Chapter 7 Reactions of Solvated Electrons

Chapter 7 Reactions of Solvated Electrons

Chapter 7 Reactions of Solvated Electrons G. HUGHES and C. R. LOBB 1. Introduction The production of the hydrated electron by the interaction of ion...

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Chapter 7

Reactions of Solvated Electrons G. HUGHES and C. R. LOBB

1. Introduction The production of the hydrated electron by the interaction of ionizing radiation with water was one of the outstanding discoveries in chemistry in the 1960s. I t thus became apparent that reactions of this species, produced as a result of cosmic bombardment of the earth’s surface, must have been occurring from primeval times. It is, however, only relatively recently that it has been possible to study reactions of the hydrated electron in the laboratory. In this review, it will be convenient first to discuss results that have been obtained on the hydrated electron and subsequently t o discuss other systems in which reactions of solvated electrons have been observed. A wealth of information concerning the hydrated electron has now been amassed and this has led to a good understanding of many of its properties. Our knowledge and understanding of the solvated electron in other systems, however, is much less complete. 2. The hydrated electron 2.1 PREPARATION OF THE HYDRATED ELECTRON

2.1.1 Radiation chemistry As a result of the earliest observations on the radiolysis of water, it was suggested that hydrogen atoms and hydroxyl radicals were the principal reactive species [ l ]. Two theories were proposed to account for the formation of these. The first [ 2 ] supposed that an electron ejected from a water molecule by the radiation is thermalized before it escapes the coulombic attraction of the parent positive ion. Its lifetime is too short for it t o exist as a significant reaction intermediate and it is recaptured t o give an excited molecule which subsequently decomposes, viz.


e- + H 2 0 +


HzO+ + e-





H 2 0 -,-..A

References p p . 4 5 8 - 3 6 I




Alternatively, Stein [ 3 ] and also Platzman [ 4 ] proposed that the hydrated electron could be an important intermediate in the radiolysis of aqueous systems. Platzman suggested that the rate of energy loss of the ejected electron was such that it would escape the coulombic field of the parent positive ion before thermalization. Subsequently the electron would be hydrated, i.e. the electron polarizes the dielectric and then is bound in a stable quantum state to it, viz. H20 H20+


e- + H 2 0



H++*OH -


(1) (4)


As will be seen subsequently, later experimental evidence verified Platzman's postulate. In the late 1950s several experiments indicated that the reducing species was not simply the hydrogen atom. In neutral solutions, the hydrogen atom formed by the reaction .OH + H2


H20 + He


reacts with oxygen much faster than with hydrogen peroxide [ 5 ] . In the y-radiolysis of aqueous solutions of H, O2 and 0, , however, the reducing species formed reacts at an equal rate with H, O2 and 0 2 .Hydrogen yields in the radiolysis of aqueous methanol were lowered by the addition of Fe3+and Cu2+[ 6 ] . This could be attributed to competition between

H. + CH,OH



+ 'CH20II


However, Fe3+ and Cu2+ suppressed the hydrogen yields much more effectively than calculations using H atom rate coefficients predicted. These experiments showed that there are at least two reducing species formed in the radiolysis of water. Further evidence for the existence of two reducing species was found in the study of the dependence of hydrogen yield on the concentration of monochloracetic acid [ 7 ] in the radiolysis of aqueous solutions of the latter. Results are shown in Fig. 1. It was suggested that hydrated electrons react with monochloracetic acid t o give chloride ions e& + ClCH,COOH


C1- + .CH, COOH





mole I-'

Fig. 1. Effect of concentration of chloracetic acid on the yields of H, and C1- at pH 1. m, G(H,); 0, G(C1-).

whereas the predominant reaction of the hydrogen atom is abstraction of hydrogen H. + ClCH2COOH




Thus indirect evidence for the existence of the hydrated electron was overwhelming by the late 1950s. The principal reducing species in the radiolysis of neutral water was shown t o have unit negative charge from studies of the kinetic salt effect on its reactions [8, 91. The dependence of the rate coefficient on ionic strength, p , for a reaction in water a t 298°K is given by the BronstedBjerrum equation

where k and k, are the rate coefficients a t ionic strengths 1.1 and 0 respectively and ZA and Z , , are the algebraic charges on the ions. The reactions of the reducing species with H 2 0 z , NO;, O2 and H 3 0 t were investigated. Assuming that the hydrated electron is involved, these reactions would be e& + H,O,


OH- + .OH


The results are shown in Fig. 2. It is clear that all the observations are entirely consistent with a reductant of unit negative charge. Finally the discovery of the optical absorption spectrum of the hydrated electron in the pulse radiolysis of water (see Vol. 1, p. 67) by R P f e r e n c e s p p ,158-461

432 O*


2 __ ,./I+

Fig. 2. Effect of ionic strength on rate coefficient ratios for reactions o f the hydrated electron. (-), K = k 5 / k 3 ; X, K = k I q / k I 3 ; [I, K = k I b / k 1 3 .

Boag and Hart [lo] and Keene [ 111,independently, removed any doubts about the existence of this species. The absorption spectrum is shown in Fig. 3. The absorption spectrum is broad with a maximum at 720 nm. Solutions of the hydrated electron are consequently blue. The spectrum is similar t o that of the solvated electron observed in the blue solutions of alkali metals in liquid ammonia (see Section 3.1). Typical electron scavengers, e.g. H,O +,N, 0, suppress the spectrum. The absorbing species was shown to have unit negative charge from studies of the effect of ionic strength on its reaction with ferricyanide ion and was thus clearly identified as the hydrated electron.

Fig. 3. Absorption spectrum of the hydrated electron.

43 3 2.1.2 Photolysis of aqueous solutions

By observation of its absorption spectrum [ 1 2 ] , the production of e,, has been demonstrated in the flash photolysis of aqueous solutions of inorganic ions, e.g. Fe(CN):- + hv


Fe(CN)i- + e&

(17) (18)

I - + h v - + I - +e;,

Efficient production of hydrated electrons can be observed in these systems using radiation of 254 nm. More recently, Boyle et al. [13] have identified e,, as a product of the direct photoionization of water at 195 nm. Photolysis of an alkaline solution saturated with hydrogen provides a good source of e;, . The reactions occurring are OH- + hv


.OH + eiq

.OH + H2’HZO H. + OH-



+ Ha


e& + H 2 0


The net reaction is Hz + 2 0 H - + h v + 2 e i q + 2 H , O


i.e. the quantum yield of production of hydrated electrons is 2. Photolysis of aqueous solutions of certain organic substrates also produces e;, . C,H, NH2 + hv



C,H, NH2

t e;lq


Studies of aromatic compounds show that photoionization in aqueous solution occurs for benzene derivatives with electron donating substituents, e.g. amines and carboxylic acids [ 141 . 2.1.3 Chemical methods

The hydrated electron can be produced chemically from the hydrogen atom H.



eaq + H 2 0


Thus when hydrogen atoms, generated by a discharge in hydrogen gas, are passed into an alkaline, aqueous solution, reactions characteristic of the hydrated electron are observed [ 1 5 ] . The conversion of hydrogen atoms Rrferences p p . 458--461

434 to hydrated electrons can also be brought about by other powerful proton acceptors [ 161, e.g. Ha + F-


eiq + HF


k 2 and k 2 4 are 1.8 x lo7 and 1.0 x lo4 1 mole-’ sec-’ respectively [17, 181. Unless stated to the contrary, all rate coefficients reported in this review are for 298°K. It has been reported [19] that a film of alkali metal in contact with water appears blue and this colour was attributed t o the hydrated electron. Subsequent attempts to repeat this observation have not, however, been successful. Bennett e t al. [20] have shown that, at low temperatures, alkali metals deposited in ice show a characteristic blue colour. ESR studies have confirmed that this is due t o the trapped electron. That the reduction of water by sodium proceeds readily and sometimes violently has been known for many years. The powerful reducing properties of this reaction are generally attributed to the production of “nascent” hydrogen. Until recently, the precise nature of “nascent” hydrogen had never been established, though it was suggested that it might be the hydrogen atom.

