value of the solvated electron in ethylenediamine is unknown, its extinction coefficient cannot be calculated. However, by assuming the same yield for eeda- as the enq- the extinction coefficient of eeda- has an approximate value of 8000 M - * em.-’ a t Amax. The absorption of eaq- was measured in concentrated electrolyte solutions. The results are summarized in Table I. I t is evident that A,, is shifted
1245
4000
I
I
6000
8000
x
,
IN
I
I
10,000
12,000
1
d
Figure 1. Spectra of electron pulse generated esq- in 15 M aqueous solution and eeda- in ethylenediamine: 0, 15 M LiC1; 0, H20(ref. 5); and X, ethylenediamine.
from 7200 A. for pure water to shorter wave lengths in the electrolyte solutions. However, the shape of the absorption bands remains similar, and their extinction coefficients are only slightly lower than that of eaQ- in pure water. The spectra of eaq- in 15 M LiCl and of eeda- in ethylenediamine are shown relative to water in Figure 1.
Table I : The Effect of Electrolytes on the Absorption Spectrum of eaqSolute
None MgCh
KF NaOH KOH NaClOd
KF LiCl
Concn., M
4.6 7.0 15 15 10 12.2 15
XmRx,
A.
7100 6500 6500 6500 6600 6050 6050 5900
Relative
c
1 .oo 0.88 0.72 0.93 0.90 0.90 0.83 0.70
X(0.5rmax),
A.
5700,8300 5000,7950 5100,7850 5050,8150 5200,81CO 4850,7450 5100,7200 4800,7200
The rates of reaction of several solutes with eaq- have been measured in 12 M K F and the specific rates are compared with those obtained in dilute solution (Table 11). In all cases, the rates of reaction are lower than in dilute solution, although a primary salt effect is ex(10) S . Gordon, E. J. 1262 (1964).
Hart, and J. K. Thomas, J . Phys. Chem., 6 8 ,
Thomas, S. Gordon, and E. J. Hart, ibid.,68, 1524 (1964). E. J. Hart, 9. Gordon, and J. K. Thomas, ibid., 68, 1271 (1964). (13) M. Anbar and E. J. Hart, J . A m . Chem. SOC.,86, 5633 (1964). (14) G. Stein and A. Treinin, Trans. Faraday SOC.,56, 1393 (1960). (15) I. Burak and A. Treinin, ibid., 59, 1490 (1963). (11) J. K.
(12)
Volume 69, Number
4 April 1966
1246
hl. ANBARA N D EDWIN J. HART
pected to increase the rates of reaction of benzoate and nitrate ions with eaq-. Whereas the rate of the reaction of eaq- with NO3- is decreased in the concentrated solution by a factor of 5 and those with acetone and N20 by a factor of 3-4, the rate of the eaq- reaction with benzoate ion is diminished by only a factor of 2.
50
e -'
-
40
a u
X
Table 11: The Reactivity of eaq- in Concentrated Electrolyte Solutions k(eaq- + R), M -1 sec. -1 HzO
Reactant (R)
12.4 M K F
CeHsCOOCHICOCHI N20 NO*-
(1.7f 0.15)X lo8 (1.6f 0.2)X 109 (2.4f 0.3) X lo8 (1.9f 0.3) x 109
a
Ref. 13.
* Ref.
12. ' Ref. 9.
(3.5f 0.3)X lo8' (5.9f 0.2) x LO@*" (8.7f 0.6)X logc (1.1f 0.1) x i01@*~
The analogy between the charge-transfer spectrum of halide ions and that of a trapped electron in the same medium had already been pointed out by Platzman and Franck in 19E14.I~ No experimental evidence was available at the time to corroborate this suggestion. The absorption spectrum of iodide ions has been investigated in the presence of other ions14*l7and in a number of solvents. 15,17 Having several absorption spectra of solvated electrons available, it was possible for us to examine quantitatively the Platzman and Franck hypothesis. The spectra of solvated electrons and of iodide ions are compared in Table 111. The re-
Table I11 : A Comparison between the Spectra of Solvated Electrons and Ions in Different Media electronsEXmax,
Solvent
Xmax.
'
A.
7,200 6,500 6,050 5,900 .6,3OOc 7,oooc 9,200 14,500' ~7,80W
HzO MgC12, 4.6 M KF, 12 M LiC1, 15 M MeOH EtOH (CHZNHdz(en) NHa 25% NHt in HzO
kcal./mole
39.7 44.0 47.2 48.4 45.4 40.8 31.0 19.7 37.1
Iodide ions EXmax,kcal./mole
126.5" 129.5* 132.4' (10 M ) 135.Ob (12 M ) 130.Id 130.I" 117.6 105.9' 123.6"
Ref. 15. Extrapolated from value in ref. 15. Ref. 6. Ref. 17. e R.ef. 2. T. R. Griffiths and M. C. R. Symons, Trans. Faraday Soc., 56, 1125 (1960). The first inflection in Figure 14,ref. 4.
