T. N. Bell, K. A. Perkins, and P. G. Perkins
2610
The Chlorination of Paraffin Hydrocarbons. Calculation of the Activation Energies and A Factors for Reactions in the Total Chlorination of Methane Thomas N. Bell, Department of Chemistry, Simon Fraser University, Burnaby 2, British Columbia, Canada
Kathryn A. Perklns, and Peter G. Perkins" Department of Pure and Applied Chemistry, University of Strathclyde, Glasgow, G I 1XL (Received May 25, 1977)
Activation energies and A factors for the series of gas-phase hydrogen transfer reactions CH(,-,,Cl, + C1. and CH(,,,Cl, + HC1 and chlorine transfer reactions CH&21, + Clz and CH(3-,)Cl(,+,)+ C1- for x = 0, 1, 2, and 3 have been calculated using the molecular orbital-bond index (MOBI) and group contribution method. Generally satisfactory agreement is found between observed and calculated activation energies, while in certain cases the A factors obtained are more reliable than currently accepted estimates.
Introduction The chlorination of long-chain paraffin hydrocarbons is an important technical process in the manufacture of solvents and fire-resistant materials and, in this reaction, the kinetic factors controlling the positions and degree of substitution are of vital importance in determining the products. Despite this, there has been little systematic kinetic work on such reactions with hydrocarbons containing more than four carbon atoms. The variety of products which may be obtained and the possibilities for substitutional isomers, together with experimental difficulties, have combined to produce this situation. Clearly, an advance could be made if a method were available to estimate the kinetic parameters for the reactions. Such a method has been described by us and in this paper, as the first part of a general theoretical investigation into the activation processes for the chlorination of paraffins, we deal with the series of reactions which bring about the total chlorination of methane. This system was chosen for the obvious reasons that it is the simplest and, moreover, kinetic parameters have been m e a ~ u r e d or ~-~ already estimated5 for all the reactions involved. The series of reactions studied were (a) the initiating step CH(,_,)CI,
+ C1.
-+
CH(,-,)Cl,
thalpy at a different reaction temperature, it is necessary to correct the value of AHo*by the integrated form of Kirchoff s law, i.e. A H " * = AH"*,,, - ACop(Tm)AT where AT is the difference between the mean experimental reaction temperature and 298 K and ACop(T,) is the difference in specific heat between the activated complex and the reactants at a mean temperature, T,. Values of C o p at different temperatures for reactants were obtained from the JANAF Tables6 and C o pfor the transition state was estimated from group contributions as described previously.2 The activation energy for a bimolecular reaction at the mean reaction temperature T is then given by' AE*, = A H " $ , -k 2RT U*may be compared directly with experimentally determined values of the activation energy. In addition to the activation energy, AE', the preexponential factor, A , was calculated for each of the reactions studied, i.e.
+ HCl
and (b) the sustaining step Back reactions were also considered. Satisfactory values for the activation energies for these steps were calculated, while the A factors obtained are probably more accurate than those previously proposed.
Method The method previously described1p2for calculating the energies of activation for some gas-phase hydrogen transfer reactions was extended to cover the reactions of present interest. As before, the heat of atomization of each system is computed in the standard state and a t 298 K at appropriate positions in the coordinate space such that the minimum enthalpy path for the reaction may be obtained. The activation enthalpy AHo*298 is then computed from for the initial and transition states a knowledge of moatam of the system. The enthalpies generated from these calculations are for reaction at 298 K. In order to obtain the activation enThe Journal of Physical Chemi@ry, Vol. 81, No. 26, 1977
where T is the mean temperature for the reaction and ASo*, is the difference in entropy between the transition state and the reactants in concentration units. A value of ASo*p (the activation entropy in pressure units) was computed for each reaction, using values of Sop from the literature,6 e.g., in the activation process CH,
+ C1.