Na + H 2 0 -+ Na’ + OH- + H.


Production of hydrogen in the system would then be via the reaction

H a + H.




Hughes and Roach [21] and Shaede and Walker [22, 231 designed experiments t o try to identify the precursor of hydrogen. Hughes and Roach studied the reaction of sodium amalgam in acid solution in the presence of various scavengers. The efficiencies of scavengers in suppressing the hydrogen yield were in good agreement with their known rate coefficients fcr reactions with hydrated electrons but did not parallel their rate coefficients for reaction with hydrogen atoms, indicating that the hydrated electron, rather than the hydrogen atom, was the important intermediate. Shaede and Walkcr studied the reaction of sodium amalgam with water in the presence of dinitrogen monoxide and methanol at various pH values. Radiolysis experiments had already established that dinitrogen monoxide is a very efficient, specific electron scavenger eaq + N 2 0 + N 2 + *O-


h , , being 8.7 x lo9 1 mole-’ sec-’ [ 2 4 ] . If hydrogen atoms are produced, they would rapidly be scavenged by methanol t o give hydrogen.

H- + CH,OH


H, + *CH,OH


435 However, yields of nitrogen were obtained in both neutral and alkaline solutions, indicating that the hydrated electron is produced in this system and is the precursor of hydrogen, viz. Na + H,O eHq + N,O eHq + e&


Na’ + e&


N, + *O-


H, + 2 OH-




Later work by Hughes and Lobb [25] on the sodium amalgam-acid system has further substantiated the hydrated electron mechanism. As discussed earlier, it is known that the hydrated electron reacts with the monochloracetate ion t o form chloride ion. eaq + ClCH, COO- + .CH2 COO- + C1-


However, when sodium chloracetate was used as a scavenger, the chloride ion yield was only 50% of that expected. It would seem that, in the amalgam reaction, the first product of reaction by the hydrated electron is localized at the amalgam surface and is therefore available for reduction by a second hydrated electron, viz. eiq + ClCH2COO-+ *CH2COO-+ C1*CH, COO- + H2 0

* *CH2COOH + OH-

*CH2COOH+ ellq



(30) (31) (32)

Similar observations [23] were made in the study of the dinitrogen monoxide system where reaction of the hydrated electron with the hydroxyl radical produced as a result of the reactions e& + N,O


N2 + -0-

*O-+H2O **OH + O H -

(27) (33)

was observed, viz. eaq + *OH+ OH-


The authors have investigated other systems [ 251 including A1-NaOH and MgHC10, and results indicate that the hydrated electron is an important intermediate in these systems. Formate ion in basic solutions can be used t o distinguish between hydrated electron and hydrogen atom mechanisms. Its reaction with electrons eaq + HCOO- -+ (HCO0)’(HCO0)’- + 2 H 2 0 + HC(OH), + 2 OHReferences p p . 458-461


436 is slow [ 2 6 ] , k 3 5 being 2.4 x lo4 1 mole-' sec-'. By contrast, it is a relatively efficient scavenger of hydrogen atoms [ 271


Ha + HCOO-


H2 + *COO-


k, being 2.5 x 10' 1 mole-' sec-' . Moreover occurrence of reaction (37) would lead t o an increased hydrogen yield. N o increase in hydrogen yield was observed in the dissolution of aluminium in sodium hydroxide in the presence of high concentrations of formate ion. The reactions occurring may be represented as

A1 + 4 OH-


Al(OH), + 3 ea,

e;, + eaq+ H, + 2 OH-

(38) (29)

It is still not clear whether the hydrated electron in amalgam reactions ever becomes entirely free of the mercury surface and further work needs t o be done. The reduction of water by U3+ and Eu2+ has also been investigated using the N, 0-CH, OH scavenging method [ 231. Results indicate that the hydrated electron may be the intermediate here also. It has been suggested [22] that the hydrated electron could be an intermediate in cathodic reduction, viz. ec;itt,ot1e +

H2 0




Walker [28] experimented with an illuminated cathode. The amount of light ( h = 632.8 nm) reflected a t a silver surface decreases when it acts as a cathode. This could be attributed to the production of the hydrated electron at the surface of the cathode but the conclusions are not unequivocal. Conway and Mackinnon [ 291 have discussed the kinetic and thermodynamic difficulties which arise from this type of mechanism involving hydrated electrons in electrode processes. 2.2 PROPERTIES O F THE HYDRATED ELECTRON

I t has already been seen that the hydrogen atom and the hydrated electron are interconvertible. Thus they form a conjugate acid-base pair.

The equilibrilim constant for the above equation is given by

2x -~

10' 16

= 1.25 x 10'

437 The acidic dissociation constant for the hydrogen atom H. + H+ + e,,


is defined by

It follows that

For water at 298"K, [H'] [OH-] 5 mole2 1-2 and [ H 2 0 ] = 55 mole 1-' , and consequently K , is 2.3 x lo-' mole 1-' . This corresponds t o a pK, of 9.6 for the hydrogen atom. The precise value of k - 4 0 is still the subject of speculation (see p. 439) and i t may be <16 1 mole-' sec-'. The value of pK, must be regarded as an upper limit. The standard free energy change for the reaction


H-+OH-+e& + H 2 0 is given by



--RT In K40


-8.34 kcal mole-'

I t is thus possible t o calculate the standard free energy change for the reaction

by considering the cycle [30]

-Haq + OHH,O

-H, Net:



+ H20

AGO (kcal mole-' -8.3














H,O+ + OH-


1 H2



H 3 0 ++ e,,

References p p . 458-461

438 Since AGO = --FEo, this corresponds t o a standard electrode potential of -2.77 V. The hydrated electron is thus an extremely potent reducing agent, much more powerful than the hydrogen atom, for which Eo is -2.10 V . The free energy and heat of hydration are -37.4 and -38.1 kcal mole-', respectively. The entropy of hydration is -1.90 cal mole-' OK-' [31]. The diffusion coefficient for the hydrated electron has been calculated from conductivity measurements [32] to be (4.9 f 0.25) x lo-' cmz sec-' . The equivalent conductance is much higher than that of all other ions except OH - and H, 0 ' but considerably lower than the ammoniated electron. The diffusion mechanism is uncertain at the present time partly because of the lack of detailed knowledge of the structure of the hydrated electron. The radius of charge distribution of the hydrated electron has been estimated by several methods. The values obtained from hydration energies, encounter radius, i.e. the radius required to accou'nt for experimental diffusion controlled rate coefficients, and Jortner's cavitycontinuum model are all in the range 0.25-0.30 nm [33]. This clearly indicates that the electron is not associated with a single water molecule only, but rather that the charge of the electron is "smeared out" over 3-4 water molecules. As mentioned previously, the absorption spectrum of the hydrated electron is broad and intense rising to a maximum at 720 nm with a molar extinction coefficient of 1.58 x l o 4 1 mole-' cm-' . This was determined from a study of the reaction of the hydrated electron with tetranitromethane [34] e& + C(NO,),


C(NO,)j + NOz


The nitroform ion has a characteristic absorption at 366 nm with an extinction coefficient of 1.02 x l o 4 1 mole-' cm-' . At this wavelength, the absorption due to the hydrated electron is weak. It is possible, therefore, to measure in the same experiment the decrease in optical density at 720 nm due to the hydrated electron with the corresponding increase in optical density at 366 nm due to the production of the nitroform ion and thus to determine the extinction coefficient of the hydrated electron. The spectrum broadens slightly and E h a x shifts towards lower energies with increasing temperature. Specific electron scavengers suppress the transient absorption obtained in the pulse radiolysis of water. By measurement of the efficiency of this process, it is possible to determine the rate coefficient for the reaction of the hydrated electron with the scavenger. Solutes are restricted to those which do not absorb significantly in some region of the hydrated electron spectrum. Computer programs have been developed for kinetic analysis of the oscilloscope traces and it is thus possible to obtain rate coefficients

