On the Reactivity of Hydrated Electrons Toward Inorganic Compounds

6 On the Reactivity of Hydrated Electrons

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Toward Inorganic Compounds M I C H A E L A N B A R and E D W I N J. H A R T 1

Chemistry Division, Argonne National Laboratory, Argonne, Ill. 60439

The rates of reaction of a number of inorganic complex ions and oxy-anions with hydrated electrons have been measured and the parameters which determine the rate of their dif fusion-controlled reactions have been evaluated and dis cussed. Ligands modify the reactivity of ions by their effect on the configuration and stability of the lower state of oxi dation, on the electronic distribution of the central atom, and on their capacity of acting as electron transfer bridges. Each of these factors contributes to the reactivity of inor ganic complexes. The reaction cross section of certain highly reactive oxidants is significantly larger than their crystallo graphic dimensions, suggesting long range electron tunnel ing. General trends in the reactivity of the hydrated electron with inorganic compounds are discussed.

Qinee the discovery of its absorption spectrum, the rate of reaction of ^ the hydrated electron with inorganic cations, anions, and complex ions has been extensively studied (2, 7, 12, 13, 14, 18, 23, 29, 36, 39) and several attempts have been made to deduce some generalizations on the mechanism of e~ reactions with inorganic compounds. Substantial differences in reactivity exist, however, between "similar" inorganic compounds, and the interpretation of their kinetic behavior is therefore much less successful than that for organic systems (1,5, 6). Our understanding of the reactivity of inorganic compounds is limited at best, to the qualitative level. In some early studies salt effects were not critically considered; consequently, some erroneous conclusions were drawn. We further evaluate these data and provide new rate constant data on some inorganic ions not heretofore studied. m

' Weizmann Institute of Science, Rehovoth, Israel. 79 In Radiation Chemistry; Hart, Edwin J.; Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

80

RADIATION CHEMISTRY

Table I.

The Reactivity of Hydrated k .


Compound 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 1.10 1.11 α b c d

plier

[Co(NH ) ](C10 ) [Co(NH ) H 0](C10 ) [Co(NH ) CN](C10 ) [Co(NH ) Cl](C10 ) [Co(NH,) Br](C10 ), [Co(NH ) N ]ClO ) [Co(NH ) (CN)H 0](C10 ) [Co(en) C0 ]C10 [en Co0 (NH )Coen ]Br [(NH ) Co0 Co(NH ) ]Br [(CN) CoO,Co(CN) ]K 3

e

3

5

3

5

3

5

4

4

4

4

5

3

3

5

2

2

2

2

3

2

3

4

5

2

2

6.7 4.9 6.1 7.3 7.7 6.3 6.1 7.2 6.2 5.9 7.0

3

4

0

3

pH

3

2

4

3

4

2

2

2

3

n

4

5

5

1

r >

5

k^corr fi

AOS

μM Χ 10

X 10~ M

5

10

8.15 8.1 6.3 6.1 6.2 6.3 5.6 4.9 9.6 8.2 2.9

6 6 4 4 4 4 4 2 10 5 10

1

seer

1

8.2 8.2 6.3 6.1 6.2 6.3 5.6 4.9 9.7 8.3 2.8

Obtained from the group of H. Taube, Stanford University. Calculated from the ionic mobility of Co(NH.,) = 1.02 X 10~ cm.Vvolts sec. (16). Estimated. Estimated from Co-NH bond length (37, 38) in Co (III) amino complexes. 3+