-+
H,C---H*---Cl
values of S o pare available over a range of temperature for CH4 and C1.; So, for the complex was estimated using the group contribution scheme.2 ASo*, was then computed from AS"*, =
-
(1- y ) R In R ' T ,
where R is the gas constant in J K-l mol-', R' is the gas constant in L atm K-' mol-l, y = 2 (Le., the molecularity of the reaction) and T, is the mean reaction temperature. The basic types of reaction of interest are discussed under two heads.
2611
Chlorination of Paraffin Hydrocarbons
TABLE I: Atom Pair Parameters Bonding parameter, kJ mol-‘
Atom pair Bonded H-Cl “Nonbonded” “Nonbonded” “Nonbonded” Bonded C-H “Nonbonded”
444.16 25.81
H-C1 CkCl C-C1
-20.49 336.70 410.79-420.66 420.66
C-H
1. Hydrogen Transfer Reactions. Reactions studied here were those between a chlorine atom and the molecules CH4, CH3C1, CH2C12,and CHCl,, together with the reverse reactions. The optimum paths for the set of reactions \
--C-H*
\
\
I
I
t C1. -+ +--H*--Cl-+ -C*
1
+ H*Cl
were obtained by calculating the heats of atomization for different geometries of the systems \
-C--H*--Cl I
The geometry of the group CH,C1(3-,, is of importance with respect to the bond properties which have to be considered in these calculations. In the isolated molecule CH(,+,~Cl(,, the geometry may be considered tetrahedral. However, the geometry of the free radical CH,C1(3,, is not established. The assumption is made in this paper that the radicals are planar, thus, in forming the transition states
c1
H \
\
c1 \
C1-C- -H* - -C1 H-C- -H*--C1 Cl-C--H* - -C1 I
H
I
c1
I
c1
the geometry of the moiety CH,C1,3-,~ was assumed to change from pyramidal to planar as its C- - -H* length increased from 0,109 to 0.5 nm concomitantly with the H*- -C1 bond length being changed from 0.5 to 0.127 nm. The atom configuration C- - -H*- - -C1 was assumed to remain colinear throughout the reaction. The atom pair bond energy paramters which afford the heats of atomization are listed in Table I. Those for the atom pairs C-C1, “nonbonded” H-C1, “nonbonded” C1-C1, and C-H remain the same as those used in previous works2 As before, the C-H parameter varies with the C-H* bond length according to whether the C-H* bond is in a planar, a “fully” pyramidal, or some intermediate environment. The parameter used for the interaction between “bonded’ H and C1 atoms was taken as 444.76 k J mol-1.2 The C-C1 bonding parameter is allowed to vary with the configuration of the carbon atom, as was previously done for the corresponding C-H quantity. The value when the C-C1 is in the “fully” pyramidal situation is 336.70 k J mol-l. Bond energy parameters for C-C1 relevant to the radicals CH,CI., CHCly, and CC13. can be determined from experimental heats of formation and SCF calculations on the species. Benson7 gives experimental values for AH: at 298 K for CHg and CClg as 142.26 and 77.40 kJ, respectively. A linear interpolation affords values of AHf’ for CH2C1. and CHC12. a t 298 K as 120.50 and 99.16 kJ, and these values may be used to compute the value of AHoatom for each radical. Thus, the C-C1 parameter appropriate to each species can be deduced. These are as follows: C-C1 (CH&I.), 288.65 kJ mol-’; C-C1 (CHCl,.), 303.05 kJ mol-l; C-C1 (CCl,.), 311.75 k J mol-I. It is now necessary to allow
