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Direct observations of solvated electrons in liquid hydrocarbons
Volume 10, number 3
DIRECT OBSERVATWNS
CHEMICAL PHYSICS LE’ITERS
OF SOLVATED
1 August 1971
ELECTRONS PN LIQUID ~YDR~~~RB~~~
*
J.T. RBXARIX Umiwsityof Salford, Sa~foniM.54WT, UK
and 3.K. THOMAS
Received 4 June 1971 Using the techniques of pulse radiolysis a short-lived transient absorbing species (Xrnax = 1500 run] has been observed in liquid hydrocarbons at 193%. This absorption has been assigned to a trapped or sofvated electron and a&olute reaction rates have been obtained for reactions of sohated electrons with several soiutcs known to be good efeo tran scavengers. Observetions have also ‘been made of the negative ions formed by electron attacfiment LO biphenyl at 193OK and it is concluded that at least two distinct species give rise to the negative ions Of biphenyt At room temperature electron transfer rates have been measured from the biphenyl anion to various acceptors, in particular oxygen. It is concluded that while oxygen reacts with solvated eledtrons in hy&omhons at 193% Witha rate constant ke-+02 = f -5x10” M-r se~-~, at room temperature where the electrons are not skated the electron pius oxygen reacnon is much slower than the reactions of electrons with biphenyt
1. Introduction Duriug the last decade e large amount of information has been obtained by radiation chemists on the yields and properties of soivated electrons in a large variety of polar liquids. In liquids such as water and, alcohols the electrons liberated in the primary ion&aGon events are eventually stabilized as soivated electrons. At room temperature the lifetimes of the soivated electrons are usually of the order of~lsec and the absorption maxima are observed in “&ewavelength region SO0 + 1000 nm [I]. However, in Liquid hydrocarbons at room temperature no direct observations of solvated eIectrons have been made, though reaotions of electrons with added solutes have been well characterized [2] _ Direct ‘observations of electrons trapped in solid hydrocarbon matrices at 77°K have been made using ~radiolysis and various competition reactions studied [33 ; recent pulse radiolysis studies. * Based on work performed under the auspicesof the U.S. AtomicEnegy Commtion. _
have shown that parts of the spectra are time dependent [4f _Since it is by no means certain that electrons in liquid hydrocarbons are in any way so&ted, perhaps due to the low diefectric constant or alternatively that they are not observed due to their short Iifetirne at room temperature, the present work describes pulse radio&&s experiments carried out on Liquid hydrocarbons at low temperatures and also on electron transfer reactions at room temperature in an attempt to obtain a clearer understanding of electron reactions in hydrocarbons.
‘Ibe pulse radiolysis apparatus utilizing a 3 MeVvan de Craaff accelerator has already been described in detaiJ [5], and sample preparations and the low temperature cell used for this workhave already been described [4] . Transient spectra irt the visr%le~ region of the _-&rum were recorded using an II?28 photomultip~cr, and in the irtfrared by using a _%ilcd (LF4200) photodiode.
Volume 10, number ‘3
CHEMKAL
PHYSICS LE+l-ERs
: The radirkon dosein thecell was monitored by using the :solvated electron absorption produced in the radiolysis of degassed ethanol. Either 12 or 33 n&c pulses were used and a dose in a 12 nsec puke was 2.4 X lero ,-eV/I.