3

3

Many of the reactions of inorganic compounds with hydrated elec trons are extremely rapid, their reactions being essentially diffusion controlled. A detailed analysis of these reactions reveals new information on the mode of interaction of hydrated electrons with reactive inorganic solutes. Experimental The preparation of solutions, irradiation of the samples, and analysis of decay curves by "Chloe" follow our previously described techniques (24). Hydrated electron scavengers were removed from the hydrogensaturated matrix by its pre-irradiation before injection of the solute i n cases where concentrations of the order of ΙΟμΜ solute were tested. The sources of supply of the chemicals used appear in Tables I and II. Results A number of inorganic compounds have been investigated for their rates of reaction with hydrated electrons. For purposes of later discus sions, we divide them into two groups : compounds which approach the diffusion controlled rate and compounds which react much slower. The data of the first group are presented in Table I which includes a number of cobalt complexes, and in Table II which contains the remaining fast reacting compounds. Here they are presented in the order of increasing atomic number of their central atom. In addition to the experimental conditions and the observed specific rate constants ( k ), these tables ohs

In Radiation Chemistry; Hart, Edwin J.; Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

6.

ANBAR AND HART

81

Reactivity of Hydrated Electrons

Electrons with Cobaltic Complexes Oxiœ cm. sec. 2

m

8

0.9 0.8 0.7 0.7 0.7 0.7 0.7 0.7" 0.6 0.6 0.6° C

C


C

C

e

C

C

e

C

e

9

a

=

C

C

i

1

_ corr kdiff :

9.2 9.0 6.3 6.2 6.2 6.3 6.2 4.4 11.6 14.3 0.12

r

e

T

1

2.6 2.6" 2.8 2.6* 2.7 2.8 2.7 3.7 4.5 4.0 5.0

h

e

k xio->° M' sec.'

r Χ 10 cm.

1

(

0.89 0.91 1.00 0.98 1.00 1.00 0.91 1.11 0.84 0.58 23.3

Estimated from (37, 38) for [Co(NH.,) Cl H 0] . Estimated from [Pt(NH ) Br ] - and [Co(en) Br ] . Obtained from J. A. Weil, Argonne National Laboratory. 3

3

2

4

2

+

2

2

2

2

+

contain the specific rate constant corrected for the salt effect according to (28) log

1.02

(k ), corr

Ζ — ^

Ζ is the charge of the substrate ion, μ is the ionic strength of the solution, μ = 1/2 2iZi Ci, and a = r/3, where τ = the sum of radii of the two reactants i n A . In these tables, fc is the diffusion limited bimolecular specific rate constant calculated from the Debye equation (20) from the sums of the diffusion coefficients ( D - + D ) and radii of e" and the substrate ion (r - -\- r ) (22). 2

(liff

e

e

x

aq

x

= 4 * N ( D - + D ) ( ν + r )Q/(e*

k

e

mt

x

x

- 1)

where Q = —Z e /ekT(r - f r ); e = the macroscopic dielectric con stant of water; 78.6 at 25 ° C ; Ν = Avogadro's number, 6.025 X 10 3 molecules/mole; k = Bolzmanns Constant, 1.38 Χ 10" erg/degree; Z = charge on substrate ion; D - = 4.7 X 10" sq. cm. sec." (33). The diffusion coefficients of the substrate ions were taken or estimated from the values of published ionic mobilities (16, 25). W e used 2.5 A . as the radius of e~ ; that of the substrate ion was calculated or estimated from known bond lengths ( 37, 38 ). x

2

e

x

2:

16

x

e

r>

1

m

Specific rate constants for less reactive compounds are presented i n Table III i n the order of increasing atomic number of the central atom. The implications of these results are discussed below.


82

RADIATION

Table II.

CHEMISTRY

Reactions of Hydrated Electrons ôbs


Compound 2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8 2.9 2.10 2.11 2.12 2.13 2.14

2.15 2.16 2.17 2.18 2.19 2.20 2.21 2.22 2.23 a 6 c d e T 9 h

Supplier

pH

_

10.2 6.6 11.0 6.7 10.0 10.5 6.6 11.0 11.0 8.5 11.0 10.6 9.0 6.5 11.0 10.9 11.0 10.6 10.0 8.5 10.6 10.2 10.6