the C-C1 bond parameters to vary with the C-H* bond length. The lower limit is selected according to which radical is involved in the reaction; the higher limit, however, remains at 336.70 kJ mol-’. Contributions to the C-C1 parameter at intermediate bond lengths stem from the parameters for the pyramidal molecule, the planar radical, and the pyramidal radical, as in the C-H case. Current evidence on the most stable geometry of the CC13radicalg suggests that, while not completely planar, the radical is only slightly bent, having a structure intermediate between CH3 and CF3. The CH3 radical has a planar geometry and it is known by calculation that the energy of reorganization from the planar to the pyramidal form is 56.41 kJ The CF3 radical, however, is known to be pyramidal and, hence, has a reorganization energy in the opposite sense. Evidence on the structure of the CHClz radical has shown it to be planar.1° In these calculations we have made the assumption that the isolated radicals CC13,CC12H,and CClH2are all planar and that the bond-energy parameter for the pyramidal form is identical with that calculated for the planar form. Since the C-C1 bond indices in the two forms of CC13differ by only 2%, this is consistent with a small reorganization energy for this radical (about 17.2 kJ mol-l). The fraction which each of the three moieties contributes to the overall C-C1 bond-energy parameter at each C-H* bond length was assumed identical with that for the C-H parameter. Thus, at 0.16 nm where each of the three types of C-C1 bond contributes equally, the overall parameter for the C-C1 bond in C1&- is
‘/3(336.70
+ 311.75 + 311.75) = 320.07 kJ mol-’
In calculating the A factor for the process CH,-H* t C1. -+ CH,- - -H*- - -C1
values of S o p a t different temperatures are listed6 for methane and the chlorine atom. Values of S o pfor “bound H” and “bound CY’ in the transition state were taken as l/$70p(H2) and 1/2Sop(C12),respectively, and a value for “bound CH,” was abstracted from Benson’s table^.^ For other reactions, in order to estimate AS0*, So (bound radical) must be estimated. Thus, for example, in the process CH(,-,)Cl,
t C1+ CH(,-,)Cl,---H*---Cl
AS” * p = Sop(CH(+,)C1,
x = 1, 2, 3
- - -H*- - -C1) -
S0p(CH,4-x,C1,) - S”p(C1)
(1)
S“,(CH,3-,)ClX - -H*- - -C1) = Sop(CH~,-,,C1, bound) Sop(H b o u n d )
+
Sop(GI bound) S“,(CH( -,.Cl, b o u n d ) = Sop(CH,,-,,C1,) Sop(H b o u n d )
+ (2)
-
(3)
Using Sop(H bound) = 1/2S0p(H2)and So,(C1 bound) = 1/2Sop(C12) then substitution of (3) into (2) and then into (1)yields =
1/2
S0,(C12) - SO,(C1)
This means that the difference ASo*, between the transition state and the reactants depends only on the difference Sop(C1bound) - So,(C1 atom) at the mean temperature. Clearly, at a specific reaction temperature the A factors for this series of reactions are the same. A factors were also calculated for the reverse reactions CH3 + HC1 and CC13 + HC1. However, no entropy data The Journal of Physical Chemistry, Vol. 81, No. 26, 1977
T. N. Bell, K. A. Perkins, and P. G. Perkins
2612 TABLE 11: C-Cl and C1-Cl Bond Energy Parameters C-H* bond length \ in -C-- -H*, 1 nm
c-Cl* bond length \ in -c- - -c1*, 1 nm
c-Cl parameter for CH,Cl*, kJ mol-'
c-Cl parameter for CHCl, ., kJ mol-'
c-Cl parameter for cc1, ., kJ mol-'
Cl-Cl parameter, kJ mol-'
0.109 0.12 0.13 0.14 0.15 0.16 0.17 0.18 0.19 0.20 0.25 0.30 0.35 0.40 0.45 0.50
0.178 0.187 0.195 0.203 0.212 0.220 0.228 0.236 0.245 0.253 0.294 0.335 0.376 0.418 0.459 0.50
336.70 321.32 316.42 310.45 307.77 304.67 302.00 299.95 297.98 295.89 291.00 289.28 288.99 288.82 288.74 288.65
336.70 326.07 322.69 318.32 316.44 314.27 312.40 310.96 309.58 308.12 304.69 303.49 303.29 303.17 303.11 303.05
336.70 328.71 326.17 323.07 321.68 320.07 318.68 317.61 316.59 315.51 312.97 312.08 311.93 311.84 311.79 311.75
0 45.46 59.91 77.55 85.46 94.56 102.53 108.62 114.42 120.57 135.02 140.08 140.94 141.45 141.73 141.96
are available for the radicals CH2Cl and CHC12, so that A factors for reactions involving these radicals could not be calculated. 2. Chlorine Transfer Reactions. Here the reactions studied were those between Clz and the radicals CH3., CH2C1., CHCl2., and CC13-,i.e. \
-c.