i Results 3.1. Irradiation at 193’K No transitory absorption spectrum which can be attributed to trapped or solvated electrons ‘is observed in the nanosecond pulse radiolysis of pure hydrocarbons at room temperature (296°K) over the wavelength range 250 + 1600 run. However, at 193°K pulse radiolysis of either n-hexane or 3-methylhexane (both of which are, mobile liquids at this temperature) produces a short-hved species having the properties of a trapped or solvated electron. The spectrum of this transient species jn 3 Ml-i is shown in fig. 1. The spectrum is markedly asymmetric showing a long tail to the blue and the transient is very short-lived, as is indicated by the oScilIogram insert which shows the decay at X = 1SP. lhe decay possesses an initial fast portion which gives way to a slower decay with a half-life of about 60 nsec. This type of decay is similar to the
1 August 1971
behaviour ‘of ion pairs in hydrocarbon solvents and has been observed for the biphenyi anion in cyclohexane. 161’.The yield of the species is G < 0.1, if an extinction coefficient of.30 000 M-l cm-l a’t A= 1.5~ is used [3 1. Addition of known electron scavengers such as CCl,+and SF, incrked the rate of decay of the spectrum given in fig. 1, while positive ion scavengers such as dissolved ammonia gas and alcohols had no significant effect. By adding small quantities of electron scavengers to the hydrocarbon and observing the increased rate of decay of the transient monitored at X= 1.5~ it was possible from the solute concentration dependence to obtain absolute rate constants for reaction of the electrons with the various added solutes, The rate constants measured at 193°K are given in table 1. Table 1 Solute
Rate constant (M” set-‘1
oxygen nitrous oxide biphenyl carbon tetiachloride
1.5 1.1 1.9 2.3
X10” x10” x 10” X 10”
In the solutions which contained biphenyl as the added solute it was also possible to monitor the production of the biphenyl anion which has a characteristic absorption at X=410 nm. Over the time scale in which the electron was observed to decay no significant growth of the anion absorption was observed; most of the absorption was produced within the resolution time of the apparatus.
I
am
I
Iuoo
I
I
1200
I 1400
!
I 16Gu
Fig. 1. Transient absorption sp+rum of the trapped election obtained in the pulse radiolysisof liquid 3-methylhexane at i93% Insert shows ut oscilloscope trace of the decay of the electrti~ abti&tion monijored at A= 1.5~.
318
:
3.2. Electron transfer reactions at 296’K In cyclohexane containing 10m3M biphenyl the negative ion transient absorption was monitored at A= 410 nm, addition of small concentrations of a second solute such as benzylchloride resulted in an increased rate of decay of the 410 nm absorption and also a decrease in the initial yield of the biphenyl anion. This’type of data can be interpreted in terms of a competition between biphenyl and benzylchloride_for electrons and also electron transfer from @; to the benzylchloride. Typical oscilIosco@e &ure data for the system are given in fig, 2. Except for a
Volume 10, number 3
CHEMICAL PHYSICS LETTERS
1 August 1971
A
% a
0.01 0.009 0.008 0.007 0.a 0.005 0.004
ABSCISSA, ORDINATE,
SWEEP SPEED TRANSMISSION
50 nsec/DiVISlON CHANGE 5%/DlVlSlON
traces obtained in the pulse ndiolysis of 10m3 M a2 in C6Ht2 solutions monitored at h = 410 m. (A) Argon saturated. (B) Containing 2.5X10” hi~cH~C1. (C) Cor;taining 5x 10m4 M cb CH2Cl. (D) Containing 10W3M d,CH2Cl. (E) Containing 10m3 hi 02.
Fii. 2.Oscilloscop~
Fig. 3. Effect of bentyl chloride on the First order decay rate of #pz on 10m3 M Q~/C~H~~ solutions monitored at A=410 run. NB.: AOD because of for@Ike& portion left, probably 0:. o 0.25 x lOa M Q C&CL 130.5x LO-3 MO CH&L A 125x io-s M Q CHzCL l 1.50x lo+ M 9 CH@.
