I
S. R. Logan
University of Strathclyde Glosgow,Scotland
The Solvated Electronthe Simplest
The existence in solution of charged particles, positive and negative, which could move about freely and independently, was first postulated in 1834 by Faraday (1). In recent years it has become recognized that in water or another polar solvent we may have the simplest ion of all, the electron itself. Although strong solvation of an ion is normal in such a liquid, the prefix "solvated" (or if in water "hydrated" or "aquated") is usually appended: this custom appears to have been established while the concept was still sufficiently novel to merit some explanation. One method of producing solvated electrons in certain non-aqueous solvents has been known for some time. The blue coloration and the very high electrical conductivity of solutions of alkali metals in liquid ammonia were attributed many years ago to the presence of ammoniated electrons (9),but since hydrogen is evolved immediately when sodium is added to water it was assumed that the aqueous counterpart did not exist. I n any case, the properties of e - ( a m ) might he altered by the presence of alkali metal ions, so that other methods of $reparation are desirable even for liquid ammonia. Prepamtion of the Solvated Electron
Three main methods have heen used to produce hydrated electrons in water. These may be employed with or adapted to other solvents. Photoionimfion of Reducing Ions
Several reducing ions, i.e., ions which readily give up an electron to an acceptor, strongly absorb ultraviolet light in a process which may he represented eqn. (I),
h"
Yz.(rtq)
Ycz"+o(aq)
+ e-(aq)
when Y may denote a halide ion (3-5), a metal ion such as the +2 state of a first-row transition element (6), or a complexed ion such as ferrocyanide (7). Similar r e sults may be achieved by illuminating aqueous solutions of aromatic compounds such as phenols (8). That the process represented by eqn' is not a simple photo-ejection is evidenced by the fact that the quantum yield is in general temperature-dependent (3) and is always less than unity. For example, a t 25'C using 2537 A radiation the quantum yield is 0.24 when y is I- (3) and 0.66 when [ F ~ ( c N ) ~ ](7). ~ - spectrascopic studies have yielded the picture (9) that the initial process on absorption of the cluantumis the promoGon of an e l e ~ t r o ~ f r oamp orbital to an orbital in the potential well formed by the oriented solvent molecules. From this state, a fraction of the electrons can au~arentlvescaDe. the ~ r o ~ o r t i ode~endine n on a number of factors, and'thus thebroducts df eqn. (1) are realized. 344
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Journal of Chemicol Education
1011and
Reagent
Rodiolyris of Water ( 1 0 )
High-energy quanta such as X- and prays produce highly energetic electrons; other ionizing radiations consist of fast electrons or ions. When a fast charged particle passes through water it imparts some of its energy to the electrons of the adjacent molecules. Excited states may he produced in cases where the acquired energy is small (HzO--ct HzO*),but some electrons will acquire more energy than is required to achieve ionization. These will travel from the site of the parent molecule a distance which will depend on the initial energy of the electron and on the energy lost in inelastic collisions. This distance is not easily calculated but it now seems that on the average it is s a ciently long that when the kinetic energy has been r e duced to kT the electron is not drawn back to the parent ion. During the period of about lo-" sec required for the thermalized electron to become hydrated, the HaO+ ion will have decomposed so that the overall process may be represented by eqn. (2).
---
H20(1)
H+(aq)
+ .OH + e-(aq)
(2)
This is now regarded as the most useful method of preparation, since it possesses several desirable features: (1) no solutes are required, (2) the presence of solutes does not interfere with the process described above, (3) it may be used over the entire pH range, and (4) equipment has been developed which produces very high dose rates and thus enables electron concentrations of up to M to be obtained. However, in using the radiation one has always to cope with chemical comp~ica~~ons arising from the generation of other reducing entities (H. and/or HzO*) and of OH radicals and from the non-homogeneous nature of the primary processes which are concentrated in particular regions, called "spurs," along the tracks of the fast charged particles, The Reodion of H Atoms with Hydroxide ions
When a stream of hydrogen atoms is passed into a solution of pH > 12, conversion to hydrated electrons proceeds with almost efiiency (11), according to eqn, (3), H OH-(aq) e-(aq) HzO (3) The pH range in which this technique might be used is thus a very small 0% but the fact that the reaction occurs is of some importance, as will be discussed later.