439 with relative ease. I t is for this reason that the hydrated electron is probably the most widely characterized kinetic intermediate. Comprehensive lists of rate coefficients are available [35-371 and will not be presented again in this review. Instead some characteristic features of the reactivity of the hydrated electron are discussed in the next section. 2.3 REACTIONS OF THE HYDRATED ELECTRON

2.3.1 Introduction

As has been seen, the hydrated electron is a very powerful reductant. I t reacts with a wide range of both inorganic and organic compounds with rate coefficients ranging from the lowest with water of 1 6 1 mole-' sec-' to the diffusion controlled limit of N 10' 1 mole- I sec- . The rate coefficient for the reaction with water


e& + IH2O


OH- + H .


has been extremely difficult to measure accurately because of impurities in the water which react with the hydrated electron at considerably greater rates. Consequently extremely pure water is required. Hart e t al. [38] obtained a value for the rate coefficient of 1 6 1 1 mole-' sec-' at 298°K by working at an estimated impurity concentration of <5 x M in a hydrogen saturated solution at pH 9. This value of the rate coefficient should be regarded as an upper limit. Because the reaction of the hydrated electron with water is so slow, its lifetime is sufficiently long for it to be possible to observe its absorption spectrum and t o study its rates of reaction with added solutes. It is interesting t o calculate the lifetime under the typical conditions of pulse radiolysis. The concentration of hydrated electrons produced isM. Under these conditions, bimolecular decay of hydrated electrons is negligible and at pH > 9, reaction with water is the only observed process. The half life is given by the equation



0.69 16 x 55



sec are required for significant obserIt is clear that pulses of vations to be made. A t lower pH, the lifetime of the hydrated electron is, of course, much less because of the high rate coefficient for the reaction with hydrogen ion e& + HjO'


Ha + H 2 0

h , , b e i n g 2 . 2 6 ~10" l m o l e - ' sec-' [391. R e f e r e n c e s pp. ,158- 4 6 1


440 TABLE 1 Reactions of eaq with intermediates produced in the radiolysis of water Reaction eiq +




+ 2 OH-

e y + . O H + OHe,q + H* + H2 + OH-

+ H~O+ + H- + H ~ O eaq + H 2 0 2 + * O H + OH eaq + HzO -+ H * + OHeaq

k ( I mole-' sec-')


4.5 lo9 3 x loLo 2.5 x 10" 2.26 x 10" 1 . 3 6 x 10" 16

17 24 24 39 40


The radiolysis of water can be represented by the equation

Now the hydrated electron reacts with all the other species produced except OH- and H , . These reactions are all very rapid, except that with water, as may be seen from Table 1. In the study of specific reactions of the hydrated electron it is therefore necessary to eliminate as many of these ,other reactions as possible. Thus, by working at high pH, it is possible t o convert all hydrogen atoms to hydrated electrons. N-+OH-+e;,



A t the same time, this eliminates occurrence of reaction with the proton. Using high pressures of hydrogen, it is possible to convert all hydroxyl radicals to hydrogen atoms. . O H + H2


H,O + H .


A t high pH, these are in turn converted to hydrated electrons. In the radiolysis of solutions at low solute concentrations, oxygen must be removed since the reaction [41]

eiq + O 2




has a diffusion controlled rate coefficient of 2.16 x 10' 1 mole-' sec-' Carbon dioxide must also be removed since it behaves similarly ea, + C 0 2





lo9 1 mole-' sec-' [ 2 4 ] . Measurements of the isotope effect on the formation of hydrogen in the radiolysis of water have been used to determine reaction mechanisms [42] and in order t o explain the high isotope effect of 3.7 for the reaction

k, being 7 . 7 x


of the hydrated electron with the proton, it is necessary t o introduce an intermediate, H 30, which decomposes t o give a hydrogen atom and water, viz. e& + H 3 0 ++ H 3 0 .


H 3 0 -+ He + H 2 0


The general reaction with acids can be written in the form e& + HA + HA-


HA-+ H. + A-


In such reactions, the hydrated electron behaves as a base. However, it has been shown that with certain carboxylic acids, it behaves as a reducing agent [ 431 , e.g.




b0 I



+ OH-


+ H 2 0 + Hd(OH), + C 0 2

It is interesting that with acetic acid, the hydrated electron behaves both as a base and a reducing agent [ 441 , viz. eaq + CH3COOH




CH3CO* + O H 2.3.2 Inorganic anions There is no absolute rule governing the reactivity of anions towards the hydrated electron. In the case of oxyanions, the availability of a vacant orbital on the central atom is an important factor. Sulphate, carbonate, perchlorate and phosphate have rate coefficients
442 TABLE 2 Rate coefficients of reaction of e& with oxyanions [45--471 Solute


( I mole-' sec-')

3.3 x 1.8 x 2.2 x 8.5 x 1.1 x

10'O 10" 1o'O 109


I t is interesting that ClO, and Nj,despite their high redox potentials, react very slowly. The rate coefficients are

The first product of the reaction of the hydrated electron with the nitrate ion is the ion (NO,)'-. This subsequently reacts with water.

The hydrated nitrogen dioxide has been identified spectrophotometrically [ 4 1 ] . The reactivity of the nitrate ion is due in part to its planar structure and, in part, to its 7r-bond character which results in the nitrogen atom having a positive charge and being a reactive centre. Phosphorus compounds, e.g. H 2 PO;, HzPO; are far less reactive than the corresponding nitrogen compounds because of the absence of r-bond character. The hydrated electron reacts with the ferricyanide ion e,,, + Fe(CN)i-




at a fairly high rate [ 2 4 ] , h , , being 3 x lo9 1 mole-' s e c - ' . I t is clear that Coulombic repulsion need not be a significant barrier to reaction between highly charged anions and the hydrated electron.

443 2.3.3 Inorganic neutral molecules The hydrated electron reacts with hydrogen peroxide and dinitrogen monoxide t o form hydroxyl radicals, viz. eBq + H 2 0 2 h,






being 1.36 x 10" 1 mole-' sec-' [40]. eiq + N 2 0



N2 0' + I l 2 0 'N20H




.N2 OH + OH-

N2 + *OH

(65) (66)

k , , being 8.7 x l o 9 1 mole-' sec-' [ 2 4 ] . Both these reagents have been used extensively in radiolysis studies in which it has been desirable t o investigate reactions of the hydroxyl radical only. Reaction (64) is specific for electrons. I t has been mainly used as a diagnostic test for the production of electrons both in aqueous and non-aqueous media. The lifetime of N,O' is still the subject of some speculation. I t is probably >lo-' sec in alkaline solutions and at high concentrations of certain solutes the direct reaction of N,O; with the solute may be observed. In acid solution, N, 0' probably reacts via

N 2 0 ; + H , O + - + N 2 0 + H,O*



H. + H,O

(67) (68)

This results in a net lower reactivity of dinitrogen monoxide towards the hydrated electron in acid solution. Both I, and CS, react at diffusion controlled rates, the rate coefficients being 5.1 x 10' and 3.1 x 10' 1 mole-' sec-' , respectively [45, 491. The reaction of C ( N 0 , ) 4 with the hydrated electron has been used t o measure the extinction coefficient of the hydrated electron (see p. 438). Its rate coefficient [34] is 4.6 x 10' 1 mole-' sec-' .