0.05 0.03
small cont~b~tjon from a very long-lived component to the absorption at 410 nm, the decays are fit order and the data for the benzylchloride system are shown in fig. 3. The initial competition is marked by the gradual decrease in the AOD intercept. Similar behaviour (i.e. increased rate of decay of @i, and reduction in yield) was obtained using sulphur hexafIuoride and pyrene as second solutes in the 10m3M 9-$&2 system. The solutes carbon dioxide and nitrous oxide when added to the biphenyl solutions, do not increase the rate of decay of $2 but do decrease the initial yield of the anion after the pulse, while addition of oxygen caused the #; absorption to decay more rapidly but did not affect the initial yield of 4, even when the oxygen concentration was much higher than the corresponrfing hiph.enyl concentration. Data for oxygen are given in figs. 2 and 4. It may be concluded that the rate of reaction of oxygen with electrons in cyclohexane at room temperature is at least t Ml time5
0.05
0.02
8 u
0.01 0.009 0.007 0.008 0.006 0.005 O.CC4 0.003
Fig. 4. Effect of o&en concedration on ffie fast order decay rate of & monitod at ii=: 410 Ml in 1O”3M&/C5Hr2
solutioqr o 0.23X104M0~. 10-4M02,~,4.56X104M~.
A E.l6XlQ* Mop. 0 2.28X
319
Volume 10, pumbef 3
CHEIS~ICAL PHYSICSLETTERS :
.leSs than the rate of reaction of electrons with biphenyl. From the concentration dependence of the fmt order rate constants in the cases where electron transfer (as measured by the increase in rate of decay of #,) did occur, the electron transfer rates given in table 2 were obtained. .Table 2 Solute
Rate of election transfer from @2 (M-1 set-‘)
benzylchloride pyrene sulphur bexafluoride oxygen nitrous oxide carbon dioxide
1.0x 10’0 3;2x10t” o.75x1o’o 2.3X 10” < lo6
4. Discussion The results indicate that at 193°K solvated or trapped electrons are formed in liquid hydrocarbons; these electrons which are only observed in small yields, react with oxygen, carbon tetrachloride, etc., at a rate which is ten times higher than that of diffusion controlled reactions of anions. It may be inferred from this that even though the electrons are trapped for a short period of time, their reactions may be mainly controlled by the thermal detrapping process to give “quasi” free electrons.. It is known that the mobility of the quasi free electron is lo3 to 104 : [7,8] greater than that of a trapped or solvated electron; hence, if the electron spends some time trapped and some free, an overall mobility which is greater than expected by classical diffusion will be achieved. The shape of the transient absorption spectrum of the trapped electron given in fig. 1 is very similar to the short time~unrclaxed spectrum obtained in the pulse radiolysis of 3 MH glasses at 77°K [4]. In the lattei case, the spectrum is attributed to the photoionization of the election from the trap and as time proceeds the solvent molecules which form the’trap relax and a transition’of the electron to a bound excited state’at E < Eli is possible. This gives rise a’
to
323
1 August 1971
symmetrical spectrum to the red of, and overlapping with the ionization~spectrum’[9] :At 193’K it appears that only, the photoiom?ation spectrum of the electron is observed..The shallow trapping of ihe electron and the evidence for detrapping suggests that the. orientation of the solvent molecules in the trap is not permanent but.that distortion of the trap is continually occurriig. Hence, the ground state transition to a bound state does not occur. The rates of electron transfer from the biphenyl anion to the various acceptors are in line with similar electron transfer reactions that have previously. been studied in more polar solvents [ lo]. We have observed three distinct classes of reactants: (1) those which undergo electron transfer reactions and also compete efficiently for the electron precursor of r#; (benzylchloride, pyrene and suiphur hexatluoride); (2) those which compete effeciently for the precursor of 92 but do not accept electrons from 42 (nitrous oxide and carbon dioxide); (3) those which do not compete efficiently for the precursor of @i but undergo electron transfer from I$; (oxygen). Of the three types, the observed reactions of oxygen are the most significant since the work at 193’K clearly shows that oxygen reacts with solvated electrons with a rate constant k= 1.5X1011 M-l set-l which is almost as fast as the rate constant with biphenyl, k = 1.9 X