+
-
+
Some Properlies of the Solvated Electron
Color. In most solvents the electron has an ahsorption within the visible or near infrared ranee of the spectrum. The wavelength of the absorption maximum depends on the solvent and on other circum-
-
stances. In liquid water, the absorption may be seen by the flash photolysis of solutions containing reducing ions (12) or by the analogous technique of pulse radiolysis (IS) where a large dose of radiation is applied in an intense beam of electrons of Mev energy and wsec duration and the absorption spectrum of the irradiated sample is measured immediately afterwards. Ih water, the maximum occurs a t 7200 A, as shown in Figure 1.
0
4000
4500
h
5000
as the relation between the rate constant, k, of the reaction a t 25°C between the hydrated electron and an ion A of charge ZAand the ionic strength, rr. The first investigations of this effect were carried out before the perfection of the pulse radiolysis technique enabled absolute rate constants to be determined with such ease. By more readily available means, radiation-chemical or photochemical, it is convenient to measure the ratio of the rate constants with two species, A and B. If B is uncharged, then eqn. (4) (with a diierent constant on the right hand side) is still applicable to the relative rate constant, k&,. Some results from these studies, first reported in 1962 UQ), are shown in Figure 2 which clearly demonstrates that in this respect the reducing entity behaves like a uni-negative ion.
6000 7000 9000
61
Figure 1. Absorption bond due to the hydrated electron in irrodiotsd aqueous solutions: (01 daoerated 0.05 M solution of Not (b) deoeroted pure water, given the same dose from the electron beam. From ref. (1301.
COi,
Hydrated electrons, not mobile but trapped in a matrix, may also be studied (14, 16) in an alkalme aqueous glass such as is produced when a 6 N solution of NaOH is rapidly cooled from room temperature to liquid nitrogen temperature. The exposure of such a low-temperature glass to -prays or, if a small amount of ferrocyanide has been added, to ultraviolet light, causes it to turn blue. This coloration may be examined a t greater leisure, since it is stable a t these temperatures except when subjected to very strong light. In this case, with the electron embedded in a non-crystalline matrix of water molecules, the absor~tionmaximum is a t 5850 A. Unpaired spin. Although it may be feasible to detect the electron spin resonance spectmm of this transient in irradiated water bv the techniaue em~lovedwith reactive intermediates"in other liqGds (16j, no results bave yet been reported. Studies of the ESR spectra of irradiated glasses bave been carried out by several workers (14, 17). A singlet peak near the free electron value is detected whenever the blue coloration is present, decreases in intensity as it is destroyed, whether by thermal or photo-bleaching, and is attributed to the trapped electron. A diierent means of producing the trapped electron for such studies has been employed by another group (IS), who detected a similar peak when water vapor and sodium vapor were alternately condensed in thin layers on a surface a t - 19G°C. Charge. The expectation that the hydrated electron would possess normal properties of a negative ion led scientists to believe that its reactions with ions would exhibit kinetic salt effects. The Br6nsted-Bjerrum theory of primary salt effects leads to eqn. (4) log,&
=
+
1.02 Z . . Z~rr'/a/(l a'/*)
+ eonst
(4)
Figure 2. Ionic strength effects on the rote of reactions between hydrated On and H + (19b); (b) A g f ( 1 9 4 and [KFeelectron, ond ( 0 )
lCNld--I19d.