2.3.4 Inorganic metal cations The hydrated electron reacts with most cations except the alkali and alkaline earth metal ions. In general, the higher the positive charge on the cation, the greater the reactivity. Inorganic cations are usually more reactive than inorganic anions. As with anions, there is no absolute rule governing reactivity of cations. The electron affinity of the metal, the nature of the ligand and the gain in free energy for the reaction

xn+ + e;q



R e f e r e n c e s PP 4 5 8 - 4 6 J



444 TABLE 3 Rate coefficients for reactions of M2+with

eiq [ 4 1 , 4 6 , 5 0 ]


k ( I mole-' sec-')


4 . 2 x 10" 3.8 x l o 7 3.5 x l o 8 1.2 x 10'O 2 . 2 x 10'O 3 x 1o'O 1 . 2 x 109

Mn" Fez+

co2+ Ni" cu2+ Zn"

are all important factors. The trend found in the rates of reaction of the hydrated electron with Sm3 +,Eu3 ' and Yb3 is the same as that of their electrode potentials but this correlation does not hold with other cations [451. The reactivity of the transition metal ions depends on the availability of an orbital and the gain in energy on addition of the electron. Rate coefficients for the reaction of some dispositive ions are shown in Table 3. The 3dsMn2 has electronic configuration t i g e l , i.e. an electron in each of the five 3d orbitals. Addition of a sixth electron would be relatively unfavourable and hence Mn2 is less reactive. The 3 d ' Zn2 has all five 3d orbitals filled and its lower reactivity is due t o the fact that addition of an extra electron must be t o a 4s orbital. The initial products of the reactions of hydrated electrons with bivalent transition metal ions are the univalent ions. +



elq + M2+





These have been identified by their transient ultraviolet absorption spectra [ 511 . They rapidly disproportionate 2 M + + M 0 +M2+


Copper is exceptional in that its monovalent cation is relatively stable. The production of unusual oxidation states in this way may lead t o the use of the hydrated electron in inorganic synthesis. The silver ion is reduced t o Ago, which has been detected spectrophotometrically [ 4 1 ] . Ag' + e;lq



k 7 is 3.2 x 10' O2 or Ag', viz. Ago + Ag+

Ago + O 2




1 mole-' sec-' . The Ago may subsequently react with Ag:


Ag'+ -0;


445 TABLE 4 Effect of ligand o n reactivity o f metal complexes with ea(l [41,50, 541


Cd" H2O NH3


5.2 x 3.1 x 3.9 x 1.4 x

10" 10''

lo8 10'


1.2 x G.5 x G1.8 x 1.8 x 1.6 x



lo6 10'


k , and k , , are 5.9 x lo9 and 3.8 x lo8 1 mole-' sec-' , respectively [52, 531. I t has been found that ligands affect the reactivity of the metal ion. Reactivity increases in parallel with the ability of the ligand t o act as a bridge for electron transfer. Some data for complexes of Cd(I1) and Zn(I1) are shown in Table 4. In general, the order of ability is OH- < CN- < NH3 < H2 0 < F- < C1- < I-. In some cases ethylenediaminetetraacetic acid decreases the rate of reaction of a metal cation with the hydrated electron relative t o that of the corresponding aquo complex by several orders of magnitude [ 541 . The ligands also affect the electron density around the cation by u and T bonding and may affect the configuration and stability of the lower oxidation state. Both these effects will be important in determining the reactivity of metal complexes towards e i s . Cr(III), Fe(lI1) and Co(1II) react at diffusion controlled rates which vary with the type of ligand. Some data are given in Table 5. TABLE 5 Rate coefficients for reaction of cobalt( 111) complexes with e i q Solute

k 8 . 2 x 10" 8.1 x 1 o ' O 4.5 x 1o'O 1.25 x 10"

5.8 x 10" 3.2 x 109 8.2 x 1 0 ' O References p p 1458-461


(1 mole-' sec-') 47 47 54 46 46 50 41

446 The lanthanides are all unreactive except for Eu3 +, Yb3 and Sm3 which have stable divalent ions. Stability problems have restricted studies of the actinides. However, the uranyl ion reacts with eaq with a rate coefficient of 7.4x 10' 1 mole-' sec-' [ 4 6 ] . +


2.3.5 Organic compounds The reactivity of organic compounds towards t h e hydrated electron depends on the availability of a low-lying vacant electron orbital. Thus saturated hydrocarbons and the corresponding amines and alcohols are unreactive. The reaction of the hydrated electron with monochloracetic acid to give chloride ion has already been discussed (see p. 430). This is only one of a general type of reaction involving haloaliphatic compounds [ 55, 561 . The reactions occurring are enq + RX RX-





R . + X-


Rate coefficients are given in Table 6 . The effect of substituent on the order of reactivity is F < C1 < Br, < I. Carbon tetrachloride and chloroform both react with the hydrated electron a t diffusion controlled limits, the rate coefficients being 3.0 x 10" and 2.5 x 10'' 1 mole-' sec- , respectively [ 491. The primary reaction of an alkene is to produce a carbanion


e,,, + RCHLCH2




The carbanion of acrylamide has been observed in the pulse radiolysis of acrylamide. I t may dimerize or react with scavengers [ 5 8 ] , viz. CH2 =CHCONH2 + eiq








(CH,CHCONH, ); +


CH, =CHCONH2 + -0;


(79 (80

The reaction of the carbanion with water has also been observed. RCH-CH,

+ H2O + RCH2CH2 + OH-


Ethylene itself is quite unreactive but the addition of electron withdrawing groups increases the reactivity markedly. Typical data are given in Table 7. Carbonyls are similar t o alkenes in that their reactivity is affected by the presence of adjoining groups. Typical data are given in Table 8.

447 TABLE 6 Rate coefficients for reactions of e& with haloaliphatic compounds [ 57 ] RX 1.2 x 109 8.5 x 109 2.0 x l o 6 6.2 x 109 1.15 X 10" 4.1 X l o 8 1.6 x 109 4.0 x 10' 2.7 x l o 9 6.6 x l o 9


TABLE 7 Rate coefficients for reactions of with elhylenic compounds ________



Ally1 alcohol Eth y I eiie Fumarate ion Maleate ion Acrylamide

<4 x l o s


( I mole-' sec-')

59 60

< 2 . 5 x 10'

7.5 x 10' 1 . 7 x 10' 1.8 x 10"

49 49 24 __

TABLE 8 \ Rate coefficients for reactions of e i y with ,C=O 62 1

[ 24, 61,


k (1 mole-' sec-') -

Formic acid Urea Acetone Tri fluoroacetone Acetic acid Acetaniide GI y c i n e Ethyl acetate Methyl trifluoroacetate

1 . 4 x 10' 3 x 105

References p p . 458-461

5 . 6 109 ~ 6.6 x l o 7 1.8 x l o 8 1.7 x 107 8.2 x l o 6 5.9 x l o 7 1.9 x l o 9

448 Electron withdrawing groups decrease the reactivity of aldehydes, ketones and carboxylic acids. This observation, at first, seems somewhat surprising but is due t o the fact that these groups cause the C=O bond t o shorten. The resulting increased T electron density makes the addition of another electron more unfavourable. The opposite effect occurs in amides and esters, where electron withdrawing groups increase the reactivity. Both -NH2 and -OR groups render the C=O bond unreactive and form new electrophilic centres for reaction with the hydrated electron. This has been attributed t o the mesomeric effect, the mesomeric form being

The presence of an electron withdrawing group R 1 enhances the electrophilic character of the new centres and thus increases their reactivity towards the hydrated electron. Aromatic compounds have been investigated in some detail. As with other reactions between hydrated electrons and organic substances, the primary product is not stable. These carbanions or radicals react t o gwe the final product, viz.

or PhX; + H 2 0 + .PhXH + OH-


Anbar and Hart [63] have correlated the reactivities of aromatic compounds towards the hydrated electron in terms of the Hammett u function and have shown that the rates are a function of the r electron density. The results are shown in Fig. 4. Hammett’s equation can be written as In this case 7) =

log [k(C,H,X + e,,)/k(C,H,

+ eiq)]

and represents the effect of a substituent on the electrophilicity of the compound. The u values are derived from aromatic substitution reactions. 2.3.6 Biological compounds

The reaction of the hydrated electron with biological molecules is of considerable importance to the radiobiologist for an understanding of the effect of radiation on living cells.


a /*4





-025 -

Carbohydrates, like alcohols, are generally unreactive [64] (12 < l o 5 1 mole-' sec-' ). Amino acids are relatively unreactive unless they contain a functional group which is reactive t o eaq [ 6 2 ] . Deamination occurs as a result of electron capture by aliphatic amino acids [65], viz.