lo1 i M-1 set-l, yet at room temperature oxygen does not compete with biphenyl for electrons even when present in a ten-fold excess. This ciearly indicates that in hydrocarbon solvents oxygen undergoes at least two distinct reactions with radiation produced electrons, one wi+h solvated electrons ?t a rate comparable with other electron scavengers and one with non-solvated electrons at a much slower rate than other electron scavengers. The lack of growth observed in the $2 absorption at 193°K even under conditions where the electron decay rate could also be monitored indicates that non-solvated electrons react much more rapidly than solvated electrons with biphenyl, since biphenyl reacts with solvated eIectrons with a rate c,onstant of about 1OL1M-l se&; this would ‘put a rate constant in excess of 1012 M-l sec’l for reactions of non-solvated or quasi free electrons in these syster&; Possible explanations for the effect of &xygen and the high rate constants will be discussed elsewhere. .-
V&m&
10, number 3
CIiEMIcAL PHYSICS fz#Tl-ERS
[Sj J.W. Nunt and Jg Thomas, Radiation Res. 32 (1967) 14%
Reference++ [l] 3.X. Thomas, Radiation Rea Rev. 1(X968) 185. [2] JM Warman, K$. Anmus and R.H. Schuler, Ad&. Chem Ser. 82,2 (1968) 25. 131 W.H. Ham& in: Radicat ions, eds. Kaiser &xl Kevan fIrMscience, New York, 1965) p. 321. [4J J.T. Richards and J.Ic Thomas,,J. Chem. Phys. 53 (1970) 218.
.LAugust 197 1
”
[6] 3.K. Thomas, K. Johnsox~,I’_ KEppert and R. Lowers, 3. CknPhys.48(t968) 1603. f7] R.f-L M&day, L.D. Schmidt and H.T. Da+, J. Chem. Phyh 50 (1969) 1473. [8f P.H. Tewari dnd G.R. Freeman, I. Chcm. Phys. 49 (1968)4394. [9] W.H.Hami& J. Chem. fbys 53 (1970) 473, [ 1OJ GE Adams; D. Michael knd R.L. %‘iEson.,AlEva%Chem. Ser. 81(1968) 289.
321
DIRECT OBSERVATWNS
CHEMICAL PHYSICS LE’ITERS
OF SOLVATED
1 August 1971
ELECTRONS PN LIQUID ~YDR~~~RB~~~
*
J.T. RBXARIX Umiwsityof Salford, Sa~foniM.54WT, UK
and 3.K. THOMAS
Received 4 June 1971 Using the techniques of pulse radiolysis a short-lived transient absorbing species (Xrnax = 1500 run] has been observed in liquid hydrocarbons at 193%. This absorption has been assigned to a trapped or sofvated electron and a&olute reaction rates have been obtained for reactions of sohated electrons with several soiutcs known to be good efeo tran scavengers. Observetions have also ‘been made of the negative ions formed by electron attacfiment LO biphenyl at 193OK and it is concluded that at least two distinct species give rise to the negative ions Of biphenyt At room temperature electron transfer rates have been measured from the biphenyl anion to various acceptors, in particular oxygen. It is concluded that while oxygen reacts with solvated eledtrons in hy&omhons at 193% Witha rate constant ke-+02 = f -5x10” M-r se~-~, at room temperature where the electrons are not skated the electron pius oxygen reacnon is much slower than the reactions of electrons with biphenyt
1. Introduction Duriug the last decade e large amount of information has been obtained by radiation chemists on the yields and properties of soivated electrons in a large variety of polar liquids. In liquids such as water and, alcohols the electrons liberated in the primary ion&aGon events are eventually stabilized as soivated electrons. At room temperature the lifetimes of the soivated electrons are usually of the order of~lsec and the absorption maxima are observed in “&ewavelength region SO0 + 1000 nm [I]. However, in Liquid hydrocarbons at room temperature no direct observations of solvated eIectrons have been made, though reaotions of electrons with added solutes have been well characterized [2] _ Direct ‘observations of electrons trapped in solid hydrocarbon matrices at 77°K have been made using ~radiolysis and various competition reactions studied [33 ; recent pulse radiolysis studies. * Based on work performed under the auspicesof the U.S. AtomicEnegy Commtion. _
have shown that parts of the spectra are time dependent [4f _Since it is by no means certain that electrons in liquid hydrocarbons are in any way so&ted, perhaps due to the low diefectric constant or alternatively that they are not observed due to their short Iifetirne at room temperature, the present work describes pulse radio&&s experiments carried out on Liquid hydrocarbons at low temperatures and also on electron transfer reactions at room temperature in an attempt to obtain a clearer understanding of electron reactions in hydrocarbons.