\
t c1,
-+
\
-c--c1*--c1-+ -c-c1* /
I
I
+ c1.
The reverse reactions were also considered. The reaction path with minimum enthalpy was again calculated. The C1-C1 bond length was varied between 0.199 (the bond length in the chlorine molecule8) and 0.5 nm. The latter corresponds to effectively total separation. The C-C1* bond length was varied concomitantly between 0.5 and 0.178 nm8 and the geometry of the +C moiety was varied from planar to pyramidal as the reactions proceeded. The method of approach to this series, where the migrating atom is chlorine, was identical with that already described and, in general, the same bond energy parameters were used. The scaled parameters for C-H and C-C1 interactions were, however, made to depend on the length of the C-C1* bond. Moreover, it was necessary to incorporate energy contributions from the two mutually bonded chlorine atoms between which the C1-C1 distance varied. Since this had not occurred in any previous calculations, a new parameter for C1-C1 was needed. The quantity appropriate to Clz was calculated from the known standard heat formation of the molecule and the C1-Cl bond index; this yielded 141.96 kJ mol-l. For transient stages in the reaction, it was found appropriate to vary the Cl-C1 parameter as a function of the C-C1* bond distance. Hence, the parameter for bonded Cl-Cl* atoms was varied from its full value, 141.96 kJ molW1,when C-C1 was 0.5 nm (i.e., when the chlorine has no interaction with the radical) to zero when C-C1* was 0.178 nm &e., as it exists in the CH,Cl4_%molecule). The C1-C1 parameter was caused to change with C-C1* bond length in a manner identical with the C-C1 parameters. Values of the C1-C1 parameter at different C-C1* distances are given in Table 11. The results also afford the activation energies for the back reactions, i.e. \
-c-Cl* I
t
\
c1. -+ -c. I
+ c1*-c1
The calculated change in enthalpy A H o * b was corrected for temperature as described above and the activation The Journal of Physical Chemistry, Vol. 8 1,
No. 26, 1977
energy for the reverse reaction obtained. In calculating the A factors for forward complex formation \
-C. /
t c1*-c1-+
\
-c- - - -c1*-- - -C] /
the relevant values of S o pfor the radicals CH3. and CC13. were taken from JANAF tables.6 For complex formation, AS"*, is given by = Sop(bound radical) - S",(radical)
since Sop(boundC1) in the transition state is assumed to be '/zS0,(C12). No entropy data for the free radicals CHzC1- and CHC12- are available, so that no A factors involving these radicals could be calculated. It was possible to calculate A factors for all four reverse (chlorine transfer) reactions, taking the values of So, for the bound radical, CH,Cl(,-,, as S0p(CH,C1,,-x)) = S0p(CH,,+,,C1,3-x)) - '/2S0p(H2)