+ e;lq -+ RCH(NH3 )CO$





R ~ H C O ;+ NH,

(85) (86)

However, amino acids containing the -SH group undergo preferential C-S bond fission. Proteins are complicated molecules having many groups present and their reactivity depends on the nature of such groups.

2.3.7 Reactions in concentrated solutions Until 1970, pulse radiolysis studies were limited t o those species with lifetime sec and observations on the hydrated electron were therefore carried out using only dilute solutions of scavenger. More recently, it has been possible t o develop pulses of = lo-' sec and consequently it has been possible t o study much earlier events in the radiolysis and also t o study the disappearance of the hydrated electron in concentrated scavenger solutions. Rate coefficients have been found t o depend on the concentration of scavenger [66-681. Thus in the competition of hydrogen ion and acetone for the hydrated electron, viz. e& + HZO


H. + H,O

eaq + CH,COCH, R e f e r e n c e s pp 4 5 8 - 4 6 1




it has been shown that the rate coefficient ratio, h , l k 6 increases from 0.48 in dilute solution (-0.01 M) to 2.4 in concentrated solution (-2 M). It has been proposed that this effect is due to the time dependence of rate coefficients [ 6 9 ] . In diffusion controlled reactions, time is required t o establish a diffusion gradient of each reactant around the other. Alternatively it has been suggested that in concentrated solutions, reactions of the “dry” electron, i.e. the electron before it has had time to be solvated in the medium, are being observed [ 7 0 ] . It is to be hoped that subsequent picosecond observations will increase our understanding of these systems.

3. The solvated electron in other systems 3.1 THE AMMONIATED ELECTRON

3.1.1 Preparation and properties

More than a century ago, Weyl [ 7 1 ] discovered that alkali metals dissolved in liquid ammonia t o give a stable, blue solution. I t was not until 1908, however, that Kraus [ 7 2 ] , as a result of electrolytic and conductimetric studies, suggested that the blue colour was due to the solvated electron. Subsequent experiments confirmed this [ 731 . One of the interesting properties of the solution of sodium metal in liquid ammonia is its high equivalent conductivity. I t is tempting t o suggest, as did early postulates, that this conductivity is due to the equilibrium Na + NH,$ + Na’ + e i H ,


This is an oversimplification since later work showed that, a t high concentrations of alkali metal, the equivalent conductivity increases with metal concentration [ 7 4 ] . Results are shown in Fig. 5. At high concen-

Fig. 5. Dependence of equivalent c o n d u c t a n c e o n concentration of s o d i u m f o r solutions of s o d i u m in liquid a m m o n i a at 239’K.

451 trations, the properties approach those of a liquid metal. Below a concentration of 0.05M, there is a decrease in conductivity with increasing concentration of metal. This and the fact that the molar paramagnetic susceptibility of the solution also decreases, indicates that ion pairing, or assembly into larger aggregates, takes place. Of the many models proposed for the diamagnetic species [75-771, the solvated electron pair model [78] seems t o be the most satisfactory, viz. eNH,





The complexity of metal-liquid ammonia solutions indicates that care must be taken in interpreting reactions in this system. Thus it is important t o be able t o distinguish between reactions of the single reducing species and those of aggregates. Several models have been proposed for the structure of the ammoniated electron. One which has shown qualitative agreement with experimental I

L i 800



1200 1600

A nm

Fig. 6. Absorption spectrum of the ammoniated electron.

results is that of Copeland et al. [ 791, who consider the electron t o be trapped in a large cavity in the solvent, stabilized by orientation of the dipoles of the ammonia molecules on the cavity surface, with a polarizable continuum beyond. Results have been fully discussed in a recent comprehensive review [ 8 0 ] . All dilute solutions of alkali and alkaline earth metals in ammonia give a single broad peak in the absorption spectrum in the near infrared region. This absorption is independent of the metal used, indicating that the absorbing species is the same in all solutions. The spectrum is shown in Fig. 6. The peak is asymmetrically broadened on the high energy side, rising to a maximum at = 1500 nm. Its exact location is slightly temperature dependent. Aggregation has little effect on the optical absorption which obeys Beer’s Law up t o a concentration of 0.03M.In liquid deutero-ammonia, lithium, sodium, potassium, calcium and barium R e f e r e n c e s p p 458- 461

452 [ 811 . X, a x is shifted t o all give the same characteristic spectrum of e i 1380 nm. The same typical absorption spectrum has been observed in both the continuous [ 821 and pulse [83] radiolysis of liquid ammonia, viz.





I t is interesting that the greater stability of the ammoniated electron makes its observation easier in continuous radiolysis than that of the hydrated electron. The ESR spectrum of dilute solutions shows a single, narrow line. The g factor is 2.0012 which is close t o the free electron value and indicates only a weak interaction between electron and solvent ~ 4 1 . Jortner [85] obtained the heat of solution of the electron in liquid ammonia using the following Born-Haber cycle. Na,



AH/kcal mole-' 26.1





-40.2 kcal mole-'


26.1 + 118.6 - 103.1 - 1.4

The value obtained is independent of the cation. Thermodynamic functions for the electron in liquid ammonia (relative t o H') have been calculated by Jolly [ 8 6 ] . The free energy and heat of formation at 298°K are 46.0 and 37.5 kcal mole-', respectively. The entropy of formation is -13.0 cal OK-' mole-' . I t is possible t o calculate that the equilibrium constant for the reaction

is - l o 6 . That the equilibrium does, in fact, exist is confirmed by the observation that a 1 M solution of potassium amide in equilibrium with 1 atm of hydrogen contains - 1 0 - 6 M ammoniated electron, as detected by ESR [ 8 7 ] .

453 3.1.2 Reactions of the ammoniated electron Solutions of sodium in liquid ammonia are stable for several months but rubidium and caesium slowly react with the solvent t o give the amide and hydrogen, e.g. Rb + NH3


RbNH2 +

4 H2


Although solutions of sodium in liquid ammonia have been used extensively as a reducing agent in organic chemistry [88, 891, the results obtained have, in the main, been qualitative. In view of the relative stability of such solutions, it is a little surprising that more quantitative data are not available. The scarcity of data is probably due t o the difficulties of interpretation since, as has been discussed earlier, species other than the simple ammoniated electron are produced. Reactions in liquid ammonia can be conveniently classified as follows. (a) Electron addition without bond cleavage, e.g. ej&, + x




ESR studies [go] have shown that carboxylic acids, ring- or N-substituted benzamides, thiobenzamides and nicotinamides can be reduced t o give the corresponding radical anions or dianions, e.g. RCO, + e",,




Other reactions that have been observed include eN:{, + O2 -+ - 0 , e i H , + MnO,



(b) Bond cleavage by the addition of one or more electrons, e.g.

Dissociative electron capture has been observed in the pulse radiolysis of solutions of methyl chloride, benzyl chloride and iodobenzene in liquid ammonia [ 91 ] , e.g.