NaOCl Na TiF NH V0 [Cr(NH ) Cl]Cl, K Cr(CN) Na [Fe(CN),NO] K FeF Ni(en) S0 K Ni(CN) KNiF(H 0), Na Se0 K Pd(CN) K Mn(CN) 2

a

6

4

d

3

s

0

5

4

i

6

J

2

3

i i I

e

2

4

2

4

5

2

2

d

4

2

i

4

4

i

6

η

K SnF„ 2

KSb0 Na TeO NaoTeU K Pt(CN) K PtCl T1 S0 KAu(CN) KJrCl KÎrCl

Ρ

3

2

a

H

a

4

2

2

1

4

1

6

2

)

4

i

2

4 t

6

6

1

X 10 M 5

10 10 10 10 X 10 10 10 100 Χ 10 10 10 10 Χ 10 10 10 10 10 10 10 10 10 X 10» Χ 10 2

4

5

s

4

5

2

4

2

3

5

3

3

2

2

2

3

3

2

3

5 5

3

^corr

X 10' M 9

J

sec." 7.0 3.5 4.9 62. 3.3 23.5 2.2 8.0 4.1 12.0 1.0 1.9 5.9 2.9 12. 1.0 15. 2.9 1.4 40. 3.5 9.3 3.0

7.2 5.8 4.9 62. 14. 23.5 10.5 7.5 5.5 7.2 1.1 2.8 25. 4.1 13. 1.1 16. 3.9 2.0 37. 4.2 20. 9.4

BDH Lab Reagent, purified. Estimated by analogy with SO/" and Fe(CN)e ". Calculated from experimental data ( 37, 38 ). BDH Analar. Estimated by analogy with C10 " and BrOa". Estimate based on planar configuration. Obtained from H. Taube's group, Stanford University. Estimated from (37, 38) Cr(CNS)4(NH ) 3

3

3

2

The effect of ionic environment on the rate of diffusion controlled e~ reactions is revealed when one compares the experimental rate con stants with the calculated values. In Table I the highly charged bispentacyano cobaltic peroxide (I) is much more reactive than expected for a pentavalent anion (1.11). It has been claimed (19) that polyvalent anions exhibit a lower effective charge in their kinetic behavior than expected from their structural formulae. W e have checked the salt effect on the reaction of I + e~ and compared it with the N C V + e~ reaction. The results presented in Table I V and Figure 1 show that nitrate ions possess a normal salt effect, a result previously obtained by competition kinetics (15, 17). O n the other hand, the salt effect of the I + e" reaction shows that this bis-pentacyano cobaltic peroxide ion has an m

m

m

m


6.

ANBAR AND HART


83

with Various Reactive Inorganic Compounds D,, X 10 cm. sec."

Γ* X JO* cm.

kdiff X -10" M " sec.'

2.2 1.1» 1.5 1.2 0.8 0.8 1.0 0.9 1.2 1.2 1.0 1.2 0.8 1.0 1.3 0.9 1.1 1.2 0.9 1.9 1.5 1.0 0.8

1.1 2.6 2.3 2.6 3.0 3.0 2.6 3.5 2.8 2.5 2.2 2.8 3.0 2.7 2.5 2.0 2.4 3.0 3.2 2.2 2.0 3.2 3.2

5.9 4.1 9.8 68. 0.66 4.8 1.4 64. 4.6 42. 3.1 4.6 0.68 4.2 10.2 2.6 3.6 5.1 5.4 46. 8.6 5.5 2.1

s

2

1

e


fc

m

J

e f

1

f

r

1

*diff 1.2 0.85 0.50 0.96 5.0 4.9 1.6 0.13 0.89 0.29 0.29 0.41 8.7 0.69 1.08 0.36 4.2 0.77 2.6 0.87 0.41 1.7 1.5

* Alfa Inorganic Inc. Flnka puriss. Estimated by analogy with Fe( CN )« which has been determined ( 37, 38 ). Estimated on basis (37, 38) of FeF*. Estimated by analogy with S0 ". Williams and Hopkins Reagent. Baker Analyzed.

j

k

a

1

m

4

2

n

p

effective charge close to unity and is far from behaving kinetically as a pentavalent anion. Discussion Diffusion Controlled Reactions. Many reactions of hydrated elec trons approach the diffusion controlled limit. This has been pointed out in previous reviews (I, 28, 34), but there has not been an extensive evaluation of a large number of these reactions. In the present study we evaluate quantitatively the parameters of diffusion controlled reactions and compare the calculated results with experimental data. The measured rates of reaction of C o (III) complexes clearly ap proach the diffusion controlled limit. A comparison between the diffusion


84

RADIATION CHEMISTRY

Table III.