NOz-,
The Interconversion of Reducing Species in Water
Evidence accumulated during the decade 193NO (20) that both oxidizing and reducing radicals were produced in irradiated water. In solution, reducing ions (e.g., Fe3+, NO,-) were shown to be oxidized and oxidizing ions (e.g., MnOa-, Ce4+,Nos-) to he reduced on radiolysis. The reducing species were normally assumed (21) to be H atoms, which was not so far off the mark since the solutions employed were frequently of low pH. Later, i t was shown (22) that, a t natural pH, the reducing species produced initially is not the same as the product of the reaction between the oxidizing radical (assumed to be .OH) and Hz: the reactions of these two entities with O2and H201had quite d i e r e n t relative rate constants and so the difficult problem arose as to which was the true H atom. That no such problem has been encountered a t low pH is largely due to the fact that here a reaction occurs, as represented in eqn. (5), by which the hydrated electron is rapidly converted to tbe H atom. e-(aq)
+ Rc(aq)
-
H.
(5)
Thus in acid solutions the reducing species which react with other solutes (normally present only in concentrations of about M ) are effectivelythe same whether produced directly or by the reaction, represented in eqn. (6), of .OH radicals with HZ. OH+Hs-H.
+Hz0
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345
Other chemical means of differentiating these two reducing species1 are now known. Nitrous oxide (93) reacts with e-(aq) and with H, a t rates which differ by a factor of about 10'. The stable molecule Nz is a product of both reactions and is easily estimated. N,O has been widely used as a selective solute, as for example in the photochemical work of Dainton and Sills ($4) or in the radiation chemical studies of Dainton and Peterson (23). In each case, the yield of nitrogen is found to decrease at pH < 3.5, since reaction ( 5 ) is now able to compete with reaction (7) e-(aq)
+ NIO
-
N1
+ 0- (-H* OH)
stream of partially atomized hydrogen was passed into chloroacetate solution are shown in Figure 4. Other Reactions of the Hydrated Electron
Two methods of investigating the rates of reactions of the hydrated electron have already been referred to, the competition method and the pulse radiolysis method. An example of the former is the photolysis of iodide
(7)
and the consequent H atom reacts only very slowly with N20.
Figure 4. The effect of p H on the ratio, CI- formed: H otornr injected using 0.01 M chioroocetote rolution. (48, data from ref. l I].
solutions containing nitrous oxide and dichloroethane (PI), where reaction (11) takes place, followed by (7) or (12).
Thus the quantum yield of nitrogen, +(N2),may he represented as that fraction of the quantum yield of hydrated electrons, +,., which react with N,O by reaction (7). This is given by eqn. (13), inversion of which leads to eqn. (14)
Figure 3. Yields of H l and of Cl-from y-radiolyred 0.1 M solutions of ehloroacetis acid or a functionofpH. From ref. (25bl.
Both of these reducing species react with chloroacetic acid but, as shown by eqns. (a), (9),and (10) they react in diierent ways (95). e-(aq) H
+ CI CHKOOH
c
+ C1 CH2COOH-
-+
HS
Cl-
C1-
+ .CHCOOH
C1 CH COOH
+ H + + .CH,COOH
which predicts that the reciprocal of the nitrogen quantum yield should be a linear function of the solute concentration ratio, [(CH2C1)J/ [N,O]. This is shown in Figure 5, from which it can be deduced that kl& =
(8)
(9) (10)
The attack of the hydrated electron produces a chloride ion whereas the H atom reacts predominantly by H atom abstraction. Thus another illustration of the effect of pH on the importance of reaction (5) is given by the product yields from radiolysis of chloroacetic acid solutions, shown in Figure 3. The same system may be used to illustrate the reverse reaction of eqn. (3) and the results obtained by Israeli workers when a
' I t may be noted that the H atom does not always act as a reducing agent. I n acid solution it oxidizes Fea+ to FeS+in a reaction of which the sloiehiometn~may be written as: H H + Fez+ Hz Fe3+.
+
+
-
+
346 / Journol of Chemical Education
Figure 5. Plot of (O(NII]against [ ( C H h l J d l N1O from eiperirnents at natural p H using o 1 0 - = M solution of Ki. From ref. (3).