Results and Discussion The activation energy and log A calculated for each of the reactions are listed in Table 111. The corresponding experimental quantities are also given; unless otherwise indicated, these were taken from Trotman-Dickenson and Milne's table^.^ Each calculated value has been corrected to the median of the temperature range quoted for the particular reaction. The two central bond lengths specifying the dimensions of the transition state are also given in Table 111. It is found that, in general, the more highly substituted the carbon atom, the more compact the transition state becomes. Hydrogen Transfer Reactions. The largest discrepancy between the experimental and calculated activation energies in this series occurs for the reaction CH4 + Cl., although even here the difference is only -10 kJ mol-'. Possible reasons for this were discussed earlier.2 In addition, it is worthy of note that the activation energy for this reaction calculated by Zavitsas and Melikian'l is also considerably greater than the accepted experimental value. The activation energies calculated for the abstraction of hydrogen by a chlorine atom from CH3C1, CI1zC1z, and CHC13 are in good agreement with experiment. For the reverse reactions, the RMS difference between the calculated and experimental activation energies is 6.3 kJ mol-l. For the reaction of HC1 with CH3. and CC13. the
Chlorination of Paraffin Hydrocarbons
2613
TABLE 111: Arrhenius Parameters for Reactions Geometry of Mean transition state, nm reaction H-Cl temp, K C-H
Reaction
?if, mol-'
kJ mol-'
Calcd logA
Exptl A E * , kJ mol-l
Exptl log A
Hydrogen Transfer Forward Reactions CH, t C1CH, t HCl CH,Cl t C1+ CH,Cl t HCl
389
0.158
0.17
-47.9
26.4
12.77
423
0.135
0.14
-55.2
13.3
12.46
CH,Cl, t C1-t CHC1, t HC1
423
0.125
0.145
-55.2
15.1
12.46
CHCl, t C1+ CCl, + HC1
423
0.125
0.14
-55.2
13.6
12.46
16.1 13.8 ~t0.4 14.1 12.9 12.5 13.1 23.0 14.0 13.9 27.2
13.42 13.5 t 0.7 13.76 13.5 13.43 13.4 14.6 12.84 13.2 14.6
12.9 20.ga
11.73 12.0'
Back Reactions CH, t HC1-t CH, t C1 CH,C1 t HCl+ CH,Cl t C1 CHCl, t HCl+ CH,Cl, + C1 CCl, + HC1CHC1, t C1
-77.5
18.1
11.46
360
0.158
0.17
400
0.135
0.14
32.6
34.3'1
12.1'1
400
0.125
0.145
38.2
46.9'
12.0'
365
0.125
0.14
47.3d 56.1'
11.65d 11.8'1
-77.2
42.9
11.18
Chlorine Transfer Forward Reactions c-c1 CH, t C1,CH,Cl t C1 CH,C1 + C1, + CH,Cl, + C1 CHCl, + C1,CHC1, t C1 CCl, t c1,CCl, + c1
c1-Cl
365
-68.2
12.ga
0.265
0.20
23.6
11.66 (est)
12.5a
12.6a
365
0.23
0.205
26.5
11.66 (est)
16.7a
12.0'1
365
0.23
0.21
30.4
11.67
25.1' 22.2
11.74 ?: 0.6a 12.86
-67.8
400
-51.0
112.8
12.63
104.6a
14.0a1b
400
0.265
0.20
-47.5
121.0
12.82
89.5'
14.0a,b
400
0.23
0.205
-38.0
111.2
13.31
87.9'
14,Oa,
400
0.23
0.21
-27.4
100.6
13.87
83.7' 79.1a A factor is assumed t o be 1014. ' Calculated from the reverse reaction.
values given by Chiltz et al.5 are somewhat higher than more recently determined values. It may be that their experimental activation energies for the reactions CH2Cl HC1 and CHC12 HC1 are also in error in the same sense. The calculated value of log A for CHC13 C1. agrees well with experiment, whereas for the remainder of the series the calculated values are rather lower than the experimentally determined quantities. The overall root mean square error is -0.8 in log units. The two A factors which it was possible to compute for the reverse reactions show satisfactory agreement with experiment. Chlorine Transfer Reactions. 1. Forward Reactions. The enthalpy calculations for the reaction
+
+
+
9.6a
365
a Values taken from ref 5. ref 4.