Some Slower Reactions of e Supplier


Compounds 3.1 3.2 3.3 3.4 3.5 3.6 3.7 3.8 3.9 3.10 3.11 3.12 3.13 3.14 α 6 c d

NaBF LioSiF KH PQ KH P0 NaH PO.> NaSO,NH.> Na S0 Na Fe(CN) NH NaAs0 NaH.,As0 NaoSe0 K Ru(CN) Na.>Sn0 K Os(CN)

a

4

2

a

4

2

e

3

a

2

2

d

c

3

3

5

2

4

d 8

c c

d

3

4

d

e

3

4

c d

e

with Inorganic Compounds M

pH 5.8 5.9 7.1 6.7 6.8 11.7 10.0 8.6 10.6 11.0 10.8 10.6 11.0 10.5

b

6

a q

1

4.0 1.5 7.7 7.2 1.1

2

5

0 0.016 0.084 0.100 0.130

0 0.025" 0.093 0.116 0.141"

/ ( l + 2.5/*ι/2).


e

α

δ

5

5

6

7

8

8

6

6

8

6


6.


ANBAR A N D HART

0

0.!

85

0.2

Figure 1. The effect of ionic strength on the reactivities of NO f and [Co(CN),0 Co(CN) ~] ~ 2

Φ NOr,a = 1.33 © tCoiCNW.CoiCN)â

5

5

= 2.5

controlled rates for these ions calculated according to Debye's formula (20) with the experimental data (Table I ) shows excellent agreement for a series of tri and divalent cations ( 1.1-1.9). The lower value for corr iff = a for pentavalent ( N H , ) o C o 0 C o ( N H , ) , > (1.10) is easily explained i f one assumes that one or two of the positive charges are permanently neutralized by gegen anions in the outer sphere (26). If two of the charges are neutralized, a equals unity within experimental error. The very large a for the pentavalent anion ( C N ) C o 0 C o ( C N ) y ~ (1.11) is most probably caused by the same effect, as has been demon strated by its behavior in the presence of inert electrolytes ( Figure 1) where a kinetically apparent charge of unity has been found. In fact, when one calculates k for [ ( C N ) C o 0 C o ( C N ) Κ ] a is equal to unity within experimental error. Another series of fifteen highly reactive cations and cationic complexes of various metals is presented i n Table V . W i t h the exception of C d ( N H ) and N i ( H 0 ) each of these com plexes reacts with e~ at a rate which is within 2 5 % of the calculated diffusion controlled rate. In view of the uncertainties involved in the parameters which determine k and k , the results indicate that the calculated k is a fairly good estimate of the diffusion controlled limit ,v

5+

2

5

r>

(Uff

8

4

2

2 +

Γ)

2

4

2

4

2 +

m

corr

(liff

dif{


t

r,

86

RADIATION

CHEMISTRY

1

of e~ reactions. This conclusion implies that the experimental value of D(e~ ) = 4.7 X 10" c m . sec." is reliable. W e now consider the ques tion as to whether or not r - ~ 2.5 A . provides a reasonable encounter radius to describe the behavior of anionic and neutral reactants. m

r>

aq

2

1

c


Table V.

aq

A Comparison Between k and & of 9 to 12 A . for these reactions instead of the "geometrical" sum of 4-6 A . Distances as long as 9-12 A . are unlikely to facilitate electron transfer by orbital overlap but electron tunneling through a free-energy barrier is expected if the energy level of the accommodated electron in the acceptor molecule matches that of its initial state in water. W h e n the gain in free energy is large enough one can envisage many available electronic states thereby enhancing the probability for tunneling in spite of the high potential barrier. Alternatively, the overall rate of the reaction may depend not only on the initial tunneling step leading to an excited state but on the electron transition from this state to a more stable one. This transition is favored by an increase in the energy difference between these states. m

e

r


88

RADIATION CHEMISTRY-

Table V I .