*
0.15 0.01 and so the ratio of the rate constants is determined. In determining rate constants by the pulse radiolysis technique (26), the optical density of the irradiated solution a t a wavelength near the solvated electron absorption maximum is measured continuously over a period of tens or hundreds of psec following the pulse of fast electrons. Since the concentration of solute, S, even if only lo-' or M, is much higher than that of electrons, the kinetics of the reaction are pseudo firstorder. The first order rate constant derived from the decay curve is a measure of k,+,[S]. If solutes which react readily with the hydrated electron are used in pulse radiolysis studies, the ratio of any two absolute rate constants may he determined. Comparison of this value with the relative rate constant obtained from competition studies shows that agreement is usually good (271, but that on occasions there may he some discrepancy. This may sometimes be due to errors of measurement, hut in other cases it seems to be due to other causes arising from the nature of the reactions and the measurements ($8). Thus the pulse radiolysis technique, although it gives more specific information and if mechanized can be much more rapid, is unlikely to become the only method of study. Since the first observation in 1962 of the absorption spectrum of the hydrated electron, numerous rate constants have been determined by the pulse radiolysis technique. A few of these are listed in the table. It Rate Constants Determined by Pulse Radiolysis Technique
Solute
Rate Constant, M-I seerL
Reference
compounds have been attempted by Hart and coworkers (54). Although there may he some uncertainty as to the process to which the rate constants refer (M), a comparison of values obtained by pulse radiolysis should be justified. It has been shown that for a halide, RC1, there is a relation between the rate constant and the electron-withdrawing capacity of the group R, as measured by Taft's 8-function (36), which would be expected if the reaction were essentially a nucleophilic attack of the electron on the halogen. I n carrying out reductions using sodium in liquid ammonia (the Birch reduction (56)), organic chemists have for many years been applying, though perhaps unwittingly, the simplest nucleophilic reagent. In this medium, aromatic compounds may be reduced to nonbenzenoid compounds and alkynes to alkeues. The fact that the product in the former case has unconjugated double bonds and in the latter case has the trans configuration are due to characteristics of the consequent radical-ion and dianion rather than of the attack of the amrnoniated electron. Thermodynamic Properties
Recent pulse radiolysis studies of highly purified, slightly alkaline (pH &9) water saturated with hydrogen after deaeration have shown that the half life of the hydrated electron is (at least) 780 pec. If it is assumed to decompose by reaction (15), e - ( 4 f HzO(1) = H(aq) OH-(ad (15) then the rate constants are known both for this reaction (87)(16 M-I sec-I) and for the reverse (38) (reaction (3), k = 1.8 X lo7M-I sec-') at 25'C. From the equilibrium constant, the standard free energy change A@ = 5.87 kcal/mole, may be calculated (39) for the standard states of pure liquid water and ideal 1 m solutions of the other species. Alternatively, treating the H atom as a very weak acid we may write the equilibrium as in eqn. (16). H ( 4 e H+(aq) e-(aq) (16) The same data lead to the conclusion that the pK of this acid is 9.7 and Jortner and Noyes (39) have also deduced that for reaction (16) AH0 = 11.1 kcal/mole and ASo = -7.1 cal mole+ deg-I. Thus, in neutral or acid solution the hydrated electron is expected to he unstable with respect to dissociation to H atoms, hut it is stable with respect to this reaction in moderately alkaline solutions. Estimates by Jortner (40) of the analogous ammoniacal system indicate that the counterpart of reaction (15) is endothermic by about 28 kcal/mole. Thus the equilibrium in eqn. (17) H(amm) e H+(amm) + e-(amm) (17)