‘Ibe pulse radiolysis apparatus utilizing a 3 MeVvan de Craaff accelerator has already been described in detaiJ [5], and sample preparations and the low temperature cell used for this workhave already been described [4] . Transient spectra irt the visr%le~ region of the _-&rum were recorded using an II?28 photomultip~cr, and in the irtfrared by using a _%ilcd (LF4200) photodiode.
Volume 10, number ‘3
CHEMKAL
PHYSICS LE+l-ERs
: The radirkon dosein thecell was monitored by using the :solvated electron absorption produced in the radiolysis of degassed ethanol. Either 12 or 33 n&c pulses were used and a dose in a 12 nsec puke was 2.4 X lero ,-eV/I.
i Results 3.1. Irradiation at 193’K No transitory absorption spectrum which can be attributed to trapped or solvated electrons ‘is observed in the nanosecond pulse radiolysis of pure hydrocarbons at room temperature (296°K) over the wavelength range 250 + 1600 run. However, at 193°K pulse radiolysis of either n-hexane or 3-methylhexane (both of which are, mobile liquids at this temperature) produces a short-hved species having the properties of a trapped or solvated electron. The spectrum of this transient species jn 3 Ml-i is shown in fig. 1. The spectrum is markedly asymmetric showing a long tail to the blue and the transient is very short-lived, as is indicated by the oScilIogram insert which shows the decay at X = 1SP. lhe decay possesses an initial fast portion which gives way to a slower decay with a half-life of about 60 nsec. This type of decay is similar to the
1 August 1971
behaviour ‘of ion pairs in hydrocarbon solvents and has been observed for the biphenyi anion in cyclohexane. 161’.The yield of the species is G < 0.1, if an extinction coefficient of.30 000 M-l cm-l a’t A= 1.5~ is used [3 1. Addition of known electron scavengers such as CCl,+and SF, incrked the rate of decay of the spectrum given in fig. 1, while positive ion scavengers such as dissolved ammonia gas and alcohols had no significant effect. By adding small quantities of electron scavengers to the hydrocarbon and observing the increased rate of decay of the transient monitored at X= 1.5~ it was possible from the solute concentration dependence to obtain absolute rate constants for reaction of the electrons with the various added solutes, The rate constants measured at 193°K are given in table 1. Table 1 Solute
Rate constant (M” set-‘1
oxygen nitrous oxide biphenyl carbon tetiachloride
1.5 1.1 1.9 2.3
X10” x10” x 10” X 10”
In the solutions which contained biphenyl as the added solute it was also possible to monitor the production of the biphenyl anion which has a characteristic absorption at X=410 nm. Over the time scale in which the electron was observed to decay no significant growth of the anion absorption was observed; most of the absorption was produced within the resolution time of the apparatus.
I
am
I
Iuoo
I
I
1200
I 1400
!
I 16Gu
Fig. 1. Transient absorption sp+rum of the trapped election obtained in the pulse radiolysisof liquid 3-methylhexane at i93% Insert shows ut oscilloscope trace of the decay of the electrti~ abti&tion monijored at A= 1.5~.