Other reactions that have been observed are

eiH, + (CH,),S References p p . 4 5 8 -461


CH3S- + $ C2H6


454 The radical produced may react with ammonia or with another ammoniated electron. In the latter case, the equation for the overall decomposition is

RX + 2 eiH,


R- + X-


Reductions of this type have been observed using Ge2H6 and N 2 0 in liquid ammonia, viz.

02+ NH,


NH; + OH-


It is interesting that although in the radiolysis of its aqueous solutions chloracetic acid undergoes one electron reduction (see p. 430), in its reduction with sodium amalgam it undergoes a similar two electron reduction. Reduction in metal-ammonia solutions often depends on the availability of proton donors. Thus benzene is not reduced unless an alcohol is present when 1,4-dihydrobenzene is formed. The following reaction scheme has been proposed [ 921



CNH~ = x-


/,+I I










H2 x

The reaction of ethanol with sodium in liquid ammonia is one of the few reactions for which quantitative data are available [ 9 3 ] . The rate of hydrogen evolution provides a convenient measure of the rate of the reaction. The following reaction mechanism was proposed [ 861 kl13


C2HsOH + N H 3

NHd + eNH3


C2H50 + NHd


NH, +

f H2



Values of k , I = 8 x 1 mole-' sec-' and k 3/kl,4 = 6 x were obtained. From estimates of h 3 , it was calculated that k ,



455 A detailed study of the reaction of water with solutions of the alkali metals in ammonia has been made [ 9 4 ] . Interpretation of the data is complicated by the large number of possible species present but it appears that the reaction of the ammoniated electron with water

is slow. As discussed earlier, reaction of the hydrated electron is similarly slow. Although undoubtedly the hydrated electron reacts with water t o give the hydrogen atom and the hydroxyl ion, thermodynamic arguments have been advanced [95] against the formation of these species for the reaction of the ammoniated electron with water and reaction (115) should therefore be regarded as speculative. in the pulse radiolysis of liquid ammonia is second The decay of e i order with a rate coefficient [96] of (1.1 0.2) x 10' 1 mole-' sec-' . By contrast, the disappearance of e i in alkali metal--ammonia solutions at concentrations < 1 0 - 3 M is first order and very slow. The ammoniated electron formed in the radiolysis probably reacts with other intermediates produced, e.g. *NH, or NH:, and does not decay via the reaction


eNH, + e i H ,


H 2 + 2 NH,


The rate coefficient of the reaction

at 238°K was found [97] to be 1.2 x lo6 1 mole-' sec-' . This value is consistent with the upper limit of 10' 1 mole-' sec-' deduced from the sodium- -liquid ammonia system. 3 . 2 THE SOLVATED ELECTRON IN AMINES

Alkali metakamine solutions are more powerful reducing agents than their ammonia counterparts. They are, however, much less stable and decomposition t o amide and hydrogen occurs quite readily. The absorption spectra consist of one or more peaks at 650, 850 and 1300 nm depending on the metal and the amine present. The major absorption in solutions of sodium in ethylene diamine at 650 nm is due t o the species Na-, the e;(,lv absorption being small, even in dilute solutions 1981. Using pulse radiolysis, Dye et al. [99] have studied the formation of Na- from e[(,lv and sodium in ethylenediamine. The reaction of caesium with water in, ethylenediamine again shows that the solvated electron reaction with water is slow [95, 1001. Kefercnces p p 4 5 8

4h I


A powerful aprotic solvent capable of dissolving alkali and alkaline earth metals is hexamethylphosphoramide. The characteristic blue paramagnetic solution is observed [ l o l l . It had been demonstrated that the ammoniated electron could be produced in the electrolysis of liquid ammonia [lo21 and recently it has been demonstrated that the solvated electron is produced in the electrolysis of hexamethylphosphoramide [103, 1041. It is, however, doubtful whether eaq can be produced in this manner [102]. 3.3 OTHER SYSTEMS

The solvated electron has been identified by its broad absorption spectrum in the pulse radiolysis of less polar solvents, e.g. alcohols [ 1051 , Typical spectra are shown in Fig. 7. Rate coefficients for reactions of some solutes with e, ,,, in liquid methanol [ 106-1081 and ethanol [ 106, 109,1101 have been measured and are summarized in Table 9.








I nm

Fig. 7. Absorption spectra of e,,l,, in alcohols.

Relative rate coefficients have been measured in 2-propanol [ 112, 1131. I t was originally thought that rate coefficients for the solvated electron in ethanol and the hydrated electron were similar but Rabani et al. [ill] have shown that for benzene and phenol the rate coefficients differ significantly. The number of solutes investigated is, as yet, very limited by comparison with the aqueous system and more data are required before the reasons for these differences can be established. The solvated electron was observed to be an intermediate in the reaction of sodium with solid alcohols at 77°K by optical and ESR techniques [ 1141. It is now known to be generated in the sodium-liquid alcohol system [115]. The formation of the solvated electron was demonstrated by using dinitrogen monoxide as the electron scavenger, when nitrogen was formed. The hydrogen is evolved via an intermediate (e&,lv)2as might be the case in water.

457 TABLE 9 Rate coefficients for reactions of e,,lv in methanol and ethanol [105,107,109-1111 Solute

k ( I mole-' sec-l) CH30H


5.2 x 10"


1.9 x


0.50x 10"


1.9 x 1o'O 0.43 x 10" 0.51 x 10" 3.5 x 1 o ' O 1.1 x 1 o 1 O 3.0 x 10" 4.0 x 10" 2.5 x l o 9 0.42x l o 7 4.5 x l o 7

The use of pulse radiolysis and scavenging techniques has identified the electron in liquid hydrocarbons [116]. In these non-polar media, interaction between the electron and the solvent is, of course, very small and it is doubtful whether such electrons can properly be described as solvated. The pulse radiolysis of hexane or 3-methyl hexane gave rise t o a short lived transient species absorbing at A, a x = 1500 nm. Rate coefficients for reaction in hydrocarbon solutions at 193°K are given in' Table 10. Baxendale et al. [117] have studied the pulse radiolysis of liquid methylcyclohexane and obtained rate coefficients for the reactions of the electron with carbon tetrachloride and pyrene at 293" K. Relative rate coefficients have been obtained by irradiation of hydrocarbons containing two electron scavengers [118]. A broad absorption band is developed in irradiated polar glasses and is attributed t o the trapped electron [119]. The distinction between trapped and solvated electrons has been recently reviewed by Dainton [ 1201. TABLE 10 Rate coefficients f o r electron reactions in hydrocarbon solutions Solute

k ( I mole-' sec-I)

0 2

1.5 x 1.1 x 1.9 x 2.3 x

N2O (C6H5 )2


References p p . 4 5 8 --461

10" 10" 10" 10"


Addition of naphthalene or biphenyl to the glass causes the broad band t o be replaced by the characteristic absorption spectrum of the naphthalene or biphenyl radical anion due t o e;m,






The addition of sodium t o aromatic hydrocarbons produces only the corresponding radical anion [ 1211 . Alkali metals, however, dissolve in aliphatic ethers, e.g. tetrahydrofuran, dioxan, producing the characteristic blue colour [122]. The solutions are diamagnetic because of the formation of higher ion pairs. Flash photolysis of solutions of sodium in ethers forms the ion pair consisting of the solvated electron and a sodium cation [123]. Three transients are formed in the flash photolysis of sodium pyrenide in tetrahydrofuran [ 1241. These have been identified as the solvated electron, e , , the ion pair, e, Na', and the sodium atom, Nao. Rate coefficients of reactions of e;,,lv with various compounds relative to the rate of reaction with N,O have been determined recently by the y-radiolysis of 2-methyl-tetrahydrofuran [ 1251 . The formation of es-oIv has also been substantiated in the pulse radiolysis of dimethyl sulphoxide [126, 1271. Conclusion It is clear from this survey that the existence of solvated electrons has now been demonstrated in a wide variety of solvents. Although, at present, detailed rate data are only available for the aqueous system, it is to be expected that in the next few years, more data on other systems will be accumulated. This should prove invaluable in aiding our understanding of the reactions of this extremely simple but powerful reducing agent.