A Comparison Between k

C0Tr

Korr


Compound 6.1 6.2

oNO "

M

X 1

M'

9

sec.'

1

10 cm. sec.' s

2

6.0 1.4

22.0 4.6

2

and kût

6.3

NO.f

8.5

1.8

6.4

s o -

10.6

0.9

6.5

cio-

7.0

2.2

6.6

vo -

4.9

1.5

6.7

Cr(CN) «-

4.0

0.9

6.8

Cr(CN) 4-

3.3

6.9 6.10 6.11 6.12 6.13 6.14 6.15 6.16

Cr(EDTA)" Cr(OX), " Cr0 CrF »Cro0 " Mn0 ~ Mn(CN) " Fe(CN) «-

6.17

Fe(CN) N0 "

2

8

2

3

e

e

26.0 18.0 10.0 2.7 33.0 22.0 5.4 3.0

3

4

2

e

2

7

4

6

4

e

a

24.0

2

5

0.8 a

0.9 0.9 1.0 1.0 0.8 1.5 0.9 0.9 1.0

6.18

FeF 8"

6.19 6.20 6.21 6.22 6.23 6.24

Fe(EDTA)OH Co(CN) Cl Co(CN),N(VCo(NO -) Co(EDTA)-

6.25

(CN) Co0 Co(CN) />-

6.26

Ni(EDTA)"

0.1

6.27

Ni(CN) "

2.7

1.2

6.28 6.29

AsF " Br ~

9.0 13.0

1.3 1.8

6.30

Pd(CN) 2-

1.9

1.1

6.31 6.32 6.33 6.34 6.35 6.36

PdCl " Ag(EDTA),Ag(CN) " Cd(EDTA) " SnOo " Sn(ÉDTA) "

7.7 1.6 1.5 0.39" 3.2 0.7

6.37

Sb(V

e

2

Co(CN) 3E

8

5

2

0

5

4

e

2

e

4

2

2

2

2

2

0.9 0.9 0.9 0.9 0.8 0.9

28.0

0.6

tt

a

a

r

2

1.0

7.1 3.6" 18.0 8.0 12.5 29.0 ffl

3

2

2.7

e

e

12.0

0.9

0.9 0.7 1.3 0.9 1.0 0.9 1.3


1

6.

ANBAR AND HART

89


H i g h l y Reactive Anions and Anionic Complexes Χ 10 cm. 8


1.0 1.1

W M

1

xv a—- • diff

x XT' seer

corr

Ref.

λ K

1

8.6 5.2

2.5 0.88 1.3

1.4

6.7

5.0

10.6

1.00

1.1

5.9

1.2

2.3

9.8

0.5

3.0

1.9

2.1

3.0

0.66

5.0

3.2 3.4 2.3 2.2 3.2 2.2 3.0 3.0

12.0 2.5 3.3 1.0 5.3 9.4 0.68 1.9

2.2 7.2 3.0 2.7 6.2 2.1 7.9 1.5

3.0

5.0

4.8

2.6

1.4

1.9

3.2 3.0 3.0 3.2 3.6 3.2

5.5 1.9 1.9 2.2 2.7 12.0

1.3 1.9 9.5 3.6 4.6 2.4 23.3 (1.3)

5.0

1.2 (21.0)

3.2

5.5

0.018

2.8

4.6

0.59

2.3 2.3

9.4 10.2

0.96 1.3

3.0

5.1

0.37

3.0 3.0 2.1 3.2 2.3 3.2

4.9 1.8 8.7 5.5 3.3 5.5

1.6 0.9 0.17 0.07 0.97 0.13

2.5

10.0

1.2

b

b

30 39 This work 39 This work This work 7 This work 36 36 13 7 39 39 2 23 This work This work 9 7 13 13 13 13 This work 9 This work 7 29 This work 7 9 7 9 7 9 This work


90

RADIATION CHEMISTRY

Table VI.