+
+
may be noted that in many cases the second-order rate constants are in excess of 10'0 M-I sec-' and thus fall within the range of diffusion-controlled reactions. Of reactions with metal ions of a particular series, the fastest normally occurs in the case of a metal where the product is a stable ion. Thus the reaction with Cu2+ is faster than with Fez+, Co2+ or Ni2+ and with EuS+ faster than with neighboring trivalent rare earth ions. However, from the study of ions such as the cobaltammines i t would seem that this criterion is applicable only to the valency state of the product and not to the actual complex. Systematic studies of the reaction of the hydrated electron with alkyl halides and other halo-aliphatic
lies well to the right, indicating that the solvated electron has a much higher stability in ammonia than in water with respect to unimolecular decomposition. However, hydrated electrons are known (41) to undergo himolecular reaction in accordance with eqn. (18) 2e-(sq) + 2Hz0 Hz + 20H(18) and the analogous reaction in liquid ammonia is estimated to he exothermic by about 50 kcal/mole (40). On this basis, the solvated electron is not truly stable in liquid ammonia and the apparent stability can only be
-
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attributed to a high activation energy for himolecular reaction or to an inhibiting effect on this reaction of the alkali metal ions. By use of the equilibrium constant for reaction (15) together with an estimate of the heat of hydration of the H atom, Jortner and Noyes also estimated that the free energy of hydration of gaseous electrons at the electrostatic potential of bulk water is given by AGO = -39.4 2 kcal/mole. This value is much less than that of other anions. For example, the heats of hydrai tion of the halide ions have been estimated (48) as 123 (F-), 89 (Cl-), 81 (Br-), and 72 kcal/mole (I-) and since the entropy changes are only a few eu the free energies of hydration will differ little from these figures. Thus the energetics indicate that in the process of hydration the electron behaves as a very large univalent anion: the effective radius has been calculated as 3.0 A. It is not clear what this figure signifies, whether the electron is dispersed over the water molecules occupying a volume of these dimensions or whether it is confined to a cavity of this size. The latter view is favored for liquid ammonia in view of the considerable expansion when sodium is dissolved. I n the more open structure of water, the actual expansion due to the electron might well be much less. Attempts have recently been made (43) to estimate the equivalent conductance of the electron in water from observations of the transient conductivity on pulse irradiation. The results indicate that at 25°C the diffusion coefficient is 4.7 X cmZ sec-' which is not as large as was earlier surmised. Thus in water the electron is more mobile than 0% (2.5 X hut less so than H+ (9.5 X lo-' cm2sec-I).
*
Conclusion
Since the solvated electron was discovered, great strides have been made in our knowledge of the radiation chemistry of water and other liquids. But this discovery has made an even greater contribution to physical chemistry in general, in ways which perhaps illustrate the nature of scientific progress. This may be exemplified by reference to the study of kinetic salt effects, which (as has been said) were so useful in identifying the electron in irradiated water. I n later work (3) it was found that a different relation between rate constant and ionic strength may apply if conditions are such that the half life of the hydrated electron is appreciably less than the ionic relaxation time of the solution, so that the reacting species lacks the usual ionic atmosphere of a negative ion. Modiying the equations in the Brensted-Bjerrum derivation to take account of the modified activity coefficients under these conditions, it was shown that the kinetic salt effect should here be half the normal. The