REFERENCES 1 J. Weiss, N a l t m (I,oridori), 1 5 3 (1944) 748. 2 J. L. Magee, Basic rticcltanisrns in radiobiology. physical and chemical effects, Natl. Res. Council Publ. 305, Washington, 1953, p . 51. 3 G . Stein, Discuss. Faraday Soc.. It' (1952) 227. 4 R. L. Platzman, Basic mechurlisms in radiobiology. physical atid chemical e f f e c t s , Natl. Rcs. Council Publ. 305, Washington, 1953, p. 22. 5 N. F. Barr a n d A. 0. Allen, J . Phys. C'hern.. 69 (1959) 928. 6 J. H. Baxendale a n d G . Hughes, 2. Phys. C h c m . (Frarihfurt am M a i n ) 13 (1958) 306. 7 E . H a y o n a n d J . Wciss, Proc. I n t . Corif. Peaceful I'ses o f A t o m i c Energy, 29 (1958) 80. 8 G . Czapski a n d H . A . Schwarz, J. P h y s . Chem.. (i6 (1962) 471.

459 9 E. Collinson, F. S. Dainton, D. R. Smith and S. TazukB, Proc. Chem. SOC. L o n d o n , (1962) 140. 10 J. W. Boag and E. J. Hart, J. A m e r . Chem. Soc., 84 (1962) 4090; N a t u r e ( L o n d o n ) , 1 9 7 (1963) 45. 11 J. P. Keene, N a t u r e ( L o n d o n ) , 1 9 7 (1963) 47. 12 M. S. Matheson, W. A. Mulac and J. Rabani, J. Phys. Chem., 6 7 (1963) 2613. 13 J. W. Boyle, J. A. Ghormley, C. J. Hochanadel and J . F. Riley, J. Phys. Chem.. 73 (1969) 2886. 14 L. I. Grossweiner and H. I. Joschek, A d u a n . C h e m . Ser., 5 0 (1965) 279. 15 J. Jortner and J. Rabani, J . A m e r . Chem. Soc., 8 3 (1961) 4868. 1 6 M. Anbar and E. J. Hart, J . Phys. Chem., 71 (1967) 3700. 17 M. S. Matheson and J. Rabani, J. Phys. Chem., 69 (1965) 1324. 1 8 M. Anbar and E. J . Hart, J. Phys. Chem., 71 (1967) 4163. 19 J. Jortner and G. Stein, N a t u r e ( L o n d o n ) , 175 (1955) 893. 20 J. Bennett, B. Mile and A. Thomas, N a t u r e ( L o n d o n ) ,201 (1964) 919. 21 G. Hughes and R. J . Roach, Chem. Commun., 1965, 200. 22 D. C. Walker, Can. J. Chem., 44 (1966) 2226. Spec. Publ., 2 2 (1967) 277. 23 E. A. Shaedeand D. C. Walker, Chem. SOC. 24 S. Gordon, E. J. Hart, M. S. Matheson, J. Rabani and J. K. Thomas, Discuss. Faraday Soc., 36 (1963) 193. 25 G . Hughes and C. R. Lobb, Unpublished results. 26 A. J. Swallow, Photochem. P h o t o b i o L , 7 (1968) 683. 27 J . Rabani, J. P h y s . Chem., 66 (1962) 361. 28 D. C. Walker, Can. J. C h e m . , 45 (1967) 807. 29 B. E. Conway and D. J. Mackinnon, J. Phys. Chem., 74 (1970) 3663. 30 J. H. Baxendale, Radiat. Res. Suppl., 4 (1964) 139. 31 J. Jortner and R. M. Noyes, J. Phys. Chem., 70 (1966) 770. 32 K. H. Schmidt and S. M. Ander, J. Phys. Chem., 7 3 (1969) 2846. 33 J. Jortner, Radiat. Res. Suppl., 4 (1964) 24. 34 J. Rabani, W. A. Mulac and M. S. Matheson, J. P h y s . Chem., 6 9 (1965) 53. 35 E. J. Hart, in M. Haissinsky (Ed.), A c t i o n s Chimiques et Biologiqucs des Radiations, Masson, Paris, Dixieme Serie, 1966, p. 1. 36 M. Anbar and P. Neta, Int. J. A p p l . Radiat. fsotop., 16 (1965) 227. 37 E. J. Hart and M. Anbar, The hydrated electron, Wiley-Interscience, New York, 1970. 38 E. J. Hart, S. Gordon and E. M. Fielden, J. Phys. Chem., 70 (1966) 150. 39 L. M. Dorfman and I. A. Taub, J. A m e r . Chem. Soc., 85 (1963) 2370. 40 J . P. Keene, Radial. R e s . , 22 (1964) 1. 4 1 J . H. Baxendale, E. M. Fielden, C. Capellos, J. M. Francis, J . V. Davies, M. Ebert, 42 43 44 45 46

C. W. Gilbert, J. P. Keene, E. J. Land, A. J . Swallow and J. M. Nosworthy, Nature ( L o n d o n ) , 201 (1964) 468. M. Anbar and D. Meyerstein, J. P h y s . Chem., 6 9 (1965) 698. 0. MiCiC and I. Draganic, l n t . J. Radiat. Phys. C h e m . , 1 (1969) 287. J . H. Baxendale, Radiat. Res. Suppl., 4 (1964) 114. J. K. Thomas, S. Gordon and E. J . Hart, J. Phys. Chem., 68 (1964) 1524. J . H. Baxendale, E. M. Fielden and J. P. Keene, Proc. R o y . Soc., Ser. A , 286

( 1965) 320. 4 7 M. Anbar and E. J. Hart, A d u a n . Chem. Ser., 81 (1968) 79.

48 49 50 51

M. Anbar and P. Neta, l n t . J. A p p l . Radiat. Isotop., 16 (1965) 227. E. J. Hart, S. Gordon and J. K. Thomas, J . Phys. Chem., 68 (1964) 1271. M. Anbar and E. J . Hart, J . Phys. C h e m . , 69 (1965) 973. G. E. Adams, J. H. Baxendale and J. W. Boag, Proc. Chem. Soc. L o n d o n , (1963)

241. 52 J. Pukies, W. Roebke and A. Henglein, Ber. Bunsenyes. Phys. C h e m . , 7 2 (1968) 842.