A Comparison Between k

corr

Χ

Korr

Compound


6.38

and k^ft of Highly

H*

9

M" sec.'

10 cm. sec."

Te0 -

15.0

0.9

1-3

20.0 7.7 11.0

1.5 1.4 1.4

4

1

2

1

s

2

6.39 6.40 6.41

i

e

;

aq

4

4

2

4

a q

2

aq

8

4

2

4

2

Ethylenediamine as well as E D T A are poorer electron conductors than H 0 . This can be seen from the reactivity of the corresponding N i ( I I ) complexes (2.8, 6.26) and N i ( e n ) (Jfc = 2 X M F M sec." ) 2

3

2 +


1

1

92

RADIATION CHEMISTRY

(31) compared with N i , as well as C d ( I I ) E D T A complex (6.34) (9) and C d ( e n ) (k = 3 X l O W sec." ) (31) compared with C d ( I I ) (7, 13). Meyer stein's data ( 31 ) shows that ethylenediamine is a poorer electron conductor than E D T A but this effect is not fully exhibited as long as some coordination sites remain unsubstituted by ethylenediamine. In other words, the presence of one or more water molecules on a given complex facilitates the incorporation of an electron even in the presence of other non-conducting ligands. The role of the ligand as electron conductor has been further elaborated in the case of the E D T A complexes of the rare earths and other trivalent cations (9). In these ions the reactivity increases with increasing stability of the complex—i.e., with increasing overlap of the metal-ligand orbitals (9). The oxide ion, O " , is an extremely poor electron conductor probably because of lack of orbital overlap with the hydrated electron in solution. Thus, the high reactivity of S n F " (2.14) compared with SnO.r" (3.13), that of TiF(i " (2.2) compared with T i 0 " (k < K F M " sec.' ) (35) or that of A s F f (6.28) (3) compared with H A s 0 " (3.10). In these fluoride complexes there is room for orbital overlap of e~ and the central atom. In the case of oxy complexes with a higher coordination number than 4 — e.g., S e 0 " (2.11) or T e 0 " (2.17) or I 0 " (7)—one observes, as expected from the A F - A F * relationship, that reactivity increases with increasing oxidation state of the central atom. In spite of the above discussed effects of ligands on the reactivity of complexes, d C r ( C N ) - (6.7), d C o ( C N ) - (6.20), C r ( E D T A ) " (6.9), Co ( E D T A ) (6.24), H g ( E D T A ) " (6.95), and P b ( E D T A ) " (6.46), as well as C r 0 (6.11), M n C V (6.14), and T e 0 " (6.38), react at diffusion controlled rates, some of them with abnormally high cross sections. These high A F reactions may involve electron tunneling which proceeds regardless of the electron conducting capacity of the ligands. Non-metallic Compounds. Certain generalizations may be reached by considering the reactivity of non-metallic elements throughout the periodic table. Phosphorus oxy-anions are non-reactive toward e~ . The comparable reactivities of H P ( V and H P O f (Table III) are probably because of their behavior as general protonic acids (32), as their pK's are comparable. H P 0 ~ , the anion of monobasic H P 0 , is non-reactive towards e~ . It should be noted, however, that the Bronsted relation between the p K values of protonic acids and their rates of reaction with e~ does not necessarily imply a proton transfer in the rate determining step (10). The observed Brônsted relation suggests that the same parameters which facilitate proton transfer to water by electron withdrawal from the hydrogen atom, X H - » X " + H 0 , enhance the incorporation of an additional electron into the vacant orbital of this atom without a 2 +

3


1

a q

2 +

1

1

a q

2

G

2

2

3

2

(

1

2

1

4

m

4

2

4

s

c

2

4

3

G

-

4

6

3

2

2

2

4

2

aq

2

2

2

2

3

2

aq

aq

3

+


6.