behavior of the hydrated electron as a very reactive ion has caused the theory of kinetic salt effects to be extended in other ways as well. The discovery (7) of very high salt effects in the reactions of electrons generated photochemically from ferrocyanide ions in solutions of fairly low ionic strength led to the consideration of a remting ion in fairly close proximity to a highly charged ion. More recently, calculations have been made (&) of the kinetic salt effects which would be predicted if reaction rates are assumed to he diffusion controlled. The results show qualitative similarity to 348
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Journal of Chemical Educafion
the effects predicted in equ. (4) and indicate that the dependence of rate constant on ionic strength is intluenced by factors which have no counterpart in the Brensted-Bjerrum formulation. Until now, the consideration of either of these problems might have been regarded as a purely hypothetical exercise. Another example of this contribution may he seen in regard to the process written in eqn. (1). Detailed studies of the absorption spectra of reducing ions were carried out about a decade ago and the interest in this subject may be said to have assisted in the discovery. The model postulated by Franck and Platzman (46) for the photoexcited ions is essentially that which led to the prediction (46) of the stability of the solvated electron, giving support to the views of Stein (47) that in an irradiated liquid a thermalized electron might polarize the medium and "dig" a potential well in it, thus decreasing the probability of the electron being pulled back by the positive ion. The challenge is now presented to elucidate the processes following the primary absorption. It is not clear whether this is basically a question of spectroscopy, of chemical kinetics, or of the structure of liquids, but attempts to solve it may be very illuminating-and may also have some success! The author would like to make particular acknowledgment to the review of this subject by Dr. F. S. Dainton (48), to the respective authors and editors for permission to reproduce Figures 1-5 and to Dr. D. Smithies for helpful criticisms of the manuscript. Literature Cited (1) FARADAY, M., Phil. T ~ a n sRov. . Soe., (London), 124, 77 (1834). C. A., J . Am. Chem. Soe., 30,1323 (1908). (2) KFCAUS, (3) DAINTON, F.S., AND LOGAN,5. R., Proc. Roy.Soc., (London), A287. 281 11965). (4) JORTNER, J., OTPOLENGKI, M., AND STEIN,G., J. Phys. Chem., 66, 2037 (1962); 68, 247 (1964). P., Proe. Roy. Soe., (London), (5) DNNTON,F. S., ANE FOWLES, A287, 312 (1965). (6) DAINTON, F. S., AND JAMES,D. G. L., Trans.Faraday Soc., 54, 649 (1958). (7) AIREY.P. L..AND DAINTON. F. S.. P ~ cROW . SOC..(London).
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DOESON, G., AND GROSSWEINER, L. I., Trans. Faraday Sac., 61, 708 (1965).
SMITH, M., AND SYMONS, M. C. R., Trans. Famday Soc., 54, 346 (1958); STEIN,G., AND TREININ, A,, Trans. Faraday Soc., 55, 1086 (1959). DAINTON, F. S., Radiutia Research, Supplement 1,1(1959); MATHESON, M. S., Ann. Rev. Phys. Chem., 13, 77 (1962); Radiation Research, Supplement 4, 1 (1964). JORTNER, J., AND RABANI,J., J. Am. Chem. Soc., 17, 388 (1961l ~-~~, (12) GROSSVEINER, L. I., SWENSON, G. W., AND ZWICKER, E. F., Scince, 141, 805 (1963); MATHESON, M. S., MULAC,W. A,, AND RABANI,J., J . Phys. Chem., 67,2613 (1963). (13) (a) HART,E. J., AND BOAQ,J. W., J . Am. Chem. Soc., 84, 4090 (1962): Nature, 197, 45 (1963); (b) KEENE,J. P., Nature, 197, 47 (1963). D.. AND EIBEN.K.. Z. Naturfwsch.. (14) . . SCHUL~FROHLINDE. l7a, 445 (1962); 18s, &I(1963). (15) GOPINATHAN, C., Ph.D. Thesis, Leeds (1966). R. W., AND SCKULER, R. H., J . Chem. Phys., (16) FESSENDEN, 39, 2147 (1963). (17) H E N ~ S E NT., , Radiation Research, 23, 63 (1964);
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AYSCOUGK, P. B., COLLINS, R. G., AND DNNTON,F. S., Natuw, 205, 965 (1965). (18) BENNETT, J. E.,MILE,B., AND THOMAS, A,, Nature, 201, 919 (1964).