460 53 J. H. Baxendale, E. M. Fielden a n d J. P. Keene, Pulse radiolysis, Academic Press, London and New York, 1965, p. 207. 54 M. Anbar a n d D. Meyerstein, Trans. Faraday SOC.,65 (1969) 1812. 55 E. Hayon and A. 0. Allen, J. Phys. C h e m . . 65 (1961) 2181. 56 M. Anbar, Aduan. C h e m . Ser., 5 0 (1965) 55. 57 M. Anbar and E. J. Hart, J. Phys. C h e m . , 69 (1965) 271. 58 K. W. Chambers, E. Collinson, F. S. Dainton, W. A. Seddon and F. Wilkinson, Trans, Faraday SOC..6 3 (1967) 1699. 59 G. Scholes and M. Simic, J. Phys. C h e m . , 68 (1964) 1731. 60 C. F. Cullis, J. M. Francis and A. J. Swallow, Proc. R o y . SOC.,Ser. A , 28 7 (1965) 15. 6 1 E. J. Hart, E. M. Fielden and M. Anbar, J. Phys. C h e m . , 71 (1967) 3993. 62 R. Braams, Radiat. R e s . , 27 (1966) 319. 63 M. Anbar and E. J. Hart, J. A m e r . C h e m . Soc., 8 6 (1964) 5633. 64 J. V. Davies, W. Griffiths and G. 0. Phillips, Pulse radiolysis, Academic Press, London and New York, 1965, p. 165. 65 R. L. S. Willix and W. M. Garrison, Radiat. Res., 32 (1967) 452. 66 E. Peled and G. Czapski, J . Phys. Chem.. 75 (1971) 2626. 67 R. K. Wolff, M. J. Bronskill and J. W. Hunt, J. C h e m . P h y s . , 5 3 (1970) 4211. 68 J. E. Aldrich, M. J. Bronskill, R. K. Wolff and J. W. Hunt, J. C h e m . P h y s . , 5 5 (1971) 530. 69 H. A. Schwarz, J. C h e m . P h y s . , 55 (1971) 3647. 7 0 W. H. Hamill, J. Phys. Cliem., 7 3 (1969) 1341. 7 1 W. Weyl, A n n . Physilz., 197 (1863) 601. 72 C. A. Kraus, J. A m e r . C h e m . Soc., 30 (1908) 1323. 73 C. A . Kraus, E. S. Carney and W. C. Johnson, J. A m e r . C h e m . SOC.,49 (1927) 2206. 74 D. S. Berns, Aduan. Chem. S e r . , 5 0 (1965) 82. 75 E. Becker, R. H. Lindquist and B. J. Alder, J. Chem. P h y s . , 25 (1956) 971. 76 M . Gold, W. Jolly and K. S. Pitzer, J . A m e r . Chem. Soc., 8 4 (1962) 2264. 77 S. Golden, C. Guttman and T. R. Tuttle, J. A m e r . Chem. Soc., 8 7 (1965) 135. 78 R. Catterall and M . C. R. Symons, J. Chem. Soc., A , (1966) 13. 79 D. A. Copeland, N. R. Kestner and J. Jortner, J . Chem. Phys., 5 3 (1970) 1189. 80 J. Jortner and N. R. Kestner (Eds.), Electrons in fluids, Springer Verlag, Berlin, 1973, p. 493. 81 D. F. Burow and J . J. Lagowski, Aduan. C h e m . S e r . , 5 0 (1965) 125. 82 R. W. Ahrens, B. Suryanarayana and J. E. Willard, J. Phys. C h e m . , 68 (1964) 2947. 83 D. M . J . Compton, Pulse radiolysis. Academic Press, London and New York, 1965, p. 13. 84 C. A. Hutchison and R. C. Pastor,J. C h e m . P h y s . , 21 (1953) 1959. 85 J. Jortner, J . C h e m . P h y s . , 30 ( 1 9 5 9 ) 839. 86 W. L. Jolly, Aduan C h e m . Ser., 5 0 (1965) 27. 87 E. J. Kirsche and W. L. Jolly, Science, 147 (1965) 45. 88 H. Smith, Organic reactions in liquid ammonia, Interscience, New York and London, 1963. 89 A. J. Birch, Quart. R e v . . 4 (1950) 69. 90 A. R. Buick, T. J. Kemp, G. T. Neal and T. J. Stone, J. C h e m . SOC., A , (1970) 2227. A , ( 1 9 6 9 ) 666. 9 1 A. R. Buick, T. J. Kemp, G. T. Neal and T. J . Stone, J. Chem. SOC., 9 2 W. L. Jolly and C. J. Hallada, Non aqueous solvent systems, Academic Press, London and New York, 1965, p. 34. 93 E. J. Kelly, H. V. Secor. C. W. Keenan and J . F. Eastham, J . A m e r . C h e m . SOC., 8 4 (1962) 3611.

461 94 R. R. Dewald and R. V. Tsina, C h e m . C o m m u n . , 1967, 647. 9 5 J. L. Dye, Accounts Chem. Res., 1 (1968) 306. 96 V. N. Shubin, V. A. Zhigunov, G. I. Khaikin, L. P. Berushashvil and P. I. Dolin, Aduan. C h e m . Ser., 81 (1968) 95. 97 J. M. Brooks and R. R. Dewald, J. Phys. C h e m . , 75 (1971) 986. 98 J. L. Dye and M. G. Debacker, J. Phys. C h e m . , 75 (1971) 3092. 99 J. L. Dye, M. G. Debacker, J. A. Eyre and L. M. Dorfman, J. Phys. C h e m . , 7 6 (1972) 839. 100 R. R. Dewald, J. L. Dye, M. Eigen and L. De Maeyer, J. C h e m . P h y s . , 39 (1963) 2388. 101 G. Fraenkel, H. S. Ellis and D. T. Dix, J . A m c r . C h e m . Soc., 8 7 (1965) 1406. 102 D. Post1 and U. Schindewolf, Ber. Bunsenges. Phys. C h e m . , 75 (1971) 662. 103 Y. Kanzaki and S. Acyagui, J. Electroanal. C h e m . Interfacial Electrochem., 36 (1972) 297. 104 H. W. Sternberg, R. E. Markby, I. Wender and D. M. Mohilner, J. A m e r . Chem. S O C . ,91 (1969) 4191. 105 L. M. Dorfman, Aduan. C h e m . S e r . , 5 0 (1965) 36. 106 I. A. Taub, D. A. Harter, M. C. Sauer and L. M. Dorfman, J. Chem. P h y s . . 4 1 (1964) 979. 107 J. H. Baxendale, Int. J. Radiat. Phys. C h e m . , 4 (1972) 113. 108 W. V. Sherman, J. Phys. C h e m . , 71 (1967)4245. 109 J . W. Fletcher, P. J. Richards and W. A. Seddon, Can. J . C h e m . , 48 (1970) 1645. 110 S. M. S. Akhtar and G. R. Freeman, J. P h y s . C h e m . , 75 (1971) 2756. 111 J. Rabani, H. B. Steen, H. Bugge and T. Brustad, C h e m . Cornmun., 1971, 1353. 112 W. V. Sherman, J. P h y s . C h e m . , 70 (1966) 667. 113 W. V. Sherman, J. Phys. C h e m . . 70 (1966) 2872. 1 1 4 J. E. Bennett, B. Mile and A. Thomas, J . C h e m . Soc., A , (1967) 1393. 115 J. W. Fletcher and P. J. Richards, Can. J . Chem., 49 (1971) 2275. 116 J. T. Richards and J. K. Thomas, C h e m . P h y s . L e t t . , 1 0 (1971) 317. 117 J . H. Baxendale, C. Bell and P. Wardman, Clhern. P h y s . L e l l . , 12 (1971) 347. 118 R. H. Schuler and P. P. Infelta, J. P h y s . C'hcm., 76 (1972) 987. 119 W. H . Hamill, J. A m e r . C h c m . S o c . , 8 4 (1962) 500. 120 F. S. Dainton, in J. C. Polanyi ( E d . ) , hl.T.P. int. rcu. s c i . . p h y s . c h e m . ser. I , Kinetics, Vol. 9, 1972, p. 271. 121 M. Szwarc, Carbanioris, living p o l y m e r s arid electron lransfcr processes, Interscience, New York, 1968, p. 297. 122 F. S. Dainton, D. M. Wiles and A . N. Wright, ,I. C h e m . S o c . , (1959) 3767. 123 J. G. Kloosterboer, L. J. Giling, R. P. H. Rettschnick and J. D. W. Van Voorst, C h e m . P h y s . L e t t . , 8 (1971) 462. 124 M. Szwarc, Proc. R o y . Soc..Scr. A , 3 2 7 (1972) 481. 125 J. Teply and I. Janovsky, Chcm. P h y s . L e t t . , 17 (1972) 373. 126 D. C. Walker, N. V. Klassen and H. A. Gillis, C h c m . P h y s . L e t t . , 10 (1971) 636. 127 R. Bensasson and E. J. Land, Chern. P h y s . Lctt., l C 5(1972) 195.