ANBAR A N D HART


93

simultaneous cleavage of the X - H bond. The non-reactivity of S i F " , H P ( V > H P 0 - , H P 0 - , S O , and SO4 " toward e\ (Table III) may be correlated with the instability of the intermediate valencies of Si ( I I I ) , P(IV),P(II),S(V),andS(III). In a given subgroup of the periodic table, e~ reactivity increases with increasing atomic number. This general trend, which has been previously pointed out ( i , 7), is confirmed by the behavior of homologous compounds of the elements i n the fifth and sixth groups, Ρ < As < Sb; S < Se < Te. c

2

2

3

2

2

a

2

2

2

q


&q

Concluding Remarks. It has recently been shown that the activation energy of many e~ -solute reactions is 3.4 ± 0.6 kcal./mole (4, 8). F r o m this result we have reasoned that the rates of many of the "slow" reactions are determined principally by their entropies of activation. These findings do not conflict with the mechanisms discussed above. If electron tunnel ing is a major contributing factor in these electron transfer reactions, the slower rates may be attributed to transmission coefficients smaller than unity. The factors which partially inhibit an electron transfer may thus be considered "geometrical" in nature, demonstrable by the entropy term of the free energy of activation, or by a decrease i n the transmission coefficient as a result of an increase in the width of the potential barrier. aq

The nature of the primary product of the e~ + X -> X " reaction remains an unsolved problem. It is probably produced i n a vibrationally or even electronically excited state ( 9 ). ( The de-excitation of this excited intermediate is still an open question (27, 35). In the absence of reliable experimental evidence on the immediate product of the rate determining step of the e" reactions, it is not practical to describe a detailed mecha nism of the interaction of e~ with a given compound. &q

aq

aq

In conclusion, we admit that our knowledge of inorganic compounds and their reactions with e~ is far from being complete. Thus, the re activity of these compounds towards e~ is less predictable than for organic compounds. However, we hope that the rate constant informa tion supplied by the e\ reactions w i l l lead to a better understanding of inorganic chemistry, and that out of the information and ideas presented in this paper some general principles w i l l develop. &q

&q

q

Literature Cited (1) Anbar, M . , A D V A N . C H E M . SER. 50, 55 (1965).

(2) (3) (4) (5) (6) (7) (8)

Anbar, M . , Chem. Commun. 1966, 416. Anbar, M . (unpublished). Anbar, M . , Alfasi, Z., Reisler, H., J. Am. Chem. Soc. 89, 1263 (1967). Anbar, M . , Hart, E . J., J. Am. Chem. Soc. 86, 5633 (1964). Anbar, M . , Hart, E. J., J. Phys. Chem. 69, 271 (1965). Ibid., 69, 973 (1965). Ibid., 71, 3700 (1967).


94

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RADIATION CHEMISTRY

1

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Schmidt, Κ. H., Buck, W. L . , Science 151, 70 (1966). Schwarz, Η. Α., Radiation Res. Suppl. 4, 89 (1965). Sutin, N . , "Exchange Reaction," p. 7, IAEA, Vienna, 1965. Szutka, Α., Thomas, J. K., Gordon, S., Hart, E . J., J. Phys. Chem. 69, 289 (1965). (37) "Tables of Interatomic Distances," Chem. Soc. (London) Spec. Publ. 11, 1958. (38) Ibid., 18, 1965. (39) Thomas, J. K., Gordon, S., Hart, E . J., J. Phys. Chem. 68, 1524 (1964). RECEIVED December 27, 1967. This work was performed under the auspices of the U . S. Atomic Energy Commission.


On the Reactivity of Hydrated Electrons Toward Inorganic Compounds

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