CH,. t C1,
11.65
Back Reactions
CH,Cl + C1+ CH, t C1, CH,Cl, + C1+ CH,Cl + C1, CHCl, + C1+ CHCl, + C1, CCl, + Cl-t CCl, t c1,
+
6.4
CH,- -Cl*- C 1 - t CH,Cl t C1.
indicate that the reaction path follows a steady decrease in energy from reactants to products. Hence, AHo*for this reaction is computed to be zero. At a mean reaction temperature of 365 K, aE* is thus 6.4 kJ mol-l, to be compared with an experimental value of 9.6 kJ mol-l.
14.0' 14.3' Values taken from
For the reactions of C12with CH2C1., CHC12., and CCl,., the experimental activation energies are given as 12.616.7, and 25.1 kJ mol-l, respectively. Of these, the data for the last reaction are probably the most accurate since, in this case, the measurements are not complicated by further chlorination of the products. Such phenomena generally lead to systematic errors in experimental determinations which are frequently difficult either to eliminate or estimate. However, for this most reliable case we calculate that the activation energy is 30.4 kJ mol-l, in good agreement with experiment. The activation energies calculated for the reactions CH2C1. and CHC12. agree less well with experiment although, over the whole series, the RMS difference is only 8 kJ mol-I. Calculation of log A was not possible for the reactions CH2C1. + Clz and CHC12. C12,since values of So are not available for the radicals CH2C1. and CHClZ.. Ifowever, evaluation of log A at 365 K was carried out for the systems CH3. + Clz and CC13. C12,yielding values of 11.65 and 11.67, respectively. Since these two values are so close, log A values for CH2C1. + C12and CHC12. + C12are estimated
+
+
The Journal of Physical Chemistry, Vol. 81, No. 26, 1977
E. Hankiewicz and D. Schulte-Frohlinde
2814
to be 11.66 by interpolation. The agreement of the calculated and experimental values of log A for CC13-+ Clz is very satisfactory. (2) Back Reactions R-Cl + Cl C12 + R. The back reactions in which a chlorine atom reacts with a chlorinated methane all have higher activation energies than any of the reactions so far studied. Little experimental work has been carried out on these systems and A factors have not been measured for the reaction of a chlorine atom with CH,Cl, CH2Cl2,and CHC13. For the reaction C C 4 C1our calculated value for log A is 13.87, to be compared with average experimental values of 14.15. Taking the activation energies first, the most reliable experimental results will again be those for the fully chlorinated reactant CC4. Here, the calculated value differs from experiment by 17 kJ mol-'. A survey of the calculated and experimental activation energies in Table I11 reveals a discrepancy in the experimental results. Whereas the RMS difference between theory and experiment for the reactions of radicals with Clz is 8 k J mol-', that for the reverse reactions is 19 kJ mol-'. This is somewhat difficult to reconcile. Since in all reactions the transition state is the same whatever the direction of reaction, we would expect the difference between the calculated and experimental activation en-
-
+
-
ergies to be very similar for both the forward and reverse reactions. It is, moreover, noteworthy that, for reactants in the reverse processes, the heats of formation have been accurately determined and so should not introduce further error to the values of &!?*forward. This discrepancy must cast some doubt on the experimental values hitherto accepted. References and Notes (1) T. N. Bell and P. G. Perkins, Nature (London), 258, 300 (1975). (2) T. N. Bell and P. G. Perkins, submitted for publication. (3) A. F. Trotman-Dickenson and G. S. Milne, Nafl. Bur. Stand. Ref. Data Ser. No. 9 (1967). (4) E. Ratajczak and A. F. Trotman-Dickenson, "Supplementary Tables of Bimolecular Gas Reactions", UWIST, Cardiff, 1969. (5) G. Chiltz, P. Goldfinger, G. Huybrechts, G. Martens, and G. Verbeke, Chem. Rev., 63,355 (1963). (6) "JANAF Thermochemical Tables", Dow Chemical Co., Mdhnd, Mich., 1965. (7) S.W. Benson and H. E. O'Neal, Natl. Bur. Stand., Circ., No. 21 (1970). (8) "Tables of In!eratomic Distances and Configuration in Molecules and Ions", Chem. SOC., Spec. Pub/., No. 11 (1958). (9) L. J. Aarons, I. H. Hillier, and M. F. Guest, J. Chem. Soc., Faraday Trans. 2, 70, 167 (1974). (10) S. P. Mishra, G. W. Neilson, and M. C. R. Symons, J . Chem. Soc., Faraday Trans. 2, 69, 1425 (1973). (11) A. A. Zavitsas and A. A. Melikian, J . Am. Chem. SOC.,97, 2757 (1975).