318
:
3.2. Electron transfer reactions at 296’K In cyclohexane containing 10m3M biphenyl the negative ion transient absorption was monitored at A= 410 nm, addition of small concentrations of a second solute such as benzylchloride resulted in an increased rate of decay of the 410 nm absorption and also a decrease in the initial yield of the biphenyl anion. This’type of data can be interpreted in terms of a competition between biphenyl and benzylchloride_for electrons and also electron transfer from @; to the benzylchloride. Typical oscilIosco@e &ure data for the system are given in fig, 2. Except for a
Volume 10, number 3
CHEMICAL PHYSICS LETTERS
1 August 1971
A
% a
0.01 0.009 0.008 0.007 0.a 0.005 0.004
ABSCISSA, ORDINATE,
SWEEP SPEED TRANSMISSION
50 nsec/DiVISlON CHANGE 5%/DlVlSlON
traces obtained in the pulse ndiolysis of 10m3 M a2 in C6Ht2 solutions monitored at h = 410 m. (A) Argon saturated. (B) Containing 2.5X10” hi~cH~C1. (C) Cor;taining 5x 10m4 M cb CH2Cl. (D) Containing 10W3M d,CH2Cl. (E) Containing 10m3 hi 02.
Fii. 2.Oscilloscop~
Fig. 3. Effect of bentyl chloride on the First order decay rate of #pz on 10m3 M Q~/C~H~~ solutions monitored at A=410 run. NB.: AOD because of for@Ike& portion left, probably 0:. o 0.25 x lOa M Q C&CL 130.5x LO-3 MO CH&L A 125x io-s M Q CHzCL l 1.50x lo+ M 9 CH@.
0.05 0.03
small cont~b~tjon from a very long-lived component to the absorption at 410 nm, the decays are fit order and the data for the benzylchloride system are shown in fig. 3. The initial competition is marked by the gradual decrease in the AOD intercept. Similar behaviour (i.e. increased rate of decay of @i, and reduction in yield) was obtained using sulphur hexafIuoride and pyrene as second solutes in the 10m3M 9-$&2 system. The solutes carbon dioxide and nitrous oxide when added to the biphenyl solutions, do not increase the rate of decay of $2 but do decrease the initial yield of the anion after the pulse, while addition of oxygen caused the #; absorption to decay more rapidly but did not affect the initial yield of 4, even when the oxygen concentration was much higher than the corresponrfing hiph.enyl concentration. Data for oxygen are given in figs. 2 and 4. It may be concluded that the rate of reaction of oxygen with electrons in cyclohexane at room temperature is at least t Ml time5
0.05
0.02
8 u
0.01 0.009 0.007 0.008 0.006 0.005 O.CC4 0.003
Fig. 4. Effect of o&en concedration on ffie fast order decay rate of & monitod at ii=: 410 Ml in 1O”3M&/C5Hr2
solutioqr o 0.23X104M0~. 10-4M02,~,4.56X104M~.
A E.l6XlQ* Mop. 0 2.28X
319
Volume 10, pumbef 3
CHEIS~ICAL PHYSICSLETTERS :
.leSs than the rate of reaction of electrons with biphenyl. From the concentration dependence of the fmt order rate constants in the cases where electron transfer (as measured by the increase in rate of decay of #,) did occur, the electron transfer rates given in table 2 were obtained. .Table 2 Solute
Rate of election transfer from @2 (M-1 set-‘)
benzylchloride pyrene sulphur bexafluoride oxygen nitrous oxide carbon dioxide
1.0x 10’0 3;2x10t” o.75x1o’o 2.3X 10” < lo6
4. Discussion The results indicate that at 193°K solvated or trapped electrons are formed in liquid hydrocarbons; these electrons which are only observed in small yields, react with oxygen, carbon tetrachloride, etc., at a rate which is ten times higher than that of diffusion controlled reactions of anions. It may be inferred from this that even though the electrons are trapped for a short period of time, their reactions may be mainly controlled by the thermal detrapping process to give “quasi” free electrons.. It is known that the mobility of the quasi free electron is lo3 to 104 : [7,8] greater than that of a trapped or solvated electron; hence, if the electron