(a) COLLINSON, E., DAINTON, F. S., SMITH,D. R., AND Tnzn~b,S., Proc. Chem. Soc., London, p. 140 (1962); (h) CZAPSKI,G., AND SCHWAFZ, H. A,, J. Phys Chem., 66,471 . (1962); (c) DUNTON,F. S., AND WATT, W. S., P ~ o e Roy. Soc., (London), A275,447 (1963). see, e.g., HART,E , J., J. CHEM.EDUC.,36, 266 (1959); Seia c e , 146, 19 (1964). WEISS,J., Nature, 153, 748 (1944). HOCHENADEL, C. J., J. Phys Chem., 56, 587 (1952); BARE, N. F., AND ALLEN,A. O., J. Phys Chem., 63, 928 (1959). DAINTON,F. S., AND PETEWN, D. B., Nature, 186, 878 (1960); Proc. Roy. Soc., London, A267, 443 (1962). DAINMN,F. S., AND SILM, 8. A., Natuve, 186,879 (1960). HATON,E., AND WEISS, J., Proc. S e m d Intern. Cmf. P m e ful Uses Atomic Energy, Geneva, 29, 80 (1958); HAYON, E., AND ALLEN,A. O., J. Phys C h a . , 65,2181 (1961). DORFMAN, L. M., AND MATHESON, M. S., Prog. R m d i m Kinetics, 3, 239 (1964). see, e.g., ALLEN,A. O., Radiation Research, Supplement 4,54 (1964); ANBAR,M., AND NETA,P., Intern. J. Appl. Radiation Isotopes, 16, 227 (1965). LOGAN.S. R., AND WILMOT.P. B.. Chem. Cwnm... . D. 558 (1966). GORDON. S.. HART.E. J.. MATHESON. M. S.. RABANI.J.. AND TEO~AS, J. k.,D$C. Fataday koc, 36, 193 (1963); J. Am. Chem. Soc., 85, 1375 (1963). ANBAR,M., AND HART,E. J., J. P h y ~ Chem., 69,973 (1965). HART, E. J., GORDON, S., AND THOMAS, J. K., J. Phys. Chnn., 68, 1271, 1524 (1964). BAXENDALE,J. H., et al., Natwe, 201,468 (1964). BAXENDALE, J. H., FIELDEN,E. M., AND KEENE,J. P.,
Proc. Roy. Soc., (London), A286.320 (1965). M., AND HART,E. J., J. Phys. Chem., 69,271 (1965); (34) ANBAR, SEUTKA,A., %OMAS,J. K., GORDON, S., AND HART,E. J., J. Phw. Chem., 69, 289 (1965). (35) T m ,R. W., "StericEfIectsinOrganic Chemistry," (Editw: NEWMAN. M. S.) John Wilev & Sons. Inc.. New York. 1956, chap. 13. ' (36) BIRCH,R. J., Quart. Rar., 4, 69 (1950). (37) HART,E. J., GORDON, S., AND FIELDON, E. M., J. Phys. Chem., 70, 150 (1966). M. S., AND RABANI, J., J. Phys Chem., 69,1324 (38) MATHESON, (1965). J., AND NOTES,R. M., J. Phys. Chem., 70, 770 (39) JORTNER, (1966). (40) JORTNER, J., Radiation Research, Supplement 4.24 (1964). L. M., AND TAW, I. A., J. Am. Chem. Soc., 85, (41) DORFMAN, 2370 (1963). (42) LATIMER, W. M., "Oxid&m Potentials," 2nd ed., PrenticeHall, Inc., Englewood Cliffs, N. J., 1952, p. 23. (43) SCHMIDT, K. H., AND BUCK,W. L., Seimee, 151,70 (1966). S. R., Tram. Faraday Soe., 62, 3416, 3423 (1966). (44) LOGAN, (45) FBANCK,J., AND PLATZMAN, R. L., "Parkas Memorial Volume," Jerusalem, 1952, p. 21; 2. Physik, 138, 411 (1954). (46) Platzman, R. L., "Physical m d Chemical Aspects of Basic Mechanisms in Radiobiology," Publ. No. 305, U.S. National Research Council, Washingtan, 1953. (47) STEIN,G., Disc. Faraday Soe., 12, 289 (1952). (48) DAINTON,F. S., J. Leeds Univ. Union Chem. Soc., 5, 3 (1963).
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