The Rate of Hydrated Electron Reaction with Neutral and Anionic Scavengers in Concentrated Salt Solutions E. Hankiewicd and D. Schulte-Frohlinde" Institut fur Strahlenchemieim Max-Planck-lnsfitut fur Kohienforschung, 0-4330 Muiheim/Ruhr, West Germany (Received July 1 1, 1977) Publication costs asslsted by Instifut fur Strahlenchemie im Max-Pianck-lnstitut fur Kohienforschung
The rate constants, kob& of the e,; reaction with nitrobenzene, IrCls2-,and Fe(CN)83-have been measured in water, and in LiCl and CsCl solutions up to a concentration of 14 M LiCl and 6 M CsCl and were found to be diffusion controlled. The viscosity of CsCl solutions is independent of CsCl concentration whereas in the case of LiCl solutions the logarithm of the viscosity decreases linearly with concentration. Since the logarithm of the rate constant, log hobs& with nitrobenzene decreases linearly with LiCl concentration the viscosity is the rate-determining parameter. In CsCl solutions the rate constants do not vary with salt concentration. Iridate and ferricyanide behave like neutral species at LiCl concentrations above 1and 6 M, respectively; above these LiCl concentrations the dependence of their rate constants and that of nitrobenzene on salt concentration is the same. This is due to the formation of undissociated scavengers and the behavior described by the Smoluchowski-Debye equation at high salt concentration.
I. Introduction Anbar and Hart'n2 have measured the rate constants for the reaction of the hydrated electron with several solutes (acetone, nitrous oxide, benzoate, and nitrate ions) in water and in aqueous solutions which were 12.4 M in potassium fluoride. They found that the rate constants are lower in the concentrated salt solution by a factor of 2-5 depending on the solute. It remained an open question whether a decrease of the activity coefficient of e, -, a decrease of its diffusion coefficient, or a viscosity clange of the salt solutions was responsible for the observed effect. A similar decrease in the reactivity of the solvated electron was observed by Pikaev a t al.3 for the reaction of eaq-with ions of transuranium elements in concentrated alkali and in carbonate solutions. These reactions were 8 Permanent Address: Institute of Applied Radiation Chemistry, Technical University, Wroblewskiego 15, 93-590 Lodz, Poland.
The Journal of Physical Chemistry, Voi. 81, No. 26, 1977
faster than diffusion controlled in contrast to the ones investigated by Anbar and Hart.' As a possible explanation for the decrease of the rate constants an increase of the viscosity of the solutions or capture of electrons in deeper traps was mentioned but not proved. A similar decrease was also observed for rate constants of the reactions ea; + eaq-and ea; + O-.4 Since concentrated aqueous salt solutions in liquid and especially in glassy states are currently under intensive investigation, we decided to study the factors which may be responsible for the decrease of the rate constant of the hydrated electron in the case of diffusion-controlled reactions. Lithium chloride was used as an added salt since the viscosity of its solutions is known and the rate of reaction of eaq-with Li' is slow. Furthermore we wanted to test whether different cations have different effects on the reaction rate constants of ea; with added scavengers, since enhancement of the rate