spends some time trapped and some free, an overall mobility which is greater than expected by classical diffusion will be achieved. The shape of the transient absorption spectrum of the trapped electron given in fig. 1 is very similar to the short time~unrclaxed spectrum obtained in the pulse radiolysis of 3 MH glasses at 77°K [4]. In the lattei case, the spectrum is attributed to the photoionization of the election from the trap and as time proceeds the solvent molecules which form the’trap relax and a transition’of the electron to a bound excited state’at E < Eli is possible. This gives rise a’
to
323
1 August 1971
symmetrical spectrum to the red of, and overlapping with the ionization~spectrum’[9] :At 193’K it appears that only, the photoiom?ation spectrum of the electron is observed..The shallow trapping of ihe electron and the evidence for detrapping suggests that the. orientation of the solvent molecules in the trap is not permanent but.that distortion of the trap is continually occurriig. Hence, the ground state transition to a bound state does not occur. The rates of electron transfer from the biphenyl anion to the various acceptors are in line with similar electron transfer reactions that have previously. been studied in more polar solvents [ lo]. We have observed three distinct classes of reactants: (1) those which undergo electron transfer reactions and also compete efficiently for the electron precursor of r#; (benzylchloride, pyrene and suiphur hexatluoride); (2) those which compete effeciently for the precursor of 92 but do not accept electrons from 42 (nitrous oxide and carbon dioxide); (3) those which do not compete efficiently for the precursor of @i but undergo electron transfer from I$; (oxygen). Of the three types, the observed reactions of oxygen are the most significant since the work at 193’K clearly shows that oxygen reacts with solvated electrons with a rate constant k= 1.5X1011 M-l set-l which is almost as fast as the rate constant with biphenyl, k = 1.9 X lo1 i M-1 set-l, yet at room temperature oxygen does not compete with biphenyl for electrons even when present in a ten-fold excess. This ciearly indicates that in hydrocarbon solvents oxygen undergoes at least two distinct reactions with radiation produced electrons, one wi+h solvated electrons ?t a rate comparable with other electron scavengers and one with non-solvated electrons at a much slower rate than other electron scavengers. The lack of growth observed in the $2 absorption at 193°K even under conditions where the electron decay rate could also be monitored indicates that non-solvated electrons react much more rapidly than solvated electrons with biphenyl, since biphenyl reacts with solvated eIectrons with a rate c,onstant of about 1OL1M-l se&; this would ‘put a rate constant in excess of 1012 M-l sec’l for reactions of non-solvated or quasi free electrons in these syster&; Possible explanations for the effect of &xygen and the high rate constants will be discussed elsewhere. .-
V&m&
10, number 3
CIiEMIcAL PHYSICS fz#Tl-ERS
[Sj J.W. Nunt and Jg Thomas, Radiation Res. 32 (1967) 14%
Reference++ [l] 3.X. Thomas, Radiation Rea Rev. 1(X968) 185. [2] JM Warman, K$. Anmus and R.H. Schuler, Ad&. Chem Ser. 82,2 (1968) 25. 131 W.H. Ham& in: Radicat ions, eds. Kaiser &xl Kevan fIrMscience, New York, 1965) p. 321. [4J J.T. Richards and J.Ic Thomas,,J. Chem. Phys. 53 (1970) 218.
.LAugust 197 1
”
[6] 3.K. Thomas, K. Johnsox~,I’_ KEppert and R. Lowers, 3. CknPhys.48(t968) 1603. f7] R.f-L M&day, L.D. Schmidt and H.T. Da+, J. Chem. Phyh 50 (1969) 1473. [8f P.H. Tewari dnd G.R. Freeman, I. Chcm. Phys. 49 (1968)4394. [9] W.H.Hami& J. Chem. fbys 53 (1970) 473, [ 1OJ GE Adams; D. Michael knd R.L. %‘iEson.,AlEva%Chem. Ser. 81(1968) 289.
321