'
The Journal of Physical Chemistry
E
x w al
30
20
Ref. 11
Discussion
--Solvated
X
Figure 2. Comparison of Exmax of solvated electrons and iodide ions.
sults, presented in Figure 2, show a linear correlation between the effects of the medium on the spectra of solvated electrons and of iodide ions.'* The charge-transfer-to-solvent spectrum (c.t.t.s) of an anion, X-, has been shown to follow the expresEx,,, = E x E, - LX - B, where E X sion15~16,19 = electron affinity of the X radical; L X = the heat of hydration of the X radical; E, = the energy involved in removing the ion from its cavity in the solvent, without disturbing the persistent polarization of the medium; and B = the binding energy of the electron to the polarized medium in its excited state. In the case of hydrated electrons it has been shownz0that Ex,,, of eaq- is equal to its energy of hydration. In this case E X = LX and B = 0; thus Ex,,, = E,. It has been suggested that the changes in Ex,,, of iodide are mainly due to changes in the E , term.15 Our results confirm this conclusion quantitatively and prove that
+
(16) R.L. Platzman and J. Franck, 2. Physik, 138, 411 (1954). (17) M.Smith and M. C. R. Symons, Trans. Faraday Soc., 54, 338, 346 (1958). (18) It should be noted that the A,, of the iodide ion was measured a t 10 M KF and 12 M LiCl,14 whereas Amax of eaq- was measured a t 12 M K F and 15 M LiC1. Correction for this difference in concentration would cause a shift of these points to the left in Figure 2 and, if the absorption spectra were measured a t the same concentrations of added electrolyte, an even better correlation between esq- and Ispectra should result. (19) G.Stein and A. Treinin, Trans. Faraday SOC.,55, 1086 (1959). (20) J. H.Baxendale, Radiation Rea. Suppl., 4, 139 (1964).
ABSORPTION SPECTRUM OF SOLVATED ELECTRONS
the effects of the medium on the spectrum of solvated electrons are the same as on that of iodide ions. In the case of concentrated electrolyte solutions it was shown that changes in E', are due to changes in the radius, r , of the cavity available to the ion in solution.14 The contraction of the cavity results in a shift to shorter wave lengths. The effect of LiC1, KF, and MgC12 on this parameter has been discussed by Stein and Treiniti,l4 and they predicted that anions which strongly bind the protons of H,O molecules will tend to decrease r. This prediction would imply a most extensive shift to the ultraviolet by OH- ions, contrary to our observations on the spectra of ea,- in concentrated NaOH and KOH solutions (Table I). As the dielectric properties of these solutions are comparable to those of KF, the observed difference in behavior of OHand I;--must be due to a difference in the magnitude of r. The discrepancy may be explained, however, by the formation of well-defined OH(H20)3- complexes,21 which leave larger cavities for e,,-, than those produced by the mutual repulsion of protons.I4 The effects of different solvents on Ex,,, are mainly due to changes in their dielectric properties which determine the value of E,.15 They affect the spectrum of the solvated electron in a manner similar to their effect on the charge-transfer spectra. I n other words, the value of Ex,,, for solvated electrons may be predicted using Treinin's c.t.t.s. ~ c a 1 e . l ~
1247
The rate of reaction of hydrated electrons with oxidizing agents must be affected by the energy required to abstract the electron from its sphere of hydration. This is confirmed by our findings that the rates of capreactions are slowed down in concentrated KF solutions (Table 11) in which an increase of 7.5 kcal./mole in Ex,,, was observed. It should be noted, however, that this decrease in rate is not equal for the different substrates examined and it appears that the more hydrated reactants, e.g., NO3-, are affected to a greater extent. I t is hard to believe that the primary salt effect on the BzO--e,,- reaction is greater than that on the NO3-ea,- interaction. This would suggest that the major effect of the environment on the rate is by changing the hydration sphere of the activated complex, and that the changes in the size of the cavity of eaq-, or the energetics of its hydration, are of relatively minor importance. This conclusion is supported by the finding that the rate of ea,- reactions in water and alcohols is e q ~ a l . ~ , ' ~ I n conclusion, it may be stated that our results have not only corroborated the quantitative theories of charge-transfer spectra but made it possible to predict the spectrum of solvated electrons in different solvents.
Acknowledgment. We gratefully acknowledge the value of our several discussions with Dr. R. L. Platzman. (21) G. Yagi1 and M. Anbar,
J. Am. Chem. SOC.,85, 2376 (1963).
Volume 69,Number 4
April 1966
